May, 1959
INTERMOLECULAR FORCES INVOLVING CHLOROFLUOROCARBONS
703
INTERMOLECULAR FORCES INVOLVING CHLOROFLUOROCARBONS BY E. BRIANSMITH, JOHN WALKLEY AND JOEL H. HILDEBRAND Contribution from the Department oj Chemistry, University of Calijornia, Berkeley, California Received October 8. 1968
Solvent properties of chlorofluorocarbons have been studied in order to fill the long interval between hydrocarbons and fluorocarbons. The solubility of iodine in 2,2,3-C@F7 is 0.1510 mole % at 25". Between 0 and 35" it accords with the equation, log x z = 13.050 log T - 35.075 where xz is mole fraction. The relation of entropy of solution to solubility previously found is again followed. In mixtures of 8.93 and 12.5 volume % of cc&with C7F16, the solubilities of 1 2 a t 25" are respectively. Partial molal volumes of Brz and 1 2 have been determined in solvents ranging 0.0388 and 0.0545 mole 70, from C7Flr J o CHBr3, including 3 chlorofluorocarbons, and in two mixed solvents, C7Fl&C14 and c-C4C1?Fe-CC14. The data are evidence that contacts between halogen molecules and chlorine atoms in the solvents have much higher than random probability.
Fluorocarbons, by greatly extending the available range of internal forces,' have been of great service in providing stringent tests for theories of non-polar liquids and solution. However, the extension has been so great as to leave a considerable gap in solvent power for iodine between f-heptane and the poorest hydrocarbon solvent investigated, 2,2-dimethylbutane. The mole fraction of iodine in the latter at 25" is 26.3 times its value in the former. We turned to chlorofluorocarbons to fill in this interval. Shinoda and HildebrandZa determined the solubility of iddine in cyclodichlorohexafluorobutane, c-C4ClzF6, in (C3F7COOCHz)4CJ and, to extend the range a t the other end, in CHBr3. We determined, also, the partial molal volume of iodine in a number of solventslsb with t'he surprising result that in CClzFCCIFz it is 67.7 cc., close to the value in many non-polar solvents, but rises to 81.2 cc. in C - C ~ C ~ ~Glew F ~ . and Hildebrand" had previously obtained the .Jalue of 100 cc. in perfluoroheptane. Because expansion contributes to the entropy of solution we have considered it important to look more closely into this factor.
The Solubility of Iodine in Mixtures of ClF16 and CC4.This was determined at 25' in two mixtures of different composition. The results are given in Table 11.
H. A. Benesi and J. H. Hildebrand, J . Am. Chsm. Soc., 70,
The data are plotted in Fig. 1 against the solubility parameters of the solvents, &. Noteworthy is the slow increase in Vz in the case of both halogens,
TABLE I1 SOLUBILITY OF IODINE IN MIXTURESOF C7F16 WITH cch MOLE% AT 25" Vol. % CCla 8.93 12.5 100x2 0.0388 0.0545 60 from xz 6.0 6.3 a0 from eq. 1 6.1 6.2 60 represents the solubility parameter of the mixed solvent calculated f i s t , from 52, second, by the equation
an =
+ b3a3
(1)
given by Hildebrand and Scott.3 In this 4 denotes volume fraction, and the subscripts 1 and 3, refer to C7F16 and CCl,; 61 = 6.0 and 63 = 8.6. The agreement is quite satisfactory. Partial Molal Volumes.-We have determined the partial molal volume of iodine in 2,2,3-C4C13F7,to .add to .those already at hand,2b but because the low solubility of iodine in fluorocarbons and fluorochlorocarbons makes it difficult to achieve the desired accuracy, we turned to bromine in order more easily to cover the solvents in this region. We used the rapid method employed in the earlier work, with all necessary precautions. The solvents were of high purity, dried as necessary over P201, distilled and run through a Solubility of Iodine in 2,2,3-Trichloroheptafluorobutane, column of silica gel. Bromine was dried with P2O6and disC4CI3F?.-We were so fortunate as to obtain a sample of this tilled under reduced pressure. The mole fraction of halogen compound from Dr. T. M. Reed, and we determined the never exceeded 0.004, therefore the values given are virtually solubility of iodine therein in order to supplement other those of the partial molal volume of the halogen, 82, a t infinite solubility data. The technique adopted was essentially dilution. Duplicate determinations agreed within 0.4 cc. that described by Glew and Hildebrand,Ic but the apparatus mole-'. The results are given in Table 111 together with was modified for use with a smaller volume of solvent, 4 4 our earlier values for iodine. cc. The solution was stirred for 20 hours in the presence of TABLE I11 excess iodine; temperatures were controlled to ztO.002'. The concentration of iodine was determined spectrophoto- PARTIAL MOLALVOLUMES OF BROMINE A N D IODINE AT 25" metrically and checked by titration. The solutions obey AND HIGHDILUTION, cc. MOLE-^ Beer's law. The results are given in Table I. The figures Solvent 61 Bra I2 in the last column have been calculated from the equation log z2 = 13.050 log T -35.111. The entropy of solution n-C7Fla 6.0 71.7 100 of solid iodine is 25.88. This places it on thd straight line c-C~H~ICF~ 6.1 71.2 ... of Fig. 1 of ref. 2a, where the entropy of solut,ion is plotted c-CaClzFe 6.8 64.3 81.2 against the solubility expressed as - R In 2 2 . The solu... 78.6 2,2,3-C4ClaF, 6.9 bility parameter of CICl3F7 calculated from its molal volume, 166 cc., and the solubility parameter of iodine a t 25" is 6.9, in CClzFCClFz 7.5 56.9 67.7 excellent agreement with the value 6.92, obtained by Reed SiCla 7.6 54.9 67.1 (private communication) from its heat of vaporization. CCl4 8.6 54. I 66.7 TABLE I CHCls 9.0 53.7 65.6 csz 10.0 49 9 62.3 SOLUBILITY OF IODINE IN C4C13F7, MOLEyo, 100x2 CHBr3 10.5 49.3 60.8 t, "C. 0.00 14.80 19.36 25.00 35.30 ... Brz 11.5 51.5 Obsd. 0.0473 0.0961 0.1184 0.1510 0.2331 12 14.1 ... 