Ind. Eng. Chem. Res. 1996, 35, 819-823
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Carbon Dioxide Hydrate Equilibrium Conditions in Aqueous Solutions Containing Electrolytes and Methanol Using a New Apparatus Pankaj D. Dholabhai, J. Scott Parent, and P. Raj Bishnoi* Department of Chemical and Petroleum Engineering, The University of Calgary, 2500 University Dr. N.W., Calgary, Alberta, Canada T2N 1N4
A knowledge of the gas hydrate formation conditions, i.e., pressure, temperature, and compositions, of the coexisting phases at equilibrium, is required in formulating processes to avoid their formation in natural gas transmission lines or processing equipment. Additives like methanol and electrolytes are often used to inhibit hydrate formation in industrial operations. Computation of the inhibition effects of these additives is necessary for the design of such operations. Development of thermodynamic methods to calculate the hydrate equilibria conditions requires accurate experimental data. In the present work experimental three-phase (aqueous liquid solution, vapor, and incipient hydrate) equilibrium conditions of CO2 hydrate formation in aqueous solutions of electrolytes and methanol are measured in the temperature range of 263279 K and pressure range of 0.9-3.0 MPa. A new “full view” sapphire tube equilibrium cell and the associated equipment are designed, fabricated, and used for the measurements. The equipment details and the measured data are also reported here. Introduction One of the problems encountered during natural gas production, gathering, processing, and transportation is the formation of gas hydrates. The industry circumvents the potential problem of hydrate formation through a variety of means like dehydration, operating outside the hydrate-forming conditions, and the use of additives like methanol that inhibit hydrate formation. Knowledge of the incipient hydrate-forming conditions, i.e., the pressure, temperature, and compositions of the coexisting phases at equilibrium, is essential for the application of these techniques. Experimental data and prediction methods validated with such data are used for this purpose. Experimental data in pure water have been reported extensively. A large number of these data have been compiled by Sloan (1990). Similarly, prediction methods for simple situations involving pure water have been available to the industry for a long time. It is only in the last few years that the data and the corrresponding prediction methods for systems containing methanol or electrolytes have been reported. Englezos and Bishnoi (1988) proposed a method to predict hydrate-forming conditions in systems containing single or mixed electrolytes. Dholabhai et al. (1991), Englezos and Bishnoi (1991), and Bishnoi and Dholabhai (1993) reported data on methane, ethane, and propane hydrates in mixed electrolytes. The agreement between the predictions generated by the use of the method and the experimental data was very good. The method of Englezos and Bishnoi (1988), however, is not suitable for gases like CO2 which have significant solubilities in the aqueous phase or for situations where another solvent like methanol is present in addition to the electrolytes. Tse (1993) and Tse and Bishnoi (1994) recently reported a method to predict carbon dioxide hydrate formation conditions in electrolyte solutions. The method was successfully employed to predict the data of Chen (1972) and Larson (1975) for CO2 in single electrolyte solutions and the recently obtained data of Dholabhai et al. (1993) and Dholabhai and Bishnoi (1994) for CO2 and two mixtures of CH4 and CO2 in mixed and single electro* Author to whom correspondence should be addressed.
0888-5885/96/2635-0819$12.00/0
lytes. There is, however, a need to develop a method to predict gas hydrate formation conditions in aqueous solutions containing electrolytes and methanol together. Formulation of a method to meet this need requires experimental data on the hydrate equilibrium in aqueous solutions containing electrolytes and methanol. In the present work experimental, three-phase (aqueous liquid solution, vapor, and incipient hydrate), equilibrium conditions of CO2 hydrate formation in aqueous solutions of electrolytes and methanol are reported. A new full view equilibrium cell made of a sapphire tube and the associated equipment, used for the measurements, is also reported. The measurements were made in a temperature range of 263-279 K and a pressure range of 0.9-3.0 MPa. Experimental Setup Experimental Apparatus. A schematic diagram of the apparatus is given in Figure 1. The main component of the apparatus consists of a sapphire equilibrium cell. It is described in detail by Parent (1993). A section of the cell is shown in Figure 2. It consists of a sapphire tube (3/4 in. i.d. × 11/4 in. o.d. × 41/2 in.) sealed at either end by 316 stainless steel flanges. Two buna-n O-rings on each flange accomplish the seal between the cell interior and the exterior. Three 9/16 in. stainless steel studs are used as spacers and also for holding the sapphire tube and the two flanges together. Each stud is about 1/16 in. longer than the sapphire tube. This clearance prevents an exertion of axial stress on the tube, allowing for an easy assembly. To assemble the cell, a magnetic stirbar is placed on the bottom flange, the sapphire tube is inserted, and the three studs are screwed into the bottom flange. The top flange is then slipped into position, and the top nuts on the studs are tightened. The bottom flange is provided with two ports suitable for 1/8 in. Swagelok fittings. The top flange has three such ports. The cell is provided with connections for charging the solution, a thermocouple in the vapor phase, gas inlet/pressure measurement, and vapor phase sampling to a HP 5840A gas chromatograph (GC). A stainless steel reservoir (1.2 L) partially filled with mercury is connected to the cell in the vapor phase. The total volume of the system (cell and the reservoir) can © 1996 American Chemical Society
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Ind. Eng. Chem. Res., Vol. 35, No. 3, 1996
Figure 1. Schematic of the experimental apparatus.
