Phase Behavior of Carbon Dioxide Hydrates: A Comparison of

Article pubs.acs.org/jced

Phase Behavior of Carbon Dioxide Hydrates: A Comparison of Inhibition Between Sodium Chloride and Ethanol Amanda Z. Guembaroski,† Moisés A. Marcelino Neto,*,† Dalton Bertoldi,† Rigoberto E. M. Morales,† and Amadeu K. Sum‡ †

Multiphase Flow Research Center (NUEM), Post-graduate Program in Mechanical and Materials Engineering (PPGEM), Federal University of Technology−Paraná (UTFPR), Curitiba, PR, Brazil ‡ Hydrates Energy Innovation Laboratory, Chemical & Biological Engineering Department, Colorado School of Mines, Golden, Colorado 80401, United States ABSTRACT: In view of the possibilities for hydrate formation caused by carbon dioxiderich fluids in the production lines of the Brazilian Pre-Salt fields, this study focus on experimental measurements to obtain fundamental insight into the phase behavior of carbon dioxide hydrate forming systems. This work considers the influence of sodium chloride and ethanol, hydrate inhibitors, on this phase behavior and its implications to practical applications. In addition, the inhibiting effect of ethanol on hydrates was compared with that of sodium chloride at the same mass fractions. The carbon dioxide hydrate phase behavior was measured using a high-pressure equilibrium cell in the temperature range of 272−279 K and pressures up to 3.9 MPa. Experimental measurements using the isothermal method were performed by monitoring the pressure response of the system with volume changes. Enthalpies of dissociation for carbon dioxide hydrates were estimated from the measured three-phase (Lw−H−V) data by applying the Clausius−Clapeyron equation. Results showed that, at the same mass fraction of the inhibitor and high-pressure conditions, the sodium chloride exhibited a superior inhibiting effect compared to that of the ethanol.

1. INTRODUCTION Clathrate hydrates are crystalline solids composed of hydrogenbonded cages water molecules. Gas clathrate hydrates may form at high pressures and/or low temperature once water and gas are present. The existence of a trapped molecule inside the water cage thermodynamically stabilizes the structure by van der Waals forces. Typical natural gas molecules, with proper sizes, including methane, ethane, propane, nitrogen, carbon dioxide, and hydrogen sulfide, are appropriate for clathrate hydrate formation. The hydrates originated from those gases possibly will form three crystal structures, cubic structure I (sI), cubic structure II (sII), or hexagonal structure H (sH).1 Carbon dioxide hydrate, CO2·nH2O (n ≥ 5.75), crystallizes as sI, where 46 H2O molecules and up to 8 CO2 molecules occupy both pentagonal dodecahedral and tetrakaidecahedral cavities. CO2 hydrate is stable over a range of elevated pressure and low-temperature conditions.2 Compared to methane and other volatile hydrocarbons, CO2 presents a higher solubility in water. This high solubility is due to the polar attractive forces that make the mixture to be more predisposed to hydrate formation.3 Gas hydrates are easily formed during the transportation of oil and gas when water is present, potentially resulting in pipeline blockages and other operational problems in the petroleum industry. During the transportation and processing, especially when the produced gas is water saturated and under cold conditions, gas hydrate formation may plug pipelines, valves, and other pieces of equipment. © 2017 American Chemical Society

