J . Phys. Chem. 1991, 95, 8714-8719
8714
Ultraviolet Absorption Spectra and Kinetics of the Self-Reaction of CH,Br and CH,BrO, Radicals in the Gas Phase at 298 K Ole J. Nielsen,* Jette Munk, Garrett Locke,+ Section for Chemical Reactivity, Environmental Science and Technology Department, Risa National Laboratory, DK-4000 Roskilde, Denmark
and Timothy J. Wallington* Ford Motor Company, SRL-E3083, P.O. Box 2053, Dearborn, Michigan 481 21 -2053 (Received: March 12, 1991; In Final Form: June 18, 1991)
The ultraviolet absorption spectra of CH2Br and CH2Br02radicals and the kinetics of their self-reactions have been studied in the gas phase at 298 K by using the pulse radiolysis technique. Absorption cross sections were quantified Over the wavelength range 220-350 nm. Measured cross sections near the absorption maxima were 0,--,~,(280 nm) = (6.26 f 1.15) X cm2molecule-l and 0,--,~(250 nm) = (7.20 f 0.83) X cm2molecule-I. Errors represent statistical errors (20) together with our estimate of potential systematic errors (10%). The absorption cross-sectional data were then used to derive the observed self-reaction rate constants for reactions 1 and 2, defined as -d[R]/dt = 2k,b[RI2 (R = CH2Br or CH2Br02)of CH2Br + CH2Br products (1)
--
CH2Br02+ CH2Br02
products
(2) kl = (2.93 f 0.60) X IO-" cm3 molecule-' s-I and kZobs= (3.26 f 0.31) X IO-" cm3 molecule-I s-l (quoted errors represent 20). These results are discussed with respect to previous studies of the absorption spectra and kinetics of peroxy radicals.
Introduction Understanding the atmospheric chemistry of halogen-containing organic compounds has assumed increased importance following recent observations of dramatic stratospheric ozone loss over the antarctic in Oct 1990.1*2This loss of ozone has been correlated with unusually high levels of the radical species CIO. Bromine is also engaged in catalytic destruction of ozone in a similar cycle to that of chlorine. Indeed on a per molecule basis, bromine is a more efficient catalyst than chlorine in removing 02one.~ Little is known about the sources of different bromine compounds. Natural sources include biological activity in the Oceans and volcanic activity. Bromine-containing compounds are used for a variety of purposes: fumigants, fire-fighting agents, anesthetics, etc. The atmospheric lifetimes of the different bromine compounds differ substantially, because of differences in the rates of reaction with hydroxyl radicals and photolytic destruction rates. It has been suggested that one of the dominant sources ofbromine in the atmosphere is methyl bromide, CH3Br.4 Data also suggest that the major part of emitted CH3Br is of anthropogenic originas Because the absorption spectrum of CH3Br has insignificant overlap with the region of the tropospheric solar flux, CH3Br is mostly removed from the atmosphere by reaction with O H radicals. Atmospheric oxidation of methyl bromide in the lower atmosphere is initiated by attack by OH radicals to generate bromomethyl peroxy radicals:
CH3Br + O H CH,Br
-
+ O2 + M
+ H20 CH2Br02+ M
CH2Br
-
(3)
(4)
The products of the oxidation of methyl bromide then depend upon the fate of the CH2Br02radicals. As part of an experimental study of the chemistry of peroxy radicals and the atmospheric fate of halogen-containing compounds, the pulse radiolysis technique was used to measure the UV absorption spectra and self-reaction kinetics of the CH2Br and CH2Br02radicals. To aid in the interpretation of the kinetic data, additional experiments were performed to investigate the products of the self-reaction of CH2Br02radicals by using an *Authors to whom correspondence may be addressed. 'Permanent address: Department of Chemistry, University College Dublin, Dublin. Ireland.
