4514
J. Phys. Chem. 1996, 100, 4514-4520
The CF3C(O)O2 Radical. Its UV Spectrum, Self-Reaction Kinetics, and Reaction with NO M. Matti Maricq* and Joseph J. Szente Research Laboratory, Ford Motor Company, P.O. Box 2053, Drop 3083, Dearborn, Michigan 48121
Gregory A. Khitrov† and Joseph S. Francisco Department of Earth and Atmospheric Sciences and Department of Chemistry, Purdue UniVersity, West Lafayette, Indiana 47907-1397 ReceiVed: NoVember 7, 1995; In Final Form: December 12, 1995X
Flash photolysis combined with time-resolved UV spectroscopy and transient infrared absorption is used to investigate the reactions of CF3C(O)O2 with itself, CF3O2, and nitric oxide. The UV spectrum of CF3C(O)O2 exhibits two bands, the stronger short wavelength component of which has a maximum cross section of 7.1 × 10-18 cm2 at 207 nm. These bands are used to monitor the disappearance of CF3C(O)O2 and the secondary +5 formation of CF3O2, yielding a self-reaction rate constant of (3.7-2 ) × 10-12 e(270(200)/T cm3 s-1. The cross reaction between CF3C(O)O2 and CF3O2 is found to be slow, having a rate constant of e2 × 10-12 cm3 s-1. Transient IR monitoring of the loss of NO and concomitant formation of NO2 leads to a rate constant of +2.2 (4.0-1.4 ) × 10-12 e(563(115)/T cm3 s-1 for the reaction between CF3C(O)O2 and NO. This result implies that CF3C(O)O2 radicals formed as intermediates in the atmospheric degradation of hydrofluorocarbons (HFCs) are rapidly converted into CF3O2 radicals, which are in turn converted into carbonyl fluoride and FNO.
I. Introduction Amongst the atmospheric degradation products of the partially halogenated hydrocarbons that are currently being used as alternatives to the chlorofluorocarbons (CFCs) are carbonyl compounds of the form CF3C(O)X, where X ) F, Cl, H. The fully halogenated species are relatively unreactive with other gas-phase atmospheric species and are primarily removed by photolysis1 or by incorporation into cloud droplets.2 Trifluoroacetaldehyde, however, is susceptible to attack by hydroxyl radicals to form CF3CO and water.3 Although CF3CO radicals collisionally dissociate4 into CF3 + CO, the rate of bond cleavage under atmospheric conditions is considerably smaller than the rate of adding oxygen;5 thus, their principal fate is conversion to trifluoroacetylperoxy radicals. Possible fates of CF3C(O)O2 induce reactions with NO, NO2, and HO2. The lower bound of k1 > 1 × 10-11 cm3 s-1 reported by Wallington et al.5 for the reaction
CF3C(O)O2 + NO f CF3C(O)O + NO2
(1)
suggests that this represents a significant removal process for the peroxy radical. The present study derives temperaturedependent rate constants for reaction 1 by monitoring in real time both the disappearance of NO and the appearance of NO2. The paper also presents results for the rate constant of the CF3C(O)O2 self-reaction, which competes under our experimental conditions with reaction 1. It is concluded that trifluoroacetaldehyde, a degradation product of some hydrofluorocarbon compounds, will not adversely affect the ozone layer. II. Experimental Section The formation of CF3C(O)O2 radicals, in concentrations of (3-10) × 1014 cm-3, is initiated by the 351-nm photolysis * Author for correspondence. † Present address: Department of Chemistry, Eastern Michigan University. X Abstract published in AdVance ACS Abstracts, February 15, 1996.
