Measurement of CO2 Solubility in NaCl Brine Solutions at Different

Jun 24, 2015 - ABSTRACT: The solubility of CO2 in brine is one of the trapping mechanisms by which the CO2 is sequestrated in aquifers. In this resear...
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Measurement of CO2 Solubility in NaCl Brine Solutions at Different Temperatures and Pressures Using the Potentiometric Titration Method Erfan Mohammadian,*,†,‡ Hossein Hamidi,† Mohammad Asadullah,† Amin Azdarpour,†,‡ Shervin Motamedi,§ and Radzuan Junin‡ †

Oil and Gas Department, Faculty of Chemical Engineering (FKK), Universiti Teknologi Mara (UiTM), Shah Alam, Selangor Malaysia ‡ Faculty of Petroleum and Renewable Energy (FPREE), Universiti Teknologi Malaysia (UTM), Johor Bahru, Malaysia § Department of Civil Engineering, Faculty of Engineering, University of Malaya, 50603, Kuala Lumpur, Malaysia ABSTRACT: The solubility of CO2 in brine is one of the trapping mechanisms by which the CO2 is sequestrated in aquifers. In this research, an unconventional method, called the potentiometric titration, was used to obtain the solubility of CO2 in distilled water and NaCl brine. The solubility data for the low salinity range are scarce in the literature. Thus, in this research, the CO2 solubility was obtained in NaCl brines of low salinity (0−15 000 ppm) at temperatures between 60 °C and 100 °C and pressures up to 25 MPa. Moreover, the salting-out effect was estimated using the Setchenov’s constant as a measure of reduction in solubility when salt is added to the solution. The solubility points obtained by the potentiometric titration method demonstrated very good consistency with those obtained by the previous methods.

1. INTRODUCTION Since the industrial revolution, burning of fossil fuels has released enormous amounts of carbon dioxide (CO2) into the atmosphere. This increase in the concentration of carbon dioxide is one of the major causes of global warming. Various methods, such as the CO2 sequestration in geological formations (depleted oil and gas reservoirs,1,2 coal beds,3 and deep saline aquifers4), ocean storage,5 and mineral carbonation6−8 have been proposed to mitigate the concentration of CO2 in the atmosphere. Large theoretical capacities as well as a common occurrence of saline aquifers make them suitable repositories for the sequestration of CO2. It has been estimated that the world’s aquifers may provide a storage capacity of up to 11 × 1012 tonnes of CO2, which is enough to store CO2 emission for several hundred years.9 Aquifers with a depth above 800 m are required, in which CO2 can be injected as a supercritical fluid to minimize the chance of leakage from the underground repository of interest. The solubility of CO2 in the formation brine is one of the trapping mechanisms for sequestrating CO2 in saline aquifers. Structural, geochemical, soluble, and capillary are other trapping mechanisms. These mechanisms have been studies for decades.4,10 Besides CO2 sequestration, the knowledge of CO2 solubility in water (or brine) has been used in a number of industries, such as food, fertilizer, oil, and gas (enhanced oil recovery methods). Therefore, there are diverse thermodynamic conditions for obtaining CO2 solubility data points. © 2015 American Chemical Society

Abundant data are available on the solubility of CO2 in brine, particularly at low pressures.11,12 A very comprehensive review of low pressure solubility data of CO2 in water was conducted by Carroll et al.,11 in which an equation of Henry’s constant was regressed based on a low pressure database that the authors had developed from the previous studies. For the sequestration studies, high pressure data are required because CO2 is injected as a supercritical fluid. One of the first studies on the highpressure solubility data was conducted by Wiebe and Gaddy.13 They obtained the solubility of CO2 in water at temperatures up to 100 °C and pressures up to 71 MPa. Since then, the solubility of CO2 in water has been studied in a variety of P−T conditions using similar measurement methods.14,15 More recently, Hou et al.16 presented highly accurate experimental data on the solubility of CO2 in water at pressures up to around 18 MPa, using gas chromatography. NaCl is considered as the major component of many formation brines. Therefore, many studies have been conducted on the solubility of CO2 in NaCl brines, which are basically a binary mixture of distilled water and NaCl. Malinin and Savelyeva17 were pioneers in the investigation of solubility of CO2 in NaCl solutions at temperatures of 25 °C, 50 °C, and 74 °C and elevated CO2 pressures using the depressurization volumetric method. Drummond18 measured more than 500 Received: December 28, 2014 Accepted: June 16, 2015 Published: June 24, 2015 2042

