J . Phys. Chem. 1989, 93, 6880-6881
6880
concentrations of the two components in the stationary phase at time t and position z (eq 1 and 2) axial dispersion coefficients of the two components (eq 1 and 2) phase ratio (eq 1) height equivalent to a theoretical plate (eq I O ) space increment in the integration (eq 8) lumped mass-transfer coefficients of the two components (eq 3 and 4) column capacity factor of component i at infinite dilution (eq 10)
Qs,,, Qs,2 concentrations of the two components in the stationary
phase, at equilibrium between the two phases of the chromatographic system, with concentrations C,,,, and Cm,2of these components in the mobile phase (eq 3 and 4) time (eq 1) duration of the rectangular injection pulse (eq 7 ) mobile phase velocity (eq 1) abscissa along the column (eq 1) time increment (eq 8)
t tP U Z T
COMMENTS Comment on the Thermal DecomposRlon of Anisole and the Heat of Formation of the Phenoxy Radlcal Sir: In a recent paper in this Journal on the thermal decomposition of anisole' a rate constant was reported for the unimolecular dissociation C~HSOCH -*~C 6 H S 0
CH3
(1)
The authors apparently overlooked an earlier measurement of this rate constant made under different conditiom2 In the earlier study anisole was pyrolyzed in a static system over the temperature range 720-795 K in the presence of a large excess of toluene to act as a radical trap, and from a kinetic analysis it was shown that the rate of formation of methane could be equated to the rate of reaction 1 . The rate constant so obtained gave a linear Arrhenius plot which was described by the equation 58000 f 2000 log kl (s-I) = 13.7 f 0.3 2.3RT While these two sets of Arrhenius parameters differ considerably, the rate constants themselves from the two studies can be fitted to a common Arrhenius plot, shown in Figure 1, and the much wider temperature range of the combined data will provide more reliable Arrhenius parameters. In the earlier work, it was recognized that the measured frequency factor was probably too low and greater weight should be given to the data from the lower temperature range of these experiments. With this in mind, the combined data are best represented by the dashed line in Figure I , which corresponds to 61000 f 2000 log kl (s-') = 14.6 0.2 2.3RT
*
The largest deviation of this rate constant from either measurement is 50%. The uncertainty in the activation energy of f 2 kcal mol-' represents a realistic estimate of the effects of possible systematic errors in the two systems. Using this activation energy of 61 kcal mol-', one may calculate the heat of formation of the phenoxy radical, following the procedures used by Colussi, Zabel, and Benson' in their calculation
0022-3654/89/2~93-6880.$01S O / O
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Figure 1. Arrhenius plot for k,: 0, ref 1; 0 , ref 2.
of the same heat of formation from similar measurements of the kinetics of dissociation of phenyl allyl ether and phenyl ethyl ether. The enthalpy of reaction 1 at 800 K is given as follows AH,-(l)8m = E,
+ R T = 61 + 1.6 = 62.6 kcal mol-'
and AHr(1)298 = AHf(1)gm
+ ACp(1)(298 - 800)
Since the average AC, for reaction 1 is approximately zero over this temperature range (estimated by additivity considerations), it follows that AHf(l)298= 62.6 kcal mol-'. Also AH,-(C,HsO) = AHf(CsHSOCH3) + D(C6HsO-CH3) - M A C H , ) Taking m r ( C 6 H 5 0 C H 3 )= -18.0 kcal mol-' and M f ( C H , ) = 35 kcal mol-',5 AHr(C6HSO) = 9.6 kcal mol-'. Colussi, Zabel, and Benson obtained values for AHAC6HSO)of 10.5 kcal mol-' from the dissociation of phenyl allyl ether and 11.4 kcal mol-' from the dissociation of phenyl ethyl ether. When the latter value is corrected for the revised heat of formation of the ethyl radical of 28 kcal mol-]: a value of 9.7 kcal mol-' is obtained. The kinetics (3) Colussi, A. J.; Zabel, F.; Benson, S. W. Int. J. Chem. Kinet. 1977, 9, 161. (4)
(1) Mackie, J. C.; Doolan, K. R.; Nelson, P. F. J . Phys. Chem. 1989, 93, 664. (2) Paul, S.; Back, M. H. Can. J . Chem. 1975, 53, 3330.
