COMMUNICATIONS TO THE EDITOR
1621
The electric moment of the solute is sensitive to the value of 1.11 used. The most current vapor-phase value listed by l\lcClellanll is used as well as the moment calculated from the Onsager equation for the pure dipole liquide3 The first of the two entries for the electric moment in Table I is obtained using the vaporphase moment of the solvent while the second is from the pure liquid value. Presumably the best value of p1 lies somewhere between the two. There is good agreement in the solute dipole moment a t the two temperatures. Examination of the literature shows that the electric moment of CopyZC1, is greater than that of the corresponding zinc complex which is consistent with the results of other analogous tetrahedral cobalt and zinc complexes obtained in nonpolar solvents.9 This agreement suggests that the electric moment of many other coordination compounds insoluble in nonpolar solvents may be determined in nitromethane.
Aclcnowledgments. The author wishes to thank C. P. Kash, in whose laboratory the dielectric constant measurements were made, and J. W. Tam for the stimulating discussions during the course of this work. The author acknowledges support of this work through Grant (PR F-2177-A3,5) from the Petroleum Research Fund of the American Chemical Society. (11) A . L. McClellan, “Tables of Experimental Dipole hloments,”
IT,H. Freeman and Co., San Francisco, Calif., 1963. DEPARTMENT OF CHEMISTRY UNIVERSITY O F CALIFORNIP, DAVIS,CALIFORNIA
ROBERTD , FARINA
RECEIVED JANUARY 14, 1969
Comment on “The Thermal Dissociation of Oxygen Difluoride.
I. Incident
Shock Waves”
Sir: Confusion has arisen from the manner in which we reported our bimolecular rate constant for the thermal decomposition of OF2.l I n the text, we pointed out that the preexponential was divided by TIi’ but this was not apparent from the dimensions of the rate expression reported. Blauer and Solomon,2 in comparing our data to more recent shock tube and static reactor data, reported that our data differed significantly from the other data. This discrepancy was in the main due to the confusion cited above in interpreting our expression. In addition, there is a calculating error in our data which is apparent in Table I. I n converting the experimentally obtained kpseudo-first-order constants, the calculated values of kbirnolecular are all off by an order of magnitude. The data given in the table are sufficient
to verify this discrepancy. With these modifications, our value of k , on the same basis as Blauer and Solomon, is
k,
-
1015*2&0’2 exp( -32’200
* ‘O0)
RT
cc/mol
This result fits well into the summary of reported OF2 data in Figure 3 of the paper by Solomon, Blauer, and Jaye. These errors were pointed out to us by Dr. -1Iing Chang Lin of the Department of Chemistry, Cornel1 University. (1) L. Dauerman, G. Salser, and Y. Tajima, J . Phys. Chem., 71, 3999 (1967). (2) J. A. Blauer and W. C. Solomon, ibid., 72, 2309 (1968). (3) W. C. Solomon, J. A. Blauer, and F. C. Jaye, ibid., 72, 2314 (1968).
DEPARTMENT OF CHEMICAL ENGINEERING NEWYORICUNIVERSITY NEWYORIC,N E WYORK
L. DAUERMAN G. E. SALSER
LOCICHEED PROPULSION REDLASDS, CALIFORNIA
Y. A. TAJIMA
RECEIVED JANUARY 27, 1969
Effect of Urea and Ethanol on the Viscosities of Several Aqueous Electrolyte Solutions
Sir: Considerable controversy exists as to the effect of urea on the structure of water. Two recent communications’,2 attempt to clarify this controversy, the first‘ presenting evidence that urea is a structure breaker and the second2 that urea has essentially no net effect on the structure. The addition of small amounts of urea3 or ethanol4 to water is known to increase the viscosity of the solution. The temperature dependences are different, in that the relative viscosities of the ethanol solutions decrease while the relative viscosities of the urea solutions increase with increasing temperature. Although it is difficult to obtain much information concerning structural effects from the absolute magnitude of the viscosity, it is possible to obtain more information from the temperature dependence because structural effects involving hydrogen bonding are relatively weak. The considerably greater viscosity of dilute ethanol solutions indicates that the alcohol is entering into some kind of hydrogen bonding with the (1) D. V. Beauregard and R. E. Barrett, J . Chem. Phys., 49, 5241 (1968), (2) S. Subramanian, D. Balasubramanian, and J. C. Ahluwalia, J . Phys. Chem., 73, 266 (1969). (3) H. M. Chadwell and G. Asmes, J . Amer. Chem. SOC.,52, 3507 (1930). (4) R. L. Kay, G. P. Cunningham, and D. F. Evans, “HydrogenBonded Solvent Systems; Proceedings,” A. K. Covington and P. Jones, Ed., Taylor and Francis Ltd., London, 1968, p 255.
Volume ‘73, Number 6
Mag 1969
