2220
Davis et al.
210-20 ns. (The rise time of Arai and Firestone's system was 40 ns.) Thus we conclude that the emission decays observed by Arai and Firestone are due to the same reaction sequence as the slow process observed by us and by Hanle et al., and which our analysis suggests involves a dissociative ion-recombinationreaction. The fast process observed by Hanle et al. is the same as the fast process observed in the present work, where the 2p levels are produced both by direct excitation and by cascade from higher electronic levels. References and Notes (1) Work performed in part under the auspices of the Divislon of Physical Research of the U S . Energy Research and Development Administration and in part under the auspices of the Australian Research Grants Committee. (2) R. F. Firestone and L. M. Dorfman, Actions Chim. Biol. Radiaf., 15, 7-46 (1971). (3) M. C. Sauer, Jr., Adv. Radiat. Chem., 5, 97-184 (1976). (4) M. C. Sauer, Jr., Int. J. Radiat. Phys. Chem., 8, 33 (1976). (5) S. Arai and R. F. Firestone, J. Chem. Phys., 50, 4575 (1969). (6) W. Hanle, E. Kugle, and A. Schmillen, Ann. Phys., 13, 252 (1964). (7) J. B. Cumming, PhD. Thesis, University of Melbourne, 1975.
(8) S. Gordon, W. Mulac, and P. Nangia, J. Fhys. Chem., 75, 2087 (1971). (9) Reference 5 above lists the sources which Identify and designate the electronic levels involved in the atomic transitions. (10) F. J. Mehr and M. A. Biondi, Phys. Rev., 176, 322 (1968). (11) H. J. Oskam and V. R. Mittelstadt, Phys. Rev., 132, 1445 (1963). (12) J. Le Calve and M. Bourhe, J . Chem. Phys., 55, 1446 (1973). (13) J. M. Warman and M. C. S a w , Jr., J. Chem. Fhp., 62, 1971 (1975). (14) J. M. Warman and M. P. de Haas, J. Chem. Phys., 63, 2094 (1975). (15) Ya. F. Verolainen and A. L. Osherovich, Opt. Spectrosc. (Engl. Trans.), 27, 14 (1969). (16) L. Allen, D. G. C. Jones, and D. G. Schofleld, J. Opt. SOC.Am., 59, 842 (1969). (17) R. S. Freiberg and L. A. Weaver, J . Appl. Phys., 36, 250 (1967). (18) J. 2. Klose, Phys. Rev., 141, 181 (1966). (19) W. R. Bennet, Jr., and P. J. Kindlmann, Phys. Rev., 149, 38 (1966). (20) J. B. Shumaker and C. H. Popenoe, J. Opt. Soc.Am., 32, 8 (1967). (21) R. Henck and R. Voltr, J. Phys. (Paris), 29, 149 (1968). (22) J. K. Bailou and C. C. Lin, Phys. Rev. A , 8, 1797 (1973). (23) V. E. Bondybey and T. A. Miller, J. Chem. Phys., 66, 3337 (1977). (24) R. L. Platzmann, Radbt. Res., Roc. Int. Congr., 3rd, 20-42 (1967). (25) T. R. Connor and M. A. Biondi, Phys. Rev. A , 140, 778 (1965). (26) C. Moore, Natl. Bur. Stand. Clrc., No. 467 (1949), Vols. I, 11, and 111. (27) The values are estlmates of rare gas ion dissociation energies which are listed with calculated and experimentally obtained values by R. S. Mulliken, J. Chem. Phys., 52, 5170 (1970). (28) L. Frommhold and M. A. Biondi, Phys. Rev., 185, 244 (1969).
A Temperature Dependent Kinetics Study of the Reactions of HCI with OH and O(3P) A. R. Ravishankara, 0. Smlth, R. T. Watson,+ and D. D. DavW Atmospheric Sciences Division, Applied Sciences Laboratory, Engineering Experiment Station, Georgia Institute of Technology, At/anta, Georgia 30332 (Received January 25, 1977) Publication costs assisted by the Georgia Institute of Technology
The flash photolysis-resonance fluorescence technique was employed to determine the temperature dependencies of the rate constants for the reaction of O(3P)and OH radicals with HC1. These reactions were studied under pseudo-first-order conditions and in the absence of interfering secondary reactions. The Arrhenius expression for each bimolecular rate constant is given as follows in units of cm3 molecule-l s-l: kl = (3.3 f 0.3) X exp[-(937 f 78) cal rnol-l/Rq 250-402 K, OH + HC1- H20 + Cl(1) and k2 = (5.2 f 1.0) X exp[-(7510 f 750) cal mol-'/RT] 350-454 K, O(3P)+ HC1- OH + C1 (2). The stratospheric implications of this new rate data are discussed.
