2889
J. Phys. Chem. 1994, 98, 2889-2898
Kinetic Study of the Reactions of CHzClOz with Itself and with HO2, and Theoretical Study of the Reactions of CH2C10, between 251 and 600 K Valbry Catoire, Robert Lesclaux,' and Phillip D. Lightfoot' Laboratoire de Photophysique et Photochimie Moltculaire, Universitt de Bordeaux I, URA 348 CNRS, 33405 Talence Cedex, France Marie-Thbrhe Rayez Laboratoire de Physico-Chimie Thtorique, Universitd de Bordeaux I, URA 503 CNRS, 33405 Talence Cedex, France Received: September 30, 1993; In Final Form: December 21, 1993"
-
+
The UV absorption spectrum, self-reaction (2CHzC102 2CH2C10 0 2 (1)) kinetics and the kinetics of the reaction with H02 (CH2C102 HO2 -products (3)) of the chloromethylperoxy radical, CH2C102, have been studied between 251 and 600 K and at 760 Torr total pressure. Chloromethylperoxy radicals were produced by the flash photolysis of C12/CH3C1/02/N2 mixtures; for the study of reaction 3, chloromethylperoxy and hydroperoxy radicals were produced simultaneously by the flash photolysis of C12/CH3Cl/CHsOH/02/Nz mixtures. Radical concentrations were monitored by UV absorption spectrometry from 205 to 290 nm. The chloromethoxy radicals formed in reaction 1 produce hydroperoxy radicals via reaction with molecular oxygen (CH2C10 0 2 HO2 CHClO (2)). These hydroperoxy radicals are removed via reaction 3 and by their self-reaction (HO2 HO2 H202 0 2 (4)). Unlike the majority of systems studied to date, the rate constant for the RO2 HO2 reaction is not significantly faster than that for the RO2 ROz reaction and significant transient concentrations of HO2 are produced. Consequently, the observed second-order rate constant for radical disappearance depends strongly on the monitoring wavelength, and UV absorption cross-section and values of kl and k3 could only be obtained by an iterative process. Sensitivity analyses with respect to all model parameters on fitted rate constants were performed. Arrhenius fits to the data on kl and k3 gave kl = (1.95 f 0.16) X 10-13 exp[(874 f 26)K/TJcm3 molecule-ls-l and ks = (3.26 f 0.61) X 10-13 exp[(822 f 63)K/TJ cm3 molecule-' s-l. Average absolute uncertainties in kl and k3 over the temperature range studied, including experimental scatter and uncertainties in the analysis parameters, were estimated to be 3 1 and 33%, respectively. The UV absorption spectrum of the chloromethylperoxy radical is typical of the spectra of alkylperoxy radicals, peaking at 230 nm, where u = (4.06 f 0.53) X 10-18 cm2 molecule-'. In sharp contrast to the behavior of fully halogenated methylperoxy radicals, there was no evidence for the elimination of a chlorine atom from the chloromethoxy radical, CH2C10, even at the highest temperature employed in this study. Semiempirical MNDO calculations and RRKM calculations were performed which suggest that unimolecular elimination of HCl from CHzClO is important at high temperatures.
+
+
-+ ++
+
Introduction Reactions of chlorine-containing organic compounds have recently attracted considerable renewed attention. Specifically, they are important in incineration of waste products' and in the production of hydrocarbons by oxidative pyrolysis,2as well as in atmospheric chemistry.3 Peroxy radicals are known to be key intermediates in such proce~ses.~ However, although a great deal of work has been done on alkylperoxy radicals, much less information is availableon halogen-substituted peroxy radicals.s*6 Yet, for example, the chloromethylperoxy radical CHzClO2 certainly plays a major role in atmospheric chemistry, since it is produced in the oxidation of chloromethane, the most important natural halogenated alkane released into the atmosphere. The main source is oceanic,' and in the marine boundary layer, where the CH$l concentration is high, CHzClO2 could react with itself and with H02.* As far as we know, there has been only one kinetic study of the self-reaction? and that was performed over a temperature range suitable for atmospheric conditions, but inappropriate for low-temperature combustion. Peroxy radical self-reactions are also of interest for their mechanism. Their rate constants generally display a negative
* Author to whom correspondence should be addressed.
7 Current address: Group Technical Center, IC1 Explosives Canada, 701
Boulevard Richclieu, McMastervillc, Quebec, Canada J3G 6N3. 0 Abstract published in Advance ACS Absrracrs, February 1, 1994.
0022-365419412098-2889$04.50/0
+
temperature dependence and up to three product channels have been identified.5.6 This behavior is rationalized in terms of the formation of a tetroxide which can redissociate to reactants or go on to yield products.1° Thus, the chloromethylperoxy radical can react as follows:
+
--
CH2C102 CH,C102 CH2C1O4CH2C1 2CH2C10 CH2C10H
-
+ 0,
(14
+ CHClO + 0,(lb)
+
CH2C100CH2C1 0,
(IC)
Here we follow the practice of defining as ct = kl,/kl the branching ratio of channel la.loJ1 Two end-product studies of the chlorineinitiated oxidation of CH$l at ambient temperature have been reported in the literature.12J3 In neither study was CHzClOH or CH2ClOOCHzC1detected as a product, suggesting that channel laisdominant. Itwasalsoshown that CHClOis themain product and that no chain reaction occurs in this system. This behavior contrasts with the oxidation of CH2Cl2 and CHCl3, where CHC12012J3 and CC130l4 radicals decompose rapidly by C1atom elimination, suggesting that the CH2ClO radical is stable and reacts with 02: 0 1994 American Chemical Society
2890 The Journal of Physical Chemistry, Vol. 98, No. 11, 1994
CH,CIO
+ 0,
-
CHClO
+ HO,
Since the latter reaction is rapid in the presence of excess oxygen,l5J6 (the l / e lifetime of CH2ClO is around 25 1.1s in 1 atm of air), the hydroperoxy radical HO2 is produced in this system and must be taken into account in the reaction mechanism by introducing the following reactions:
-
+ HO, HO, + HO,
CH,ClO,
CH2C100H + 0,
(3)
+ 0,
(4)
-
H,O,
C1,
(2)
Because the rate constants for the reactions of organic peroxy radicals with HO2 are generally at least 1 order of magnitude greater than those for self-reactions, most studies of the selfreactions until now have monitored the overall rate (koh) of disappearance of the peroxy radicals according to a second-order kinetic law and deduced the true rate constant kl by employing therelationk, = k,,b/(l+ a).5However,whentherateconstants for the two reactions 1 and 3 are of comparable magnitude and when a is large, as is the case here, the kinetics are more complex and there is no simple relation between koh and kl.’ Thus, the complete reaction mechanism must be taken into account in order to extract kl from experimental data. Reactions of peroxy radicals with HOz display a negative temperature dependence of their rate constants, suggesting the reversible formation of an association complex. There is a lack of information even more severe than for the self-reactions concerning the kinetics of the reactions with H02, although these are usually more significant in the atmosphere, because of higher concentrations of HO2 than R02.5J7 The only reactions studied to date are those involving alkylperoxy, acylperoxy, and hydroxyalkylperoxy radicals, but no data exist for halogenated peroxy radicals.5~6Correlations between structure and reactivity are far from being established, although a general trend of increasing rate constant with increasing radical size has been observed for alkylperoxy r a d i c a l ~ . ~One ~J~ of the purposes of the present study was to address the following question: does CH2C102 react like CH3O2, which seems to be the least reactive of the peroxy radicals toward HOz, or does it react faster, like HOCH202, which also possesses an electron-withdrawing substituent?*O This paper presents a kinetic study of the self-reaction of the chloromethylperoxy radical and of its reaction with H02. The experiments were performed using the flash photolysis technique, with ultraviolet absorption spectrometry as the means of radical detection. The rate constants for the self-reaction and reaction with H02 of CH2ClOz were determined over a wide range of temperature (251-588 K), at atmospheric pressure. Because absoluteconcentrations must be known precisely in order to study second-order reactions, we also report measurements of the ultraviolet absorption cross-sections of CH2ClO2 as a function of temperature. At the highest temperatures employed in this work, no decomposition of the chloromethoxy radical CH2ClO by C1-atom elimination could be observed, in contrast to the behavior of other chlorine substituted methoxy radicals which generally react in this way, even at room temperature. Semiempirical and RRKM calculations were performed in order to explain this result. Experimental Section The flash photolysis/UV absorption setup has already been described in detail elsewhere10 and is discussed only briefly here. Chloromethylperoxy radicals were generated by flash photolysis, at wavelengths longer than the Pyrex cutoff, of slowly flowing C12/CH3C1/02/Nz mixtures:
+ hv(X > 280 nm)
-
C1+ CH3C1
CH,CI
+ 0, + M
-
Catoire et al.
