I n d . Eng. Chem. Res. 1988, 27, 872-874
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A Calorimetric Study of Naturally Occurring Gas Hydrated Y. Paul Handa Division of Chemistry, National Research Council of Canada, Ottawa, Ontario, Canada K I A OR9
An automated Tian-Calvet heat-flow calorimeter was used to determine compositions, enthalpies of dissociation, and heat capacities of two naturally occurring gas hydrates: a structure I hydrate recovered from the Middle America Trench slope sediment off Guatemala and a structure I1 hydrate recovered from the Green Canyon area of the northern Gulf of Mexico. The properties of the naturally occurring gas hydrates were found to be similar to those of the laboratory-synthesized hydrocarbon hydrates. The clathrate compounds formed by water when the guests are above their normal boiling points at room temperature may be called gas hydrates in the truest sense. Most hydrates exist as one of two structures called structure I, and structure 11. Structure I is generally formed by molecules in the size range 410-550 pm and structure I1 by molecules either smaller than 410 pm or larger than 550 pm in size (Davidson et al., 1987). Each structure contains two kinds of cages in which guest molecules can be accommodated. An excellent review on gas hydrates is given by Davidson (1973). Numerous investigations of three-phase equilibria of the type hydrate-ice-gas and hydrate-aqueous solution-gas were reported during the period mid-1930s to mid-1960s (Davidson, 1973; Hiza et al., 1982)due primarily to a report by Hammerschmidt (1934) that the cause of plugging of natural gas transmission pipelines was the formation of gas hydrates. Vast deposits of natural gas in the form of gas hydrates have been reported (Kvenvolden and McMenamin, 1980; Cox, 1983) to occur under the ocean bed and in the permafrost regions, and several schemes (Katz, 1975; Elliot et al., 1982; McGuire, 1982) have been proposed for recovering the gas from these deposits. This had led to a revival of interest in the study of thermodynamic properties of gas hydrates. However, because of unavailability of the naturally occurring samples, these studies have been limited to the laboratory-synthesizedgas hydrates (Cherskii et al., 1982; Groisman et al., 1985; Handa, 1986a,b). Recently, two naturally occurring gas hydrate samples have been recovered and were made available to us by Dr. W. F. Lawson of Morgant~wnEnergy Technology Center, U.S. Department of Energy, Morgantown, WV. This provided an opportunity to study the thermodynamic properties of the natural samples, to compare the results with those of the laboratory-synthesized samples, and to assess the effect of the sedimentary material on the thermodynamic properties. Compositions and enthalpies of dissociation of gas hydrates have almost always been determined by treating the hydrate phase equilibrium temperature-pressure data in terms of the Clapeyron equation. The results thus derived are generally associated with large uncertainties (Handa, 1986a). On the other hand, direct composition and calorimetric measurements on gas hydrates are difficult to make because these compounds are stable only either at low temperatures or under high gas pressures. Recently, a calorimetric technique was reported which overcomes the problems associated with sample handling and which, from a single loading of the sample, determines the composition, enthalpy of dissociation, and heat capacities of gas hydrates (Handa, 1986a,b). The accuracy of the calorimetric measurements was found to be better than those of the Issued as NRCC No. 28574. 0888-5885/88/2627-0872$01.50/0
results derived from phase equilibrium data. The calorimetric technique is ideally suited for naturally occurring gas hydrates and has been used in this work to measure their compositions and thermal properties.
