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Rate constants for the reaction of OH radicals with 1-chloroalkanes at 295 K
Volume 189, number 2
CHEMICAL PHYSICS LETTERS
31 January 1992
Rate constants for the reaction of OH radicals with l-chloroalkanes at 295 K Frank Marker? and Ole John Nielsen Section for Chemical Reactivity, Environmental Science and Technology Department, Rise National Laboratory, DK-4000 Roskilde, Denmark Received 4 October 199 1
The rate constants for the reaction of OH radicals with a series of 1chloroalkanes were measured at 295 K and at a total pressure of 1 atm. The rate constants were obtained by using the absolute technique of pulse radiolysis combined with kinetic UV-spectroscopy. The results are discussed in terms of reactivity trends.
1. Introduction
In recent years, much work has been done to build up the database for the reaction of OH radicals with organic compounds [ 11. In the years up to 1986, Atkinson [ 2 ] developed a prediction model for reactions of OH with a wide range of organic compounds. It is now possible to predict rate constants for those reactions, also in a wide temperature range, with good accuracy [ 3,4]. The basic assumption using this method is that only neighboring groups influence the H abstraction. However, long-range effects of different substituents have been shown to exist, i.e. for, ketones, ethers, alcohols, nitro-compounds and others [ 5-101. In this work, we have used the technique of pulse radiolysis combined with kinetic UV-detection to measure the rate constants for the reaction of OH radicals with Cl-ethane, l-Cl-propane, l-cl-butane, 1-Cl-pentane, and I-Cl-hexane. The results are compared with the rate constants for the reaction of these I-Cl-alkanes with chlorine atoms [ 111 and with the predicted rate constants derived by Atkinson’s structure-activity relationship (SAR) method [ 21.
2. Experimental All experiments have been carried out by using the
technique of pulse radiolysis combined with kinetic UV-spectroscopy detection. This technique has been described in detail elsewhere [ 12,131. Briefly, a 1 P stainless-steel gas cell was directly mounted on the radiation source, and a Febetron 705B lield-emission accelerator producing a 30 ns pulse of 2 MeV electrons, was filled with a certain partial pressure of the compound to be investigated, adding 15 mbar of water vapor and argon to a total pressure of 1 atm. The partial pressures were measured by a MKS Baratron 170 absolute membrane manometer with a resolution of 1O-’ bar. Pulse radiolysis of such mixtures produce OH radicals: 2 MeV e-+Ar+Ar*, AS+H20+0H+H+Ar. The decay of OH radicals was followed by monitoring the transient absorbance of rotational lines at 309 nm. The gas cell was equipped with a set of White mirrors providing a pathlength of 120 cm. Furthermore, the analyzing light source, a 150 W high-pressure xenon arc lamp, was pulsed. The other components of the signal detection system were a Hilger and Watts grating spectrograph (resolution of 0.8 nm for 1 mm slitwidth), a Hamamatsu photomultiplier and a Biomation 8100 waveform digitizer. For storage of raw data and data handling, a PDPl 1 minicomputer was used.
0009-2614/92/S 05.00 0 1992 Elsevier Science Publishers B.V. All rights reserved.
