Article pubs.acs.org/jced
Phase Equilibria of Carbon Dioxide Hydrates in the Presence of Methanol/Ethylene Glycol and KCl Aqueous Solutions Majid Dastanian, Amir Abbas Izadpanah,* and Masoud Mofarahi Department of Chemical Engineering, Faculty of Petroleum, Gas and Petrochemical Engineering, Persian Gulf University, Bushehr, Iran
ABSTRACT: In this study, the experimental data for dissociation conditions of carbon dioxide hydrates in the presence of 0.05 and 0.1 mass fraction KCl solution + 0.1 and 0.2 mass fraction methanol and ethylene glycol were measured and then reported at different temperatures and pressure ranges not available in the related literature. The phase equilibrium curves were drawn using an isochoric pressure-search method. To validate the used apparatus and the experimental findings of the current study and also to show the inhibition effects of the aqueous solutions used in this study, the experimental values were compared with some selected experimental data from the literature about the dissociation conditions of carbon dioxide hydrates in the presence of pure water and aqueous solution with 0.05 mass fraction KCl. Finally, to examine the inhibitory effect of various inhibitors and their synergies on each other, the suppressed temperature for hydrate formation was evaluated in the presence of different inhibitor solutions. This value showed that the rate of suppressed temperature for hydrate formation for each solution has been almost constant in various pressures. The synergy effect of KCl with methanol or glycol at low concentrations is negligible indicating that these two inhibitors have no impact on each other. It was also shown that, by increasing the concentration of the inhibitors, this rule was violated, the inhibitors were affected by each other, and the amount of inhibition effect increases. This synergy is of utmost importance for oil and gas pipelines and also for the industrial equipment that naturally contain some salt, in which alcohol or glycol will be added to prevent hydrate formation.
1. INTRODUCTION
The usual methods for preventing gas hydrates formation are classified as one of the following methods or a combination of these methods: (1) the injection of inhibitor to the system, (2) dehydrating the flow of gas, and (3) maintaining temperature and pressure conditions of gas pipeline in a zone which is outside the hydrate stability zone through insulation or heating the pipeline. It should be noted that among these methods the inhibitor injection is used as the most practical method in the oil industry.3 The selection of hydrate inhibitor type is of paramount important in the oil and gas industry. In this regard, inhibitors should be easily accessible and economical; also, they should be completely soluble in water, be easily recyclable from water, and be able to lower the hydrate formation temperature as much as possible. Furthermore, inhibitors should not react with components available in the gas stream or be deposited,
Gas hydrates are complex nonstoichiometric crystalline structures that are produced usually at a low temperature and high pressure from a mixture of water (host) and lowmolecular-weight gases (guest).1 After the discovery of hydrate clathrates by Sir Humphry Davy in 1810, Hummer Schmidt in 1934 found that the obstruction of gas pipelines was due to hydrate formation. Since then, preventing hydrate formation in the oil and gas industries and networks of transmission lines has become one of the favorable domains of study for the researchers. Due to the stability of gas hydrates at temperatures above the freezing water temperature, these substances can cause obstruction in the pipelines, nozzles, valves, trays of the distillation towers, and other industrial equipment.2 Studies showed that hydrate formation in pipelines and chemical processes happens in the presence of four conditions: (i) presence of water, (ii) presence of gas, (iii) low temperature, and (iv) high pressure. © XXXX American Chemical Society
Received: February 8, 2017 Accepted: April 3, 2017
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DOI: 10.1021/acs.jced.7b00146 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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nor should they be flammable or increasing the flammability nature of the gas; also their viscosity and freezing point should not be low. Finally, their pressure steam should be low to prevent the inhibitor from being wasted.4 The most practical and the most common method for preventing the hydrate formation is making use of thermodynamic inhibitors.1 Thermodynamic inhibitors are soluble chemicals such as methanol, ethanol, monoethylene glycol (MEG), and diethylene glycol (DEG) available in water that prevent hydrate formation by shifting the hydrate dissociation curve toward lower temperatures and higher pressures. These compositions generally make hydrogen bonds with water molecules, creating disorders in the hydrogen bonds between water molecules generating crystal; in this way they prevent the formation of gas storage shelves.3 