Experimental Section Two naturally occurring gas hydrate samples were studied. The sample NGHl was recovered during leg 84 of the Deep Sea Drilling Project at site 570 from the Middle America Trench slope sediment off Guatemala. Before use, the sample was ground, at liquid nitrogen temperature, to a fine powder. The sample NGH2 was recovered from the Green Canyon area of the northern Gulf of Mexico by Texas A&M University research personnel. The sample was made up of small, porous particles and was used as such. All samples were stored in liquid nitrogen until further use. The calorimeter unit used was a Tian-Calvet heat-flow type built by S e t ” of Lyon, France. As purchased, the calorimeter was suitable for measurements on samples under ambient pressures only (Handa et al., 1984). With a simple modification, it was adapted for study of samples under pressure as is required in the case of gas hydrates. The details of automating the control and operation of the calorimeter (Handa et al., 1984) and of the study of samples under pressure (Handa, 1986a) are given elsewhere. The calorimeter has been used previously for measuring compositions, enthalpies of dissociation, and heat capacities of hydrates of krypton and of xenon (Handa, 1986a) and of methane, ethane, propane, and isobutane (Handa, 1986b). The accuracy of the measurements was estimated to be *l%(Handa, 1984, 1986a). Sample NGHl was placed in the calorimeter at 78 K, and any liquid nitrogen sticking to it was pumped off. Kvenvolden et al. (1984) analyzed a part of the original hydrate sample and reported it to be almost pure CH, hydrate. Consequently, an appropriate amount of CHI was introduced in the system such that over the temperature range of measurement the gas pressure in the system was always higher than the dissociation pressure of the hydrate. The amount of C H I to be introduced was based on its P-V-T properties, volume of the calorimetric system available to the gas, and temperature dependence of the dissociation pressure of CH, hydrate. The sample was then heated from 80 K to about 285 K at the rate 9 K h-l to determine heat capacities, and from the thermal anomaly observed at 273 K, the amount of ice in the sample was also determined. The calorimeter was then cooled back to 78 K, the CH, gas introduced previously was pumped off, and sample was heated again at the rate 9 K h-l. During this run, two thermal anomalies were observed: one due to the dissociation of the hydrate into ice and gas and the other due to the melting of all of the water in the sample. The amount of gas released by the sample was
Published 1988 by the American Chemical Society
Ind. Eng. Chem. Res., Vol. 27, No. 5, 1988 873 Table I. Hydrocarbon Composition of Gas Enclathrated in DSDP Sample NGHl component methane ethane propane isobutane
comDosition. mol % this work Kvenvolden et al. (1984) 99.93 99.4 0.01 0.2 0.01 0.05
determined by making on-line P-V-T measurements. The two calorimetric runs together with the P-V-T measurements gave the composition, enthalpy of dissociation, and heat capacities of the hydrate. The amount of sedimentary material in the sample was determined at the end of the calorimetric measurements. The composition analysis of the gas released by the sample was conducted by using a Finnigan 4010 gas chromatograph-mass spectrometer for identifying the components and a Perkin-Elmer 3920 gas chromatograph for quantitative analysis. In the case of sample NGH2, the composition of the enclathrated gas was not known prior to calorimetric measurements, and thus a slightly different procedure was adopted. The sample was placed in a precooled calorimeter cell which in turn was placed in a thermally insulated bottle and the sample mass determined. The sample was then placed in the calorimeter and heated from 78 K at the rate 9 K h-' in the absence of any gas and heat capacities up to the dissociation temperature and the enthalpy of dissociation determined. At the end of the calorimetric run, the masses of the liquid water and of the sedimentary material were determined. The gravimetric method is not as accurate as the calorimetric one for determining the composition of the hydrate because of problems with handling the hydrate and condensation of a small amount of moisture onto the weighing bottle. The accuracy of this method is estimated at &5%. The composition analysis of the enclathrated gas was done as described above.
Results and Discussion DSDP Sample NGH1. The composition of the enclathrated gas determined on an air-free basis is given in Table I. Kvenvolden et al. (1984) also analyzed a part of the original hydrate sample; their results are also given in Table I. They also reported a COz content of about 0.4% and trace amounts of C3H8and i-C4Hip Since the enclathrated gas is almost pure CH,, the hydrate sample NGHl was most likely of the structure I type, and for the purpose of the calorimetric study, the enclathrated gas was assumed to be pure CHI. The sample mass was 2.747 g and consisted of 0.021 g of sediment, 0.296 g of enclathrated gas, and 2.430 g of H20 of which 0.467 g was present as ice. The possible sources of ice could be dissociation of some of the hydrate during the various handling procedures the original sample went through since it was recovered in February 1982 (Kvenvolden et al., 1984) and any seawater associated with the recovered hydrate core. From the above analysis, the composition of the hydrate studied in this work was determined to be CH4.5.91H20. This compares well with the composition of CH4.6.00H20determined for the laboratory-synthesized CHI hydrate. In a previous study on laboratory-synthesized hydrates (Handa, 1986a), it was found that when a finely powdered sample was heated from 80 K it released all of its gas below 273 K and thus underwent a one-step dissociation process. However, if the hydrate was in the form of large crystals, then it released a major part of the enclathrated gas below
