602
COMMUNICATIONS TO THE
EDITOR
Table I
:-
’
[ : : , I
BE111_:,SE,F
0
- CYCLOHEXANE,
A
- MUTUAL (
1.2
,
D~FFU,SION,
SELF DIFFUSION
7
DIFFUSION
0.1
1.88 1.85 1.79 1.72 1.66 1.63 1.64 1.68 1.78 1.90 2.09
1.88 1.86 1.83 1.79 1.77 1.78 1.80 1.83 1.90 1.98 2.09
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0
R O D W I N , HARPST 8 LYONS’
I. 0
0
1.88 1.86 1.83 1.81 1.79 1.79 1.80 1.84 1.90 1.99 2.10
0.3 0.4 0.5 0.6 0:7 0.0 MOLE FRACTION CYCLOHEXANE
0.2
0.9
1.0
Figure 1. The concentration dependence of self-diffusion coefficients in benzene-cyclohexane solutions, 25”.
in terms of the “density defect.” These quantities are roughly proportional to one another although the “density defect” is about an order of magnitude smaller than the (‘viscosityd e f e ~ t . ” ~
Acknowledgment. We are indebted to D. C. Douglass and B. Ottar for stimulating discussions of this work. absolute accuracy of better than 5% with the pure liquids but the errors can be expected to be greater in solutions dilute in protons. This is particularly true in dilute solutions of CaH6 in CBDIP. The results are shown in Figure 1 together with the mutual diffusion results of Rodwin, Harpst, and Lyons.’ It is clear that the spin-echo results (which are absolute and have not been adjusted to fit the tracer data) exhibit the correct limiting behavior, i.e. lim D = D,
(6) R. J. Bearman, J . Phys. Chem., 65, 1961 (1961). (7) J. D.Birkett and P. A. Lyons, ibid., 69, 2782 (1965). (8) J. A. Dixon and W. Schiessler, ibid., 58, 430 (1954).
(9) NOTEADDED IN PROOF. Rodwin, Harpst, and Lyons’ do not suggest that the parameters they have deduced ( L e . , those corresponding to our Db and D,) are self-diffusion coefficients. This interpretation has been made by Bearman.6 Dr. Bearman has called our attention to a recent tracer diffusion study of this system by D. A. Collins and H. Watts, Australian J. Chem., 17, 516 (1964). We have benefited greatly from comments contributed by Professor P. A. Lyons and Professor R. J. Bearman. BELLTELEPHONE LABORATORIES HILL, N i w JERSEY MURRAY
zo+o
and
DAVIDW. MCCALL ERNESTW. ANDERSON
RECEIVED DECEMBER 14, 1965
Table I shows a comparison of experimental D values with D values calculated from eq 1. It is clear that the agreement is poor. Bearmad has discussed eq 1 from a theoretical viewpoint and casts some doubt upon the general validity of the relation. Thus, the discrepancies observed are not too surprising. A simple equation that describes the data well is
D
= (ZcDb
+
ZbDc)P/(ZbVb
+
Z c d
(4)
where 7 is the solution viscosity. Table I also contains D values computed from eq 4. The agreement is well within experimental error. We have not corrected the viscosities for the isotope effect.'^^ Equation 4 can be regarded as the simplest mixture formula corrected for the “viscosity defect,” [l r1/(xCrlc4- Z b v b ) ] . The discussion could also be phrased The Journal of Phyeical Chemietrl
Heats of Mixing of Benzene with Hexafluorobenzene, Pentafluorobenzene, and 1,2,4,5-Tetrafluorobenzene’
Sir: As part of our research program on fluorocarbon solutions, we have measured the molar heat of mixing (ie., the molar excess enthalpy RE)for each of the
+
+
+
Systems CeHa CeFa, CsHa CaFsH, and CeH6 1,2,4,5-tetrafluorobenxeneover a wide mole fraction
(1). Contribution No. 1905 from the Department of Chemistry, University of California, Los Angeles, Calif. This work was supported in part by the U. 8. Atomic Energy Commission and in part by the University Grants Committee, New Zealand (Postgraduate Scholarship in Science to D. V. F.).
