Solubility of carbon dioxide in binary mixtures of dimethyl sulfoxide and ethylene glycol: LFER analysis

J. Chem. Thermodynamics 141 (2020) 105968

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Solubility of carbon dioxide in binary mixtures of dimethyl sulfoxide and ethylene glycol: LFER analysis Ali Reza Harifi-Mood Department of Physical Chemistry, Faculty of Chemistry, Kharazmi University, Tehran, Iran

a r t i c l e

i n f o

Article history: Received 1 August 2019 Received in revised form 20 September 2019 Accepted 23 September 2019 Available online 24 September 2019 Keywords: Carbon dioxide Solubility Henry’s constant Linear free energy relationship

a b s t r a c t The solubility of carbon dioxide was determined in dimethyl sulfoxide, ethylene glycol, and their binary mixtures at the temperatures ranging from 298.15 Kto 328.15 K and the pressures ranging from (0 to 6) bar (0 to 0.6 MPa). It was shown that the solubility of CO2 increased with the increasing of pressure while decreased with the increasing of temperature. The Henry’s constants and thermodynamic properties of solution were calculated from the solubility data. The solubility and Henry’s constant of CO2 in the solvent mixtures increased with increasing the mole fraction of dimethyl sulfoxide. Finally, the dependence of CO2 solubility on solvent properties was investigated using the linear Gibbs energy relationship method. Results showed that solute-solvent interactions in the scale of polarity/polarizability and hydrogen bond acceptor ability of the solvent were the main parameters which control the solubility of CO2 in the solvents. Ó 2019 Elsevier Ltd.

1. Introduction Global carbon emissions from human activities have significantly increased in recent decades. Carbon dioxide is the primary greenhouse gas emitted through these activities. While carbon dioxide emissions come from a variety of natural sources, human-related emissions are responsible for the increase that has occurred in the atmosphere since the industrial revolution. The main human activity that emits carbon dioxide is the combustion of fossil fuels. One of the operational methods to reduce the large carbon dioxide emissions that are occurring for example during fossil fuel combustion are post-combustion technologies. Some of post-combustion technologies for carbon dioxide removal include physical and chemical absorption [1–3], adsorption [4–6], membrane processes [7], cryogenic separation [8] and electrochemical methods [9]. The most commonly method for carbon dioxide removal is the absorption onto solvents [10–12]. The solubility of gases in solvents may be divided into two types. In the first type, the gas molecules react chemically with solvent molecules while in another molecules do not. Since there is always solute–solvent interaction in solution, a solvent is never a substance which merely acts as an inert media for gases. Therefore, there is no sharp boundary for this classification and only in such cases where this interaction is strong enough to produce new

E-mail address: [email protected] https://doi.org/10.1016/j.jct.2019.105968 0021-9614/Ó 2019 Elsevier Ltd.

chemical species is it known chemically solubility such as aminebased solvents [11]. Understanding the solvent effects on the gases solubility based on physical interactions is essential for a rational selection of solvents. This purpose will be achieved by the investigation of the solubility of a selected gas in different solvents. It’s a necessity that all solvents behave chemically inert and the solubility involves solute–solvent interactions. The preparing of many solvents with similar behaviour might be both difficult and high cost. Solvent mixture systems provide a high potential to investigate the solvent effects on physical and chemical processes [13]. Solute-solvent interactions in term of solvent polarity have been as an effective factor which was used synonymously with the power to solvate solute charges. It is expected that solubility in the binary and ternary mixtures of the solvents is much more complex than in pure solvent due to solvent–solvent interactions. These interactions affect the solubility and therefore the capacity of the solvent to remove the gases from atmosphere. Because of the application of carbon dioxide in supercritical anti-solvent technology and the use of very polar organic solvents such as DMSO to precipitate of pharmaceuticals [14–16] and polymers [17,18], an individual contribution has been reported about the solubility of carbon dioxide in DMSO in high pressure [19]. In this contribution, we have reported experimental data on the solubility of carbon dioxide and its calorimetric properties in binary systems of dimethyl sulfoxide (DMSO) and ethylene glycol (EG) in the temperature range of 298–328 K. Enthalpy and entropy of

