670
N. THORP AND ROBERT L.Scow
Vol. 60
ature. Physically adsorbed ethylene or oxygen will decrease as the temperature is raised so we must be dealing with chemisorbed ethylene or oxygen. Ethylene appears notiJ3 to be chemisorbed on silver so that the amount of chemisorbed oxygen must increase with temperature. This can happen
if the rate of oxygen adsorption is comparable with the rate of reaction and the energy of activation for oxygen adsorption is greater than the true E A for the reaction. This mechanism is the one proposed by Twigg’ from his kinetic measurements and in part can firmed by others. 4, 16
(13) We have been able to measure the exchange between CzDl and Ha over a polycrystal Ag sheet after reduction but the rate a t 210° is considerably slower than the oxidation rate suggesting that ethylene chemisorbed to Ag plays no part in the oxidation reaction.
(14) A. Orzechowski and K. E. MacCormack, Canadian J . Chem., SP, 388,415 (1954). (15) S. Z. Roginskii and L. Y . Margolis, Doklady Akad. Nauk S.S.S.R., 89, 515 (1953).
149
FLUOROCARBON SOLUTIONS A T LOW TEMPERATURES. I. THE LIQUID MIXTURES CF4-CHF3, CF4-CH4,CF4-Kr, CH4-Kr BY N. THORP AND R. L. SCOTT Contribution from the Department of Chemistry of the University of California, Los Angeles, California Received November 86, 1866
The liquid-liquid binary systems, CHFa-CE, CHF+2rFa, CF4-CH4, CF4-Kr and CH4-Kr, have been studied at low temperatures (105-140’K.). Fluoroform is completely miscible with perfluoroethane but forms two hases with perfluoro-, methane below a consolute temperature of 130.5”K. These results are in reasonable agreement wit{ solubility parameter theory; the polar fluoroform has a much higher cohesive energy density than the fluorocarbons. The vapor pressures of methane-krypton, methane-perfluoromethane and krypton-perfluoromethane mixtures have been measured at temperatures near 110°K. The first system is nearly ideal, in agreement with the small difference in solubility parameters, but the latter two systems show abnormally large positive deviations from Raoult’s law, much greater than can be explained easily. The anomalous deviations in the CH4-CFd mixture are analogous to those found previously in other fluorocarbon-hydrocarbon systems, but the similar behavior of the CF4-Kr system was completely unexpected.
Introduction energies of vaporization per ml., be empirically In recent years, considerable interest has been increased by about 0.6 unit in order to fit the. aroused by the unusual solvent properties of fluoro- data. These two suggestions have been shown to carbons and related fluorochemicals. I n 1948 be mutually exclusive,12 and neither seems very Scott1 concluded that the unusually low mutual satisfactory for some systems.8 Simons and Dunlap3-13included corrections in solubilities of fluorocarbons and standard organic solvents were a direct result of their low solubility the theory of regular solutions to take into account parameters (5.7-6.0 cal.’/zcm.-8/z), and the experi- volume changes which occur on mixing. This mental data then available’s2 seemed to be in good treatment has been extended by ReedI4 to allow properly for differences in the ionization potential agreement with predictions. In the last five years, however, several hydro- of the molecules; the harmonic mean of the two potentials appears in the London formula carbon-fluorocarbon solutions have been s t ~ d i e d , ~ -ionization ~ and abnormally low mutual solubilities have been for dispersion forces, but in the past has usually uniformly observed, in disagreement with the been approximated by the geometric mean. With values predicted from the solubility parameters of this correction he was able to calculate partial molar free energies of mixing which were in accordhydrocarbons and fluorocarbons.1° As an explanation of this anomaly, Simons and ance with the experimental values for the hydroDunlap3 suggested an abnormally close interaction carbon but not the fluorocarbon in the systems between the C-H groups of adjacent hydrocarbon studied by Simons and c o - w ~ r k e r s . ~ ~ ~ In continuation of a general program of research molecules (“interpretation”) which