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Figure 1. Schematic diagram of experimental setup used for solubility experiments. (1) CO2 cylinder, (2) valve, (3) ISCO pump, (4) pressure gauge, (5) thermocouple probe, (6) reactor, (7) dip tube, (8) magnetic bead, (9) magnetic stirrer, (10) floating piston sampler, (11) syringe pump, (12) exhaust line.

accurate results. These methods are often time-consuming; therefore, efforts have been made to accelerate and improve the measurement techniques.23 For the first time, Portier and Rochelle27 used the potentiometric titration to measure the solubility of carbon dioxide. However, the data point obtained using this method was limited to specific formation brine, that is, that of the Utsira pore water. Therefore, their work did not cover a wide range of pressures, temperatures, and salinity. In this research, the unconventional method of potentiometric titration was used to obtain the solubility of CO2 in brine. The aforementioned method has been widely used in chemical engineering applications, but very rarely used in the CO2 sequestration projects. Moreover, since there is very limited data in the literature on the solubility of CO2 in the low salinity brines, the solubility was obtained in NaCl brines of low salinity (0 wt % to 1.5 wt %) in the current study. The saltingout effect is discussed in details as a measure of reduction in solubility by increasing the salinity.

data points of CO2 solubility in NaCl brines of different salinities. This work could have the most comprehensive data sources in the solubility database, but because too many rough assumptions were used in Drummond’s calculations, the accuracy of this data is doubtful. Since then, many researchers have studied the solubility of CO2 in NaCl brines in a wide range of pressures, temperatures, and salinity.19−24 However, there are still gaps in the research that should be addressed. For instance, although a wide range of salinities has been covered, there is limited data on the solubility in the low salinity range, for example, in the range of 0 ppm to 15 000 ppm (0 mol·kg−1 to 0.258 mol·kg−1) of brine salinity. One example of low salinity formation brines is the one encountered in the Malay basin, offshore Peninsular Malaysia, with an average salinity of ∼1 wt %.25 Few methods have been applied to measure CO2 solubility in water (or brine). Pepper and Dohren26 extensively reviewed the sampling techniques to measure CO2 solubility. With the exception of the works by Koschel et al.20 and Portier and Rochelle,27 other studies used either of these two methods, that is, the volumetric method and the combination method. The volumetric method is based on measuring weight or volume of a subsample of CO2 solution withdrawn from a high pressure reactor. The samples often undergo several stages of depressurization, and the solubility is estimated based on the cumulative CO2 released throughout the process. In the second method, which is very similar to the first one, a visual cell is used to observe the phase changes and help obtain more

2. MATERIALS AND METHODS 2.1. Materials. CO2 (99.99 % purity) supplied by the SIG was used throughout the experiments. Distilled and deionized water (Milli-Q) with resistivity of 18.2 Ω was used to prepare brines of various salinities. NaCl with mass fraction purity of 0.995 was supplied by Systerm. All chemicals were utilized without further purification and their purities were obtained from suppliers. 2043

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2.2. Methods. An experimental setup was established to investigate effects of pressure, temperature, and salinity on the solubility of CO2 in brine. The combined uncertainty value of data was treated with analysis of variance (ANOVA). The effect of salinity concentrations (1000, 10 000, and 15 000) ppm (equivalent to molality values of (0.0172, 0.1720, 0.2580) mol· kg−1, respectively) is the variance “among groups”. The variance “within groups” covers only the repeatability of the measurement. The experimental setup is illustrated in Figure 1. It mainly consisted of an ISCO pump made by Teledyne (model 100 DX), CO2 cylinder, and a 100 mL autoclave reactor equipped with a magnetic stirrer. The reactor was enclosed in an electric heater. A 1/8 in. dip tube was connected to a floating piston sampler fabricated locally, which was controlled by a syringe pump manufactured by Kdscientific (model 53100). The dip tube was used for the sampling of the CO2 saturated brine solution. The pressure in CO2 cylinder was 8.25 MPa. The pressure of the experiments was adjusted via the ISCO pump. The autoclave reactor was placed in a temperature-adjustable electric oven with precision of 0.1 °C according to the device manual. Despite the 0.85 cm thickness of the reactor base, proper “coupling” between the magnetic stirrer and the stirrer bead was achieved and resulted in a well-mixed solution. The brine was injected into the reactor and preheated to the desired temperatures. The CO2 was then injected at the pressure of interest into the preheated reactor, 70 mL of which was filled with the brine. The inlet and outlet valves of the reactor were then closed, and the solution was stirred for 3 h until the equilibrium was reached. The duration to reach equilibrium has been reported to be as low as 10 min up to 24 h in the previous studies.23,26−28 Afterward, the valve at the bottom of the reactor was opened very carefully in order to diminish the pressure changes in the system. A sample of CO2-saturated brine was withdrawn from the reactor via the dipping tube all the way to the sampling chamber, which was equipped with a floating piston. As the sample reached into the sampling chamber, it reacted with 0.5 M NaOH that had filled half of the chamber, approximately 3.5 mL, to preserve any sort of dissolved carbon species in the sample. The dipping tube and sampling chamber were designed in a manner to minimize the possibility of the gas flow into the sampler. The other half of the sampling chamber was filled with the distilled brine, which was gradually withdrawn to allow the CO2 saturated brine reacting with the NaOH. The withdrawal of brine was achieved by draining, using a syringe pump set at a constant rate of 1 mL/min. The reaction between the solution and the NaOH was given 10 min before it was withdrawn. During the withdrawal, no gas bubble was generated because all types of the carbon species (including HCO−3 and H2CO3) were dissolved in the NaOH and converted to CO2− 3 as long as excess amounts of NaOH were available.27,29 The NaOH preserved samples were then analyzed through the potentiometric titration method. The titrant, hydrochloric acid (HCl) was added to the analyte (NaOH and CO32− mixture) until the equivalence point was achieved. The equivalence point of titration can be identified via various methods, such as potentiometric (as is the case in this research), indicators, and conductivity methods. As the end point was reached, the volume of consumed reactant was measured and used to determine the concentration of analyte through the following correlation:

Ca =

C t × Vt × N Ma

(1)

where Ca is the concentration of the analyte, which is equal to the CO2 solubility in brine, in molal units (mol·kg−1); Ct is the concentration of the titrant (HCl), typically in mol·L−1 units; Vt is the volume of the titrant used, in mL; N is the mole ratio of the analyte and reactant from the balanced chemical equation; and Ma is the mass of the sample that is titrated in grams. The advantage of using mass of solution (Ma) as opposed to the volume of solution Va, which has been more commonly used in the previous researches, is that the mass of solution is not a function of pressure or temperature. Therefore, the uncertainty in determining the solubility is lower. Then, a 5 mL sample reacted with 0.5 M HCl, which was used as the titrant. The sample’s pH was observed as a function of volume of added titrant, and the titration continued until a pH of less than 2 was achieved. The equivalence points were obtained using the derivative curves of titrant volume versus pH. One of the main advantages of titration method for measuring the solubility over the previous methods is that preservation of the samples prevents CO2 loss due to the degassing during the depressurization phase as well as technical simplicity. Moreover, unlike in the previous studies, the need for estimation of parameters, that is, the fugacity and the density via correlations is eliminated.

3. RESULTS AND DISCUSSION Despite several researches on the solubility of CO2 in brine solution, the solubility data in low salinity brines is still limited.

Figure 2. Comparison among the solubility results of this study (×) with those of Duan and Sun29 (■) and Li et al.21 (▲) in distilled water (zero salinity) at 60 °C. The solid line represents the regression line fitted to the solubility results of the current study.

Therefore, in this study, the effects of pressure change on the solubility of CO2 were investigated under various conditions. The experiments were conducted in pressure ranges of 1 MPa to 25 MPa and temperature ranges of 60 °C to 100 °C. Moreover, the experiments were conducted in brine solution of different salinity to ensure consistency of results in conditions more representatives of saline aquifers. Unusual data, such as abnormally high or low data points were considered as outliers and omitted using the data analysis software (SPSS 18). The higher or lower solubility values were mainly caused by early opening of the sampling valve, resulting in the pressure drop in the solution or breakthrough of supercritical CO2. To test reproducibility of the results, the CO2 solubility was measured for three times in 1000 ppm brine at 60 °C and 10 2044

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Figure 3. Solubility of CO2 in distilled water at 60 (⧫), 80 (■), and 100 °C (▲) versus pressure.

Figure 6. Solubility of CO2 in 15 000 ppm of NaCl brine at temperatures of ⧫, 60 °C; ■, 80 °C; and ▲, 100 °C versus pressures.

Figure 4. Solubility of CO2 in 1000 ppm of NaCl brine at temperatures of ⧫, 60 °C; ■, 80 °C; and ▲, 100 °C versus pressures.

Figure 7. Solubility of CO2 in NaCl brine of various salinities at 80 °C. ⧫, 1000 ppm brine; ■, 10 000 ppm; and ▲, 15 000 ppm.

