INDUSTRIAL AND ENGINEERING CHEMISTRY
662
All salt determinations were made in duplicate, and for those salts which exhibited only a slight solubility in furfural, the difference in the two varied nearly 100 per cent a t times. This seems to be rather large, but when converted to the percentage of salt dissolved in the furfural, it is quite small. Generally duplicates checked within 25 per cent of each other. The mutual solubilities of petroleum ether (boiling at 63-69' C.) and furfural were determined by sealing known quantities of the two in test tubes and then noting the temperature a t which complete miscibility occurred. The results obtained are as follows: Wt. 7, Petroleum ether
Temp. of Complete Miscibility, C.
96.5 18.7 9.9
48 55 63
I n general, the results of this study indicate that inorganic salts are quite insoluble in furfural a t ordinary temperatures. Exceptions are the hydrates of barium hydroxide, calcium nitrate, and ferric chloride. Anhydrous zinc chloride was
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found to be very soluble. Sulfates were insoluble, but corresponding chlorides exhibited slight solubility, with the exception noted above. A number of the acids listed in Table I exhibit a marked temperature solubility coefficient which suggests a possible method of purification of those acids by recrystallization from furfural. Such a procedure would have particular value in the separation of two acids, one of which possessed a large solubility coefficient Jvhile the other did not. A further suggestion is the possibility of using furfural as an analytical reagent to differentiate between aromatic and saturated aliphatic hydrocarbons. This follows when i t is remembered that saturated aliphatic hydrocarbons exhibit only a limited solubility in furfural while the aromatic hydrocarbons are soluble in all proportions.
Literature Cited (1) Eichwald, E., U. S. Patent 1,550,523(Aug. 18, 1926). (2) Groth, B.,Svensk Kern Tid., 43,23 (1931). (3) Kaiser, H. E., and Hancock, R. S., U. S. Patent 1,715,085 (May 28, 1929).
NATURAL GAS HYDRATES WILLARD I. WILCOXl, D. B. CARSON, m D D. L. KATZ University of Michigan, Ann Arbor, Mich.
N
ATURAL gases under pressure unite with water to form crystalline hydrates at temperatures considerably above 32 O F. The accumulation of liquid water in a natural gas pipe line and favorable conditions of pressure and temperature may cause the line t o become plugged with a solid phase. Early work on hydrates of gaseous substances carried on during the period of theoretical studies of the fluid substances by Villard (12, IS, IQ), de Forcrand (Q), and others was reviewed by Schroeder (11). Hammerschmidt (6, 6,7 ) called the attention of the natural gas industry to the early work on hydrates and reported on the problems of hydrate formation in natural gas transmission systems. Deaton and Frost ( 1 , 2, 3 ) studied the pipe line problems and reported laboratory data on the conditions a t which hydrates will form for several gases. Recent papers (2,7 ) are concerned more with remedial measures for preventing hydrate formation than with the phase relations of natural gases, water, and hydrates. The usual method of preventing hydrate formation is the dehydration of the gas. This paper reports experimental determinations of equilibrium temperatures and pressures of hydrate formation up t o 4000 pounds per square inch for three natural gases. The equilibrium between a propane-rich fluid, a water-rich fluid, and crystalline hydrate was measured. A discussion of the phase relations and vapor-solid equilibria is included.
Phase Relations A conventional phase diagram, Figure 1 (11), may be used t o represent the behavior of a substance such as carbon dioxide which forms a hydrate (3,8) and is capable of condensation t o a liquid phase at the temperatures under consideration. E C F represents the equilibrium between liquid and vapor 1 Died
Beptember 11. 1939
carbon dioxide a t a slightly higher pressure than for the pure substance due to the presence of water. Above CF are the condensed or liquid phases of water and carbon dioxide, and below GBCF are liquid water (containing dissolved carbon dioxide) and vapor (carbon dioxide with a small amount of water). This vapor and liquid when cooled below C B change to a crystalline hydrate of carbon dioxide with 6 molecules of water with liquid water remaining if there is excess water. Further cooling below BH, which deviates from 32" F. to the extent that the dissolved carbon dioxide and pressure lowers the freezing point, causes the liquid water to change t o ice. Curve AB, which may exist, is the boundary between the two solid phases and ice, with a vapor phase composed predominately of carbon dioxide. DC is the equilibrium between the solid hydrate and liquid (mostly carbon dioxide), and i t approximates a vertical line. Points C and B are quadruple points; the system is invariant with four phases present.