Figure 2. Cross section of the equilibrium cell.
be changed by varying the quantity of mercury in the reservoir. A 250 mL capacity Ruska piston pump is used for this purpose. The quantity of mercury in the reservoir was adjusted such that the volume of the vapor phase in the cell and the reservoir was 545 mL, with the piston of the Ruska pump at the 0.0 L position. This corresponded to a vapor phase volume of 500 mL in the reservoir. A minimum of 50 mL of vapor phase volume was always maintained in the reservoir. The equilibrium cell and the reservoir are immersed in a constanttemperature bath. A refrigerated ethylene glycolwater solution is used as the circulating coolant. An external magnet mounted on the shaft of a motor under the bath couples with the stirbar in the cell to provide stirring in the cell. The cell pressure is measured with (i) a 0-15 MPa Heise gauge and (ii) a gauge pressure transmitter (DP). The data acquisition system consists of a Digitrend 235 DORIC data logger connected to a PC. Cell pressure and temperature were sampled at predetermined intervals and stored on the PC. The thermocouple was calibrated by immersing it in a constant-temperature bath along with a mercury-in-glass thermometer graduated to 0.1 K. The calibration was performed in the range of 264-290 K. The standard deviation of the linear least-squares fit was 0.13 K. The span of the DP was 12 MPa, with a combined accuracy of (0.25% of the span ((30 kPa). It was calibrated against a dead weight tester. The standard deviation of the linear least-squares fit was 5.4 kPa.
Materials and Preparation of Solutions. The electrolytes used were certified ACS grade supplied by Fisher Scientific Co. Deionized water was distilled in the laboratory before use. Coleman Instrument grade carbon dioxide (Union Carbide) in cylinders with a certified purity of 99.99% was used. Glass-distilled methanol with a stated minimum purity of 99.9% was supplied by BDH. Appropriate quantities of the electrolytes, methanol, and water were weighed into a 50 mL stoppered flask on a top-loading Mettler balance with a readability of 0.01 g. The mixture was shaken at room temperature to dissolve the electrolytes. Analysis of the Vapor Phase. Vapor mixtures of CO2, methanol, and water of known compositions were prepared in our laboratory. The relative response factors between (1) water and methanol and (2) CO2 and methanol were obtained by analyzing these mixtures on the GC. Vapor phase samples from equilibrated mixtures of CO2 and methanol were analyzed, and their compositions were calculated using the response factors. The average absolute deviation between these compositions and those generated using the computer program MEGHA was 5.05% in the mole fraction of methanol. Vapor phase samples from equilibrated mixtures of CO2, methanol, and water were also analyzed. The average absolute deviation between the compositions calculated using the response factors and those predicted using the computer program was 4.1% in the mole fraction of methanol and 9.58% in the mole fraction of water. The temperatures and the pressures of the above-equilibrated mixtures were close to the conditions expected in the present work. The concentrations of methanol in the vapor phase at the three-phase (vapor-aqueous solution-incipient hydrate) equilibrium at representative conditions of temperature, pressure, and mass percent methanol in the feed were computed using the computer program. These are shown in column E in Table 1. As expected, the concentrations of methanol and water in the vapor phase of the above mixtures are very small. It is also seen from Table 1 that the calculated amount of methanol in the vapor phase is so small that its effect on the equilibrium pressure is negligible. Vapor phase samples from equilibrated mixtures of CO2 and a solution containing 10 mass % NaCl and 10 mass % methanol were analyzed on the GC. As was the case with the nonelectrolyte solutions, the concentrations of methanol and water in the vapor phase were negligibly small. In view of the above, the vapor phase analysis was not carried out during the experiments reported in this work. The concentrations of methanol reported refer to those in the feed aqueous solutions. Charging Fresh Solution. Before charging a solution, the cell was rinsed with deionized-and-distilled water to wash out electrolytes from the previous experiment. The cell was then rinsed once with the experimental solution. The piston of the Ruska pump was brought to its 0.0 mL position. The cell and the gas reservoir were repeatedly flushed with CO2 from the cylinder. A sample was analyzed on the GC to ensure the absence of air. About 12 mL of the solution was charged into the cell and the vent valve closed. For each solution generally three experiments were performed at increasing temperatures. Experimental Procedure. The piston of the Ruska pump was brought to its 50% position (125 mL). This resulted in an increase in the pressure in the system. The pressure was further increased to the expected equilibrium value by introducing gas from the cylinder.