It should be noted that in many cases, especially in subsea wells, the extracted oil and gas contain some salty water leading to a reduction in the hydrate formation tendency. The reason behind this is the existence of some electrolytes such as NaCl, KCl, and CaCl2 in the salted water. Due to the presence of strong electrostatic charge, the sea salts act as natural inhibitors of hydrate formation.1 In cases where the inhibition power of the electrolyte is not sufficient, methanol or ethylene glycol solutions are added to the system. Therefore, it is the essential to have data with high accuracy obtained from hydrate dissociation equilibrium conditions with the simultaneous presence of electrolyte and methanol or ethylene glycol.5 Most of data reported in the literature was related to the systems only containing glycol, alcohol, or saline solution, and there were only few data related to systems containing a mixture of salt and alcohol/glycol.6−8 Furthermore, in most cases the data available in the literature are related to the systems containing a mixture of salt and alcohol/glycol for methane.9−12 To the best of our knowledge in this domain, there are not much data available about systems containing a mixture of salt and alcohol/glycol for CO2 which as one of the principal components of natural gas indicates a high tendency to be dissolved in the water phase due to high polarity (as opposed to methane). The first research that studied the effects of alcohol and electrolytes simultaneously on the balance point of hydrate formation was conducted by Dholabhai et al. (1996). In 1996, for the first time they found the equilibrium point of carbon dioxide hydrate formation for some solutions containing methanol along with NaCl or KCl.13 They conducted the same experiments in their subsequent studies for CaCl2 and methanol, even for ethylene glycol and NaCl or CaCl2 and for methane and H2S.14,15 Mohammadi and Richon (2009 and 2012) using a constantvolume laboratory apparatus examined the equilibrium point of hydrates dissociation for some gases like methane, carbon dioxide, and hydrogen sulfide in the presence of ethylene glycol or methanol and salt. To examine the validity of the obtained data, Mohammadi and Richon (2009 and 2012) made use of some thermodynamic models. They stated that the existing models for the solutions containing salt gave acceptable deals only in the absence or in the presence of a small amount of alcohol or glycol, and by increasing the concentration of alcohol/glycol the accuracy of the models (even for solutions without salt) decreases extremely.10,16 Moreover, Najibi et al. (2013 and 2014) using a constant-volume laboratory apparatus examined the equilibrium point of carbon dioxide hydrate dissociation in the presence of ethylene glycol or methanol
along with salt.5,17 They compared their experimental data with the thermodynamic models available in the literature;18,19 this comparison indicated that there is a difference between the experimental data and the results obtained from modeling especially for inhibitor mixtures containing ethylene glycol. Table 1 indicates the equilibrium data of CO2 hydrate dissociation in the presence of KCl + methanol or ethylene Table 1. Equilibrium Data of CO2 Hydrate Dissociation in the Presence of KCl + Methanol or Ethylene Glycol solution ID
KCl (mass %)
methanol (mass %)
K10Me5 K10Me10 K10Me20 K10EG20
10 10 10 10
5 10 20
ethylene glycol (mass %)
no. equilibrium points
ref
20
4 3 4 4
13 13 5 5
glycol available in the literature. As it can be seen, there are only a few data in this regard, and the inability of the available thermodynamic models in predicting equilibrium conditions of CO2 hydrate dissociation can be attributed to this factor. In the current study, using a constant-volume apparatus the equilibrium data of CO2 hydrate dissociation in the presence of KCl salt, as one of the salts available in seawater was measured at two concentrations of 5 and 10 mass percent along with ethylene glycol or methanol at concentrations of 10 and 20 mass %. To validate the measured experimental data, the data on CO2 hydrate dissociation conditions were measured in the presence of pure water and KCl; then these data were compared with the data available in the literature on KCl salt20 and pure water.6 By comparing the equilibrium data of CO2 hydrate dissociation conditions in the presence of KCl and alcohol or glycol with some selected experimental equilibrium data on the dissociation conditions of carbon dioxide hydrates in the presence of pure water,6 the inhibition effects and inhibitory values of the solutions used in this study were examined and evaluated. Moreover, by comparing equilibrium data on CO2 hydrate dissociation conditions for solutions containing KCl + ethylene glycol or methanol with the equilibrium data available in the literature on CO2 hydrate dissociation conditions for solutions containing only methanol6,13 (or ethylene glycol6) and only KCl,20 the synergy effect value of these two inhibitors (electrolyte + alcohol or glycol) on each other was also investigated.