Pure Hydrocarbons I n the early work (12, IS) Villard reported the data of Figure 2 on methane hydrates, giving the critical temperature above which methane hydrate would not form as 70.7" F. Similar critical temperatures were given for ethylene, acetylene, and ethane hydrates without mention of the fact that methane and ethane were in different positions with regard to their liquid-vapor critical temperatures. If the gaseous substance were below its vapor-liquid critical temperature along B C of Figure 1, a quadruple point C would break the continuity of the line B C . The effect of pressure on equilibrium along CC (Figure 1) would be expected to be negligible, and hence the determination of hydrate critical temperatures for those substances below their vapor-liquid critical temperature became essentially the determination of t h r temperature of quadruple point C.
INDUSTRIAL AND ENGINEERING CHEMISTRY
May, 1941
Figure 2 shows the data of Villard (18, IS) and of Deaton and Frost (S). The good agreement on methane and the differences in the ethane values indicate that Villard's ethane probably was impure. No data on the hydrocarbons were available to show the position of the equilibrium between the compressed liquid hydrocarbon and the hydrate, CD,of Figure 1. Deaton and Frost (S) stated that the hydrate would form from compressed liquids of ethane and propane.
HYDRATE MELTING POINTS TABLE I. PROPANE Pressure, Lb./Sq. In. Abs.
Temp., O F.
Pressure Lb./Sq. In. kbs.
Temp., F
117.0 188.0 255.0 295.0
42.3 41.7 41.7 42.1
421.0 616.0 887.0
42.8 42.1 42.8
D
t
663
Natural gases under pressure have been known to form crystalline hydrates with water. This paper gives data on the temperatures and pressures, up to 4 0 0 pounds per square inch, at which three natural gases form hydrates. The phase data on hydrates of pure hydrocarbons were1,)'8nlarged by the equilibrium tempera*res and pressures for the hydrate formation from compressed liquid propane. The concept was introduced of handling vapor-solid equilibria in the same manner as the vaporliquid equilibria data have been used to predict the phase relations of fluid hydrocarbon systems.
S - H20 S-HYDRATE
Natural Gas Hydrates
/
Several natural gases were examined for the temperature a t which hydrates would form a t pressures up to 1400 pounds per square inch (3, 6). These pressure-temperature curves of the natural gas hydrates parallel the curve for pure methane hydrate and have higher equilibrium temperatures at any pressure, somewhat in proportion to the amount of
/
' E
/
mPONFNTS: H 0 (EXCESS) C~~RBON DIOXIDE
d DEATON AND FRosTB)/ AAUTHORS I Id X
FIGURE I. PHASE HYDRA
rURE --c >IAGRAM ILLUSTRATING : FORMATION
An apparatus similar to that described by Deaton and Frost ( 1 ) was used to study propane in the condensed region. The best procedure to form hydrates was to cool the two compressed fluids to about 25 " F. and gradually warm through 32' F. until the hydrate began to melt. This procedure did not always form hydrates, but successful runs were made. The results (Table I) were corrected for a measured lag in the temperature of the cell below the temperature of the bath. On one occasion the cell was filled with finely crushed ice, and propane was injected in a cold room at 15" F. When gradually warmed to 32" F. and above, this mixture changed definitely in appearance and had a melting point of 42" E'. When plotted on Figure 2, these data indicate that the equilibrium between compressed liquid propane and the hydrate was little affected by pressure. Data on the formation of hydrates of pure n-butane are not available, although n-butane does affect the equilibrium pressure of hydrate formation for a gaseous mixture (S). The procedure of cooling liquefied butane and water to 25" F., and gradually heating to 32' F. and above was followed several times. On only one occasion was there evidence of hydrate formation. The crystals were similar in appearance to the propane hydrate and had a melting point of 33.8' F. after correction for the measured lag in the cell temperature. This point on Figure 2 indicates the probable position of the *butane hydrate curve.
+/--
+%--
4!l!g.
IO
FIGURE 2 . HYDROCARBON
HYDRATES
constituents (ethane, propane, and butane) present. Three natural gases were examined, and the pressure-temperature curves were determined to 4000 pounds per square inch.
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INDUSTRIAL AND ENGINEERING CHEMISTRY
664
The apparatus shown in Figuie 3 was used t o measure the pressures and temperatures at which the hydrate would form or melt. The cell, J , had a glass window, A , and was surrounded by a bath, the temperature of which could be changed from 32' F. to any desired point. D was the cooling coil. Pressures were measured by a calibrated Bourdon tube gage, and temperatures were obtained by the bath thermometer, B , and by the thermocouple introduced into the working space of the hydrate cell. The precautions and methods of manipulation were similar to those of revious investigators (1, 6). The procedure was to fill the cef; with natural gas to the desired pressure and inject a small amount of water from chamber C. The cell mas cooled until the crystals of hydrate began to form, and then the bath temperature was gradually raised, usually 2-6" F., until the frost or crystals started to melt. The temperatures given by the bath thermometer and by the thermocouple were recorded for the constant pressure. The melting or freezing of the hydrate was easily reversible, once seed crystals were pyesent.
