J. Phys. Chem. 1996, 100, 6587-6593
6587
Thermal Decomposition of CF3O2NO2 A. Mayer-Figge, F. Zabel,* and K. H. Becker Bergische UniVersita¨ t-Gesamthochschule Wuppertal, Physikalische Chemie/FB 9, 42097 Wuppertal, Germany ReceiVed: October 30, 1995; In Final Form: January 25, 1996X
The unimolecular decomposition rate constant of CF3O2NO2 has been measured in detail as a function of temperature, pressure, and collision partner (M ) N2, O2, NO). Temperatures were between 264 and 297 K, and total pressures ranged from 3 to 1013 mbar. The first-order decay of CF3O2NO2 in the presence of excess NO was followed in a temperature-controlled DURAN glass chamber by long-path IR absorption, using the absorption bands at 1768 and 1303 cm-1. At 1013 mbar, the first-order decomposition rate constants are best represented by the Arrhenius expression k3 ) 5.7 × 1015 exp{(-97.7 ( 1.0) kJ mol-1/RT} s-1 (2σ). The temperature and pressure dependencies of k3 are well reproduced by the equation log(k3/k3,∞) ) log{(k3,0/k3,∞)/(1 + k3,0/k3,∞)} + log(Fc){1 + [log(k3,0/k3,∞)/Nc]2}-1, Nc ) 0.75-1.27 log(Fc) with the parameters k3,0/[N2] ) 2.4 × 10-5 exp(-78.4 kJ mol-1/RT) cm3 molecule-1 s-1, k3,∞ ) 1.49 × 1016 exp{(-99.3 ( 1.3) kJ mol-1/RT} s-1, Fc ) 0.31, and k3,0(M)O2) ≈ k3,0(M)N2). By combining the present decomposition rate constants with recombination rate constants k-3 from Caralp et al., the following thermochemical data for the equilibrium CF3O2NO2 S CF3O2 + NO2 (k3,k-3) are derived from second- and third-law evaluations: ∆H°r,298 ) 102.7 ( 2.0 kJ mol-1, ∆S°r,298 ) 163 ( 7 J mol-1 K-1. The temperature dependence of the equilibrium constant between 200 and 300 K is described by the expression Kc ) k3/k-3 ) 3.80 × 1027 exp{(-12140 ( 240)K/T} molecules cm-3. Consistency of the data on k3 (this work) and k-3 is shown by comparing experimental and theoretical limiting low-pressure rate constants, which lead to the reasonable value βc ) 0.17 for the collision efficiency of N2. The present data confirm that CF3O2NO2 is thermally quite stable in the upper troposphere and lower stratosphere and that its lifetime is probably limited by photolysis in these regions of the atmosphere.
Introduction
Experimental Section
Trifluoromethyl peroxynitrate (CF3O2NO2) is a potential intermediate in the atmospheric degradation of fully halogenated hydrocarbons (e.g., CF4, CF3Cl, CF3Br, CF3I) and of substitutes replacing these compounds (CHF3, CH3CF3, CH2FCF3 (R 134a), CHF2CF3, CHCl2CF3, CHCl2CF2CF3, etc.). Pure CF3O2NO2 has been synthesized and characterized by Hohorst and DesMarteau.1 It was identified by its IR spectrum as an intermediate in the photolysis of CF3NO in the presence of O2 and NO2 by Chen et al.2 and Wallington et al.3 Its lifetime in the atmosphere depends mainly on the rates of thermal decomposition and photolysis. So far, experiments have revealed only a lower limit for the thermal decomposition rate constant of CF3O2NO2 at room temperature and atmospheric pressure.3,4 Experiments on the thermal stability of other chloro(fluoro)methyl peroxynitrates5-10 show that the energies of the weakest bond in these molecules are on the order of 100 kJ mol-1. From these data it may be estimated that CF3O2NO2 is thermally fairly longlived at the temperatures of the upper troposphere and lower stratosphere. In fact, in these regions of the atmosphere the lifetime may be limited by photolysis.11,12 However, the UV absorption coefficients of CF3O2NO2 have been published only up to wavelengths of 280 nm,11 and the extrapolation to longer wavelengths introduces considerable uncertainties to the photolysis lifetimes derived from these data. For this reason, kinetic data on both thermal decomposition and photolysis are needed to reliably estimate atmospheric lifetimes of CF3O2NO2. In the present work, thermal decomposition rate constants of CF3O2NO2 have been measured in detail as a function of temperature, pressure, and collision partner (M ) N2, O2, NO), and the data are evaluated in terms of unimolecular reaction rate theory.
The experiments were performed in a temperature-controlled 420 l Duran glass photoreactor (chamber I) which was described elsewhere13 in more detail. CF3O2NO2 was prepared in situ by photolyzing CF3I/O2/NO2/N2 mixtures. Typical initial concentrations were 1 × 1015 molecules cm-3 CF3I, (1-2) × 1014 molecules cm-3 NO2 which was added in several portions during the photolysis time in order to reduce the photolysis of NO2, 1.3 × 1016 molecules cm-3 of O2, and 1000 mbar of N2. Since the photolysis rate of CF3I was very low in chamber I (λ g 300 nm), the photolysis was performed at 253.6 nm (σ253.6(CF3I) ) 4.3 × 10-19 cm2 molecule-1 (ref 14) in a second 420 L chamber (chamber II) equipped with three internal 40 W lowpressure mercury lamps. CF3O2NO2 is formed by the following mechanism:
X
Abstract published in AdVance ACS Abstracts, March 15, 1996.