2. EXPERIMENTAL SECTION 2.1. Chemicals. Table 2 indicates the purity value and the manufacturer of chemicals used in the current study. The required solutions were prepared by a gravimetric method using a digital balance with an uncertainty value of ±0.001 g. Table 2. Purities and Suppliers of Chemicalsa
a
B
chemical
supplier
mole fraction purity
carbon dioxide methanol ethylene glycol KCl
Persian gas Merck Merck Merck
0.999 0.999 0.99 0.995
Deionized water was used in all experiments. DOI: 10.1021/acs.jced.7b00146 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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2.2. Experimental Apparatus. Figure 1 shows a schematic diagram of the apparatus used in this study. Briefly, the main
Figure 1. Schematic picture of the apparatus used in this work.
solution, commensurate to the capacity of our vessel, was shed into the sample vessel, and then the air inside the vessel was evacuated using a vacuum pump. Next, the sample vessel was immersed into a glycol bath having the capabilities for the glycol to be circulated and temperature to be controlled, and then CO2 was introduced into the vessel from a high pressure cylinder and the valve between the cylinder and the vessel was closed. In order to form the hydrate, the temperature was decreased step by step. Hydrate formation within the sample vessel is detected through a drastic reduction in pressure and increasing temperature of the vessel. Subsequently, the temperature was increased step by step with 0.2−1 K. Reaching the equilibrium conditions in every step can be detected through pressure stability in the sample vessel. As a result, a pressure−temperature (P−T) diagram will be obtained for each experimental run. The point at which the slope of the P−T curve suddenly changes is reported as the equilibrium point of hydrate dissociation3. In the other words, the intersection between the cooling and heating curves was considered as the equilibrium point of hydrate dissociation as shown in Figure 2.
part of this apparatus is a cylindrical vessel with a volume of about 800 cm3 which can withstand the pressure up to 10 MPa. Also, a mechanical stirrer has been used to ensure sufficient agitation to the contents of sample vessel and to facilitate reaching equilibrium as well. The sample vessel is immersed into the circular glycol bath with a maximum cooling temperature of −22 °C. The temperature of the glycol bath is controlled by a PID controller (Autonics TZ4M). The accuracy of the controller for maintaining a constant temperature is 0.1 K. The temperature and pressure of the sample vessel were measured by a Pt100 and a BD pressure transmitter sensor, respectively. The indicators Autonics TZ4M and Autonics TZ4W were used to display and record temperature and pressure data, with a resolution of 0.1 K and 0.01 MPa and the uncertainties of 0.1 K and 0.03 MPa, respectively. 2.3. Experimental Procedure. In the current study, for measuring the equilibrium point of the hydrate dissociation conditions, the isochoric pressure-search method was used. The details of this procedure can be found in the literature.5,10 Briefly, based on this method, at first 200 cm3 of aqueous
3. RESULTS AND DISCUSSION To check the reliability of the results obtained from the experimental apparatus used, three experimental data from dissociation point of carbon dioxide hydrate in the presence of pure water and four equilibrium points in the presence of the 0.05 mass fraction KCl aqueous solution were measured and compared with some selected data from literature. The results of this comparison are indicted in Figures 3 and 4. As can be observed, there is an acceptable accordance between our experimental data and data reported in the literature.6,8,20 Table 3 shows the composition of the components in aqueous solutions used in this study. The experimental data obtained from the equilibrium point of hydrate dissociation for the solutions mentioned in Table 3 are reported in Tables 4 and 5 and drawn in Figures 5 and 6. Furthermore, some of the equilibrium data of carbon dioxide hydrate dissociation in the presence of pure water reported in the literature6 are also presented in these figures. The reason behind this was to show the inhibitory power of the solutions used in the experiment. It should be noted that the meaning behind inhibitory power is a change in the equilibrium conditions of hydrate dissociation
Figure 2. An example of the pressure−temperature diagram taken from (P−T) trace of CO2 hydrate in an aqueous solution of 0.05 mass fraction KCl + 0.1 mass fraction ethylene glycol shows hydrate formation and the dissociation process, with the determination of the equilibrium point. C
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Table 4. Experimental Dissociation Data of Carbon Dioxide Clathrate Hydrates in the Presence of KCl + Methanol Aqueous Solution Ta (K)
Pb (MPa)
0.05 Mass Fraction KCl + 0.10 Mass Fraction Methanol 275.7 3.59 274.4 3.07 273.5 2.72 271.9 2.25 269.9 1.77 0.05 Mass Fraction KCl + 0.20 Mass Fraction Methanol 267.4 3.23 266.0 2.68 264.7 2.26 263.3 1.90 a
Measured with 0.1 K uncertainty. uncertainty.