0022-3654/91/2095-8714$02.50/0
FTIR spectrometer coupled to an atmospheric reactor. The results from both experimental systems are reported herein. Experimental Section Two distinct experimental systems were used as part of the present work. A pulse radiolysis transient UV absorption spectrometer was used to study the UV absorption spectra and selfreaction kinetics of CH2Brand CH2Br02radicals, while a Fourier transform infrared (FTIR) spectrometer was used to investigate the products of the self-reaction of CH2Br02 radicals. Both experimental systems have been described in detail in previous publication^^^ and will only be discussed briefly here. Pulse Radiolysis System. Radicals were generated by the irradiation of SF6/CH3Bror SF6/02/CH3Br gas mixtures in a I-L stainless steel reactor with a 30-11spulse of 2-MeV electrons from a Febetron 705B field emission accelerator. sF6 was always in great excess and was used to generate fluorine atoms: sF.5 + eSF6* (5) sF6* SFS F (6) Three sets of experiments were performed by using the pulse radiolysis system. First, the rate of reaction of F atoms with CH3Br was determined by observing the rate of formation of CH2Br radicals in the presence of known concentrations of CH3Br. F + CH3Br CH2Br H F (7) Second, by using our measured value for the rate constant k,, optimal experimental conditions for the study of CH2Br radicals were determined and the absorption spectrum and kinetics of the self-reaction of this radical were measured. Third, the value of k7 from this work was combined with our recent measurement of k8 to determine optimal experimental +
-+
-
+
(1) Farman, J. D.; Gardiner, B. G.; Shanklin, J. D. Nurure 1985,315, 207. (2) Solomon, S.Nature 1990, 347, 6291 and references therein. (3) Wofsy, S.C.; McElroy, M. B.; Yung, Y. L.Geophys. Res. Lerr. 1975, 2, 215. (4) Rasmussen, R. A.; Khalil, M.A. K. Geophys. Res. Lerr. 1984,II, 433. ( 5 ) Penkett, S.A.; Jones, B. M. R.; Rycroft, M.J.; Simmons, D. A. Nature 1985, 318, 550. (6) Hansen, K. B.; Wilbrandt, R.; Pagsberg, P. Reu. Sci. Insrrum. 1979, 50, 1532. (7) Nielsen, 0. J. Rise Nurl. Lab., [Rep.] 1984, Rise-R-480. (8) Wallington, T. J.; Japar, S. M. J. Amos. Chem. 1989, 9, 399. (9) Wallington, T. J.; Nielsen, 0. J. Inr. J. Chem. Kiner. 1991, 23, 785.
0 1991 American Chemical Society
The Journal of Physical Chemistry, Vol. 95, No. 22, 1991 8715
Self-Reaction of CH2Br and CH2Br02Radicals
c l . 0
.
,
.
,
.
.
.
. I . . microseconds
0.0
1 .o
0.5
1.5
[CHJB~], (mbar)
18
Figure 1. Absorption at 280 nm following the pulsed radiolysis of a mixture of 0.8 mbar of CH3Br and 599 mbar of SF+ Single pulse, no signal averaging. The smooth line represents a first-order fit to the data.
Figure 2. Observed first-order rate constant for the appearance of absorption at 280 nm as a function of CH3Brconcentration before (0)and after ( 0 )correction for the effect of the CH2Br self-reaction; see text.
conditions for generating the peroxy radical CH2Br02. The absorption spectra and kinetics of the self-reaction of CH2Br02were then measured.