0022-3654/96/20100-6587$12.00/0
CF3I + hν w CF3 + I
(1)
CF3 + O2 (+M) w CF3O2 (+M)
(2)
CF3O2 + NO2 (+M) S CF3O2NO2 (+M) (3,-3) During photolysis, a small amount of CF2O was formed in side reactions, probably initiated by the slow photolysis of NO2. CH3ONO2 which could complicate the kinetic analysis of reaction 3 was observed as a major product in the photolysis of CF3I/ O2/NO/NO2/N2 mixtures with initial [NO]/[NO2] ratios of ≈1 but was not detectable under our standard photolysis conditions. Subsequent to photolysis, a valve was opened between chamber II and the evacuated chamber I which was kept at the desired reaction temperature. The gas mixture in chamber I was then pressurized with N2 to the total pressure of the respective experiment. At the selected reaction temperatures, the concen© 1996 American Chemical Society
6588 J. Phys. Chem., Vol. 100, No. 16, 1996
Mayer-Figge et al.
Figure 1. IR absorption spectrum of CF3O2NO2; the spectrum is strongly perturbed by imperfectly subtracted absorption from the precursor CF3I between 1178 and 1192 cm-1.
trations of CF3O2 present in the equilibrium (3,-3) were so small that the self-reaction of CF3O2,
CF3O2 + CF3O2 w 2CF3O + O2
(4)
(k4,295K ) (1.8 ( 0.5) × 10-12 cm3 molecule-1 s-1 (ref 15)), was unimportant. Thermal decomposition was initiated by addition of 5 × 1015 molecules cm-3 of NO, thereby withdrawing CF3O2 radicals from the equilibrium (3,-3) via reaction 5:
CF3O2 + NO w CF3O + NO2
(5)
The CF3O2 loss in reaction 5 is irreversible because CF3O rapidly reacts with NO to form the stable product CF2O:16,17
CF3O + NO w CF2O + FNO
(6)
(k6,295K ) 1.7 × 10-11 cm3 molecule-1 s-1 (ref 17)). In the course of the reaction, the [NO]/[NO2] ratio decreased by about a factor of 2 due to the NO2 formed in reactions 3 and 5. However, at the end of the reaction time [NO]/[NO2] was always larger than 10. Since k5 is higher than k-3 by at least a factor of 1.6,10 the relationship k5[NO] . k-3[NO2] was valid throughout the reaction. The relative CF3O2NO2 concentrations were monitored as a function of time by long-path IR absorption (optical path length 50.4 m), using a built-in White mirror system and an FTIR spectrometer (Nicolet Magna 550). The effective first-order rate constants for CF3O2NO2 loss after NO addition, keff, were corrected for wall loss. Individual wall loss rate constants kw were determined in each experiment from the slow first-order loss rate of CF3O2NO2 before NO addition. Numerical simulation of the reaction mechanism using the Larkin kinetic modeling program18 established that CF3O2 loss by self-reaction was not responsible for the slow decrease of CF3O2NO2 before the addition of NO. For this reason, the difference keff - kw was identified as the gas-phase unimolecular decomposition rate constant k3. kw amounted to (0-3) × 10-5 s-1, showing a small systematic increase with increasing temperature and decreasing pressure.
Research grade CF3I (Aldrich, 99%) and NO2 were used as supplied. Research-grade NO was trapped at liquid nitrogen temperature and purified by distillation. Pressures were measured using capacitance pressure gauges. The temperature in the reaction chamber was measured with a platinum resistance gauge. The temperature was constant within (0.1 K during an experiment. Results and Data Reduction IR Absorption Spectrum of CF3O2NO2. The IR absorption spectrum of CF3O2NO2 is shown in Figure 1. This spectrum was obtained from the product spectrum of the 253.6 nm photolysis of a mixture of 0.053 mbar of CF3I, 1.3 mbar of synthetic air, 0.008 mbar of NO2, and 1011 mbar of N2 by spectrally subtracting absorptions of CF3I, CF2O, and NO2. The band maxima/double maxima were at the following positions (in cm-1; see Figure 1 for comparison): 3057, 1768/1762, 1303, 1250/1245, ≈1190, 959, 792, 710, ≈675. These band positions are in good agreement with all the bands denoted to be “very strong” and “strong” by Hohorst and DesMarteau1 and with the work of Chen et al.2 The 1190 cm-1 band is strongly perturbed due to imperfectly subtracted absorption of the precursor CF3I. IR absorptions at 1768 and 1303 cm-1 were used to monitor the concentration of CF3O2NO2 as a function of time. Kinetic Results on the Thermal Decomposition of CF3O2NO2. The time dependence of the IR absorbances of CF3O2NO2 at 1768 and 1303 cm-1 was analyzed according to a firstorder rate law (Figure 2). The resulting first-order rate constants were identical for both wavelengths within error limits and were averaged. A total of 56 experiments were performed in N2 at temperatures between 264 and 297 K and total pressures between 3.5 and 1013 mbar. At low pressures, five additional experiments were performed in O2 and two experiments with much higher than normal mixing ratios of NO in order to estimate the collision efficiencies of these buffer gases. The temperature dependence of the first-order rate constants k3 at different total pressures is shown in Figure 3. Rate constants obtained at total pressures between 3.6 and 3.9 mbar
Thermal Decomposition of CF3O2NO2
Figure 2. IR absorbance of CF3O2NO2 at 1768 cm-1 as a function of time before (9) and after (b) NO addition; T ) 276.3 K, p ) 3.9 mbar, M ) N2.