Figure 3. Experimental dissociation conditions of carbon dioxide clathrate hydrates in the presence of pure water. Symbols represent the experimental hydrate dissociation data: (●) this work; (△) ref 6 and (□) ref 8.
b
Table 5. Experimental Dissociation Data of Carbon Dioxide Clathrate Hydrates in the Presence of KCl + Ethylene Glycol Aqueous Solution Ta (K)
Table 3. Composition of Aqueous Solutiona K5Me10 K5Me20 K5EG10 K5EG20 K10EG10 a
a
Measured with 0.1 K uncertainty. uncertainty.
KCl (mass %) methanol (mass %) ethylene glycol (mass %) 5 5 5 5 10
Pb (MPa)
0.05 Mass Fraction KCl + 0.10 Mass Fraction Ethylene Glycol 277.4 3.54 275.6 2.87 273.8 2.25 271.8 1.73 270.2 1.38 0.05 Mass Fraction KCl + 0.20 Mass Fraction Ethylene Glycol 273.4 3.61 272.2 3.22 270.6 2.66 268.7 2.03 266.9 1.65 0.10 Mass Fraction KCl + 0.10 Mass Fraction Ethylene Glycol 275.8 3.65 274.6 3.12 272.8 2.44 270.6 1.88 268.4 1.48
Figure 4. Dissociation conditions of carbon dioxide clathrate hydrates in the presence of 0.05 mass fraction KCl. Symbols represent the experimental hydrate dissociation data: (●) this work; (□) ref 20.
solution ID
Measured with 0.03 MPa
10 20 10 20 10
b
Measured with 0.03 MPa
Furthermore, the measured experimental data were compared with predicted values from a thermodynamic model proposed in the literature.21,22 Figure 7 indicates the results obtained from the comparison of solutions containing KCl + methanol, and Figure 8 shows the results obtained from the comparison of experimental data and modeling for solutions containing KCl + ethylene glycol. The results obtained from both of these comparisons indicate that the so-called model does not accurately predict the high concentrations of methanol or ethylene glycol in such systems. This finding is consistent with other research which has compared the experimental data of hydrate dissociation in the presence of electrolyte + organic inhibitor with the available thermodynamic models.5,10 Much of the uncertainty in predicting the equilibrium point of hydrate dissociation using thermodynamic models can be attributed to the inability of such models to calculate the interactions among the electrolyte, alcohol/glycol, and water.10
All data measured with 0.001 g uncertainty.
toward higher pressures and lower temperatures due to the presence of inhibitors. Moreover, Figures 5 and 6 indicate the equilibrium curve of carbon dioxide hydrate for solutions containing 0.1 mass fraction KCl + 0.2 mass fraction ethylene glycol (K10EG20), 0.1 mass fraction KCl + 0.1 mass fraction methanol (K10Me10), and 0.1 mass fraction KCl + 0.2 mass fraction methanol (K10Me20). The reason behind this was to observe and compare the inhibitory power of these solutions with the solutions used in this study. As indicated in these figures, by adding an inhibitor to the system, equilibrium curves are shifted toward the inhibitory region. Also, by increasing the concentration of the inhibitors, this tendency becomes more evident. D
DOI: 10.1021/acs.jced.7b00146 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Figure 5. Experimental dissociation conditions of carbon dioxide hydrates in the presence of KCl + methanol aqueous solutions. The symbols represent the experimental dissociation conditions: (○) pure water;6 (□) K5Me10, this study; (■) K5Me20, this study; (△) K10Me10;13 (▲) K10Me20.5
Figure 8. Comparison between experimental data and modeling results21 of equilibrium conditions of carbon dioxide gas hydrate in the presence of KCl + ethylene glycol.