2C1
HC1+ CH,C1
-
CH,C10,
+M
(5)
(6) (7)
Molecular chlorine concentrations were varied over the range (1.0-8.0) X 1016 molecules cm-3, resulting in initial radical concentrations of (3.6-26) X 1013 molecules ~ m - ~Methyl . chloride and oxygen concentrations were chosen so that the sequence of conversion of chlorine atoms into peroxy radicals was very rapid (the lifetime of C1 was always less than 5 ps) on the time scale of the reactions studied and dominated all other loss processes for C1 atoms at all temperatures:21y22[CH3Cl] = (1-10) X lO17moleculescm-3, [O,] = (6.6-240) X 1017molecules cm-3. Because of the high concentrations of oxygen used, the possible reverse reaction -723 and reaction 824were avoided even at the highest temperatures. CH,Cl
+ C1,
-
CH,Cl,
+ C1
(8)
For the study of the reaction between CH2ClO2 and HO2, hydroperoxy radicals were produced via the reaction of chlorine atoms with methanol:
-
+ CH,OH CH,OH + 0,
Cl
+ CH,OH HO, + CH,O
HCl
(9) (10)
CH3OH was therefore added in sufficient concentration ((3.011.0) X loi5 molecules ~ m - to ~ )generate HO2 and CH2C102 simultaneously.21 No systematic dependence of the extracted cross-sections or rate constants on the C12, CH3Cl, 0 2 , and CH30H concentrations were noticed within the ranges specified above, suggesting the absence of unforeseen chemistry. Radical concentrations were monitored by time-resolved UV absorption spectrometry. The resulting decay traces were averaged, typically over 20-50 shots, and stored on a microcomputer for analysis. A program, which numerically integrated a set of differential equations describing the possible kinetic processes, enabled the decay traces to be simulated by adjusting selected parameters (rate constant, crosssection, initial concentration), using nonlinear least-squares fitting. The residence time of the flowing gas mixture was such that the gas mixture was replaced after every one or two flashes, thus preventing secondary reactions involving reaction products. Oxygen, nitrogen, syntheticair, methane (AGAGaz Sp&iaux, >99.995%), chlorine (AGA Gaz SpQiaux, 5% in nitrogen, >99.9%), methylchloride (Mathesonor L’Air Liquide, >99.5%), and methanol (Merck Spectroscopic grade, >99.7%) were all used without further purification. Synthetic air was employed to bring the total pressure to 1 atm. However, in the experiments in which the oxygen concentration was varied, nitrogen was used instead. Results
The CHzCl02 UV Absorption Spectrum. The present kinetic study required an accurate determination of the ultraviolet absorption cross-sections for CH2C102. The UV spectrum was obtained by extrapolating the transient signal back to the time origin to measure the initial absorbance precisely a t each wavelength, since the scattered light from the flashes overloads our electronic system and renders the first 200-300 ps of the trace unusable. Thus, initial radical concentrations were kept low enough (ca. 5 X 1013 molecules cm-3) to prevent too fast a decay of radicals during the dead time, and time scales were chosen to be relatively short (5 ms) compared with the total decay time, so that extrapolation errors were minimized. Even with these favorable experimental conditions, a simple second-order
Reaction Kinetics of CHzClO2 with Itself and with H 0 2
TABLE 1: Values of Cross-Sections and Rate Constants Used in the Analysis absorption cross sections (10%/(cmz molecule-')) x HOP H202' CHzClOOH' CHClOd 210 4.17 0.37 0.31 0.00 0.26 0.15 220 3.47 0.00 230 2.30 0.19 0.10 0.02 240 1.22 0.12 0.06 0.04 250 0.54 0.09 0.04 0.05 0.05 0.03 0.06 260 0.18 Reaction Parameters kl/(cm3 molecule-' s-l) = 1.95 X 10-13exp(874KlT)e
= 1f k3/(cm3molecule-' s-I) = 3.26 X exp(822KlT)' k4/(cm3molecule-1s-*) = 3.0 X 1O-I' exp(580KIT)+ 9.4 X 1P3'[M] exp(ll50KlT)g Taken from ref 5; recentmeasurements in our laboratoryhave shown the excellent agreement of ~lo(H02)calibrated against ~m(CH302)= 4.58 X 10-18cm2 molecule-1 with the value from ref 5; the slight temperature dependenceof the cross-sectionwas alsotaken intoaccount.27 Taken from ref 28. Assumed to be identicalto that of CHsOOH taken from ref 28. Taken from ref 29. e Derived from this work. /Derived from this work and from end products studies. 8 Recommended expression from ref 25 and corrected for the revised HOz cross-sections above. CY
kinetic analysis did not describe the CH2C102removal well: the observed rate constant kob (d[R02]/dt = -2k,b[R02]2) for the self-reaction of the radical increased regularly with the wavelength (by more than a factor of 4 over the wavelength range 210-280 nm). In addition, the cross sections that were found did not give the usual shape of the spectra of alkylperoxy radicals+ we did not find a peak at around 240 nm, but rather a flat spectrum between 205 and 240 nm. These observations were interpreted by the presence of H02 produced by reaction 2. The HO2 radical has a broad and structureless UV spectrum with a maximum at 205 nm, which monotonically decreases with increasing wavelength. That H02 contributed significantly to the absorption profile arises because it was not consumed rapidly in reaction 3 on the time scale of reaction 1. In contrast to systems studied previously where the rate constants for ROz + HO2 reactions were much greater than those for RO2 + RO2 both rate constants kl and k3 in this system are almost of the same order of magnitude, as will be shown below. Consequently, the simulation of decay traces had to be performed using the whole chemical system, which means that cross-sections and kinetic parameters had to be determined simultaneously. All parameters used in simulations are given in Table 1 and typical traces with the calculated fits are shown in Figure 1a,b. The procedure used for obtaining both rate constants kl and k3 is detailed in the paragraphs that follow. At each wavelength, an adjustment of the cross-section to optimize the fit enabled the UV spectrum of CH2C102 to be determined together with a first estimate of the rate constant kl for the self-reaction, that was independent of wavelength within experimental uncertainties. The presence of H02 thus provides a satisfactory explanation of the variation of kob with wavelength. The initial CH2C102concentration was determined by replacing CHsCl with CH4 before and after each determination of a spectrum, under the same experimental conditions. The photolysis of a Clz/CH4/02/N2 mixture converts C1 atoms quantitatively into methylperoxy radicals CHs02, whose cross-section, 0240(CH302) = 4.58 X 10-'8 cm2 molecule-1 at room temperature, is now well establi~hed.~ Above ambient temperature, the slight temperature dependence of the CH302 spectrum was also taken into account.27 Three spectra were recorded, between 205 and 290 nm: at room temperature, at 464 K, and at 588 K. The so-called roomtemperature spectrum was the average of several determinations a t 273,298, and 393 K, because no significant difference in the cross-section values was noticed in this temperature range. The
