796
M. S.B. MUNSON
Solutions of Fluorochemicals and Hydrocarbons
by M. S. B. Munson Humble Oil and Relining Company, Research and Development, Baytown, Texas (Received October 7, 1965)
Critical solution temperatures for several paraffins and mixtures of paraffins in fluorochemicals can be correlated well with AEv even though the Hildebrand-Scatchard solution theory is inadequate. A single curve of mole fraction composition us. T / T owill reproduce the solubility data for an isomeric series of compounds with (C4F9)3N,and a solubility curve of volume fraction vs. reduced temperature, T/T,, will adequately represent the solubility data for pentanes, hexanes, and heptanes in (C4Fg)3N. From the correlations of critical solution temperatures of a series of hydrocarbons with different fluorochemicals one would expect similar behavior for other noncomplexing fluorochemical-hydrocarbon systems. Chemical shifts for paraffins in (C4F9)3Nshowed no correlation with AE, and gave no indication of any solution anomalies.
Introduction Solutions of fluorochemicals and hydrocarbons furnish examples of highly nonideal mixtures (without the complication of complex formation) whose components may be varied through a wide range of chemically similar materials. It has been observed, however, that solutions of fluorochemicals and hydrocarbons do not behave as one would expect from HildebrandScatchard solution theory. Although extensive comment has been forthcoming, no satisfactory reconciliation of theory with experiment has been achieved.’ The critical solution temperature is a reasonable, though indirect, measure of solution nonideality and one which is relatively easily obtained. Critical solution temperatures for various compounds with fluorochemicals have been measured and it has been found for many fluorochemical-hydrocarbon systems that predictions based upon the Hildebrand-Scatcliard equations (including a correction for inequalities in molar volumes)
+
RT, E (61 - 62)2(2V~Vz)/(V1”z Vz1i/z)2 do not agree with the experimental values. v is the molar volume; 6 is the solubility parameter, (AE,/V)‘”, in which AE, is the molar energy of vaporization; R is the gas coiistant; and To is the critical temperature in O K ‘ Disagreement between and experimellt should not be Surprising Since To iS related The Journal of Physical Chemistry
to the second derivative of the excess free energy of mixing8 and drastic approximations are made in the development of an equation for the excess free energy of mixing in solubility parameter theory. The parameters of the solution theory, 6 or AE and V , have proved useful in describing “weakly interacting” systems and should be helpful with the noncomplexing fluorochemical-hydrocarbon systems.
Results and Discussion Solubilities of several hydrocarbons were determined with portions of fluorochemical taken from one carefully purified sample. The purity of this sample of fluorochemical (C4Fg)3N, perfluorotributylamine from the Minnesota Mining and Mfg. Co., was estimated at >95% by gas chromatographic analysis. The impurities were close-boiling fully fluorinated compounds. The purification procedure involved passing through silica or alumina, refluxing with Ki\iInO4 HzS04, refluxing over KOH, and distilling. Of these procedures, only extremely careful fractionation had any effect on the distribution of fully fluorinated materials. The hydrocarbons were Phillips Research grade. A comparison of the critical solution tempera-
+
(1) R . L. Scott, J . Phys. Chem., 6 2 , 136 (1958). (2) J. H. Hildebrand, B. B. Fisher, and H. A. Benesi, J . A m . Chem. Soc., 7 2 , 4348 (1950). (3) J. 5. Rowlinson, “Liquids and Liquid Mixtures,” Butterworths Publications, Ltd., London, 1959, Chapter 5 .
SOLUTIONS OF FLUOROCHEMICALS AND HYDROCARBONS
797
Table I : 1Vlixturt:s of Paraffins with (C4FB)sN to, OC.