J. Phys. Chem., Vol. 100, No. 16, 1996 6589
Figure 4. Falloff curves for k3 at 273, 280, and 288 K; full lines are calculated from molecular parameters and shifted to fit to the data points (see text); statistical errors for the individual data points are smaller than the size of the symbols.
TABLE 1: Arrhenius Paramters of k3 at Different Total Pressures
Figure 3. Arrhenius plots for k3 at different total pressures; straight lines are least-squares fits to the data points for M ) N2.
were converted to 3.7 mbar using the empirical relationship k3 ∝ [N2]1/2 which is deduced from the slope of the falloff curves for [M] ≈ 1 × 1017 cm3 molecule-1 (see Figure 4 for comparison). For larger total pressures, the measured rate constants were directly included in the Arrhenius plots. The scatter of the data points is very low, due to the absence of both interfering side reactions and overlapping IR absorptions from either main or side products. The Arrhenius parameters for different total pressures (1013 ( 1, 104 ( 2, 10.6 ( 0.3, 3.7 mbar for M ) N2 and 3.6 ( 0.2 mbar for M ) O2) are presented in Table 1. The activation energies decrease with decreasing total pressure, in agreement with unimolecular reaction rate theories. The pressure dependence of k3 at 288, 280, and 273 K is shown in double-logarithmic plots in Figure 4. The data points were derived from the Arrhenius parameters in Table 1. The full lines were calculated in a reduced form by application of eq I:19
log(k/k∞) ) log{(k0/k∞)/(1 + k0/k∞)} + log(Fc){1 + [log(k0/k∞)/Nc]2}-1 (I) Nc ) 0.75 - 1.27 log(Fc) k0, which is proportional to the gas density [M], and k∞ are the
M
p [mbar]
A [s-1]
Ea [kJ mol-1]
N2 N2 N2 N2 O2
1013 104.3 10.6 3.7 3.6
5.7 × 1015 9.1 × 1014 1.05 × 1014 1.85 × 1013 3.1 × 1013
97.7 ( 1.0 94.2 ( 1.0 90.8 ( 1.1 87.8 ( 1.6 89.0 ( 3.3
first-order rate constants in the low- and high-pressure limit, respectively, and Fc is a parameter which determines the broadening of the falloff curve and can be calculated from molecular parameters of the decomposing molecule (Fc ) 0.31, see below). The resulting reduced falloff curves were then fitted to the data points under the constraint that the limiting lowand high-pressure rate constants k0 and k∞ obey the Arrhenius equation. The temperature dependencies of k0 and k∞ derived from Figure 4 are represented by the expressions k3,0/[N2] ) 2.4 × 10-5 exp(-78.4 kJ mol-1/RT) cm3 molecule-1 s-1 and k3,∞ ) 1.49 × 1016 exp{(-99.3 ( 1.3) kJ mol-1/RT} s-1. Fc was calculated using vibrational frequencies of CF3O2NO2 from Hohorst and DesMarteau1 (1760, 1305, 1292, 1244, 1186, 953, 880, 783, 702, 669, 600, 562, 485, 436, 371, 284, and 253 cm-1), Nielsen et al.20 (263 cm-1 with CF3CHCl2 as the model compound), and estimated values for the torsional vibrations (120 cm-1 using CH3O2NO221 as a model compound; 40 and 35 cm-1 (refs 9 and 22)). The selected wavenumbers of the low-frequency vibrations representing the three hindered rotations are based on a combined treatment22 of experimental data on the thermal decomposition of CCl3O2NO2, CCl2FO2NO2, and CClF2O2NO2 and the corresponding recombination reactions RO2 + NO2 (+M) w RO2NO2 (+M) (see below). Since it was not clear if it is more appropriate to approximate the hindered rotations by free rotations (Fc ) 0.36, including weak collision corrections23) or torsional vibrations (Fc ) 0.26), the mean value 0.31 was adopted for Fc. For the experiments performed in 3.6 mbar of O2, there was no significant change of either the absolute value of k3 or the activation energy as compared to M ) N2 (see Figure 3). For the experiments at a total pressure of ≈3.6 mbar with NO as a major collision partner (pNO ≈ 1.5 mbar as compared to 0.2 mbar in the other experiments), there was no notable change of k3. Hence
k3(MdNO) ≈ k3(MdO2) ≈ k3(MdN2) for p, T ) constant
6590 J. Phys. Chem., Vol. 100, No. 16, 1996
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TABLE 2: Equilibrium Constants Kc ) k3/k-3 Calculated from k3 (This Work) and k-330 [M] [1017 molecules cm-3]
T [K]
k3 [10-3 s-1]
k-3 [10-12 cm3 molecule-1 s-1]
Kc ) k3/k-3 [109 molecules cm-3]
Kc(T) [109 molecules cm-3]
0.65 1.95 3.25 0.65 1.95 3.25 0.65 1.95 3.25
273 273 273 280 280 280 288 288 288
0.236 0.399 0.495 0.627 1.07 1.35 1.81 3.12 3.95
0.121 2.09 2.57 1.12 1.88 2.29 0.102 1.68 2.03
0.195 0.191 0.193 0.562 0.572 0.587 1.77 1.85 1.94
0.193
According to the previous discussion, k3 can be calculated for all atmospheric applications by using eq I and the following parameters:
k3,0/[N2] ) 2.4 × 10-5 exp(-78.4 kJ mol-1/RT) cm3 molecule-1 s-1 k3,∞ ) 1.49 × 1016 exp{(-99.3 ( 1.3) kJ mol-1/RT} s-1 Fc ) 0.31 k3(MdO2) ≈ k3(MdN2)
(II)
