Experimental Study on the Solubility of Carbon Dioxide in Systems

Article pubs.acs.org/jced

Experimental Study on the Solubility of Carbon Dioxide in Systems Containing Ethane-1,2-diol + Water + Salt (Sodium Chloride or Calcium Carbonate) Fabiane Santos Serpa,† Reginaldo Scariot Vidal,† Joaõ Henrique Bernardo Amaral Filho,† Alexandre Ferreira Santos,‡ Jailton Ferreira do Nascimento,§ Camila Machado de Senna Figueiredo,§ Giancarlo Richard Salazar-Banda,† Gustavo Rodrigues Borges,† Cláudio Dariva,† and Elton Franceschi*,† †

Núcleo de Estudos em Sistemas Coloidais - PEP/PBI/ITP, Universidade Tiradentes, Av. Murilo Dantas, 300, Farolândia, Aracaju, Sergipe Brazil, 49032-490 ‡ Department of Chemical Engineering, Federal University of Parana (UFPR), Polytechnic Center, 82530-990, Curitiba, Paraná Brazil § PETROBRAS/CENPES/PDEP/TPP, Centro de Pesquisa e Desenvolvimento Leopoldo Américo Miguez de Mello, 21941-915, Rio de Janeiro, Brazil S Supporting Information *

ABSTRACT: The influence of salt concentration on the solubility of gases in aqueous solutions is of interest in many processes of the oil and gas processing industry. Here, we show the influence of the presence of different concentrations of sodium chloride or calcium carbonate on the solubility of carbon dioxide (CO2) in mixtures of monoethylene glycol (MEG) and water by using a gas expansion method. From the experimental data, the solubility and Henry’s constant were measured as a function of pressure (from 0.1 to 0.5 MPa) and temperature (from 298 to 333 K), in pure water, pure MEG and in its equimolar mixture, using concentrations of salt up to saturation for all the studied solvents. The experimental data indicate that the increase in temperature decreases the CO2 solubility in the liquid mixture, whereas pressure shows a positive linear influence on the dissolution of gas. The addition of salt reduced the CO2 solubility in all mixtures due to the salting-out effect.

1. INTRODUCTION

glycols originates an unwanted effect, although glycol is a better solvent for some gases than water.6−9 Therefore, the knowledge of the phase behavior of systems containing gases and electrolytes in the presence of MEG is vital for the correct design, operation, and improvement of MEG regeneration systems. It is necessary to know the thermodynamic behavior of carbon dioxide with solutions found in MEG regeneration to prevent mineral scale formation.10 In particular, the knowledge of the amount of carbon dioxide dissolved in MEG + water + salts mixtures is crucial, since its presence produces carbonate formation (salt precipitation) and decreases pH (corrosion).1,3,10 Carbonates may precipitate when the pressure decreases and CO2 is released. The presence of salts in formation waters influences the amount of CO2 dissolved in the aqueous phase and thus affects the solid−liquid equilibrium behavior of carbonates.11,12 Several investigations have been carried out to measure the CO2 solubility in aqueous NaCl or CaCO3 solutions12−17 and

The dissolution of gaseous compounds in aqueous electrolytes is an important issue for petroleum exploration and production operations. For example, carbon dioxide (CO2) decreases the solubility of salts in produced water, such as sodium chloride (NaCl) present in large amounts and calcium carbonate (CaCO3) which has low solubility in water. Thus, gas solubility studies are very important to ensure proper specifications and to prevent possible problems related to flow assurance due to scale formation and deposition of salts during the production of oil and associated water, as well as during the regeneration step of monoethylene glycol (MEG) in natural gas plants1 During the production of oil and gas, thermodynamic or kinetic inhibitors are injected to avoid the formation of hydrates. Gas hydrates are crystalline solids formed by water molecules (host) which trapped natural gas molecules (guest) under specific thermodynamic conditions.1,2 The use of inhibitors such as monoethylene glycol (MEG) for hydrate prevention is a common practice, particularly in deepwater operations; however, this practice enhances the possibility of scale formation and salts deposition during the production of oil; thus inhibiting proper flow assurance (salting out).3−5 Furthermore, the use of © XXXX American Chemical Society