2,2-Dimethylbutane 2,3-Dimethylbutane 2-Methylpeatane 3-Methylpen tane n-Hexane
6.71 6.97 7.02 7.13 7.27
133.7 131.2 132.9 130.6 131.6
6.024 6.367 6.545 6.642 6.947
(XHCh
26.29 37.42 45.04 48.28 59.90
132.0 f 1 . 0 n-Pentane 2-Methylbutane
7.02 6.75
116. 1 117.4
7.43 6.96
147.6 149.9
5.723 5.344
36.94 22.83
)aN
5.9a
358.0
0.807 0.798
0.802 f 0.005
8.142 7.267
81.02 50.63
0.764 0.763
__
0.764i0.001
148.7 f 1 . 2 ((288
__
0.785i0.005
116.8 f 0 . 7 n-Eeptane 2,4-Dimethy:lpentane
0.780 0.783 0.781 0.795 0.786
13.2a
G. J. Rotariu, R. J. Hanrahan, and R. E. Fruin, J . Am. Chem. Soc., 76, 3752 (1954).
tures reported hlere and others in the literature affords fair agreement, which is all that one may reasonably expect. The solubility curves were determined in the accustomed manner by changing the temperature of a mixture of known Composition (with agitation of the sample vessels) until two phases appeared (or disappeared). The thermometers were calibrated and the accuracy of the temperature measurements is f0.02'. Critical solution temperatures and compositions were obtained by plotting the average hydrocarbon concentration, 'lz(XHc(hydrocnrbon phase) XHc(fluorochemical phase)j , against temperature and noting the intersection with the solubility curve. The lines
exhibited only a slight curvature and were easily extrapolated. The values of the critical solution temperatures and compositions, together with pure component properties, are given in Table I. For hydrocarbons of the same molar volumes, the critical concentrations are constant and decrease as the molar volume of the hydrocarbon increases toward that of the fluorochemical. From a plot of vs. ~ H C(which for the hexanes is not a good fit to the data) one obtains average values for the solubility parameter of the fluorochemical (CdF&N: 4.87 from the pentanes, 4.46 from the hexanes, and 3.24 from the heptanes. This agreement is not good and the values thus obtained do not agree
Figure 1. Critical ;solution temperature us. energy of vaporization.
Figure 2. Mole fraction solubility vs. reduced temperature.
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Volume 68,Number 4
April, 1,964
798
M. S.B. MUXSON
1.000
0.950
s
0
0
9.900
@ HC
Figure 3. Volume fraction solubility us. reduced temperature.
with that calculated from the properties of (C4F,),I\J. If one disregards the specific equations of the Hildebrand-Scatchard theory but still considers the energy of vaporization and molar volume to be measures of the interaction energy and size of the molecules, one may hope to find an empirical equation relating the critical solution temperature and these parameters. Figure 1 shows a plot of t, us. AE,(298OK.) which gives a good fit to a straight line for the hexanes. The slopes of the lines are substantially the same for the homologs and the intercepts decrease with increasing c aphydrocarbon volume. A plot of t, us. 6 ~ shows preciably more scatter for the hexanes than does the plot of t, us. AE, and does not allow the data for pentanes, hexanes, and heptanes to be plotted on one curve. A plot of t, us. 62 shows slightly more scatter of the data than indicated in Fig. 1 and does not allow the points to be considered as one curve. Although the quantitative description of the theory is poor, the qualitative agreement is good in that the critical concentrations of the isomeric species are the same and an increase in (or AE,) is paralleled by an increase in to. Similar behavior was noted for the hexanes with “perflu~roheptane’~~ and “perfluoropentane.”6 By an obvious (although shaky) analogy to the corresponding states treatment of gases, one would expect a one-parameter equation to correlate the solubilities of isomers in a given solvent, since the molar volumes (which are perhaps the same as the sizes) are the same. Figure 2 shows the solubilities for the heptanes, hexanes, and pentanes plotted as functions of the reduced temperature, 6’ = T/T,. In this figure, the curve for the heptanes has been displaced upward by 0.02 and the curve for the pentanes downward by 0.02 to improve legibility. The data fit remarkably well to these curves. The solubility curves for the homologs are The Journal of Physical Chemistry