The Journal of Physical Chemistry, Vol. 98, No. 11, 1994 2891 spectra are presented in Figure 2 and the corresponding crosssections given in Table 2. Like all other UV spectra of peroxy radicals, they appear as broad bands with no obvious vibrational structure. Overall 1u uncertainties on the cross-sections listed in Table 2 are estimated to be 13%,including experimental scatter and a 11% error in absolutevalues of the cross-sections for CH302 used for calibrati~n.~ The CHzC102Self-Reaction. Kinetic measurements for kl were carried out at temperatures ranging from 251 to 588 K, a t atmospheric pressure, and at monitoring wavelengths systematically varied from 210 to 260 nm for each experimental temperature. As we mentioned above, the comparable order of magnitude of both rate constants kl and k3 made their determinationdifficult. In this series of experiments, where only CH2C102 was generated, HO2 was present at a detectable concentration as a secondary product of the self-reaction 1, according to the mechanism CH,ClO,
+ CH2C10,
-
+ 0, (la) CH,ClOH + CHClO + 0, (1b) CH2C100CH,C1 + 0, (IC) CH,C10 + 0, CHClO + HO, (2) 2CH,C10
-
so that the reactions CH,ClO,
+ HO,
HO,
-
+ HO,
CH2C100H
-
H,O,
+ 0,
+ 0,
(3) (4)
had to be included in the chemical system used for the analysis of decay traces. A typical set of decay traces collected at 588 K is shown in Figure 1; the difference in the forms of the decays obtained at 220 and 240 nm is very apparent. This was attributed to the formation of significant quantities of HO2 through reactions l a and 2, which produced a distortion of the CHzClO2 decay traces from second-order kinetics, particularly at short (A < 230 nm) wavelengths. This is shown in Figure la, where HOz has about the same absorption cross-section at CH2C102. At longer wavelengths, the relative contribution of H02 to the total signal was less, but still significant: an example is given a t 240 nm (Figure lb). Decay traces recorded at longer reaction times (Figure lc,d) were analyzed for derivation of the kinetic parameters. The rate constants kl and k3 were determined in parallel at each temperature by using the present results and those obtained in the kinetic study of reaction 3 described in the next section. The experimental data were analyzed by the computer program which integrated the differential equations corresponding to reactions 1,3, and 4, and returned optimized values for the selfreaction rate constant kl. The rate constants and absorption cross-sections used in the analysis are given in Tables 1 and 2. A value of the branching ratio, a = kl,/kl, had to be introduced in the kinetic scheme in order to take into account the possible channels of reaction 1. At all temperatures and all wavelengths, good fits of kinetic traces were obtained by using a > 0.8, and no differences in the fits could be observed for 0.8 < a < 1.O. We were thus not able to derive a more precisevalue for this branching ratio from the analysis of the present data. This result is to be compared with those of the end-product studies by Sanhueza and Heicklen,l2 and Niki et al.,13 of the same chemical system as the one used in the present work. No chloromethanol was detected in the products, and it was therefore concluded that only channel l a was significant,since channel IChas never been characterized so far. Consequently, we decided to use a = 1.0 in the whole study.
Catoire et al.
The Journal of Physical Chemistry, Vol. 98, No. 11, 1994
2892
1
1
i]
2.857E-0002
OD max
a
1
OD max
b
3.174E-0002
1
i
4I
Total Absorption
1 1
3.3iEE-0002
OD wax
C
1
D.D.
max
3.216E-0002
d
J
b
Total Absorption
’
I
OD max
‘
ii
3.318E-0002
e
i
1 I
’ I
,
O.D.
ma%
f
3.21KE-0002
Total Absorption
h
Total Absorption
Decompositon of CHZCIO :
90%
50%
20%
0%
Fipre 1. Experimental decay traces and fits at 588 K and 1 atm of total pressure and at 220 and 240 nm for the determination of the cross-section and kl. For each curve, the maximum optical density (natural logarithm) is given. Smooth lines represent the simulated total absorption and the contributionsfrom each absorbing species: (a) 220 nm, full time scale 5 ms, [CH2C102]i = 6.9 X 10” molecules ~ m - 145 ~ , Torr of 0 2 ; (b) 240 nm, full time scale 5 ms, [CH2C102]i = 6.9 X 10” molecules ~ m - 145 ~ , Torr of 0 2 ; (c) 220 nm, full time scale 102 ms, [CH2C102]i 7.2 X lOI3 molecules 1211113,735Torr of 0 2 ; (d) 240 nm, full time scale 102 ms, [CH2C102]i = 7.2 X 10” molecules c ~ I -735 ~ , Torr of 0 2 ; (e) 220 nm, full time scale 102 ms, [CH2C102]i= 7.2 X 1013 molecules cm-3, 50 Torr of 0 2 ; ( f ) 240 nm, full time scale 102 ms, [CH2C102]i = 7.2 X 10” molecules ~ m - 50 ~ ,Torr of 02,simulated curves showing decomposition of fractions of CHzClO of 0,20,50,and 90% relative to reaction with 0 2 are also drawn; see text for
details. experiments carried out at 298 and 588 K, the molecular oxygen It was assumed in the preceding analysis that all the chlopartial pressure was varied between 40 and 740 Torr, and no romethoxy radicals CHzClO were converted into H02 by reaction change in the shape of the experimental traces was observed: the 2, when 0 2 was in excess. That seems to be the case at room temperature, according to the previous product s t u d i e ~ . ~ ~ J ~fitting procedure gave an optimized kl which was constant regardless of the O2 concentration. In particular, the same However, dichloro- and trichloromethoxy radicals are known to distortion of the decay traces collected at the shortest wavelengths decompose easily at room temperature by C1-atom eliminati0nl2-l~ (210-220 nm) due to the formation of H02 was always observed, as shown in Figure lc,e. Furthermore, no indication was observed CHC1,O M CHClO C1+ M (1 1) of the Occurrence of a chain reaction, which would have resulted from the decomposition of the CHzClO radical by chlorine atom CC1,O M CC1,O C1+ M (12) elimination, Since the present study was performed over a large range of CH,CIO M CH20 C1+ M (13) temperature, up to 588 K, we had to consider the stability of the even at the highest temperatures of the present study. The chain CH2C10 radical at the highest temperatures. In two series of
+
+
-
-
+ +
+
-
+
Reaction Kinetics of CH2C102 with Itself and with HO2
The Journal of Physical Chemistry, Vol. 98, No. 11, 1994 2893
TABLE 2 W Absorption Cross-Sectionsof the C H P O Z Radical from 205 to 290 nm ~(251-393K)"b
wavelength (nm)
-
0
200 220 240 260 280 300 Wavelength
/
nm
Figure 2. UV spectrum of the CHzClOz radical determined a t 298,464, and 588 K. The solid line represents the rescaled spectrum obtained by Dagaut et aL9 The other lines represent optimized Gaussian fits to the experimental data from this work; see the Discussion for details.