59.0 Calcd. 0.0479 0.0959 0.1181 0.1509 0.2347 4
(1) (a)
3978 (1948); (b) J. H. Hildebrand, H. A. Benesi and L. M. Mower, ibid., 72, 1017 (1950); (c) D. N. Glew and J. H. Hildebrand, THIS JOURNAL, 60, 616 (1956). (3) J. H. Hildebrand and R. L. Scott, "Solubility of Non-eleotro(2) (a) K. Shinoda and Joel H. Hildebrand, ibid., 62, 292 (1958); (b) 62, 293 (1958). lytes," Reinhold Publ. Corp., 1950,eq. 11, p. 201.
E.BRIANSMITH,JOHN WALKLEY AND JOELH. HILDEBRAND
704
Vol. 63
of a pure fluorocarbon. However, in C4C13F7and C4CIzFs, chloride atoms, on the same basis, occupy 49 and 38%, respectively, and the halogen molecules are more frequently forced to be under the much weaker attract.ive potential of fluoride atoms. We evidently have in these solutions molecules for which it is appropriate to recognize different attractive potentials in different parts of their surfaces, as was done years ago by Langmuir.6 He assigned parameters for attractive potentials between molecules and groups present in molecules, and invoked the Boltzmann factor to give the probabilities of contacts of various energies. We undertook to throw additional light upon this matter by determining the partial molal volume of bromine in two mixed solvents, permitting the ratio of chloride to fluoride areas to be varied continuously. We used two mixtures, one, C7FI6 with cc1.1, which has a miscibility gap, the other, c-C4ClzF6 with CCL, whose composition can be varied through I I the entire range. The results are plotted in Fig. 2 against mole fraction of CC14 in the solvent. The striking feature is the rapid drop in ;Jz with increasing mole fraction of CC14. It is much steeper if plotted against volume fraction. Our interpretation is: MISCIBILITY GAP- I n the mixture C7F16-CC14,when only a little CC14 is present, it can have but little effect in lowering the value of ’i7, for Brp, because contacts bew tween them are rare, but as more CC14 is added, they increase rapidly, permitting Vz to drop rapidly 50 toward its value in pure Tq. Therefore V 2 is less 0.6 0.8 1.0 0 0.2 0.4 than an additive function of its values in the two MOLE FRACTION OF C C 1 4 I N M I X T U R E . pure solvents. Fig. 2. Tn the mixture, C4Cl2F6-CCL,in which there is no as 6, decreases all the way from CHBr3 to CCl2F- gap, V2 again drops rapidly toward its value in pure CC1F2,followed by rapid rise in both cases from the CC14. This curve is not convex upward probably because chloride atoms are already present in large latter solvent to C7F16. A smooth increase was to be expected on the basis excessin the pure chlorofluorocarbon, and the further drop is roughly proportional to their increase. It of the relation4 seems more important to present the evidencefor the V9 - vz = PORTIn y~ (2) concept qualitatively than, a t this stage at least, to where Po is the compressibility of the solvent and yz make a quantitative calculation by Langmuir’s the activity coefficient of the solute. The points method. We remark only that these ternary soluso calculated are shown in Fig. 1. tions are obviously not strictly regular (we use The sharp break in the experimental points, how- “strictly” here in its proper, adjectival sense) but ever, calls for further attention. We explain it as that preferred contacts occur with higher than follows. purely random probability. The major part of the surface area of a molecule We are determining values of (bP/bT)vfor these of CClzFCCIFzconsists of chloride atoms. On the solvents in order to study the effects of expansion basis of a ratio of 1.33 for the van der Waals radii of upon the entropy of solution. chloride to fluoride, their cross-sections are in the We gratefully acknowledge the kindness of Dr. ratio of 1.8, therefore approximately 64% of the T. M. Reed in supplying the C4C13F7, and the surface area of this compound consists of chloride, Atomic Energy Commission and National Science and its behavior toward bromine and iodine with Foundation for the grants that have supported the respect to both solvent power and partial molal work. volume is much closer to that of CC1, than to that (5) I. Langmuir, “Colloid Symposium Monograph,” 1925, p. 48. See also, ref. 4, pp. 162-166.
(4) Cf. ref. 3. p. 141
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