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The Journal of Physical Chemistry, Vol. 82, No. 2, 1978
(31) See, e.g., L. N. Sidorov and E. N. Kolosov, Russ. J. Phys. Chem., 42, 1382 (1968). (32) E. W. Dewing, Metall. Trans., 3, 2699 (1972). (33) 0.J. Kleppa and M. E. Melnichak, R o c . 4th Conf. Int. Thermodyn. Chim., 3, 148 (1975).
Evans and E. Tschuikow-Roux
(34) P. D. Greene, P. Gross, and C. Hayman, Trans. faraday SOC., 64, 633 (1966). (35) E. W. Dewing, Metall. Trans., 1, 2211 (1970). (36) "JANAF ThermochemicalTables", Dow Chemical Co., Midland, Mich. 1965.
Thermodynamic Properties of Tetrafluorohydrazine and the Difluoroamino Radical P. J. Evans and E. Tschulkow-Roux" Department of Chemistry, University of Calgaty, Calgary, Alberta T2N 1N4, Canada (Received May 18, 1977)
The equilibrium dissociation of tetrafluorohydrazinevia the reaction NZFd(g) * 2NFz(g) has been studied behind incident shock waves over the temperature range 375-1500 K. At temperatures above 500 K, the N2F4is completely dissociated and absorption measurements in this region yielded the absorptivity of NF2 as a function of temperature. Absorption measurements in the temperature range 375-462 K yielded equilibrium constants for the above reaction. On the basis of second and third law analyses of the data, a value of 20.9 f 0.4 kcal mol-' for the heat of dissociation of NzF4 was obtained. In addition, previously reported measurements of the equilibrium constants for the reaction NF3(g) * NF2(g)+ F(g) have been reanalyzed and a value of 57.2 f 1.0 kcal mol-' was obtained for the heat of dissociation of NF3. The above results were then used to obtain revised values for the heats of formation of N2F4 (AHfozg8[NZF4(g)] = -5.3 f 1.4 kcal mol-l) and NF2 (AHfozg8[NFZ(g)] = 7.8 f 1.0 kcal mol-').
1. Introduction In a recent study on the thermal decomposition of nitrogen trifluoride,' we reported measurements of the equilibrium constants of the reaction NFs(g) + NFz(g) + F(g) (1) as a function of temperature. These measurements were in disagreement with the equilibrium constants calculated from the JANAF Thermochemical Tablesa2We attributed the difference to uncertainties associated with the thermodynamic data of the difluoroamino radical.' An examination of the thermodynamic properties of NFz(g) indicates that the proposed value for the heat of formation of this species, A&ozg8[NFz(g)],is the most likely source of this uncertainty. AHfOz98 for the difluoroamino radical has usually been determined2-6from the heat of formation of tetrafluorohydrazine, AHfozg8[NZF4(g)], and its heat of dissociation. The former has been measured by Armstrong et alS6who = -2.0 f 2.5 kcal mol-'. report a value of AHfo298[N2F4(g)] This appears to be the only experimental determination of this quantity. The heat of dissociation of NzF4 has been determined4v7-l0from measurements of the equilibrium constant of the reaction NzF,(g) + 2NFz(g) (11) as a function of temperature. The results of these studies are summarized in Table I together with the values recommended by several reviewer~.~,~J' It can be seen from this table that the measured values of AH are not in good agreement and that this is reflected in the values recommended by the various reviewers. Thus our earlier assignment of the value of mfo2,,[NF2(g)] as the most likely source of error in the thermodynamic data for reaction I appears to be justified. In the present report, reaction I1 is studied as a function of temperature by the shock tube technique. At temperatures above 500 K, the equilibrium expressed by reaction I1 lies far to the right and absorption measurements in this region are used to determine the absorptivity of NF2 as a function of temperature. This information is 0022-3654/76/2082-0182$0 1.0010
TABLE I: Summary of AH Values for Reaction I1 Temp range, AH,^ kcal K mol-' Method Ref 333-451 298 373-423 340-435 243-298 423-523 298.15 298.15 360-425
21.5 ?: 1.6 21.7 19.85 i 0 . 2 19.3 .i 1.0 1 9 . 8 i 1.3 20.5 i 0.15 22.26 2 1 . 0 i 0.9 20.1 i 0.3
Mass spectroscopy Spectrophotometry (dP/dTX EPR ESR
(dP/dT), b c d
4 7 7 8 9 10 2 5 11
The values in this column refer to the listed temperature range. Correction to a common temperature is necessary for them to be directly comparable. Based Analysis of on recalculation of measurements in ref 4. literature. Weighted mean of ref 4, 7, and 8. then combined with lower temperature (e460 K) absorption measurements to yield values of the equilibrium constant, K,, as a function of temperature. The latter are then analyzed by both the second and third law" methods to obtain the heat of dissociation, i w 0 2 9 8 , for reaction 11. In addition, our earlier equilibrium constant results for reaction I' have been recalculated using more extensive NF2 absorptivity measurements and the results of this analysis are presented.