Received: May 10, 2016 Accepted: October 28, 2016

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DOI: 10.1021/acs.jced.6b00381 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

salt solubility in aqueous CO2 solutions.18 However, the different conditions presented in the literature make difficult the reproduction of experimental data. Sandengen (2006)1 reported CO2 solubility in MEG (up to 0.99 mass fraction), NaCl (concentration 0.0−0.7 mol/kg H2O) solution at 298−363 K and 1 bar. Kan and Lu (2010)19 reported experimental data of carbon dioxide partitioning in MEG + H2O + NaCl containing MEG (from 0 to 99 mass%), NaCl (from 0 to 6 mol/kg H2O), pressures up to 0.27 atm of CO2 partial pressure at 276−343 K. Pssarou et al. (2011)10 studied the solubility of carbon dioxide in systems containing MEG (from 60 to 98 mass%), and at low CO2 partial pressure from 353 to 413 K. Studies based on the solubility of carbon dioxide in aqueous NaCl and CaCO3 solutions containing, or not, MEG have been reported in the literature.13,17,19 However, these papers mainly focused on high temperature and pressure conditions, but the gas solubility at moderate temperature (298−333) K and pressure (0.1−0.5) MPa also affects the MEG reclaiming and regeneration process. Data to determine these equilibria at moderate temperature and CO2 pressure, high MEG content, containing salts (NaCl and CaCO3) are not available in the literature. Thus, we report the effect of the kind of electrolyte (sodium chloride and calcium carbonate) on the experimental solubility of carbon dioxide in aqueous solutions in the presence of MEG in temperatures from 298 to 333 K, and pressures ranging from 0.1 to 0.5 MPa.

connected to the mobile part, a vacuum is made in all lines and gas reservoir, and the mobile part is inserted into the thermostatic bath at the temperature of the experiment. After this, the mobile part (lines and CO2 reservoir) is pressurized with CO2 at a specified pressure. As pressure is stabilized, CO2 is admitted to the equilibrium cell by opening a ball valve located at its entrance. Then CO2 solubilizes into the liquid solution until the pressure of the system stabilizes, indicating the equilibrium condition. Details of the experimental apparatus and procedure are presented in Serpa et al. (2013).9 All solubility data were measured in a temperature range of 298 to 333 K, molar fractions of MEG + H2O of 0.0, 0.5, and 1.0, and different salt concentrations in the solutions. Regarding the CO2 solubility measurements in solutions containing CaCO3, every care was taken to avoid the precipitation of this salt. As the addition of MEG into the aqueous salt solution can lead to deposition of CaCO3, a visual observation of the equilibrium cell content was made during all procedure prior to the connection of the equilibrium cell containing the solution to the mobile part ensuring no salt precipitation. This procedure was conducted in a short time, but enough to ensure the quality of the experiment. After the equilibrium cell connection to the mobile part, CO2 was admitted to the equilibrium cell and dissolved into the liquid solution leading to an increase of the CaCO3 solubility in the system due to a decrease in pH of the solution. Hence the solutions become undersaturated. The thermodynamic model used was described in our previous paper,9 and schemes are presented in the literature. The validity of Henry’s law to the low solubility component is admitted, since CO2 has limited solubility in mixed solvents of interest.20−22 The corresponding state theorem was used to determine the partial molar volume of carbon dioxide dissolved in the liquid phase.23,24 Thermodynamic properties such as vapor phase fugacity coefficient, entropy, and enthalpy in excess of the mixture were also calculated using straightforward thermodynamics.25,26

