ORIGIN OF UNDERGROUND CARBON DIOXIDE
61
4. Within the accuracy of the experiments, no interionic effects were found between potassium ions associated with the clays and those added in true solution as potassium salts. REFERENCES (1) BAVER,L.D.:Missouri Agr. Expt. Sta. Research Bull. 129 (1929). (2) BRADFIELD, R.: J. Am. Chem. SOC.U,2669 (1923). (3) FREUNDLICH, H.,SCHMIDT, O., AND LINDAU,G.: Kolloid-Beihefte 38, 43 (1932). G.S.: Trans. Faraday SOC.S1, 31 (1935). (4) HARTLEY, (5) HENDRICKS, S.B., AND Ross, C. S.: Z. Krist. mineral. Petrog. lWA, 2.51 (1938). (6) HOFFMAN, W.F., AND GORTNER, R. A,: J. Biol. Chem. 66,371 (192.5). C. E.: J. Phys. Chem. 49,1155 (1939). (7) MARSHALL, (8) MARSHALL, C. E., AND BERGMAN, W. E.: J . Am. Chem. 900.89, 1911 (1941). (9) MARSEALL, C. E., AND GUPTA,R. S.: J. Soo. Chem. Ind. 62,433T (1933). (10) NUT TIN^, P. G.: J. Franklin Inst. 224, 339 (1937). (11) PAULI,W.,AND MATULA, J.: Kolloid-Z. 31, 49 (1917). (12) RABINOVITCH, A. F., AND KARGIN, V. A.: Trans. Faraday SOC. S1,M (1936). G.:Kolloid-Z. 61, 49 (1930). (13) WIEGNER,
THE ORIGIN OF UNDERGROUND CARBON DIOXIDE' FRANK E. E . GERMANN
AND
HERBERT W. AYRES
Department of Chemistry, University of Colorado, Boulder, Colorado Received August 18, 1941 I. INTRODUCTION
Many theories of the origin of subterranean carbon dioxide have been advanced (4), and it is a reasonable assumption that there is truth in a number of them. It is the purpose of the present paper to study only one of these, Le., thermal dissociation, and to show that actual underground conditions may lead to quite different results from those obtained under idealized conditions realized in a modern laboratory. The first careful study of the equilibrium established between calcium carbonate, calcium oxide, and carbon dioxide was reported by Debray (2), whose results are incorrect principally because of incorrect values assigned to the temperatures of boiling cadmium and zinc. Since that time numerous studies of the thermal dissociation of pure calcium carbonate have been reported and the dissociation pressures over a wide range of temperatures are accurately known. Carbon dioxide has been encountered in various deep oil wells in the United States and Mexico. Volumes estimated as high as fifty million cubic feet per day have been reported from wells the measured closed-in pressures of which Presented a t the Eighteenth Colloid Symposium, which waa held a t Cornel1 University, Ithaoa, New York, June 1!3-21,1941.
62
FRANK E . E. GERMANN AND HERBERT W. AYRES
were as high as 1000 pounds per square inch. These wells range in depth from 1500 to 3000 feet. According to tables compiled by C. E. Van Orstrand (lo), temperatures a t these depths vary greatly depending on the location. Thus a t Carlsbad, Xew Mexico, a temperature of 24'C. was recorded a t 1500 feet and one of 30°C. a t 3000 feet, while a t the same depths a t Kettleman Hills, California, temperatures of 37' and 52'C., respectively, were recorded. Since the critical temperature of carbon dioxide is 31.loC., it is clear that a t least in some localities the gas could appear as a liquid. Moreover, the pressures encountered strongly suggest that the gas does appear as a liquid, since they range downward from the critical pressure (73 atmospheres or 1073 pounds per square inch). In order to account for such pressures on the basis of the thermal dissociation of calcium carbonate, the studies of Smyth and Adams (8) show that a temperature of 1275'C. would be necessary. Selden (6) dismisses the possibility of thermal decomposition of limestone on the basis of the fact that a t the assumed temperature of 100O'C. of the intruded igneous rocks, the pressure of carbon dioxide would be only 4 atmospheres absolute or about 45 pounds gauge. Moreover, he doubts whether the mass of intruded rock is great enough t o heat an appreciable thickness of limestone. Guilleaume (3) calls attention to the close relation between the existing deposits and evidence of rather recent volcanic activity. Tammann and Seidel (9) suggest that marl strata, which consist of calcium and magnesium carbonates, iron and aluminum silicates, silicic acid, and water, and which are widespread, may well give rise to gas pressures such as are encountered in nature. They pulverized three samples of dolomitic marl taken from shell limestone outcraps near Gottingen, Germany, and placed them in closed containers both with and without admixture of water. After heating for 20 hr. a t 100'C. they observed pressures of 3 cm. of mercury for dry samples, and 5 cm. for wet samples. In this connection it is well to remember that dry pure calcium carbonate shows a pressure of only 0.1 mm. a t 500'C. Robinson, Smith, and Briscoe (5) state that the lowest temperature at which hydrolysis of calcium carbonate occurs with low-pressure steam is 44OoC., so the carbon dioxide pressures reported by Tammann cannot have been due to simple hydrolysis. Since this latter work was rather sketchy, and no evidence was given to establish definitely that the gas liberated was carbon dioxide, we undertook to reinvestigate the problem, using as our natural carbonate oolitic limestone from Utah. Although any natural calcium carbonate would have served the purpose, the uniform quality of this variety as well as its porous nature made it appear particularly desirable. The sample chosen contained, in addition to the calcium carbonate, a trace of alumina, small quantities of magnesium, iron, and manganese, and considerable silica. It was pulverized to 100 mesh and thoroughly mixed. 11. EXPERIMENTAL METHOD