A consistent analysis of thermal decomposition rate constants and thermochemical properties of peroxynitrates9,22 suggests that for most of these compounds the activation energy at 1013 mbar is only a few tenths of a kilojoule below the limiting highpressure value. Hence, the error limits of Ea,∞ were identified in the present work with the error limits of the experimental activation energy at 1013 mbar, slightly enhanced by 0.3 kJ mol-1 allowing for the additional uncertainty introduced by the extrapolation to the high-pressure limit. For Ea,0, error limits are uncertain due to the large extrapolation necessary to attain the low-pressure limit. However, in the range 1-1000 mbar, the rate constants calculated by using eqs I and II are uncertain by no more than 10% at room temperature and by less than 50% over the whole range of tropospheric and stratospheric temperatures. Reaction Products. According to reactions 3, 5, and 6, NO2, CF2O, and FNO are expected to be the main reaction products of the thermal decomposition of CF3O2NO2 in the presence of NO. In fact, the concentrations of NO2 and CF2O which were already present before the addition of NO, increased during the thermal decomposition of CF3O2NO2. FNO was easily detected as a product by its prominent Q branches at 1844 and 766 cm-1.16,24 However, [FNO] rapidly reached a steady-state value, then decreased with a time constant close to that of CF3O2NO2. The loss reaction of FNO is unclear; however, it is wellknown25 that FNO readily reacts with both water, leading to HF and HONO (which, actually, were also seen in the product spectra), and glass surfaces, leading to SiF4. SiF4, if formed, could not be detected by its IR spectrum26 in the present experiments due to the strong overlap by absorption from CF3I. In one experiment, 5 × 1013 molecules cm-3 CH4 was added to check for possible interference by F atoms which would rapidly react with methane, forming CH3 and HF. However, there was no evidence from the IR spectra for the presence of any side products originating from CH3 radicals. CF3ONO2 and CF3OH could be formed via reactions 7 and 8:
CF3O + NO2 (+M) w CF3ONO2 (+M)
(7)
0.574 1.86
CF3O + HX w CF3OH + X (e.g. X ) OH, alkyl from impurities) (8) However, there was no indication from the IR spectra that either CF3ONO22 or CF3OH27-29 was produced under the reaction conditions of the thermal decomposition experiments. In addition, there was no change of the product distribution with total pressure. Discussion Kinetics and Thermochemistry. Wallington et al.3 obtained a lower limit for k3 of 0.019 s-1 (ref 4) at 295 K and atmospheric pressure which is consistent with the present value 0.029 s-1 derived from Table 1. The present results on k3 can be combined with recombination rate data of Caralp et al.30 These authors measured k-3 in the pressure range 1.3-13 mbar at three different temperatures (233, 298, and 373 K). They analyzed their data by RRKM theory and presented the temperature and pressure dependence of k-3 in terms of eq I. In Table 2, recombination rate constants k-3 obtained by interpolation of the data of Caralp et al. (Table 1 in ref 30) are listed together with the dissociation rate constants k3 derived from the present work (eq I and parameter set II), in the common temperature and pressure range of both studies. The resulting equilibrium constants Kc ) k3/k-3 are also included in Table 3 and shown as a van’t Hoff plot in Figure 5. The temperature dependence of Kc between 273 and 288 K is represented by the expression Kc ) (1.40 ( 0.23) × 1027exp{(-11860 ( 240)K/T} molecules cm-3, where the error limits reflect the uncertainties of both k3 and k-3. 0 ) 100.9 ( 2.0 kJ mol-1 The data in Table 2 lead to ∆Hr,280 (2σ). Using the above-mentioned vibrational frequencies for CF3O2NO2, CF3O2, and NO2, the reaction enthalpy can be converted from 280 to 298 K, resulting in 0 (CF3O2NO2wCF3O2+NO2) ) ∆Hr,298
100.8 ( 2.0 kJ mol-1 (“second law”) Applying a third-law analysis based on Kc(280 K) ) (5.74 ( 0.92) × 108 molecule cm-3, the frequencies from above for CF3O2NO2 and from Butler and Snelson31 for CF3O2, and the moments of inertia of CF3O2NO2 from Hohorst and DesMarteau1 and of CF3O2 from Caralp et al.:30 0 (CF3O2NO2wCF3O2+NO2) ) ∆Hr,298
104.6 ( 2.4 kJ mol-1 (“third law”) is derived where the error limits reflect the error of Kc and, in particular, the uncertainty of the frequencies of the three 0 are in agreement torsional vibrations. Both values of ∆Hr,298 within the combined error limits, demonstrating the consistency of the available data on k3 and k-3. Since the error limits of 0 values from the second- and third-law treatments the ∆Hr,298
Thermal Decomposition of CF3O2NO2
J. Phys. Chem., Vol. 100, No. 16, 1996 6591
TABLE 3: Comparison of Selected Kinetic and Thermochemical Data from Reaction 3 with Literature Data Caralp et al.30 Destriau, Troe34 k3 (280 K, 60 mbar N2) [s-1] k3 (298 K, 1013 mbar N2) [s-1] k3 (220 K, 50 mbar N2 + O2) [s-1] Kc (280 K) [molecules cm-3] 0 E0 ) ∆Hr,0 [kJ mol-1] 0 ∆Hr,298 [kJ mol-1]
3.4 ×
10-3
3.3 ×
10-3a
this work (2.3 ( 0.3) × 10-3
7.1 × 10-2
8.2 × 10-2
4.2 × 10-2
3.5 × 10-8
3.1 × 10-8
2.9 × 10-8
6.8 ×
1.19 ×
(5.74 ( 0.92) ×
108
109
(101.6)b
(101.6)c
100.5 ( 2.0
(105.0)d
(105.0)c
102.7 ( 2.0
Kc ) k3/k-3 ) 108
See also ref 37. E0(CF3O2NO2) ) E0(CF2ClO2NO2) assumed. 0 Value from ref 30 adopted. d Average of calculated ∆Hr,298 values for reactions 3 and 11-13. a