2. EXPERIMENTAL SECTION 2.1. Chemicals. All chemicals employed in the study of gas−liquid solubility, and their purity, are displayed in Table 1. All were used without additional treatment. Table 1. Chemicals, Their Purity and Suppliers, Employed in the Study of Gas-Liquid Solubility chemical water carbon dioxide sodium chloride calcium carbonate monoethyene glycol

purity (wt %) ultrapure 0.999 0.999 0.999 0.995

supplier

3. RESULTS AND DISCUSSION The solubility of CO2 in MEG + H2O + NaCl and MEG + H2O + CaCO3 mixtures was determined. The mixtures containing dissolved amounts of salt were prepared gravimetrically with a composition conveniently chosen. The experimental procedure used for measurements of solubility of CO2 in MEG + H2O + salt mixtures was based on our previous report.9 Table 2 displays the solubility data for CO2 in MEG + water + NaCl and MEG + water + CaCO3 systems in diverse mole fractions of MEG + H2O (0.0, 0.5, and 1.0 in CO2), from 298 to 333 K, using different salt concentrations in the solution and pressures up to 0.5 MPa. Figure 1 (A and B) depicts the solubility of CO2 in aqueous NaCl solutions, in the 0.1−0.5 MPa pressure range and at different temperatures (from 298 to 333) K. The salt concentrations in the solution were 0.6 and 1.1 mol of salt per kg of water. When pure MEG was used as solvent (x5 = 1.0) the same concentrations of salt were used and correspond, respectively, to 50% and 90% of NaCl saturation in pure MEG.27 Figure 2 presents the experimental data for carbon dioxide solubility in aqueous CaCO3 solutions, at pressures up to 0.5 MPa, and temperatures from 298 to 333 K. The salt concentration in the solution was 0.00012 mol of salt per kg of water, and corresponds to concentration of CaCO3 saturation in pure water at 293 K.

White Martins Merck Merck Merck

2.2. Apparatus and Procedures. The experimental methodology employed here to measure the gas−liquid solubility data with electrolytes was already described by Dalmolin et al. (2006)20 and Serpa et al. (2013).9 Thermodynamic laws were used to obtain the solubility data.21,22 Briefly, the experimental apparatus is composed of a “mobile part”, CO2 cylinder, and a vacuum pump (Edwards, model RV3). The carbon dioxide cylinder is connected to a pressure regulating valve, which enables regulation of the pressure of the gas that is admitted to the mobile part of the experimental apparatus. The mobile part of the unit comprises a pressure transducer (Ter−Press, model TP/ST18, accuracy of 0.015 bar), a gas reservoir, and a borosilicate glass cell, besides a set of specific pipes and valves. For temperature control the mobile part remains inserted into a thermostatic water bath (Julabo, model MC) during the entire experiment. Experimental procedure begins by degassing the solutions. The mass of the solution is computed as the difference of the mass of the equilibrium cell with solution before and after degassing. After the solution is degassed, the equilibrium cell is B



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Table 2. Experimental Solubility (x) and Henry’s (H) Constants for CO2 (1) in mixtures of water (2) + NaCl (3) + MEG (5) and water (2) + CaCO3 (4) + MEG (5). x5 Values in the Solvent Mixture on a Free CO2 Basisa 298 K P/MPa mNaCl = 0.6 mol·kg

x5 = 0.0

x5 = 0.5

308 K x1 × 102

water H1,m/MPa = 287.5 0.1172 0.0243 0.2152 0.0677 0.3152 0.1193 H1,m/MPa = 260.8 0.1231 0.0272 0.2228 0.0787 0.3206 0.1315

H1,m/MPa = 304.0 0.1220 0.0243 0.2238 0.0673 0.3210 0.1141

mNaCl = 1.1 mol·kg−1 water H1,m/MPa = 0.1161 x5 = 1.0 0.2143 0.3069 0.4065 mCaCO3 = 0.00012 mol·kg−1 water H1,m/MPa = 0.1472 x5 = 0.0 0.2385 0.3220 0.4217 H1,m/MPa = 0.1130 x5 = 0.5 0.2130 0.3155 0.4145 H1,m/MPa = 0.1234 x5 = 1.0 0.2197 0.3151 0.4214 a