-1
-
The experiments were performed in a stirred flow reactor over the temperature range 850-1000 K, and kl was calculated from the disappearance of anisole. An Arrhenius expression for k , was reported as ( R = 1.987 cal mol-' K-I) 64000 f 600 log kl (s-I) = 15.46 f 0.13 2.3RT
Benson, S. W. Thermochemical Kinetics, 2nd ed.;Wiley: New York,
1976. ( 5 ) McMillen, D. F.; Golden, D. M. Annu. Reu. Phys. Chem. 1982, 33, 493. (6) Cao, J.-R.; Back, M. H. Int. J . Chem. Kinet. 1984, 16, 961.
0 1989 American Chemical Society
The Journal of Physical Chemistry, Vol. 93, No. 18, 1989 6881
Comments of the decomposition of anisole, phenyl allyl ether, and phenyl ethyl ether are all therefore in good agreement regarding the heat of formation of the phenoxy radical, even though the studies were performed by different techniques and in different temperature ranges. These values are also in agreement with estimates of approximately 9 kcal mol-' from a wide variety of measurements, discussed in more detail in ref 2, and confirm the considerable resonance stabilization of this radical. Acknowledgment. The author thanks Dr. R. A. Back for helpful discussions. Registry No. C6H50CH3,100-66-3; C6H50,2122-46-5.
Ottawa-Carleton Chemistry Institute Department of Chemistry University of Ottawa Ottawa, Canada K I N 6N.5
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M. H. Back
Received: March 20, 1989 0 0
Mass Spectrometrlc Studies of CF,' Sir: I was intrigued to read the paper by Hagenow, Denzer, Brutschy, and Baumgartel on the formation of CF4+ from CF4 clusters.' The decomposition of CF4 is indeed important and interesting work. Due to the nature of their experimental setup, it is difficult to compare much of their data directly to other studies of the pressure dependence of the relative yields of the CF4+ and CF3+ ions. Direct comparison with our data is limited by our ability to compare their vacuums and molecular beam arrangement with our own. Nonetheless, since the CF4+ion has been previously observed and measured? it is interesting to compare the data where possible. In particular, sufficient data are presented to compare the results on the ratio of the CF4+yield to CF3+yield as a function of electron impact energy. The electron energy dependence of this ratio Hagenow et al. report at a stagnation pressure of 2.5 bar is quite similar to data reported previously2 for a niass spectrometer pressure of 3 X lo-' Torr (see Figure 1). It is encouraging to see that the results obtained by Hagenow, Denzer, Brutschy, and Baumgartel are consistent with our results. Their reported ionization potential (of 15.77 eV) may be smaller than (1) Hagenow, G.; Denzer, W.; Brutschy, B.; Baumgartel, H. J . Phys. Chem. 1988, 92, 6487.
10 20 Electron
30
40
50
Impact Energy (eV)
Figure 1. The ratio of the CF3'/CF4+ ion intensities as a function of electron impact energy. The pressures of 3 X lo-' Torr (-), 3 X 10-6 Torr ( 0 ) ,and 4.5 X lod Torr (+) from ref 1 are uncorrected for ion gauge cross section. Data for a stagnation pressure po = 3.5 bar ( X ) is from ref 2.
ours (16.2 eV) either because of commonly observed differences between photon and electron impact ionization mass spectroscopy or because of signal contributions from processes associated with clusters (CF&' and ion-molecule collisions.* Earlier authors3 have inferred the existence of the CF4+ ion by extrapolating to zero pressure the (linear) relationship between the dissociation ion product (CF,+) and the background gas pressure. (2) Kime, Y. J.; Driscoll, D. C.: Dowben. P. A. J . Chem. Soc.. Faradav Trans. 2 1987,83,403. ( 3 ) Deutsch, H.; Leiter, K.; Mark, T. D. Inf. J. Mass Specfrom. Ion Processes 1985, 67, 191.
Department of Physics Syracuse University Syracuse, New York 13244- 1 1 30
Yolanda J. Kime* Peter A. Dowben
Received: March 17, 1989; In Final Form: June 14, 1989