0 0 ) 102.7 ( 2.0 kJ mol-1, ∆Hr,0 ) E0 ) from ref 30, ∆Hr,298 0 -1 -1 100.5 ( 2.0 kJ mol , ∆Sr,298 ) 163 ( 7 J mol K-1. Based on these data, the Kc values in the atmospherically important temperature range 200-300 K are approximated by the following simple van’t Hoff-type expression:
b
3.80 × 1027 exp{(-12140 ( 240)K/T} molecules cm-3 (200-300 K) The internal consistency of this data analysis can be tested by comparing the experimentally derived low-pressure value of k3 ()k3,exp(pw0) ≡ k3,0, see Figure 4) with the value calculated using the equation32
c
k0 ) βck0(strong collisions) ) βc[M]ZLJFvib,harm(E0)Qvib-1 exp(-E0/RT)RTFEFanhFrot (III)
Figure 5. van’t Hoff plots of Kc ) k3/k-3; data points: mean Kc values from Table 2; broken line: least-squares fit of the experimental data from Table 2; full line: “best” straight line through the slightly curved representation of the equilibrium constants calculated for the temperature range 200-300 K from the recommended set of values for 0 0 ∆Hr,298 , ∆Sr,298 , and the vibrational frequencies; error bars: estimated total errors of Kc, the major part of these errors is probably temperature independent.
are comparable, the mean value 0 ∆Hr,298 (CF3O2NO2wCF3O2+NO2) )
102.7 ( 2.0 kJ mol-1 (final value) has been adopted for the further discussion. To obtain a consistent set of molecular parameters and thermochemical quantities, the three low-frequency vibrations were adjusted so that both the experimental value Kc ) 5.74 × 0 ) 108 molecules cm-3 at 280 K and the mean value ∆Hr,298 -1 102.7 kJ mol obtained from second- and third-law treatments of the experimental data are reproduced at the same time. The experimental data on reaction 3 are thus best described using the following set of parameters: frequencies (in cm-1) CF3O2NO2: 1760, 1303, 1292, 1244, 1186, 953, 877, 783, 702, 669, 600, 562, 485, 436, 371, 287, and 264 (from ref 1); 263 (from ref 20); 75, 34, 30 (from this work, best fit to the data); CF3O2: 1303, 1260, 1172, 1092, 870, 692, 597, 580, 448, 286, 190, 120 (from ref 31); NO2: 1666, 1358, 757 (from ref 21) moments of inertia (in amu Å2); CF3O2NO2: 129.7, 558.9, and 600.5 (from ref 1); CF3O2: 93.0, 161.5, 162.8 (from ref 30); NO2: 2.10, 38.6, 40.8 (from ref 21), Kc (280 K) ) (5.74 ( 0.92) × 108 molecules cm-3 with k3 from this work and k-3
where ZLJ ) Lennard-Jones collision number, Fvib,harm(E0) ) vibrational density of states at E ) E0 for harmonic oscillators, Qvib) vibrational partition function, FE ) correction factor allowing for the energy dependence of Fvib,harm at E > E0, Fanh ) correction factor for anharmonicity of vibrations, Frot ) correction factor for contributions of external rotations. Defining βc by k0,exp/k0(strong collisions), βc(M)N2) was found to be in the range 0.06-0.9 for a large number of reactions at room temperature.32,22 For the decomposition reactions of a total of 12 peroxynitrates where comprehensive data were available for both the dissociation and recombination reactions, βc(M)N2) was in the range 0.10-0.85.9 Since, for theoretical reasons, βc is expected to be very similar for different decomposition reactions in the same buffer gas, most of the scatter of these βc values is due to the large number of uncertainties inherent in both the theoretical and the experimental quantities entering eq III; usually, a value of βc ≈ 0.20.3 is assumed for M ) N2. Accordingly, the consistency of the present data analysis can be checked by calculating βc from eq III, with k0 being set equal to k3,exp(pw0) and k0(strong collisions) being calculated from molecular parameters. For this purpose, the above vibrational frequencies of CF3O2NO2 were used to calculate Fvib, Qvib, FE, Fanh, and Frot; ZLJ was estimated by a method described in the literature21,32,33 which is based on the boiling point of CF3O2NO2.1 With the present values E0 ) 100.5 kJ mol-1 and k3,0(M)N2) ) 5.7 × 10-20 cm3 molecule-1 s-1 at 280 K (see eq II and Figure 4), βc(M)N2) ) 0.17 is derived for reaction 3 which is well within the expected range of values. Data on the thermal decomposition of the analogous peroxynitrates CCl3O2NO2, CCl2FO2NO2, and CClF2O2NO2 from Ko¨ppenkastrop and Zabel7 and on the reverse reactions from Caralp et al.30 have been used by Caralp et al.30 and by Destriau and Troe34 to accomplish a self-consistent theoretical analysis of the following group of reactions:
CCl3O2NO2 (+M) S CCl3O2 + NO2 (+M)
(11,-11)
CCl2FO2NO2 (+M) S CCl2FO2 + NO2 (+M)
(12,-12)
CClF2O2NO2 (+M) S CClF2O2 + NO2 (+M)
(13,-13)
CF3O2NO2 (+M) S CF3O2 + NO2 (+M)
(3,-3)
Since there were no experimental data for k3 at that time, k3 was estimated from the results on reactions 11-13 by analogy. Essential features of these treatments were the following: Caralp
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TABLE 4: Comparison of Selected Kinetic and Thermochemical Data for Reactions 3 and 11-13 CCl3O2NO2 CCl2FO2NO2 CClF2O2NO2 CF3O2NO2 (ref 7) (ref 7) (ref 7) (this work) kdis(298 K, 1013 mbar N2) [s-1] Ea(≈1000 mbar N2) [kJ mol-1] 0 [kJ mol-1] ∆Hr,0 k/k∞ (298 K, 1013 mbar N2)