In the petroleum industry, operating the production lines outside the hydrate zone is a common practice to prevent hydrates from forming. There are several methods to either prevent, halt, or even to revert hydrate formation in circumstances they are susceptible to form. In some situations, a combination of controlling the produced fluid temperature and pressure is sufficient to prevent hydrate formation. Water removal operations can be efficient as well. However, there are situations where the application of one of those techniques may not be cost-effective. Consequently, the use of thermodynamic inhibitors, such as alcohols and salts, is often an option due to their inherent properties to suppress hydrate formation. The dissolved inhibitor molecule or ion decreases the water activity through hydrogen bonding (for alcohols) or via Coulombic forces (for salt ions).2 Gas hydrate phase equilibrium data inhibited with sodium chloride and ethanol is therefore essential to assess the boundaries of the gas hydrate formation zone. Equilibrium conditions of CO2 hydrate in pure water have been reported in a relatively abundant number of publications.4−28 Experimental equilibrium data for CO2 hydrate in aqueous solutions of electrolytes have been partially reported, and the development of thermodynamic methods to calculate the hydrate equilibria conditions requires accurate experimental Received: May 22, 2017 Accepted: August 8, 2017 Published: August 18, 2017 3445

DOI: 10.1021/acs.jced.7b00463 J. Chem. Eng. Data 2017, 62, 3445−3451

Journal of Chemical & Engineering Data

Article

data involving electrolytes. Vlahakis et al.29 first published hydrate phase equilibrium data of CO2 in aqueous solutions of NaCl. Dholabhai et al.15 also reported data for CO2 hydrates in aqueous solutions containing NaCl. Sabil et al.1 investigated the competing effects of sodium chloride and tetrahydrofuran on carbon dioxide hydrate formation through phase equilibrium measurements. Ruffine and Trusler30 reported new measurements for hydrate equilibrium curves in aqueous systems containing either pure carbon dioxide or carbon dioxide with 4.3 wt % sodium chloride. More recently, Godishala et al.40 carried out experimental studies on carbon dioxide hydrates inhibited with tetra-n-butylammonium bromide (TBAB) and NaCl. Three mass fractions of TBAB (0.05, 0.10, and 0.20) and three mass fractions of NaCl (0.035 and 0.10) in an aqueous system were measured. While there are a number of studies published for NaCl as an inhibitor for CO2 hydrates, the need for more experimental data is still current. Notwithstanding the number of publications on CO2 hydrate phase equilibrium inhibited with alcohols, there is a lack of data for CO2 hydrate phase equilibrium inhibited with ethanol. Mohammadi et al.31 published experimental carbon dioxide hydrate phase equilibrium data inhibited with ethanol (5 and 10 wt %) for temperatures ranging from 271.3 to 280.1 K and pressures up to 4.54 MPa. Maekawa32 measured equilibrium conditions for carbon dioxide hydrates inhibited with ethanol in the 264.1− 283.1 K temperature range and pressures up to 3.45 MPa. Ferrari et al.33 measured carbon dioxide hydrate phase equilibrium data inhibited with ethanol (5 and 10 wt %) in the 275.65−281.65 K temperature range and pressures up to 3.5 MPa. Sami et al.41 published experimental gas hydrate dissociation data for methane and carbon dioxide mixture with (0.1 mass fraction methanol + 0.03, 0.1 mass fraction MgCl2) and (0.1, 0.2 mass fraction ethylene glycol + 0.1 mass fraction MgCl2) aqueous solutions at temperatures ranging from 263.74 to 280.54 K and pressures ranging from 0.98 to 8.02 MPa. This article follows on our previous study33 in which the dissociation conditions of CO2 hydrates in the presence of aqueous ethanol solutions was reported. This article reports new data on CO2 hydrates inhibited with sodium chloride (0.02, 0.05, 0.10, and 0.15 mass fraction) in liquid water− hydrate−vapor (Lw−H−V) phase equilibrium. In addition, new equilibrium conditions of Lw−H−V for carbon dioxide hydrates with ethanol (0.02 and 0.15 mass fraction) were measured. The experimental data are compared with literature data on dissociation conditions of carbon dioxide hydrates in the presence of pure water to show the inhibiting effects of the aforementioned inhibitors in aqueous solutions.

