Reaction of NO2 with
2507
03
(11) J. D. McDorsald, P. R. LeBreton, Y. T. Lee, and D. R. Hershbach, J. Chem. Phys., 56, 769 (1972). (12) K. C. Kim, Ph.D. Thesis, Kansas State University, Manhattan, Kan. 1973. (13) K. Karplus, 1-3. N. Porter, and R. D. Sharma, J. Chem. Pbys., 43, 3529 (1965). (14) D. L. Bunker, Methods Compot. Phys., 1 0 , 287 (1971). (15) (a) A. Ralston and H. S. Wilf, "Mathematical Methods for Digital Computers," Wiley-lnterscience, New York, N. Y. (1960); (b) A. Ralston, MTAC, 16, 431 (1962). (16) P. J. Kuntz, E. M. Nemeth, J . C. Polanyi, S. D. Rosner, and C. E. Young,J. Ch8.m. Phys., 44, 1168 (1966). (17) (a) The data for the various F f HR systems were taken from " J A N A F Thermochemical Tables" (U. S. Government Printing Office). (b) The data for the diatomic pairs of the D IC1 system were taken from B. Rosen, "Selected Constants Spectroscopic Data Relative to Diatomic Molecules," Pergamon Press, New York, N. Y . , 1970.
+
(18) W. A. Chupkaand J. Berkowitz,J. Chem. Phys., 54,5126 (1971). (19) G. Herzberg, "Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules," Van Nostrand, Princeton, N. J . , 1950, pp 108,101. (20) J. C. Polanyi, Accounts Chem. Res., 5, 161 (1972). (21) (a) R. Foon and N. S. McAskil, Trans. Faraday SOC.. 65, 3005 (1969); R . Foon and G. P. Reid, ibid., 67, 3515 (1971); (b) R. F. Walker and M. A. A. Clyne, University of London, Queen Mary College, private communication, 1972. (22) L. M. Raff, private communication. (23) 1.M. Raff and M. Karplus, J. Chem. Phys., 44, 1212 (1966). (24) D. L. Bunker and M. C. Pattengi1,J. Chem. Phys., 53,3041 (1970). (25) C. A. Parr, J . C. Polanyi, and W. H. Wong, J. Chem. Phys., 58, 5 (1973). (26) B. A. Hodgson and J. C. Polanyi, J. Chem. Phys., 55,4745 (1971) (27) D. L. McFadden, E. A . McCullough, Jr., F. Kalos, and J . Ross, J . Chem. Phys., 59, 121 (1973). (28) L. M. Raff,J. Chem. Phys., 44, 1202 (1966); 50, 2276 (1969).
The Reaction of Nitrogen Dioxide with Ozone C.
H. Wu, E. D.Morris, Jr.,* and H. Niki
Scientific Research Staff, Ford Motor Company, Dearborn, Michigan 4812 1 (Received April 6, 19731 Publication costs assisted by the Ford Motor Company
+
-
The mechanism and rate constant of the reaction NO2 O3 NO3 + 0 2 were investigated at the partsper-million concentration level. A chemiluminescent detection method was used to monitor the decay of O3 in excess NO2. From these data a rate constant of 4.4 X 1O-I' cm3 molecule-1 sec-I at 26" was determined with an uncertainty of &15%. Long-path infrared was used to measure the overall stoichiomeNO3 ~2 N205, the observed try (ANOz/A03). Since the primary reaction is rapidly followed by NO2 stoichiometry ranged from 1.65 to 2.00.
+
The reaction of NO2 with O3 plays an important role in photochemical smog and acts as a sink for both NO2 and O3.I The mechLanism proposed2 for this reaction is NO, 0, NO, 0, (1)
+
+
+
NO, NO, e N,05 (2) In the atmosphere, NzO5 is subsequently removed by reactions involving water and solid surfaces.3 In 1922, Wulf, Daniels, and Karrer4 measured the stoichiometry (A1\102/A03) of this reaction by titrating NO2 with 0 3 until the brown color of NO2 disappeared. From the initial concentration of NO2 (10-50 Torr) and the amount of 0:;added, they concluded that two molecules of NO2 were consumed for every molecule of 0 3 . A few of their data points gave results significantly less than 2, but they were attributed to a faulty ozone analysis. The most extensive kinetic data on this reaction are those of Johnston and They followed the concentration of NO2 by its absorption at 4400-4800 b, in a stopped flow reactor. 'The ozone concentration during the reaction was determined assuming a stoichiometry of 2. From the decay of 0 3 , they report a rate constant at 21" of 6.1 x cm3 molecule-I sec-1 with a standard deviation of 10% over the range of 1 to 10 Torr initial concentrations. An activation energy of 7.0 f 0.6 kcal/mol was obtained over the temperature range of 13 to 29".