Introduction According to present atmospheric models, hydrogen chloride is predicted to be one of the principal chlorine containing species in the strat0sphere.l Recent measurements of the HC1 concentration in the lower stratosphere now support this predictiona2 In the stratosphere the formation of HC1 proceeds through the reaction of chlorine atoms with RH species (Le., CH4, HQ,HzOt, and/or HOz), thus removing reactive chlorine from the catalytic cycle: c 1 + 0, c10
-+
c10
+ 0,
+ 0 -+ c1 + 0,
0 + 0,
-+
20,
Present Address: Jet Propulsion Laboratories, Building 183-601, 48?0 Oak Road Drive, Pasadena, Calif. 91103. This author acknowledges the partial support of this research by the National Aeronautics and Space Administration. Part of this work was carried out while this author was at the Department of Chemistry, University of Maryland, College Park, Md. 20742. The Journal of Physical Chemistry, Vol. 8 1, No. 24, 1977
The chlorine atom in HC1, however, can be reintroduced into the catalytic ozone destruction cycle via reaction of OH or O(3P)with HC1 and/or by photolysis of HC1, e.g. OH
+ HCl+k l
H,O k,
O(3P)+ HC1hv
HCl+ H
+ C1
+
OH
C1
(1)
+ C1
(2) (3)
Since HC1 is the dominant "temporary sink" for C1 atoms in the stratosphere, reliable rate constants for reactions 1 and 2 are essential for stratospheric modeling calculations. Five measurements of the rate constant kl have been reported. Wilson et aL4measured klat high temperatures, while Takacs and Glass5and Hack et a1.6 obtained room temperature values. In addition, there have been two measurements of kl over an extended temperature range, one by Smith and Zellner' and the other by Zahniser et al.* All measurements are in very good agreement at 300
Reactions of HCI with OH and O(3P)
2221
K; however, at stratospheric temperatures of 225 K, the klvalue reported by Zahniser et al. is approximately 20% higher than that of Smith and Zellner. Several investigations of reaction 2 have been reported over an extended temperature range where the activation energy could be evaluated. Balakhnin et al. were the first to study the kinetic behavior of reaction 2. These authors reported an activation energy of 4.52 kcal/mol. Brown and Smithzolater obtained an E value of 5.95 kcal/mol, a number which agrees well with some recent work reported by Singleton and Cvetanovic." Hack et aL6and Wong and Belles,125owever, have reported values for reaction 2 of 6.44 and 7.2 kcal/mol-l, respectively. Thus, at the present time the activation energy for process 2 must be considered to be in only fair shape. Reported here are the results of a new study on both reactions 1and 2, the purpose of which was to further test the reliability of the earlier rate data using a completely different experimental technique-flash photolysis-resonance fluorescence.
Experimental Section Detailed descriptions of the flash photolysis-resonance fluorescence technique employed in this investigation have been described p r e v i o ~ s l y . l ~ -In ~ ~this manuscript, therefore, we have pointed out only those experimental features necessary for an understanding of the present study. In this investigation, an all-Pyrex cell with an internal volume of -150 cm3 was used to study both reactions 1 and 2. The reaction mixture was maintained a t a known constant temperature by circulating either methanol (250-300 K) or silicone oil (300-450 K) from a thermostated circulating bath through an outer jacket of the reaction cell. The temperature of the reaction cell was measured with an iron-constantan thermocouple. The transient species OH and O(3P) were formed by the photolysis of a suitable photolyte using a nitrogen spark discharge. The technique for detection of OH radicals by resonance fluorescence has been documented e1~ewhere.l~ The OH radicals in this study were produced either by directly photolyzing H20or by photolyzing a mixture of 50 mTorr of O3 and 100 mTorr of H2. In the latter case, OH was formed through the sequence of reactions: 0,
hv
O('D) t 0,