Figure 5. Solubility of CO2 in 10 000 ppm of NaCl brine at temperatures of ⧫, 60 °C; ■, 80 °C; and ▲, 100 °C versus pressures.

Figure 8. Salting-out effect of NaCl brine of ⧫, 1000 ppm; ■, 10 000 ppm; and ▲, 15 000 ppm (equivalent to molalites of (0.0172, 0.1720, and 0.258) mol·kg−1) at 80 °C.

MPa ((0.929 ± 0.017) mol·kg−1). Similarly, the CO2 solubility in distilled brine was tested at a pressure of 15 MPa and temperature of 80 °C. The solubility at this condition was (1.007 ± 0.020) mol·kg−1. Low standard deviation in the first and second cases (ur(X) = 0.017 p and 0.020 p, respectively) demonstrates that the potentiometric titration method is quite reproducible. To verify the experimental method, CO2 solubility data points obtained in this study were compared to those of previous studies in similar conditions as shown in Figure 2. It is apparent that the obtained results in this research are close to those obtained by the previous studies and more commonly used methods, such as depressurization21 and combination.28 The solid line in the figure depicts the regression line calculated

in this research. The error bars illustrate 5 % deviation from the solubility data obtained in the current research. It should be noted that the data points obtained by Li et al.21 were measured at 28 °C. The slight difference between the data obtained in this research and those of Li et al.21 is most probably due to the small difference in the experimental temperature of this research, that is, 30 °C as opposed to 28 °C in Li et al.21 By comparing the data points obtained in the current study to those of Duan and Sun,29 a very good consistency can be observed, specifically at higher pressures. Therefore, it can be concluded that the potentiometric titration method is both precise and reliable. 2045

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Table 1. Experimental Values of CO2 Solubility x0 in Distilled Water at Pressure P and Temperature Ta P/MPa

T/°C

x0/(mol·kg−1)

0.1 2.1 4.1 6.2 8.3 10.3 13.8 15.8 17.2 19.3 21.3 0.1 2.1 4.1 6.2 8.3 10.3 13.8 15.8 17.2 19.3 21.3

60 60 60 60 60 60 60 60 60 60 60 80 80 80 80 80 80 80 80 80 80 80

0.014 0.315 0.572 − 0.929 1.040 1.144 1.174 1.193 1.221 − 0.007 0.245 0.455 0.630 0.750 0.884 1.007 1.077 1.097 1.140 1.170

Table 2. Experimental Values of CO2 Solubility x in Aqueous NaCl Solutions and Calculated Setchenov Constant Ke at 60 °C and Pressure P, where m Denotes Molality of NaCl in Watera

Combined expanded uncertainties are U(T) = 0.05 °C, U(P) = 0.02 MPa, and U(x0) = 0.050 mol·kg−1.

a

3.1. Effects of Pressure and Temperature. The solubility of CO2 in distilled water is illustrated in Figure 3 at three different temperature series, namely (60, 80, and 100) °C. The error bars represent 5 % deviation from the base (60 °C). Regarding the pressure, it is apparent from the figure that the rise in pressure increases the solubility of CO2 in brine solution, regardless of the temperature. It can be seen that pressure dependency of carbon dioxide diminishes as pressure increases in all three temperature series. However, in this research, no point was observed on which the solubility became completely insensitive to pressure. Effects of pressure on the solubility can be explained from the Henry’s law of solubility (partial pressures) as well as la Chatelier’s principle.29 Regarding the temperature, it can be observed that an increase in temperature causes a reduction in the solubility of CO2 in brine. For instance, the solubility below 17 MPa and 60 °C is 1.193 mol·kg−1, whereas under the same pressure and at 80 °C and 100 °C, the solubility is 1.097 mol·kg−1 and 1.037 mol·kg−1, respectively. In other words, there are 8 % and 13 % reductions in the solubility as temperature increases to 80 °C and 100 °C from an initial value of 60 °C. A decrease in the solubility with increase in temperature has been reported by previous researchers as well.23,29,30 The solubility of CO2 in (1000, 10 00, and 15 000_ ppm ((0.0172, 0.1720, 0.2580) mol·kg−1, respectively) NaCl solutions is presented in Figures 4, 5, and 6. The error bars in these three figures illustrate 10 % deviation from the baseline (60 °C). It can be seen that in all three figures, regardless of the salinity and temperature, that increased pressure results in higher amounts of CO2 dissolving in the brine. In all three cases ((1000, 10 000 and 15 000) ppm), the pressure dependency of CO2 solubility was reduced at higher pressures. On the other hand, an increase in the temperature reduced the solubility. For instance, in 10 000 ppm brine (Figure 5) as