171
Volume 189, number
2
CHEMICAL
PHYSICS
The OH rotational lines were narrow compared to the experimental resolution resulting in a nonlinearity in Beer’s law. Therefore, a modified version of Beer’s law A = (dc)” was used. The value of n was determined from the function log A= n log( E/C), by varying the OH concentration through dose variation and carrying out a linear leastsquares-fit analysis. For a spectral band pass of 0.08 nm, n was found to be 0.70?0.04. The observed transient absorbance is a direct measure of the OH concentration at any time during the decay. The experimental error in the determination of n introduces an error of about 3% in the determination of the pseudo-first-order rate constant [ 121. In all experiments, the OH concentration was estimated to be in the order of lOI molecules cmp3 based on previous experiments from this laboratory. In pure water vapor, the decay of the OH radical is governed by 0H+OH+M-+H202+M,
(1)
H+OH+M+HIO+M,
(2)
where M is Ar and H20. When a reactive compound system, the following reaction
- F
=2k,
[OH]‘[M]
In the pure H,O/Ar system, the OH radical decay was nonexponential. However, after adding l-chloroalkanes, the decay became clearly exponential over at least three half-lives. Plot of k’ versus the l-chloroalkane concentration are shown in fig. 1. In repeating the experiments, an amount of scatter in the bimolecular rate constants was observed. Therefore, the average of these bH values was used and listed in table 1. The given uncertainty represents the standard deviation of the mean values. To minimize errors in the rate constants induced
+k, [OH] [H] [M] (4)
The experimental conditions in this work were pseudo-first-order, where k3[RH] is the all dominant part of eq. (4). Therefore, eq. (4) reduces to =k’
[OH],
with k’=k,
[RH],
7
20.0 111
k f +o -
10.0
(5)
which is easy to integrate. The bimolecular rate constants k3 were obtained by measuring the OH decays at different concentrations of RH, fitting the data to a first-order decay and plotting the pseudo-first-order rate constants k’ against the concentration of RH. To predict the rate constants following the SAR method [ 21, the following equation was used: 172
(6)
9
3. Results
by the following
+k, [RH] [OH].
- y
1992
k,,=O.838, both in units of with ~~i=O.144, lOI* cm3 molecules- ’ s- ’ , and F(-Cl) =0.38, F(-CH,Cl) =0.57, F(-CH3) = 1.00, F(-CH2-) = 1.29. The substances, RH, used in this work were chloroethane (Gerling Holz & co, purity 99.7%), l-chloropropane (Aldrich, 99%), 1-chlorobutane (Aldrich, HPLC 99.9%), 1-chloropentane (Aldrich, 99% + ) and 1-chlorohexane (Aldrich, 99%). The argon diluent (AGA) had a purity of 99.998%. The water was triple distilled. The gaseous chloroethane was used directly. The other compounds were checked by GC-MS and used directly after degassing them by freeze-pump-thaw cycles.
(3)
The decay of OH can be described differential equation:
31 January
km,= 1 kpriF(X)+ C kecJ’(X)F(Y)
RH is added to the also takes place:
RH+OH+R+H*O.
LETTERS
0.0
0.0
1 .o
2.0 Concentration.
\
Fig. 1. The plot shows k’ in mbar. The bimolecular of these plots. The points I-EtCI; (V) I-PrC1; (v)
3.0
4.0
50
mbor
in s-’ versus compound concentration rate constant is derived from the slope represent all our measurements. (0 ) I-BuCl; (Cl) I-PeCl;, (m) I-HeCl.
CHEMICAL PHYSICS LETTERS
Volume 189, number 2
31 January 1992
Table 1 Rate constants br., for the reaction of OH radicals with l-chloroalkanes at 295 + 2 K in units cm3 molecules-’ s-’ Compound
lo’* kOH
1-chloroethane
0.39 f 0.07
No. of exp.
I141 1151 It61
0.39kO.05 0.41 f 0.07 l-chloropropane I-chlorobutane I-chloropentane I-chlorohexane
Ref.
0.43+0.05
2
0.82?0.14 1.5kO.4 3.1f0.3 3.8kO.2
2 4 2 2
this this this this this
work work work work work
by secondary products [ 171, the sample was renewed after every pulse. Reaction of OH radicals with impurities in either the 1-chloroalkanes or H,O/Ar buffer gas cannot be discounted for by the above analytical procedure.