0.20
0.070
0.041
0.042
96.8 ( 1.4
100.3 ( 1.8
98.7 ( 1.8
97.7 ( 1.0
100.1 ( 3a 0.70
102.8 ( 3a 0.66
103.9 ( 3a 0.76
100.5 ( 2.0 0.72
a Based on third-law analysis of dissociation7 and recombination30 data.
et al.30 used RRKM calculations with MNDO calculated vibrational frequencies and moments of inertia for the peroxynitrates and the corresponding peroxy radicals, a single rate constant for k-12 from their own work at 273 K and 10 mbar, a k12 value from Ko¨ppenkastrop and Zabel35 at the same conditions, and the assumption that the reaction enthalpies of reactions 11, 13, and 3 are equal to that of reaction 12. On the basis of these presuppositions, Caralp et al.30 were able to satisfactorily fit their own data on the recombination rate constants k-11, k-12, and k-3 as a function of temperature and pressure by RRKM calculations, with only a few parameters to be slightly adjusted for the individual reactions. As a result of these calculations, k3 values were also predicted which agree with the present results within 60% over the complete range of experimental conditions. Destriau and Troe34 used high-pressure recombination rate constants k-3, k-11, k-12, and k-13 as calculated from SACM theory36 and preliminary data on k11, k12, and k13 from Ko¨ppenkastrop and Zabel7 to derive kinetic and thermochemical data for reactions 11-13, and 3. The results on reaction 3 of Caralp et al.30 and Destriau and Troe34 are compared with the experimental data of the present work in Table 3. Table 3 shows that the most accurate kinetic result of this work (the k3 value at 280 K, 60 mbar, i.e., in the center of the temperature and pressure ranges covered by the present experiments) is quite well reproduced by the previous estimates which were deduced from kinetic results on the decomposition of other peroxynitrates by analogy. The general agreement between the k3 data estimated earlier30,34 and the experimental data from this work shows that kinetic and thermochemical data can be estimated quite reliably using existing theories when sufficient kinetic and thermochemical data are known for similar molecules. For k3, the slightly different temperature dependence of Kc leads, by accident, to even closer agreement between the earlier estimates and the present results for the conditions close to the tropopause (220 K, 50 mbar) which are of particular interest for atmospheric applications. Since there is a large body of accurate data available on both k3 and k-3, several more general conclusions shall be drawn with respect to the evaluation of kinetic and thermochemical data. The present data on k3 may be compared with earlier results from this laboratory on k11, k12, and k13.7 The former experiments were carried out in the pressure range 10-800 mbar (M ) N2) and at temperatures between 260 and 293 K. Slight extrapolation to 1013 mbar and 298 K leads to the rate constants collected in Table 4. Since the experimental activation energies of reactions 11-13 and 3 agree within the combined error limits (see Table 4), it is difficult to deduce a trend in the OO-N bond energies of CX3O2NO2 when Cl atoms in the methyl group are successively replaced by F atoms. Bearing in mind that the absolute values of the rate constants are determined mainly
by the bond energy and to a lesser extent by the preexponential factor in a series of analogous compounds (see, e.g., ref 22), and that k11, k12, k13, and k3 are close to the high-pressure limit such that there is no notable difference in falloff (see Table 4), the absolute values of the rate constants are more appropriate to derive a trend in bond energies than is the activation energy. Inspection of Table 4 shows that there is a distinct drop in the rate constants when going from CCl3O2NO2 to CFCl2O2NO2 and a smaller decline from CFCl2O2NO2 to CF2ClO2NO2, whereas the rate constants for CF2ClO2NO2 and CF3O2NO2 are about the same. When the difference in rate constants is totally ascribed to different bond energies, E0 increases by 2.6, 1.3, and 0 kJ mol-1 when Cl is successively replaced by F in CX3O2NO2 (X ) Cl, F). The present data thus show that it might be inadequate to linearly extrapolate a trend of kinetic data within a series of compounds when very accurate values are required. When the second- and third-law analysis of kinetic data of a dissociation-recombination system leads to different values for the enthalpy of reaction, the third-law value is usually preferred (see, e.g., work on the reactions ClO + ClO w ClOOCl,38 CH3C(O)O2 + NO2 w CH3C(O)O2NO2,39 F + O2 w FO240). This is certainly adequate when the vibrational frequencies of the dissociating molecule and the products are well-known and/ or the experimental data on the temperature dependence of Kc are not very accurate; however, the accuracy of third-law treatments may be overestimated in other cases. The lowfrequency vibrations and, in particular, the very low torsional oscillations are often unknown, and estimates are subject to large uncertainties. In the present case of the unimolecular decomposition of