reaction (reactions 13,6,7, and 1) which regenerates CH2C102 would have resulted in a slower decay of the radical when decreasing the 0 2 concentration, because of the competition between reactions 2 and 13. This hypothetical behavior is illustrated in Figure If: an experimental trace at low oxygen pressure is shown together with simulated traces, showing how the shape of the traces would havevaried in case of decomposition of fractions of CHzClO of 0,20,50, and 90%, relative to reaction with 02. Using the same value for kl as found with experiments using high oxygen concentrations (Figure Id), a value of 0% for the fraction that decomposes gave the best fit to the experimental decay trace. Moreover, this test allowed us to estimate that no more than 20% (if any) of the CH2ClO radical decomposes according to reaction 13, at 588 K and 40 Torr of oxygen. All data were therefore analyzed according to thereaction mechanism in which all CH2ClO radicals react with oxygen to produce HO2. The stability of the CH2C10 radical will be discussed later. The results of all our experiments on the CH2ClOz self-reaction are summarized in Table 3. The small experimental scatter in the values for the rate constant kl at each temperature shows clearly that the system is well described by the set of parameters given in Tables 1 and 2 for all experimental conditions: no systematic dependence of the results on 0 2 concentration or on wavelength was found. A weighted least-squares fit of all the data yields the following Arrhenius expression for kl in the temperature range 251-588 K. k, = (1.95 f 0.16) X
exp[(874 f
26)K/T] cm3molecule-' s-' (I) giving kl = (3.7 f 0.4) X 10-'2 cm3 molecule-1 s-l at 298 K. The corresponding Arrhenius plot is shown in Figure 3. The quoted errors are 1u, representing only experimental scatter. IncludingthecovarianceuZ(A;E/R)= 3.14 X l@13cm3moleculd s-1 K-1, the experimental uncertainties on kl are then 27% at 251 K, falling to 24% at 298 K and 18% at 588 K. As the rate constant kl determined in this study is extracted from a relatively complex chemical system, it is essential to quantify its sensitivity to the parameters used for analysis. A sensitivityanalysis h a s been carried out as described previously.10 Artificial decay traces were generated using the kinetic scheme described, with [CHzClOz]idtid = 1 X 1014 molecules cm-3, at monitoring wavelengths of 220, 240, and 260 nm, and for temperatures of 251,298,393, and 588 K. The more important analysis parameters were then varied by *lo or i 2 0 % and the values of kl were returned from the data analysis program. It is clear that kl is sensitive to u(H02), a,and k3, particularly at the shorter-wavelengths: H02 absorbs strongly there, the amount predicted to be formed depends directly on the branching ratio
205 210 215 220 225 230 235 240 250 260 270 280 290
2.93 f 3.43 f 3.78 f 3.94 4.00 4.06 3.99 3.92 3.40 f 2.72 f 1.83 f 0.97 f 0.61 f
*
0.14# 0.28 0.18 0.19 0.19 0.22 0.21 0.25 0.20 0.20 0.09 0.10 0.15
u(464K)u
u(588K)d
* 0.25
2.61 i 0.15 2.85 f 0.12
3.58 f 0.24
3.16 f 0.17
3.71 f 0.26
3.19
3.43 f 0.22 2.88 0.09 2.31 f 0.05
3.19 f 2.80 f 2.30 f 1.63 f 1.05 f
3.35
* 0.13 0.13 0.18 0.15 0.09 0.06
a Unitsof 1018cm2molecule-1. b Valua derived from 13 absolutecrosssection measurements a t 240 nm and 6 determinations of the relative spectrum. Values derived from 2 absolute cross-section measurements of the whole spectrum. Values derived from 8 absolute cross-section measurements at 240 nm and 3 determinations of the relative spectrum. a Erron are lu; uncertainty in u (CH302) is not included, see text for details.
TABLE 3: Experimental Values of the Rate Constant for the CH&lO2 SeIf-Re&i~n temp no.of temp no.of &la (K) determinations &la (K) determinations 25 1 10 6.3 f 0.3 393 14 2.1 f 0.2 213 13 4.4f 0.3 461 11 1.4 f 0.1 298 22 4.2 f 0.4 570 5 0.86 f 0.06 323 9 3.1 f 0 . 3 588 10 0.81 f 0.06 Units of cm3 s-l; errors are lo, based only on experimental scatter; see text for details.
-
10-13 1.5
2.0
2.5 3.0 1000 K
/
3.5 T
4.0
1
5
Figure 3. Arrhenius plots of the rate constants for the reactions of CH2C102 with itself (filled squara) and with H 0 2 (open circles). a,and its disappearance depends on the rate of its reaction with CH2C102. However, it must be noted that, even though the wavelength was varied between 210 and 260 nm at each experimental temperature, most of experiments were performed at 240 nm, where the signal-to-noise ratio was maximum and where HO2 has less influence on the shape of the decay traces. Whatever the temperature, a variation of 10%in u(HO2) results in a variation of 20% in kl measured at 220 nm, but only 6% at 240 nm and 2% at 260 nm. Similarly, a decrease of 20% in a results in a decrease of 15%in kl measured at 220 nm, but in a negligible decrease at 240 or 260 nm. A variation of 20% in k3 results in an opposite variation of 25% in kl at 220 nm, decreasing to 11% at 240 nm and 7% at 260 nm. The sensitivity of k1 to kd (the rate constant for the H02 HO2 reaction) was negligible under all experimental conditions. Using errors of approximately 10 and 13% in absolute values of cross-sectionsfor H 0 2 5and CHzClOz (see above), respectively,
+
Catoire et al.
2894 The Journal of Physical Chemistry, Vol. 98, No. I I , I994
I
O.D.
7.088E-0002
max
I
TABLE 4 Experimental Values of the Rate Constant for the Reaction CHzcl02 + H02 temp temp no.of no.of kab (K) determinationsa ksb (K) determinationsa
a
255 273 298 307
/'"-
N /
\
CHSIOI
9 8 11 4
9.3 f 1.4 7.3 0.8 4.9 f 0.6 5.OkO.6
*
323 390 460 588
4 3 8 5
3.1 f 0.4 2.6 f 0.2 1.9 f 0.2 1.6 0.2
*
a 1.e. pairs of experiments. b Units of 10-12 cm3 molecule-' rl;errors are la, based only on experimental scatter; SCC text for details.