273 K and the remaining above 273 K and thus underwent a two-step dissociation process. The sample NGHl underwent one-step dissociation to give ice and gas. This is probably because the sample was finely powdered before use. The dissociation took place between 160 and 220 K, which supports the speculation that the hydrate was a structure I type (Handa, 1986b). The enthalpy of dissociation of the hydrate in ice and gas, AH(hig), at 273 K and 1 bar was obtained by using the heat capacity change, ACJhig), of 4 J K-' mol-' for the synthetic CHI hydrate reported previously (Handa, 1986b). For NGH1, iVi(hig) was determined to be 17.50 kJ mol-', which compares well with the value 18.13 f 0.27 kJ mol-' obtained for the synthetic hydrate CH4.6.00H20 (Handa, 1986b). During the calorimetric scan under gas pressure, there is always the possibility that some of the ice in the sample will be converted to hydrate. A number of scans were made in which finely powdered ice was heated from 80 to 285 K at 9 K h-' in the presence of CHI gas with pressures as high as about 2.5 times the dissociation pressure. The maximum conversion of ice or liquid water to hydrate was found to be 3 % by mass. Thus, due to the relatively large ice content of NGH1, the error due to the ice-gas or liquid water-gas reaction can be expected to be of the order of 3%. Consequently, the compositions and enthalpies of dissociation of the synthetic and the naturally occurring CHI hydrates can be taken to be nearly the same. Precise heat capacity measurements could not be made because of the relatively large ice content of the sample. In general, the heat capacity of the sample was found to be very much icelike. Gulf of Mexico Sample NGHB. A detailed study of sample NGHB has been recently reported (Davidson et al., 1986),which showed, from X-ray powder diffraction, the hydrate to be structure 11. This was also reflected in the high propane content, of about 15 mol %, of the enclathrated gas (Davidson et al., 1986). The composition of the enclathrated gas has been reported previously (Davidson et al., 1986) and was found to be very similar to that reported by Brooks et al. (1984) who also analyzed a sample from a similar but not identical source. Methane was the major component of the enclathrated gas, constituting about 60 mol %. The enclathrated gas also contained ethane, isobutane, carbon dioxide, and hydrogen sulfide in the range of about 1-3 mol % and higher hydrocarbons, up to C5, in amounts less than 0.1 mol %. The sample mass was 0.907 g, made up of 0.730 g of HzO, 0.123 g of the enclathrated gas, and 0.054 g of sedimentary material. It was not possible to determine the ice content of the sample. On the basis of the composition of the enclathrated gas reported previously (Davidson et al., 1986), the average molar mass of the gas was calculated to be 24.9 g. This gives 8.2 mol of water per mole of gas. If all the water were present as hydrate, then a composition of M.8.2H@ corresponds approximately to 70% occupancy for both kinds of cages. Such occupancy factors are too low for the hydrate to be stable (Davidson, 1973). It is possible that some of the water originally present as hydrate was converted into ice during the various handling procedures that the sample went through at Morgantown and during sample manipulation at Ottawa, as NMR, dielectric, X-ray powder diffraction, and finally calorimetric studies were conducted on the same sample (Davidson et al., 1986). Another source of ice can be any seawater associated with the recovered hydrate core. If the original sample had full cage occupancy, then the maximum amount of water present as ice in the sample used for calorimetric study was about 20% by mass.
Ind. Eng. Chem. Res. 1988, 27, 874-882
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The enthalpy of dissociation and heat capacity of the hydrate could not be determined precisely because of uncertainty in the composition of the hydrate. Since this sample was used without grinding to a fine powder, it dissociated in the two-step process discussed above. Most of the dissociation occurred between 220 and 260 K. The specific enthalpy of dissociation of hydrate into ice and gas was found to be 27.8 J g-l based on all the water present as hydrate and 33.1 J g-l based on 20% water present as ice. The specific heat of the hydrate in the range 100-210 K was found to be about 10% lower than that of ice if all the water is assumed to be present as hydrate and about 10% higher than that of ice if 20% water is assumed to be present as ice. Conclusion The small sample sizes and relatively large ice content or uncertainty about the ice content made it difficult to make a thorough study of the samples. A sample size of about 10 g of an intact, naturally occurring gas hydrate is highly desirable. Then a payt of the sample can first be dissociated to determine the composition of the enclathrated gas, and the hydrate can then be scanned in the calorimeter under the pressure of gas of the appropriate composition. Moreover, a larger sample size would allow determinations of the effect of sedimentary material on the thermal properties and the effect of grain size of the hydrate on its dissociation characteristics. The present work, however, does indicate that the properties of the naturally occurring hydrates are broadly the same as those of the laboratory-synthesizedhydrates, and likewise it may be possible that the thermodynamic properties of the sediment-consolidated hydrates are similar to those of sediment-consolidated ice.
Registry No. CH3CH3, 74-84-0; H3CCH2CH3,74-98-6; (H3C)*CHCH,, 106-97-8.