4. Discussion The rate constants for reactions of OH radicals with I-Cl-alkanes obtained at 1 atm total pressure and 295 ? 2 K in this work are compared with available literature data in table 1. However, only the OH + chloroethane reaction has been investigated previously. In this case, the agreement between the result of this investigation and the other three studies [ 14- 16 ] is excellent. The rate constant for the raction of OH with l-Clalkanes is shown as a function of the number of CHI groups in fig. 2a. Also in fig. 2a, the rate constants are compared to the rate constants for the reaction of OH with alkanes minus a number representing the a-CH3 reactivity [ 2 1. The a-CH, reactivity is taken as f&+ethane [ 21. It is seen that the Cl substitution has a deactivating effect that extends at least to the b-carbon atom. There could be several reasons for this deactivating effect. The C-H bond dissociation energies in all the 1-Cl-alkanes have not been determined experimentally. However, for CH,CH&l it was found that the bond energy of the primary C-H bond was the same as in ethane within the estimated uncertainty [ 181. This indicates that the decreased reactivity is not due to an increase in C-H bond strength. The reason for the reduced reactivity must be sought elsewhere. It has been suggested in other
0.0
,
0
1
2
Number
3
I
4
5
6
of CH, groups
Fig. 2. (a) Comparison between the rate constants for OH with i -Cl-alkanes ( 0 ) and alkanes ( 0 ) versus number of CH, groups. (b) Comparison between the rate constants for Cl with I-Cl-alkanes ( 0 ) and alkanes ( 0 ) versus number of CH2 groups.
investigations of substituted n-alkanes [ lo,19 ] that an inductive effect in the transition state involving polar repulsion between the electrophilic OH radical and the abstracted H atom may be the reason. In fig. 2b, the rate constant for the raction of Cl atoms with 1-Cl-alkanes [ 111 is compared to the rate constants for Cl-atom reaction with alkanes [20] minus the a-CH, reactivity, calculated as above. It can be seen that in this case, the deactivating effect is almost constant over the series of 1-Cl-alkanes. Again, the reduction cannot be due to increased C-H bond-dissociation energies. Besides the polar effect, there could also be a steric hindrance for the smaller chloroalkanes. OH radicals and Cl atoms are both electrophilic species and correlation between the respective rate constants for reactions having the same mechanism is expected. Fig. 3 shows a free-energy plot for the OH and Cl reactions with 1-Cl-alkanes. As expected, 173
Volume 189, number
2
CHEMICAL
PHYSICS
LETTERS
31 January
1992
Acknowledgement Thanks are due to Jette Munk for technical assistance and Helge Egsgard (both from Riss National Laboratory) for checking the purity of samples, and to Dr. T.J. Wallington for communicating his data prior to publication. The authors are grateful to the Commission of the European Communities for financial support. lo-l3 10-l'
I
,,....,I
10s k,,
I
I
lo-"
10-12 (cm3molecu1es
Fig. 3. Free-energy plot for the reaction OH and Cl radicals.
7
.,<,.,.
10-l'
-I
s
-I
References
)
of l-chloroalkanes
I
I
with
I
v
5.0 m
7
i
4.0
7 nii
3.0
i
2.0
B A
1.0
N 0 -
0.0
v
@ 8
Q 1
ETCL
PRCL
BUCL
/
PECL
I
HECL
Fig. 4. Comparison between calculated and measured rate constants. (V ) Calculated by SAR, (0 ) experimental values.
this plot is essentially linear indicating that both reactions proceed via H-atom abstraction. In fig. 4, the experimental rate constants are compared with those calculated by Atkinson’s SAR method; we found that ours agree well for Cl-ethane, but are lower for the higher homologous l-Cl-alkanes. This leads to the conclusion that the factors F(-Cl) and F( -CH,Cl) in Atkinson’s model might be recalculated and probably a new factor F(-CH2CH2Cl) should be taken into account, describing the long-range effect over more than two carbon atoms.