I
A.
J
1i
I -
1 0 ms
\, I
O.D.
3.726E-0002
m0x
b
TdalAb+arMIo~on
i
an analysis parameter, and we deduced a first rough value for k3; we then reanalyzed the decay traces of the 'self-reaction experiments" to obtain a new kl, and we proceeded by iteration until neither rate constant kl nor k3 varied in the analysis of the self-reaction and cross-reaction experiments. In addition, experiments with HO2 alone were conducted regularly in order to check the total radical density and to check the value of the rate constant k4 used in the analysis. No systematic dependence of k3 on the total radical concentration or on the ratio of initial radical concentrations were noted. The results are summarized in Table 4. A weighted leastsquaresfit of all the data yields the following Arrhenius expression for k3 in the temperature range (255-588 K):
k, = (3.26 i 0.61)
X
exp[(822 f 63)K/Tj cm3 molecule-' s-' (11)
Figure 4. Typical pair of decay tram and fits for an experimental determination of k3 at 220 (a) and 250 nm (b). The two experiments were performed under identical conditions of temperature (298 K), total pressure (1 atm), full time scale (10 ms), initial concentrations ([CHzC102li 6.2 X lo1>molecules cm-3, [H02]i = 7.6 X 10" molecules cm-3). The simultaneousfitting of these two tram gives k3 = 4.9 X 10-12 01313
molecule-1 s-1.
20% in a,and 28% in k3 (see below), the estimated systematic uncertainty for kl determined at 240 nm is 21%, essentially independent of temperature. Combining this uncertainty with the experimental uncertainties results in overall estimated 1u uncertainties on kl of 34% at 251 K, decreasing to 32% at 298 K and 28% at 588 K. The CH2C102 HO2 Reaction. The rate constant k3 was determined between 255 and 588 K. All experiments were performed at atmospheric pressure. As explained in the experimental section, HO2 and CHzClOzwere generated simultaneously and in comparable concentrations so that the most important reaction in the analysis of decay traces was the cross-reaction 3 between these two peroxy radicals. That the H02 UV spectrum is displaced to shorter wavelengths, in comparison with that of CH2C102, enabled the absolute initial concentrations of both radicals to be determined from pairs of experiments at different wavelengths, as was done in earlier studies on other RO2 H02 reactions.18,19,25,26 Values for the rate constant k3 were extracted from pairs of decay traces: one at 210 (or 220) nm, where u(HO2) and u(CH2C102) are about the same, and the other at 240 (or 250) nm, where u(CHzC102) is at least three times greater than u(HO2), under otherwise identical experimentalconditions. The nonlinear least-squaresoptimization program enabled the two traces to be treated together, and the HO2 and CH2C102 initial concentrations, plus the cross-reactionrate constant k3, to be returned. A typical pair of decay traces and fits is shown in Figure 4. On average, six pairs of experiments were performed at each temperature, using different initial [CH2C102]/ [H02] ratios between 1.6 and 0.6. This range is the most appropriate for extracting values of k ~ . ~The 5 parameters used for the fitting procedures are given in Tables 1 and 2: kl was thus entered as
+
+
giving k3 5: (5.2 f 0.6) X cm3 molecule-1 s-1 at 298 K. The correspondingplot is shown in Figure 3. The quoted errors are 1u,representingonly experimentalscatter. With consideration of the covariance u2(A$/R) = 4.98 X W3cm3 molecu1e-l s-l K-l, the experimental uncertainties on k3 are then 31% at 251 K, falling to 28% at 298 K,and 21% at 588 K. As for the self-reaction rate constant kl, it is important to assess the effects on the rate constant k3 of possible changes in the parameters used for analysis. Pairs of artificial decay traces were generated with [ C H ~ C ~ O J = ~M [HO&u = 5 X 10'3 molecules X cm-3, at monitoring wavelengths of 220 and 250 nm. On average, over the range of temperature255488 K,a decrease of 20% in the branching ratio a results in a decrease of 5% in k3, and variations of 10% in k4 and 20% in kl result in opposite variations of 4 and 9% in k3, respectively. In addition, k3 depends on the relative absorptions of HO2 and CHzClO2 at 210 (or 220) nm and 240 (or 250) nm: for example, a 10%greater value for u(H02) and a 10% lower value for u(CH2C102) give an increase in k3 of 10%. Using average errors of 10 and 13%for absolute cross sections of H02Sand CH2C102 (see above), respectively, and 10% for their relative cross-sections, 24% for kl (see above), 10% for k4, and 20% for a, the estimated systematic uncertainty for k3 is 20% approximately independent of temperature. Combining this uncertainty with the experimental uncertainties results in overall estimated la Uncertainties of 36% at 255 K, 34% at 298 K, and 29% at 588 K. Discussion The CH&lO2 UV Abgorption Spechum. The chloromethylperoxy radical absorption spectrum is typical of the spectra of alkylperoxy radicals. It appearsas a broad band in the wavelength range 200-300 nm, with a maximum cross-section at X 230 nm: u2m(CH2C102) = 4.06 X 10-18 cm2 moleculd at room temperature, 10%lower than that of the methylperoxy radicals and blue shifted by ca. 10 nm. This shift to shorter wavelengths seemsto be general behavior for halogenated peroxy radicals.5.30-33 It is interesting to note that the effect of substitution by a chlorine atom of a hydrogen atom in the series CH3022 CH2C102, CHC1202,34 and cC130235is to decrease regularly the maximum cross-section and to widen the absorption band.