COMMUNICATIONS TO THE EDITOR
400
603
t
IO0
80 60 40
,OOt
100
20
0
+
0
-aJ
-2 0
-100
-40
I
3
-0
\\
-60< W
1 1 -8 0 -I 0 0
-120
-140
-600 0.0
0.2
0.4
0.6
0.8
1.0
x2
Figure 1. T h e excess enthalpies of CeHa at 25' (0)and 45' (A), C& -I- C&H at 25' (0)and 42' and Cd& f 1,2,4,5-tetrafluorobenzeneat 25' (+) and 39 ( X ) ; 2 2 is the mole fraction of the fluorochemical in each case.
(o),
range and a t two temperatures (Figure 1). The calorimeter used is similar to that described by Larkin and McGlashan2 and is described elsewhereO3 The C6H6 CBFP, results near x2 = 0 and those for CeH6 C6FsHnear z2 = 1 have been carefully checked and the sign change in the heat of mixing has been definitely established in each case. Since runs a t these composi-, tions jnvolve no electrical compensation, the sign of the temperature change unequivocally determines the sign of RE. The S-shaped curves found in the system l-hydron-perfluoroheptane acetone4 were interpreted in terms of the combination of a symmetric exothermic "chemical" interaction arising from the formation of a 1: 1 hydrogen-bonded complex and a skewed endothermic "physical" interaction arising from the mixing of hydrocarbon and fluorocarbon groups. The former contribution was obtained independently from nmr data and the latter inferred from the system CTF16 acetone. We believe the S-shaped curves for C6H6 C6F6 and C6H6 CeFsH may be similarly explained.
+
There is strong evidence suggesting charge-transfer complex formation between C6H6, acting as donor, and C6F6, acting as acceptor. The freezing point diagrams shows the formation of a 1: 1 solid complex. The predominantly exothermic results that we obtain for this system can be understood in terms of such a complex; the positive temperature dependence of BE must be due to a decrease in complex formation with increasing temperature. The more endothermic results for C6H6 C6F6Hpresumably arise from a smaller exothermic contribution; i.e., C6FsH is a poorer acceptor than C6F6 and forms a weaker complex. Any temperature dependence in this case is small, less than the experimental error. With C ~ H B4- 1,2,4,5-tetrafluorobenzene, the further increase in RE again reflects further decrease in complex formation. The negative temperature dependence for this system corresponds to the normal behavior6 of systems with positive RE; indeed, there is no clear evidence for any complex formation in this case. Additional measurements with other fluorine sub~, CaHsF) are stituted benzenes (e.g., C B H ~ FCdU?", underway. (2) J. A. Larkin and M. L. McGlashan, J. Chem. Soc., 3425 (1961). (3) J. A. Larkin, D. V. Fenby, T. S. Gilman, and R. L. Scott, to be published. (4) D. L. Anderson, R. A. Smith, D. B. ,Myers, S. K. Alley, A. G. Williamson, and R. L. Scott, J.Phys. Chem., 6 6 , 621 (1962). (5) F.L.Swinton, private communication. (6) M. L. McGlashan, Pure Appl. Chem., 8 , 157 (1964).
DEPARTMENT OF CHEMISTRY UNIVERSITY OF CALIFORNIA Los ANGELES,CALIFORNIA90024
DAVIDV. FENBY IANA. MCLURE ROBERT L. SCOTT
RECEIVED DECEMBER 27, 1965
+
+
+
+ +
Bond Lengths in Iron Pentacarbonyl
Sir: On the basis of an electron diffraction study of gaseous iron pentacarbonyl, Davis and H. P. Hanson recently reported' that the axial Fe-C bonds were shorter than the trigonal by 0.045 A, the values being 1.797 and 1.842 A, respectively. They stated that the existence of the shorter axial bonds was corroborated by X-ray crystallographic studies of A. W. Hansoq2 whose results gave axial bonds of 1.785 and 1.807 A and trigonal bonds of 1.827, 1.827, and 1.837 A. Davis and H. P. Hanson were apparently unaware that the crystal structure refinement by A. W. Hanson (1) M.I. Davis and H. P. Hanson, J . Phys. Chem., 69,3405 (1965). (2) A. W.Hanson, Acta Cryst., 15, 930 (1962).
Volume 70, Number 8 February 1966