Ford, Doyle, and Endow5 measured the rate of the NO2
+ O3 reaction in a 50-1. stirred-flow reactor using concentrations in the ppm range. (Throughout this paper the unit ppm refers to 760 Torr at 25" so that 1 ppm = 2.45 x 1013 molecules/cm3.) Ozone was determined in situ by an ultraviolet photometric method, while NO2 was measured a t the exit of the reactor by a continuous wet chemical method. The analysis of their data is based in the mechanism (1)-(2) which requires a stoichiometry of 2. However, the stoichiometry calculated from the NO2 and O3 concentrations given in this paper range from 0.9 to 4.8. Although the individual determinations scatter over a factor of 4, these authors derived an average rate constant of 3.3 x 10-17 cm3 molecule-1 sec-1. This is in reasonable agreement with the value of Johnston and Yost obtained at much higher concentrations. Scott, Preston, Andersen, and Quick6 monitored the decay of NO2 produced by the flash photolysis of O3 in NzO. From their high-pressure data (>400 Torr), they estimated k l as 1.3 x cm3 molecule-1 sec-l. Experiments performed at 100 Torr yield a value for lzl some 5-6 times smaller. This suggests the mechanism for the formation of NO2 even at higher pressures is not entirely correct. In a recent paper from this laboratory, Stedman and Niki7 used a chemiluminescent detector to follow the The Journalof Physical Chemistry, Vol. 77, No. 21, 1973
C. H. WU, E. D. Morris, Jr., and H. Ni ki
2508
first-order decay of small amounts of O3 (0.1-0.01 ppm) in the presence of 1-8 ppm of NOz. Under these conditions a rate constant can be determined on the assumption that 0 3 is consumed only in the primary step. They reported a cm3 molecule-1 sec-I. rate constant for hl of 6.5 X This was based on rather preliminary data and the uncertainties in their rate constant judged to be f50%. Thus there is considerable uncertainty in the reported rate constant for the reaction of NO2 with 0 3 . In addition, no mechanistic studies have been done under the conditions which were used for kinetic measurements. In the present work, more extensive kinetic data have been gathered using chemiluminescent detection of ozone to obtain a more precise rate constant. Experiments were also performed to investigate the reaction mechanism using longpath infrared spectroscopy.
Experimental Section The kinetics and mechanism of the N02-03 reaction were studied in the ppm range using two experimental methods. The kinetics was studied at atmospheric pressure in Nz in a 45-1. Pyrex bell jar. The jar was sealed to a Teflon-coated stainless steel base plate which was connected to a mechanical pump through a 1-in. line and a liquid nitrogen cooled trap. The reactor was evacuated below Torr for a minimum of 30 min between runs. The leak rate under vacuum was less than 1 0 - 3 Torr/min. Nitrogen dioxide was prepared by adding excess oxygen to prepurified NO. The NO2 was then frozen out and excess 0 2 was pumped off. The NO content was less than 0.1%. Ozone was prepared by passing high-purity oxygen through a silent discharge ozonizer. As a precaution, the 0 2 was first passed through heated CuO to remove hydrocarbon impurities. The O3 produced was trapped in silica gel at -78" and excess 0 2 was pumped off. The purity of the ozone was found to be 98 f 2% by decomposing a small portion and measuring the pressure rise. Ozone could be stored in the cold trap for up to a week without significant decomposition. Ultrahigh-purity nitrogen ( 30, a pseudo-first-order decay of 0 3 was observed. Figure 1 shows typical plots of In [ 0 3 ] us. time for several different NO2 concentrations between 1 and 10 ppm. The depen-
+
0 0
m
0
Time (rnin) Figure 1. Logarithmic decays of ozone at various initial NO2 concentrations. [No~lo/[o,]o > 30.
dence of the 0 3 decay slope, d In [03]/dt, is shown as a function of [NOz] in Figure 2, curve A. This plot is linear with a zero intercept. From the slope a bimolecular rate constant of 4.3 x l O - I 7 cm3 molecule-1 sec-1 was determined. The experimental precision of the slope measurements in these experiments was better than f 0 . 2 in the same units. To check the effect of wall treatment on the rate constant, a second set of experiments were carried out by replacing the 0 3 pretreatment of the vessel with The Journal of Physical Chemistry, Vol. 77, No. 27, 7973
c . H. Wu, E. D. Morris, Jr., and H. Niki
251 0
0.I
0.06 0.04
Figure 2. Decay of ozone as a function of [NO210 for the wall conditions: A , = 03 treatment; B, = NO2 treatment; C, = evacuation only.