than the expected values. We have corrected our data for the effect of the self-reaction of CH2Br radicals by detailed modeling of each experimental data point. The corrected data are shown in Figure 2. As seen in Figure 2, the corrected pseudo-first-order rate was observed to be linear with the CH3Br concentration with an intercept at the origin (within the experimental uncertainties). Variation of the radiolysis dose from a quarter to a half of the maximum dose had no effect on the observed kinetics. Linear least-squaresanalysis of the data in Figure 2 yields a rate constant k7 = (4.5 f 0.4) X 1O-Il cm3 molecule-' s-'. Quoted errors represent 2u. In view of the corrections applied to account for the effect of the CH2Br self-reaction, we choose to add an additional 10% uncertainty, resulting in our final value of k7 = (4.5 0.9) X lo-" cm3 molecule-' s-I. There are two sets of possible products following the reaction of CHJBr with F atoms: CH2Br + H F or CH3 + FBr. To establish the importance of the production of CH3 radicals, experiments were performed in which the absorption at 216.4 and 220 nm was monitored following radiolysis of mixtures of 5 mbar of CH3Br and 595 mbar of SF6. Methyl radicals absorb strongly cm2 molecule-' ''*I2) but do at 216.4 nm (u(CH,) = 4.12 X not absorb at 220 nm. No difference in the initial maximum absorption at these two wavelengths was observed, enabling an upper limit of 3% to be derived for the yield of CH3 radicals from the reaction of F atoms with CH3Br. Absorption Spectra of CH2Br and CH2Br02. Measurement of the absolute absorption cross sections for the CH2Br and CH2Br02 radicals requires calibration of the initial F atom concentration. Additionally, experimental conditions have to be chosen such that there is stoichiometric conversion of F atoms to the appropriate radical. The yield of F atoms was established by monitoring the transient absorption at 216.4 nm due to methyl radicals produced by radiolysis of SF6/CH4 mixtures (as described in detail previ0us1y'~J~).In the present series of experiments, the yield of F atoms at lo00 mbar of SF6 was measured to be 2.8 X lOI5cm-3 at full irradiation dose (using a value of 4.12 X cm2 molecule-' for u(CH3) at 216.4 nm'1q12). This calibration has an estimated accuracy of &lo%. To work under conditions where the F atoms are converted stoichiometrically into either CH2Br or CH2Br02radicals, it is necessary to consider potential interfering secondary chemistry. Potential complications include (i) competition for the available F atoms by reaction with molecular oxygen F + 0 2 M --c F02 M (8) (ii) reaction of CH2Br radicals with CH2Br02radicals CH2Br + CHzBrO2 CH2BrO + CH2BrO (9)
F+02+MdFO2+M
(8)
In all cases,the radical species F02, CH2Br,and CH2Br02were monitored by using their transient UV absorption. The output of a pulsed 150-W xenon arc lamp was multipassed through the reaction cell using intemal White cell optics. Path lengths of either 80 or 120 cm were used. The reagent concentrations used were SF,, 105-820 mbar; 02,25-100 mbar; and CH3Br, 0.4-10 mbar. All experiments were performed at 298 K, as measured by a chromelalumel thermocouple located in the center of the reaction cell. Ultrahigh-purity o2was supplied by L'Air Liquide; SF6 (99.9%) and CH3Br (99%) were obtained from Gerling and Holz and used as received. FI'IR System. The FTIR system was interfaced to a 150-L Pyrex reactor. Radicals were generated by the UV irradiation of mixtures of CH3Br and C12 in air at 700-Torr total pressure and 298 K using the output of 22 blacklamps (GE-BLB-40). The loss of reactants and the formation of products were monitored by Fourier transform infrared spectroscopy, using an analyzing path length of 26.6 and a resolution of 0.25 cm-'. Reference spectra were acquired by expanding known volumes of a reference material into the long path length cell. Initial concentrations of the gas mixtures were approximately 8 Torr of CH3Br and 2 Torr of C12 diluted in 700 Torr of ultrapure synthetic air at 295 K.
Results Reaction F + CH3Br. The rate constant for the reaction of fluorine atoms with CH3Br, k,, was measured by observing the rate at which absorption at 280 nm increases following the pulsed radiolysis of SF6/CH3Brmixtures. This absorption was attributed to the formation of CH2Br radicals from reaction 7. Radiolysis of SF6 in the absence of CH3Br produced no observable absorption at 280 nm. Figure 1 shows the observed increase in absorption at 280 nm following the radiolysis of a mixture of 600 mbar of SF6 and 0.8 mbar of CH3Br. The smooth line represents a first-order fit to the data. Experiments were performed using CH3Br concentrations ranging from 0.4 to 1.2 mbar. In all cases, within the experimental errors, the rise in absorption was observed to follow first-order kinetics. The pseudo-first-orderrate constant for the appearance of absorption at 280 nm increased linearly with the CH3Br concentration as shown in Figure 2. To investigate the cause of the nonzero intercept in Figure 2, the evolution of CH2Br radicals in our system was simulated by using the Acuchem program'O with a mechanism consisting of reactions 1 and 7. The model results were then converted into a plot of absorption as a function of time, and a first-order rise was fit. The simulated data were closely fit by first-order kinetics but the pseudo-first-order rate constants were uniformly higher (IO) Braun, W.;Hcrron, J. T.; Kahancr, D. K.Int. J. Chem. Kinet. 1988,
20, 51.