not superimposable. The marked assymetry of the curves is that usually observed in systems with a large disparity of molecular sizes. These data also seem to indicate that there is no pronounced effect of molecular shape within this range. Since a one-parameter equation was so successful for molecules of the same size, one is led to attempt to reduce all of these solubility data to a single curve. It has been observed repeatedly that the pronounced assymetry in solubility curves for mixtures of molecules of different size is reduced by using volume fractions instead of mole fractions. Figure 3 shows a plot of volume fraction solubility of hydrocarbon against reduced temperature. The volume fractions are calculated at 25” assuming no volume change on mixingthe volume change on mixing is small, though not zero.6 The curve is much more nearly symmetric than those in the previous figure, but the critical solution composition is at I$HC = 0.572 f 0.002 (for pentanes, hexanes, and heptanes) and not +Hc = 0.500. As indicated by the legend in Fig. 3, the solubilities of nine paraffins (two pentanes, five hexanes, and two heptanes) are reproduced by one curve. The observed correlations may perhaps be fortuitous or perhaps the range of hydrocarbon volumes is not great enough to test it accurately, but the representation of the present data is good. A hasty search of the literature indicates no other available tabulation of data on solubility of paraffin isomers in fluorochemicals. Some studies have been made but they are not extensive enough to be of utility in verifying a reduced equation of solubilities. Taking all of the available data on solubilities of various compounds with different fluorochemicals and plotting as volume fraction solubility against reduced temperature, one can obtain a “semiquantitative” fit to zt single curve, with an appreciable scatter of the data points. Considering the general discordance of solubility data, one should not expect much more. There are several other ways of plotting solubility data which may yield information of interest. These data are not accurate enough to add anything of value to the discussions of whether or not the solubility curve has an absolutely flat top,’ since either a discontinuous or smooth curve can be drawn through the points in the critical region with equal confidence. ~~
~
(4) J. B. Hiokman, J . A m . Chem. Soc., 77, 6154 (1955). (5) R. Dunlap, R. Digman, and J. Vreeland, Abstrarts, 124th National Meeting of the American Chemical Society, Chicago, Ill., September, 1953. (6) See, for example, R. G. Bedford and R. D. Dunlap, J . A m . Chem. Soc., 80, 282 (1958). (7) See, for example, 0. K. Rice, J . Chem. Phys., 23, 164 (1955).
SOLUTIONS OF I?LUOROCHEMICALS AND
Figure 4. Log ( X a c / ~- x H C ) us. (IT, for the hexanes with (C4FQ)&.
-
HYDROCARBONS
799’
T)’/x
Cox and Herrington8 have reported that solubility data may be represented by the equations
( T , - T ) = [A’ log ( ~ ’ /-l x.’) and
(To- T )
=
[A” log
(Z”/l
-
x”)
+ B‘I3 + B”]a
in which ’ and “ refer to the two coexisting phases. They found for Fieveral imixtures that A’S A ” . Rudd and Widomg in another examination of one of the systems for which Cox and Herrington found A’ # A” observed that within the experimental precision for pure ethylene glycol monoisobutyl ether with water A’ = A”, but, with impure materials A’ # A ” . Figure 4 shows a plot of log (XI1 - X)vs. (Tc - T)l’* for the hexanes with (GF9)3IS. The ]points for ( T , T)’18less than about 0.7 are not reliable because any reasonable value for the uncertainty in the critical temperature produces an appreciable uncertainty in these values. All of the points, however, fall close to the curves indicating that this method would probably be an excellent one for determining the critical temperature. The plots are good straight lines over a fairly wide temperature range, up to about ( T , - T) = 25”, where the precision of the data becomes bad, and intersect at ( T , - T ) almost zero. The intersection at ( T , -- T)’18< 0.1 corresponds to only a trivial discrepancy in the critical temperature. As in the previous reduced equation (Fig. 2) ad1 of the hexanes fit on the same curve. The slopes of the two curves are not the Sam(’ and cannot be considered the same within experimental error. Whether this disagreement of slopes is t o be considered as characteristic of fluorochemical-hydrocrtrbon mixtures or the results of known impurities in the fluorochemical is unknown. The slopes of the curves are different for the pentanes and heptanes as well: 2.97 and 3.41 flor the pentanes
AE,,,,
k col/mole
Figure 5. Critical solution temperature us. energy of vaporization.
