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J . Phys. Chem. 1989, 93, 5458-5461
5458
+
+ + NO, Rate Constant by Infrared Diode
Determination of the NO3 NO, NO 0, Laser and Fourier Transform Spectroscopy
J. Hjorth,* F. Cappellani, C. J. Nielsen,+ and G. Restelli Commission of the European Communities, Joint Research Centre-Ispra Establishment, I - 21020 Ispra (VA), Italy (Received: August 24, 1988; In Final Form: February 8, 1989)
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The rate constant, k l , at 296 K for the reaction NO, + NO2 NO + NO2 + 0, has been determined relative to the rate constant, k2, for the reaction NO3 + N O 2N02 by direct measurement of [NO] and [NO,] in a steady-state situation where the effect of the thermal decomposition of NO3 was negligible. NO3was generated in the presence of a large excess of NO2. The [NO,] was determined by FTIR while the [NO] was measured with a tunable diode laser spectrometer operated in second-harmonic detection mode. From a series of 18 experiments performed at 50- and 15-Torr total pressures, the rate constant k l was calculated at 296 K as k , = ((1.5 f 0.28) X 10-,9/k2 cm3 molecule-l s-I. Assuming the rate constant k2 = (3.0 i 0.9) X lo-'' cm3 molecule-' s-l, a kl value at 296 K equal to (5.1 f 1.8) X cm3 molecule-' s-' is calculated.
Introduction One of the reactions that destroys NO3 in the troposphere at night is NO, NO2 N O NO2 0 2 (1)
+
+
+
+
Although this reaction is not the most important sink for NO,, it is far from negligible. An extensive review of the N 2 0 5 / N 0 3 kinetics has been published in ref 1. The first attempt to evaluate the rate constant, k l , was made by Schott and Davidson,2 who performed shock tube experiments in the temperature range 750-934 K. In these experiments, N2OSwas completely dissociated into equal concentrations of NO2 and NO, and the rate of decrease of [NO,] and increase of [NO,] was measured. The decay of NO3 under these conditions was assumed to be due either to reaction 1 followed by the very rapid reaction NO3 + N O 2N02 k 2 = 3.0 X lo-'' cm3 molecule-] s-' (2)
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or to the self-reaction 2NO,
-
2N02
+ O2
(3)
This leads to the following expression for the NO3 decay: -d[NO,]/dt = 2 k l [ N 0 2 ] [ N 0 3 ]+ 2k3[N0,I2
(a)
As previously mentioned, the concentrations of NO3 and NO2 were initially equal ( t = 0) while the final concentration of NO3 approached zero with a corresponding increase in the NO2 concentration ( t = m). Applying eq a to the results extrapolated to t = 0 yielded kl
+ k3 = 3.63 X
exp(-3231/T)
At long reaction times ( t = a) when [NO,] approaches zero, the contribution of reaction 3 becomes negligible and k , is estimated as k, = 6.4 X
exp(-4562/T)
Further, the concentration ratio [NO,]/ [NO,] was varied between 0.1 and 1 by adding NO2 to N2O5 before it was decomposed by the shock wave. Extrapolating these results to t = 0 and the contribution of reaction 3 to zero resulted in k l = 4.28
X
exp(-2222/T)
Very different values are obtained for kl from the two expressions above when the results are extrapolated to tropospheric temperatures, e.g., at 296 K: k l = 1.29 X cm3 molecule-' s-' ( t = 0) k , = 2.35 X cm3 molecule-' s-' ( t = m ) On leave from the Department of Chemistry, University of Oslo, P.O. Box 1033 Blindern, N-0315 Oslo 3, Norway. 0022-3654/89/2093-5458$01.50/0
Graham and Johnston4 derived the equilibrium constant, K , for the reaction N205 + M
;=t
NO3 + NO2 + M
(4)
from an analysis of the N 2 0 5 / 0 3system kinetics. The temperature dependence of K was found as
K = (8.4 i 1.8)
X
exp[(-11180 f lOO)/q
Prior to these studies, Johnston and TaoS had determined the value for the product K k , in the temperature range 338-396 K by measuring the conversion of N2O5 to NO2. By combining the results from these studies, k l was determined as4
k l = (2.5 f 0.5) X 1 0 - l ~exp[(-123O f l O O ) / q cm3 molecule-Is-, At 296 K, k , is calculated as 3.92 X 10-l6 cm3 molecule-' s-l. As recently shown by Johnston et al.' a strong discrepancy appears when this expression is used to calculate k , in the temperature range investigated by Schott and Davidson.2 This discrepancy has been, however, discussed and accounted for' by the contribution of the thermal unimolecular decomposition of NO3 in the system NO,
.-+N O
+ 0,
(5)
In the present study, k l was evaluated from a simultaneous determination of [NO] and [NO2] in a steady-state situation after mixing NO2 and N205 in a large reaction chamber. The steady-state condition for N O is given by k5[N031 + ki [NO21[NO31 = k2[NOI [NO31
(b)
A comparison between k5 as evaluated in ref 1 and the product k l [NO,] under our experimental conditions shows that with k l =5X cm3 molecule-l s-I reaction 5 has negligible (
