SOLUBILITY O F HYDROCARBOXS 1 3 NONPOLAR LIQUIDS
6G5
T H E SOLUBILITY OF SOLID ETHANE, ETHYLESE, AND PROPYLEXE I X LIQUID XITROGEN A S D OXYGES ANNA LUCILE COX1 A N D THOMAS DE VRIES
Department of Chemistry, Putdue University, Lafayette, Indiana Received August 1 1 , 19.49
Hildebrand ( 5 ) has contributed to the understanding of intermolecular forces in the liquid state by developing an equation which he has successfully applied to many solutions of nonpolar substances. Such solutions have a random arrangement, because any tendencies toward molecular association are sufficiently weak to be overcome by thermal agitation. These solutions are designated ”regular” solutions. The treatment of these solutions is more complicated than for ideal solutions, since consideration must be given to the differences in molal volumes and also to the differences in internal pressures of the two componcnts. Relatively few data are available on solutions of nonpolar solids in nonpolar liquids at low temperatures. Solutions of krypton and xenon ( l 5 ) , of argon ( i ) , and of ethane (6) in liquid oxygen have been investigated, as well as solutions of acetylene and carbon dioxide (4, 9), ethylene and propylene ( 6 , 16), and methane (3) in liquid nitrogen and in liquid oxygen. The purpose of the present investigation has been to obtain more solubility data on solutions at low temperatures and to study their deviations from Hildebrand’s equation for regular solutions. The systems investigated were ethane in liquid nitrogen, and ethane, ethylene, and propylene in liquid oxygen. EXPERIMENTAL
The oxygen used as solvent was prepared by heating potassium permanganate ; it was purified by passing it first over soda lime and then over phosphorus pentoxide (A, B, C in figure 1). The nitrogen used as solvent was obtained from a cylinder of Linde dry nitrogen. It was bubbled through concentrated sulfuric acid, passed over hot copper wire, and then over soda lime and phosphorus pentoxide. The ethane, ethylene, and propylene used as solutes were of technical grade and, after preliminary treatment with sulfuric acid and 50 per cent potassium hydroxide, were subjected to repeated fractional condensation until their vapor pressures did not change on further fractionation. Solubilities were determined at different temperatures by measurement of the lowering of the vapor pressure of the solvent in a saturated solution in the presence of excess solute. This method is based upon the assumption, which \vas experimentally verified, that the solution obeys Raoult’s law and that the vapor pressure of the solute is negligible. The differential manometer, D in figiire 1. registered the difference between the vapor pressure of the pure solvent, conThis paper is an abstract of the thesis submitted by A . Lucile Cox to the Faculty of the Graduate School of Purdue University as partial fulfillment of the requirements for t,he degree of Doctor of Philosophy, August, 1946.
666
ANNA LUCILE COX AND THOMAS DE VRIES
tained in E, and that of the solution in the solubility cell, G. The oxygen vaportension thermometer, P, gave both the temperature of the bath and the vapor pressure of the pure solvent. A rotating magnet, R, operated the bath stirrer, Q. All readings were made with a Gaertner cathetometer reading to 0.05 mm. and were corrected for capillary depression of mercury in glass BS well as the usual corrections for the thermal expansion of mercury. Readings were made after 3 hr. of intermittent stirring of the solution. Each reading was preceded by a period of 20 min. of stirring and then 15 min. standing with the stirrer held above the solution by a small permanent magnet, I. This procedure was necessary,
I3 I
FIG.1. Vapor preasure apparatus
since the cell stirrer conducted heat to the solution causing its vapor pressure to be slightly too high. The concentration of solute was sufficiently dilute so that the vapor pressure of the solvent would obey Raoult’s law, but experiments were conducted to confirm this fact. For this part of the work, known but increasing amounts of the solvent were added to a known amount of the solute to produce unsaturated solutions of different concentrations. The amounts of solute and solvent were determined by pressure-volume measurements in a gasometer (K in figure 1). The known mole fraction was compared with that calculated from the lowering of the vapor pressure. The data tabulated in table 1 show that, within the limits of experimental error, the vapor pressure of the solvents obeyed Raoult’s law.
667
SOLUBILITY O F HYDROCARBONS IN NONPOLAR LIQUIDS RESULTS
In order to make a quantitative test of Hildebrand’s equation for regular solutions, observed solubility data mere used to calculate (AEIV)’”, the internal pressure, which has also been called the “solubility parameter” by Hildebrand TABLE 1 Test of validity of Raoult’s law for the solvent
I
SOLUTlOS
‘K.
I
Ethane in ox) gen Ethjlene
111
85 96
oxkgen
Ethane in nitrogen
10.93 8.62 i.35
77 24
Logarithm of ideal solubili ______-‘K.