*
+
+
+
(1 1) Parkes, D. A. Inr. J. Chem. Kinet. 1977, 9, 45 I. (12) Macpherson, T.; Pilling, M.J.; Smith, M. J. C. J . Chem. Phys. 1985,
89, 2268. (13) Ellermann, T. R i m N d . fub., [Rep.] 1991, Rim-M-2932. (14) Cox, R. A.; Munk, J.; Nielsen, 0. J.; Pagsberg, P.; Ratajczak, E. Chem. Phys. Left. 1990, 173, 206.
8716 The Journal of Physical Chemistry, Vol. 95, No. 22, 1991
Nielsen et al.
%
0.8
-
0.6
-
0.4
-
0.2
-
n
v
a 0
! P-
3 .C
-
200
0
400
so0
800
0.0 0
1000
[SF61 (mbar)
Figw 3. Maximum transient absorption at 280 nm following the pulsed radiolysis of CH3Br/SF6mixtures as a function of the sF6 concentration. Open symbols represent full radiolysis dose experiments; filled symbols represent half-dose. To compare both sets of data, the sF6concentration
employed in the lower dose experiments has been divided by 2.
and (iii) reaction of F atoms with CH2Br and/or CH2Br02 radicals F + CH2Br products (10) F CH2Br02 CH2Br0 + FO (11)
+
-
The rate constant for the reaction of F atoms with CH3Br was measured as part of the present work. The rate constant for reaction 8 has been measured previously in our laboratory9 ( k , 2.4 X cm3 molecule-' s-' at 600 Torr of SF6diluent). By using the rate constant ratio k7/ksand the [02]/[CH3Br]concentration ratio, the importance of FOz formation in the present series of experiments can be calculated. With regard to the second potential complication, as far as we are aware, there are no kinetic data for the reaction of CH2Br with CH2Br02;however, by analogy to the corresponding reactions of CH3 with CH30211J5 and CH3COCH2with CH3COCH202,14 k9 was estimated to be approximately 5.0 X lo-" cm3molecule-' s-I. By using this rate constant, the potential importance of the reaction of CH2Br with CH2Br02was assessed. To minimize the importance of FOz radical formation in the present experiments, it was desirable to use low molecular oxygen concentrations. On the other hand, to minimize the importance of the reaction of CH2Br with CH2Br02,it was desirable to use high molecular oxygen concentrations. Clearly a compromise needed to be struck. It was decided to operate the experiments with [CH3Br] = 2.5-10 mbar and [O,] = 25-100 mbar. To our knowledge, there are no kinetic data available on the reactivity of F atoms toward the radical species CH2Br or CH2Br02. To check for the presence of complications in our experiments caused by unwanted radical-radical reactions, two series of experiments were performed. Firstly, the transient absorption at 280 nm was observed in experiments using [CH,Br] = 5 mbar, with the SF6 pressure varied over the range 100-1OOO mbar. The maximum in the transient absorption of CH2Br at 280 nm observed in these experiments is plotted as a function of the SF6 concentration in Figure 3. Secondly, the transient absorption at 250 nm was observed in experiments using [CH3Br] = 5 mbar and [O,] = 50 mbar and with the SF6pressure vaned between 100 and 1000 mbar. In Figure 4, the maximum of the transient absorption of CH2Br02at 250 nm is plotted as a function of SF6 concentration. In both sets of experiments, the pulse dose was varied by a factor of 2. To compare both sets of data, the SF6 concentration employed in the lower dose experiments has been divided by 2. For experiments employing the full radiolysis dose and SF6concentrations greater than 600 mbar, use of [CH3Br] = 5 mbar and [O,] = 50 mbar resulted in initial absorptions that were 10-20% lower than expected based upon extrapolation of the data obtained at lower SF6concentrations. This behavior is ascribed to secondary chemistry at high F atom concentrations, resulting in incomplete (15) Pilling, M. J.; Smith, M. J. C. J . Phys. Chem. 1985. 89. 4713.