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solution have been determined as well as the excess Henry’s coefficient as a function of solvent composition. Due to the advantages of low-pressure gas solubility over high-pressure equilibrium data such as possibility to adopt various approximations in evaluating thermodynamic parameters, all experiments were performed at pressures up to 6 bar (0.6 MPa). DMSO and EG were chosen as common high polar-aprotic and moderate polar-protic solvents, respectively. Since EG is an abundant and low cost solvent, the mixture of DMSO with EG is more affordable than using the pure DMSO. Also, the presence of EG can tune the solubility of carbon dioxide. On the other hand, this binary system can demonstrate the influences of solute–solvent interactions in addition to solvent–solvent interactions on the solubility of carbon dioxide.

2. Experimental Fig. 1. Schematic diagram of the experimental apparatus. (1) gas cylinder; (2) gas reservoir; (3) equilibrium cell; (4) and (5) pressure transducers; (6) vaccum pump; (7) magnetic stirrer; (8) water bath; V1, gas regulator ; and V2 – V5, ball valves.

2.1. Materials Carbon dioxide with maximum purity of 99.9% in liquid phase was purchased from Spadana Gas, Iran. Analytical reagent grade DMSO (
99.9%) and ethylene glycol (
99.8%) were purchased from Merck. Solvents were dried by 3 Å molecular sieve. The water content of the dried compounds was determined by coulometric Karl Fischer titration yielding < 200  106 ppm residual water for ethylene glycol and DMSO. The summary of the chemical used, their purities, and sources are summarized in Table 1.

2.2. Apparatus and procedure The apparatus and procedures used were similar to those reported elsewhere [20]. The equipment used to obtain the gas solubility data in pure DMSO, EG, and their mixtures was made of stainless steel and schematically presented in Fig. 1. It was composed of CO2 cylinder, water bath, CO2 gas equilibrium cell with magnetic stirrer, gas reservoir, pressure transmitters and a vacuum pump. The pressures were monitored using pressure transmitter (WIKA 1/400 NPT 0 to 600.0 kPa). A charge of carbon dioxide was introduced into the gas reservoir equipped with pressure transmitter through first ball valve (V2). Second valve (V3) connected the reservoir gas to equilibrium cell whose volume was precalibrated. The amount of the gas introduced was determined using ideal gas equation state by pressure measurements. Third ball valve (V4) was used to purge the gas and introduce a given amount of solvent into the equilibrium cell. Finally, the fourth valve (V5) connected the manifold and other parts of equipment to vacuum pump. The equilibrium cell was movable and was connected to fixed parts by a clamp which could undergo high pressure up to 10 bar (1 MPa). The equilibrium cell and gas reservoir were immersed in a temperature controlled water bath where its temperature is monitored continuously by a PT-100 thermocouples with a precision of 0.1 K. The volume of equilibrium cell was carefully determined by recording the pressure of equilibrium cell and reservoir tank in present and in absent of a standard metallic ball into equilibrium cell. The volume of equilibrium cell along with its connection was exactly obtained 75 cm3.

Before the solubility measurements, the apparatus was purged with CO2 for half an hour. Binary mixtures of DMSO and EG over the entire range of mole fractions were gravimetrically prepared before each solubility measurement. Solvent or solvent mixture degassing was performed by freezing (using liquid nitrogen) and melting (at room temperature) the solvent under moderate vacuum, repeatedly, until no gas bubbles were observed in the melting process. In order to measure gas solubility a known mass of the solvent or binary solvent mixture was loaded into the equilibrium cell placed in water bath and connected to the rest of equipment. The equilibrium cell was purged with incubated CO2 in desired temperature for 10 min when V4 was open. Then V4 was closed and the gas was immediately injected into the equilibrium cell at a known pressure. The solution was agitated by magnetic stirring to reach equilibrium after approximately 30–60 min, corresponding to the temperature and depending on the type of solvent. The vapour pressure of the solvent and the mole fraction of the solvated gas were measured by using the initial and final (equilibrium) pressure readings of the equilibrium cell corresponding to introduced method by Nitta et al. [21]. Densities of all the pure solvents and mixtures under ambient pressure and different temperatures are used from the literature [22]. The volume of liquid solution in equilibrium cell was directly obtained from the mass and density of solvent at different temperatures. The volume expansion of the liquid in equilibrium cell because of the dissolution of CO2 was very small and neglected.