certainties on a mole fraction basis are estimated as being less than 0.01. Degassed and deionized water (MILI-Q) was used in all experiments. Four mass fractions of sodium chloride were selected: 0.02, 0.05, 0.10, and 0.15. Higher concentrations were not tested since the ranges of temperature and pressure were limited by two factors: the solubility threshold of sodium chloride in water (∼26%) could not be reached, and at higher pressures, the condensation of a CO2-rich liquid phase could be observed. Figure 1 illustrates a representation of the experimental facility used, which was described in detail in our previous study.33 According to Dohrn et al.,34 this is a static-synthetic apparatus. An equilibrium cell (7), shown in Figure 1, was used to study the conditions in which hydrates dissociate. The cell is a horizontal, cylindrical vessel made of stainless steel with a total volume of 25 cm3. The limit working pressure of this cell is 25 MPa, and the temperature range is 233−473 K. The cell has two sapphire windows (b) which allow phase transition observations. A magnetic stirrer (with a variable speed) was used to reduce the time to reach equilibrium. A platinum resistance thermometer with an accuracy of ±0.17 K (95% confidence level) was used to read the temperature of the equilibrium cell. The absolute pressure measurements were made with an uncertainty of ±0.30% of the absolute reading (95% confidence level). The temperature of the equilibrium cell (7) was controlled by a thermostatic bath (3), whereas the pressure is controlled with a syringe pump (2). A digital camera (5) with a resolution of 1080p recorded the phase equilibrium images. 2.2. Procedure. The three-phase (LW−H−V) data were measured following an isothermal procedure, as described below. At first, before assembling the apparatus, the equilibrium cell (7) is carefully cleansed. The cell is fed with a specified amount of aqueous solution (measured with a precision balance) containing inhibitor at ambient pressure and pressurized carbon dioxide is injected using the syringe pump (2). The temperature and pressure of the hydrate region are reached through the thermostatic bath (3) and syringe pump (2), respectively. After the hydrates are formed in the cell, the system reaches equilibrium at constant temperature and pressure approximately 3 h later. Next, the cell is gradually depressurized in gradual steps (increasing the syringe pump volume) to dissociate the hydrates. These steps would initially be 0.1 MPa every 20 min, i.e., at a depressurization rate of 0.005 MPa/min. Between two steps, the volume of the pump was kept unchanged for 10 min, to verify any pressure increase that might indicate hydrate dissociation. If no pressure increase is detected, the pressure is further decreased. In the case of increase of pressure, indicating that hydrates are dissociating, the volume of the pump is kept constant for pressure stabilization. For this situation, a step decrease of 0.05 MPa was allowed to perturb the equilibrium once again for the future pressure response to return the system to the level achieved before. This process is repeated, with gradually decreasing pressure drops, until the hydrates are totally dissociated. The dissociation point (hydrate equilibrium) is the last pressure where hydrates are even noticeable. This procedure is repeated for each temperature until the hydrate equilibrium pressure is found. Figure 2 shows the visual states of the contents of the cell along the experimental procedure.

2. EXPERIMENTAL SECTION 2.1. Materials and Apparatus. Table 1 contains the purities and suppliers of the chemicals used in the experiments described in this work. Aqueous solutions were prepared according to a gravimetric method, using an accurate analytical balance (mass uncertainty ±0.0001 g). Consequently, unTable 1. List of Chemicals Used for the Experimental Measurements chemical

supplier

mole fraction purity

carbon dioxide sodium chloride ethanol

White Martins Biotec Vetec Chemicals

0.99995 0.99000 0.99800 3446



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Figure 1. Schematic diagram of the experimental apparatus for hydrate equilibrium measurements.

