Baron eiehloride and hydrogen dlso reacted with
pounds. Although the effect of hydrogen in inhibiting the decomposition of hydroborons a t room temperature is well known, i t well may be that this effect is not operative at elevated temperatures and the hydrogen may be effective largely by diluting the diborane product. In addition it may be possible that more stable monoboron compounds, e. g., BHClo, B a C l , may preparative method, however. be the primary products of a stepwise reaction and that these transient compounds later undergo Discussion rearrangement into diborane and boron trichloIt indee ride when out of l%e reactor hot zone. thermally In the runs where boron fluoride was used no under con inflammation was noticed even though large rather rapid pyrolysis of this material. Several factors, however, may explain why so little py- quantities of diborane were present a t times in the rolysis actually does occur and why the losses of exit gas. This may indicate either that diborane monofluoride (BzHaF) is not formedin the reaction boron are small. The actual contact time during which a unit of or is not spontaneouslyin%ammableif it is formed. Reasoning from the experiments with metal gas remained a t elevated temperatures was in most cases rather short (Table I). This would hydrides, the reduction of boron halides with tend to minimize p sis losses, Also, there hydrogen over metals similarly may appear at generally has been p in the reaction zone a first glance to involve the initial formation of relarge excess of hy r the amount actually active hydrides on the surface of the metal folneeded for the formation of boron-hydrogen com- lowed by replacement reactions. There is, however, little direct evidence to support this over TABLE1 other possible mechanisms, and at the present time the exact nature of the reaction is not known. DIBORANE PREPARATIONS T,
Condensed product,
n1/
*C. 1 360 2 320 3 360
BCls 4-1 3-1 3-1
C. T.” 5“ 12” 12”
-40
4 350 5 476 6 600
6-1
15”
6-1
-40 -30
6-1
6” 4’
460
6-1
6”
7
&ize
-12 -40
-30 -30
mole per cent. HCl BCla 2.9 4.1 93 2 5 1.3 95.2 5 8 0 94.2 8.8 1 7 89.5 12 4 i 7 85 9 6 7 0 9 92.4 17.2 3.9 78.9 ByHe
Conver s1on.
Summary
%b
6 6.7
11 16.4 20.4 12.5
80
*C.T. Calculated contact time in hot zone, 1.” dia.
reactor. Per cent. BCia converted to R“6, assuming no boron loss in the reactor.
Dibarane and diborane monohalides have been produced by the reduction of boron halides with hydrogen over active metals a t elevated temperatures. The reaction of gaseous boron halides with active metal hydrides also has been found to produce diborane. SCHBNECTADY. KEW YORK
RECEIVED JUNE
a,1948
[CONTRIBUTION FROM THE CHEMICAL LABORATORY OF THE UNIVERSITY OF C A L I ~ R N I A ]
Liquid-Liquid Solubility of Perfluorornethylcyclohexane with Benzene, Carbon Tetrachloride, Chlorobenzene, Chloroform and Toluene BY J. H. HILDEBRAND SND D. R. F. COCI-IRAN properties of fluorolight during the last
tion, does riot nvcrc;tmin existing solubility theory (1) R. L.Scott, Txrs J O U R N ~ L , 70, 4090 (1948) (2) H A Renew and T Ti Nildebrand. zhcd 70. 1978 (198h~
for regular solution^.^ This instance of the adequacy of theory to cope with an extreme departure from Raoult’s law encouraged us to measure other fluorocarbon solubilities in the expectation that they would have much more than merely em pirical value and could be used in connection with refinements in the general theory. The large molal volumes of the fluorocarbons compared with the corresponding hydrocarbons leads to unusually low “internal pressures” or “cohesive energy densities,” while their high (3) Cj’ J I-I Hildebrand and R. L. Scott, “Solubility of NonPTrcfrol>tc, ” 3rd rditioii lq‘$‘~ Reinbnld Pnbl rorp , Nrrv S o r t
\T
’:
LIQUID-LIQUID SOLUBILITY OF PERFLUOROMETHYLCYCLOHEXANE 23
Jan., 1949
TABLE I CONSOLUTE TEMPERATURES, ‘c., OF
SOLUTIONS OF
PERFLUOROMETHYLCYCLOBEEXANE AT VARIOUS VOLUhfE FRACTIONS
62,AND MOLEFRACTIONS, x2, OF THE OTHERCOMPONENT
c
n
0.909 .goo .833
0.953
.800 .769 .714 .700 .667 .625 .600 .500 .400 .333 .300 ,286 .231 ,200 .167
... .910 ...
