Phase Equilibria of CO2 Hydrate in CaCl2–MgCl2 Aqueous Solutions

Article pubs.acs.org/jced

Phase Equilibria of CO2 Hydrate in CaCl2−MgCl2 Aqueous Solutions Zhen Long, Li Zha, Deqing Liang,* and Dongliang Li Key Laboratory of Renewable Energy, Guangzhou Institute of Energy Conversion, and Guangzhou Center for Gas Hydrate Research, Chinese Academy of Sciences, Guangzhou 510640, People’s Republic of China ABSTRACT: Phase equilibrium data for CO2 hydrate in binary CaCl2−MgCl2 aqueous solutions are determined at four different concentrations (2 wt % CaCl2-8 wt % MgCl2, 8 wt % CaCl2-2 wt % MgCl2, 5 wt % CaCl2-15 wt % MgCl2, and 15 wt % CaCl2-5 wt % MgCl2) in the temperature range of (257.14 to 276.89) K and at pressures ranging from (0.94 to 3.38) MPa. An isochoric pressure-search method is applied in all of the measurements. A comparison is made between our experimental results with some selected experimental data from the literature for CO2 hydrate in pure water. All of the solutions are observed to have an inhibiting effect on the CO2 hydrate formation.

1. INTRODUCTION Gas hydrates are ice-like nonstoichiometric compounds that are made up of host molecules (water) and guest molecules (e.g., CH4, C2H6, C3H8, etc.) under appreciate conditions of low temperatures and high pressures. Water molecules are linked together through hydrogen bonds to form different cavities, where guest molecules or ions are trapped and stabilized by van der Waal interaction forces.1 Depending on difference in cavity size and shape, the hydrate structures are divided into three common types: sI (structure I), sII (structure II), and sH (structure H).2 The formation of gas hydrates is unpleasant and dangerous for the petroleum industry, because it can cause the blockage of pipeline or equipment and lead to safety concerns during hydrocarbon production and transportation. At the same pressure, the formation temperature of a CO2 hydrate is relatively higher than either CH4, H2, or N2. For the systems rich of CO2, they are therefore at a greater risk of hydrate formation. The precise amount of such additives as electrolytes and alcohols injected into the pipeline, where the formation of hydrates can be effectively inhibited at the lowest cost, depends on the stability shift in the hydrate formation.3 On the other hand, CO2 hydrate formation is beneficial to environmental protection. The global greenhouse effect is mainly attributed to CO2 emission, of which one-third is generated from fossil-fueled power plants. There are many methods for separating and fixing CO2, such as absorption,4 adsorption,5 membrane separation,6,7 and underground injection and direct ocean dump.8,9 Disposal of CO2 in the form of hydrates in the ocean also has been proposed10−13 to reduce the quantities of CO2 released into the atmosphere and mitigate the global warming effect. Owing to larger density of CO2 hydrate than that of seawater, the hydrate particles can sink toward the deep sea bottom and would be stable for a long time. Various salts, including NaCl, KCl, CaCl2, and MgCl2 © XXXX American Chemical Society

relatively widely dissolved in seawater, however, change thermodynamics stability of CO2 hydrate, resulting in hydrate decomposition and hindering CO2 sequestration.13 Hence it is very necessary to investigate the effects of the systems containing electrolytes on the CO2 hydrate stability conditions. Phase equilibrium data for CO2 hydrate in pure water1,14,15 are available in the literature together with the systems containing single composition electrolyte aqueous solutions. The formation conditions of CO2 hydrate in NaCl, KCl, CaCl2, NaCl-KCl, and NaCl-CaCl2 were first published by Dholabhai et al.15 Clark et al.16 reported the experimental values for CO2 hydrate formation conditions in aqueous solutions of KNO3, MgSO4, and CuSO4 and observed a promoting effect of CuSO4 at high concentrations. Three phase (hydrate−aqueous solution−vapor) equilibrium data for CO2 hydrate in the MgCl2-containing systems were obtained together with quadruple points at the MgCl2 concentration ranging from 3 to 10 wt %, and it was found that MgCl2 has a stronger inhibition effect than CaCl2.17 However, to our knowledge, data for CO2 hydrate in the bivalent mixture, herein CaCl2−MgCl2, have not been provided elsewhere, except those data in the mixture of mono- and divalent salts (NaCl-MgCl2) sourced from our previous study.18 Those data are beneficial for comparing the influence of different ions on the equilibrium conditions and hence helpful for further research on salts’ inhibition mechanism. On the other hand, the simple, explicit, predictive correlations available for quick estimation of hydrate inhibition effect of thermodynamic inhibitors are based on limited data and only applied under limited conditions. As a follow-up to the aforementioned study,18 the present work aims Received: May 6, 2014 Accepted: June 20, 2014

