COMMUNICATIONS TO THE EDITOR
2090 the sample cell (quartz tube, wall strength 0.1 cm, outer diameter 0.4 cm) was immersed in liquid nitrogen in a quartz dewar. After a relatively short period, ca. 5 min, the phosphorescence lifetimes did not change significantly from a constant value. The lifetimes were calculated without consideration of second-order effects in accordance with the procedure of Martin and Kalantar. It is now possible to compare the data of Martin and Iialantar with the curves obtained here. Thus a change of the phosphorescence lifetimes of benzene from 4.70 to 5.75 sec in 31IP and from 7.80 to 8.45 sec in EPA, as reported by these authors, indicate temperature differences of 3 and 4.5"K, respectively. Compared to the temperature difference between 77°K and room temperature (300°K) these differences represent, respectively, 1.3 and 1.8% change of temperature. Depending on the experimental device, e.g., if the sample is cooled indirectly, the equilibration of temperature at 77°K may well take a long period of time. Unfortunately, these authors did not mention their way of sample cooling in the description of their experimental device. The incoherence of reported phosphorescence lifetimes for a single substance a t 77°K in the same solvent must be attributed to temperaturedependent effects.2 Combining these results, the dependence of the phosphorescence lifetime upon the time of cooling can in this case be explained by temperature effects without the use of relaxation phenomena described by Martin and Kalantar.
Acknowledgment. The author wishes to acknowledge the generous financial support of the Robert A. Welch Foundation during the course of this work. (4) Eastman Kodak Co., Rochester N. Y.
14650.
DEPARTMEKT O F CHEMISTRY
INGO
H. LEUBr\-CR4
TEXAS CHRISTIllN UNIVERSITY
RECEIVED DECEMBER 16, 1968
Acetone Formation in the Pyrolysis of Isopropoxyl Radical
Batt' has presented a quantitative treatment of acetone formation in the pyrolysis of acetaldehyde in terms of a displacement reaction involving formation and decomposition of the isopropoxyl radical (i-Pro
Siy:
a )
. CH,
+ CH3CH0
b
The Journal of Physical Chemistry
+ CH3CHO
A
CHsCOCH,
+ He
Liu and Laidler2 have used a similar but more qualitative treatment. Batt has shown that (c) is 10 to 50 times slower than (b) at 523" using the experimental high-pressure activation energy3of (b), an estimated activation energy (based on thermodynamic considerations) of (c), and assuming pressure independence of both reactions and equal preexponential factors. Assuming A , 'v 1013.6 see-l, the calculated rate of acetone formation in the pyrolysis of acetaldehyde at 523" and 195 mm agreed well with that reported by Laidler and L ~ u whose ,~ results were also shown to fit a general equation for acetone production derived on the above basis. It was concluded that the displacement reaction via the isopropoxyl radical was the only route to acetone. It appears therefore that the assumptions made by Batt are substantially correct; in particular, (a), (b), (c), and (d) are not pressure dependent in the range used by Laidler and Liu (10-200 mm). However, pressure dependence of (b) (and therefore (a)) has been demonstrated by three independent groups3J'f6 using different isopropoxyl radical sources and competing reactions, and in one case3inert gas effects have been demonstrated. The pressure, p/,, a t which ICb is half the limiting high-pressure value, is 100 mm at 200" using isopropyl nitrite and 103 mm a t 150" using diisopropyl peroxide, with 20-24 Slater oscillators effective in energy transfer. According to the approximate relationship' A log pl/, = (n/2) log (T2/T1), p l / , will be at least l o 4 mm a t 523", so that, a t the pressures used by Laidler and Liu, kb will be near the second-order value. No definite information is available on possible pressure dependence of (c) and (d); Batt's estimate1 of E, 23.5 kcal mol-' suggests that (c) mill probably be pressure dependent up to high pressures since E,/RT is relatively small. Battl suggests that k b might decrease more rapidly than IC, a t low pressures, and that therefore our results on the pressure dependence of Ab, established3 by comparing acetaldehyde yields with acetone derived from the disproportionation reaction i-Pro.
Acetaldehyde. Decomposition Reactions of the
i-PrO.
'CH3
N
FORTWORTH,TEXAS 76129
8.
rather than the concerted reaction
-
0 A
d
CH3COCH3
+ H.
+ XO
A
HNO +(CH&CO
(1) L. Batt, J . Chem. Phys., 47, 3674 (1967). (2) M. T . H. Liu and K . J. Laidler, Can. J . Chem., 46,479 (1968). (3) D. L. Cox, R . A. Livermore, and L. Phillips, J. Chem. Soc., B , 245 (1966). (4) K. J. Laidler and M. T . H . Liu, Proc. Roy. Soc., A297, 365 (1967). (5) P. Gray, R. Shaw, and J. C. J. Thynne, Progr. Reaction Kinetics,
4,95 (1967). (6) M. J. Yee Quee and J. C. J. Thynne, J. Phys. Chem., 72, 2824 (1968). (7) A. F. Trotman-Diekenson, "Gas Kinetics," Butterworth and Co. Ltd., 1955, p 62.