Literature Cited Brooks, J. M.; Kennicutt, M. C.; Fay, R. R.; McDonald, T. J. Science (Washington,D.C.) 1984,225, 409. Cherskii, N. V.; Groisman, A. G.; Nikitina, L. M.; Tsarev, V. P. Dokl. Akad. Nauk SSSR 1982,265, 185. Cox, J. L., Ed. Natural Gas Hydrates: Properties, Occurrence and Recouery; Butterworth, Boston, 1983. Davidson, D. W. In Water: A Comprehensive Treatise; Franks, F., Ed.; Plenum: New York, 1973; Vol. 2, Chapter 3. Davidson, D. W.; Desando, M. A.; Gough, S. R.; Handa, Y. P.; Ratcliffe, C. I.; Ripmeester, J. A.; Tse, J. s. J. Zncl. Phenom. 1987,5, 219. Davidson, D. W.; Garg, S. K.; Gough, S. R.; Handa, Y. P.; Ratcliffe, C. I.; Ripmeester, J. A.; Tee, J. S. Geochim. Cosmochim. Acta 1986, 50, 619. Elliot, G. R. B.; Vanderborgh, N. E.; Barraclough, B. L. U.K. Patent Appl. 2 093 503, 1982. Groisman, A. G.; Sawin, A. Z.; Barachov, S. P.; Tsarev, V. P. Izu. Sib. Otd. Akad. Nauk SSSR, Ser. Khim. Nauk 1985,1, 44. Hammerschmidt, E. G. Ind. Eng. Chem. 1934, 26, 851. Handa, Y. P. J . Chem. Thermodyn. 19868, 18, 891. Handa, Y. P. J . Chem. Thermodyn. 1986b, 18,915. Handa, Y. P.; Hawkins, R. E.; Murray, J. J. J . Chem. Thermodyn. 1984, 16, 623. Hiza, M. J.; Kidnay, A. J.; Miller, R. C. Equilibrium Properties of Fluid Mixtures; IFI/Plenum: New York, 1982; Vol. 2. Katz, M. L. U S . Patent 3916993, 1975. Kvenvolden, K. A.; McMenamin, M. A. US'. Geol. Suru. Circ. 1980, 825, 1. Kvenvolden, K. A.; Claypool, G. E.; Threlkeld, C. N.; Sloan, E. D. Org. Geochem. 1984,6, 703. McGuire, P. L. In Alternate Energy Sources;Veziro, T. N., Ed.; Ann Arbor Science: Ann Arbor, MI, 1982; Vol. 4; p 203.
Received for review August 24, 1987 Accepted December 11, 1987
Virial Equations for Light and Heavy Water Philip G. Hill* and R. D. Chris MacMillan Department of Mechanical Engineering, University of British Columbia, Vancouver, British Columbia, Canada V 6 T 1W5
New fundamental equations have been developed for the vapor states of H 2 0 and D20. For H20, the PVT data, supplemented by numerous accurate throttling data (isenthalpic and isothermal), and the Osborne determinations of saturation vapor enthalpy and density permit formulation with statistically significant third and fourth virial coefficients and a representation of all available data up to a temperature of 300 O C . Based on more limited experimental data, and in recognition of the similarity of H 2 0 and D20 states, a corresponding formulation has been developed for D20 which may be used to determine the equilibrium vapor states of D20 up to 300 "C. 1. Introduction
The subcritical temperature vapor states of H 2 0 are defined by abundant experimental data, Besides PVT data (which are difficult to use alone in establishing virial coefficients), there are numerous isenthalpic and isothermal throttling coefficients, some specific heat and speed of sound data, and accurate determinations by Osborne et al. (1937, 1939a,b) of the saturated vapor density and enthalpy. The use of these data in a selective regression analysis which retains only statistically significant terms makes it possible to establish the second, third, and fourth coefficients of a virial expansion in the form
-P-
- 1 + B2p
+ B3p2 + B*p3
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and to represent thereby all experimental data within what is believed to be the accuracy of those data. For DzO, experimental data on the vapor states are few in addition to PUTdata. However, the close similarity of the DzO and HzO vapor states, coupled with the direct experimental determinations of the small differences in the second virial coefficient of DzO and HzO, makes it possible to derive a corresponding virial equation for DzO from the HzO virial equation. As was shown by Hill and MacMillan (1980), such an equation is required (in conjunction with the quite accurately known vapor pressures of DzO) to establish the saturation vapor densities and enthalpies of DzO. The Hill and MacMillan (1980) derivation of a virial equation for H 2 0 provided values within the estimated 0 1988 American Chemical Society