323 K x1 × 102

P/MPa

333 K x1 × 102

P/MPa

x1 × 102

−1

H1,m/MPa = 89.9 0.1149 0.0246 x5 = 1.0 0.1989 0.1768 0.3004 0.3590 0.3936 0.5351 mNaCl = 1.1 mol·kg−1 water H1,m/MPa = 312.2 0.1230 0.0201 x5 = 0.0 0.2227 0.0658 0.3196 0.1089

x5 = 0.5

P/MPa H1,m/MPa = 409.6 0.1212 0.2124 0.3098 0.4077 H1,m/MPa = 327.8 0.1143 0.2137 0.3094 0.4097 H1,m/MPa = 110.8 0.1135 0.2142 0.3060 0.4090 H1,m/MPa = 456.0 0.1146 0.2084 0.3108 0.4123 H1,m/MPa = 358.2 0.1177 0.2132 0.3120 0.4116

104.8 0.0227 0.1697 0.3137 0.4507

H1,m/MPa = 128.5 0.1162 0.2087 0.3083 0.4072

161.6 0.0864 0.1497 0.1931 0.2655 80.83 0.1374 0.2533 0.3896 0.5189 50.15 0.2334 0.4349 0.6470 0.8685

H1,m/MPa = 216.2 0.1877 0.2584 0.3303 0.4248 H1,m/MPa = 141.7 0.1294 0.2205 0.3239 0.4235 H1,m/MPa = 60.31 0.1303 0.2183 0.3169 0.4195

0.0082 0.0415 0.0794 0.1160 0.0083 0.0521 0.1004 0.1454 0.0256 0.1645 0.2933 0.4173

0.0069 0.0368 0.0713 0.1049 0.0073 0.0496 0.0917 0.1281

0.0242 0.1415 0.2448 0.3637

0.0802 0.1140 0.1507 0.2090 0.0875 0.1503 0.2274 0.2999 0.2066 0.3598 0.5317 0.7241

H1,m/MPa = 586.4 0.1425 0.2482 0.3459 0.4244 H1,m/MPa = 412.3 0.1380 0.2243 0.3209 0.4129 H1,m/MPa = 144.4 0.1214 0.2177 0.3162

H1,m/MPa = 629.2 0.1311 0.2210 0.3048 0.4033 H1,m/MPa = 515.2 0.1459 0.2376 0.3280

H1,m/MPa = 146.1 0.1272 0.2151 0.3169 0.4148 H1,m/MPa = 303.5 0.1615 0.2412 0.3262 0.4182 H1,m/MPa = 236.9 0.1492 0.2309 0.3292 0.4264 H1,m/MPa = 77.53 0.1374 0.2350 0.3239 0.4235

0.0069 0.0363 0.0569 0.0789 0.0073 0.0451 0.0744 0.1181 0.0149 0.1357 0.2460

0.0056 0.0288 0.0503 0.0675 0.0065 0.0414 0.0665

0.0117 0.1292 0.2161 0.3314

0.0471 0.0709 0.1074 0.1439 0.0589 0.0915 0.1373 0.1765 0.1536 0.3095 0.4387 0.5955

H1,m/MPa = 625.1 0.1336 0.2256 0.3226

0.0053 0.0327 0.0488

H1,m/MPa = 532.2 0.1251 0.2206 0.3265

0.0068 0.0348 0.0656

H1,m/MPa = 161.6 0.1269 0.2256 0.3133

0.0112 0.1262 0.2167

H1,m/MPa = 756.9 0.1272 0.2245 0.3225 0.4192 H1,m/MPa = 588.2 0.1367 0.2232 0.3227