COMMUNICATIONS TO THE EDITOR
209 1
in the isopropyl nitrite-nitric oxide system, might be questionable, since some acetone (and therefore hydrogen) might come from (c), especially at our highest temperature (200”). We have repeated our earlier work at 200” at the extremes of pressure used and estimated hydrogen by analyzing the gases volatile at - 212”, and acetaldehyde and acetone, with the results shown in Table I. Carbon monoxide formation, via radical attack on acetaldehyde, is negligible.
Comment on the Communication “Acetone Formation in the Pyrolysis of Acetaldehyde”
Sir: Phillips,’ who has correctly implied that the rate constant for the decomposition of the isopropoxy radical (i-Pro.) will be pressure dependent at 523’ and 195 Torr, questions the suggestion by Batt2 that acetone formation occurs in the pyrolysis of acetaldehyde by a displacement mechanism 6H3
+ CH3CH0
a
0
i-PrO.* b
CH3COCH3 d
+H
Table I
(1)
Initial pressures, Product yields, mm mma (C€Ia)z- CHa-
IPNo=NOo
100 20.5
CO
CHO
Hz
6.22 0.62 0.004 0.65 0.09 0.0002
[CHa-
[(CHa)zCOI [CHaCHOI[NOla
0.100(0.104b) 0.362(0.363b)
CHOl [Hzl
155 450
a I P N = isopropyl nitrite. Our previous results (ref 3). Pressures are in 154-mlreactor at 200’.
It is evident that acetone formation by (c) is insignificant in our system. Acetaldehyde-H2 ratios increase markedly with decreasing pressure showing that, provided all hydrogen is derived from decomposition of the i-Pro radical, k, falls off more rapidly than kb which is contrary to Batt’s suggestion. It is virtually certain that (c) and (d) will be pressure dependent under Laidler and Liu’s conditions. Similar observations have been made for the two possible decomposition paths of the sec-butoxyl radical; we have previously reported pressure dependence for C-C bond fission.8 Batt’s conclusions on the mechanism of acetone formation in the pyrolysis of acetaldehyde must therefore be questionable; the good agreement between calculated and experimental values for acetone production is probably fortuitous and does not prove that it is formed solely by a displacement process involving the isopropoxyl radical. A concerted displacement process and/or the mechanism suggested by Bensong 9
+ CH2CO + ‘CHZCOCH, *CHzCOCHa + CHICHO + CH3COCH3 + CH3CO *CH3
Pressure dependence means that k, and k, can be expressed in the form A ( E - E*/E)S-’, where A and E* are the Arrhenius parameters for the unimolecular decomposition of i-Pro., and E is the total energy content. S is taken to be two-thirds of the total internal degrees of freedom in i-Pr0..3 Using the thermodynamic data in ref 2, we find that k b is 101o.l sec-I, and k, is lo8.’ sec-’ if E, is 22 kcal/mol.2 At 195 Torr, spontaneous decomposition is much faster than quenching, unless deactivation occurs on every collision which is very unlikely. Hence the rate of formation of acetone is given by d(acetone)/dt = k,(i-Pro Once again, the data in ref 2 may be used to calculate (i-Pro.). This leads to a value for dmol/ml sec compared to the (acetone)/dt of 40 X mol/ml sec14still suggesting, in observed 57.5 X agreement with Benson13that most, if not all, of the acetone is produced by a concerted, displacement process. At this stage, however, acetone formation by the two bimolecular steps below cannot be excluded. , CH3 CH2CO IfCH2COCH3 a ) .
’
+
CH2COCH3
+ CHSCHO +CH3COCHs
(1) L. Phillips, J . Phys. Chem., 73, 2090 (1969).
(2) L. Batt, J . Chem. Phys., 47,3674 (1967). (3) S. W. Benson, “Thermochemical Kinetics,” John Wiley & Sons, Inc., New York, N.Y., 1968. (4) M .T. H. Liu and K . S. Laidler, Can.J . Chem., 46,479 (1968).
DEPARTMENT OF CHEMISTRY UNIVERSITY OF ABERDEEN OLDABERDEEN, SCOTLAND
L. BATT
RECEIVED APRIL1, 1969
cannot be excluded. (8) R. L. East and L. Phillips, J . Chem. Soc., A , 1939 (1967). (9) S. W. Benson, “Thermochemical Kinetics,” John Wiley & Sons, Inc., New York, K.Y., 1968, p 142.
EXPLOSIVES RESE.4RCH AND DEVELOPMENT ESTABLISHMEXT MINISTRY OF TECHNOLOGY WALTHAM ABBEY,ESSEX, USITEDKINGDOM RECEIVED JANUARY 27, 1969
Identification of the Competitive Elementary Reaction of Hydrogen Atoms with Chlorine Monoxide by Infrared Chemiluminescence
L. PHILLIPS
Sir: From a recent study’ of the reaction of hydrogen atoms with ClzO in a fast-flow system in which the (1) C. G. Freeman and L. L. Phillips, J . Phys. Chem., 72, 3031 (1968).
Volume ‘78,Number 6
June 1969
