Peter Biggs

Investigation into the pressure dependence between 1 and 10 Torr of the reactions of NO2 with CH3 and CH3O

The kinetics and pressure dependence of the reactions of NO2 with CH3 and CH3O have been investigated in the gas phase at 298 K, at pressures from 1 to 10 Torr. A low-pressure discharge-flow laser-induced fluorescence (LIF) technique was used. In a consecutive process, CH3 reacted with NO2 to form CH3O, CH3+ NO2→ CH3O + NO (1), which further reacted with NO2 to form products, CH3O + NO2→ products (2). Reaction (1) displayed no discernible pressure dependence over the pressure range 1–7 Torr, and k1 was calculated to be (2.3 ± 0.3)× 10–11 cm3 molecule–1 s–1. Reaction (2) displayed a strong pressure dependence and an RRKM analysis yielded the following limiting low- and high-pressure rate constants in He, k0= 5.9 × 10–29 cm6 molecule–2 s–1 and k∞= 2.1 × 10–11 cm3 molecule–1 s–1. It is unrealistic to quote errors for this type of analysis. Parametrisation in the standard Troe form with Fc= 0.6 yielded k0=(5.3 ± 0.2)× 10–29 cm6 molecule–2 s–1 and k∞=(1.4 ± 0.1)× 10–11 cm3 molecule–1 s–1. Atmospheric implications and possible reaction mechanisms are discussed.

Investigation into the kinetics and mechanism of the reaction of NO3 with CH3O2 at 298 K and 2.5 Torr: a potential source of OH in the night-time troposphere?

The kinetics of the reaction CH3O2+ NO3→ CH3O + NO2+ O2(1) have been studied at 298 K and at pressures between 2 and 3 Torr of helium using the discharge-flow technique combined with laser-induced fluorescence detection of the methoxyl radical and measurements of the NO3 radical using visible absorption. Numerical modelling of the concentration–time profile of CH3O with or without NO3 and NO as a titrant has allowed us to show that CH3O is a product of reaction (1) and to derive a rate constant k1=(1.0 ± 0.6)× 10–12 cm3 molecule–1 s–1, at 95% confidence limits. A comparison of the reactivities of NO3 and NO2 towards the species R, RO and RO2, where R = H or CH3, is given. The implication of reaction (1) in the possible production of OH in the atmosphere at night is discussed.

Investigation into the kinetics and mechanism of the reaction of NO3 with CH3 and CH3O at 298 K between 0.6 and 8.5 Torr: is there a chain decomposition mechanism in operation?

The reactions CH3+ NO3→ products (1), and CH3O + NO3→ products (2), have been studied using a flow system at T= 298 K and at pressures between 0.6 and 8.5 Torr. The laser-induced fluorescence (LIF) technique was used to detect CH3O and multi-pass optical absorption to detect NO3. The chemical systems were studied as a pair of consecutive reactions; however, a simple analytical treatment was not sufficient to describe them because CH3O2 was formed as one of the products in the major channel of reaction (2). This species also reacts with NO3 regenerating CH3O. Use of a numerical model to correct for this regeneration process allowed rate parameters of k1=(3.5 ± 1.0)× 10–11 cm3 molecule–1 s–1 and k2=(2.3 ± 0.7)× 10–12 cm3 molecule–1 s–1 to be determined at 2.4 Torr. There is no pressure dependence observed for reaction (1) between 1 and 2.4 Torr, but the possibility of a slight pressure dependence for reaction (2) exists. These pressure effects are examined using the semi-empirical quantum RRK method.

Rate constants for the reactions of C2H5, C2H5O and C2H5O2 radicals with NO3 at 298 K and 2.2 torr

A discharge-flow system equipped with a laser-induced fluorescence cell to detect the ethoxyl radical and an optical absorption cell to detect the nitrate radical has been used to measure the rate constants for the reactions C2H5+ NO3→ products (1), C2H5O + NO3→ products (2), C2H5O2+ NO3→ products (3), at T= 298 K and P= 2.2 Torr. The major products of these reactions are C2H5O and NO2 for reaction (1), C2H5O2 and NO2 for reaction (2) and C2H5O, O2 and NO2 for reaction (3). Reactions (2) and (3) are therefore highly coupled and, in order to determine the rate constants k2 and k3 for these reactions, two types of experiment were performed. In the first set of experiments, C2H5O was generated in situ via reaction (1) and allowed to react with NO3 and, in the second set of experiments, C2H5O2 was generated separately and allowed to react with NO3. A numerical model was then used to derive the following rate constants : k1=(4.0 ± 1.0)× 10–11 cm3 molecule–1 s–1, k2=(3.5 ± 1.0)× 10–12 cm3 molecule–1 s–1 and k3=(2.5 ± 1.5)× 10–12 cm3 molecule–1 s–1. The errors quoted are not statistically derived, but rather represent ranges over which the numerical simulations fit the experimental data adequately. The mechanisms for reactions (1) and (2) are discussed in terms of the QRRK treatment and the atmospheric implications of reaction (3) are considered.

Efficiency of formation of CH3O in the reaction of CH3O2 with ClO

A discharge-flow apparatus was used to determine the branching ratio α for the channel of the reaction of ClO with CH3O2 that leads to the formation of CH3O. The experiments were performed at 2 Torr pressure and at room temperature. ClO was the excess reactant and was measured mass spectrometrically; CH3O was detected by laser-induced fluorescence. The value of α was shown to depend on the fraction β of the reaction of CH3O2 with Cl atoms that leads to the formation of CH3O. For β = 0, 0.5 and 1, the derived values of α are 0.28±0.07, 0.32±0.08 and 0.42±0.09, respectively. Comparison with other published results suggests the possibility of three significant product channels for the reaction of CH3O2 with ClO.