CH2Br02 at 250nm
200
400 600 [SFgl W a r )
0
800
1000
Figwe 4. Maximum transient absorption at 250 nm following the pulsed
radiolysis of CH3Br/o2/SF6 mixtures as a function of the sF6 concentration. Open symbols represent full radiolysis dose experimetlts; filled symbols represent half-dose. To compare both sets of data, the sF6 concentration employed in the lower dose experiments has been divided by 2.
TABLE I: UV Absorption Cross !%ctioas Measured in This Work 1 P u , cm2 molecule-' wavelength, nm CHiBr CHiBr0i 220 174 346 230 210 398 240 332 699 250 414 720 260 389 706 270 435 663 280 626 736 290 523 722 300 425 639 310 343 456 320 231 305 330 I54 201 340 96 136 88 350 62
conversion of F into CH2BrOZ. From Figures 3 and 4, it can be seen that the initial absorption was found to be linear with the SF6(and hence initial F atom) concentration for [CH3Br] = 5 mbar and [SF6] < 600 mbar. This linearity suggests that, under these experimental conditions, reactions 9-1 1 are of negligible importance. Consequently, the majority of our experiments were performed using [SF6] < 600 mbar. The solid lines in Figures 3 and 4 represent the linear leastsquares fits to the data (>600 mbar full dose experiments excepted), which have slopes of (9.14 f 0.76) X 10" and (10.1 f 0.2) X IO4, respectively, errors represent 2a. Combining these values with the calibrated yield of F atoms of 2.8 X lOI5 ~ 1 1 at3 ~ 1000 mbar of SF6 with full irradiation dose yields cm2 molecule-' U,-~,~,(280 nm) = (6.26 f 1.15) x acH,Bf12(250nm) = (7.20 f 0.83) x
cm2 molecule-'
Errors represent statistical errors ( 2 4 together with our estimate of potential systematic errors (&lo%). The value quoted for includes a small correction (4%) to account for the formation of F02 radicals in our system; calculation of this correction is discussed below. In order to map out the absorption spectra of CHzBr and CHZBrO2radicals, experiments were performed to measure the initial absorption between 220 and 350 nm following the pulsed irradiation of SF6/CH,Br or SF6/CH3Br/02 mixtures with sF6 = 550 mbar. Initial absorptionswere then scaled to those at either 280 (CH2Br) or 250 nm (CH2Br02),and the appropriate measured absorption cross section was then used to convert the observed initial absorption into an absorption cross section. The values so obtained are given in Table I and shown in Figures 5 and 6. For measurements in the presence of 02, corrections were applied to account for the formation of F 0 2 radicals. The concentrations of F02 radicals in our experiments were calculated using the [CH3Br]/[02]concentration ratio and the rate constant
~
The Journal of Physical Chemistry, Vol. 95, No. 22, 1991 8717
Self-Reaction of CH2Br and CH2Br02Radicals
TABLE 11: Measured Kinetic Data for the Radical Self-Reactions 1 O6k/u, cm2 IOIlk, cm3 wavelength, nm molecule-' molecule-I s-I CH2Br CH2Br Products 5.57 2.30 250 260 7.72 3.00 270 6.82 2.97 280 4.66 2.92 5.1 1 2.67 290 300 5.91 2.5 1 310 8.86 3.04 3.15 320 13.63 330 21.6 3.32 340 32.9 3.16 350 51.1 3.17 mean = (2.93 i 0.60) X IO-" cm3 molecule-I s-I
-
+
180
260
220
300
340
h , nm
Figure 5. Absorption cross-sectional data for CH2Br measured in this work together with data for CHICl reported by Lightfoot et al?'
-
+
N
E
600
-
400
-
200
-
0 N
E
CH,Br02 CH2Br02 Products 240 4.15 2.90 250 4.54 3.27 260 4.66 3.29 270 5.06 3.35 280 4.43 3.26 290 4.60 3.32 300 5.34 3.41 mean = (3.26 & 0.31) X IO-" cm3 molecule-' s-I
X
b
250
200
350
300 h , nm
Figure 6. Absorption cross-sectional data for CH,Br02 ( 0 )and CH,Br (0) measured in this work.