174
[ 1 ] R. Atkinson, J. Phys. Chem. Ref. Data, Monograph No. 1 (1989). [2] R. Atkinson, Chem. Rev. 85 (1985) 185. [ 31 R. Atkinson, Intern. J. Chem. Kinetics 18 (1986) 555. [ 41 R. Atkinson, Intern. J. Chem. Kinetics 19 ( 1987) 799. [ 51 T.J. Wallington, P. Dagout, R. Liu and M.J. Kurylo, Intern. J. Chem. Kinetics 20 (1988) 541. [ 61 L. Nelson, 0. Rattigan, R. Neavyn, H. Sidebottom, J. Treaty and O.J. Nielsen, Intern. J. Chem. Kinetics 22 ( 1990) 1111. [ 7) J.T. Wallington and M.J. Kurylo, J. Phys. Chem. 91 ( 1987) 5050. [8] R. Atkinson, S.M. Aschmann, W.P.L. Carter and J.N. Pitts Jr., Intern. J. Chem. Kinetics 14 (1982) 839. [9] T.J. Wallington and M.J. Kurylo, Intern. J. Chem. Kinetics 19 (1987) 1015. [lo] O.J. Nielsen, H.W. Sidebottom, D.J. O’Farrell, M. Donlon and J. Treaty, Chem. Phys. Letters 156 ( 1989) 3 12. 1” T.J. Wallington, L.M. Skewes and W.O. Siegl, J. Phys. Chem. 93 (1989) 3649. 1’2 O.J. Nielsen, 0. Jorgensen, M. Donlon, H.W. Sidebottom, D.J. O’Farrell and J. Treaty, Chem. Phys. Letters 168 (1990) 319. [13 K.B. Hansen, R. Wilbrandt and P. Pagsberg, Rev. Sci. Instr. 50 (1979) 1532. ]14 C.J. Howard and K.M. Evenson, J. Chem. Phys. 64 ( 1976) 4303. D.L. Singleton and R.S. Irwin, J. Phys. 115 G. Paraskevopoulos, Chem. 85 (1981) 561. 1’6 J.H. Kasner, P.H. Taylor and B. Dellinger, J. Phys. Chem. 94 (1990) 3250. Acta, Band 2 Heft 2 ( 1963) [17 R.N. Schindler, Radiochimica 62. [ 181 K. Miyokawa and E. Tsuikow-Roux, J. Chem. Phys. 94 (1990) 715. [ 191 O.J. Nielsen, H.W. Sidebottom, M. Donlon and J. Treaty, Chem. Phys. Letters 178 (1991) 163. [ 201 R. Atkinson and S. Aschmann, Intern. J. Chem. Kinetics 17 (1985) 33.
CHEMICAL PHYSICS LETTERS
31 January 1992
Rate constants for the reaction of OH radicals with l-chloroalkanes at 295 K Frank Marker? and Ole John Nielsen Section for Chemical Reactivity, Environmental Science and Technology Department, Rise National Laboratory, DK-4000 Roskilde, Denmark Received 4 October 199 1
The rate constants for the reaction of OH radicals with a series of 1chloroalkanes were measured at 295 K and at a total pressure of 1 atm. The rate constants were obtained by using the absolute technique of pulse radiolysis combined with kinetic UV-spectroscopy. The results are discussed in terms of reactivity trends.
1. Introduction
In recent years, much work has been done to build up the database for the reaction of OH radicals with organic compounds [ 11. In the years up to 1986, Atkinson [ 2 ] developed a prediction model for reactions of OH with a wide range of organic compounds. It is now possible to predict rate constants for those reactions, also in a wide temperature range, with good accuracy [ 3,4]. The basic assumption using this method is that only neighboring groups influence the H abstraction. However, long-range effects of different substituents have been shown to exist, i.e. for, ketones, ethers, alcohols, nitro-compounds and others [ 5-101. In this work, we have used the technique of pulse radiolysis combined with kinetic UV-detection to measure the rate constants for the reaction of OH radicals with Cl-ethane, l-Cl-propane, l-cl-butane, 1-Cl-pentane, and I-Cl-hexane. The results are compared with the rate constants for the reaction of these I-Cl-alkanes with chlorine atoms [ 111 and with the predicted rate constants derived by Atkinson’s structure-activity relationship (SAR) method [ 21.