Ind. Eng. Chem. Res., Vol. 35, No. 3, 1996 821 Table 1. Change in Computed Equilibrium Pressure as a Result of the Assumption Methanol (MeOH) Concentration in the Aqueous Phase Remains Unchanged (Based on a Maximum Vapor Phase Volume of 545 mL) predicted MeOH predicted mass % in equilibrium conditions predicted mol % estimated actual % error if equilibrium feed aqueous MeOH in MeOH mass % in mass % MeOH pressure (kPa) for % change in solution T/K P/kPa vapor phase aqueous phase reported as B mass % MeOH in F equil. pressure
system A CO2-MeOH-water
B 10 20 20
C 269 263 272
D 1241 998 3022
E 0.036 95 0.052 42 0.052 11
F 9.97 19.97 19.87
(B - F)/B ) G 0.34 0.17 0.63
H 1240 996 2998
(D - H)/D ) I 0.08 0.2 0.8
Table 2. Experimental Carbon Dioxide Hydrate Equilibrium Conditions solution ID
T/K
P/MPa
pure water
275.11 279.49 269.90 269.83 271.31 270.50 271.11 272.36 269.51 271.13 275.37 270.63 273.53 276.23 263.39 265.94 269.20 265.53 268.79 272.36 274.71
1.560 2.620 1.380 1.400 1.610 1.500 1.620 1.870 1.370 1.650 2.820 1.480 2.139 3.038 1.271 1.736 2.707 0.910 1.381 2.103 2.900
Me10-1 Me10-2 Me10-3 Me10-4 Me5Na5 Me5Na15 Me5K10
solution ID Me5Ca10 Me5Ca15 Me10Na10 Me10K10 Me10Ca10 Me15Na5 Me15Ca5
T/K
P/MPa
265.94 271.04 274.19 265.23 267.84 270.74 264.00 267.37 270.83 265.58 268.90 271.85 264.72 267.28 270.93 264.76 267.77 270.83 264.72 267.77 270.76
0.985 1.807 2.809 1.328 1.827 2.715 1.184 1.818 2.872 1.248 1.828 2.767 1.137 1.543 2.552 1.243 1.807 2.730 1.160 1.658 2.456
An estimate of the equilibrium pressure for nonelectrolyte solutions was obtained using the computer program. For solutions containing electrolytes, the pressure was estimated using the data at lower temperatures as described by Dholabhai et al. (1993). The contents of the cell and the reservoir were then allowed to cool. Once the desired temperature was attained, the pressure in the system was increased by introducing additional mercury into the reservoir using the Ruska pump. As soon as hydrates formed, they were decomposed by reducing the system pressure to a value lower than the expected equilibrium pressure. This was accomplished by withdrawing the mercury from the reservoir. This procedure of forming and decomposing the hydrates was repeated one more time. The hydrates were formed for the third time, the pressure was adjusted to the expected equilibrium value, and the system was left to equilibrate. If the temperature and pressure of the system remained constant for 4-6 h with a very small quantity of the hydrates present in the solution, the constant temperature and pressure were taken as the equilibrium conditions. If the system pressure was higher than the equilibrium value, then the pressure would continuously decrease. In such a case the pressure was readjusted to a lower value. On the other hand, if all the hydrates had decomposed, then the hydrates were formed once again as described above, the pressure was adjusted to a value slightly higher than that at which hydrates had decomposed, and the observations continued. Once the equilibrium conditions were established, the system pressure was reduced by about 50 kPa to verify that all the hydrates decomposed at the lower pressure and the experiment was terminated. The temperature of the system was changed, and the procedure was repeated to obtain another equilibrium condition. Experimental Results and Discussion Validation of the Apparatus. To validate the apparatus, equilibrium conditions were obtained for CO2
Figure 3. Experimental equilibrium conditions for CO2 hydrates in pure water: comparison of data obtained in this work with those in literature. The solid curve was obtained from a best fit regression of all the data plotted. The dashed curve represents the predictions obtained using MEGHA.