CF3O2NO2, evaluation of the same kinetic data with two different sets of vibrational frequencies and moments of inertia, i.e., (i) the analysis by Caralp et al.30 with MNDO calculated parameters, and (ii) the analysis in the present work with experimentally determined parameters and estimated values based on thermochemical data of analogous reactions, leads to 0 which are different by 3.1 kJ mol-1. third law values of ∆Hr,298 30 It has been argued that MNDO calculations give frequencies which are high by 10-20% for stretching vibrations but well approximate deformation vibrations and torsions. However, Dewar et al.41 suggested that MNDO calculations underestimate the torsional frequencies. In their paper, a total of nine MNDO calculated torsional frequencies were compared with experimental values, and the average deviation between both sets of data was νobs/νcalc ≈ 1.33. On the other hand, there can also be a large scatter in experimental values. With respect to the present case, experimental values for the wavenumbers of the torsions of both CF3O2NO2 and CF3O2 are not known. Following a suggestion of Butler and Snelson,31 CF3OF can be used as a model compound for the CF3O2 radical. The MNDO calculated value for the torsion in CF3OF is 80 cm-1.42 By multiplying by 1.33, 106 cm-1 results as a “best” MNDO value. On the other hand, three experimental values of 56,43 120,44 and 127,45 cm-1 exist for the torsion in CF3OF. If these wavenumbers are adopted for CF3O2, their range translates to Kc values different by a factor of 2 which corresponds to an 0 of 1.6 kJ mol-1. In CF3O2NO2, there uncertainty of ∆Hr,298 are three of these torsions, and the uncertainty of the derived thermochemical values is correspondingly larger. Caralp et al.30 suggested that MNDO-calculated frequencies represent a consistent set of values, implying that systematic errors in these frequencies apply to both the dissociating molecule and the products and largely cancel in the calculation of Kc. In the case of CF3O2NO2 and CF3O2, however, two of the three torsions in CF3O2NO2 which strongly contribute to the vibrational partition function disappear when the decomposition products CF3O2 and
Thermal Decomposition of CF3O2NO2 NO2 are formed and directly translate into errors of the calculated equilibrium constants. Hence we conclude that whenever the experimental error of a reaction enthalpy derived from a van’t Hoff plot is a matter of only a few kJ mol-1, the “second-law” analysis could be preferable or at least equivalent to the “third-law” analysis. In these cases, it appears to be wise to consider the low-frequency vibrations merely as adjustable parameters which are useful to derive a self-consistent set of thermochemical data. Atmospheric Implications. In a recent publication by Ko et al.,12 the implications of the CF3 radical chemistry for stratospheric ozone were thoroughly discussed. Ko et al. introduced CF3O2NO2 as a possible temporary reservoir of CF3Ox which would be the major CF3X compound (CF3X ) CF3, CF3O, CF3O2, CF3OH, CF3OOH, CF3ONO2, CF3O2NO2, CF3OOCl) in the lower stratosphere if its photolysis rate constant were e10-4 s-1. From eq I and parameter set II it is estimated that the thermal lifetime of CF3O2NO2 is very short at room temperature and atmospheric pressure (ca. 1 min) but reaches a lifetime of ca. 1 year at the temperatures and pressures of the tropopause. For this reason, the lifetime τ of CF3O2NO2 in the tropopause and the lower stratosphere is probably limited by photolysis (τhν ≈ several days11 or even less46). Acknowledgment. Financial support of this work by the EC (European Commission) under Contract EV5V-0024 and by the Bundesminister fu¨r Bildung, Wissenschaft, Forschung und Technologie (BMBF) is gratefully acknowledged. References and Notes (1) Hohorst, F. A.; DesMarteau, D. D. Inorg. Chem. 1974, 13, 715. (2) Chen, J.; Young, V.; Zhu, T.; Niki, H. J. Phys. Chem. 1993, 97, 11696. (3) Wallington, T. J.; Hurley, M. D.; Schneider, W. F. Chem. Phys. Lett. 1993, 213, 442. (4) The value for k(CF3O2NO2 w CF3O2 + NO2) in Table 1 of ref 3 should read 0.017 s-1 (Wallington, T. J., private communication, 1993); the value g0.019 s-1 given here corresponds to the stated upper limit of ref 3 for the lifetime of CF3O2NO2 in the presence of NO. (5) Simonaitis, R.; Heicklen, J. Chem. Phys. Lett. 1979, 62, 473; 1979, 68, 245. (6) Simonaitis, R.; Glavas, S.; Heicklen, J. Geophys. Res. Lett. 1979, 6, 385. (7) Ko¨ppenkastrop, D.; Zabel, F. Int. J. Chem. Kinet. 1991, 23, 1. (8) Xiong, J. Q.; Carr, R. W. J. Phys. Chem. 1994, 98, 9811. (9) Zabel, F. Z. Phys. Chem. 1995, 188, 119. (10) 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. (11) Morel, O.; Simonaitis, R.; Heicklen, J. Chem. Phys. Lett. 1980, 73, 38. (12) Ko, M. K. W.; Sze, N.-D.; Rodrı´guez, J. M.; Weisenstein, D. K.; Heisey, C. W.; Wayne, R. P.; Biggs, P.; Canosa-Mas, C. E.; Sidebottom, H. W.; Treacy, J. Geophys. Res. Lett. 1994, 21, 101. (13) Barnes, I.; Becker, K. H.; Fink, E. H.; Reimer, A.; Zabel, F.; Niki, H. Int. J. Chem. Kinet. 1983, 15, 631.