2. Experimental All experiments have been carried out by using the
technique of pulse radiolysis combined with kinetic UV-spectroscopy detection. This technique has been described in detail elsewhere [ 12,131. Briefly, a 1 P stainless-steel gas cell was directly mounted on the radiation source, and a Febetron 705B lield-emission accelerator producing a 30 ns pulse of 2 MeV electrons, was filled with a certain partial pressure of the compound to be investigated, adding 15 mbar of water vapor and argon to a total pressure of 1 atm. The partial pressures were measured by a MKS Baratron 170 absolute membrane manometer with a resolution of 1O-’ bar. Pulse radiolysis of such mixtures produce OH radicals: 2 MeV e-+Ar+Ar*, AS+H20+0H+H+Ar. The decay of OH radicals was followed by monitoring the transient absorbance of rotational lines at 309 nm. The gas cell was equipped with a set of White mirrors providing a pathlength of 120 cm. Furthermore, the analyzing light source, a 150 W high-pressure xenon arc lamp, was pulsed. The other components of the signal detection system were a Hilger and Watts grating spectrograph (resolution of 0.8 nm for 1 mm slitwidth), a Hamamatsu photomultiplier and a Biomation 8100 waveform digitizer. For storage of raw data and data handling, a PDPl 1 minicomputer was used.
0009-2614/92/S 05.00 0 1992 Elsevier Science Publishers B.V. All rights reserved.
171
Volume 189, number
2
CHEMICAL
PHYSICS
The OH rotational lines were narrow compared to the experimental resolution resulting in a nonlinearity in Beer’s law. Therefore, a modified version of Beer’s law A = (dc)” was used. The value of n was determined from the function log A= n log( E/C), by varying the OH concentration through dose variation and carrying out a linear leastsquares-fit analysis. For a spectral band pass of 0.08 nm, n was found to be 0.70?0.04. The observed transient absorbance is a direct measure of the OH concentration at any time during the decay. The experimental error in the determination of n introduces an error of about 3% in the determination of the pseudo-first-order rate constant [ 121. In all experiments, the OH concentration was estimated to be in the order of lOI molecules cmp3 based on previous experiments from this laboratory. In pure water vapor, the decay of the OH radical is governed by 0H+OH+M-+H202+M,
(1)
H+OH+M+HIO+M,
(2)
where M is Ar and H20. When a reactive compound system, the following reaction
- F
=2k,
[OH]‘[M]
In the pure H,O/Ar system, the OH radical decay was nonexponential. However, after adding l-chloroalkanes, the decay became clearly exponential over at least three half-lives. Plot of k’ versus the l-chloroalkane concentration are shown in fig. 1. In repeating the experiments, an amount of scatter in the bimolecular rate constants was observed. Therefore, the average of these bH values was used and listed in table 1. The given uncertainty represents the standard deviation of the mean values. To minimize errors in the rate constants induced
+k, [OH] [H] [M] (4)
The experimental conditions in this work were pseudo-first-order, where k3[RH] is the all dominant part of eq. (4). Therefore, eq. (4) reduces to =k’
[OH],
with k’=k,
[RH],
7
20.0 111
k f +o -
10.0
(5)
which is easy to integrate. The bimolecular rate constants k3 were obtained by measuring the OH decays at different concentrations of RH, fitting the data to a first-order decay and plotting the pseudo-first-order rate constants k’ against the concentration of RH. To predict the rate constants following the SAR method [ 21, the following equation was used: 172
(6)
9
3. Results
by the following
+k, [RH] [OH].
- y
1992
k,,=O.838, both in units of with ~~i=O.144, lOI* cm3 molecules- ’ s- ’ , and F(-Cl) =0.38, F(-CH,Cl) =0.57, F(-CH3) = 1.00, F(-CH2-) = 1.29. The substances, RH, used in this work were chloroethane (Gerling Holz & co, purity 99.7%), l-chloropropane (Aldrich, 99%), 1-chlorobutane (Aldrich, HPLC 99.9%), 1-chloropentane (Aldrich, 99% + ) and 1-chlorohexane (Aldrich, 99%). The argon diluent (AGA) had a purity of 99.998%. The water was triple distilled. The gaseous chloroethane was used directly. The other compounds were checked by GC-MS and used directly after degassing them by freeze-pump-thaw cycles.