3. Results and discussion Prior to the measurement of the solubility of CO2 in the binary mixtures of DMSO and EG, the values of the solubility of the gas in pure DMSO and EG were measured for validating and testing the reproducibility of our experimental work. The solubility of the gas in pure DMSO at 298.15 K and several pressures were first measured and compared with data reported in previous researches. Although there are few data for solubility of CO2 in

Table 1 Provenance and mass fraction purity of chemicals used in the study. Chemical

Abbreviation

Mass fraction purity by the supplier

Supplier

Purification Method

Degassing of solvents

Carbon dioxide Dimethyl sulfoxide Ethylene glycol

CO2 DMSO EG

>0.999 >0.999 >0.998

Spadana Gas Co., Iran Merck Merck

Molecular sieve water content < 150  106 Molecular sieve water content < 200  106

Freezing and melting in liquid N2 Freezing and melting in liquid N2

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DMSO at low pressure, Fig. 2 shows that the experimental results in this work follow the suggested trend in the literature [19,23]. The similar comparison of the solubility of CO2 in EG has been performed with the literature. The collected data in Fig. 3 demonstrate good agreement between our experimental results and those from the literature [24–27]. Among the reported data of CO2 solubility in EG at 318.15 K, it has been found a series of results with noticeable incompatibility [28]. Although the authors have claimed good agreement between their own experimental results and those from the literature [29], data obviously represent a trend away from reality (Fig. 2 in [28]). This can be attributed to the impurities of EG such as water content. So that, the solubility values reported at 318.15 K correspond to the data obtained at 298.15 K (Fig. 3). Consequently, these comparisons confirm the correctness of our experimental method and the operation of the equipment. Table 2 presents experimental carbon dioxide solubility data in pure DMSO, EG, and their binary mixtures at temperatures ranging from 298.15 K to 328.15 K and pressures up to 6.00 bar (0.600 MPa). It is obvious that carbon dioxide is much more soluble in DMSO than in EG. Less solubility in EG can be attributed to the enhancing of hydrogen bond cages and aggregate structures of adjacent EG molecules in present of a non-polar solute such as carbon dioxide. This would certainly limit the location of solute molecules within the structured part of solvent [30]. It will be further discussed in term of specific and non-specific solute–solvent interactions (see 3.3). 3.1. Henry’s constants Henry’s law constant, H, based on mole fraction is often used to assess the solubility of CO2 in solvent. It was evaluated from the experimental data measured in this work as follows:

Hi ¼ lim

xli !0

xv u P fi ¼ lim i l i l l xi xi !0 xi

ð1Þ

where fi and ui are the fugacity and fugacity coefficient of the dissolved carbon dioxide in solution. xli and xvi are, respectively, the vapour and liquid phase mole fractions of carbon dioxide. At experiment conditions, the fugacity coefficient of gas was approximately unity and the fugacity was equal to equilibrium pressure of CO2

Fig. 3. Comparison of experimental values of solubility of CO2 in EG at 298.15 and 318.15 K with literature data: (e, /, ▻) [24,25,27] at 298.15 K; (s, D, h) [25,26,28] at 318.15 K; (d, j) this work at 298.15 and 318.15, respectively; and lines, smoothed.

because of the relative low equilibrium pressure. Fig. 4 shows representative isothermal CO2 solubility at five different compositions of binary solvent mixtures at T = 298.15 K. According to Eq. (1), Henry’s constants were determined from the slope of the isotherm created from the linear fit of CO2 mole fraction versus equilibrium pressure and are listed in Table 3. It is clear that the Henry’s constant values of CO2 in the DMSO-EG system increases with increasing temperature. At all temperatures, Henry’s constant is the highest in pure EG and its value decreases rapidly with increasing mole fraction of DMSO. Versteeg et al. showed that Henry’s constant varied non-linearly with temperature according to proposed Eq. (2) [31]:

lnHCO2 ¼ aH þ

bH RT

ð2Þ

where aH and bH are the constants of the relationship. Evaluated Henry’s constants were fitted to the Eq. (2) and the values of the parameters aH and bH along with the standard deviation of the fits are summarized in Table 4. Clearly, the slope, bH =R, increases sharply and then decreases smoothly (through a maximum atxEG = 0.1– 0.4) to its value in pure EG. Physical dissolution of a gas in binary mixtures demonstrates a pseudo-exponential variation for Henry’s constant respect to the solvent composition, according to O’Connell et al. report [32]. This thermodynamic model is based on Eq. (3) which gives Henry’s constant in the mixed solvent as a function of the two Henry’s constants in the pure solvents.