0022-3654/96/20100-4514$12.00/0
(∼300 mJ in a 1-cm2 cross section beam) of either F2 or Cl2 in a slowly flowing gas mixture containing trifluoroacetaldehyde, oxygen, nitrogen, and, when required, nitric oxide. F2 is preferred over Cl2 as a radical source, because fluorine atoms react more rapidly with CF3CHO than do chlorine atoms; however, F2 cannot be used in conjunction with NO because of the spontaneous reaction
F2 + NO f F + FNO Following photolysis of the halogen, the peroxy radicals are formed via
X + CF3CHO f CF3CO + XH
(2)
CF3CO + O2 + M f CF3C(O)O2 + M
(3)
where X ) F, Cl. Sufficient oxygen (∼15 Torr) is added to the gas mixture to ensure that reaction 3 overwhelms CF3CO dissociation and that it is essentially instantaneous on the >100µs time scale of the subsequent peroxy radical reactions. With fluorine initiation (k2(F) ) 2.3 × 10-11 cm3 s-1 at 295 K),5 enough trifluoroacetaldehyde can be added to ensure that reaction 2 is also rapid on this time scale. This is more difficult to achieve for the slower chlorine initiation (k2(Cl) ) 2.4 × 10-12 cm3 s-1 at 295 K);6 even with ∼15 Torr of CF3CHO, a slight delay in CF3C(O)O2 formation is noted. The delay is taken into account by including reaction 2 in the model used to fit the data. Both time-resolved UV spectroscopy7,8 and transient IR spectroscopy9 are used to interrogate in real time the reaction progress. More complete descriptions of the two experimental setups are available in refs 7 and 8 and in ref 9, respectively. The loss of CF3C(O)O2 radicals via
CF3C(O)O2 + CF3C(O)O2 f 2CF3C(O)O + O2
(4)
and the secondary generation of CF3O2 radicals and their © 1996 American Chemical Society
The CF3C(O)O2 Radical Reaction with NO
J. Phys. Chem., Vol. 100, No. 11, 1996 4515
subsequent loss via
CF3C(O)O f CF3 + CO2
(5)
CF3 + O2 + M f CF3O2 + M
(6)
CF3C(O)O2 + CF3O2 f CF3C(O)O + CF3O + O2 (7) CF3O2 + CF3O2 f 2CF3O + O2
(8)
are followed by observing the changing contribution of these peroxy radicals to the UV spectrum of the reaction mixture. Concentration vs time profiles of the individual peroxy species are determined by fitting the spectra at various times following the photolysis pulse to the expression
Abs(λ,t) ) σCF3C(O)O2(λ)l[CF3C(O)O2]t + σCF3O2(λ)l[CF3O2]t + c(t)Abs(λ,∞) (9) where σn(λ) represents the optical cross section of species n and l is the path length. The last term in eq 9 accounts for the increasing contributions of stable products to the absorbance as they accumulate. c(t), which is proportional to the concentration of products, is varied along with [CF3C(O)O2]t and [CF3O2]t in order to achieve the best fit to the absorbance. Its affect on the peroxy radical concentrations is included as part of the 2σ fitting error. The reaction between CF3C(O)O2 radicals and NO is studied by measuring NO loss and NO2 formation using transient IR absorption. This procedure was chosen in lieu of using UV spectroscopy to monitor CF3C(O)O2, because it is more specific to reaction 1. In contrast, peroxy radical loss occurs both via reaction 1 and by the competing peroxy radical self-reaction. NO is monitored via lines in the P branch of the V ) 0 f 1 transition in the vicinity of 1850 cm-1. Absorption features in the R branch of the ν3 asymmetric stretch near 1630 cm-1 are used to probe NO2. Radiation at the required IR wavelengths is provided by a lead-salt diode laser, which is frequency locked to the desired absorption line and directed through the reaction vessel counter to the photolysis pulse. The IR light passes through a monochromator for mode selection and to filter stray light and is then detected by a HgCdTe detector with a 0.2-µs response time. CF3CHO was prepared by the method of Berney.10 It was purified by several freeze-pump cycles, checked by comparing its IR spectrum against those reported by Berney10 and by Dodd et al.11 and stored at -196 °C. During the experiments, it was kept at 0 °C to minimize its polymerization. Nitrogen (99.999%) and oxygen (99.8%) were obtained from Michigan Airgas. Nitric oxide (99%) was obtained from Matheson as a 10.1% mixture in N2. Nitrogen, oxygen, and NO gas flows were controlled by Tylan flow controllers, whereas the flows of CF3CHO, F2, and Cl2 were set by needle valves. Flow rates were determined by measuring the rate of pressure increase into a constant volume. The total pressure was monitored at the entrance and exit of the reaction cell. The cell temperature was maintained by a Neslab ULT-80dd temperature regulator, with the gas mixture precooled/preheated prior to entering the cell. III. Results A. UV Spectrum. The formation of CF3C(O)O2 radicals proceeds in two steps: attack of CF3CHO by F atoms followed by O2 addition to the resultant trifluoroacetyl radical. A quantitative measurement of UV absorption cross sections requires that these formation steps be fast compared to the loss
Figure 1. UV absorption spectrum of CF3C(O)O2. The spectrum (solid line) consists of two bands: a strong feature with an absorption maximum of 7.1 × 10-18 cm2 at 210 nm and a weaker band with an onset of about 300 nm and a maximum that is obscured by the strong band. The dashed line represents a fit of the spectrum to two Gaussianshaped absorption bands (see text).