-
The Journal of Physical Chemistry, Vol. 98, No. 11, 1994 2895
Reaction Kinetics of CH2C102 with Itself and with H02
TABLE 5 Gaussian Fi Temperature (K)for the Darameter 298 K
at Given 8 Parameters Spectrum 464 K
4.16 f 0.Mb 30.3 k 1.7 228.6 0.7
,JmU(V
4T ) e
*
hu”u
588 K 3.33 0.07 24.3 1.9 228.0 f 1.0
* *
3.71 f 0.04 23.0 1.8 225.4 0.4
*
Units of lo-’*cm2molecule-l. Errors are 1u statisticallimits based solely on the quality of the fitting. Unitlcss parameter of equation 111. d Units of nanometers. a
10’
1
7h 8
I
D
io5 15
20
25 30 1000 K
35
40
45
/ T Figure 5. Arrhenius plots of kh/o(250 nm) for the CHZC102 selfreaction: filled squares and dashed line, this work solid line, Dagaut et a1.9 The UV spectrum of CH2C102 was previously determined by Dagaut et a1.,9 also using the flash photolysis of a Ch/CH3Cl/ 02/N2 mixture. Their cross-sections, obtained by calibration against the loss of C12, are lower than those obtained in the present work 3.75 X 10-18 cm2 molecuie-l against 4.06 X 10-l* cm2 molecule-1 at the maximum (230 nm). This is not surprising since the same technique gave them a(CH302) = 3.26 X 10-18 cm2molecule-1 at 250 nm, avalue which is lower than thecurrently accepted a(CH302) = 4.09 X lO-l8 cm2 m ~ l e c u l e - ~Rescaling .~ their spectrum accordingly gives a general better agreement between the two spectra, but the peak intensity is greater than ours by 161, as shown in Figure 2. The unstructured shapes of the UV absorption bands of alkylperoxy radicals are a result of the absorptions being due to electronic transitions from the ground state to low dissociative excited states.36 Consequently,the absorptionsare generally well described by a Gaussian distribution f u n c t i ~ n : ~ J ~ . ~ ~
dT,V = umax(T)exPI-a(T)[1n~L,/VI2)
(111)
where(a, 7‘)represents the maximumvalue of the temperaturedependent absorption cross-section, a( 7‘) is a parameter characteristic of the broadness of the band, and ,A, is the position in wavelength of the maximum. The optimized Gaussian fits to the experimental CH2ClO2 spectra at room temperature, 464 K, and 588 K, are shown in Figure 2 and the fitting parameters are given in Table 5. The CH2Cl02 Self-Reaction and the Reactions of the CH2CIO Radical. The CH2C102 Self-Reaction. The observed rate constant kob(d[CH2C102]/dt = -2kh[CH2C102]2) for the selfreaction has been already measured in a previous study by Dagaut et a1.,9 over the temperature range 228-380 K and at a single monitoring wavelength (A = 250 nm). In order to compare our results with this work, our data were reanalyzed using a simple second-order kinetic scheme and values of k0b/a(250 nm) were plotted in an Arrhenius form over the same range of temperature: Figure 5 clearly shows the excellent agreement between the two studies. However, unlike the self-reactions of other peroxy
radicals, whose rate constants are much smaller than those of their reactions with H02 or whose subsequent reactions do not produce HO2, this observed rate constant cannot be simply linked to the real rate constant by the factor 1 a.5 Distortion of the decay traces due to HO2 absorption still exists at 250 nm. Numerical modeling of the whole reaction system is required to extract kl. In addition, kob decreases markedly with A below 250 nm, as explained earlier, owing to the contribution of H02 to the total absorption at these wavelengths. This change of koa with X is a clear indication that a simple second-order analysis is inappropriate. The rate constant kl was extracted from a relatively complex chemical scheme. However, this work, together with the endproduct studies,12J3 demonstrates that the reaction scheme is well-defined: reaction l a is the only important channel for the self-reaction. It produces the chloromethoxy radical, CHzClO, which then reacts with oxygen to generate H02. The removal of the latter Occurs by both reaction with CH2ClO2 (reaction 3) and with itself (reaction 4), the rate constants for which are of the same order of magnitude, as is that for the self-reaction (reaction 1). That kl and k3 are of the same order of magnitude is consistent with the analysis of end-products of Niki et al.,” who pointedout that when HO2 was initially added in their system, the relative yield of CHzClOOH did not markedly increase and the formation of CHClO was not suppressed completely, as expected if the rate constant for the reaction CH2ClO2 HO2 hadbeen muchgreater than the rateconstant for theself-reaction. Reactionsof the CHSlOradical. In contrast with the behavior of the CHCl2O and CC130radicals, which readily decompose at room temperature by chlorine-atom elimination, according to reactions 11 and 12.12-14 the present results show that the CH2C10 radical is stable and reacts quantitatively with 0 2 to produce HO2 (reaction 2) under ambient conditions, in agreement with the results of previous studies by Heicklen et aL12 and Niki et al.13 The result is also consistent with calculations reported by Rayez et aL38 who used the MNDO semiempirical method to calculate the potential barriers for chlorineatom elimination from CC130, CHC120, and CH2C10. Rayez et al.’* concluded that the large difference between these potential barriers makes splittingof the C-Cl bond much easier for CC130 (9.7 kcal mol-’) and CHC120 ( 11.8 kcal mol-1) than for CH2ClO (20.5 kcal mol-1). However, using the parameters quoted above and assuming Arrhenius preexponential factors values of about 5 X 1013 s-1,39 it can be evaluated that the CHzClO radical should decompose very rapidly at 588 K by chlorine-atom elimination, with a rate constant of about lo6 s-l, This is much larger than the rate of reaction with oxygen which, in contrast, is not very dependent on temperature: around 104s-1 for the lowest 0 2 concentration used (6.6 X 1017molecules cm-)), according to the values of the rate constants for reactions of alkoxy radicals with 0 2 recommended by Atkinson in recent reviews.15J6 Consequently, we should have observed the presence of chlorine atoms in the system, which would have resulted in the Occurrence of chain reactions at 588 K, as emphasized above. On the contrary, the behavior of the system was still consistent with the predominance of the reaction of CH2C10 with 0 2 . There are two possible explanationsfor this observation. Either the decomposition of CHzClO is much slower than we have evaluated, owing to a pronounced fall-off effect, or there exists an alternativedecompositionpathway of CHzClO producing H02, which would also explain the experimental observations. The reaction
+
+
CH2CI0
+M
-
HCO
+ HCl + M
(14)
could be such an alternative pathway. This reaction is exothermic ( c L H o 3= ~ -5.1 ~ kcal mol-’) in contrast with reaction 13 (cLHom~
2896
Catoire et al.
The Journal of Physical Chemistry, Vol. 98, No. 11. 199'4 r
= +6.6 kcal mol-') as calculated below and thus is expected to be energetically more favorable. The HCO radical would then rapidly react with 0 2
HCO + 0,
+
HO, + CO
I?
I
(15)
to produce the HO2 radicals that we observe experimentally.40 The elimination of an H atom from CH&lO, which would also have resulted in the formation of H02, is endothermicand involves a much higher barrier than the C1 atom e l i m i n a t i ~ n . ~The ~ analysis of our experimentalresults led us to explorethe reactivity of the CH2ClO radical in more detail, by performing new calculations,for 300and 600 K, in order to evaluate the possibility of Occurrence of the alternative reaction pathway 14. The molecular propertiesand potential barriers for dissociation reactions were calculated using the MNDO method at the SCF half-electron leve141.42 and with the AMPAC 4.0pr0gram.~3The search for transition states was carried out using the chain method implemented in the p r ~ g r a m . The ~ . ~corresponding ~ saddlepoints were characterized, in the optimizationprocedure, as having only one negative second-order derivative of the energy, i.e. one imaginary vibrational frequency in the normal mode analysis. The properties of the transition states and of the reactants (geometries, moments of inertia, and vibrational frequencies) were used to derive the activation entropies ASo*, corresponding to reactions 13 and 14,from statistical thermodynamic calculations. The MNDO semiempirical method calculates directly the heats of formation of compounds and the enthalpy barriers at room temperature,which can belinked to the activationenergies E, derived from experimental results after addition of an energy term RT. A correction was applied to the calculated enthalpies at 600 K. The entropies were directly calculated at different temperatures, and the Arrhenius preexponential factors A were derived from the classical relationshipof the transition state theory for unimolecular reaction:
A = ( e k T / h )e x p ( U o S / R )
1
:.. .
.. .. .: ... .,. .. ',. ...',
,...
i
.. '..