Kinetics of the reactions ofCF3O2 with OH, HO2 andH

The kinetics of the reactions: CF 3 O 2 +OH→products (1), CF 3 O 2 +HO 2 →products (2) and CF 3 O 2 +H→CF 3 O+OH (3), have been investigated using a discharge fast-flow tube, at T=296±1 K and P=ca. 2 Torr. Two detection systems were used: laser-induced fluorescence of NO 2 produced from the titration of CF 3 O 2 with NO, and resonance fluorescence of OH. The rate constants (in cm 3 molecule -1 s -1 ) were found to be, k 1 =(4.0±0.6)×10 -11 , k 2 ⩽3×10 -12 and k 3 =(1.1±0.3)×10 - 10 . Experimental evidence suggests that the likely products of the reaction of CF 3 O 2 with OH are HO 2 and CF 3 O.

Kinetics of the self reaction of CF3O2 radicals and their reaction with O3 at 298 K

Two low-pressure discharge-flow systems, one equipped with laser-induced fluorescence (LIF) detection of NO 2 and the other mass spectrometric detection of NO 2 , have been employed to determine the rate constant for the self reaction CF 3 O 2 +CF 3 O 2 →products, at T=298 K and P=13 Torr. The phenomenological rate constant is k 1 =(2.0±0.1)×10 -12 cm 3 molecule -1 s -1 ; however, after consideration of secondary chemistry, numerical modelling suggests that the true rate constant is k 1 =(1.2±0.3)×10 -12 cm 3 molecule -1 s -1 . In addition, an upper limit has been derived for the reaction CF 3 O 2 +O 3 →products of k 3 <3×10 -14 cm 3 molecule -1 s -1 . The mechanism of the self reaction and the atmospheric implications of these results are discussed briefly.

Kinetics of the reaction of F atoms with CH3ONO and CH3O, and the reaction of CH3O with a number of hydrocarbons

A discharge-flow system equipped with a laser-induced fluorescence cell to detect the methoxyl radical and a quadrupole mass spectrometer to detect products has been used to determine the kinetics of the reactions: CH 3 ONO + F → products (1), CH 3 O + F → products (2), CH 3 O + CH 3 O → products (3), CH 3 O + n-C 4 H 10 → products (4), CH 3 O + i-C 4 H 10 → products (5), CH 3 O + C 2 H 4 → products (6), at T = 295 K and P = 2–3 Torr. The rate constants determined are: k 1 = (8.0 ± 0.5) × 10 −11 cm 3 molecule −1 s −1 , k 2 = (1.5 ± 0.5) × 10 −10 cm 3 molecule −1 s −1 , k 3 = (1.3 ± 0.4) × 10 −11 cm 3 molecule −1 s −1 , and upper limits for the rate constants k 4 ⩽ 1.0 × 10 −14 cm 3 molecule −1 s −1 , k 5 ⩽ 1.7 × 10 −14 cm 3 molecule −1 s −1 and k 6 ⩽ 1.1 × 10 −15 cm 3 molecule −1 s −1 . In addition, a summary of the sources of CH 3 O developed for use in low-pressure flow systems in this and other groups is presented.

Rate constants for the reaction between OH andCH3ONO2 andC2H5ONO2 over a range ofpressure and temperature

A discharge-flow apparatus equipped with resonance-fluorescence detection for OH radicals has been used to measure rate coefficients for the two reactions over the pressure range 1–20 Torr at T=298 K and ca. 1–3 Torr over the temperature range 300–400 K. The rate constants, k 1 and k 2 , were found to be invariant with pressure, but were found to increase with increasing temperature. In light of these findings, we conclude that at low pressure the reaction mechanism is consistent with an abstraction process. The rate constants derived are k 1 =(4.1±0.8)×10 -13 exp(-604±121/T) cm 3 molecule -1 s -1 and k 2 =(3.30±0.66)×10 -12 exp(-699±140/T) cm 3 molecule -1 s -1 . The existence of pressure-dependent (association) channels for reactions (1) and (2) is briefly discussed. Model calculations are presented that imply that photolysis, and not reaction with the OH radical, is the dominant loss process in the atmosphere for CH 3 ONO 2 and C 2 H 5 ONO 2 . Lifetimes of CH 3 ONO 2 and C 2 H 5 ONO 2 are derived for the lower atmosphere; they vary from a few years to several days depending upon season and location. Possible sources of these alkyl nitrates in the atmosphere are discussed.

Kinetics and mechanism of the reaction of CH3 and CH3O with ClO and OClO at 298 K

A discharge-flow system equipped with a laser-induced fluorescence cell to detect the methoxyl radical and a quadrupole mass spectrometer to detect both the chlorine monoxide and chlorine dioxide radicals has been used to measure the rate constants for the reactions CH3+ ClO → products (1), CH3O + ClO → products (2), CH3+ OClO → products (3), CH3+ OClO → products (4), at T= 298 K and P= 1–3 Torr. The observed products of these reactions are CH3O for reaction (1), HOCl for reaction (2), and CH3O and ClO for reaction (3). For reaction (4), CH3OCl is a possible product. The rate constants derived for reactions (1)–(4) are: k1=(1.3 ± 0.4)× 10–10 cm3 molecule–1 s–1; k2=(1.3 ± 0.3)× 10–11 cm3 molecule–1 s–1; k3=(1.6 ± 0.3)× 10–11 cm3 molecule–1 s–1; k4=(1.5 ± 0.5)× 10–12 cm3 molecule–1 s–1. The likely mechanisms for reactions (1)–(4) are briefly discussed.