the presence of pure water and the temperature of hydrate formation in the presence of selected inhibitors. For a better comparison of the suppressed temperatures of hydrate formation, these values were compared to each other at pressures of 1.5, 2, 2.5, and 3 MPa. To evaluate the equilibrium temperature in these pressures, a polynomial of degree 3 was fitted to experimental equilibrium data, and then this function was used to calculate equilibrium temperature at each pressure. The “sum ΔT” symbol in this table represents the sum of reductions in the temperature of hydrate formation in the presence of KCl salt alone and reductions in the temperature of hydrate formation in the presence of methanol or ethylene glycol alone. For example, the sum ΔT for the K5Me20 solution is equal to the reductions in the temperature of hydrate formation for KCl 5% solution + reductions in the temperature of hydrate formation for 20% methanol solution. As shown in Table 6, the suppressed temperature of hydrate formation for each solution has been almost constant in different pressures, and it can be considered independent of the pressure. The synergy value for KCl along with methanol or glycol is almost zero in low concentrations indicating that these two inhibitors have no impact on each other. This phenomenon can be explained in this way that organic inhibitors like methanol and ethylene glycol prevent hydrate formation by making hydrogen bonds with water molecules and influencing guest− host interactions, whereas the salts prevent hydrate formation by decreasing the host concentration through a phenomenon named as “salting-out”. By increasing the inhibitors concentration, this rule will be violated, and the inhibitors will influence each other; as a result, the inhibitory value increases. Moreover, Lee et al.1 found similar results in their hydrate formation experiments on synthesized natural gas in the presence of NaCl + methanol/ethylene glycol. They also reported that the pressure and low concentration of inhibitors had no effect on the synergy of salt and organic inhibitor. In their experiments, a deviation from this rule was also observed at high concentrations of salt and organic inhibitors; the inhibitors had impacts on each other, and the effect of synergy was significant.1
Figure 6. Experimental dissociation conditions of carbon dioxide hydrates in the presence of KCl + ethylene glycol aqueous solutions. The symbols represent the experimental dissociation conditions: (○) pure water;6 (□) K5EG10, this work; (■) K5EG20, this work; (△) K10EG10, this work; (▲) K10EG20.5
Figure 7. Comparison between experimental data and modeling results22 of equilibrium conditions of carbon dioxide gas hydrate in the presence of KCl + methanol.
In this study, to examine the inhibitory effect of different inhibitors as well as their synergy effects on each other, the suppressed temperature of hydrate formation was obtained in the presence of different inhibitors as shown in Table 6. The suppressed temperature of hydrate formation means the difference between the temperature of hydrate formation in E
DOI: 10.1021/acs.jced.7b00146 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 6. Suppression Temperature of Carbon Dioxide Hydrate Formation at Various Inhibitor Concentrations experimental ΔT
solution pressure (MPa) 20