..
..
'.,.
-
'i.C H z O
". j
CH,ClO
+ -
:. HCI
HCO
14)
(-On
(-5.1)
F w e 6. Energetics and structures of the transition states for the decompositions of CH2ClO by C1 elimination (reaction 13) and HCl elimination (reaction 14) calculated using the MNDO semiempirical method. Figuresin bracketscorrespond toenthalpiea (units of kcal mol-').
TABLE 6: Calculated Parameters Characterizing tbe Decompositions of CHflO by CI Elimination (Reaction 13) or HCI Ellmination (Reaction 14)
reaction 13 CHzCIO + M CH20 + CI + M
-+
(IV)
Finally, rate constants were evaluated using the calculated Arrhenius parameters, in order to verify whether an alternative reaction channel such as reaction 14, consistent with our experimental observations, may occur at high temperatures. Calculations were performed for the decomposition of CH2ClO into C1and CH20 (reaction 13) and into HC1 and HCO (reaction 14). Transition states have been located for both reactions channels,and their structures are drawn in Figure 6. The results, which are summarized in Tables 6 and 7, lead to several conclusions. The activationenergy for HCI eliminationis slightly lower than that for C1 elimination by 1.5 to 1.7 kcal mol-', depending on the temperature. As expected, the preexponential factor A is smaller for reaction 14,since this reaction involves a tightertransitionstatethanthatofreaction13 (Figure6),resulting in a smaller entropy of activation. However, this difference in A factors is compensatedby the smaller potential barrier, so that, according to calculations, reactions 13 and 14seem to have similar rate constants of about 106 s-1 at around 600 K. The preexponentialfactors A13 and A14 were calculated using transition-state theory, and the values of k13 and k14 obtained, therefore, correspond to their high-pressure limits. In order to evaluate the fall-off effect on the rate constants of these reactions for the conditions of our experiments, i.e. at 1 atm of pressure and at 300-600 K, we have performed RRKM calculations for reactions 13 and 14,using the procedure developed by The parameters used in calculations (moments of inertia, vibrational frequencies, and threshold energies) were those calculated by the MNDO method. The collisional efficiency & was taken to be 0.3,which is a reasonable value for N2 as buffer gas. The result is that the fall-off effect is indeed significant and that therateconstantsaremarkedlydecreasedat 1 atmofpressure,
+ C1 ( W o n 13)
(6.6)
parameter T/K
ASo *a/(cal mol-' K-1) S-1) ~0/(1013
E,a/(kcal mol-') a
300 1.97 4.58 21.2
600 1.66 1.84 21.6
reaction 14 CHzClO + M HCO + HCI + M -+
300 0.63 2.33 19.5
600 0.62 4.65 20.1
Calculations were performed with the MNDOsemiempiricalmethod; Discussion for details.
see the
compared to their values at the high-pressure limit, particularly at temperatures of around 600 K. In the case of reaction 13 (Figure 7), k/k, is about '/4 at room temperature while it is about '/13 at 600 K. A similarfall-off effect was found for reaction 14. Finally, the rate constants k13 and k14 which were calculated, and the pseudo-first-orderrate constant for reaction 2 which has been estimated's (units of s-I; 40 Torr of 02) are at 300 K, 1 atm: k,[O,] = 1 X lo4; k,, = 4.3 X at 600 K,1 atm: k2[o21= 1.1 x
k,, = 4.0 X
lo4; k13 = 7.5 x io4; k,, = 18 x io4
Therefore, it is clear that, at low temperatures, the reaction with oxygen is largely dominant, in agreement with the experimental results obtained in the present work and in previous endproduct a n a l y ~ e s . ~On ~ J the ~ contrary, at 600 K, all calculated rate constants fall within the uncertainties of the calculations, and thus no definite conclusion can be drawn. However, the comparable heights of the calculated potential barriers for reactions 13 and 14 show that the elimination of HCl from CHIC10 is indeed a possible process at elevated temperatures.
Reaction Kinetics of CHzClOz with Itself and with H 0 2
The Journal of Physical Chemistry, Vol. 98, No. 11, 1994 2897
TABLE 7: Geometries of the Reactant and Transition States for the Decomposition Reactions 13 and 14 Calculated Using the MNDO Semiempirical Method CHzClO CHI' CH2' CW CCP HICClb H2CClb OCCP HlCClOC HzCClOC 1.11 1.11 1.34 1.83 reactant 106.3 106.3 111.9 120.7 -1 20.7 1.24 TS 13d 1.11 1.1 1 2.27 98.1 98.2 84.9 122.1 -122.1 TS 14e 1.11 1.18 1.23 2.04 104.7 67.0 114.6 139.7 -113.0 a Bond lengths in angstroms; labels 1 and 2 of H atoms refer to Figure 6. Planar angles in degrees. Dihedral angles in degrees. d Transition state for reaction 13. * Transition state for reaction 14. cyclic transition state where the alkyl group does not participate directly.49 Another reaction channel has been proposed,13 giving chloroformaldehyde, water, and oxygen
loo
10-1
CH,C102
8
Y
\
10-2
I
Y
10-3
10-4 io16 io1'
iola 1019 io20 io21
1023
Density (cm-')
Figure 7. Pressuredependence of k13 showing the fall-off effect at 1 atm of pressure, at 300 K (open square) and at 600 K (open circle), calculated by the RRKM procedure.