P/MPa

m/mol·kg−1

x/mol·kg−1

Ke

0.1 2.1 8.3 10.3 13.8 17.2 19.3 21.3 0.1 2.1 4.1 6.2 8.3 13.8 15.8 17.2 19.3 21.3 0.1 2.1 4.1 6.2 8.3 13.8 15.8 17.2 21.3

0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.2580 0.2580 0.2580 0.2580 0.2580 0.2580 0.2580 0.2580 0.2580

0.007 0.250 0.924 1.037 1.140 1.189 1.216 1.225 0.006 0.238 0.549 0.722 0.866 1.104 1.111 1.151 1.174 1.188 0.004 0.201 0.530 0.674 0.740 0.998 1.009 1.011 1.161

16.95 5.83 0.15 0.07 0.09 0.08 0.10 − 2.08 0.71 0.10 − 0.18 0.09 0.14 0.09 0.10 − 20.72 7.56 1.28 − 3.83 2.30 2.55 2.79 -

a Combined expanded uncertainties are U(P) = 0.02 MPa, U(T) = 0.05 °C, U(m) = 0.0020 mol·kg−1, U(x) = 0.050 mol·kg−1 and U(Ke) = 0.03.

the temperature increased to 100 °C from the initial value of 60 °C, between 5 % and 24 % reductions in the solubility were observed along the pressure range of the experiments. At the salinity of 15 000 ppm (Figure 6), reduction in the solubility for the same increment of temperature (from 60 °C to 100 °C) along with the pressure range of the experiments, was between 5 % and 18 %. Overall, it is apparent that pressure and temperature sensitivity is higher in conditions where CO2 is in subcritical (liquid or gas) rather than supercritical state. A rise in temperature leads to an increase in the kinetic energy, resulting in more rapid motion between the molecules and breakage of intermolecular bonds. This enables molecules to escape to the gas phase.31 Therefore, in these sets of experiments, increased temperature causes reduced solubility, regardless of the brine’s salinity and pressure. The experimental data of this section can be compared to those of Nighswander et al,32 obtained by the depressurization method in 1 wt % NaCl brine. A very good consistency among the data points can be seen at 80 °C. However, the comparison was not possible at higher isotherms because of the difference in the experimental temperatures (nonidentical temperature points) of the current study and those of Nighswander et al.32 3.2. Effects of Salinity on Solubility of CO2. Figure 7 illustrates the solubility of CO2 versus pressures at different salinities at 80 °C. It is apparent that solubility of CO2 decreases with increased salinity of the brine under different 2046

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Table 3. Experimental Values of CO2 Solubility x in Aqueous NaCl Solutions and Calculated Setchenov Constant Ke at 80 °C and Pressure P, where m Denotes Molality of NaCl in Watera P/MPa

m/ mol·kg−1

x/mol·kg−1

Ke

0.1 2.1 4.1 8.3 10.3 13.8 15.8 17.2 21.3 0.1 2.1 4.1 8.3 13.8 17.2 21.3 0.1 2.1 4.1 8.3 10.3 15.8

0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.0172 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.1720 0.2580 0.2580 0.2580 0.2580 0.2580 0.2580

0.004 0.225 0.402 0.710 0.831 0.973 1.040 1.059 1.147 0.004 0.215 0.394 0.694 0.961 1.041 1.112 0.002 0.184 0.380 0.657 0.766 1.019

14.13 2.15 3.13 1.38 1.56 0.87 0.88 0.89 0.50 1.41 0.33 0.36 0.20 0.12 0.13 0.13 21.09 4.82 3.03 2.23 2.41 0.93

The salting-out effect can be calculated using the following correlation: S‐O/% =

(x 0 − x ) ·100 x0

(2)

where x0 is the solubility of CO2 in distilled water, x is the solubility of CO2 at any brine and S-O/% is the percentage of salting-out effect. The values of x0 and x are reported in Tables 1 to 3. The salting-out percentage (S-O %) is shown nominally on the graph and defined as the ratio of the difference between the solubility of carbon dioxide in distilled water with the solubility at any type of brine divided by the solubility in distilled water. It can be seen from Figure 8 that at the lowest salt concentration, that is, 1000 ppm, no substantial decrease is observed in the solubility. Therefore, the salting-out effect is negligible in brines with low total dissolved solids. However, the salting out effect is intensified as the brine’s solid content is increased. The maximum salting-out effect can be seen in the data series of 15 000 ppm, in which the effect is up to 25 % at lower pressures. The lowest S-O % is observed for the brine of 1000 ppm, in which the fluctuations of the S-O are between 2 % and 7 %. The overall decline of S−O with pressure is evident in the graphs, but in order to quantify the pressure dependency, further investigations are required. Another measure of salting-out effect is the Setchenov’s constant,33 which is defined as follows:

a Combined expanded uncertainties are U(P) = 0.02 MPa, U(T) = 0.05 °C, U(m) = 0.0020 mol·kg−1, U(x) = 0.050 mol·kg−1 and U(Ke) = 0.03.