gives rise to greater heats of mixing than that calculated from on fluorocarbon solutions, we have investigated the the &values. On the other hand, Hildebrand“ binary liquid systems of CHF3-CF4, CHF3-C2Fe, suggested that the solubility parameters of the CF4-CH4, CF4-Kr and CH4-Kr a t temperatures hydrocarbons, normally calculated from their between 105-140°K. In general, our vapor pressure measurements give deviations from Raoult’s (1) R. L. Scott, J . A m . Chem. Soc., TO, 4090 (1948). law which are much greater than would be expected (2) J. H. Hildebrand and D. R. F. Cochran. ibid., 71, 22 (1949). (3) J. H.Simons and R. D. Dunlap, J. Chem. Phys., 18,335 (1950). from the difference in solubility parameters. (4) J. H. Hildebrand, B.. B. Fisher and H. A. Benesi, J. A m . Chem. Experimental Soc., TB, 4348 (1950). (5) J. H. Simons and J. W. Mausteller, J . Chem. Phys., 8 0 , 1516 (1952). (6) J. H.Simons and M. J. Linevsky, J . A m . Chem. Soc., 74, 4750 (1952). (7) G. J. Rotariu, R. J. Hanrahan and R. E. Fruin, ibid., 76, 3752 (1954). (8)3.L.Scott and E. P. McLaughlin, ibid., 76,5276 (1954). (9) J. A. Neff and B. Hickman, THISJOURNAL, 69, 42 (1955). (10) J. H. Hildebrand and R. L. Scott, “Solubility of Non-electrolytes,” 3rd Edition, Reinhold Publ. Gorp.. New York, N. T.,1950. (11) J. H. Hildebrand, J. Chcm. Phys., 18, 1337 (1950).
Perfiuoromethane, perfluoroethane and fluoroform were obtained from the Jackson Laboratoreis of E. I . du Pont de Nemours and Company, Incorporated, and were further purified by repeated passage over activated charcoal held a t solid COZ temperatures. Before being admitted to the storage bulbs, traces of air were removed by repeatedly freezing, pumping while frozen, and melting. Semi-quan(12) R. L. Scott, J . Chem. Ed., S O , 542 (1953). (13) R. L). Dunlay, J. Chem. f’hys., 81, 1293 (1953). (14) T. M. Reed, 111, THIBJ O U R N ASS, L , 425 (1955).
FLUOROCARBON SOLUTIONSAT Low TEMPERATUUV
May, 1956
67 1
titative mass spectra were taken of the purified products and they were found to contain less than 1% of impurities. Phillips Research Grade methane and Airco krypton were used directly. Apparatus.-The general form of the apparatus, designed for use when only small quantities of gases are available, is shown diagrammatically in Figs. 1 and 2, the latter showing in greater detail the vapor pressure vessel and temperature control system. The vapor pressure vessel has a volume of about 2 cc. and is closed by a click gage sensitive to a small pressure differential which is utilized to measure the total vapor pressures above the liquid mixture. This vessel is placed in a Dewar type container and is held at constant temperature by a bath of refluxing liquid. The boiling point is held constant by means of a solenoid operated valve, which admits liquid nitrogen to the cold finger whenever the pressure rises. Both the bath and the experimental mixture are stirred by means of magnetic stirring bars. The cold finger is filled with copper turnings to aid the heat transfer from the condensin vapors. By this arrangement it was possible to controk the bath temperature to 0.2-0.3'K., below 112OK. by using methane, above 112'K., by using methane-propane mixtures. Procedure.-A sample of gas was drawn from the storage bulbs into the Toepler pump, where its volume was measured, and then transferred to the vapor pressure vessel, where it was condensed. Small amounts of the second component were then added in a similar manner. The mole fraction composition of the liquid mixture was taken to be the ratio of the gas volumes, since the volume of the vapor pressure vessel was less than 1% of the volume of gases added. To obtain mutual solubility data, the temperature of the bath was varied until two phases separated. This process wae repeated several times until the unmixing temperature was determined to about 0.5"K. Vapor pressure readings were taken after the mixture had been allowed to equilibrate for about five minutes. The pressure on the upper side of the click gage was varied until i t operated. From the calibration of the gage, the vapor pressure of the mixture can be calculated with an estimated accuracy of about 1 mm.