of peroxy radicals. This was ensured by examining the absorbance of the reaction mixture 20 µs after F2 photolysis as a function of CF3CHO and O2 concentrations. The absorbance increased with the addition of each reactant, reaching a plateau for > 1 Torr of CF3CHO and > 10 Torr of O2, at which concentrations the peroxy radical formation has an expected half-life of ≈1 µs. Loss of CF3C(O)O2 by self-reaction occurs with a rate constant of 7.9 × 10-12 cm3 s-1 at 295 K (see section IIIB), implying that after 20 µs approximately 9% of the radicals are converted to CF3O2, given a typical initial concentration of 3 × 1014 cm-3 CF3C(O)O2. Using the known UV spectrum12 of CF3O2, the measured absorbance at 20 µs is corrected to zero time in order to extract the CF3C(O)O2 spectrum. The formation of FO2 by the addition of O2 to fluorine atoms (k ) 4.4 × 10-33 cm6 s-1)13 poses a possible source of interference in the determination of the CF3C(O)O2 UV spectrum, since it absorbs strongly, about 3 times that of a typical peroxy radical, over the 190-225-nm region. This interference is minimized by the use of moderate total pressures (≈170 Torr) and sufficient quantities of CF3CHO (>2 Torr), under which conditions [FO2]/[CF3C(O)O2] < 0.005. Absolute absorption cross sections are assigned by comparing the CF3C(O)O2 absorbance, corrected to zero time, to the absorbance of ethylperoxy radicals8 recorded under identical conditions, except for the substitution of ≈2 Torr of ethane for ≈2 Torr of CF3CHO. This radical is ideal for intensity calibration because it is rapidly formed, but its removal via selfreaction is very slow.14 Furthermore, three recent diode array spectra of C2H5O2 yield cross sections in very good agreement with each other;8,15,16 thus, its contribution to the error in the intensity of the CF3C(O)O2 band is estimated at 5 × 104 s-1 b k ) 4.0 × 10-12(T/300)-1 cm3 s-1 k ) (0-2) × 10-12 cm3 s-1 b k ) 1.8 × 10-12 cm3 s-1 20 k ) 1.4 × 10-11 cm3 s-1 20 k < 2 × 10-14 cm3 s-1 b k ) 1.4 × 10-11 cm3 s-1 21
a
with the latter two reactions accounting for the CF3O2 decay. Reaction 6, the addition of O2 to CF3, is sufficiently fast at the concentrations of O2 employed in the experiments to be essentially instantaneous. Although only a lower limit5 has been reported for the rate of CF3C(O)O decomposition, a rough estimate suggests that it, too, is sufficiently rapid to appear instantaneous on our experimental time scale. Ab initio computations by Francisco22 reveal the dissociation to proceed with an activation energy of 4.9 kcal/mol, which, when combined with an A factor of 2 × 1013 s-1, leads to an estimated
Reaction numbers correspond to those used in the text. Rate constants are taken from ref 13 unless otherwise specified. b Measured in the present study.