H1,m/MPa = 181.9 0.1192 0.2099 0.3160

H1,m/MPa = 454.9 0.1657 0.2577 0.3334 0.4323 H1,m/MPa = 354.4 0.1370 0.2267 0.3390 0.4432 H1,m/MPa = 95.52 0.1423 0.2192 0.3195 0.4264

0.0043 0.0209 0.0466 0.0621 0.0053 0.0319 0.0586

0.0098 0.1084 0.1854

0.0201 0.0601 0.0853 0.1248 0.0241 0.0655 0.1066 0.1484 0.1334 0.2208 0.3541 0.4934

u(T) = 0.1 K, u(x) = 0.0001, u(P) = 0.02 MPa, u(m) = 0.00001 mol.kg−1 water; u is the standard uncertainty in the measurement.

The presence of NaCl or CaCO3 decreases the solubility of carbon dioxide in the solvent mixture. In other words, a saltingout effect occurs in the systems presented in Figures 1 and 2. As expected, the presence of salt decreases the solubility of the gas, once the water molecules interact with the ions dissociated in solution.12,14,16,28−31 Moreover, carbon dioxide is more soluble in systems containing CaCO3 than in NaCl solutions. The effect of the partial pressure of CO2 and, consequently, the pH

Figures 1 and 2 show that solubility increases with the increase of pressure for the studied systems for all temperatures and salt concentrations, as expected. Moreover, the CO2 solubility in solutions gradually decreases with increasing concentration of the salt for all temperatures and pressures.14−17 An inspection of Table 2 reveals that the same trends were observed when MEG is present in the solution (x5 of 0.5 and 1.0) for both NaCl and CaCO3 salts. C



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Figure 1. Experimental P−x data for the system CO2 (1) + water (2) + NaCl (3) at (□) 298 K, (◊) 308 K, (△) 323 K, and (○) 333 K: (A) [NaCl] = 0.6 mol·kg−1 water and (B) [NaCl] = 1.1 mol·kg−1 water.

Figure 2. Experimental P−x data for the system CO2 (1) solubility in water (2) + CaCO3 (4) as a function of temperature: (□) 298 K, (◊) 308 K, (△) 323 K, (○) 333 K, and CaCO3 concentration of 0.00012 mol·kg−1 water.

Consequently, the gas molecules release from the liquid phase. Furthermore, the addition of ionic compounds reduces the solubility of carbon dioxide in the solvent mixture. However, the solubility of carbon dioxide is more pronounced in

can modify the distribution of the different ionic species, thus favoring the dissolution of calcium carbonate.18 The higher the temperature of the system, the lower the gas solubility as a function of the kinetic energy increase of gas molecules. D



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Figure 3. Dependence of the solubility of CO2 (1) in water (2) + CaCO3 (4) + MEG (5) mixture at 298 K, and CaCO3 concentration of 0.00012 mol·kg−1 water. Mole fractions of MEG (x5) in the solvent mixture MEG + water are expressed on a CO2 free basis. (◊) x5 = 0.0, (□) x5 = 0.5, (△) x5 = 1.0.

Figure 4. Effect of NaCl and MEG on carbon dioxide solubility for the system CO2 (1) in water (2) + NaCl (3) + MEG (5) mixture at temperatures of 298.15 K (A), 308.15 K (B), 323.15 K (C), and 333.15 K (D). Mole fractions of MEG (x5) in the solvent mixture MEG + water are expressed on a CO2 free basis and NaCl concentration is in mol·kg−1 water. x5 = 0.0 (△, [NaCl] = 0.6; ▲, [NaCl] = 1.1); x5 = 0.5 (◊, [NaCl] = 0.6; ⧫, [NaCl] = 1.1); x5 = 1.0 (□, [NaCl] = 0.6; ■, [NaCl] = 1.1).