A

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determined. In this way, the point at which the slope of P−T curve changes sharply is considered to be the dissociation point, at which all the hydrate crystals have dissociated. Briefly, for the system containing the mixture of CaCl2 and MgCl2, the hydrate dissociation points are determined by an isochoric step-heating pressure-search method.18−20 When changing the initial system pressure, the next hydrate phase equilibrium point can be obtained by repeating the above procedure. The reliability and reproducibility of the data reported via this procedure is validated to be good.18−20

to measure those data for CO2 hydrate in the presence of CaCl2−MgCl2.

2. EXPERIMENTAL SECTION 2.1. Chemicals. The chloride salts, calcium chloride CaCl2· 2H2O (0.99 mass fraction pure) and magnesium chloride MgCl2·6H2O (0.99 mass fraction pure), used for the present work are supplied by Sigma-Aldrich Co., Ltd., and Eastman Kodak Laboratory Chemicals of Belgium, respectively. Four different concentrations of mixed solution of CaCl2 and MgCl2 (2 wt % CaCl2-8 wt % MgCl2, 8 wt % CaCl2-2 wt % MgCl2, 5 wt % CaCl2-15 wt % MgCl2, and 15 wt % CaCl2-5 wt % MgCl2) are prepared. The compositions and notation of the electrolyte aqueous solutions are listed in Table 1. Highly pure CO2 gas

3. RESULTS AND DISCUSSION CO2 hydrate equilibrium conditions in the presence of CaCl2− MgCl2 electrolyte solutions are shown in Table 2 and Figure 1.

Table 1. Compositions of the Aqueous Solutions

Table 2. Three Phase Equilibrium (Hydrate−Aqueous Solution−Vapor) Data for CO2 Hydrate in Aqueous Electrolyte Solutionsa

composition, wt % on wet basis solution

CaCl2

MgCl2

Ca2Mg8 Ca8Mg2 Ca5Mg15 Ca15Mg5

2.0 8.0 5.0 15.0

8.0 2.0 15.0 5.0

(0.9999 volume fraction pure) is obtained from Foshan Kede Gas Co., Ltd. Deionized water made in the laboratory with a resistivity of 18 MΩ is used to prepare the aqueous solutions of the chloride salts. Considering the strong moisture absorption, the solution preparation process is undertaken in a dry environment with a dehumidifier. The required amounts of each component are weighed by an electronic analytical balance with a reading uncertainty of ± 0.1 mg. Once being completely dissolved and ready, the solution is inhaled into the cell by the vacuum pump. All chemicals are used without further purification. 2.2. Experimental Apparatus and Procedure. A detailed experimental setup is described in our previous work.18 The main part of the apparatus is a cylindrically shaped equilibrium cell, which is made of stainless steel with an internal volume of about 25 cm3. The materials in the cell are mixed by a magnetic stirrer which is coupled with a magnet located outside the cell. The cell is immersed into a high and low temperature test chamber to keep a constant temperature of the system. The temperature in the cell is measured by a platinum resistance thermometer (PT100) with ± 0.1 K accuracy. A pressure sensor (CYB-20S) ranged to (0−20) MPa with an uncertainty of ± 0.02 MPa is used to detect the pressure in the cell. The data of temperature and pressure are recorded and displayed on a computer via an Agilent online data acquisition system. The experimental procedure in the present work is the same as our previous work.18 First, the cell is washed with deionized water and dried with air oven. Next is to flush the cell and all the connected pipes with CO2 gas. Then the cell is evacuated before introducing aqueous solutions and gas. When the desired initial pressure and temperature are stable, the temperature is slowly decreased to form hydrate. Once detecting an abrupt pressure drop and an evident temperature increase, it can judge the hydrate is formed. The system temperature is then increased with steps of 0.1 K. At every temperature step, the temperature keeps constant for about (4 to 6) h until a steady equilibrium state is achieved. Consequently, the P−T curve is obtained for each experimental run, from which the hydrate dissociation point can be