3.0
...
xr 0.960
1
rr
1
26.3
0.956
61.1
17.5
...
... .916 ...
23.4 25.5
.890 .858
...
”..
... ...
.. ,. ..
.752 ,669 .574
26.8 26.7 25.2
.784 .708 .618
50.2 50.1 49.0
.870
I834
... .
.
I
...
.. ...
...
... ...
.447 .377
19.9 14.4
...
,288
.loo
...
.091
.168
...
4.0
... -8.0
..
41.0
..
... .916 ...
46.6 48.9
.876 .845
81.6 83.9
.814 .785 .767 .687 * 593 .523
84.7 85.1 85.3 84.9 82.7 79.2
..
...
.. .. 38.6 ..
.327
28.9
...
... ... .422
...
.195
..
...
...
.467 .397
...
,305
..
...
7.2
.190
molecular weights aid in maintaining them in the liquid state; the net result of which is to make many of their solutions with ordinary liquid hydrocarbons deviate from Raoult’s law to such an extent as to form two liquid phases, a state of affairs comparatively rare among non-polar liquids. Since liquid-liquid solubilities are easily determined, we have used t h i s means of further comparing theory with experiment. Materials.-We are indebted to the Jackson Laboratory, du Pont de Nemours Co., for the supply of perfluoromethylcyclohexane. The benzene, “reagent grade, thiophene free,” showed no color with concentrated sulfuric acid. It was dried with calcium chloride and distilled. The carbon tetrachloride, “reagent grade,” was washed with concentrated sulfuric acid, then with water, and dried over calcium chloride. The chloroform, also “reagent grade,” was washed successively with concentrated sulfuric acid, sodium carbonate aq., water, and dried over calcium chloride. The toluene was that used by Benesi and Hildebrand.2 The chlorobenzene, “reagent grade,” was dried over phosphorus pentoxide and distilled. All the above distilled within 0.1’. Procedure.-The two liquids were measured by means of a carefully calibrated pipet made of 2-mm. capillary tubing into tubes 10 cm. long and 7 mm. diameter through capillary stems a t right angles. The liquids were frozen in liquid nitrogen after introduction and sealed in. These tube; were heated above the consolute temperatures in a water-bath for 0 to 90’ and in sulfuric acid for the range 90 to 130’. On slowly lowering the temperature while rocking a tube, the unmixing temperature of a solution could be accurately determined.
xa
..
76.1
..
..
.. 63.7 .. 83.1 .. ..
...
..
.733 ,647 .550
88.9 88.6 87.2
.440
82.0
58.9
..
. . I
.905
... -.. ...
86.9
...
110.4
... .. ...
... ...
...
.742 .657 .561
126.8 126.7 123.5
...
... ...
..
”..
...
68.7
...
.277
42.8
...
92.9
.169
.161
65.7
...
35.0
..
0.950
,314
...
..
87.7
1
XI
... ... ... 108.1 ...
...
75.7 69.3
1
... 0.943 ... .880 ... ,.. .810 ...
...
..
ChHrCI
CEHKCHI
.. ..
..
*.
...
.365
1 . .
Results.-Table I gives the consolute temperatures of the different solutions together with the corresponding volume fractions of the fluorocarbon as measured a t 2 5 O , and the corresponding mole fractions. These temperatures are plotted against mole fraction in Fig. 1 and against volume fraction in Fig. 2. In Table I1 are given the critical mixing temperatures read from the maxima of these curves, also the critical compositions determined by aid of the assumption of rectilinear diameter, as in the analogous case of liquid-gas equilibrium.
0
0.5
1.0
e. Fig. 1.-Consolute temperatures of pertluoromethylcyclohexane solutions vs. mole fraction of second component.