K5 K108,20 Me106 Me206,13 EG106 EG206 K5Me10 K5EG10 K5Me20 K5EG20 K10Me1013 K10EG10 K10Me205 K10EG205
sum ΔT
1.5
2.0
2.5
3.0
1.65 3.41 4.52 10.01 2.61 5.63 6.19 3.91 13.46 8.62 7.57 6.21 15.34 10.61
1.60 3.34 4.36 9.94 2.52 5.65 6.13 4.23 13.43 8.56 7.44 5.97 16.26 10.65
1.62 3.34 4.45 10.10 2.54 5.71 6.14 4.44 13.51 8.74 7.64 6.00 16.81 10.53
1.70 3.32 4.61 10.24 2.61 5.79 6.15 4.47 13.58 8.84
14.97
synergy
1.5
2.0
2.5
3.0
1.5
2.0
2.5
3.0
6.17 4.26 11.66 7.28 7.93 6.02 13.42 9.04
5.96 4.12 11.54 7.25 7.70 5.86 13.28 8.99
6.07 4.16 11.72 7.33 7.79 5.88 13.44 9.05
6.31 4.31 11.94 7.49 7.93 5.93 13.56 9.11
0.02 −0.35 1.80 1.34 −0.36 0.19 1.92 1.57
0.17 0.11 1.89 1.31 −0.26 0.11 2.98 1.66
0.07 0.28 1.79 1.41 −0.15 0.12 3.37 1.48
−0.16 0.16 1.64 1.35
■
4. CONCLUSIONS In this study, new experimental data have been measured and reported for carbon dioxide hydrate dissociation conditions in the presence of methanol + KCl and ethylene glycol + KCl aqueous solutions under various inhibitor concentrations and in different temperature and pressure conditions. The phase equilibrium curves were drawn using an isochoric pressuresearch method. To validate the used apparatus and the experimental findings in this study and also to show the inhibitory power of the used solutions, these experimental values were compared with the experimental data on the dissociation conditions of carbon dioxide hydrates related to the pure water6,8 and 0.05 mass fraction KCl aqueous solution available in the literature.20 The results show that, in the systems like ethylene glycol/ methanol + KCl + H2O + CO2, by increasing the inhibitors concentration the hydrate equilibrium curve is shifted toward higher pressures and lower temperatures. Furthermore, it has been observed that methanol has a higher inhibitory effect than the ethylene glycol at the same mass fraction which this finding is in agreement with the findings reported in the literature.16,23 Furthermore, the measured experimental data were compared with the predicted results from the thermodynamic model.21,22 This comparison indicates the models’ lack of high accuracy in predicting the equilibrium conditions of hydrate dissociation and the need to expand and develop more accurate thermodynamic models for predicting the equilibrium conditions of hydrate dissociation in the simultaneous presence of electrolyte + alcohol/glycol in high concentrations of alcohol/ glycol. In order to examine the inhibitory effect of different inhibitors as well as their synergy effects on each other, the suppressed temperature of hydrate formation was obtained in the presence of different inhibitors and also compared at pressures of 1.5, 2, 2.5, and 3 MPa. The suppressed temperature for each solution has been almost constant in various pressures so that it can be considered to be independent of the pressure. The synergy rate of KCl with methanol or glycol at low concentrations is nearly zero indicating that these two inhibitors have no impact on the each other. By increasing methanol or ethylene glycol concentration, this rule will be violated, and the inhibitors will influence each other; as a result, the inhibitory value increases.
0.18
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Amir Abbas Izadpanah: 0000-0003-2051-5119 Masoud Mofarahi: 0000-0001-8583-7923 Notes
The authors declare no competing financial interest.