FIGURE 3. APPARATUS FOR
OF
HYDRATE
PRESSURE GAGE
0BS ERVAT1 0N FORMATION THERMOCOURE LEADS
U
TABLE Ir. ANALYSESOF NATURAL GASES 7
0 0064
Methane Ethane Propane Isobutane n-Butane Pentanes+
0.8641 0,0647 0.0357 0,0099 0.0114 0 0078
coz
a
Mole Fractions Gas C 0 0043 0,0051 0,9320 0,0428 0.0161
..
Gas D 0 . SSiw 0.0682 0.0254 0.0038 0.0089 0.0101
Includes nitrogen.
Vapor-Solid Equilibria A pure gaseous constituent and water have two components and three phases a t the formation of a hydrate and hence have only one degree of freedom. Methane with water in excess could then be completely changed to hydrate at the pressures corresponding to the temperatures given by the curve of Figure 2 . A two-component gas with water has two degrees of freedom if the hydrates are in a solid solution or single-phase state. Likewise a three- (or more) component gaseous system and water would have the same degree of freedom for the equilibrium between vapor and solid with liquid water in excess as the multicomponent vapor-liquid systems have for equilibrium between vapor and liquid. Hydrate formed a t 35' F. and 1000 pounds per square inch from gas C, when separated from the remaining gas by mercury displacement, decomposed from about 41" F. and 450 pounds to 49" F. and 200 pounds per square inch. Gas given off during decomposition had a gradually increasing density. These data indicate that natural gas hydrates are likely to be solid solutions. This analogy between vapor-liquid and vaporsolid equilibria should assist the prediction of hydrate formation and their compositions. ASD PRESSURES FOR FORMA'IYON OF TABLE 111. TEMPERATURES NATURAL G A S HYDRATES
MERCURY PUMP Temp., O
The analyses of the three gases for which the pressuretemperature curves of hydrate formation were measured are given in Table 11. The pressures and the lower temperatures of the two readings at which a very small quantity of the hydrate melted are recorded in Table 111. This difference in temperatures of the bath and cell averaged 0.3" F. and never exceeded 1.0" F. The data are plotted on Figure 4 along with the curve for methane and two natural gases previously reported ( 3 ) . The curves for the three gases a t pressures under 1000 pounds per square inch are much as would be expected from the previous reports on natural gas hydrates. The change in the slope of the curves a t about 66" F. is significant in that the equilibrium temperature a t 4000 pounds is lower by 5-10' F. than the value obtained by extrapolating the lowpressure data as a straight line. The change in the slope is indicative of some change in the hydrates formed, such as lowering the number of water molecules per mole of gas. The determination of the number of molecules of water per molecule of natural gas constituent is not known precisely. Early work (4) indicated from 6 to 8 moles of water per mole of gaseous constituent in the crystals. Preliminary trials on the decomposition of hydrates in the cell after displacing the gas with mercury gave values of 4.3, 4.8, and 5.3 moles of water per mole of gas obtained during decomposition of the hydrates.
Gas B
Compound Nitrogen
F.
42.2 49.5 60.0 66.2 69.6 71.5 73.3 75.0 77.3
Gas B _~ Pressure, it./sq.