3. RESULTS AND DISCUSSION The accuracy of the experimental apparatus and procedure was verified in our previous study.33 The LW−H−V equilibrium lines of carbon dioxide−pure water were measured and compared with literature data.4,10,20,21 Hydrate equilibrium conditions (LW−H−V) were obtained for CO2 in aqueous solutions of 0.02, 0.05, 0.10, and 0.15 NaCl mass fraction and 0.02, 0.15 ethanol mass fraction. The data obtained in this study are given in Table 2 and are also plotted in Figure 3. The equilibrium curves for LW−H−V for the pure CO2 and CO2 inhibited with 0.05 and 0.10 ethanol mass fraction (from our previous study33) were also included in Figure 3 for comparison. Figure 3 shows that, for a given pressure, the hydrate formation temperature decreases as the inhibitor concentration increases; i.e., adding sodium chloride and ethanol inhibits the carbon dioxide hydrate formation process and that inhibiting effect is dependent on concentration. Also shown in the figure is the CO2 vapor pressure line indicating the superior limit of the LW−H−V equilibrium. The latent heat required to dissociate the hydrate into its liquid and gas phase constituents is another important property, that is, the enthalpy of dissociation (ΔH) per mole of gas. The

Figure 2. Pressure during time trace for pure CO2 hydrate dissociation (measured hydrate equilibrium dissociation point is 1.65 MPa at 275.65 K). Images of the content of the cell along the experiment: (a) hydrate formation, (b) hydrate dissociation, and (c) phase equilibrium point.

Table 2. Measured Dissociation Points of CO2 Hydrates Inhibited with Sodium Chloride (NaCl) and Ethanol (EtOH) inhibitor

2 wt % NaCl

5 wt % NaCl

10 wt % NaCl

15 wt % NaCl

2 wt % EtOH

15 wt % EtOH

T (K)

P (MPa)

P (MPa)

P (MPa)

P (MPa)

P (MPa)

P (MPa)

2.751 3.011 3.194 3.579

1.351

2.181 2.469

272.15 272.65 273.15 273.65 274.65 275.15 275.65 276.15 276.65 277.15 277.65 278.65 279.15

1.436

1.463 1.567 1.825

2.016 2.187 2.501

1.715

1.637 2.912

1.931

2.247

2.789 1.839

3.395 2.159

3.182 2.066

2.712

3.919

2.647

3.689 2.514

3.268 3447



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It is known that the hydrate stability region shifts to lower temperature and higher pressure in the presence of inhibitors. To evaluate the inhibition strength of sodium chloride and ethanol on the formation of hydrates, the experimental data of hydrate equilibrium were analyzed in the presence of different inhibitors concentrations at constant pressures. Figure 5a shows that the CO2 hydrate dissociation temperature decreases with increasing NaCl concentration. Sodium chloride ionizes in solution and has strong electrostatic interactions with the water molecules. The water solvating the dissolved ions are unable to form hydrates, and as such, a lower temperature is needed to order the water molecules with the CO2 molecules for hydrates to form.2 As a secondary effect, the presence of NaCl in aqueous solutions results in the salting-out effect. This effect is based on the electrolyte−nonelectrolyte reciprocal influence, where the nonelectrolyte may possibly be not as much as soluble at high salt concentrations. Consequently, the solubility of CO2 decreases with increasing NaCl concentration, as experimentally observed by Sun et al. 36 As the salt concentration is increased, more water molecules are tied up by the salt ions, decreasing the number of water molecules available to form the hydrogen bonds that cause clustering around the CO2 molecules. Salting-out effect and ion clustering act together to require considerably more subcooling to bring out the structural changes and to trigger hydrate formation. The action of ethanol as a hydrate inhibitor is different than that of the sodium chloride. The hydrogen bonding of the hydroxyl group in ethanol with water molecules is the only molecular interaction that ties up water molecules making them unavailable to form hydrates. Additionally, the hydrocarbon segment of ethanol causes a clustering effect on water molecules similar to that caused by CO2.2 As for the hydrate cluster formation, both effects compete with the solubilized CO2 molecules. However, the hydrogen bonding of the hydroxyl group is more pronounced. In both sodium chloride and ethanol, the water activity is altered in the arrangement for the formation of the hydrate, and stronger driving forces are therefore required to induce the formation.2