lnHCO2 ;

mixed

¼ xDMSO lnHCO2 ;

EG  aD:E xDMSO xEG

Fig. 2. Comparison of experimental data trend of solubility of CO2 in DMSO at 298.15 K (s) with literature data (j) [19], (▲) [23].

DMSO

þ xEG lnHCO2 ; ð3Þ

In this equation, xDMSO and xEG are the mole fractions of DMSO and EG, respectively, and aD:E refers to Margules model parameter [33] and is so-called as an interaction term. The characteristic of this equation is that Henry’s constant for the solute in the mixed solvent shows an exponential rather than a linear function of the solvent composition even when aD:E ¼ 0. Eq. (3) was used to correlate Henry’s law constant and the results are shown in Fig. 5. The interaction term aD:E show temperature dependence and is given by Eq. (4).

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Table 2 Experimental solubility data of CO2 in DMSO, EG, and their binary mixtures at different pressures and temperatures.a T = 298.15 K

T = 308.15 K 2

T = 318.15 K 10  xCO2

10 P/MPa

10  xCO2

10 P/MPa

102  xCO2

0.792 1.015 1.270 1.686 2.052 2.972 3.179

0.953 1.587 2.125 2.729 3.234 3.860 4.294

0.559 0.938 1.294 1.672 1.960 2.280 2.592

1.280 1.834 2.371 2.980 3.612 4.196 4.763

0.629 0.945 1.202 1.536 1.820 2.111 2.452

1.022 1.640 2.402 2.862 3.549 4.271 4.883

0.489 0.756 1.090 1.319 1.635 1.954 2.245

xEG ¼ 0:1b 1.210 1.766 2.208 2.883 3.541 4.015

0.863 1.210 1.526 2.017 2.439 2.781

0.930 1.560 1.943 2.673 3.234 3.87 4.483

0.542 0.879 1.110 1.560 1.848 2.190 2.551

1.197 1.720 2.342 2.943 3.788 4.270 5.027

0.519 0.740 1.069 1.335 1.721 1.977 2.296

1.045 1.457 2.034 2.755 3.096 3.956 4.782

0.402 0.580 0.819 1.098 1.273 1.600 1.906

xEG ¼ 0:2 0.923 1.640 2.274 2.823 3.533 3.911 4.213

0.538 0.971 1.362 1.690 2.150 2.370 2.516

1.143 1.824 2.338 3.013 3.767 4.112 4.741

0.538 0.903 1.157 1.511 1.865 2.080 2.345

1.231 1.853 2.367 3.054 3.760 4.476 5.293

0.515 0.813 0.990 1.278 1.531 1.873 2.242

1.068 1.457 2.145 2.676 3.567 4.310 4.853

0.378 0.530 0.764 1.009 1.297 1.538 1.765

xEG ¼ 0:3 0.930 1.732 2.154 2.896 3.354 3.880 4.243

0.501 0.925 1.177 1.583 1.833 2.080 2.319