and 2.90 and 3.38 for the heptanes. The intercepts are slightly different, so that the curves are essentially parallel. A plot using volume fractions would make the curves for pentanes, hexanes, and heptanes superSmposable but the slopes for the two phases would still be different, From Fig. 2 or 3 one can see that the ratios of the temperatures of miscibility for two paraffins in (C4F9)3N are independent of concentration and that the temperature of miscibility of an equal volume mixture of fluorochemical and hydrocarbon is very close to the critical solution temperature. Consequently, since it was easier to measure the miscibility temperature of equal volume mixtures, to’, than to determine the critical solution temperatures, these quantities, t,’, were determined for mixtures of hydrocarbons with another fluorochemical for comparison : C8F160from the Minnesota Mining and Mfg. Co., purified in the same manner as (CeFg):,N. These data are listed in Table 11. Figure 5 shows a plot of 1,’ vs. AEv for hexanes with C8F160. A good linear relationship is observed (the dotted line is drawn from Fig. 1 for comparison). The points on this curve represent the hexanes and (8) J. D. Cox and E. F. G . Herrington, Trans. Faraday Soc., 52, 926 (1956). (9) DeF. P. Rudd and B. Widom, J . Chem. Phys., 3 3 , 1816 (1960).
Volume 68, Number .4 A p r i l , 2964
14.S. B. MUNSON
800
Table I1 : Miscibility Temperatures of Equal Volumes of Fluorochemical and Hydrocarbon c-----ta',
(OFdsN
2,2-Dimethylbutane 2,3-Dimethylbutane 2-Methylpentane 3-Methylpentane n-Hexane 2,2,4Trimethylpentane n-Heptane Methylcy clohexane
53.3
OC.----
CaFiaO
-5.0 6.0 11.8 16.0 26.3 18.2 45.6 65.0
methylcyclohexane, which has essentially the same molar volume as the hexanes. Mixtures of hexanes and methylcyclohexane, even 2,2-dimethylbutane and methylcyclohexane whose miscibility temperatures differ by about 70°, fall close to this curve if plotted against average AE". Other data on solubilities of mixtures of paraffins show that that data can be well represented by a solubility curve as a function of reduced temperature as indicated in Fig. 2 and 3. To illustrate the similarity in solution behavior for other noncomplexing fluorochemical-hydrocarbon systems, Fig. 6 shows a plot of t, or t,' for one set of hydrocarbons with a fluorochemical against to' for the same set of hydrocarbons with (C4F9)3N. The data 50 40
-r--
30
20 10
0
- 10 -20
30
Figure 6. Correlation of critical solution temperature of a series of hydrocarbons with different fluorochemicals.