0.0229 0.01i7 0.00973
ETHANE
0 0224 0 0181 0 00969 0 00668 0 00667 0 00403
1
1 03 0 95 1 003 1 017 0 95 0 91
a*, as inole fraction 1
ETEYLEhT
PlOPYLEVE
-
70 75
-0.643
80
-0.382 -0.275 -0.179 -0.0926
-0.501
85 90
95 T,, OK. . . , . . . . , . . . . . . . . . AH,, cal. per mole.. . .. . .
1Ol.P 6Mt
* Reference
t
-0.813 -0.646 -0.502 -0.372
I
-0 -0 -0 -0
457 307 176 061
-0.258
-0.155 103.7*
~ 1 1 $ Reference 1. 0 Reference 14.
12. Reference 17.
and denoted by 6, and such values were compared with values calculated from energies of vaporization data. Hildebrand’s equation is
ET In
= Ti:!
(82
-
where N 1 and ‘VT~ are the mole fractions of solvent and solute, respectively, and V1 and V:!are the corresponding molal volumes. The activity, Q, of the solute is given by the well-known equation In a2 = A H / R (l/T, - l/T), in which AH is the heat of fusion of the solute and 2’1 is its melting point. The results of such a calculation are given in table 2 for the ideal solubilities of ethane, ethylene, and propylene together with the data required for the calculations. The calculated values of the “solubility parameters” of the solutes and solvents are tabulated in table 3. For these calculations the molal volumes were calculated from densities
668
ANNA LUCILE COX AND THOMAS DE VRIES
of the liquids (2) which, in some cases, had to be extrapolated over more than a 100" range. The energies of vaporization were obtained from various sources. For ethane, the expression A E = 4753 - 9.927' was obtained, using the A H values of Porter (13) at 184.5"K. and heat capacities of the liquid and vapor given by Wiebe (17) and by Kistiakowsky (10) and their coworkers. For ethylene, TABLE 3 Molal volumes and values of (AE/V)'12 MOLAL V O L r n S
~-
Ethane Ethylene Propylene Nitrogen Oxygen
__
43.20 39.75 52.25 33.28 25.79
,
43.61 40.14 52.67 34.20 26.28
SOLUBILITY PAPAXXTEP
__ -- __ __ 85'K. 7O.K. 80.K. - IS'K. __
WK. ~
44,02 40,54 53,09 35,18 26.80
-
44.44 40.95 53.51 36.22 27.34
9.70 9.97 10.65 6.14 7.94
-_
9.58 9.86 10.52 5.93 7.75
9.48 9.74 10.42 5.72 7.57
9.38 9.63 10.31 5.51 7.39
TABLE 4 Ezperimenlal solubilities and solubility parameters __
-
SOLUTION
'K.
AP
Ethane in oxygen.. . . . . . ,
77.3 80.0 81.1 85.3
2.55 6.14 7.83 24.80
-1.784 -1.564 -1.526 -1.250
-0.445 -0.382 -0.350 -0.264
11.4 10.8 10.9 10.4
3.38 3.28 3.30 3.22
1.87 1.91 1.93 1.99
Ethylene in oxygen.. , . , .
81.1 82.3 86.9 87.6 89.9
1.17 1.79 6.13 7.03 9.95
-2.350 -2.230 -1.939 -1.903 -1.870
-0.471 -0.438 -0.326 -0,310 -0.257
17.4 16.9 16.2 16.1 16.7
4.17 4.12 4.03 4.02 4.08
2.19 2.20 2.27 2.8 2.31
Propylene in oxygen. . . . .
83.0 83.3 86.8 87.4
0.46 0.88 3.23 3.70
-2.85 -2.64 -2.21 -2.18
-0.106 -0.105 -0.023 -0.010
19.6 18.3 16.6 16.6
4.43 4.28 4.07 4.07 '
2.89 2.90 2.95 2.95
Ethane i n nitrogen.. . . . . .
77.3 78.1 78.6
5.87 6.62 8.37
-2.11 -2.09 -2.04
-0.445 -0.426 -0.414
13.7 13.8 13.6
3.70 3.72 3.69
LOO a1
(62
61):
mm.
.
~
__
I ~
3.70 3.72 3.73
the expression AE = 4677 - 10.31T was derived from the extensive thermodynamic data of York and White (18). For propylene, the expression A E = 7011 - 15.437' was obtained from the AH value of Powell and Giauque (14), and the heat capacity data of Kistiakowsky and coworkers (11). For nitrogen and oxygen, the expressions A E = 1973 - 10.3T, and A E = 2217 - 8.5T were used, obtained from heats of vaporization given in the literature (8).