11. Results and Discussion A. Absorptivity of NFz. The absorptivity of NF2 in the temperature range 500-1500 K was measured in a series of shock wave experiments on a 0.25% N2F4-99.75% Ar mixture. A detailed description of the apparatus and method used for performing these experiments has been reported previous1y.l In the above temperature range, the NzF4 is completely dissociated and the absorptivity of NFz, a, was calculated by means of the relation
u = { [NF,]l}-' log (lo/l)
(1)
where Io and I are the incident and transmitted light intensities respectively, [NF,] is the concentration in moles
0 1976 American Chemical Society
Equilibrium Dissociation of Tetrafluorohydrazine 6
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The Journal of Physical Chemistry, Vol. 82,No. 2, 1978 103 i
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TABLE 11: Experimental Equilibrium Constants for the Reaction NF, + NF, t F
Tg9 7 500-
5
c
-al
-
E
. 0
z 4000
300-
1000
500
1500
T2 / K
Flgure 1. Plot of absorptivity a vs. incident shock temperature, TP: (open circles) present work: (solid circles) ref 1; (solid diamond) ref 7.
per liter, and I is the sample path length (cm). The results of the present experiments together with those originally reported in ref 1 are shown in Figure 1 as a plot of a vs. temperature behind the incident shock, Tz. The application of the data in this figure to the present analysis will be discussed in the following sections. Previously,' we discussed in some detail other published values of the absorptivity of NF2. The more extensive measurements of Figure 1 provide additional support for most of that discussion. However, slight revisions to that discussion are necessitated by the results of the present work and these will be discussed in the relevant sections. B. Thermodynamic Properties of NF2. The transmitted light intensities from our previously reported1 equilibrium measurements of reaction I have been recalculated using the more extensive NF2 absorptivity measurements of the preceding section to yield equilibrium constants, Kp' For this purpose, the NF2 absorptivities in the temperature range 1000-1500 K were fitted to a straight line by .the method of least squares. This analysis yielded the equation
a(L mol-' cm-') = 537.1 - 0.1708T
(2)
which is represented in Figure 1 by the broken line. Kp values were then determined from the experimental results by means of eq 2 and the following relations: [NF2Ieq = [FIeq = (ai)-' log ( I o I I e q ) [NF,I,
= P ~ i , e g [ N F a li [NF21,
pA = [ A P T
( A = NF3, NF2, F )
Kp =PNF2PFIPNF,
(3)
(4) (5)
(6)
where p21,eqis the density ratio across the shock front a t equilibrium, pA(atm)is the equilibrium partial pressure of A, K, is the equilibrium constant, subscripts 1 and eq refer to initial (ahead of the shock front) and post-shock equilibrium conditions, respectively, and all concentrations have units of moles/liter. The K , values calculated by the above procedure are listed in Table I1 and a plot of -In Kp vs. l/Teqis presented in Figure 2. The least-squares line through the data points is represented by the equation
In K,
= 18.20 - 2.890 X 104/T,,
(7)
The slope of this line is equal to -AH/R for reaction I and from this a value of AHo, = 57.4 f 1.7 kcal mol-l is obtained. The error associated with this quantity and all subsequent errors are standard deviations. If the above value is assumed to refer to the midpoint of the experi-
1129 1148 1175 1186 1199 1214 1220 1228 1266 1278 1294 1345 1354 1362 1374
-In Kp 7.156 7.158 6.374 6.184 5.876 5.610 5.491 5.347 4.784 4.312 4.373 3.395 2.928 3.177 2.584
-
- N G o T H",,,)/T? cal mol-' K-' 35.914 35.917 35.921 35.922 35.924 35.925 35.926 35.927 35.930 35.931 35.932 35.935 35.936 35.936 35.936
kcal mol-' 56.60 57.56 57.09 57.18 57.07 57.15 57.14 57.17 57.52 56.87 57.74 57.41 56.53 57.54 56.43
a Values of the Gibbs free-energy function for the reaction were calculated from the tabulations of ref 2.