high-pressure limiting rate of 5 × 109 s-1 at 295 K. The present data are consistent with this estimate; a lower limit of 5 × 104 s-1 can be established for the dissociation rate at 150 Torr of N2 from the lack of an observable delay in the CF3O2 rise relative to the CF3C(O)O2 decay. Temporal profiles of the peroxy radicals, such as those depicted in Figure 3, were compared to predictions from the
The CF3C(O)O2 Radical Reaction with NO
J. Phys. Chem., Vol. 100, No. 11, 1996 4517
TABLE 3: Self-Reaction Rate Constants conditions temp, K
CF3CHO, Torr
F2, Torr
O2, Torr
Ptot, Torr
[F]0, 1014 cm-3
results k4 10-11 cm3 s-1
210 213 233 253 273 295
3.2 3.7 3.6 4.3 4.4 3.3
3.3 3.2 3.5 3.9 3.9 3.9
8.3 11 9.0 9.2 9.4 9.2
151 148 159 170 172 175
3.8 3.6 3.3 7.8 7.7 3.6
1.2 ( 0.5 1.4 ( 0.4 1.1 ( 0.3 1.2 ( 0.3 1.1 ( 0.3 0.8 ( 0.2
model of Table 2, treating the rate constants for CF3C(O)O2 self-reaction and its cross reaction with CF3O2 as adjustable parameters. In all cases, the best fit to the data is obtained with a rate constant for the cross reaction of k7 ) 0. After accounting for possible systematic errors, an upper limit of k7 e 2 × 10-12 cm3 s-1 is reported. Rate constants for the self-reaction, k4, are tabulated in Table 3 along with the experimental conditions. The overall 2σ error bars range from 20% to 40%. The major fraction (15-30%) arises from noise in the data (fitting error), which originates from separating the overlapping contributions of CF3C(O)O2 and CF3O2 to the overall absorption of UV light by the reaction mixture. Smaller contributions to the error bars of about 10% each arise from uncertainties in the CF3C(O)O2 and CF3O2 UV cross sections and from the initial radical concentrations. As described in section IIIA, the latter is ascertained by calibration against the production of ethylperoxy radicals. The temperature dependence of the CF3C(O)O2 self-reaction rate constant is illustrated in Figure 4. It has at best a slight +5 ) × negative temperature dependence given by k4 ) (3.7-2 10-12 e(270(200)/T cm3 s-1. Both the magnitudes of the rate constants and their variation with temperature are typical of peroxy radical self-reactions.14 For comparison, the roomtemperature rate constant is about 50% smaller than that of the analogous CH3C(O)O2 self-reaction.17,18 C. Reaction of CF3C(O)O2 with NO. The reaction between CF3C(O)O2 and NO was followed by IR probes of NO loss and NO2 formation. Figure 5 illustrates the concentration vs time dependence of these species at 324 K and for two initial NO concentrations. The NO and NO2 traces were obtained in “back-to-back” experiments run under identical conditions; only the frequency of the diode laser was adjusted. As expected, the rate of NO2 formation matches the rate of NO loss, and the rates are faster at the higher initial NO concentration. Of note are the relative values of the initial radical concentration, the amount of NO lost, and the NO2 formed. At the higher [NO]0 level in Figure 5, [NO2]∞ > [Cl]0 and ∆[NO]∞ = -2[Cl]0; clearly additional chemistry besides reaction 1 plays a role in determining the NOx levels. Nitric oxide is involved in two secondary reactions. The CF3C(O)O product of the primary CF3C(O)O2 + NO reaction rapidly dissociates into CF3 and CO2. The former fragment adds oxygen to generate a secondary peroxy radical which itself converts NO to NO2:
CF3O2 + NO f CF3O + NO2
(12)
Furthermore, the trifluoromethoxy radical that is formed subsequently reacts with nitric oxide
CF3O + NO f CF2O + FNO
(13)
to generate carbonyl fluoride and FNO. Thus, one expects from reactions 1, 12, and 13 to observe a [Cl]0:∆[NO]∞:[NO2]∞ ratio of 1:-3:2. Instead, the observed ratio is closer to 1:-2.3:1.3. Possible reasons are (a) removal of peroxy radicals via their
Figure 4. Variation of CF3C(O)O2 self-reaction rate constant with temperature. Error bars and dotted lines represent deviations of 2σ.
Figure 5. Time dependence of NO loss and NO2 formation for the reaction of CF3C(O)O2 with NO. Both sets of data are fit simultaneously, treating k1 and k5 as adjustable parameters.
self-reaction and cross reactions, (b) loss of peroxy radicals by reaction with chlorine atoms, (c) reactions of the peroxy radicals CF3C(O)O2 and CF3O2 with the NO2 formed by reactions 1 and 12, and (d) a competition between the dissociation of CF3C(O)O and its reaction with NO. The first two possibilities involve reactions that compete with the reaction between CF3C(O)O2 and NO. The second of these,
Cl + RO2 f RO + ClO
(14)
is effective when peroxy radical formation is limited by Cl attack of the precursor, in this case CF3CHO, as opposed to the O2 addition step, reaction 3. When this occurs, peroxy radicals and chlorine atoms are present simultaneously and are removed via reaction 14, which is typically very fast.28 Both a and b can explain a decrease in ∆[NO]∞ and [NO2]∞ relative to [Cl]0; however, neither can effectively rationalize a higher than expected ratio of ∆[NO]∞ to [NO2]∞. Furthermore, since these reactions compete with reaction 1, they should become less effective at influencing the NO and NO2 levels as [NO]0 is raised; this is not observed. Possibilities c and d can explain a smaller than expected NO2 yield as compared to NO loss. According to c, NO2 already
4518 J. Phys. Chem., Vol. 100, No. 11, 1996
Maricq et al.