Since it was intended to measure an increase of pressure above a constant vapor pressure of water, the adaptation of some standard method of vapor-
ORIGIN OF UNDERGROUND CARBON DIOXIDE
63
pressure determination was deemed desirable. The static isoteniscope of Smith and Menzies ( 7 ) , which uses mercury as a confining liquid, was adopted and served the purpose admirably. The general arrangement of the apparatus is shown in figure 1. The U-tube, made of 7-mm. Pyrex tubing, was lengthened to permit a greater range of pressures and the bulb was blown to a diameter of about 3 cm. A 2-liter bottle served as a pressure reservoir and w-as connected to an open-tube manometer, a source of air pressure, and an aspirator. The large volume of the bottle made it possible to vary the pressure to within 0.1 mm. by a proper manipulation of the stopcocks. The isoteniscope was immersed in a well-insulated oil bath equipped with a thermoregulator set a t 98°C. and an efficient stirrer. The temperature never
FIG.1 . Diagram of apparatus
varied more than 0.2OC. between readings, but at the time readings were taken the temperature was adjusted to within O.0loC., as recorded on a Beckman thermometer. The pressure of water vapor and carbon dioxide in the bulb was measured by varying the pressure on the outside of the apparatus until the mercury levels in both limbs of the U-tube appeared to be identical. The pressure noted on the manometer plus the barometric pressure then gave the pressure produced by the water vapor and carbon dioxide in the isoteniscope. The error made in observing the levels of mercury in the U-tube did not exceed 0.2 mm. The bulb was filled with a mixture of the pulverized carbonate and distilled water previously boiled to expel any dissolved air. Fifty to 75 g. of limestone were taken for each determination. Sufficient water was present to make a
64
FRANK E. E. GERMANN AND HERBERT W. AYRES
thick suspension of the solid, which upon settling left a thin layer of clear water on its surface. The mixture was then heated to boiling several times to expel any air that might have been trapped with the mixture, and the bulb was sealed to the isoteniscope. It was not essential that all the air be removed from the apparatus, but it was thought well to evacuate so that the space between the confining liquid and the carbonate would contain only water vapor and carbon dioxide. In this way the pressure developed was principally due to the vapor pressure of water. To this end the mercury was drawn up into the bulb on the limb of the U-tube and evacuation was effected by opening the connection to the aspirator until ebullition of the suspension showed that considerable water.vapor was passing out. The mercury was then restored to the U-tube by tipping the apparatus, and the isoteniscope, connected with the manometer as described, was placed in the oil bath. As the vapor pressure of the water rose, increased pressure was maintained in the reservoiir until the pressure within the bulb of the isoteniscope became constant. It was found experimentally that the system reached thermal equilibrium within 15 or 20 min., but generally the initial reading was not made for 45 to 60 min. in order to make certain that the vapor pressure of water had reached a constant value. In some of the early trials this time was considerably longer, and it is possible that the pressure increases noted were slightly less than those actually produced as a result of carbon dioxide evolved, since it was found that the gas was produced most rapidly a t the outset. Readings were taken a t regular intervals varying from 2 to 24 hr. depending on the rapidity with which the internal pressure increased. It was desired to follow the course of the reaction by noting the increase of pressure-presumably due to carbon dioxide-with time. The initial reading was that of the vapor pressure of water plus that of any slight amount of air which had not been washed out of the apparatus by evacuation. It is apparent that the slope of the curve that results from plotting pressure increase against time is of no particular significance, as it depends upon the free space between the reacting substance and the mercury in the U-tube as well as upon the amount of carbonate present. Since the apparatus had to be broken to introduce each new sample, this space was kept only approximately constant. The value sought,-that is, the final pressure of carbon dioxide when equilibrium was attained,-was assumed to be independent of the amount of carbonate present as long as there was an excem. I t was expected that the final pressures obtained due to carbon dioxide evolved would be constant and reproducible. 111. EXPERIMENTAL RESULTS
The pressure increase was plotted against time; typical results obtained are shown in the accompanying figures. Curve A of figure 2 was obtained using chemically pure calcium carbonate; the slight pressure observed was apparently due to the displacement of adsorbed air. To check this a sample of pure talc was run using the same procedure; an identical curve was obtained.