solution

T/K

P/MPa

solution

T/K

P/MPa

Ca2Mg8

270.33 272.24 273.79 275.02 275.99 269.74 271.98 274.55 275.77 276.89

1.64 2.06 2.51 2.95 3.38 1.31 1.71 2.22 2.83 3.33

Ca5Mg15

257.14 259.53 262.34 264.37 265.5 257.49 261.23 263.91 265.74 266.33

1.03 1.46 1.98 2.45 2.74 0.94 1.45 1.97 2.42 2.76

Ca8Mg2

a

Ca15Mg5

Standard uncertainties u are u(T) = 0.1 K, u(P) = 0.02 MPa.

Figure 1. Three phase equilibrium (aqueous solution + hydrate + vapor) data for CO2 hydrate in CaCl2−MgCl2 aqueous solutions. ■, Ca2Mg8; ▲, Ca8Mg2; ★, Ca5Mg15; ▼, Ca15Mg5; ○, pure water, ref 14; △, pure water, ref 18; +, pure water, ref 1; ▽, pure water, ref 15; •••, saturated vapor pressure of CO2.

The hydrate equilibrium data in pure water1,14,15,18 are also presented for illustrating the inhibition effect of CaCl2−MgCl2. In Figure 1, the lower right region of the dotted line represents the conditions at which CO2 is present in vapor phase. As shown, it is clear that the CaCl2−MgCl2 at each mass concentration composition performs an evident inhibiting effect. Moreover, as the mass concentrations of CaCl2 and MgCl2 contained in the mixed aqueous electrolytes both increase at the same time, the inhibition effect increases. The same finding is also obtained in the single CaCl2 and MgCl2 B

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Figure 2. Relationship between (a) Cl− and (b) cation molarity with the shift in equilibrium conditions of CO2 hydrate in various electrolyte solutions. ■, NaCl, ref 15; ●, CaCl2, ref 15; ▲, MgCl2, ref 16; ☆, NaCl−CaCl2, ref 15; ◊, NaCl−MgCl2, ref 18; ▽, CaCl2−MgCl2, this work; solid line, the regression line of data in NaCl; dot line, the ☆ regression line of data in CaCl2; dash line, the regression line of data in MgCl2; dash-dot line, the regression line of data in NaCl−CaCl2; dash-dot-dot line, the regression line of data in NaCl−MgCl2; short dot line, the regression line of data in CaCl2−MgCl2.

aqueous solutions which inhibit CH4,21,22 C2H620,22 and CO215,16 hydrates formation. The hydrate suppression temperatures for the solutions Ca8Mg2 and Ca2Mg8 are found very close, equal to 4.1K and 5.3K on the average, respectively. A similar behavior is also observed for the two other solutions Ca15Mg5 and Ca5Mg15, where hydrate formation temperatures relative to those required in pure water depress 13.2K and 14.8K on the average, respectively. This leads one to infer that the inhibition strengths of CaCl2 and MgCl2, on a weight basis, are very close. The shift in the CO2 hydrate phase equilibrium conditions relative to that in pure water is plotted against Cl− and cation molar concentration in single and binary electrolyte solutions of NaCl, CaCl2, and MgCl2, as shown in Figure 2. Here, T0 stands for hydrate formation temperature of CO2 in pure water and is calculated at a given pressure by using CSMHYD thermodynamic model.1 All the concentrations of electrolyte solutions have been converted to molarity from original mass fraction. The densities taken from the model proposed by Laliberté et al.23 are used for calculation of the molarities. As seen from Figure 2a, the shift in equilibrium temperature (T0−T) can be well linearly regressed in terms of Cl− molar concentration in solution NaCl, CaCl2, MgCl2, NaCl−CaCl2, NaCl−MgCl2, and CaCl2−MgCl2. If the cation presents a much stronger influence in hindering the gas hydrate formation, the regressed lines should be much different. On the contrary, the slopes of the regression lines are observed to be very near, though some small differences exist among various chloride salts containing different cation. The same findings have been revealed for C2H6 hydrate20 and C3H8 hydrate24 in the chloride salts. An explanation for this phenomenon has been given by Lu et al.24 in that ion hydration leads to the decreases in the number of free water available for hydrate formation and thereby lowers water activity, and hydration cation’s influence on the water molecules is limited to the first and second hydration layer while anion has a significantly superior hydration behavior which affects the whole ambient water molecule network. According to the ion charge balance, the temperature shift (T0−T) also shows a good linear relationship with cation molarity in single solutions NaCl, CaCl2, and MgCl2, as depicted in Figure 2b. The divalent cations (Ca2+, Mg2+, and