TABLE 4: CF3C(O)O2 + NO Reaction Mechanism reactiona
rate constant
Initiation Cl + CF3CHO f CF3CO + HCl CF3CO + O2 + M f CF3C(O)O2 + M
k ) 6.1 × 10-11e-955/T cm3 s-1 23 k ) 7.3 × 10-13 cm3 s-1 5
Self-Reactions CF3C(O)O2 + CF3C(O)O2 f 2CF3C(O)O + O2 CF3C(O)O f CF3 + CO2 CF3 + O2 + M f CF3O2 + M CF3C(O)O2 + CF3O2 f CF3C(O)O + CF3O + O2 CF3O2 + CF3O2 f 2CF3O + O2 CF3O2 + CF3O f CF3OOOCF3 CF3O + CF3CHO f CF3CO + CF3OH CF3O + CF3O f products
k ) 3.7 × 10-12e270/T cm3 s-1 b k > 5 × 104 s-1 b k ) 4.0 × 10-12(T/300)-1 cm3 s-1 k ) (0-2) × 10-12 cm3 s-1 b k ) 1.8 × 10-12 cm3 s-1 20 k ) 1.4 × 10-11 cm3 s-1 20 k < 2 × 10-14 cm3 s-1 b k ) 1.4 × 10-11 cm3 s-1 21
13.
NOx Reactions CF3C(O)O2 + NO f CF3C(O)O + NO2 CF3O2 + NO f CF3O + NO2 CF3C(O)O + NO f products CF3C(O)O2 + NO2 f CF3C(O)O2NO2 CF3O2 + NO2 f CF3O2NO2 CF3O + NO f CF2O + FNO
k ) 4.0 × 10-12e(563)/T cm3 s-1 b k ) 1.5 × 10-11(T/298)-1.2 cm3 s-1 20 k ) 3 × 10-11 cm3 s-1 24 k ) 6.6 × 10-12(T/298)-1 cm3 s-1 25 k ) 8.9 × 10-12(T/298)-0.72 cm3 s-1 20 k ) 4.2 × 10-11e120T cm3 s-1 26,27
14. 14.
RO2 + Cl Cl + CF3C(O)O2 f CF3C(O)O + ClO Cl + CF3O2 f CF3O + ClO
k ) 1 × 10-10 cm3 s-1 (see text) k ) 1 × 10-10 cm3 s-1 (see text)
2. 3. 4. 5. 6. 7. 8. 11.
1. 12. 15.
a Reaction numbers correspond to those used in the text. Rate constants are taken from ref 13 unless otherwise specified. b Measured in the present study.
TABLE 5: CF3C(O)O2 + NO Reaction Constants conditions temp, K
CF3CHO, Torr
Cl2, Torr
O2, Torr
Ptot Torr
220 220 254 254 295 295 324 324
20 20 17 17 16 16 12 12
0.29 0.29 0.28 0.28 0.28 0.28 0.34 0.34
12 12 12 12 12 12 13 13
104 104 100 100 102 102 101 101
formed is removed by reaction with peroxy radicals, thereby lowering its yield relative to ∆[NO]∞. By explanation d, the secondary peroxy radical CF3O2 is not formed. Instead, CF3C(O)O is removed by reaction with NO, whereby the ratio of ∆[NO]∞ to [NO2]∞ is -2:1 rather than -3:2. Both mechanisms lead to good fits of the ∆[NO]t and [NO2]t data, yielding essentially the same values of k1. However, the rate constant required for the reaction between CF3C(O)O2 and NO2 in order to obtain a good fit to the data using explanation c varies from 5 to 30 times the accepted value. We find that the observed NOx changes are best explained by d. Thus, we postulate that the reaction
CF3C(O)O + NO f products
(15)
competes with CF3C(O)O dissociation and fit simultaneously the NO loss and the NO2 formation data to the model in Table 4, treating k1 and k5 as fitting parameters. [Cl]0 is measured separately by substituting ethane for CF3CHO, as explained in section IIIA. The results are collected in Table 5, and representative fits are illustrated by the solid lines in Figure 5. The slight delay observed in NO2 formation may be due to a fraction being formed in excited vibrational states, which rapidly relax via collisions to the ground state. However, the ability to fit successfully both the NO and NO2 traces with the same values of k1 and k5 adds confidence to the reaction model adopted in Table 4. Because it is based on an assigned rate constant for reaction 15,24 k5 must be considered as a relative as opposed to an absolute value. However, as reference to Table
results [Cl]0
1014 8.4 8.4 6.8 6.8 5.7 5.7 7.5 7.5
cm-3
NO,
1015 6.7 12.4 5.7 11.0 5.2 9.7 4.8 9.3
cm-3
k1,
10-11
cm3 s-1
5.6 ( 1.0 5.3 ( 1.3 3.4 ( 0.6 3.3 ( 0.8 2.7 ( 0.5 2.8 ( 0.5 2.6 ( 0.5 2.1 ( 0.4
k5, 105 s-1 1.3 0.9 1.9 0.9 2.3 2.5 2.2 1.9