Ke =

log(x0 − x) m

(3)

where Ke is the Setchenov coefficient (dimensionless), x0 is the solubility of CO2 in distilled water, x is the solubility at any aqueous solution and m is the total dissolved solids in brine, all in molality units (both in mol·kg−1). Low affinity (for a particular solvent−salt pair) means salting-out of that particular solvent. The values of x0 can be obtained from Table 1 while values of x at 60 °C and 80 °C can be obtained from Tables 2 and 3, respectively. The value of m (salinity units) in the denominator of eq 3 is equal to 0.0172, 0.1720, and 0.2580 in molality units (equivalent to (1000, 10 000 and 15 000) ppm, respectively). Tables 1 and 2 summarize the values for Setchenov constant (Ke) in NaCl solution at various conditions of pressures and salinities at 60 °C and 80 °C, respectively. The values in the cells marked with “−” could not be obtained, because the solubility data for that specific point was not available. The uncertainties for temperature, pressure, and salinity were 0.05 °C, 0.02 MPa, and 0.002 mol·kg−1, respectively. The uncertainty for solubility was 0.001 mol· kg−1 as mentioned in Tables 2 and 3. These values of uncertainties are in line with those reported in the previous researches on the solubility of CO2 various brines using different measurement techniques such as the depressurization method.28,30,32 It can be observed from both Tables 2 and 3 that an increase in the salinity of aqueous solution decreases Ke; therefore more salting-out is expected. By comparing the values of points with similar pressure and salinity, one can conclude that increased temperature raises the salting-out effect. Regarding effects of the pressure, it can be observed that Ke is increasing at higher pressures; thus, the salting-out effect is reduced at higher pressures. It should be noted that because of the presence of a

pressures and temperatures. The increase in the brine’s salinity leads to a reduction of CO2 dissolution, regardless of its pressure and temperature. Under the experimental condition, the solubility of CO2 was reduced by 1 % to 3 % when the salinity increased from 0 to 1000 ppm, whereas a 4 % to 10 % reduction in the solubility was observed when the brine concentration was increased by 10 times (from 1000 ppm to 10 000 ppm). The decreased CO2 solubility in a brine solution can be explained by the fact that when NaCl is added to water, the ions from the salt attract the water molecules to “solvate”; thus, less water is attracted by CO2. In other words, involvement of water molecules in ions solvation reasonably decreases the weak affinity of CO2 molecules to water and drives the dissolved CO2 out of the polar water. In general, the solubility of CO2 in water is affected significantly by the presence of other solutes like NaCl. The solubility, in fact, decreases as the salinity increases due to the increased salting-out effect. It can be concluded that the solubility mechanism becomes less effective in deeper and more saline formations from the point of view of the effect of salinity on the solubility of CO2. 3.3. Salting-out Effect. The salting-out effect is the reduction in solubility of carbon dioxide as the concentration of solid particles increases in the solvent (brine). It is important to understand this effect for the estimation of the reduction in solubility as the salinity increases. The solubility of CO2 in distilled water was used as the reference and it was compared with the solubility of CO2 in brine solution with different salinities. Figure 8 illustrates the salting-out effect in NaCl solutions of (1000, 10 000, and 15 000) ppm under the pressure range of 1 MPa to 28 MPa and temperature of 80 °C. 2047