1.025 1.854 2.623 3.144 3.810 4.443 4.971

0.420 0.792 1.110 1.390 1.660 1.930 2.208

1.112 1.675 2.439 3.060 3.677 4.155 4.783

0.373 0.623 0.921 1.120 1.374 1.548 1.740

1.022 1.650 2.376 3.008 3.760 4.307 5.008

0.325 0.568 0.771 0.977 1.227 1.399 1.640

xEG ¼ 0:4 0.896 1.643 2.190 2.776 3.354 3.896 4.373

0.439 0.805 1.050 1.340 1.660 1.900 2.140

0.954 1.630 2.321 2.870 3.229 3.932 4.672

0.395 0.659 0.938 1.170 1.330 1.589 1.888

0.928 1.543 2.195 2.934 3.435 4.231 4.903

0.297 0.508 0.718 0.968 1.126 1.370 1.608

1.231 1.745 2.238 2.884 3.564 4.286 5.098

0.348 0.494 0.637 0.809 1.020 1.200 1.470

xEG ¼ 0:5 0.943 1.467 2.165 2.765 3.334 3.954 4.420

0.407 0.633 0.942 1.230 1.443 1.680 1.913

1.106 1.544 2.231 2.678 3.543 4.102 4.635

0.391 0.546 0.788 0.943 1.252 1.460 1.636

0.832 1.685 2.400 2.980 3.764 4.365 5.153

0.235 0.482 0.712 0.876 1.090 1.260 1.514

1.110 1.622 2.322 3.123 3.880 4.441 5.12

0.281 0.419 0.595 0.794 0.995 1.146 1.317

xEG ¼ 0:6 0.945 1.572 2.249 2.846 3.651 4.231 4.859

0.355 0.567 0.813 1.090 1.352 1.550 1.795

1.065 1.663 2.326 2.967 3.547 4.105 4.728

0.332 0.524 0.745 0.912 1.140 1.279 1.471

1.118 1.774 2.168 3.014 3.835 4.539 5.218

0.265 0.451 0.578 0.785 0.978 1.160 1.359

1.012 1.639 2.188 2.893 3.679 4.536 5.265

0.224 0.385 0.476 0.659 0.838 1.020 1.200

xEG ¼ 0:7 1.607 1.987 2.576 3.245 4.302 5.012 5.858

0.512 0.630 0.858 1.055 1.383 1.609 1.871

1.153 1.673 2.364 3.007 3.673 4.184 4.878

0.315 0.463 0.662 0.812 1.040 1.146 1.335

1.078 1.758 2.334 2.923 3.856 4.561 5.293

0.262 0.401 0.537 0.692 0.876 1.070 1.207

1.241 1.857 2.379 3.021 3.759 4.544 5.378

0.257 0.371 0.491 0.603 0.771 0.907 1.080

xEG ¼ 0:8 1.043 1.691 2.456 3.345

0.321 0.491 0.696 1.002

0.945 1.523 2.245 2.965

0.213 0.386 0.536 0.708

1.028 1.782 2.552 3.321

0.203 0.335 0.524 0.646

0.893 1.658 2.438 3.356

0.152 0.276 0.436 0.594

10  xCO2

Pure DMSO 0.956 1.188 1.502 2.050 2.540 3.915 4.181

2

T = 328.15 K

10 P/MPa

10 P/MPa

b

2

5

A.R. Harifi-Mood / J. Chem. Thermodynamics 141 (2020) 105968 Table 2 (continued) T = 298.15 K