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The Journal of Physical Chemistry, Vol. 82, No. 2, 1978
TABLE IV: Thermodynamic Properties of Gauche N,F,(g)
TABLE 111; Physical Parameters for Gauche and Trans N,F, Gauche Point group Vibrational freauencies. cm-I
Symmetry number, u rNNf A
hF,
LF,NF,, deg LF,NN, deg LF,NN, deg Dihedral angle
CZ 116, 242, 284, 300. 423. 518; 590; 737, 934, 946,1010, 1023 2 1.489 1.375 105.1 101.1 104.3 67.1
Trans
Ref
131, 252, 354 467. 494. 542; 601; 719, 873, 962, 999, 1039 2
20
1.489 1.375 102.9 100.6 100.6 180.0
17 17
1.3660
C
Czh
17 17 17 17
Bond length. Product of the moments of inertia, For the gauche conformer: D. R. Lide, Jr., and D. E, Mann, J. Chem. Phys., 31, 1129 (1959); for the trans conformer: calculated from data in ref 17.
is in excellent agreement with the second law value. On the basis of second and third law analyses of our experimental data for reaction I, we propose a mean value of = 57.2 f 1.0 kcal mol-' for this reaction. The above value differs significantly from m 0 2 9 8 = 60.4 kea1 mol-l calculated from the JANAF' preferred enthalpies of formation for the species involved in reaction I. In keeping with our earlier conclusions,l we attribute most of this difference to uncertainties associated with iWf0298[NF2(g)] and propose a revised value of AHfO298[NF,(g)] = 7.8 f 1.0 kcal mol-'. C. Thermodynamic Properties of N2F4. An examination of the data listed in Table I indicates that the heat of dissociation of N2F4(g)via reaction I1 is an area of uncertainty. The most systematic examination of the literature values of this quantity appears to be that presen.ted in JANAF Thermochemical Tables2 where consideration was given to the first five values of Table I. The JANAF compilers, in accordance with their stated policy of examining literature data for m0298by both the second and third law methods, found that only their recalculation2 (see discussion presented with N2F4 and NF2 tables) of Herron and Dibeler's data4 yielded consistent second and third law values for m0298[II].However, for reasons that will become apparent shortly, this agreement is fortuitous. The heat capacity and entropy values for N2F4 listed in the JANAF Tables were calculated from fundamental vibrational frequencies measured by Durig and Lord.13 The latter authors analyzed their results by assuming that only the gauche isomer of NzF4was present. However, since the work of Durig and Lord,13extensive spectroscopic studies on N2FlP2Ohave established that both the gauche and trans isomers of this molecule are stable. Furthermore, a t room temperatures, the gas consists of an equilibrium mixture of the two forms even though there is still some uncertainty about the exact composition of this mixtUre.14J7,19,21
Fundamental vibrational frequencies for both the gauche and trans forms of N2F4 have been reported15J6*18v20 and the assignments of different investigators are in good agreement. The minor differences that do exist have been discussed.20,22 For the present work, the fundamental frequencies of Durig and MacNameeZ0were used. Values measured from
298 300 400 500 600 7 00 800 900 1000
Go
s"
-(GOT - H",,,)/T
21.096 21.162 24.173 26.217 27.598 28.550 29.223 29.713 30.079
75.245 75.376 81.903 87.531 92.441 96.771 100.630 104.102 107.252
75.245 75.246 76.115 77.848 79.880 81,990 84.084 86.119 88.077
TABLE V: Thermodynamic Properties of Trans N,F,(g)
Cpo T, K 298 300 400 500 600
700 800 900
1000
S"
-(GOT -U"z,,)IT
cal mol-' K-' 20.438 20.510 23.773 25.955 27.415 28.415 29.120 29.632 30.013
73.814 73.940 80.318 85.873 90.743 95.049 98.892 102.353 105.495
73.814 73.814 74.663 76.361 78.361 80.443 82.513 84.530 86,471
TABLE VI: Equilibrium Constants for the Reaction N,F, + 2NFz
375 385 386 388 389 408 414 415 417 421 427 430 432 432 437 458 462
6.665 5.880 5.874 5.644 5.667 4.502 4.291 3.866 3.972 3.871 3.399 3.190 3.074 2.792 2.757 1.621 1.318
42.794 42.785 42.784 42.782 42.781 42.761 42.754 42.753 42.751 42.746 42.739 42.735 42.732 42.732 42.725 42.695 42.689
21.01 20.97 21.02 20.95 21.02 21.10 21.23 20.93 21.12 21.23 21.13 21.10 21.10 20.86 21.06 21.03 20.93
Based on 50% gauche-50% trans mixture of N,F, isomer and JANAF, values for NF,( 8).