Results Table I gives some of the pertinent physical properties of the compounds studied. TABLE I
SOMEPHYSICAL PROPERTIES OF THE COMPOUNDS STUDIED CF4 M.p. (OK.) B.p. (OK.) Heat of vaporization, koal./mole Molal vol. at b.p.
CHFB CHI 113.0 91.7 189.0 111.7
89.5 145.1 3.01
4.23
54 2
1.95
Kr 115.9 119.9
2.16
CzFa 173.1 197.8 3.86
( A B V / V ) ' h = d a t b.p. d at 130°K.,oal.'h
7.1
47.8 9.0
38.0 6.8
34.0 7.5
86.3 6.4
am. Electron polarizability X 1024,cm.8 Ionization potential, e.v.
7.3
10.3
6.6
7.3
7.7
4.02 17.8
.. . ...
2.58 13.16 14.5
2.54 13.93
(00.)
... .. .
21
In
sm[a
y1
+
I Mo gnrt ic
Strrrar
Figures 3, 4 or 5 show the measured total vapor pressures versus the mole fraction of one of the constituents in the mixture. The circles are the experimental points, and the full curves are those of the theoretical total vapor pressures, calculated by assuming that the concentration dependence of the excess free energies of mixing can be described by the equations of Scatchard.'6 A@/RT =
r
+ z2In
@(XI
-
y2
52)
=
+
451
-342
+ ...1
(1)
Considering the first two terms only, we obtain (15) G.Soatchsrd, Chem. Reus., 64, 7 (1949).
Fig. 2.-Experimental arrangement for measuring vapor pressures at low temperatures.
for the partial molal free energies of mixing
RT = In
y1
=
z221a
+ o(3rL - sz)l
AKE RT = In Y Z = me[a- 8 ( 3 t ~- a ) ]
(2a) (2b)
The total vapor pressure P of a two-component mixture may be expressed in terms of the vapor pressures of the pure components py and pg as P = pPa1
+ p;ar
= 'pP",Y,
+ p:z2YP
(3)
N. THORP AND ROBERT L. SCOTT
672
170
Vol. 60
the non-ideality of the vapor phase were made. These were found t o have a negligible effect on the value of the constants a and p. Table I1 gives the value of these constants for the systems CF4-Kr, CH4-Kr and CF4-CH4. TABLE I1 VALUESOF THE CONSTANTS AND 8 BY THE METHODOF LEASTSQUARES FOR THE THREE SYSTEMS (Y
t
\
CFC-CHI, CF4-Kr, CHrKr Equations In y1 = [CY p(3z1 - 22)]22z In y2 = [CY- 8(3z2 - zl)]zl*
+
w
i20
lo L I I I 1 , I I , , 10 CH4 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 CF4 $2.
pressures of CH4-CF4 mixtures.
Fig. 3.-Vapor
System
Subscript 1 denotes
a
B
CHa-CF4 at 110.5”K. CH4-CF4 at 108.5”K. C&-CF(at 105.5”K. CF4-Kr at 117.1”K. C&Kr at 115.5’K.
CHI CH4 CH, Kr Kr
1.56 1.61 1.78 1.21 0.25
0.36 0.42 0.30 0.32 0.0
The solubility of a c,omponent of a solution of non-polar non-electrolytes may frequently be explained with the aid of the simple regular solution equationlo APE = (ZITI 22r22)(&- &)2+~+z (6) where APE is the excess free energy of mixing; XI, x2 are mole fractions; c $ ~ c, $ ~are volume fracare molar volumes; al, a2 are the tions; “solubility parameters” of the pure components and are the square roots of the energy of vaporiza; subscripts refertion per ml. ( i e . , A E v / V ) ’ / ~ the ring to components 1 and 2. An alternative equation is obtained by using the Flory-Huggins entropy which attempts to correct for the d,ifference in the molecular size of the two components
+
I
1
I
I
L
I
L
I
I
!o
Kr 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 CF4 $2.