a b

T = 308.15 K

T = 318.15 K

T = 328.15 K

10 P/MPa

102  xCO2 b

10 P/MPa

102  xCO2

10 P/MPa

102  xCO2

10 P/MPa

102  xCO2

3.987 4.763 5.673

1.101 1.342 1.591

3.643 4.230 4.938

0.881 1.010 1.177

4.078 4.673 5.335

0.803 0.920 1.065

4.154 4.894 5.568

0.708 0.851 0.972

xEG ¼ 0:9 1.156 1.669 2.367 3.228 3.828 4.628 5.610

0.312 0.451 0.622 0.804 0.951 1.128 1.343

1.056 1.892 2.564 3.367 4.012 4.657 5.142

0.226 0.375 0.514 0.697 0.811 0.937 1.051

0.986 1.754 2.478 3.006 3.754 4.461 5.315

0.147 0.289 0.422 0.521 0.634 0.760 0.908

1.112 1.878 2.678 3.465 4.087 4.967 5.733

0.170 0.267 0.411 0.517 0.643 0.745 0.866

Pure ethylene glycol 0.546 1.543 2.434 2.763 3.550 4.570 5.350

0.139 0.319 0.513 0.602 0.733 0.934 1.076

1.134 1.794 2.564 2.934 3.674 4.670 5.073

0.195 0.293 0.451 0.502 0.652 0.795 0.872

1.012 1.892 2.564 3.078 3.875 4.653 5.761

0.148 0.298 0.39 0.452 0.576 0.723 0.876

1.231 1.760 2.345 3.054 3.854 4.765 5.711

0.168 0.239 0.286 0.403 0.514 0.619 0.768


Standard uncertainties u are ur xCO2 ¼ 0:05, ur ðxEG Þ ¼ 0:01, u (T) = 0.05 K, u(P) = 0.0001 MPa. xCO2 is the mole fraction of solvated CO2 in solvent and xEG is the mole fraction of ethylene glycol in binary mixture of solvents.

aD:E ¼ 3:6411 þ

1233:5 K T

ð4Þ

This equation shows that the interaction term decreases with increasing temperature. Solvent-solvent interactions consisting of specific and non-specific interactions decrease with the increasing temperature owing to the increasing intermolecular distances. Therefore, the variation of Henry’s constant versus solvent composition is leaded to an ideal behaviour when the temperature increases (Fig. 5). 3.2. Thermodynamic properties Thermodynamic properties of solutions of CO2 in solvent can be calculated by relating to the Henry’s law constants as follows [34]:

  @lnHCO2 DH ¼ R @ ð1=T Þ

ð5Þ

  @TlnHCO2 @T

ð6Þ

DG ¼ RT lnHCO2 ¼ DH  T DS

ð7Þ

DS ¼ R

The DH, DS, and DG are the enthalpy, entropy, and standard Gibbs energy changes of CO2 solution at the standard state, respectively. The average thermodynamic property changes at temperatures ranging from 298.15 K to 328.15 K and standard state were shown in Table 5. Data has been compared with some literature results for pure solvents [26,35]. The enthalpy change of solution is related to strength of interaction between the liquids and the gas molecules while the entropy change of solution shows the degree of ordering of the liquid–gas mixture. In these solutions, the negative value of DH indicates that the dissolution of CO2 in solvents is exothermic, on the other word, is favourable in terms of enthalpy. The highest values of DH are observed around the mole fraction of 0.1–0.4 with respect to EG. It can be attributed to stronger solute–solvent interactions which can be affected by high solvent–solvent interaction (see below). Since the DS is related to the solvent molecules organization surrounding the dissolved CO2, larger negative entropy around the mentioned solvent composition indicates stronger solvent-CO2 interactions. Table 3 The experimental determined Henry’s constants of CO2 (H) in DMSO, EG, and their binary mixtures at various temperatures.a xEG

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 Fig. 4. The plot of linear relationship between equlibrium pressure and mole fraction of dissolved CO2 in binary solvent mixtures at T = 298.15 K.

10  H/MPa T = 298.15 K

T = 308.15 K

T = 318.15 K

T = 328.15 K

137.5 145.4 164.2 183.3 203.6 232.7 270.3 313.5 363.0 434.8 506.8

165.8 176.5 197.9 222.3 247.8 282.1 320.7 363.7 418.8 493.5 576.8

194.8 213.3 239.3 268.0 305.7 340.9 382.8 438.0 500.7 572.2 652.2

219.2 247.1 275.7 307.7 348.2 387.4 438.9 501.4 571.8 653.7 746.9

a Relative standard uncertainty ur for xEG and the standard uncertainty u of T are as defined in Table 2.

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Table 4 The values of coefficients, aH , bH , and

a

r (standard deviation) obtained from Eq. (2).a

xEG

aH

 bH = J  mol

102  r

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1

7.7625 8.5328 8.5323 8.6327 9.0003 8.8047 8.6545 8.6703 8.6436 8.2608 8.1382

12717.1 14507.1 14206.1 14174.5 14825.5 14005.8 13278.3 12967.3 12538.3 11144.9 10453.2

1.63 1.1 1.27 1.36 2.21 1.77 1.06 1.74 1.65 1.31 0.97

1

Relative standard uncertainty ur of xEG is as defined in Table 2.