(3)
The decay of OH can be described differential equation:
31 January
km,= 1 kpriF(X)+ C kecJ’(X)F(Y)
RH is added to the also takes place:
RH+OH+R+H*O.
LETTERS
0.0
0.0
1 .o
2.0 Concentration.
\
Fig. 1. The plot shows k’ in mbar. The bimolecular of these plots. The points I-EtCI; (V) I-PrC1; (v)
3.0
4.0
50
mbor
in s-’ versus compound concentration rate constant is derived from the slope represent all our measurements. (0 ) I-BuCl; (Cl) I-PeCl;, (m) I-HeCl.
CHEMICAL PHYSICS LETTERS
Volume 189, number 2
31 January 1992
Table 1 Rate constants br., for the reaction of OH radicals with l-chloroalkanes at 295 + 2 K in units cm3 molecules-’ s-’ Compound
lo’* kOH
1-chloroethane
0.39 f 0.07
No. of exp.
I141 1151 It61
0.39kO.05 0.41 f 0.07 l-chloropropane I-chlorobutane I-chloropentane I-chlorohexane
Ref.
0.43+0.05
2
0.82?0.14 1.5kO.4 3.1f0.3 3.8kO.2
2 4 2 2
this this this this this
work work work work work
by secondary products [ 171, the sample was renewed after every pulse. Reaction of OH radicals with impurities in either the 1-chloroalkanes or H,O/Ar buffer gas cannot be discounted for by the above analytical procedure.
4. Discussion The rate constants for reactions of OH radicals with I-Cl-alkanes obtained at 1 atm total pressure and 295 ? 2 K in this work are compared with available literature data in table 1. However, only the OH + chloroethane reaction has been investigated previously. In this case, the agreement between the result of this investigation and the other three studies [ 14- 16 ] is excellent. The rate constant for the raction of OH with l-Clalkanes is shown as a function of the number of CHI groups in fig. 2a. Also in fig. 2a, the rate constants are compared to the rate constants for the reaction of OH with alkanes minus a number representing the a-CH3 reactivity [ 2 1. The a-CH, reactivity is taken as f&+ethane [ 21. It is seen that the Cl substitution has a deactivating effect that extends at least to the b-carbon atom. There could be several reasons for this deactivating effect. The C-H bond dissociation energies in all the 1-Cl-alkanes have not been determined experimentally. However, for CH,CH&l it was found that the bond energy of the primary C-H bond was the same as in ethane within the estimated uncertainty [ 181. This indicates that the decreased reactivity is not due to an increase in C-H bond strength. The reason for the reduced reactivity must be sought elsewhere. It has been suggested in other
0.0
,
0
1
2
Number
3
I
4
5
6
of CH, groups
Fig. 2. (a) Comparison between the rate constants for OH with i -Cl-alkanes ( 0 ) and alkanes ( 0 ) versus number of CH, groups. (b) Comparison between the rate constants for Cl with I-Cl-alkanes ( 0 ) and alkanes ( 0 ) versus number of CH2 groups.
investigations of substituted n-alkanes [ lo,19 ] that an inductive effect in the transition state involving polar repulsion between the electrophilic OH radical and the abstracted H atom may be the reason. In fig. 2b, the rate constant for the raction of Cl atoms with 1-Cl-alkanes [ 111 is compared to the rate constants for Cl-atom reaction with alkanes [20] minus the a-CH, reactivity, calculated as above. It can be seen that in this case, the deactivating effect is almost constant over the series of 1-Cl-alkanes. Again, the reduction cannot be due to increased C-H bond-dissociation energies. Besides the polar effect, there could also be a steric hindrance for the smaller chloroalkanes. OH radicals and Cl atoms are both electrophilic species and correlation between the respective rate constants for reactions having the same mechanism is expected. Fig. 3 shows a free-energy plot for the OH and Cl reactions with 1-Cl-alkanes. As expected, 173
Volume 189, number
2
CHEMICAL
PHYSICS
LETTERS
31 January
1992
Acknowledgement Thanks are due to Jette Munk for technical assistance and Helge Egsgard (both from Riss National Laboratory) for checking the purity of samples, and to Dr. T.J. Wallington for communicating his data prior to publication. The authors are grateful to the Commission of the European Communities for financial support. lo-l3 10-l'
I
,,....,I
10s k,,
I
I
lo-"
10-12 (cm3molecu1es
Fig. 3. Free-energy plot for the reaction OH and Cl radicals.