Y
I
I
I
I I 0
I
I microseconds
I 40
Figure 7. Typical CHzBr decay curve monitored in absorption at 330 nm with [CH,Br] = 10 mbar and [SF,] = 590 mbar. Single pulse, no signal averaging. The solid line represents a nonlinear least-squares secondorder fit.
ratio, k7/k8. The following absorption cross sections for F0, radicals were used: 220 nm, 1.34 X lO-I7; 230 nm, 6.9 X 10-l8; 240 nm, 3.0 X lo-? and 250 nm, 1.60 X 10-18.16J7Corrections ranged from 3 to 13%. Kinetic Data for Self-Reactions. Figure 7 shows typical transient absorption data obtained for the self-reaction of CH2Br radicals, together with a nonlinear least-squares second-order fit. As discussed in previous publi~ations,'~J~ kinetic analysis of second-order decays can be complicated by the formation of products that absorb at the monitoring wavelength, leading to a postflash transmitted light intensity that is lower that the preflash value. Accordingly, the experimental data were fit to the three-parameter expression 1 / B - 1/Bo 7 2kt/~,L (16) Pagsberg, P.; Ratajczak, E.; Sillesen, A. Chem. Phys. Lrr.1987,141, 88. ( 17) Pagsberg, P. Private communication, 1990. (18) Kurylo, M . J.; Ouellette, P. A.; Laufer, A. H. J . Phys. Chem. 1986, 90,437. (19) Sander, S.P.; Watson, R. T. J . Phys. Chem. 1981,86, 2960.
where B = In (I,/Z), Bo = In ( Z J Z ' ) , k is the second-order rate constant for the self-reaction of the radicals, UR is the absorption cross section of R a t the monitoring wavelength, L is the monitoring UV path length (80 or 120 cm); I' is the minimum transmitted light intensity following the radiolysis pulse, I, is the final light intensity, and I is the transmitted light intensity at time t.
The decay of the transient absorption from CH2Br and CHIBr0, radicals was monitored at a variety of wavelengths over the range 250-350 nm. At all wavelengths, the decay was well represented by the second-order least-squares fit. The results, expressed as k/u values, are given in Table 11. These results can be converted to rate constants by multiplying the kob/u values, derived from the aforementioned analysis, by our measured absorption cross sections at the appropriate wavelengths. The rate constants thus derived are listed in Table 11. Systematic variation of each of the following parameters from the base case ([SF,], 545 mbar; [CH,Br], 5 mbar; [O,], 50 mbar; dose, maximum) had no observable effect on the values of k/u derived: [SF,], 1 W 9 5 0 mbar; [CH3Br], 2.5-10 mbar; [O,], 25-100 mbar; pulse dose, by a factor of 3. As seen from Table 11, the observed rate constants for reactions 1 and 2 were independent of wavelength, suggesting that our kinetic analysis is not complicated by the presence of other absorbing species. No effect of total pressure over the range 150-1000 mbar of SF6diluent on either kl or k2 was observed. Consequently, the data in Table I1 can be averaged to yield k l h = (2.93 f 0.60) X lo-" and k20bs= (3.26 f 0.31) X lo-" cm3 molecule-' s-I at 298 K. Quoted errors represent 2u. Product Study of the Self-Reaction of CH2Br02Radicals. In view of the rapidity of reaction 2 measured in this work, a product study of the self-reaction of CH2Br02radicals was conducted. Specifically we were interested in the potential generation of Br atoms from the reaction sequence CH2Br02+ CH2Br02 CH2Br0 CH2Br0 0, (2a) CH2Br0 M HCHO Br (12)
-+
+
-+
+
+
Bromine atoms could then react with CH2Br02radicals via a chain mechanism, so Br + CH2Br02 CH2Br0 + BrO (13)
-+ -
BrO + CH2Br02 Br02
M
CH,BrO Br
+ Br0,
+ O2
(14)
(15) Such a chain mechanism, if operative, would considerably complicate the kinetic analysis of the observed CH2Br02decay.
Nielsen et al.