5 shows, the values of k5 derived in this manner are consistent with the lower limit of 5 × 104 s-1 found above from the CF3C(O)O2 self-reaction experiments and the lower limit of 6 × 104 s-1 reported by Wallington et al.5 Contributions to the error in k1 arise from signal noise, the measurement of [Cl]0, uncertainties in σir(NO) and σir(NO2), and uncertainties in the reaction model, primarily in k2, k4, and k14. Noise in the absorption vs time traces introduces approximately a 4-20% error into k1. A 5% uncertainty in [Cl]0, arising from the ethylperoxy UV cross section, contributes a 9% error to the rate constant. Uncertainty in the IR cross section of NO adds another 10% to the error in k1, whereas the contribution from σir(NO2) is about 3% (it has a larger effect on k5). A delay in CF3C(O)O2 formation owing to the slow attack of CF3CHO by Cl was found to lead to underprediction of k1. For example, reduction of CF3CHO from 16 to 4.3 Torr leads to an apparent decrease of k1 from 2.8 × 10-11 to 1.5 × 10-11 cm3 s-1 at 295 K. At the higher CF3CHO concentration, a 50% uncertainty in k2 leads to an error of 9.9 × 10-12 cm3 s-1 reported from the pulsed radiolysis work of Wallington et al.5 As illustrated by Figure 6, the reaction exhibits a negative temperature dependence with k1 ) +22 ) × 10-12 e(563(115)/T cm3 s-1. This dependence sug(4.0-1.4 gests that the reaction proceeds via complex formation followed by rapid rearrangement and dissociation. While an increase in rate constant with declining temperature is generally observed for peroxy radical-NO reactions, the typical dependence, with EA ∼ 250 K, is more modest than found here.14 IV. Atmospheric Implications The atmospheric oxidation of HFCs and HCFCs initiated by the reaction of these compounds with OH radicals generates a variety of halogen-substituted aldehydes and acid aldehydes, amongst them CF3CHO. This molecule is itself subject to OH attack, yielding a series of CF3COx radicals. The purpose of this, and a previous paper,4 has been to investigate the kinetics of reactions relevant to the atmospheric fate of these radicals. The CF3CO radical formed by the reaction of OH with CF3CHO collisionally dissociates with an activation energy of approximately 12 kcal/mol. Thus, the dissociation rate constant falls from a value of about 105 s-1 at the earth’s surface (T ) 295 K, P ) 760 Torr) to roughly 300 s-1 at an altitude of 20 km (T ) 215 K, P ) 45 Torr). The dissociation process competes with oxygen addition to CF3CO to produce the corresponding peroxy radical. While the latter rate constant is unknown, oxygen addition to alkyl radicals is typically of the order 4 × 10-12 cm3 s-1. This reaction, too, is temperature and pressure dependent, but both dependencies are expected to be modest compared to those for the dissociation process. Furthermore, the expected negative temperature dependence should cancel the effect of decreasing pressure, rendering the rate constant relatively independent of altitude. A comparison of the dissociation rate to the O2 addition rate reveals that ∼99.5% of the CF3CO radicals are converted to peroxy radicals at 0 km and >99.9% are converted at 20 km. Peroxy radicals are removed from the atmosphere by reactions with NO, NO2, HO2, and other peroxy radicals. Assuming a tropospheric NO concentration29 of 4 × 108 cm-3 leads to an atmospheric lifetime of ∼1.5 min for CF3C(O)O2 with respect to removal by NO; the lifetime falls to about 0.3 min at 20 km, with [NO] = 109 cm-3 and k1 ) 5 × 10-11 cm3 s-1. For comparison, the reported rate constant for the reaction of CF3C(O)O2 with NO2 is 6.6 × 10-12 cm3 s-1 at 295 K and 1