Fig. 4.-Vapor pressures of Kr-CF4 mixtures.
r1,r2
A@ = R T [ In ~ (+1/a)
+
22
+ + z~VZ)(&-
In (+2/z2)1
(ziri
6~)~+1+2
(6)
where the symbols have the same meaning as equation 5. a1 -
1
I
I
I
I
I
Kr 0.1 0.2 0.3 0.4 0.5
I30 0.6 0.7 0.8 0.9 CH4 I
t
I
I
$2.
Fig. 5.-Vapor pressures of Kr-CH4 mixtures.
Combining this equation with 2a and 2b we get P = p y a exp Z P * [ ( Y
+ 8(3m - 4 1+
PPZZ exp m Z [ a- 8(3n - a ) ] (4) where a1 and u2 are activities, xl,x2 mole fractions of components 1 and 2, respectively, and CY and P are constants. From the initial slopes of the total vapor pressure curve, a rough extimate of the value of the constants CY and p was made. The final value of these constants was obtained by a method of least squares. l6 From the second virial coefficient and the equations of Scatchard and Raymond1’ corrections for (16) J. A. Barker, Aualml. J . Chem., 6 , 207 (1953). (17) G.Saatchard and C. L. Raymond, J . A m . Chem.,Soc., 60, 1278 (1938).
6z
TABLE I11 DIFFERENCES FOR THE THREE SYSTEMS CF4-CH4, CF4-Kr AND CH4-Kr Calcd.
System
61’
6io
Id1
-
681
Exptl. A6 Eq. 5 Eq. 6
CFd-CH, 7.6 6.8 0.8 2.9 CF4-Kr 7.6 7.6 0.0 2.7 CH4-Kr 6.8 7.6 0.8 1.3 a Solubility parameters estimated at 110°K.
2.6 2.3
1.1
Table I11 gives the “thermodynamic” solubility parameters of the three substances, evaluated from the heats of vaporization and molar volumes of the pure components at 110°K. The difference of these 6’s is then compared with the “experimental” values obtained by fitting equations 5 and 6 to the experimental excess free energies. Table IV gives the excess free energies of mixing a t x = 1/2, for the systems studied, calculated from (a) the experimental value of the activity coefficient and (b, c) calculated from the ‘‘thermodynamic” 6 values of the pure liquids, assuming (b) ideal entropy of mixing and (c) the FloryHuggins entropy of mixing.
FLUOROCARBON SOLUTIONS AT Low TEMPERATURE
May, 1956
TABLEIV EXCESS FREEENERGIES OF MIXINQ, CAL/MOLE AT XI = xz = I / I CFeCHd at
CFd-Kr at
110.SoK. 117.1’K.
AiF
Obsd. 86 Calcd. from “thermodynamic” 6’s Eq. 5 7 Eq. 6 5
at
75
14
0 -5
6 6
Figure 6 shows the miscibility curve for CHFa and CF4. The critical solution temperature is 130.5OK., and the critical mole fraction of CHFa is about 0.43by the method of rectilinear diameters. The system CHFa-C2Fe is miscible in all proportions. The simplified exprewion for the critical solution temperature, To, is given by the regular solution theory1° as 4RTo
(TI
+ rd61-
62)’
133.2
- 140
128.2
- 145
CHcKr 116.6’K.
673
Y 8 123.2
-150
-@
9 $ 42
k
B
118.2
- 155 H
113.2
- 160
(7)
where the P’s and 6’s have the same meaning as in equations 5 and 6.