J. Phys. Chem., Vol. 100, No. 16, 1996 6593 (14) Solomon, S.; Burkholder, J. B.; Ravishankara, A. R.; Garcia, R. R. J. Geophys. Res. 1994, 99, 20929. (15) Nielsen, O. J.; Ellermann, T.; Sehested, J.; Bartkiewicz, E.; Wallington, T. J.; Hurley, M. D. Int. J. Chem. Kinet. 1992, 24, 1009. (16) Chen, J.; Zhu, T.; Niki, H. J. Phys. Chem. 1992, 96, 6115. (17) DeMore, W. B.; Sander, S. P.; Golden, D. M.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R.; Kolb, C. E.; Molina, M. J. JPL Publ. 1994, 94-26. (18) Deuflhard, P.; Nowak, U. Ber. Bunsen-Ges. Phys. Chem. 1986, 90, 940. (19) Troe, J. J. Phys. Chem. 1979, 83, 114; Ber. Bunsen-Ges. Phys. Chem. 1983, 87, 161. (20) Nielsen, J. R.; Liang, C. Y.; Smith, D. C. J. Chem. Phys. 1953, 21, 1060. (21) Patrick, R.; Golden, D. M. Int. J. Chem. Kinet. 1983, 15, 1189. (22) Zabel, F. Thermischer Zerfall Von Peroxynitraten; habilitation thesis, Wuppertal, 1993. (23) Gilbert, R. G.; Luther, K.; Troe, J. Ber. Bunsen-Ges. Phys. Chem. 1983, 87, 169. (24) Woltz, P. J. H.; Jones, E. A.; Nielsen, A. H. J. Chem. Phys. 1952, 20, 378. (25) Ruff, O.; Menzel, W.; Neumann, W. Z. Anorg. Allg. Chem. 1932, 208, 293. (26) Jones, E. A.; Kirby-Smith, J. S.; Woltz, P. J. H.; Nielsen, A. H. J. Chem. Phys. 1951, 19, 242. (27) Klo¨ter, S.; Seppelt, K. J. Am. Chem. Soc. 1979, 101, 347. (28) Chen, J.; Zhu, T.; Niki, H.; Mains, G. J. Geophys. Res. Lett. 1992, 19, 2215. (29) Sehested, J.; Wallington, T. EnViron. Sci. Technol. 1993, 271, 46. (30) Caralp, F.; Lesclaux, R.; Rayez, M. T.; Rayez, J. C.; Forst, W. J. Chem. Soc., Faraday Trans. 2 1988, 84, 569. (31) Butler, R.; Snelson, A. J. Phys. Chem. 1979, 83, 3243. (32) Troe, J. J. Chem. Phys. 1977, 66, 4758. (33) Reid, R. C.; Sherwood, T. K. The Properties of Gases and Liquids; McGraw-Hill: New York, 1958. (34) Destriau, M.; Troe, J. Int. J. Chem. Kinet. 1990, 22, 915. (35) This rate constant for k3 was presented by Reimer and Zabel at the 9th International Symposium on Gas Kinetics, Bordeaux, 1986, but is virtually identical to that of ref 7. (36) Quack, M.; Troe, J. Ber. Bunsen-Ges. Phys. Chem. 1974, 78, 240. (37) It should be noted that a reevaluation of the method applied in ref 34 in the most recent IUPAC evaluation (J. Phys. Chem. Ref. Data 1992, 21, 1125) led to k3 (280 K, 60 mbar of N2) ) 1.6 × 10-3 s-1 which is a factor of 2 lower. (38) Nickolaisen, S. L.; Friedl, R. R.; Sander, S. P. J. Phys. Chem. 1994, 98, 155. (39) Bridier, I.; Caralp, F.; Loirat, H.; Lesclaux, R.; Veyret, B.; Becker, K. H.; Reimer, A.; Zabel, F. J. Phys. Chem. 1991, 95, 3594. (40) Campuzano-Jost, P.; Croce, A. E.; Hippler, H.; Siefke, M.; Troe, J. J. Chem. Phys. 1995, 102, 5317. (41) Dewar, M. J. S.; Ford, G. P.; McKee, M. L.; Rzepa, H. S.; Thiel, W.; Yamaguchi, Y. J. Mol. Struct. 1978, 43, 135. (42) We thank Dr. G. Hirsch (Wuppertal) for the calculation of the vibrational frequencies of CF3O2, using a MNDO program provided by Prof. W. Thiel, Zu¨rich. (43) Wilt, P. M.; Jones, E. A. J. Inorg. Nucl. Chem. 1968, 30, 2933. (44) Buckley, P.; Weber, J. P. Can. J. Chem. 1974, 52, 942. (45) Smardzewski, R. R.; Fox, W. B. J. Fluorine Chem. 1975, 6, 417. (46) Libuda, H. G.; Zabel, F. Data presented at the International Conference on Ozone in the Lower Stratosphere, May 15-20, 1995, Halkidiki, Greece.
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