NOTES
504
630[
Vol. 67
lend some confidence to this number. These experiments give an intermolecular distance of 2.92 A. in water a t room temperature, compared to 2.76 A. in ice. The radius of a water molecule might be expected to be 1/2(2.76) = 1.38 8., while the radius of the "cavity" in which it finds itself in the liquid might well be about 1/2(2.92) = 1.46 8. For comparison, the value of a obtained from 0nsag:r's approximation,12 47raa/3 = 1/N, a t 2 5 O , is 1.93 A. This assumption states that the molecules fill all space, and its use with E = n2reduces eq. 1 to the Lorens-Lorentz equation
'\\\
n2
6
2
7
'
k
' ,340 I '
- 2. 2nl + 1
I
'
'
342
'
'
'
344 I '
2n'
Fig. 1.-Application of Bottcher's method to Tilton and Taylor's data a t X = 5460.7 d. from 0 t o 60'.
- 1 n/r
47r _ _
..
The curves shown in Fig. 2 illustrate the variation 9f polarizability with temperature a t X = 5460.7 A. The curves for a = 1.44 and 1.36 8.were calculated from eq. 1. The curve for CY from the Lorens-Lorentz equation was obtained from eq. 2. h'ote that CY is very sensitive to a, while the temperature dependence is not greatly different for the three cases. It is thought that the curve for a = 1.44 A. gives the best representation of the data, although the assumption of a constant value of a over the interval may not be valid. Values of a a t 25' for a = 1.44 A.are shown in Table I1 for wave lengths of interest. TABLE I1 POLARIZABILITY OF WATERAT 25' x (AJ 4358.3 5460.7 5769.6
a X IO24 (ml./molecule)
1.363018 1.345611 1.342301
x (A.) 5780 5875.6 5892.6
a X 1024 (nil./moleoule)
1.342197 1.a41264 1.341102
The present analysis, unlike the earlier onej2does not indicate of itself whether the Bottcher or LorenzLorentz method is more appropriate, although theoretical con~iderationsl~ seem to favor Bottcher's equation. Acknowledgment.-Support of the National Science Foundation under Grant No. G 15655 is gratefully acknowledged.
11, \- 1 ,
0
IO
,
y?z13E;
20
30
40
50
60
T ("C.).
TABLE I ESTIMATE OF CAVITY RADIUS
- t i ("C.) 60-55 55-50 50-45
a'
(A.)
1.44 1.43 1.40
SOLUBILITIES OF LIQUIDS I N 1-HYDROPERFLUOROHEPTBSE BY J. 0. KOSECNY AND C. H. DEAL Shell Development Company, Eneryvzlle, Calzfornza
Fig. ?.-Variation of polarizability with temperature a t A = 5460.7 A. for three cases. Values a t 25' are shown on each curve. Kote that the ordinate is greatly compressed betm-een curves.
ta
(12) L. Onsager, J . Am. Chem. Soc., 68, 1486 (1936). (13) W. F. Brown, Jr., "Handbuoh der Physik," Band XVII, SpringerVerlag, Berlin, 1956,pp. 1-154.
t2
- t l (".) 45-40 40-35 35-30
a'
(A,)
1.37 1.33 1.29
be most strongly suggested, although this conclusion might be modified if higher temperature data were available. The small angle X-ray scattering (10) J. Morgan and B. E. Warren, J. Chem. Phyq., 6,666 (1938). (11) G. W. Brady and W. J. Romanow, ibid.. 32, 306 (1960).
Received August 6 , 1068
In their investigation of fluorocarbon solutions, McLaughlin and Scott' determined the solubility of phenanthrene in 1-hydroperfluoroheptane and 13dihydroperfluorooctane and found it to be in fair agreement with the regular solution theory. This fact, together with the abnormally low solubility of phenanthrene in the perfluorocarbon-like solvents, perfluorotributylamine and perfluoropropylpyran,2 was first taken to indicate that the hydrocarbon-fluorocarbon "anomaly" disappears when one or two hydrogens are (1) E. P. McLaughlin and R. L. Scott, J . Phys. Chem., 60, 6741 (1956). (2) E. P. RlcLaughlin and R. L. Scott, J . Am. Chem. Sac., 76, 5276 (1954).