Figure 3. CO2 hydrate equilibrium (LW−H−V) in pure water and aqueous solutions with sodium chloride and ethanol. Dashed lines are the best fit and are used to guide the eye.

enthalpy of dissociation of the hydrates can be obtained by either direct or indirect methods. In the direct method, the data are typically obtained by calorimetry. The indirect method is a simpler form to calculate the enthalpy of dissociation and can be calculated from the Clausius−Clapeyron equation. Figure 4 shows a semilogarithm plot of the dissociation pressure versus the inverse of the absolute temperature. As described by the Clausius−Clapeyron equation, a linear trend for the data plotted in this way not only can verify the thermodynamic consistency of the data but also gives the heat of dissociation, which is directly calculated from the slope of the lines.2,35 The slope is equal to −ΔH/zR, where z is the compressibility factor of the gas and R is the universal gas constant.1 The average heat of dissociation calculated by the Clausius−Clapeyron equation are 68.87, 70.04, 72.12, and 74.08 kJ/mol for 0.02, 0.05, 0.10, and 0.15 sodium chloride mass fraction, respectively, and 65.60 and 69.81 kJ/mol for 0.02 and 0.15 ethanol mass fraction.

Figure 4. Semilogarithmic plot of the dissociation pressure versus inverse absolute temperature. (a) Sodium chloride aqueous solution. (b) Ethanol aqueous solution. 3448



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Figure 5. CO2 hydrate equilibrium conditions at constant pressure for aqueous (a) sodium chloride and (b) ethanol solutions.

Table 3. Hydrate Depression Temperature (ΔTH) and Freezing Point Depression (ΔT) ΔTH (K) system

inhibitor molar fraction

ΔT (K) P = 0.1 MPa

P = 2 MPa

P = 3 MPa

P = 4 MPa

P = 5 MPa

P = 6 MPa

CO2 + H2O + NaCl

0.034 0.086 0.171 0.257 0.043 0.108 0.217 0.326

1.190 3.050 6.560 10.915 0.810 2.090 4.470 7.370

0.770 2.150 4.310 5.891 0.470 0.735 3.267 5.229

0.970 2.108 4.661 7.495 0.650 1.405 2.872 4.253

1.030 2.007 5.044 9.058 1.205 2.301 3.255 4.607

1.790 1.974 5.393 10.613 2.420 3.115 3.678 5.025

1.830 2.006 5.715 11.969 3.270 3.857 4.090 5.439

CO2 + H2O + EtOH

Figure 6. Hydrate depression temperature for (a) sodium chloride and (b) ethanol.

To quantify the inhibiting effects of those two inhibitors on carbon dioxide hydrate equilibria, the hydrate depression temperature, ΔTH = T0 − TI, between carbon dioxide hydrate equilibria in pure water (T0) and that in the aqueous inhibitor solution (TI) was calculated. The hydrate depression temperature was calculated assuming regression curves at fixed

pressures for each set of inhibitor and concentration. Table 3 shows the hydrate depression temperature (ΔTH) for each concentration at five different pressures, and these values are compared to the freezing point depression (ΔT) values of aqueous sodium chloride and ethanol solutions.37 Figure 6 3449



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T0 TI ΔT ΔTH xi

compares the inhibiting effects of the sodium chloride and ethanol at specific pressures. In Figure 6, it can be observed that, as the pressure increases, increasing NaCl concentrations in aqueous solutions result in higher depression temperature. This can be understood by the decrease in the CO2 solubility in the inhibited aqueous phase as the salt concentration increases (salting-out effect). The Coulombic forces and the salting-out effect seem to make NaCl a better inhibitor than ethanol. It is also observed in Figure 6 a reversing trend in the hydrate depression temperature at higher pressures for lower inhibitor concentration, which shows that ethanol has a higher inhibition effect than sodium chloride. This can be attributed to the selective solubility caused by the pressure effect and such behavior possibly can be attributed to the salting-in effect for NaCl, that is, the solubility of CO2 increases at low salt concentrations. The salting-in effect at low concentrations is explained by the Debye−Huckel theory.38,39