glycol as can be seen for the CO2 + water + MEG system.9 This effect is a result of the higher affinity of CO2 to MEG than to water.10 Solubility values of carbon dioxide in MEG + water + CaCO3 (0.00012 mol of salt per kg of water) system at 318 K, and MEG + water + NaCl (0.6 and 1.1 mol of salt per kg of water) systems in temperature range from 298 to 333 K are presented in Figures 3 and 4A−D, respectively, at different molar concentrations of MEG in water (0.0, 0.5, and 1.0 in CO2 free basis),

and pressures up to 0.5 MPa. Therefore, the higher the MEG concentration is, the higher is the CO2 solubility for both systems. The salt addition reduced the solubility of CO2 in the mixture. The addition of MEG dilutes the ionic strength directly.19 Moreover, CO2 is more soluble in systems containing CaCO3 (Figure 3) when compared to systems in the presence of NaCl (Figure 4). The amount of CaCO3 present in the solution is smaller encouraging the dissolution of carbon dioxide. MEG behaves ideally since the pressure is linear with composition. E



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For the system with NaCl, the increase in the MEG concentration into solution to 0.5 does not influence the CO2 solubility. But, it was observed a pronounced increase on the CO2 solubility when MEG molar fraction was increased from 0.5 to 1.0. According to Lu et al. (2010)32 the increase of MEG content in the aqueous salt solution causes a decrease of CO2 solubility to a minimum value and then, the same increases with increasing MEG content into solution, surpassing the minimum value, due to a dilution of the ionic strength. To clearly verify this effect more experimental data on the CO2 solubility in these systems further experiments need to be performed in different concentrations of MEG and NaCl into mixture of those investigated in this work. Data presented in Table 2 indicate that, as expected, the solubility of CO2 decreased with increasing temperature. The effect of temperature can also be observed through the Henry’s law constant. Table 3 summarizes the calorimetric properties of the systems. The enthalpy and entropy data for CO2 + MEG + water + salt solution were calculated from the relations described in a

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Δh̅/J·mol−1 −1

*E-mail: [email protected]. Funding

The authors thank to CENPES/PETROBRAS, FAPITEC, CAPES, CNPq (303630/2012-4) and UNIT for financial support and scholarships. Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Sandengen, K. Prediction of Mineral Scale Formation in Wet Gas Condensate Pipelines and in MEG (Monoethylene Glycol) Regeneration Plants. Ph.D. Thesis, Norwegian University of Science and Technology, Trondheim, Norway, 2006. (2) Sloan, E. D.; Koh, C. A. Clathrate Hydrates of Natural Gases; CRC Press: Boca Raton, FL, 2007. (3) Masoudi, R.; Tohidi, B.; Anderson, R.; Burgass, R. W.; Yang, J. Experimental Measurement and Thermodynamic Modeling of Clathrate Hydrate Equilibria and Salt Solubility in Aqueous Ethylene Glycol and Electrolyte Solutions. Fluid Phase Equilib. 2004, 219, 157− 163. (4) Amy, T. K.; Gongmin, F.; Malene, A. W.; Mason, B. T.; Effect of hydrate inhibitors on oilfield scale formation and inhibition, Int. Symp. Oilfield Scale, SPE-74657-MS, Aberdeen: United Kingdom, January 30−31, 2002.10.2118/74657-MS (5) Matthews, P. N.; Subramanian, S.; Creek, J. High Impact, Poorly Understood Issues with Hydrates in Flow Assurance. In Proc. 4th Int. Conf. Nat. Gas Hydrates, Yokohama, Japan, May 19−23, 2002. (6) Wang, L. K.; Chen, G. J.; Han, G. H.; Guo, X. Q.; Guo, T. M. Experimental Study on the Solubility of Natural Gas Components in Water with or Without Hydrate Inhibitor. Fluid Phase Equilib. 2003, 207, 143−154. (7) Galvão, A. C.; Francesconi, A. Z. Experimental study of methane and carbon dioxide solubility in 1,4 butylene glycol at pressures up to 11 MPa and temperatures ranging from 303 to 423 K. J. Supercrit. Fluids 2009, 51, 123−127. (8) Galvão, A. C.; Francesconi, A. Z. Solubility of Methane and Carbon Dioxide in Ethylene Glycol at Pressures up to 14 MPa and Temperatures Ranging From (303 to 423) K. J. Chem. Thermodyn. 2010, 42, 684−688. (9) Serpa, F. S.; Vidal, R. S.; Amaral Filho, J. H. B.; Nascimento, J. F.; Ciambelli, J. R. P.; Figueiredo, C. M. S.; Salazar-Banda, G. R.; Santos, A. F.; Fortuny, M.; Franceschi, E.; Dariva, C. Solubility of Carbon Dioxide in Ethane-1,2-diol-Water Mixtures. J. Chem. Eng. Data 2013, 58, 3464−3469. (10) Psarrou, M. N.; Josang, L. O.; Sandengen, K.; Ostvold, T. Carbon Dioxide Solubility and Monoethylene Glycol (MEG) Degradation at MEG Reclaiming/Regeneration Conditions. J. Chem. Eng. Data 2011, 56, 4720−4724. (11) García, A. V.; Thomsen, K.; Stenby, E. H. Prediction of Mineral Scale Formation in Geothermal and Oilfield Operations Using the Extended UNIQUAC Model. Part II. Carbonate-Scaling Minerals. Geothermics 2006, 35, 239−284. (12) Yan, Y.; Chen, C. C. Thermodynamic Modeling of CO2 Solubility in Aqueous Solutions of NaCl and Na2SO4. J. Supercrit. Fluids 2010, 55, 623−634. (13) Kiepe, J.; Horstmann, S.; Fischer, K.; Gmehling, J. Experimental Determination and Prediction of Gas Solubility Data for CO2 + H2O Mixtures Containing NaCl or KCl at Temperatures Between 313 and 393 K and Pressures up to 10 MPa. Ind. Eng. Chem. Res. 2002, 41, 4393−4398.