NO2 pretreatment. Ozone again exhibited first-order decay a t various NO2 concentrations. A plot of d In [03]/dt us. [NO21 for data taken under these conditions is given in Figure 2 , curve B. This line has a slope corresponding to a cm3 molecule-I sec-I and a rate constant of 4.5 x small positive intercept. This slope is slightly greater than that found with 03-treated walls but is within experimental uncertainty. To study the wall effect further, kinetic data were obtained after subjecting the vessel to overnight pumping. When these data are plotted in the form d In [Oa]/dt us [NO21 (Figure 2, curve C), the slope gives a cm3 molecule-1 sec-l. This rate constant of 4.9 x value is greater than that observed with either O3 or NO2 pretreat men t . The dependence of the O3 decay rate on the initial ozone concentration was investigated a t a fixed [NO210 = 4.9 ppm. In this experiment [ 0 3 ] 0 was varied from 0.1 to 1.0 ppm. When [KO21 >> [ 0 3 ]d, In [O,]/]dt us. time is linear with a constant slope as shown in Figure 3. At higher 0 3 concentrations, these plots become curved as shown in this figure. The points are experimental data taken from the continuous recording of the ozone detector. The lines are from a computer analysis using reactions 1 and 2 (stoichiometry = 2) and a rate constant for h l = 4.4 x 10-17 cm3 molecule-1 sec-l. The experimental data do not start at time zero since it requires 1-2 min to fill the cell. Accordingly, the time scale of the experimental points has been shifted to take this into account. The calculation gives an excellent fit to all the data including those a t the higher 0 3 concentrations where NO2 is no longer in large excess. A number of calculations were performed assuming various stoichiometries between 1.0 and 2.0. The data show a good fit to calculations with a stoichiometry between 1.7 and 2.0. Below 1.7 the calculation exhibits significant deviation from the experimental data. These results are in agreement with the stoichiometric determinations made in the long-path infrared cell. Discussion
In all previous studies of this reaction the mechanism (1)-(2) and the corresponding stoichiometric value of 2 were assumed in the data analysis. Before proceeding t o
kinetic measurements. we wished to confirm the validity of this mechanism at the ppm concentration level. The stoichiometry measured in the long-path infrared cell was The Journal of Physical Chemistry, Vol. 77, No. 21, 1973
0.021
O.Ol0
'
2
I 4
6
I
8
1
0
Reaction Time ( m i n ) Figure 3. Concentration vs. time profiles of ozone at various initial concentrations with fixed [ N 0 2 ] o = 4.9 ppm. The points are experimental data while the lines are computed results using reactions 1 and 2 and k l = 4.4 X IO-" cm3 molecules-1 sec- l . 1.88 f 0.15 with ozone excess and 1.68 f 0.15 for NO2 excess. It was not possible to determine whether this difference was caused by a change in mechanism or by a systematic error in the experiments. A similar stoichiometric value could be inferred from 0 3 decay in nonpseudo-firstorder conditions. These measurements thus indicate the reaction is proceeding largely via the proposed mechanism. However, the measured stoichiometries tend to indicate other reactions may be occurring to reduce the ratio of IN02/-103 below 2.0 as predicted by the simple mechanism. The following reactions can be incorporated into the mechanism NO, NO2 NO NO2 0, (3)
+
+ + -
+
KO NO, 2N02 NO + 03 + NO, + 02
(4)
(5) The overall stoichiometry is reduced by competition of reaction 3 with reaction 2. The NO formed in reaction 3 can react either uia reaction 4 or 5 . A computer integration of the mechanism (1)-(5) was carried out using rate constants suggested by Johnston.11 The predicted stoichiometry is 1.95. In addition, an alternative path for reaction 1
NO, + O3
-
20,
+ NO
(6) cannot be ruled out entirely. If this reaction is occurring, it would have a strong effect on the stoichiometry. For example, if h l / h 6 = 20 the overall stoichiometry is calculated to be 1.76. The present stoichiometric data indicate the reaction proceeds greater than 90% via mechanism (1)-(2) at ppm concentrations. This is supported by the observation of an N z 0 5 yield approximately one-half the initial NOz. A stoichiometry of 2.0 is an upper limit. Several possible minor reactions tend to reduce this value under the present experimental conditions.