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(7) Azdarpour, A.; Assadullah, M.; Junin, R.; Manan, M.; Hamidi, H., Mohammadian, E. Direct carbonation of red gypsum to produce solid carbonates Fuel Process. Technol. 2014, volume 126, 429−434. (8) Azdarpour, A.; Assadullah, M.; Mohammadian, E.; Junin, R.; Manan, M.; Hamidi, A. Review on carbon dioxide mineral carbonation through pH-swing process. Chem. Eng. J. DOI: 10.1016/ j.cej.2015.05.064. (9) Fuller, R. C.; Prevost, J. H.; Piri, M. Three-phase equilibrium and partitioning calculations for CO2 sequestration in saline aquifers. J. Geophys. Res. 2006, volume. 111, b06207. (10) Qui, R.; LaForce, T.; Blunt, M. J. Design of carbon dioxide storage in aquifers. Int. J. Greenhouse Gas Control. 2009, Volume 3, 195−205. (11) Carroll, J. J.; Slupsky, J. D.; Mather, A. E. The solubility of carbon dioxide in water at low pressure. J. Phys. Chem. Ref. Data. 1991, volume 20. 1201−1209. (12) Zheng, D.-Q.; Guo, T.-M.; Knapp, H. Experimental and modeling studies on the solubility of CO2, CHC1F2, CHF3, C2H2F4 and C2H4F2 in water and aqueous NaCl solutions under low pressures. Fluid Phase Equilib. 1997, 129, 197−209. (13) Wiebe, R.; Gaddy, V. L. The solubility in water of carbon dioxide at 50, 75 and 100 °C at pressures to 700 atm. J. Am. Chem. Soc. 1939. Volume 61, 315−318. (14) King, M. B.; Mubarak, A.; Kim, J. D.; Bott, T. R. The mutual solubilities of water with supercritical and liquid carbon dioxides. J. Supercrit. Fluids 1992, 296−302. (15) Bamberger, A.; Sieder, G.; Maurer, G. High-pressure (vapor + liquid) equilibrium in binary mixtures of (carbon dioxide + water or acetic acid) at temperatures from 313 to 353 K. J. Supercrit. Fluids. 2000, 17, 97−110. (16) Hou, S.; Maitland, G. C.; Martin Trusler, J. P. Measurement and modelling of the phase behaviour of the (carbon dioxide + water) mixture at temperatures from 298.15 to 448.15 K. J. Supercrit. Fluids. 2013, 73, 87−96. (17) Malinin, S. D.; Savelyeva, N. I. The solubility of CO2 in NaCl and CaCl2 solutions at 25, 50, and 75 °C under elevated CO2 pressures. Geokhimiya. 1972, 6, 643−653. (18) Drummond, S. E. Boiling and mixing of hydrothermal fluids: chemical effects on mineral precipitation. Ph.D. Thesis, Pennsylvania State University,1981. (19) Rumpf, B.; Nicolaisen, H.; Ocal, C.; Maurer, G. Solubility of carbon dioxide in aqueous solutions of sodium chloride: experimental results and correlation. J. Solution Chem. 1994, 23, 431−448. (20) Koschel, D.; Coxam, J. Y.; Rodier, L.; Majer, V. Enthalpy and solubility data of CO2 in water and NaCl (aq) at conditions of interest for geological sequestration. Fluid Phase Equilib. 2006, 247, 107−120. (21) Li, Z. W.; Dong, M. Z.; Li, S. L.; Dai, S. L. Densities and solubilities for binary systems of carbon dioxide plus water and carbon dioxide plus brine at 59 °C and pressures to 29 MPa. J. Chem. Eng. Data 2004, 49, 1026−1031. (22) Liu, Y.; Hou, M.; Yang, G.; Han, B. Solubility of CO2 in aqueous solutions of NaCl, KCl, CaCl2 and their mixed salts at different temperatures and pressures. J. Supercrit. Fluids. 2011, 56, 125−129. (23) El-Maghraby, R. M.; Pentland, C. H.; Iglauer, S.; Blunt, M. J. A fast method to equilibrate carbon dioxide with brine at high pressure and elevated temperature including solubility measurements. J. Supercrit. Fluids. 2012, 62, 55−59. (24) Mao, S.; Zhang, D.; Li, Y.; Liu, N. An improved model for calculating CO2 solubility in aqueous NaCl solutions and the application to CO2−H2O−NaCl fluid inclusions. Chem. Geol. 2013, 347, 43−58. (25) Heavyseage, R. G. Formation evaluation of fresh water shaly sands of the Malay basin, offshore Malaysia. Annu. Logging Symp. 2002, SPWLA 43ed. (26) Peper, S.; Dohrn, R. Sampling from fluid mixtures under high pressure: Review, case study and evaluation. J. Supercrit. Fluids. 2012, 66, 2−15. (27) Portier, S.; Rochelle, C. Modelling CO2 solubility in pure water and NaCl-type waters from 0 to 300 °C and from 1 to 300 bar:

logarithm in eq 3, the values in Tables 2 and 3 are smoother than those obtained from eq 2.