■
REFERENCES
(1) Lee, J.-W.; Kang, S.-P. Phase equilibria of natural gas hydrates in the presence of methanol, ethylene glycol, and NaCl aqueous solutions. Ind. Eng. Chem. Res. 2011, 50, 8750−8755. (2) Moradi, S.; Haghtalab, A.; Fazlali, A. Prediction of hydrate formation conditions in the solutions containing electrolyte and alcohol inhibitors and their mixtures using UNIQUAC-NRF models. Fluid Phase Equilib. 2013, 349, 61−66. (3) Tariq, M.; Rooney, D.; Othman, E.; Aparicio, S.; Atilhan, M.; Khraisheh, M. Gas Hydrate Inhibition: a review of the role of ionic liquids. Ind. Eng. Chem. Res. 2014, 53, 17855−17868. (4) Makogon, I. U. r. F.; Cieslewicz, W. J. Hydrates of natural gas; PennWell Books: Tulsa, OK, 1981. (5) Najibi, H.; Kamali, Z.; Mohammadi, A. H. Phase equilibria of carbon dioxide clathrate hydrates in the presence of methanol/ ethylene glycol+ single salt aqueous solutions: experimental measurement and prediction. Fluid Phase Equilib. 2013, 342, 71−74. (6) Maekawa, T. 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. (7) Mohammadi, A. H.; Richon, D. Gas hydrate phase equilibrium in the presence of ethylene glycol or methanol aqueous solution. Ind. Eng. Chem. Res. 2010, 49, 8865−8869. (8) 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. (9) Masoudi, R.; Tohidi, B.; Anderson, R.; Burgass, R. W.; Yang, J. Experimental measurement and thermodynamic modelling of clathrate hydrate equilibria and salt solubility in aqueous ethylene glycol and electrolyte solutions. Fluid Phase Equilib. 2004, 219, 157−163. (10) Mohammadi, A. H.; Richon, D. Methane hydrate phase equilibrium in the presence of salt (NaCl, KCl, or CaCl2)+ ethylene glycol or salt (NaCl, KCl, or CaCl2)+ methanol aqueous solution: Experimental determination of dissociation condition. J. Chem. Thermodyn. 2009, 41, 1374−1377. F
DOI: 10.1021/acs.jced.7b00146 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
Journal of Chemical & Engineering Data
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(11) Najibi, H.; Chapoy, A.; Haghighi, H.; Tohidi, B. Experimental determination and prediction of methane hydrate stability in alcohols and electrolyte solutions. Fluid Phase Equilib. 2009, 275, 127−131. (12) Lafond, P. G.; Olcott, K. A.; Sloan, E. D.; Koh, C. A.; Sum, A. K. Measurements of methane hydrate equilibrium in systems inhibited with NaCl and methanol. J. Chem. Thermodyn. 2012, 48, 1−6. (13) Dholabhai, P. D.; Parent, J. S.; Bishnoi, P. R. Carbon dioxide hydrate equilibrium conditions in aqueous solutions containing electrolytes and methanol using a new apparatus. Ind. Eng. Chem. Res. 1996, 35, 819−823. (14) Majumdar, A.; Mahmoodaghdam, E.; Bishnoi, P. R. Equilibrium hydrate formation conditions for hydrogen sulfide, carbon dioxide, and ethane in aqueous solutions of ethylene glycol and sodium chloride. J. Chem. Eng. Data 2000, 45, 20−22. (15) Eichholz, C.; Majumdar, A.; Clarke, M.; Oellrich, L.; Bishnoi, P. Experimental investigation and calculation of methane hydrate formation conditions in the presence of ethylene glycol and sodium chloride. J. Chem. Eng. Data 2004, 49, 847−851. (16) Mohammadi, A. H.; Richon, D. Phase equilibria of hydrogen sulfide and carbon dioxide simple hydrates in the presence of methanol,(methanol+ NaCl) and (ethylene glycol+ NaCl) aqueous solutions. J. Chem. Thermodyn. 2012, 44, 26−30. (17) Najibi, H.; Amiri, E.; Mohammadi, A. H. Experimental measurement and prediction of dissociation condition for carbon dioxide clathrate hydrates in the presence of methanol/ethylene glycol + sodium chloride/magnesium chloride aqueous solution. Fluid Phase Equilib. 2014, 363, 70−73. (18) Javanmardi, J.; Moshfeghian, M. A new approach for prediction of gas hydrate formation conditions in aqueous electrolyte solutions. Fluid Phase Equilib. 2000, 168, 135−148. (19) Nasrifar, K.; Moshfeghian, M.; Maddox, R. Prediction of equilibrium conditions for gas hydrate formation in the mixtures of both electrolytes and alcohol. Fluid Phase Equilib. 1998, 146, 1−13. (20) 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. (21) Osfouri, S.; Azin, R.; Gholami, R.; Izadpanah, A. A. Modeling hydrate formation conditions in the presence of electrolytes and polar inhibitor solutions. J. Chem. Thermodyn. 2015, 89, 251−263. (22) Osfouri, S.; Azin, R.; Gholami, R.; Moshfeghian, M. Wilson Non-random Factor Reference State Based Model for Prediction of Gas Hydrate Formation Conditions in the Presence of Electrolyte and/or Alcohol in Solution. J. Solution Chem. 2015, 44, 1382−1406. (23) 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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