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AUTHOR INFORMATION

Corresponding Author

*Phone/fax: + 55 41 3310 4907; e-mail: [email protected]. ORCID

Moisés A. Marcelino Neto: 0000-0001-5492-6640 Amadeu K. Sum: 0000-0003-1903-4537 Funding

The authors acknowledge the financial support from ANP and FINEP through the Human Resources Program to Oil and Gas segment PRH-ANP (PRH 10-UTFPR) and from TE/ CENPES/PETROBRAS. Notes

The authors declare no competing financial interest.

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ΔH P R T

REFERENCES

(1) Sabil, K. M.; Witkamp, G.-J.; Peters, C. J. Phase Equilibria of Mixed Carbon Dioxide and Tetrahydrofuran Hydrates in Sodium Chloride Aqueous Solutions. Fluid Phase Equilib. 2009, 284, 38−43. (2) Sloan, E. D.; Koh, C. A. Clathrate Hydrates of Natural Gases; CRC Press/Taylor & Francis: Boca Raton, 2008. (3) Adisasmito, S.; Frank, R. J.; Sloan, E. D. Hydrates of Carbon Dioxide and Methane Mixtures. J. Chem. Eng. Data 1991, 36, 68−71. (4) Deaton, W. M.; Frost, E. M. Gas Hydrates and Their Relation to the Operation of Natural-gas Pipe Lines; Tech. Rep, United States Department of the Interior, 1946. (5) Unruh, C. H.; Katz, D. L. Methane Hydrate at High Pressure. JPT, J. Pet. Technol. 1949, 86, 83−86. (6) von Stackelberg, M. Solid Gas Hydrates. Die Naturwissenschaften 1949, 36, 359−362. (7) Larson, S. D. Phase Studies of the Two-component Carbon-dioxidewater System Involving the Carbon Dioxide Hydrate. Ph.D. Thesis, University of Illinois, Urbana, IL, 1955. (8) Takenouchi, S.; Kennedy, G. C. Dissociation Pressures of the Phase CO2·53/4 H2O. J. Geol. 1965, 73, 383−390. (9) Miller, S. L.; Smythe, W. D. Carbon Dioxide Clathrate in the Martian Ice Cap. Science 1970, 170, 531−532. (10) Robinson, D. B.; Mehta, B. R. Hydrates in the Propane-Carbon Dioxide-Water System. J. Can. Pet. Technol. 1971, 10, 33−35. (11) Vlahakis, J. G.; Chen, H.-S.; Suwandi, M. S.; Barduhn, A. J. The Growth Rate of Ice Crystals: The Properties of Carbon Dioxide Hydrate, a Review of Properties of 51 Gas Hydrates; Tech. Rep, United States Department of the Interior, 1972. (12) Falabella, B. J. A Study of Natural Gas Hydrates. Ph.D. Thesis, University of Massachusetts, Amherst, MA, 1975. (13) Ng, H.-J.; Robinson, D. B. Equilibrium Phase Composition and Hydrating Conditions in Systems Containing Methanol, Light Hydrocarbons, Carbon dioxide and Hydrogen sulfide; Gas Processors Association, 1983. (14) Song, K. Y.; Kobayashi, R. Water Content of CO2 in Equilibrium with Liquid Water and/or Hydrates. SPE Form. Eval. 1987, 2, 500−508. (15) Dholabhai, P. D.; Kalogerakis, N.; Bishnoi, P. R. Equilibrium Conditions for Carbon Dioxide Hydrate Formation in Aqueous Electrolyte Solutions. J. Chem. Eng. Data 1993, 38, 650−654. (16) Ohgaki, K.; Makihara, Y.; Takano, K. Formation of CO2 Hydrate in Pure and Sea Waters. J. Chem. Eng. Jpn. 1993, 26, 558−564. (17) Komai, T.; Yamamoto, Y.; Ikegami, S. Equilibrium Properties and Kinetics of Methane and Carbon Dioxide Gas Hydrate Formation/dissociation. Fuel Chemistry Symposium on Gas Hydrates, San Francisco, CA, 1997; pp 568−572. (18) Yoon, J.-H.; Lee, H. Clathrate Phase Equilibria for the Water− Phenol−Carbon Dioxide System. AIChE J. 1997, 43, 1884−1893. (19) Nakano, S.; Moritoki, M.; Ohgaki, K. High-Pressure Phase Equilibrium and Raman Microprobe Spectroscopic Studies on the CO2 Hydrate System. J. Chem. Eng. Data 1998, 43, 807−810. (20) Fan, S.; Guo, T. M. Hydrate Formation of CO2-Rich Binary and Quaternary Gas Mixtures in Aqueous Sodium Chloride Solutions. J. Chem. Eng. Data 1999, 44, 829−832. (21) Wendland, M.; Hasse, H.; Maurer, G. Experimental Pressure Temperature Data on Three- and Four-Phase Equilibria of Fuid, Hydrate and Ice Phases in the System Carbon Dioxide-Water. J. Chem. Eng. Data 1999, 44, 901−906. (22) Fan, S. S.; Chen, G.-J.; Ma, Q. L.; Guo, T. M. Experimental and Modeling Studies on the Hydrate Formation of CO2 and CO2-rich Gas Mixtures. Chem. Eng. J. 2000, 78, 173−178.