infrared spectra of gaseous N2F4 were selected if available and values from gas-phase Raman studies used in all other cases. Using these fundamental frequencies, heat capacities, entropies and Gibbs free-energy functions for the temperature range 298-1000 K have been Calculated for both gauche and trans NzF4 Values of the physical parameters used in these calculation are listed in Table I11 and the thermodynamic functions are presented in Tables IV and V. A comparison of our calculated ( G O T - Ho2,,)/T values with those presented in the JANAF Tables2 for N2F4 indicates that the absolute magnitude of the former are larger by -2-4 cal mol-l K-l. Thus the use of JANAF values for the calculation of A(GoT - H0298)/Tfor reaction I1 will cause the third law values of to be too high. This fact is the basis of our earlier statement describing the agreement between the second and third law analyses obtained by the JANAF compilers as fortuitous.
Equilibrium Dissociation of Tetrafluorohydrazine
The Journal of Physical Chemistry, Vol. 82,No. 2, 1978
A series of experiments on a 0.25% NzF4-99.75% Ar
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mixture were performed for post-shock temperatures in the range 375-462 K. In this temperature range, the NzF4 is not completely dissociated and the NF2 absorption measurements will yield equilibrium concentrations of NF2 provided the absorptivity is known. For the present work, the latter information was obtained in the following manner. The absorptivities in the temperature range 500-800 K (see Figure 1)were fitted to a straight line by the method of least squares. The expression resulting from this analysis is
a(L mol-' cm-') = 638.5 - 0.2671T
(9 1 and this is represented by the solid line in Figure 1. NF2 absorptivities in the temperature range 375-462 K were then obtained by linear extrapolation of eq 9 (dashed line in Figure 1). It may be argued that expressions other than a straight line would provide a better description of the data in Figure 1over the whole temperature range. A least-squares analysis was performed for several higher order polynomials but most of these proved to be very unreliable when extrapolated beyond the region for which data were available. A quadratic expression was the only one of those tested that was stable when extrapolated to lower temperatures. However, the use of this expression in the subsequent analysis did not show the interconsistency obtained with eq 9 and for this reason the latter was preferred and adopted. Values of the equilibrium constant, K,, for reaction I1 were calculated from eq 3, 5, and 9 in conjunction with
K, = PNF221PN,F, where poNZF4 is the initial partial pressure behind the shock front and all other symbols retain their earlier definitions. A plot of the In K p values (refer Table VI) against the inverse of the temperature is shown in Figure 3. From the slope of this plot, we obtained a value, after adjustment to 298 K, of iWozg8[II]= 20.9 f 0.4 kcal mol-'. A third law analysis was also performed on the experimental Kp values. Since the Gibbs free-energy functions for the gauche and trans forms of N2F4are slightly different and the entropy of mixing has to be taken into account, the relative proportions of the two isomers present in the gas phase will effect the value of AHoZg8[II]determined by the third law method. Two groups of investigators have concluded that the relative amounts of the two isomers are approximately equa1.14J7 However, Gilbert et al.I9 concluded that the composition of the mixture was 70% trans-30% gauche. Several studies14~15~18~zo~z1 have examined the variation of the gauche/trans ratio with temperature. Only in the work of Selig and HollowayZ1where the sample was heated to 523 K was a change in the relative intensities of the gauche and trans forms observed. In this case, the trend was toward the gauche isomer but the change was quite small. In the absence of a reliable value for the gauche/trans ratio, we have assumed that N2F4 consists of a mixture of 50% gauche + 50% trans. Furthermore, the gauche component of this mixture is assumed to consist of equal amounts of d, 1 optical rotamers.14 In calculating the entropy values for this mixture from the data in Tables IV and V, it is necessary to include a term 1.5R In 2 to account for the entropy mixing. In Table VI, A ( G o-~Hozg8)/T values for reaction I1 are listed for each experimental temperature and these were
4 g
1
4.0
1
d
3.0
20
i lo I
I
I
I
22
I
1
24 103 KIT
I
I
J
26
Figure 3. Temperature dependence of the measured equilibrium constant K, for the reaction N2F4+ 2NF2.