J. Phys. Chem., Vol. 100, No. 11, 1996 4519 atm of SF6.25 Taking this as an upper limit to the rate constant under atmospheric conditions and assuming [NO2] ) 4 × 108 cm-3 provides a lifetime of >6 min for CF3C(O)O2 with respect to reaction with NO2; assuming the rate constant to remain constant with altitude, the lifetime at 20 km is about 2 min. The rate constant for the reaction between HO2 and CF3C(O)O2 is unknown; as an estimate, we assume at 295 K the value of 1.4 × 10-11 cm3 s-1 from comparison to CH3C(O)O2. Using an HO2 concentration29 of 2 × 108 cm-3 leads to a lifetime of about 6 min. If the temperature dependence of CF3C(O)O2 + HO2 mimics that of CH3C(O)O2, then at 20 km, where the HO2 concentration has fallen to 5 × 106 cm-3, the lifetime of CF3C(O)O2 with respect to HO2 increases to 60 min. The only CF3C(O)O2-peroxy radical reaction that has been measured is the self-reaction. If the rate constant of ∼1.5 × 10-11 cm3 s-1 is taken as typical and if the atmospheric RO2 concentration is assumed to be comparable to the HO2 concentration, then the lifetime for removal by RO2 reactions will also be roughly 6 min. Comparing the above atmospheric lifetimes indicates that, in the lower troposphere CF3C(O)O2 radicals will mainly be removed by reaction with NO, but the reactions with NO2, HO2, and RO2 will contribute significantly. Based on the presently available data, approximately 55% of the removal will be by NO, 15% by NO2, 15% by HO2, and 15% by RO2. In contrast, the contributions from HO2 and RO2 will decrease significantly at 20 km, where 85% of the removal is expected to by via NO reaction and 15% by NO2. The CF3C(O)O radical that is produced in reaction 1 rapidly dissociates. Even at 233 K, the self-reaction data and the NO reaction experiments are consistent with a dissociation rate of >5 × 104 s-1. This can be compared with a possible reaction between CF3C(O)O and NO, for which, even assuming a rate constant of 5 × 10-11 cm3 s-1, the atmospheric rate would be of the order 10-2 s-1. Thus, in the atmosphere, NO effectively converts CF3C(O)O2 radicals into CF3O2 radicals and CO2. The chemistry of CF3O2, and the CF3O radical that is subsequently formed, has itself been the subject of intense scrutiny over the past few years due to suggestions that these radicals can participate in ozone depletion cycles.20 However, subsequent work has shown that CF3O2 radicals are principally converted, via CF3O, to CF2O and FNO;26,27 thus, the major fate of CF3C(O)O2 is conversion into CO2, CF2O, and FNO. V. Conclusion As typical of peroxy radicals, CF3C(O)O2 exhibits a strong UV absorption in the range 190-300 nm. The spectrum consists of two Gaussian-shaped bands, such as is observed for CH3C(O)O2; however, in this case, the band centers are almost coincident at 207 and 209 nm, but with the latter band much broader than the first. The long-wavelength tail of the CF3C(O)O2 spectrum allows it to be distinguished from CF3O2 produced by self-reaction of the former radicals. Time-resolved UV spectroscopy provides a means of monitoring the progress of the self-reaction and of the cross reaction between CF3C(O)O2 and CF3O2 radicals. The self-reaction exhibits a slight negative temperature dependence typical of peroxy radicals. In contrast, the cross reaction is significantly slower than expected. As observed by both the loss of NO and the production of NO2, CF3C(O)O2 reacts rapidly with NO, with rate constants ranging from 2.4 × 10-11 cm3 s-1 at 324 K to 5.5 × 10-11 cm3 s-1 at 220 K. Consideration of the atmospheric concentrations of the principal species involved in peroxy radical removal reactions, NO, NO2, HO2, and RO2, along with the available