ORIGIN OF UNDERQROUND CARBON DIOXIDE
'
65
In every case an increase in pressure was observed when the natural carbonate was heated with water. The time required for the curves to flatten out was considerably greater than that reported by Tammann, owing possibly to differences in the apparatus and in the limestone used. Curves B, C, D, and E of figure 2 are typical and represent only a small part of those actually obtained. A number of the curves reached a maximum pressure of 57 to 60 mm. after 200 hr. when starting conditions were similar. Figure 3 shows the pressure developed with a much larger sample of limestone. In place of the usual 50 to 75 g., a sample of 200 g. was taken. The bulb w&s placed in such a way that the gas evolved could not escape into the U-tube, the water above the carbonate mixture entirely filling the space between the bulb and the mercury. Under these conditions the gas was evolved in
72
50
3;
48
z
a" 24 Ib 8
lJ.1
I L.1 I I 1 I I
1 I I IT11 I I I I 1 Oo
40
m
IZQ m T I M IN WE5
maom TIME IN W R S
FIQ.2. Carbon dioxide pressures from
FIQ.3. Carbon dioxide pressure from
50- to 75-g. samples
200-g. sample
minute pockets within the carbonate mixture. As is evident from the curve, the pressure was rising linearly with time up to about 40 hr., a t which time it had risen to about 12 cm. of mercury. This curve shows definitely that we are not dealing with a simple hydrolytic decomposition of calcium carbonate, since the maximum pressure is seen to depend on the amount of sample taken. According to the phase rule, the maximum pressure in such a system should be a function of the temperature, but not of the amount of the solid component. The dependence of pressure on the amount of sample taken can be explained if we w u m e that the carbon dioxide does not result from a hydrolytic reaction such as CaCOI
+ 2Hz0= Ca(OH)* + HzCOs
(1)
66
FRANK E . E. GERMANN AND HERBERT W. AYRES
but from a chemical reaction with some impurity such as CaC03
+ Si02 + H20= CaSiOt + H2COs
(2)
+ Al1O3 = 3Ca0.A1203+ 3c02(g)
(3)
or 3CaC03
If such a reaction were to go to completion as a result of the large amount of carbonate, the final pressure would depend on the initial amount of sample taken, since in any sample there would be a large excess of the carbonate associated with only a small amount of the silica or aluminum oxide. The heat of formation A H o for the reaction represented inequation 3, according to the data compiled by Biochowsky and Rossini (I), is -132.5 k.-cal. per mole. Although the reaction does not take place a t ordinary temperatures when the component solids are brought together, the large amount of heat liberated makes the reaction seem reasonable. Data to calculate the change in free energy are not available. When the finely divided, colloidal carbonate is heated to 100°C. with colloidal aluminum oxide in the presence of water there is every reason to believe that the reaction should take place. Synthetic mixtures of h e l y divided pure calcium carbonate and silica have been heated a t 100°C. in the presence of water, but positive results of the liberation of carbon dioxide have as yet not been obtained. I t is believed that we have not obtained mixtures of the proper degree of colloidality and possibly not of the proper components. Figure 4 illustrates a case of the natural carbonate developing a pressure which was of the usual magnitude. When the gas was pumped off and heating continued, the pressure was not reestablished, as shown by the lower curve; this result shows that the mixture had changed. Had we been dealing with pure thermal and hydrolytic dissociation, the original curve should have been reproduced. Although it seemed improbable that the gas liberated was anything other than carbon dioxide, it was nevertheless deemed advisable to establish this fact conclusively. To this end a tube of potassium hydroxide solution was sealed to the apparatus and separated from the isoteniscope by a thin glass membrane. When the desired pressure was attained, this membrane w&s broken by the impact of an iron rod drawn into a magnetic field, as shown in the inset in figure 1. There was an immediate drop in pressure when the potassium hydroxide bulb was opened, as a result of three factors: (1) The potassium hydroxide bulb and the isoteniscope were evacuated separately and no way was provided to equalize the pressures on the two sides of the membrane. ( 2 ) The vapor pressure of the potassium hydroxide solution was less than that of the water-carbonate mixture. As a result of these two factors there would be a pressure change even though there were no gas other than water vapor in the isoteniscope. (3) The total space above the carbonate mixture was approximately tripled
67
ORIGIN OF UNDERGROUND CARBON DIOXIDE
upon opening the potassium hydroxide chamber. The partial pressure of carbon dioxide accordingly fell to about one-third of its former value or a distance A B as shown on the curve (figure 5). This distance is an estimate, since it was not possible t o distinguish this effect from that resulting from the first two factors described above. The gradual drop in pressure resulting from absorption of carbon dioxide is then shown beginning at the point B . Although the distance A B is an estimate and the result, therefore, cannot be considered strictly quantitative, yet the gradual drop in pressure due to absorption of carbon dioxide was accurately measured, and it is significant that this pressure decrease is about one-third of the value of the original pressure due to the displaced gas.