CF,
0.2
0.4
0.6
0.8 CHFs
22.
Discussion
Fig. 6.-Immisoibility curve for CF4-CHF8.
The immiscibility of fluoroform and carbon tetrafluoride near the boiling point of the latter has been reported by Hadley and Bigelow.’* From the heats of vaporization a t the boiling points,’g values a t other temperatures were estimated by methods described elsewhere.lO With an accuracy of 0.1-0.2 6 unit, the solubility parameters of CF4 and CHFa a t 130’K. are 7.3 and 10.3 (cal./cm.S)’/s respectively. The difference of these two is 3.0, and this compares favorably with a value of 3.3, derived from the experimental results and equation 7. Thus this system is in reasonable agreement with theoretical predictions. The great difference between CF4 and CHFa, as evidenced by the difference in solubility parameters, is also reflected in the large difference in the heats of vaporization and boiling points, in contrast with the great similarity in properties of CC14 and CHCla; an appealing explanation of the abnormal behavior of fluoroform lies in the possibility of hydrogen bonding. C2Fs and CHFs do not form two phases. The estimated (61 - 62) difference is 2.6 a t 190°K. This gives To, by equation 7, about llO°K., well below the freezing point of C2Fs. Therefore the miscibility of C2FrCHFa mixtures is in agreement with theory. The solubility parameters at 11O0K. for CF4, Kr and CH4 are calculated as 7.6, 7.6 and 6.8, respectively, with an accuracy of 0.1-0.2 6 unit. The (61 - 62) differences from these figures gives 0.0 for CFd-Kr, 0.8 for CFrCH4 and 0.8 for KrCHd. The experimentally determined values for the same systems are 2.7, 2.9 and 1.3, respectively. The difference between the calculated and observed excess free energy of mixing a t x1 = xp = 1/2, Table IV, ‘is about 8 calories for CHd-Kr mixtures, which is only 10% of a measure of thermal energies, 1/2 RT. Small effects of this magnitude may be due to any number of minor causes, which
are ignored in the derivative of the regular solutionsolubility parameter equations. The experimental error in the measurement of excess free energy of mixing is certainly of this order. I n this light, therefore, the CH4-Kr system is unexceptional and in good agreement with theory. However, this is not the case for the CHd-CF, and Kr-CF4 systems. Here APE at z = 1/2 is approximately 86 cal./mole (0.39RT) at 110°K. and 75 cal./mole (0.32RT) at 117OK., respectively. Such a result was not unexpected for CH4-CF4 mixtures in view of anomalous behavior found by previous workersaeg for fluorocarbon-hydrocarbon pairs. The large deviations from ideality found for Kr-CF4 mixtures were not anticipated; solutions of fluorocarbons with liquids other than hydrocarbons have been found to conform well to solubility parameter theory. Since the differences in molar volume of the various substances are not very large, use of the Flory-Huggins entropy of mixing does not alter the situation appreciably. Increasing the 6 value of the hydrocarbon by about 0.6 unit, as suggested by Hildebrand,ll only makes matters worse, for then the 6 value of CH4 almost equals that of both CF, and Kr. No information is available on the volume changes of mixing for these systems, so we are unable to make the corrections of Simons and Dunlap.a There is little agreement on the value of the ionization potential for CF4, so it has not been possible to apply the corrections of Reed14with any success. Acknowledgment.-This work was supported by the Atomic Energy Commission under Project 13 of Contract AT(l1-1)-34 with the University of California. We wish to thank the Jackson Laboratories of E. I. du Pont de Nemours and Company for their generous gift of the CF,,CHF, and CaFe used in these experiments, and Dr. B. B. Fisher for the construetion of the major part of the apparatus.
(IS)E. H. Hadley and L. A. Bigelow, ibid.. 82, 3302 (1940). (le) “Selected Values of Chemical Thermodynamic Properties,” United Stater Government Printing Office, Wsahington, D. C.