in. abs. 182 279 59s 1010 1404 1779 2443 3000 3963
Gas C Pressure, lb./sq. O F. in. abs. 1st Series 40.2 232 81.3 492 60.8 1010 66.0 1523 68.3 2049
Temp.,
2nd Series 42.7 279 47.6 384 56.4 699 62.8 1180 64.5 1375 68.3 1999 71.3 2939 70.0 2489 24.3 3989 12.3 3335
Gas D I'PBSSU~B, lb./sq. O F. in. abs. 1st Serios 38.5 175 47.2 317 53.2 510 60.0 790 64.5 1190 68.1 1740 2nd Series 60.8 895 66.0 1350 69.4 2040 71.3 2640 74.7 3850 73.0 3270
Temp.,
A consideration of the initial solidification data by Deaton and Frost for binary gases and water indicated that vaporsolid equilibrium relations might be estimated. The pressures for the formation of hydrate a t 35" F. for pure gases were plotted at K v - s (mole fraction in vapor/mole fraction in solid) equal to unity (on Figure 5), similar to ideal equilibrium relations of vapor and liquids (9). A 45' line was drawn through the methane point, the ideal vapor-liquid relation. Using these values for methane K y - s , the K V F s was computed for ethane and propane from the binary hydrocarbon initial solidification points using methods in coinmon use for vapor-liquid equilibria ( 9 ) . It was found that
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INDUSTRIAL AND ENGINEERING CHEMISTRY
May, 1941
the ethane and propane K v - s values were symmetrical but different in slope from the 45" methane curve. Accordingly the methane was changed to have the same approximate slope as the ethane and propane curves with the recalculation of the ethane and propane values as shown on Figure 5. The ethane K v - s values reverse with increased pressure similar t o the behavior of the less volatile constituents in a vapor-liquid system (IO). Villard showed in 1896 (1.4) that solid substances such as iodine and camphor vaporized with increased gas pressure, but the pressure for such a reversal in vaporization tendency might be expected to be at a higher value. The absolute values of the K v - s equilibrium constants {vapor-solid) in Figure 5 may be completely incorrect, but in the absence of compositions of hydrates in equilibrium with gases, they will permit a discussion of the value of accurate equilibrium values. These constants do compute the initial solidification point of gas C as 173 pounds per square inch as compared to the value from the curve of 163 pounds a t 35" F. The uncertainty of the butane curve and of the butane content of the gases for which the pressures of hydrate formation have been reported prevent a check on the proposed curves. It was indicated that gases containing nitrogen may be handled with the same procedure as a gas completely made up of hydrocarbons by using the nitrogen K as infinity.
a
TEMPERATURE
-
sa.
IN. ABS:--
"E
The concept of computing the hydrate point, the percentage solidified, and the solidification point for a natural gaswater system in the same manner as computations are made of the dew point, the percentage vaporized, and the bubble point of a fluid system (9) would simplify the handling of hydrate prediction. Such computations when using the K v - s of Figure 5 show that the hydrate in equilibrium with gas C at 173 pounds and 35' F. is 17 mole per cent methane hydrate, 20 mole per cent ethane hydrate, and 63 mole per cent propane hydrate. At 250 pounds and 35" F., 3.6 mole per cent of the gas may be computed to have solidified, with the solid and vapor having the following mole per cent composition : Solid 0.348 0.335 0.317
PRESSURE LBS?PER
APPROXIMATE NATURAL GAS HYDRATE EQUILIBRIA AT 35OE
The computations of the equilibrium solid show a changing composition with increased pressure. If the solid were all formed a t the equilibrium conditions, this relation would probably hold. However, if solid is formed under changing conditions, its ability to reach a new uniform equilibrium composition is not comparable to a liquid phase and may be practically impossible to obtain by rearrangement in the solid state. If complete and reliable vapor-solid equilibrium data were available, the entire phase diagram could be estimated. It is possible that border curves for hydrates of gaseous mixtures will bear a relation to the curves for the pure constituents similar to that given by vapor-liquid border curves as compared to vapor pressure curves for the pure substances. I n that case the critical temperatures and pressures for hydrates of gaseous mixtures might be expected to exceed the values for the more volatile pure constituent, and the natural gas curves reported would be in correct positions with regard to the critical point for the methane hydrate.
FIGURE 4. CONDITIONS FOR FORMATION OF NATURAL GAS HYDRATES
Methane Ethane Propane
I"
FIG 5 .
Vapor 0.981 0.014 0.005
Acknowledgment The natural gases used were kindly supplied by the Phillips Petroleum Company and the Humble Oil and Refining Company.
Literature Cited Deaton, W. M., and Frost, E. M., A m . Gas Assoc.. Natural Gas Dept., Proc., 1937,23; A m . Gas Assoc. Monthly, 19, 219 (1937). Deaton, W. M., and Frost, E. M., A m . Gas Assoc., Natural Gas Dept., Proc., 1938, 112. Deaton, W. M., and Frost, E. M . , Am. Gas Assoc. reprint (1940). Fororand, M. R. de, Compt. rend., 135, 959 (1902). Hammerschmidt, E. G., A m . Gas Assoc. Monthly, 18, 278 (1936). Hammerschmidt, E. G., IND. ENQ.CHEM., 26, 851 (1934). Hammerschmidt, E G . , Oil Gas J.,39, No. 2, 61 (1940). International Critical Tables, Vol. VII, p. 224 (1930). Katz, D . L., and Brown, G. G., IND. ENG. CHEM.,25, 1373 (1933). Kats, D. L., and Kurata, F., Ibid., 32, 817 (1940). Schroeder, W., in Ahrens' "Sammlung chemischer and chemischtechnischer Vortraae". Vol. XXIX. DU. 1-98 (1927) (12) Villard, M., Compt. r e n d , 106, 1602 (1888). (13) Ibid., 107, 395 (1888). (14) Villard, M.. J . phys., 5 , 453 (1896).