7
.,<,.,.
10-l'
-I
s
-I
References
)
of l-chloroalkanes
I
I
with
I
v
5.0 m
7
i
4.0
7 nii
3.0
i
2.0
B A
1.0
N 0 -
0.0
v
@ 8
Q 1
ETCL
PRCL
BUCL
/
PECL
I
HECL
Fig. 4. Comparison between calculated and measured rate constants. (V ) Calculated by SAR, (0 ) experimental values.
this plot is essentially linear indicating that both reactions proceed via H-atom abstraction. In fig. 4, the experimental rate constants are compared with those calculated by Atkinson’s SAR method; we found that ours agree well for Cl-ethane, but are lower for the higher homologous l-Cl-alkanes. This leads to the conclusion that the factors F(-Cl) and F( -CH,Cl) in Atkinson’s model might be recalculated and probably a new factor F(-CH2CH2Cl) should be taken into account, describing the long-range effect over more than two carbon atoms.
174
[ 1 ] R. Atkinson, J. Phys. Chem. Ref. Data, Monograph No. 1 (1989). [2] R. Atkinson, Chem. Rev. 85 (1985) 185. [ 31 R. Atkinson, Intern. J. Chem. Kinetics 18 (1986) 555. [ 41 R. Atkinson, Intern. J. Chem. Kinetics 19 ( 1987) 799. [ 51 T.J. Wallington, P. Dagout, R. Liu and M.J. Kurylo, Intern. J. Chem. Kinetics 20 (1988) 541. [ 61 L. Nelson, 0. Rattigan, R. Neavyn, H. Sidebottom, J. Treaty and O.J. Nielsen, Intern. J. Chem. Kinetics 22 ( 1990) 1111. [ 7) J.T. Wallington and M.J. Kurylo, J. Phys. Chem. 91 ( 1987) 5050. [8] R. Atkinson, S.M. Aschmann, W.P.L. Carter and J.N. Pitts Jr., Intern. J. Chem. Kinetics 14 (1982) 839. [9] T.J. Wallington and M.J. Kurylo, Intern. J. Chem. Kinetics 19 (1987) 1015. [lo] O.J. Nielsen, H.W. Sidebottom, D.J. O’Farrell, M. Donlon and J. Treaty, Chem. Phys. Letters 156 ( 1989) 3 12. 1” T.J. Wallington, L.M. Skewes and W.O. Siegl, J. Phys. Chem. 93 (1989) 3649. 1’2 O.J. Nielsen, 0. Jorgensen, M. Donlon, H.W. Sidebottom, D.J. O’Farrell and J. Treaty, Chem. Phys. Letters 168 (1990) 319. [13 K.B. Hansen, R. Wilbrandt and P. Pagsberg, Rev. Sci. Instr. 50 (1979) 1532. ]14 C.J. Howard and K.M. Evenson, J. Chem. Phys. 64 ( 1976) 4303. D.L. Singleton and R.S. Irwin, J. Phys. 115 G. Paraskevopoulos, Chem. 85 (1981) 561. 1’6 J.H. Kasner, P.H. Taylor and B. Dellinger, J. Phys. Chem. 94 (1990) 3250. Acta, Band 2 Heft 2 ( 1963) [17 R.N. Schindler, Radiochimica 62. [ 181 K. Miyokawa and E. Tsuikow-Roux, J. Chem. Phys. 94 (1990) 715. [ 191 O.J. Nielsen, H.W. Sidebottom, M. Donlon and J. Treaty, Chem. Phys. Letters 178 (1991) 163. [ 201 R. Atkinson and S. Aschmann, Intern. J. Chem. Kinetics 17 (1985) 33.