505
NOTES
Feb., 1963 introduced into a perfluorocarbon molecule. However, Scott himself later pointed out that data on CH4CF3H and other systems do not support this idea3 and that there is considerable evidence for hydrogen bonding in CF3H-CF2H2 mixture^.^ I n view of these results it appeared worthwhile to examine other binary solutions containing hydrofluorocarbons. In the present work we report liquid-liquid miscibilities and critical solution temperatures for mixtures of C7F15H with paraffins, alkylbenzenes, and a few polar solvents. The behavior of C7F15H in solution is compared with that of perfluorocarbons. Experimental 1-Hydroperfluoroheptane was prepared by the decarboxylation of perfluoro-n-octanoic acid as described by LaZerte and his coworkers.6 The product waa dried over phosphorus pentoxide and distilled on a 10 plate microhelipod column a t a reflux ratio of 5 : l . The heart, cut, boiling a t 95' a t 751 mm. and having a density of 1.718 a t 25", was used for the investigation. The proton magnetic resonance spectrum obtained in chloroform solution had the expected multiplet structure, namely a triplet of triplets. The chemical shift was 5.91 p.p.m. from tetramethylsilane and the spin-spin coupling constants were 52.2 and 5.1 C.P.S. to the a- and P-CF, groups, respectively. Sitromethane was distilled in the same manner. The product boiled a t 101" a t 762 mm. Acetonitrile ("spectroscopic grade") was used as supplied. The hydrocarbons, of the purest grade available, also were used as supplied. Their purity, determined by g.1.c. analysis, was higher than 99.6% with the exception of n-decane (99.1%), n-dodecane (98.3y0), and n-tetradecane (92.301,). The solubilities were determined by the cloud point method in sealed Pyrex tubes. All values reported are an average of a t least three determinations agreeing within 0.2' or better. The volume fractions of the components a t 25" were calculated irom their weights and densities.
C14H30
120
100
9
80
$ 2 60
40
20
0
0
Fig. 1.-The
0.2
0.4 0.6 Volume fraction of hydrocarbon.
0.8
1.0
mutual solubility of C7FljH and normal paraffins.
Results &d Discussion The liquid-liquid solubilities of C7F15H and paraffins, alkylbenzenes, and the polar liquids, expressed in volume fractions of these components, are shown graphically in Figs. 1, 2, and 3, respectively. The critical solution temperatures t, are summarized in Table I. TABLE I COMPARISON OF SOLUBILITY PARAMETERS BASEDON HEATSO F VAPORIZATION AND SOLUTION DATA Liquid
to,
C7H16 CsHu CioHzz
81,
OC.
31.1 48.0 78.2 102.0 127.0 41.3 34.0 67.1 99.7 43.8
C12H26
C14H30 CsHe CsHbCHs C&C3H7 CHgN02 CH3CN
80
ca1.'/2 ~ m . ~ /81*, 2 ca1.'/2 cm.-'/2
8.9 8.9 8.9 8.9 8.9 9.2 9.1 9.1
7.5 7.6 7.7 7.8 7.9 9.2 8.9 8.7 12.6 11.9
The values in the third column of Table I are the solubility parameters of the hydrocarbons a1 calculated from the approximate equation3 4RTc
= ('Vi
+ 'V2)(81 -
(1) by substituting the experimental values of the critical 82)'
(3) R. L. Scott, J . Phys. Chenz., 62, 136 (1955). (4) A t the time of this writing a detailed investigation of hydrogen bonding in the system CTFiaH-CHsCOCHs has been reported in the literature (D. L. Anderson, R. A. Smith, C. B. Myers, S. K. Alley, a n d R. L. Scott, zbzd., 66, 621 (1962)). (5) J. D. LaZerte, L. J. Hals, Soc.. 76, 4525 (1953).
T. S. Reid, a n d G. H. Smit, J . Am. Chenz.
0
Fig. 2.-The
0.2 0.4 0.6 Volume fraction of hydrocarbon.
0.8
1.0
mutual solubility of C?R,H and alkylbenxenes.
solution temperature T , (A.)) the molar volumes of the components 'V (at 25') and the solubility parameter of C,HbH,l 82 = 6.3 a t 25'. The values of 81*, given in
NOTES
506
I-oi. G i MOLECULAR CORIPLEXES OF NETHOXYBESZESES BYARXOLD ZWEIG
Chemacal Research Department Central Research Dzurszon, Amerzcan Cyanamzd Company Stamford Conn.
100 -
Receawd August 6 , 1962
80
-
9
60-
d
-
0
Fig. 3.-The
I
CH,CN
0.2 0.4 0.6 Volume traction of polar component.