SOLUBILITY O F HYDROCARBONS IN NOKPOLAR LIQLTIDS
669
Representative results of the experimental work are given in table 4, together with a graph of the solubility of ethane in liquid oxygen and in nitrogen in figure 2. Our experimental results agree with those obtained by Tsin (16) for the solubility of ethylene in liquid oxygen, using an analytical method, but they differed greatly from his results for the solubility of propylene in liquid oxygen. He reported a value of 0.079 mole fraction a t 83’K. I n the present work it was noted that two liquid layers were present throughout the experiments with saturated solutions. Tsin’s method for determining solubilities depends upon filtering a mixture of the saturated solution and excess solute, followed by analysis of the filtrate. The presence of two liquid layers would give erroneous results in such a procedure. The quantity ( 6 2 - 6 1 ) was calculated from the solubility data by using Hildebrand’s equation; such values are given and compared in the last two columns of
*-”I
z
0
s
- 2.0
FIG.2. Solubility of ethane in liquid nitrogen and oxygen
table 4 with values based on thermodynamic data (see table 3). The results indicate that a solution of solid ethane in liquid nitrogen appears to be a “regular” solution. The agreement for the (62 - 61) values is better than one would expect.2 However, the solubilities in liquid oxygen gave values larger than those calculated: namely, differences of 1.4 for ethane, 1.8 for ethylene, and 1.3 for propylene solutions. Such a value means that the solubilities are less than that given by Hildebrand’s equation, a deviation which may be ascribed to strong intermolecular repulsions. Hildebrand’s derivation assumed that such repulsions were negligible. Because the molar volume of liquid oxygen is appreciably less than that of liquid nitrogen, the molecules in liquid oxygen solution would be 8 The solubility of ethylene was too small to measure a t the boiling point of nitrogen. Propylene gave a lowering of 0.25 mm., which corresponds t o 0.00033 mole fraction, a value too small t o be reliable.
670
ANNA LUCILE COX AND THOMAS DE VRIES
closer together and this might lead to stronger intermolecular repulsions. It has been observed (3) that the solubility of methane in liquid oxygen deviated more from ideal solubility than that of methane in liquid nitrogen. It also seems logical that larger deviations would be observed a t low temperatures when the kinetic energy of the solvent molecules might be unable to overcome the intermolecular forces between the solute molecules. I n fact, the data show such a trend of smaller deviations a t higher temperatures. An extended theory of solubility ought to consider the effect of intermolecular forces. SUMMARY
The solubility of solid ethane, ethylene, and propylene in liquid oxygen is less than that calculated from Hildebrand’s equation for regular solutions. The solubility of solid ethane in liquid nitrogen agrees well with theory. An explanation is offered, based on the concept of intermolecular repulsions between the solute molecules. REFERENCES (1) EQAN, C. J., AND KEMP,J. D.: J. Am. Chem. SOC.69, 1264 (1937). (2) EOLOFF,G. : Physical Constants of Hydrocarbons, Vol. I, Reinhold Publishing Corporation, New York (1939); International Critical Tables, Vol. 111,pp. 20, 21, McGraw-Hill Book Company, Inc., New York (1928). (3) FASTOVSKI?, V. G., AND KRESTINSKI~, Yu. A , :J. Phys. Chem. (U.S.S.R.) 16,525 (1941). (4) FEDEROVA, M. F.: J. Phys. Chem. (U.S.S.R.) 14, 422 (1940). (5) HILDEBRAND, J. H . : Solubility of Nonelectrolytes, 2nd edition. American Chemical Society Monograph Series, Reinhold Publishing Corporation, New York (1936). (6) HUNTER,M. A.: J. Phys. Chem. 10, 330 (1906). (7) INGLIS, J. K. H.: Phil. Mag. [6] 11, 656 (1906). (8) International Critical Tables, Vol. V, p. 135. McGraw-Hill Book Company, Ino., New York (1929). (9) ISHKIN, I. P., AND BURBO,P. Z.: J. Phys. Chem. (U.S.S.R.) 13, 1337 (1939). (10) KISTIAKOWSKY, G. B., AND RICE,W. W.: J. Chem. Phys. 7,281 (1939). (11) KISTIAKOWSKY, G. B., RICE, W. W., LACHER,J. R., AND RANSOM, W. W.: J. Chem. Phys. 8,610, 970 (1940). (12) MAASS,O., A N D WRIGHT,C. H . : J. Am. Chem. SOC.43, 1098 (1921). (13) PORTER,F.: J. Am. Chem. SOC.48,2055 (1926). (14) POWELL, T. M., AND GIAUQUE, W. F.: J. Am. Chem. SOC.61, 2366 (1939). (15) STACKELBERG, M. v.: 2. physik. Chem. A170, 262 (1934). (16) TSIN,M. N.: J. Phys. Chem. (U.S.S.R.) 14,418 (1940). (17) WIEBE,R., HUBBARD, K. H., AND BREVOORT, M. J.: J. Am. Chem. SOC.62,615 (1930). (18) YORK,R., AND WHITE,E. F.: Trans. Am. Inst. Chem. Engrs. 40, 227 (1944).