251 1
Kinetics of Gas-Phase Reactions of Ozone with Olefins
The equation for O 3 decay has been demonstrated in this study to be d[O3]/dt = - h [ N O ~ ] [ 0 3 ] from the decay of O3 and the dependence on NO2 where h = 4.4 x cm3 molecule-1 sec-l. The scatter in the individual determinations is less than 5%. Possible systematic errors in the rate consttint could result from the absolute calibration of reactant concentrations. The NO and NO2 calibrations were based on a standard mixture of NO which was analyzed independently by two laboratories reporting results within f4%. The absolute calibration of ozone is not necessary for the rate constant determination. However, the calibration was done for purposes of cross checking and investigation of the second-order region. The temperature of the reaction mixture was monitored and found to be 26 f lo,which results in a maximum error of 3% assuming 7 kcal/mol for the activation energy. The stoichiometric value does not affect the rate constant derived from 0 3 decay unless reactions such as (6) which consume extra ozone are occurring. In that case the rate constant measured here would be too large by the fraction k 6 / k l . However, this ratio is not likely to be greater than 0.05. The rate constant obtained from results in the "untreated" reactor is larger than that obtained with either 0 3 or NO2 treatment. This is likely a major cause of the larger value reported by Stedman and Niki. Based on these considerations, the overall uncertainty in the rate constant is estimated to be f15%. The present value of k l = 4.4 f 0.6 x cm3 molecule-I sec-l ILSlarger than that determined by Ford, et ~ l . and , ~ that estimated by Scott, et aL6 However, both of these studies exhibited large uncertainty. Johnston and cm3 molYost report a rate constant of 6.1 f 0.6 X
ecule-1 sec-l a t 21O.2 Using the reported activation energy of 7.0 f 0.6 kcal/mole, their value a t 25" is 7.2 f 1 x The rate constant derived in the present work is somewhat smaller than this. However in view of the vastly different experimental conditions used in their work, this should be considered in good agreement with the present value. After the completion of this work another report of h i appeared. Ghormley, Ellsworth, and Hochandel12 observed the decay of NO2 produced in the flash photolysis of O3 in N20. Due to an error in their rate equation, the rate constant as printed is incorrect. With the proper kicm3 netic equation, their data yield a value of 3.2 X molecule-1 sec-l for k1.13
References and Notes (1) H. Niki, E. E. Daby, and B. Weinstock, Advan. Chem. Ser. No. 113, 16 (1972). (2) H. S. Johnston and D. M. Yost, J. Chem. Phys., 17,386 (1949). (3) E. D. Morris, Jr., and H. Niki,J. Phys, Chem., 77, 1929 (1973). (4) 0. R. Wulf, F. Daniels, and S. Karrer, J . Amer. Chem. Soc.. 44, 2398 (1922). ( 5 ) H. W. Ford, G. J. Doyle, and N. Endow, J . Chem. Phys., 26, 1336 (1957). (6) P. M. Scott, K. F. Preston, R. J. Andersen, and L . M. Quick, Can. J. Chem., 49, 1808 (1971). (7) D. H. Stedman and H. Niki, "Kinetics and Mechanism for the Photolysis of NO? in Air," Ford Motor G o . , Scientific Research Staff ReportJan 17, 1973. (8) D. H. Stedman, E. E. Daby, F. Stuhl, and H . Niki, J . Air Pollut. Contr. Ass., 22, 260 (1972). (9) L. P. Breitenbach and M. Shelef, J . Air Pollut. Contr. Ass.. 23, 128 (1973). (IO) F Cramarossa and H S Johnston, J Chem Phys 43, 727 (1965) (11) H . S . johnston, "Project Clean Air," University of California, 1970. (12) J. A. Ghormley, R. L . Ellsworth, and C. J. Hochandel, J. Phys Chem.. 77. 1341 11973). (13) Private communication'.
Kinetics of Gas-Phase Reactions of Ozone with Some Olefins D. H. Stedman, C. H. Wu,* and H. Niki Scientific Research Staff, Ford Motor Company, Dearborn, Michigan 48727
(Received February 22, 1973)
Publication costs assisted by the Ford Motor Company
Gas-phase reactions of ozone with ethylene, propylene, trans-2-butene, and 1-hexene were studied a t 299 f 2°K using a 45-1. static reactor and low reactant concentrations (-mTorr) a t 760 Torr total pressure. Ozone was analyzed by the NO/O3 chemiluminescence method, while hydrocarbons were monitored by gas chromatography. The kinetics was studied in both hydrocarbon and ozone excess conditions in the presence and absence of 0 2 . Absolute values of the second-order rate constant were determined to be 1.55 f 0.15, 12.5 f 1, 275 f 23, and 11 f 1.5 X 10-lS cm3 molecule-1 sec-1 for ethylene, propylene, trans-2-butene, and 1-hexene, respectively. For propylene and trans-2-butene, a stoichiometry of unity was observed in both nitrogen and air diluents.
Introduction Ozonolysis of simple olefins plays a key role in the formation of photochemical smog.lJ As a result, the kinetics of the gas-phase reactions of ozone with olefins has been g available kinetic the subject of numerous ~ t u d i e s . ~ -The data on ethylene, propylene, trans-2-butene, and 1-hexene
are summarized in Table I. Most of the studies listed here were carried out a t low reactant concentrations (--Torr) in air near atmospheric pressure. While the rate constants for 1-hexene show a general agreement, there is considerable discrepancy in rate constants and stoichiometries for the other olefins. In particular, an order of magnitude variation exists among the reported values for trans-2-buThe Journal of Physical Chemistry, Vol. 77, No. 2 1, 7973