J. N. HILDEBRAND AND D. R. F. COCHRAN
24
Vol. 71
inurn occurs at equal volume fractions, i. e., 61 = 6, = 0.5, as shown in Fig. 2. The curves are so %at at the top that this may suffice for ordinary pur100
poses. It has been pointed out, further, that eqn. 2 yields values of T, not seriously different from a greatly simplified equation 4RTc =
(VI
+
v2)(61
- &)*
(3)
In Table I1 can be seen comparisons between obc, served and calculated values of the critical compow sitions and temperatures. The “observed” values of x2 were taken from the solubility curves by aid of the assumption of rectilinear diameter, as in gas-liquid density curves. The curves are so flat around the maximum that these values are subject to an error of at least *0.02. The “calculated” values were obtained from Equation 1 by substituting the molal volumes for 2 5 O , given in the first column of figures. It would be preferable to use the molal volumes at the critical tempera0 0 .-I 1 [I tures, but since we lack some of the necessary dens47. ity data and in view of the uncertainty in fixing Fig. 2.--Consolute temperatures of perfluoromethylthe “observed” values, a closer check is hardly to cyclohexane solutions DF. volume fraction of second combe desired. ponent. We have found that values of 61 - 6 2 calculated T A B L E ii from T, by eqn. 3 agree within a few per cent., COMPARISOX BET WEE^ OBSERVED AND CALCULATEDi. e., well within the limit of the unavoidable errors and uncertainties of the method, with those \.ALUES OF CRITICAL COMPOSIIIOXS AND TEMPERATURES derived by aid of the more complicated eqns. 1 and f’l3K SOLriTIOUS O F PERFLUOROMETHYLCYCLOKEYAVE (COM2, therefore eqn. 3 has been used in the following POYFNT 1’ v calculations. Substituting the observed values of cc., Crit. x4 To dl T, and the v’s for 25’ gives 82 61, which, sub23O ohs. calcd. ohs. calcd. 6. calrd tracted from the &values in the next to the last C FuCFt 195 column gives the calculated values of 61 for perCCll 97 0.70 0.74 300.0 269 8 6’ 5 8 fluoromethylcyclohexanegiven in the last column. CIICla 81 .71 .78 323 5 325 9 0“ 6.0 In so calculating we have neglected the effect of GH6 89 .73 .76 358 5 360 9.15l D 0 temperature upon both 61 and 62 and on this acC,H&Ha 102 . 7 1 362 0 324 8 91 5.8 count alone we should not expect uniform values C HeCl 107 1 .72 300 0 464 9 .i’ A 3 for 61 “calculated.” In addition, the dipole moThis value is consistent with the solvent power of ments of chloroform and chlorobenzene might be \ chloroform for other substances;‘ the value of ~ A E/v)’/z is 9.3. See J. H. Hildebrand, A Critique of thq Theory expected to cause some disturbance. It is sufof Solubility of Won-eirctrolytr~.” Chem. KPTJ. 111 prcss. ficient to find that these values of 61 for the fluorocarbon derived from solubility data agree well Comparison with Theory.-The regular solu- with the value 6.0 a t 25’ set by Scott‘ long before tion equations4 for liquid-liquid solubility are as this work was undertaken. Reversing the calcufollows, for the mole fraction at the critical point lation, we have calculated T, for each solution and the critical temperature, respectively using this value of 61 in combination with the 82values given in the table. It will be seen that while the result is far from precise it has value, nevertheless, in giving the general region in which and two liquid phases can coexist. The high critical temperature with toluene is not what its &value, less than that of benzene, would lead us to expect. where the v’s denote molal volumes, x’s mole A similar reversal has been noted by Hildebrand fractions, and 6’s are what we have been calling and Jenks6 in the solubilities of sulfur in these two the “solubility parameters,” which are ordinarily liquids. An explanation of such facts must await the square root of the energies of vaporization per refinements in the theory as well as additional cc., i. e., 6 = ( A E v / ~ ) 1 / 2In . this paper we shall data, such as partial molal volumes. distinguish the fluorocarbon by the subscript 1. Aclmow1edgments.-We wish to express our It has been pointed out that eqn. 1 yields values appreciation to Dr. J. M. Tinker, Director of the of the critical composition not far from those ob- Jackson Laboratory, E. I. du Pont de Nemours tained by the simpler assumption that the maxiJ 11 Iiildrhratid and 1 Tenks, Tnrs TOURNAL, 43, 2172
-
Ii)
(4) Cf. ref. 3, Chap XVI
(I‘J21
I