4. CONCLUSIONS New equilibrium data for carbon dioxide hydrates inhibited with sodium chloride and ethanol were experimentally reported at temperatures ranging from 272.15 to 279.65 K and pressures up to 3.91 MPa using the isothermal method. The experimental three-phase equilibrium data (LW−H−V) were compared with those in the literature, showing consistent agreement. The Clausius−Clapeyron equation was used to estimate the enthalpy of dissociation of the hydrates in those systems, and it was found that these enthalpies do not differ so much in the inhibitor presence, mainly because there is no change in hydrate structure. The hydrate depression temperature was used to compare the inhibiting effect of the sodium chloride and ethanol in the carbon dioxide hydrate equilibrium. Two trends were observed. For higher inhibitor concentrations, with increasing pressure, higher depression temperature were observed for sodium chloride, indicating that it is a stronger hydrate inhibitor than ethanol at the same mole fraction. Yet, for lower inhibitor concentrations, ethanol exhibited a higher inhibition effect than sodium chloride. Those two trends possibly can be attributed to the selective solubility caused by the pressure effect and the salting-out and salting-in effect for the CO2 solubility.

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Hydrate equilibrium temperature (K) Hydrate equilibrium temperature with inhibitor (K) Freezing point depression (K) Hydrate depression temperature (K) Molar fraction of component i (mol mol−1)

NOMENCLATURE Enthalpy of dissociation (J mol−1) Pressure (Pa) Universal gas constant [= 8.314] (J mol−1 K−1) Temperature (K) 3450