based on the 50% gauchedo% trans mixture. Using these values, a third law analysis yielded AH0298[II]= 21.05 f 0.10 kcal mol-I and this is in very good agreement with the second law value. However, in view of the uncertainties in the relative proportions of the gauche and trans forms of N2F4we give slight preference to the second law value. The Gibbs free-energy changes, AGoz98[II],obtained from the above second and third law values for AHO~~~[II] and the entropy for the 5050 mixture are 8.13 f 0.4 and 8.28 f 0.10 kcal mol-', respectively. It is instructive to consider the effect of the value of AGOzg8 on the room temperature extinction coefficient measurements. In Figure 1,the absorptivity of NF2 a t 298 K as measured by Johnson and Colburn7has been plotted. Their value of 565 L mol-' cm-' is based on A G O z g 8 = 8.26 kcal mol-I (Le., K , = 8.8 X atm). The recalculation of Johnson and Colburn's value using = 8.23 kcal molt1 results in a = 549 L mol-' cm-'. Similarly, if we assume that AGOzg8 = 8.29 kcal mol-', a value of a = 577 L mol-I cm-l is obtained. These examples illustrate the sensitive dependence of room temperature absorptivities on the equilibrium constant. A further illustration of this point is obtained from the work of Makeev et alez3who used Pankratov's5 equilibrium data (AGoB8[II]= 7.45 kcal mol-'). These authors determined the absorptivity to be 241 f 12 L mol-' cm-'. This value is more than a factor of 2 lower than that indicated by the extrapolation of our high temperature results which are independent of the equilibrium constant. In view of the good agreement between the extrapolation of our data and the value of Johnson and Colburn, we propose AGOzg8 = 8.25 f 0.20 kcal mol-' for reaction 11. A comparison of our values of Aff0298[II] with those listed in Table I indicates that the present results are in reasonable agreement with those of ref 4, 7, and 10. However, the AH values reported by von Ellenrieder et d.O ' do not appear to have been adjusted to 298 K. When this correction is made, their result is in excellent agreement with the present study. Finally, by using the values of AH0298[II]= 20.9 kcal mol-' and A"fozg8[NFz(g)]= 7.8 kcal mol-' determined in
186
The Journal of Physical Chemistry, Vol. 82, No. 2, 1978
this study, a revised value of AHfo29s[NzF4(g)]= -5.3 f 1.4 kcal mol-’ is obtained. This value differs significantly from that obtained by Armstrong et aL6 and no satisfactory explanation of the disagreement is presently available.
Acknowledgment. The support of the National Research Council of Canada is gratefully acknowledged. The authors thank Professor C. B. Colburn for a very helpful communication, Dr. H. Wieser for several informative discussions, and Dr. J. Duckett for calculating the moment of inertia of N2F4. References and Notes (1) P. J. Evans and E. TschuikowRoux, J. Chem. Phys., 65,4202 (1976). (2) D. R. Stull and H. Prophet, Natl. Stand. Ref. Data Ser., Natl. Bur. Stand., No. 37 (1971). (3) A. Kennedy and C. B. Colburn, J . Chem. Phys., 35, 1892 (1961). (4) J. T. Herron and V. H. Dibeler, J . Res. Natl. Bur. Stand., Sect. A , 85, 405 (1961); J . Chem. Phys., 35, 747 (1961). (5) A. V. Pankratov, Russ. J. Phys. Chem., 43, 214 (1969). (6) G. T. Armstrong, S. Marantz, and C. F. Coyle, Natl. Bur. Stand. Rep., No. 6584 (1959). (7) F. A. Johnson and C. B. Colburn, J. Am. Chem. SOC.,83, 3043 (196 1).