properties of solute and solvent will influence the solubility of compound in solvents. Scientists have attempted to define and obtain compound descriptors that would model solvation in various solvent phases. Work has been carried out to introduce the special descriptors of solute–solvent interactions. Kamlet and Taft, and co-workers have shown that it is indeed possible to define particular descriptors for solvents as distinctive characteristic of the solute–solvent interactions [36,37]. The most famous parameters those are well-known as ‘‘solvatochromic parameters” include polarity/polarizability, p , hydrogen bond donor, a, and hydrogen bond acceptor, b, abilities. These solvent properties can be combined into a linear free energy relationship (LFER) where dependent variable such as physical or chemical behaviour of a system shows multi-dependence on media parameters. A similar argument has been reported for standard Gibbs energy of the transfer of compound in the two phases, which in turn is related to the solute properties [38]. In a different way, we have investigated the correlation of solubility of CO2 or equilibrium constant of solubility, K, with the solvent properties, through the LFER shown in Eq. (8).

logK ¼ c þ pp þ aa þ bb

ð8Þ

The p, a, and b refer to the tendency of the solvent phase to interact with solute through polarity/polarizability, hydrogen bond acidity, and hydrogen bond basicity, respectively. Most of the effects of Eq. (8) have been really quite beneficial to not only the predicting of theoretical solubility in a known solvent but also the analysing of relationship term-by-term in order to isolate and to quantify the particular interactions that influence the solubility of gas. The solvatochromic parameters for binary mixtures of DMSOEG are available as a function of solvent composition from our previous report [39]. The general LFER, Eq. (8), was applied to the values of standard Gibbs energy changes of CO2 solution in Table 5 to yield Eq. (9), Fig. 5. The variation of Henry’s constant as a function of DMSO mole fraction in DMSO-EG binary mixtures. Points are the obtained values from experimental data and lines show fitted data based on Eq. (3).

Table 5 Calculated standard enthalpy (DH), entropy (DS), and Gibbs energy (DG) changes of solutions at 0.1 MPa and T = 298.15 K.a

a b c

1

xEG

DH=kJ  mol

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1

12.72 (12.43b) 14.51 14.21 14.17 14.83 14.01 13.28 12.97 12.54 11.14 10.45 (11.78c)

DS/J∙K1∙mol

1

83.51 (80.0b) 90.01 90.01 90.81 93.83 92.20 91.05 91.25 91.04 87.90 86.88 (91.6c)

DG=kJ  mol

1

12.18 (11.42b) 12.33 12.63 12.91 13.15 13.48 13.87 14.24 14.60 15.07 15.45

Relative standard uncertainty ur of xEG is as defined in Table 2. Ref. [35]. Ref. [26].

3.3. Correlation of gas solubility with solute-solvent interactions The equilibrium constant of solubility is controlled by the standard Gibbs energy changes of the compound dissolution in the solvent, which in turn is related to the Henry’s constant of the gaseous compound in the solvent (Eq. (7)). Certainly, physiochemical

DG ¼ 27:350ð0:601Þ  8:271ð0:913Þp  6:592ð0:642Þb

n ¼ 11; r2 ¼ 0:991; s ¼ 0:119;

F ¼ 432:67

ð9Þ

where n, r2, s, and F are the number of data points, square of regression coefficient, standard deviation, and statistical Fischer number, respectively. Eq. (9) is the best multi-parameter regression in which the hydrogen bond donor ability of the solvent shows no effect on solubility. The LFER investigation indicates that the polarity/polarizability and hydrogen bond acceptor ability of the solvent are the main factors in determining solvent effects on the CO2 solubility in this binary mixture. The negative sign of these parameters in the equation shows that the standard Gibbs energy of CO2 solution decreases with the increasing polarity/polarizability and hydrogen bond acceptor of media. Hence the solubility of CO2 is best in a solvent with high polarity/polarizability or hydrogen bond acceptor ability. On the other hand, according to preferential solvation discussion in our previous report [31], the presence of high solvent–solvent interactions in DMSO-EG mixtures cause that the solvent mixture as a new media is more polarizable around the mole fraction of 0.1–0.4 with respect to EG compared with each pure solvent. These observations are compatible with the present results and it confirms that the solute–solvent interactions based on Kamlet-Taft parameters can thermodynamically control the solubility of CO2.