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(23) Yang, S. O.; Yang, I. M.; Kim, Y. S.; Lee, C. S. Measurement and Prediction of Phase Equilibria for Water+CO2 in Hydrate Forming Conditions. Fluid Phase Equilib. 2000, 175, 75−89. (24) Servio, P.; Englezos, P. Effect of Temperature and Pressure on the Solubility of Carbon Dioxide in Water in the Presence of Gas Hydrate. Fluid Phase Equilib. 2001, 190, 127−134. (25) Hachikubo, A.; Miyamoto, A.; Hyakutake, K.; Abe, K.; Shoji, H. Phase Equilibrium Studies on Gas Hydrates Formed from Various Guest Molecules and Powder Ice. The Fourth International Conference on Gas Hydrates, Yokohama, Japan, vol. 19−23, 2002; pp 357−360. (26) Zhang, Y. Formation of Hydrate from Single-phase Aqueous Solutions. Master’s Thesis, University of Pittsburgh, 2003. (27) Mohammadi, A. H.; Anderson, R.; Tohidi, B. Carbon Monoxide Clathrate Hydrates: Equilibrium Data and Thermodynamic Modeling. AIChE J. 2005, 51, 2825−2833. (28) Chapoy, A.; Haghighi, H.; Burgass, R.; Tohidi, B. On The Phase Behaviour of the (Carbon Dioxide + Water) Systems at Low Temperatures: Experimental and Modeling. J. Chem. Thermodyn. 2012, 47, 6−12. (29) Vlahakis, J. G.; Chen, H.-S.; Suwandi, M. S.; Barduhn, A. J. The Growth Rate of Ice Crystals: The Properties of Carbon Dioxide Hydrate; a Review of Properties of 51 Gas Hydrates; Data Base Office of Saline Water, United States Department of the Interior, 1972. (30) Ruffine, L.; Trusler, J. P. M. Phase Behaviour of Mixed-gas Hydrate Systems Containing Carbon Dioxide. J. Chem. Thermodyn. 2010, 42, 605−611. (31) Mohammadi, A.; Afzal, W.; Richon, D. Experimental Data and Predictions of Dissociation Conditions for Ethane and Propane Simple Hydrates in the Presence of Distilled Water and Methane, Ethane, Propane, and Carbon Dioxide Simple Hydrates in the Presence of Ethanol Aqueous Solutions. J. Chem. Eng. Data 2008, 53, 73−76. (32) Maekawa, T. J. Equilibrium Conditions for Carbon Dioxide Hydrates in the Presence of Aqueous Solutions of Alcohols, Glycols, and Glycerol. J. Chem. Eng. Data 2010, 55, 1280−1284. (33) Ferrari, P. F.; Guembaroski, A. Z.; Marcelino Neto, M. A.; Morales, R. E. M.; Sum, A. K. Experimental Measurements and Modelling of Carbon Dioxide Hydrate Phase Equilibrium with and without Ethanol. Fluid Phase Equilib. 2016, 413, 176−183. (34) Dohrn, R.; Peper, S.; Fonseca, J. M. High-pressure Fluid-Phase Equilibria: Experimental Methods and Systems Investigated (2000− 2004). Fluid Phase Equilib. 2010, 288, 1−54. (35) Smith, J. M.; Ness, H. C. V.; Abbott, M. M. Introduction to Chemical Engineering Thermodynamics; McGraw Hill: Boston, 2004. (36) Sun, Q.; Tian, H.; Li, Z.; Guo, X.; Liu, A.; Yang, L. Solubility of CO2 in Water and NaCl Solution in Equilibrium with Hydrate. Part I: Experimental Measurement. Fluid Phase Equilib. 2016, 409, 131−135. (37) Lide, D. R. Handbook of Chemistry and Physics; CRC Press: Boca Raton, 2005. (38) Debye, P.; Hückel, E. Zur Theorie der Elektrolyte. Physikalische Zeitschrift 1923, 24, 185−206. (39) Englezos, P.; Bishnoi, P. R. Prediction of Gas Hydrate Formation Conditions in Aqueous Electrolyte Solutions. AIChE J. 1988, 34, 1718−1721. (40) Godishala, K. K.; Sangwai, J. S.; Sami, N. A.; Das, K. Phase Stability of Semiclathrate Hydrates of Carbon Dioxide in Synthetic Sea Water. J. Chem. Eng. Data 2013, 58, 1062−1067. (41) Sami, N. A.; Das, K.; Sangwai, J. S.; Balasubramanian, N. Phase Equilibria of Methane and Carbon Dioxide Clathrate Hydrates in the Presence of (Methanol + MgCl2) and (Ethylene Glycol + MgCl2) aqueous solutions. J. Chem. Thermodyn. 2013, 65, 198−203.

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