L. H. Blanco C. and N. 0. Smith (6) L. H. Piette, F. A. Johnson, K. A. Booman, and C. B. Colburn, J. Chem. Phys., 35, 1481 (1961). (9) H. E. Doorenbos and B. R. Loy, J . Chem. Phys., 39, 2393 (1963). (10) G. von Ellenrieder, E. Castellano, and H. J. Schumacher, Z. Phys. Chem., 55, 144 (1967). (11) D. A. Armstrong and J. L. Holmes, Compr. Chem. Kinet., 4, 143 (1972). (12) G. N. Lewis and M. Randall, “Thermodynamics”, McGraw-Hill, New York, N.Y., 1961, p 177. (13) J. R. Durig and R. C. Lord, Spectrochim. Acta, 19, 1877 (1963). (14) C. 8. Colburn, F. A. Johnson, and C. Haney, J . Chem. Phys., 43, 4526 (1965). (15) J. R. Durig and J. W. Clark, J . Chem. Phys., 48, 3216 (1968). (16) D. F. Koster and F. A. Miller, Spectrochim. Acta, ParfA, 24, 1487 (1968). (17) M. J. Cardillo and S. H. Bauer. Inoru. Chem.. 8. 2086 (1969). (18) A. Oskam, R. Elst, and J. C. Dulnker, Spectrochlm. Acta, b a r t i , 28, 2021 (1970). (19) M. M.Gilbert, G. Gundersen, and K. Hedberg, J . Chem. Phys., 58, 1691 (1972). (20) J. R. Durig and R. W. MacNamee, J . Raman Spectrosc., 2, 635 (1974). (21) H. Selig and J. H. Holloway, J. Inorg. Nucl. Chem., 33, 3169 (1971). (22) J. R. Durig, B. M. Gimarc, and J. D. Odom in “Vibrational Spectra and Structure”, Vol. 2, J. R. Durig, Ed., Marcel Dekker, New York, N.Y.. 1975. D 35. (23) G. N.’ Makee;, V. F. Sinyanskii, and B. M. Smirnov, b k l . Akad. Nauk USSR, 222, 452 (1975).
The High Pressure Solubility of Methane in Aqueous Calcium Chloride and Aqueous Tetraethylammonium Bromide. Partial Molar Properties of Dissolved Methane and Nitrogen in Relation to Water Structure‘ LUISH. Blanco C. and Norman 0. Smlth” Department of Chemistry, Fordham University, New York, New York 10458 (Received September 6, 1977)
The solubility of methane in 1 m CaClz at 25-125 “ C , and in 1 m Et4NBr at 25-71 “C, from 100 to 600 atm, has been measured using a direct sampling technique. Methane is salted out by the CaC12 but salted in by the Et4NBr. Salting coefficients are presented. The isobaric Henry’s law applies throughout most of the range of measurement. A more detailed study, confined to a pressure of 200 atm, was made of the methane solubility in 1m Et4NBr over the same temperature range. Henry’s law constants, partial molar volumes, and entropies and heat capacities of solution were derived from the solubility data and interpreted in terms of the structure-making qualities of methane. Similar data are presented for nitrogen. Hepler’s suggested criteria of structure are examined and applied to the present results. Methane appears making and breaking based on (acp0/aP)~ to be a better structure maker in the presence of dissolved NaCl than in water alone, and a poorer structure maker in the presence of dissolved CaC12or Et4NBr.
Introduction This study is a continuation of work already reported2i3 on measurements of the high pressure solubility of gases in water and aqueous salt solutions. Attempts were made there to interpret the thermodynamic data derived from the solubilities in terms of structural considerations. This paper extends the work on methane in water and aqueous sodium chloride2 to include calcium chloride and tetraethylammonium bromide, where interest is attached to (1) methane dissolved in natural brine (of which sodium and calcium chlorides are major components) and ( 2 ) the salting-in property of tetraalkylammonium halides. It is pertinent to inquire whether, in these solvents, Henry’s law holds for methane, whether its partial molar volume depends on the total pressure, and whether there is any relation between structure in solution and the temperature dependences of partial molar volume and solubility. With data now available for comparison from this and previous 0022-365417812082-0 186501.OO/O
studies some answers to these questions can be attempted. Earlier quantitative solubility measurements of methane in aqueous calcium chloride are surprisingly scarce. Michels, Gerver, and Bij14 reported values in 2.7 N CaC12, but only a t 25 “C. A preliminary study in this laboratory5 extended to only 60 atm, and was also limited to room temperature. No previous solubilities of methane in tetraethylammonium bromide have been published except for a 1-atm study at 5-35 “C by Wen and Hung.6 Tiepel and Gubbins7 have measured the partial molar volumes of several gases, including methane, in water and various electrolyte solutions, but their data are confined to 25 “C and do not include the systems here reported.
Experimental Section Solubilities were determined in a stirred, thermostated 1-gal stainless steel autoclave, samples of liquid phases being withdrawn into a thermostated buret system for 0 1978 American Chemical Society