The Journal of Physical Chemistry
for C6F12 are the critical solution temperatures from ref. 5 and the data for C7F1, are from ref. 2 and 4. The similarities in behavior are striking and one can expect similar behavior for solutions of other noncomplexing fluorochemical-hydrocarbon systems. Simple nuclear magnetic resonance studies were undertaken on solutions of fluorochemicals and hydrocarbons to see if any information could be obtained with regard to any anomalous solution behavior. A solvent shift in excess of that due to the differences in magnetic susceptibility has been observed for several systems for gas going into solution and has been considered as due to dispersion forces. The nuclear magnetic resonance data were obtained using a 60 nlc. Varian spectrometer with proton-controlled magnetic field. Benzene sealed into a melting point capillary was used as an external standard. The benzene standard and all samples were degassed by several freezing, evacuating, and melting cycles. The concentration of all samples was -2%. Line separations were determined using a modification of the side-band superposition method. l 2 Instead of superimposing a modulation side band from the standard directly on the peak to be measured, the modulation frequency was adjusted to place the benzene side band adjacent to the unknown peak. This region of the spectrum was then scanned repetitively in both increasing and decreasing field directions. The separation between the side band and the unknown peak was determined from the horizontal scale which had been calibrated previously using the side-band technique. This separation was added to the audio oscillator frequency to give the chemical shift data reported here. The frequency of the oscillator was monitored with a Hewlett-Packard Model 522B frequency counter. The solvent shift measurements were made at room temperature 29 I 3 O . l 3 For 2,2,4-trimethylpentane and 2,2-dimethylbutane, the peak observed was the unresolved --C;€13 peak; for the other compounds the composite -CH2- peak was observed. Table 1x1 shows the measured solvent shifts; the column headed A represents the experimental difference, in parts per million, of the chemical shift for the hydrocarbon in a 2% solution in the fluorocarbon ( C ~ F S ) ~ S minus the value for the pure hydrocarbon with reference to benzene as an external standard. The column (10) A. A. Bothner-By, J. Mol. Spectry., 5 , 62 (1960). (11) A. D. Buckingham, T. Schaefer, and W. F. Schneider, J. Chem. Phya., 32, 1227 (1960). (12) J. T. Arnold and M. G. Packard, ibid., 19, 1608 (1951). (13) The author is grateful t o Mr. R. K. Saunders and Dr. F. C. Stehling for performing the nuclear magnetic resonance experiments and for discussions concerning the interpretations thereof.
SOLUTIONS OF FLUOROCHEMICALS AND HYDROCARBONS
80 1
Table 111: N.m.r. Solvent Shifts for Several Hydrocarbons h a
Vsss
AHm
-K1"
Cy clopentane n-Decane
8.10 7.72
95 195
8.54 12.28
0.629 0.612
-0.047~Oo.022 -0.056 f 0.027
+ O . 171 +O. 195
2,2,4-Trimethylpentane n-Octane
6.85 7.57
166 163
8.40 9.92
0.596 0.595
-0.172f0.033 -0.086 f O . O 1 l
+O. 203
Methylcy clohexane n-Hexane 2,2-Dimethylbutane
7.82 7.27 6.71
128 132 132
8.45 7.54 6.62
0.618 0.567 0.577
-0.025 -0.180 -0.144
+0.215 +O. 168 +O. 182
(C4Fg)aN
5.9
358
13.2
A'
$0.170
0.733b
S. Broersma, J . Chem. Phys., 17, 873 (1949); volume susceptibility, all values X 10-6. Humble Oil and Refining Co., Baytown, Texas.
headed A' represents the solvent shift corrected for differences in magnetic susceptibility. A reasonable estimate of the precision of the data, as shown, indicates that the excess solvent shifts are equal within experimental error. The positive shift is to be expected in going from hydrocarbons to fluorochemicals ( ~ H C> ~ F ;C hydrocarbon intermolecular forces greater than fluorochemical intermolecular forces) as would be predicted from other dataloand is of a reasonable value. The expected trend with 6 was not observed. The solvent shifts (Table 111, column A) show a variation with the difference in volume susceptibility of the fluorocarbon and hydrocarbon, but are not proportional to it. Further, the ratio of A/(Kz - K1), the ratio of
A
b
Determined by Dr. J. T. Richardson,
solvent shift to magnetic susceptibility difference shows no constancy for hydrocarbons of the same volume and no trend with the volume of the hydrocarbon.14 A further comparison was made of the nuclear magnetic resonance spectrum of 2,3-dimethylbutane as pure liquid and in solution in (C4F9)3N a t various temperatures (25-90'). The spectra were essentially identical under all conditions. There is no evidence for any solution anomalies from the nuclear magnetic resonance measurements.
(14) R. E.Glick, D. F. 31, 667 (1969).
Kates, and S. J. Ehrenson, J . Chem. Phvs.,
Volume 68,Number 4 April, 1964