Ca2+ + Mg2+) display stronger inhibition than mixture of monoand divalent cations (Ca2+ + Na+ and Mg2+ + Na+). Keeping Cl− molar concentration constant, for example, equal to 1000 mmol·dm−3, when 500 mmol·dm−3 Mg2+ is substituted by 500 mmol·dm−3 Ca2+, the values of (T0−T) are about (2.24 and 1.97) K, respectively. It seems smaller Mg2+ ion has very slightly superior inhibition capability compared to Ca2+, following the order of inhibition strength among chloride salts suggested by Berecz et al.25 and Englezos et al.26 As a summary, the effective inhibition of ions on the hydrate formation is not only related with the ion concentration, but also the ion type. Besides limited data and limited ions, there is need for plentiful of data to elucidate the mechanism how the salts affect the hydrate formation.

4. CONCLUSIONS In this work, thermodynamic experiments are conducted to study the effects of the mixed CaCl2−MgCl2 aqueous solutions on the CO2 hydrate formation conditions at four mass concentrations in the temperature and pressure ranges from (257.14 to 276.89) K and (0.94 to 3.38) MPa, respectively. Measurements are made by utilizing the isochoric pressuresearch method. CaCl2−MgCl2 solutions are observed to inhibit CO2 hydrate formation. Moreover, the higher Cl− concentration, the inhibition effect becomes stronger. A comparison of the inhibition effect among several chloride salts on the CO2 hydrate formation reveals that the stability shift is mostly determined by the anion Cl−, while cations such as Na+, Ca2+, and Mg2+ have a relatively weaker influence.

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AUTHOR INFORMATION

Corresponding Author

*Phone: +86 20 8705 7669. Fax: +86 20 8705 7669. E-mail: [email protected]. Funding

This work is financially supported by the National Natural Science Foundation of China (No. 51176192), the National HiTech Research and Development Program of China (863, No. 2012AA061403-03), and the Chinese Academy of Sciences Key Development Program (No. KGZD-EW-301). C

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Notes

(22) Mohammadi, A. H.; Afzal, W.; Richon, D. Gas Hydrates of Methane, Ethane, Propane, and Carbon Dioxide in the Presence of Single NaCl, KCl, and CaCl2 Aqueous Solutions: Experimental Measurements and Predictions of Dissociation Conditions. J. Chem. Thermodyn. 2008, 40, 1693−1697. (23) Laliberté, M.; Edward Cooper, W. Model for Calculating the Density of Aqueous Electrolyte Solutions. J. Chem. Eng. Data 2004, 49, 1141−1151. (24) Lu, H.; Matsumoto, R.; Tsuji, Y.; Oda, H. Anion Plays a More Important Role than Cation in Affecting Gas Hydrate Stability in Electrolyte Solution?A Recognition from Experimental Results. Fluid Phase Equilib. 2001, 178, 225−232. (25) Berecz, E.; Balla-Achs, M. Gas Hydrates; Elsevier: Amsterdam, 1983. (26) Englezos, P.; Hall, S. Phase Equilibrium Data on Carbon Dioxide Hydrate in the Presence of Electrolytes, Water Soluble Polymers and Montmorillonite. Can. J. Chem. Eng. 1994, 72, 887−893.

The authors declare no competing financial interest.

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REFERENCES

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dx.doi.org/10.1021/je500400s | J. Chem. Eng. Data XXXX, XXX, XXX−XXX