This conclusion is supported by results recently reported by Shi et al.47who show that the 1-chloroethoxy radical CH3CHClO dissociates at room temperature into CH3C0 and HCl, Le., according to the same type of process as reaction 14. Furthermore, it has been brought to our attention since the submission of this paper that Kaiser and Wallington48have obtained experimental evidence for the eliminationof HCl from CH2C10, in competition with the reaction with oxygen at room temperature. Their result indicates that the barrier heights calculated in our work are too high, which is generally the case in semiempirical calculations, and that the HCl elimination process is even faster than we have estimated. Obviously, a new reaction pathway for halogensubstituted alkoxy radicals has been revealed by the experimental data obtained both in this work and in recent studies4'a48 and by the calculations reported in the present work, on two very similar radicals. The CH2ClOz HOz Reaction. As far as we are aware, the present work representsthe first determinationof the rate constant for the CH2C102 HO2 reaction. Moreover, it is the first determination of a rate constant for reactions between a halogenated peroxy radical and H02. No direct comparison with literaturevalues is possible, but a comparisonof k3 with the value for the CH3O2 + H02 reaction can be made. There is a great similarity between the reactivity of CH302 and CHzClO2 toward H02: the rate constants are about the same and present the same negative temperature dependence. Thus, the chlorinesubstitution seems to have no influence on the reactivity of the methylperoxy radical toward H o t , while it has a large effect on the self-reaction. This conclusion confirms that reactions of the type RO2 + RO2 and RO2 H02 have different mechanisms.5 More generally, all alkylperoxy HO2 reactions studied so far have rate constants with negative temperature dependences and of the same order of magnitude ((5-20) X 1 W c m 3molecule-1 s-l at room temperature). This observation indicates that the mechanism for RO2 + H02 reactions is analogous: the initial reversible formation of a complex intermediate, presumably a tetroxide, would lead to the transfer of a hydrogen atom via a
+
+
+
+
+ HO,
-
CHClO
+ H 2 0 + 0,
(16)
This channel would involve a six-membered transition state in which the alkyl group does participate directly, as for the R02 + RO2 reactions.50 Consequently, larger variation of the magnitude of rate constants with substituents on the alkyl group would have been expected. However, this channel cannot be excluded, more especially as an excess of H2O has been detected in the product studies.12J3 In the present work, no particular channel could be identified as operating because of the weak absorption cross-sectionsof CHClO and CHzC100H. Thevalues of k3, however, were not affected by this uncertainty as both reaction channels 3 and 16 yield molecular products, which do not influence significantly either the absorption decay traces or become involved in the secondary reactions assumed to describe the kinetic behavior of the absorbing species. Atmospheric Implications. The largest part of CH3Cl released by Oceans undergoes OH attack in the troposphere, resulting in the formation of CHzC102. According to the rate constants measured in this work for the self-reaction and for the reaction with HO2, and those that can be estimated for the reactions with NO and N02,5 1.2 X 10-11 and 0.5 X 10-" cm3 molecule-' s-1, respectively, we can evaluate the fate of CH2C102 in the marine boundary layer. Taking the concentrations given by Donahue and Prinn* for NO, NO2, and H02, =2.0 X lo8, 2.0 X 108 and 1.0 X lo8, molecules cm3, respectively, 10-1596 of CH2C102 radicals react with HO2, presumably resulting in the production of the hydroperoxide CH2C100H. Modeling the marine tropospheres has shown that CH302 may reach concentration levels as high as those of H02, so that the cross-reaction CH2C102 CH302 may also play a significant role in the fate of CHzC102 radicals. However, more data are needed on the rate constants of such cross-reactions in order to assess their importance. In low-temperature combustion, the present study suggests that the CHzClO radical reacts by direct elimination of HC1.
+
Conclusions The present experiments clearly demonstrate the difficulties often encountered in obtaining peroxy radical self-reaction rate constants in the presence of secondarychemistry arising from the production of alkoxy radicals in the initial step. This problem is particularly severe in the present case as the rate constant for the reaction between the chloromethylperoxy radical and the hydroperoxy radical and those of the self-reactions of the peroxy radical and of the hydroperoxy radical are very similar in magnitude. It is clear from the very significant change in its value with monitoringwavelength that the observed second-order rate constant for disappearance of the UV absorbance cannot be equated with the rate constant of the self-reaction. Indeed, this has never strictly been the case for peroxy radical self-reactions, and product analysis should be seen as an essential complement for the vast majority of kinetic studies. The corollary of the complications introduced into kinetic measurements by alkoxy radical chemistry is that, when carefully done, kinetic measurements on peroxy radical self-reactions can
2898
The Journal of Physical Chemistry, Vol. 98, No. I I, I994
Catoire et al.
(14) Lesclaux, R.; Dognon, A. M.; Caralp, F. J . Phorochem. Photobiol. provide valuable information on the alkoxy radical chemistry A.: Chem. 1987, 41, 1. itself. In the present case, the hypothesis of chlorine atom (15) Atkinson, R. Atmos. Enuiron. 1990, 24A, 1. elimination from the chloromethoxyradical could be rejected on (16) Atkinson, R.; Carter, W. P. L. J. A m o s . Chem. 1991, 13, 195. the basis of the time-resolved experiments alone. (17) Finlayson-Pitts, B. J.; Pitts. J. N., Jr. Atmospheric Chemistry: Fundamentals and Experimental Techniques; John Wiley and Sons: New The very close similarity between the rate constants for the York, 1986. CHsO2 + HO2 and the CH2C102 + HO2 reactions over a wide (18) Rowley, D. M.; Lesclaux, R.; Lightfoot, P. D.; Hugues, K.; Hurley, range of temperature suggests that chlorine atom substitution at M. D.: Rudv. S.: Wallinnton. T. J. J. Phvs. Chem. 1992. 96. 7043. the carbon atom center does not significantly affect the rate (19) Feiter, F. F.; Citoire, V.; Leschx, R.; Lightfdot, P. D. J . Phys. Chem. 1993, 97, 3530. constant. Indeed, preliminary results in this laboratory indicate (20) Veyret, B.; Lesclaux, R.; Rayez, M. T.; Rayez, J. C.; Cox, R. A,; that the room-temperature rate constant for the CC1302+ HO2 Moortgat, G. K. J. Phys. Chem. 1989, 93, 2368. reaction is again of the sameorder of magnitude. However, more (21) Demore, W. B.;Sander,S.P.;Golden,D.M.;Hampon,R. F.;Kurylo, M. J.;Howard,C.T.;Ravishankara,A.R.;Kolb,C.E.;Molina,M. J.Chemical chlorine-containingsystemsneed to be studied before generalities Kinetics and Photochemical Data for Use in Stratospheric Modeling, can be made. Eualuarion Number 10; NASA JPL Publication 92-20; Jet Propulsion The independence of the rate constant for the RO2 + HO2 Laboratory: Pasadena, CA, 1992. reaction on chlorine substitution does not extend to that for the (22) Fenter, F. F.; Lightfoot, P. D.; Caralp, F.; Lesclaux, R.; Niiranen, J. T.; Gutman, D. J. Phys. Chem. 1993, 97, 4695. peroxy radical self-reaction: kl is approximately an order of (23) Russell, J. J.; Seetula, J. A.; Gutman, D.; Melius, C. F.; Senkan, S. magnitude faster than the rate constant for the methylperoxy M. 23rd Symposium on Combustion; The Combustion Institute: Pittsburgh, self-reaction at room temperature, emphasisingthe much greater PA, 1990; p 163. (24) Seetula, J. A,; Gutman, D.; Lightfoot, P. D.; Rayez, M. T.; Senkan, sensitivity of the rate constants of RO2 R02 reactions on the S.M. J . Phys. Chem. 1991, 95, 10688. nature of R than the rate constants for ROz H 0 2 reactions. (25) Lightfoot, P. D.; Veyret, B.; Lesclaux, R. J. Phys. Chem. 1990, 94, Finally, elimination of a chlorine atom from the CH2C10radical 708. does not appear to be important, even at elevated temperature, (26) Rowley, D. M.; Lesclaux, R.; Lightfoot, P. D.; Nozike, B. N.; Wallington, T. J.; Hurley, M. D. J. Phys. Chem. 1992, 96, 4889. in sharp contrast with fully halogenated alkoxy radicals. The (27) Lightfoot, P. D.; Jemi-Alade, A. A. J. Photochem. Phorobiol. A: results of semiempirical MNDO and RRKM calculations Chem. 1991,59, 1. presented here, confirmed by recent experimental r e s ~ l t s , ~ ~ * ~ (28) ~ Vaghjiani, G. L.; Ravishankara, A. R. J . Geophys. Res. 1989, 94, 3487. suggest that unimolecular elimination of HCl is an important (29) Libuda, H. G.; Zabel, F.; Fink, E. H.; Becker, K. H. J. Phys. Chem. fate for the CH2C10 radical, particularlyat elevated temperatures.
+
+
1990, 94, 5860.
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