4. CONCLUSIONS Because of the scarcity of data in the low salinity range, CO2 solubility was studied at low salinity synthetic brines in the current research, through the potentiometric titration method in distilled water and NaCl solution. The current method was proven to be reproducible and accurate in comparison to the previous conventional methods. Measurements were taken at various temperatures, pressures, and salinities, and the results exhibited a good consistency with those of previous researches under the same thermodynamic conditions. Regarding the pressure, it can be concluded that increase in pressure enhances the solubility of CO2 in the distilled water as well as in the brine. However, the pressure dependency of the solubility reduces as the pressure increases. The effect of temperature on the solubility is adverse; an increase in the temperature substantially reduces the solubility, regardless of the brine’s salinity. Lastly, an increase in the salinity also exerts an adverse effect on the solubility. At low salinities, however, the effect was not that noticeable. Yet, reduction in the solubility increased considerably by almost 13 % as the salinity increased to 1.5 % from an initial value of 0 (distilled water).

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Funding

This research was financially supported by the UiTM Research Excellence Fund (Grant No. 600-RMI/DANA 5/3/RIF 547/ 2012). The authors also wish to thank the Faculty of Chemical Engineering at Universiti Teknologi MARA in Shah Alam, Malaysia for the provision of laboratory facilities for this work. Also, sincere thanks are expressed for the funding support received from the HIR-MOHE, University of Malaya, under the Grant No. UM.C/HIR/MOHE/ENG/34. Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Zhao, X.; Liao, X. Evaluation method of CO2 sequestration and enhanced oil recovery in an oil reservoir, as applied to the Changqing Oilfields, China. Energy Fuels. 2012, 26 (8), 5350−5354. (2) Godec, M. L.; Kuuskraa, V. A.; Dipietro, P. Opportunities for using anthropogenic CO2 for enhanced oil recovery and CO2 storage. Energy Fuels. 2013, 27 (8), 4183−4189. (3) Zhang, D.; Gu, L.; Li, S.; Lian, P.; Tao, J. Interactions of supercritical CO2 with coal. Energy Fuels. 2013, 27 (1), 387−393. (4) Bachu, S.; Adams, J. Sequestration of CO2 in geological media in response to climate change: Capacity of deep saline aquifers to sequester CO2 in solution. Energy Convers. Manage. 2003, 44, 3151− 3175. (5) Tsouris, C.; Szymcek, P.; Taboada-Serrano, P.; McCallum, S. D.; Brewer, P.; Peltzer, E.; Walz, P.; Adams, E.; Chow, E.; Johnson, W. K.; Summer, J. Scaled-up ocean injection of CO2−hydrate composite particles. Energy Fuels. 2007, 21 (6), 3300−3309. (6) Azdarpour, A.; Assadullah, M.; Mohammadian, E.; Junin, R.; Manan, M.; Hamidi, H.; Mohd Daus, A. R. Mineral carbonation of red gypsum via pH-swing process: Effect of CO2 pressure on the efficiency and products characteristics. Int. J. Chem. Eng. 2014, 28 (9), 5953− 5958. 2048

DOI: 10.1021/je501172d J. Chem. Eng. Data 2015, 60, 2042−2049

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Article

Application to the Utsira Formation at Sleipner. Chem. Geol. 2005, 217, 3−4 187−199.. (28) Yan, W.; Huang, S.; Stenby, E. H. Measurement and modeling of CO2 solubility in NaCl brine and CO2−saturated NaCl brine density. Int. J. Greenhouse Gas Control. 2011, 5, 1460−1477. (29) Duan, Z. H.; Sun, R. An improved model calculating CO2 solubility in pure water and aqueous NaCl solutions from 273 to 533 K and from 0 to 2000 bar. Chem. Geol. 2003, 193, 257−271. (30) Wang, L.; Shen, Z.; Hub, L.; Yu, Q. Modelling and measurement of CO2 solubility in salty aqueous solutions and application in the Erdos Basin. Fluid Phase Equilib. 2014, 15, 45−55. (31) Zumdahl, S. S. Chemical Principles, 4th ed.; Houghton Mifflin Company: Boston, MA, 2002. (32) Nighswander, J. A.; Kalogerakis, N.; Mehrotra, A. K. Solubilities of carbon dioxide in water and 1 wt % NaCl solutions at pressures up tp 10 MPa and temperatures from 80 to 200 deg C. J. Chem. Eng. Data 1989, 34, 355−60. (33) Lioyde, L. L. A molecular theory of Setchenov’s salting-out principle and applications in mixed-solvent electrolyte solutions. Phase Equilibr. 1997, 131, 67−82.

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DOI: 10.1021/je501172d J. Chem. Eng. Data 2015, 60, 2042−2049