8718 The Journal of Physical Chemistry, Vol. 95, No. 22, 1991
0.0
0.2
0.1
0.3
0.4
0.5
~([CHqlt,J[CH4lt)
Figure 8. Plot of the observed decay of CH,Br in the presence of CI atoms versus that of CHI. Open symbols were obtained in air diluent; filled symbols were obtained in nitrogen diluent.
To check for such a complication, experiments were performed to search for the formation of HCHO, and hence Br atoms, following the self-reaction of CH2Br02radicals. In this product study, photolysis of molecular chlorine in the presence of CH3Br/air mixtures was used as a source of CH2Br02radicals. To assess the magnitude of possible loss of any HCHO product through secondary reactions with CI atoms, it is necessary to know the reactivity of CI atoms toward both CHI& and HCHO. While the kinetics of the reaction of C1 with HCHO are well established,Mto our knowledge there is no kinetic information available on the reaction of CI atoms with CH3Br. Hence, it was necessary to first investigate the kinetics of reaction 16. C1+ CH,Br
-
HCI
+ CH2Br
(16)
A relative rate technique was used to measure the ratio kI6/kl7.
C1+ CH4
+
HCI
+ CH3
(17)
The experimental procedure has been described recently.2' Figure 8 shows a plot of the observed decay of CH3Br as a function of that of CH4 following the photolysis of CH3Br/CH4/CI2mixtures in either air or nitrogen diluent at 700 Torr of total pressure and
295 K. As seen from Figure 8, there was no observable difference between separate experiments performed in air or N2. A linear least-squares analysis of the data in Figure 8 yields a rate constant ratio kI6/kl7= 4.00 f 0.17. Using the literature value of kI7= 1.0 X ICi3an3molecule-' then leads to kl6 = (4.00 0.17) X cm3 molecule-' s-I. Quoted errors represent 2u.
*
The literature value for the rate constant of the reaction of C1 atoms with HCHO is 7.3 X 10-I' cm3 molecule-' s-Ie20 Thus, HCHO is a factor of 180 times more reactive toward CI atoms than CH3Br. Hence, to avoid significant consumption of HCHO by CI atoms in our product study, it was necessary to have large amounts of CH3Br (a8 Torr) and operate at small conversions ( F > C1. Before comparing the bimolecular rate constant observed for reaction 2 in our work with literature data for CH302,CH2F02, and CH2C102,it should be noted that the value measured, k*, may be an overestimate of the true bimolecular rate constant. This potential overestimation arises from the possible reaction of CH2Br02 with H 0 2 radicals produced from the reaction of CHzBrO radicals with molecular oxygen: CH2Br02+ CHZBrO2 CH2Br0 + CH2Br0 + O2 (2a) CH2Br0 + O2 H 0 2 + HC(0)Br (19) CH2BrOZ+ H 0 2 CH2BrOOH + 0 2 (20)
- -
--
-
cm3 molecule-I s-1.22
(20) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Molina, M. J.; Hampson. R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R. NASA-JPL Publication 90-1, 1990. (21) Wallington, T. J.; Andino, J. M.; Ball, J. C.; Japar, S. M. J . Atmos. Chem. 1998, IO, 301. (22) Dolson, D. A.; Leone, S. R. J . Phys. Chem. 1987, 91, 3543.
(23) Yanvood, G.; Green, M.; Niki, H.XIXth Informal Conference on Photochemistry, Ann Arbor, MI, June 1990. (24) R o u ~ c lP. , 8.;Lightfoot. P. D.; Caralp, F.;Catoh, V.; Lesclaux, R.; Font, W.1. Chem. Soc., Faraday Trans. 1991,87, 2361. (25) Dagaut, P.; Kurylo, M. J. J . Photochem. Photoblol. A: Chem. 1990, 51, 133. (26) Dagaut, P.; Wallington, T. J.; Kurylo, M. J. Int. J. Chem. Klwt. 1986, 20, 815.
J. Phys. Chem. 1991, 95,8719-8726 To correct for loss of CH2Br02by reaction with H 0 2 , it is necessary to have two pieces of information, namely, the branching ratio kk/k2 and the rate constant km Since neither of these pieces of information are available, it is not posible to correct for this effect in the present work. However, in view of the rapidity of reaction 2, any correction is likely to be small (