0.8
1.0
mutual solubility of CiFljH and polar liquids.
the last column, are the solubility parameters of the hydrocarbons and polar liquids a t 2.5’ based on the heats of vaporization and taken from Hildebrand and Scott.6 (The values for decane, dodecane, and tetradecane were interpolated by means of the linear plot of 61” us. the reciprocal carbon number of the paraffins n, the limiting value for l/n = 0 being 8.45.) Inspection of the results shows that the behavior of C7F16H in paraffins and alkylbenzenes does not differ markedly from that of perfluorocarbons. The paraffin solutions display the usual fluorocarbon “anomaly,” 61 being appreciably larger than &*. In alkylbenzene solutions the values of 6lagree quite well with the values of al*, which is also true in alkylbenzene-perfluorocarbon mixtures. However, in solutions containing a polar component marked differences become apparent. The critical solution temperatures of the systems C7FI5H-CH3KO2and C7F15-CH3CN are surprisingly low considering the solubility parameters of the liquids. Moreover, further experiments with polar solvents have shown that C7F15H-in contrast to perfluoropropylpyran-is completely miscible a t room temperature with methanol, ethanol, and acetone. On the basis of these results we may conclude that the behavior of C7FlBHin paraffins and alkylbenzeiies is dominated by the fluorocarbon chain. In polar solvents, however, specific interactions involving the highly polarized CH group of C7FI5Hstrongly influence or even dominate solution behavior. Acknowledgments.-The authors are indebted to Dr. C. A. Reilly for the nuclear magnetic resonance measurements. (6) J. H . Hildebrand and R L. Scott, “The Solubiiity of Non-Electrolytes,” Reinhold Publ. Corp., Neir I‘ork, S . Y., 1955.
Recent success in correlation of the orientation effects in the Birch reduction of methoxybenzenes with molecular orbital theory’ has suggested the examination of some other possible correlations of properties of methoxybenzenes with the simple Huckel molecular orbital (HAIO) scheme.2 AIolecular r-complexes of organic compounds have been known since 1858,3although their spectral propertied were not fully appreciated until the work ot Benesi and Hildebrand.4 hIulliken6 worked out a detailed quantum-mechanical theory for complexes where one molecule acts as an electron donor and another as an acceptor. In this theory, the charge-transfer absorption is ascribed to a transition from a ground state which is mostly an uncharged aggregate, DA, to an excited state which is mostly an ion-pair aggregate D+-A-. Dewar and Lepley6have shown that if the interactions between donor and acceptor are weak, that is, the ground state has little charge-transfer character, the transition energy, AEo, for the first charge transfer band for a series of donors with a given acceptor should be a linear function of the energies of the highest filled molecular orbitals (HFAIO) of the donors, assuming that the ionization potential is equal to the energy of the HFMO. Similar results with ionization potentials had been reached by Bhattacharya and Basu.’ Examination of charge-transfer absorption of catacondensed hydrocarbons with i ~ d i n e ,1,3,5-trinitro~ benzene,6 and tetracyanoethylenes have confirmed the linear relationship of AEo with the ionization potential and the HFMO of the hydrocarbon donor. Results and Discussion The major results of this investigation are summarized in Table I. The energies of the HMO’s mere obtained for each of the methoxybenzenes by assuming a planar delocalized system that includes six electrons from the benzene ring and two electrons from each attached oxygen atom. The carbon-oxygen exchange integral was taken as O.8pocand the oxygen 2pcc.2 coulomb integral was taken as CY, The secular determinant was factored when possible and the eigenvalues of the residual determinant were determined by an iterative procedure on a Burroughs 205 computer. Plotting the frequency of the chargetransfer absorption maximum against the HFhIO, as shown in Fig. 1, gave a very good correlation for seven of the methoxybenzenes and scattered points for the five others. From the slope, the indicated value of p is - 3 . 3 e.v. or -76 kcal./mole. Extinction coefficients and equilibrium constants
+
(1) H E. Zimmerman, Tetrahedron, 16, 169 (1961). (2) A . Streitwieser, “iUolecular Orbital Theory,” John J T h y & Sons, Inc , New York, 3’ Y., 1961. (3) J. V. Fritsche. J. prakt Chem , 73, [ l ] 282 (1858). (4) H A. Benesi and J H. Hildebrand, J . Am Chem Soc , 71, 2703 (1949). (j) R S Mulliken, z b z d , 74, 811 (1952). (6) AI J S Dewar a n d A. R Lepley zbzd , 83, 4560 (1961) (7) R Bhattachar,a and S Basu Trans Faraday Soc , 64, 1286 (1‘358) ( 8 ) 11 J. S D e u a r and H Rogers, J A m Chem S o c , 84, 393 (1962)