A.R. Harifi-Mood / J. Chem. Thermodynamics 141 (2020) 105968

4. Conclusions Solubility studies of carbon dioxide in binary mixtures DMSOEG at different pressure and temperature showed that the solubility of CO2 in the solvent increase with the increasing of mole fraction of DMSO. The thermodynamic properties of the formation of gas solution, such as enthalpy, entropy, and standard Gibbs energy changes, confirmed that the process was exothermic, along with decreasing entropy, and thermodynamically non-spontaneous. The solubility of CO2 in term of Henry’s constant or Gibbs energy was correlated using LFER analysis. Obtained results indicated the direct effects of polarity/polarizability and hydrogen bond acceptor ability of solvent on the solubility of CO2. It seems that the solvent media consisting of high polarity/polarizability and hydrogen bond acceptor ability can prepare a favourable media for dissolving the carbon dioxide. Declaration of Competing Interest The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper. Acknowledgment The author gratefully acknowledge the financial support of Kharazmi University. (Grant Number: D/2043). References [1] A. Meisen, X. Shuai, Research and development issues in CO2 capture, Energy Convers. Manag. 38 (2002) S37–S42, https://doi.org/10.1016/s0196-8904(96) 00242-7. [2] M. Wang, A. Lawal, P. Stephenson, J. Sidders, C. Ramshaw, Post-combustion CO2 capture with chemical absorption: a state-of-the-art review, Chem. Eng. Res. Des. 89 (2011) 1609–1624, https://doi.org/10.1016/j.cherd.2010.11.005. [3] M. Wang, A.S. Joel, C. Ramshaw, D. Eimer, N.M. Musa, Process intensification for post-combustion CO2 capture with chemical absorption: a critical review, Appl. Energy. 158 (2015) 275–291, https://doi.org/10.1016/j. apenergy.2015.08.083. [4] A. Samanta, A. Zhao, G.K.H. Shimizu, P. Sarkar, R. Gupta, Post-combustion CO 2 capture using solid sorbents: a review, Ind. Eng. Chem. Res. 51 (2012) 1438– 1463, https://doi.org/10.1021/ie200686q. [5] W.A. Akber Hassan, X. Jiang, Corn ethanol growth in the USA without adverse foreign land-use change: define, Greenh. Gases Sci. Technol. 2 (2012) 408–418, https://doi.org/10.1002/ghg. [6] N. Gargiulo, F. Pepe, D. Caputo, CO2 adsorption by functionalized nanoporous materials: a review, J. Nanosci. Nanotechnol. 14 (2014) 1811–1822, https://doi. org/10.1166/jnn.2014.8893. [7] K. Simons, Membrane technologies for CO2 capture, J. Memb. Sci. 359 (2010) 115–125. [8] G. Manzolini, S. Campanari, P. Chiesa, A. Giannotti, P. Bedont, F. Parodi, CO2 Separation From Combined Cycles Using Molten Carbonate Fuel Cells, J. Fuel Cell Sci. Technol. 9 (2011), https://doi.org/10.1115/1.4005125 011018. [9] H.W. Pennline, E.J. Granite, D.R. Luebke, J.R. Kitchin, J. Landon, L.M. Weiland, Separation of CO2 from flue gas using electrochemical cells, Fuel. 89 (2010) 1307–1314, https://doi.org/10.1016/j.fuel.2009.11.036. [10] X. Gui, Z. Tang, W. Fei, CO2 capture with physical solvent dimethyl carbonate at high pressures, J. Chem. Eng. Data. 55 (2010) 3736–3741, https://doi.org/ 10.1021/je1002708. [11] T. Filburn, J.J. Helble, R.A. Weiss, Development of Supported Ethanolamines and Modified Ethanolamines for CO 2 Capture, Ind. Eng. Chem. Res. 44 (2005) 1542–1546, https://doi.org/10.1021/ie0495527. [12] E.D. Bates, R.D. Mayton, I. Ntai, J.H. Davis, CO 2 Capture by a Task-Specific Ionic Liquid, J. Am. Chem. Soc. 124 (2002) 926–927, https://doi.org/ 10.1021/ja017593d. [13] Y. Marcus, Solvent Mixtures: Properties and Selective Solvation, CRC Press, 2002.

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JCT 2019-621