o TIME IN W U R S
35
io
105 YO 175 TIME IN HOURS
20
FIG 4 FIG.5 FIG.4 Carbon dioxide pressures on original sample and on same sample after pumping off gas liberated in first run. FIG.5. Liberation of gas, followed by absorption in potassium hydroxide.
The development of these large pressures of carbon dioxide from natural carbonates at 98%. should clearly demonstrate the fallacy of drawing the conclusion that the high underground pressures of this gas cannot result from a thermal decomposition of carbonates, merely because it is assumed that the temperatures prevailing in the neighborhood of igneous intrusions are too low to produce such pressures. This assumption is based on calculations of the dissociation pressure of pure calcium carbonate, rather than on the impure mixtures which actually occur underground. If we consider the fact that pure calcium carbonate at 98°C. yields a pres\sure of only about 0.01 mm., whereas our impure mixture with water gives pressures up to 12 or 15 cm. at this temperature, it is obvious that these same impure mixtures underground in the presence of water vapor could yield pressures much higher than those which
68
FRANK E. E. GERMANN AND HERBERT W. AYRES
have been previously calculated. Moreover, the pressure of carbon dioxide over magnesium carbonate, even in the pure state, is much greater than it is over calcium carbonate. Since we have large deposits of dolomitic limestone, it is reasonable to assume that it may play a considerable r81e in the develop ment of the observed pressures. The chemist in his laboratory works with idealized systems of pure materials and obtains physical constants which have a validity when properly used. When these constants are used to interpret phenomena in nature, one may easily go astray if he is not on guard to be sure that the conditions in nature exactly parallel those in the laboratory. For years chemists in the analytical laboratory cautiously avoided colloids, so that later the student received a shock when he learned that life itself as well as most industrial processes are concerned chiefly with colloidal phenomena. The experimental results described no doubt will find their answer in the fact that we are not dealing with ideal pure substances, but with natural colloids originating partly from organic l i e such as the shells which make up a large part of natural limestone, the siliceous residues of diatoms, and the numerous colloids laid down by various inorganic processes. It is one of the tasks of the colloid chemist to be ever on the alert to see to it that the phenomena of life and of nature are interpreted in the light of information which applies to colloid systems rather than to idealized laboratory conditions. REFERENCES (1) BICHOWSKY, F. R.,
ROSSINI,F. D.: Thermochemistry of Chemical Substances. Reinhold Publishing Corporation, New York (1936). (2) DEBRAY, H.: Compt. rend. 64, 603 (1887). (3) GUILLEAUYE, CHAS.:Rev. universelle mines 10, 169 (1926). (4) MELLOR, J. W.: A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. YI, p. 1. Longmans, Green and Company, New York (1925). (5) P. L.. SMITH.H. C.. AND BRISCOE.H. V. A,: Proc. Univ. Durham Phil. ~ -ROBINSON. , SOC.7, 161 i1928). ' (6) SELDEN, R. F.: U. S.Bur. Mines Rept. Investigations8153 (1934). (7) SMITH,ALEXANDER, AND MENZIES, A.W. C.: J. Am.Chem. Soc.91, 1420 (1910). (8) SMYTH, F. H., AND ADAMS,L. H . : J. Am. Chem. SOC.46, 1167 (1923). (9) TAMMANN, G., AND SEIDEL, K.:'Z. anorg. allgem. Chem. 206,209 (19321. (10) VANORSTRAND, C. E.: Physics of the Earth, Internal Constitution of the Earth (edited by Beno Gutenberg), Vol. VII, p. 125. McGraw-Hill Book Company, New York (1939).
AND