hydrates in pure water and compared against data already available in the literature. The data obtained using our new apparatus are given in Table 2. These data, along with those available in the literature, are plotted in Figure 3. As seen from this figure, our new data match the literature data very well. The solid curve was obtained from a “best fit” regression of all the data plotted in the linearized form of
ln(P/MPa) ) A/(T/K) + B
(1)
where P is the pressure and T is the temperature. The standard error of estimate of the pressure was 73.64 kPa for all the data and 73.90 kPa for all the data excluding those from this work. The dashed curve in Figure 3 represents the predictions obtained using the computer program MEGHA. There is a very good agreement between the predictions and the experimental data. Solutions with Inhibitors. The compositions of the solutions with inhibitors used in the experiments are shown in Table 3. The experimental data are shown in Table 2. It may be noted that, for each condition in Table 2, the composition of the aqueous phase in equilibrium with the other two phases is shown to be the same as that of the solution fed into the cell. It is well-known that electrolytes do not enter the hydrate phase and that their presence in the vapor phase can be safely neglected. As indicated earlier, the concentrations of methanol and water in the vapor phase were also found to be negligibly small. Therefore, the vapor phase was assumed to be pure CO2. It was also pointed out in the description of the experimental procedure that the quantity of hydrates at equilibrium was negligibly small. Consequently, it is assumed that the amounts of methanol and electrolytes relative to that of water in the aqueous phase at equilibrium are identical to those in the feed shown in Table 3. The experimental data with a 10 mass % methanol solution are also shown in Figure 4 along with those of
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Figure 4. Experimental equilibrium conditions for CO2 hydrates in 10 mass % methanol solutions (Me10-1, Me10-2, Me10-3, and Me10-4). Table 3. Compositions of Aqueous Solutions (All Mass % on Wet Basis) solution ID
methanol mass %
Me10-1 Me10-2 Me10-3 Me10-4 Me5Na5 Me5Na15 Me5K10 Me5Ca10 Me5Ca15 Me10Na10 Me10K10 Me10Ca10 Me15Na5 Me15Ca5
10.06 10.00 10.04 10.00 5.00 5.00 5.00 4.96 5.00 9.96 10.00 9.95 15.14 15.03
NaCl mass %
KCl mass %
Figure 5. Experimental equilibrium conditions for CO2 hydrates in pure water and solutions containing methanol and sodium chloride. Each solid curve is a “visual” fit of the data for a solution and is drawn to clarify the trend for the solution. The compositions of the solutions are given in Table 3.
CaCl2 mass %
5.00 15.00 10.01 10.00 14.97 9.96 10.01 9.93 5.00 5.02
Ng and Robinson (1985). As is seen from this figure, there is some disagreement in the two sets of data, especially at higher temperatures. However, it should be noted that the temperature scale is highly expanded. The maximum disagreement is 1.6 K at 2.8 MPa. In order to check the repeatability and reliability of our data, a total of nine experiments was performed for this methanol concentration. These experiments involved four different solutions (Me10-1, Me10-2, Me10-3, and Me10-4). Experiments with Me10-2 and Me10-3 were performed following the “temperature search” procedure instead of the “pressure search” procedure described earlier. Both the procedures are expected to yield essentially identical results. The temperature search procedure followed was similar to that described by Bishnoi and Dholabhai (1993). The data for the 10 mass % methanol solution obtained in this work were fitted to eq 1. The standard deviation of the linear leastsquares fit was 38.4 kPa, much lower than the standard deviation of 73.64 kPa reported earlier for pure water experiments from a number of different laboratories. The excellent repeatability of these experiments leads us to conclude our data are reliable. The experimental data for the rest of the solutions are plotted in Figures 5-7. Each solid curve in these figures is a “visual” fit of the data for a solution and is drawn to clarify the trend for the solution. They do not represent predictions obtained using any model. Observations Regarding Relative Inhibition Strengths of NaCl and CaCl2. A measure of the inhibiting effect is the hydrate depression which is the difference between the hydrate equilibrium temperatures in pure water and a solution at a given pressure. Predictions for pure water are also shown in Figures 5-7 to illustrate the inhibiting effect of the solu-
Figure 6. Experimental equilibrium conditions for CO2 hydrates in pure water and solutions containing methanol and calcium chloride. Each solid curve is a visual fit of the data for a solution and is drawn to clarify the trend for the solution. The compositions of the solutions are given in Table 3.