−1 −Δs/J·(mol·K) ̅

CO2 + Water + MEG + NaCl (0.6 mol·kg Water) 0.0 18648.3 128.9 0.5 16145.8 119.7 1.0 14832.2 99.8 CO2 + Water + MEG + NaCl (1.1 mol·kg−1 Water) 0.0 20460.8 135.5 0.5 16253.9 121.4 1.0 12088.6 98.1 CO2 + Water + MEG + CaCO3 (0.00012 mol·kg−1 Water) 0.0 23437.2 140.5 0.5 33912.8 169.6 1.0 14990.1 102.3

previous work.9 The values of these properties for carbon dioxide solubility in pure water are greater than for pure MEG.9 The CO2 solubility in MEG results from opposite effects of enthalpy and entropy, while these effects decrease by increasing temperature.7−9 However, for the system CO2 + MEG + water + CaCO3 it was verified that the enthalpy and entropy values do not follow the same trend observed for the other systems. This is because there are two related effects: the dissolution of CO2 favors the dissolution of CaCO3 in the mixture, and the addition of MEG induces precipitation of the salt.

4. CONCLUSIONS Solubility data for CO2 as a function of temperature, pressure, ratios of MEG + water, and salt concentration as well as the kind of salt were investigated in this study. Pressure and temperature have positive and negative effects, respectively, on the solubility of carbon dioxide in the liquid phase. The addition of MEG in the liquid mixture increases the amount of CO2 dissolved in the solution, whereas the solubility of CO2 decreases with increasing salt concentration in the mixture (salting-out effect).

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AUTHOR INFORMATION

Corresponding Author

Table 3. Values of the Enthalpy (Δh̅) and Entropy (−Δs)̅ of Solution for CO2 in MEG + Water + Salt Mixtures in the Temperature Range from (298 to 333) K xMEG

Dependence of carbon dioxide solubility in function of pressure, temperature and composition of NaCl (PDF)

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.6b00381. F



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Experimental Study on the Solubility of Carbon Dioxide in Systems

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