4520 J. Phys. Chem., Vol. 100, No. 11, 1996 rate constants (or estimates), suggests that NO plays the major role in the fate of CF3C(O)O2 in the lower troposphere and dominates its removal in the stratosphere. References and Notes (1) Maricq, M. M.; Szente, J. J. J. Chem. Phys. 1995, 99, 4554. (2) George, Ch.; Saison, J. Y.; Ponche, J. L.; Mirabel, Ph. J. Phys. Chem. 1994, 98, 10857. (3) Dobe, S.; Kachatryan, L. A.; Berces, T. Ber. Bungenges. Phys. Chem. 1989, 93, 847. Scollard, D. J.; Treacy, J. J.; Sidebottom, H. W.; Balestra-Garcia, C.; Laverdet, G.; Lebras, G.; MacLeod, H.; Teton, S. J. Phys. Chem. 1993, 97, 4683. (4) Maricq, M. M.; Szente, J. J.; Khitrov, G. A.; Dibble, T. S.; Francisco, J. S. J. Phys. Chem. 1995, 99, 11875. (5) Wallington, T. J.; Hurley, M. D.; Nielsen, O. J.; Sehested, J. J. Phys. Chem. 1994, 98, 5686. (6) Wallington, T. J.; Hurley, M. D. Int. J. Chem. Kinet. 1993, 25, 819. (7) Maricq, M. M.; Szente, J. J. J. Phys. Chem. 1992, 96, 10862. (8) Maricq, M. M.; Wallington, T. J. J. Phys. Chem. 1992, 96, 986. (9) Maricq, M. M.; Szente, J. J.; Kaiser, E. W. J. Phys. Chem. 1993, 97, 7970. (10) Berney, C. V. Spectrochim. Acta, Part A, 1969, 25, 793. (11) Dodd, R. E.; Roberts, H. L.; Woodward, C. A. J. Chem. Soc. 1957, 2783. (12) Maricq, M. M.; Szente, J. J. J. Phys. Chem. 1992, 96, 4925. (13) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Hampton, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R.; Kolb, C. E.; Molina, M. J. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling. JPL Publication 92-20, 1992. (14) Lightfoot, P. D.; Cox, R. A.; Crowley, J. N.; Destriau, M.; Hayman, G. D.; Jenkin, M. E.; Moortgat, G. K.; Zabel, F. Atmos. EnViron. 1992, 26A, 1805.
Maricq et al. (15) Bauer, D.; Crowley, J.; Moortgat, G. K. J. Photochem. Photobiol. A: Chem. 1992, 65, 329. (16) Fenter, F. F.; Catoire, V.; Lesclaux, R.; Lightfoot, P. D. J. Phys. Chem. 1993, 97, 3530. (17) Moortgat, G.; Veyret, B.; Lesclaux, R. J. Phys. Chem. 1989, 93, 2362. (18) Maricq, M. M.; Szente, J. J. J. Phys. Chem. 1996, 100, 4507. (19) Joens, J. A. J. Phys. Chem. 1994, 98, 1394. (20) Francisco, J. S.; Maricq, M. M. AdV. Photochem. 1995, 20, 79. (21) Taken as “typical” alkoxy self-reaction rate constant. (22) Francisco, J. S. Chem. Phys. Lett. 1992, 191, 7. (23) Based on the ab initio activation energy of 1.9 kcal/mol from Dibble, T. S.; Francisco, J. S. Chem. Phys. Lett. 1993, 215, 409, and on the room-temperature rate constant reported by Wallington and Hurley.6 (24) Taken by comparison with the rate constant for FC(O)O + NO: Maricq, M. M.; Szente, J. J.; Dibble, T. S.; Francisco, J. S. J. Phys. Chem. 1994, 98, 12294. (25) Wallington, T. J.; Sehested, J.; Nielsen, O. J. Chem. Phys. Lett. 1994, 226, 563. (26) Turnipseed, A. A.; Barone, S. B.; Ravishankara, A. R. J. Phys. Chem. 1994, 98, 4594. (27) Dibble, T. S.; Maricq, M. M.; Szente, J. J.; Francisco, J. S. J. Phys. Chem. 1995, 99, 17394. (28) Maricq, M. M.; Szente, J. J.; Kaiser, E. W.; Shi, J. J. Phys. Chem. 1994, 98, 2083. (29) Cantrell, C. A.; Shetter, R. E.; Calvert, J. G.; Parrish, D. D.; Fehsenfeld, F. C.; Goldan, P. D.; Kuster, W.; Williams, E. J.; Westberg, H. H.; Allwine, G.; Martin, R. J. Geophys. Res. 1993, 98, 18355.
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