Figure 7. Experimental equilibrium conditions for CO2 hydrates in pure water and solutions containing methanol and potassium chloride. Each solid curve is a visual fit of the data for a solution and is drawn to clarify the trend for the solution. The compositions of the solutions are given in Table 3.
tions. Table 4 lists the hydrate depression for some of the solutions used in this work at three different pressures. The constants A and B in eq 1 for each of the solutions in Table 4 were estimated by regressing the experimental data. The equation was then applied to compute the equilibrium temperature at a given pressure. This temperature was subtracted from the corresponding temperature for pure water to obtain the hydrate depression. Dholabhai et al. (1993) reported the values of A and B for CO2 hydrates in 20 mass %
Ind. Eng. Chem. Res., Vol. 35, No. 3, 1996 823 Table 4. Hydrate Depression (in Degrees K) in Solutions Containing NaCl, CaCl2, and Methanol solution ID
P ) 1.5 MPa
P ) 2.0 MPa
P ) 2.5 MPa
Me5Na15 Me5Ca15 Me10Na10 Me10Ca10 Me15Na5 Me15Ca5 Na20a Ca20a
10.0 8.5 8.9 7.8 8.5 7.9 11.96 12.70
10.2 8.7 9.1 8.0 8.7 8.0 12.32 12.80
10.3 8.9 9.2 8.2 8.8 8.0 12.60 12.88
a Taken from Dholabhai et al. (1993). The hydrate depression computations for these two solutions were performed using the values of A and B for eq 1 given in that publication.
NaCl and CaCl2 solutions. Table 4 also lists the hydrate depression values for these solutions computed using the published values of A and B. It has been observed in the past (Dholabhai et al., 1991, 1993; Dholabhai and Bishnoi, 1994) that the inhibiting effects of NaCl and CaCl2 on a mass basis are very close. This is apparent from the depression values for Na20 and Ca20 in Table 4. The difference between the depressions is 0.7 K at 1.5 MPa and 0.3 K at 2.5 MPa. It is also seen that CaCl2 is marginally a stronger inhibitor than NaCl. However, there are two very interesting observations that one could make from the information given in Table 4 when pairs of solutions containing the same mass percent methanol are compared: 1. The hydrate depression in solutions containing methanol and CaCl2 is lower than that in solutions containing methanol and NaCl. This suggests that the inhibiting strength of CaCl2 gets “diluted” in the presence of methanol. 2. The difference between the inhibiting strengths of NaCl and CaCl2 in the presence of methanol is the greatest when the methanol concentration is the lowest. For example, the difference between the depression values in Me5Na15 and Me5Ca15 is 1.5 K, whereas the difference is only 0.59 K between Me15Na5 and Me15Ca5. No attempt is made here to explain these observations. However, one may be able to get an insight into the behavior once modeling of the data is successfully completed. Conclusions A new fixed volume sapphire tube equilibrium cell was designed and constructed. It was converted to a variable volume cell by connecting it to a reservoir, which was in turn connected to a mercury displacement pump. The apparatus was validated with equilibrium data on CO2 hydrates in pure water. Experimental equilibrium data for CO2 hydrates in solutions containing methanol and electrolytes were obtained in the temperature range of 263-279 K and the pressure range of 0.9-3.0 MPa. It was observed that, in the presence of methanol, CaCl2 is a weaker inhibitor on mass basis than NaCl. This is the reverse of an earlier observation that CaCl2 was a marginally stronger inhibitor than NaCl on mass basis in the absence of methanol. It was also observed that the difference in the inhibiting strengths of CaCl2 and NaCl was the greatest at low concentrations of methanol.
Calgary, Alberta, Canada, for the use of the computer program MEGHA. This work was carried out as a part of the “Solids Deposition in Hydrocarbon Systems” Project, Contract No. 5091-260-2138, funded by the Gas Research Institute, Chicago, IL, and the Gas Processors Association, Tulsa, OK. Nomenclature A, B ) coefficients in eq 1 P ) pressure (MPa) T ) tempurature (K)
Literature Cited Adisasmito, S.; Frank, R. J., III; Sloan, E. D., Jr. Hydrates of Carbon Dioxide and Methane Mixtures. J. Chem. Eng. Data 1991, 36, 68. Bishnoi, P. R.; Dholabhai, P. D. Experimental Study on Propane Hydrate Equilibrium Conditions in Aqueous Electrolyte Solutions. Fluid Phase Equilib. 1993, 83, 455-462. Chen, H.-S. The Properties of Carbon Dioxide Hydrate. The Growth Rate of Ice Crystals: The Properties of Carbon Dioxide Hydrate. A Review of Properties of 51 Gas Hydrates. Research and Development Progress Report No. 830; Office of Saline Water, U.S. Department of Interior: Washington, DC, 1972, p 1. Deaton, W. M.; Frost, E. M., Jr. U.S. Bur. Mines Monograph 1946, 8. Dholabhai, P. D.; Bishnoi, P. R. Hydrate Equilibrium Conditions in Aqueous Electrolyte Solutions: Mixtures of Methane and Carbon Dioxide. J. Chem. Eng. Data, 1994, 39, 191-194. Dholabhai, P. D.; Englezos, P.; Kalogerakis, N.; Bishnoi, P. R. Equilibrium Conditions for Methane Hydrate Formation in Aqueous Mixed Electrolyte Solutions. Can. J. Chem. Eng. 1991, 69, 800-805. Dholabhai, P. D.; Kalogerakis, N.; Bishnoi, P. R. Equilibrium Conditions for Carbon Dioxide Formation in Aqueous Electrolyte Solutions. J. Chem. Eng. Data 1993, 38, 650-654. Englezos, P.; Bishnoi, P. R. Prediction of Gas Hydrate Formation Conditions in Aqueous Electrolyte Solutions. AIChE J. 1988, 34, 1718-1721. Englezos, P.; Bishnoi, P. R. Experimental Study on the Equilibrium Ethane Hydrate Formation Conditions in Aqueous Electrolyte Solutions. Ind. Eng. Chem. 1991, 30, 1655-1659. Larson, S. D. Phase Studies of the Two-Component Carbon Dioxide-Water Systems Involving the Carbon Dioxide Hydrate. Ph.D. Thesis, University of Illinois, 1955. Data taken from tables reproduced in Sloan (1990). Ng, H.-J.; Robinson, D. B. Hydrate formation in Systems Containing Methane, Ethane, Propane, Carbon Dioxide or Hydrogen Sulfide in the Presence of Methanol. Fluid Phase Equilib. 1985, 21, 145-155. Parent, J. S. Investigation Into the Nucleation Behaviour of the Clathrate Hydrates of Natural Gas Components. M.Sc. Thesis, Department of Chemical and Petroleum Engineering, The University of Calgary, Calgary, Alberta, Canada, July 1993. Robinson, D. B.; Mehta, B. R. Hydrates in the Propane-Carbon Dioxide-Water System. J. Can. Pet. Technol. 1971, 10, 33-35. Sloan, E. D. Clathrate Hydrates of Natural Gases; Mercel Dekker, Inc.: New York, 1990. Tse, C. Predictions of Gas Hydrate Formation Conditions for Carbon Dioxide in the Presence of Electrolytes. M.Sc. Thesis, Dept. of Chemical and Petroleum Engineering, The University of Calgary, Calgary, Alberta, Canada, 1993. Tse, C. W.; Bishnoi, P. R. Prediction of Carbon Dioxide Gas Hydrate Formation Conditions in Aqueous Electrolyte Solutions. Can. J. Chem. Eng. 1994, 72, 119-123.
Received for review February 27, 1995 Revised manuscript received November 13, 1995 Accepted December 7, 1995X IE950136J
Acknowledgment The authors acknowledge the assistance provided by Amit Majumdar in obtaining the data in the laboratory. The authors thank Canadian Gas Hydrates Group Inc.,
Abstract published in Advance ACS Abstracts, February 1, 1996. X
