Obi et ai.
2292 (9) T. T. Tuttle, Jr., and S. I. Weissman, J. Am. Chem. Soc., 80, 5342 (1958). (10) L. K. Pattersonand J. Lilie, Int. J. Radiat. fhys. Chem., 8, 129 (1974). (1 1) R. H. Schuler, P. Neta, H. Zemel, and R. W. Fessenden, J. Am. Chem. Soc., 98, 3825 (1976). (12) J. Lilie and R. W. Fessenden, J. Phys. Chem., 77, 674 (1973). (13) The 380-nm peak is slightly higher In Figure 2 b n in Figure 3 because in the former case G = 2.8 was used in the calculation, neglecting
the conversion of H into eaq-at high pH. (14) J. Murto, Acta. Chem. Scand., 18, 1043 (1964). (15) R. 0. Bates, ”Determination of pH, Theory and Practice”, Wiley, New York, N.Y., 1965, p 195. (16) A. Spernol and K. Wirtz, Z . Naturforsch. A , 8 , 522 (1953). (17) See, e.g., P. Neta, Adv. Phys. Org. Chem., 12, 223 (1976). (18) E. T. Strom, 0. A. Russell, and R. Konaka, J. Chem. fhys., 42, 2033 (1965).
Rate Constant Measurements for the Reactions of HCO with NO and O2 in the Gas Phase Karuhiko Shibuya, Takayukl Ebata, Klnlchl Obi, * and Ikuro Tanaka Department of Chemistry, Tokyo Institute of Technology, Ohokayama, Meguro-ku, Tokyo 152,Japan (Received May 3 I, 1977)
The behavior of the HCO radical produced by the flash photolysis of gaseous acetaldehyde has been studied with and without so-called radical scavengers at room temperature. The analysis of the results gives rate constants of (8.5 f 1.0) X and (5.6 f 0.9) X cm3molecule-l s-l for the reactions of the HCO radical with NO and 0 2 , respectively.
Introduction The HCO radical is a very important triatomic free radical which participates in many chemical reactions. Especially it is thought to be an intermediate species in the photochemistry of the polluted atmosphere1and in the combustion of most hydrocarbons.2 In spite of its important role in many oxidation phenomena, kinetic data on the elementary processes have been mainly based on product analyses. The first report of direct kinetic measurements on HCO was made by Washida, Martinez, and B a y e ~ .By ~ their flow experiment with a photoionization mass spectrometer, they reported rate constants for the reactions of the HCO radical with atomic and molecular oxygen of (2.1 f 0.4) X and (5.7 f 1.2) X cm3 molecule-l s-l, respectively. Studies of flash photolysis and kinetic absorption spectroscopy present additional complementary information on HCO radical reactions. We report rate constants of the HCO radicals with NO and O2in this paper. Experimental Section The HCO radical was generated by the flash photolysis of acetaldehyde with irradiation at wavelength longer than 200 nm. The reaction vessel was a cylindrical quartz cell of 10 cm diameter and 1m length. The cell contained a multiple reflection mirror s y ~ t e r nwhich ,~ increased the length of the absorption path up to a 54 m. Two flash photolysis lamps were mounted parallel to the cell and filled with 150 Torr of argon and 1.5 Torr of hydrogen. The electric circuit for the flash system was of the same design as reported by Welge et al.5 The capacitor bank consisted of eight capacitors (8 pF each charged to 6.4 kV) gave a 1300-5, 10-ps half-width light pulse. Decay of the HCO radical concentration was measured by the variation of the optical absorbance around 6J3.8 nm assigned to the (0, 9, 0)-(0, 0, 0) band of A2A”-X2A’ system? The absorption spectrum of the HCO radical was recorded by a Nikon P-250 spectrograph. The spectroscopic flash lamp was filled with argon at pressure of 1atm and the discharge energy was 50 J with a 8-ps half-width duration. All absorption spectra were recorded on Kodak The Journal of Physical Chemistty, Vol. 81,No. 24, 1977
2475 recording films and the measurements were made over the linear part of the characteristic curve of the plates. Acetaldehyde was purified by many thaw-freeze-pump cycles. NO and O2 (Takachiho Chemical Ind., 99.99%) were used without further purification. Helium was used as a diluent gas after passing through activated charcoal in a cold trap at liquid nitrogen temperature. The reaction mixtures were left in a 20-L glass bulb for about 4 h before experiments to assure complete mixing.
Results and Discussion The photolysis of acetaldehyde has been intensively studied in the gas p h a ~ e . ~ -Three l~ different primary processes occur with varying degrees of relative importance which depend on experimental conditions.11J2 -+
CH,CHO t hv
CH, + HCO CH, t CO CH,CO + H
1 -+
Conclusive evidence for primary process 1was found by flash photolysis12and process 2 was proved by the use of scavengers.’ The main dissociation process occurs via process 1in the photolysis of acetaldehyde at 313 nm and the quantum yield of process 2 increases with excitation energy. Hydrogen is a minor product from dissociative excitation in the wavelength region 235-350 nm which corresponds to its (T* n) absorption band.l0!l1 Primary process 3 is therefore certainly of little importance. Acetaldehyde (10, 20, and 30 Torr) and a mixture of acetaldehyde (20 Torr) and helium (510 Torr) were flash photolyzed and the concentration of HCO radicals was measured at various delay times. Figure l a shows the typical decay of the HCO radical in the photolysis of the mixture of acetaldehyde and helium. The decay curve apparently displays second-order disappearance as shown in Figure lb. If the reaction of the HCO radical with acetaldehyde was dominant as regards the disappearance of the HCO radical, the decay curve should be exponential. Therefore, the HCO radical reacts very slowly with the parent molecule, acetaldehyde, but is consumed by radical-radical reactions.
-
Reactions of HCO with NO and
h
n L
O2 in the Gas
2293
Phase
(b)
41
I I I I I I I I I I I l . l l l l j
OO
100
200
300
TIME ( p s e c ) Flgure 1. Time variation of (a) HCO concentration and (b) its inverse in the flash photolysis of a mixture of CHBCHO(20 Torr) and He (510 Torr). The solid lines represent the numerical simulation using the same initial concentration of 6 X l O I 3 molecule for HCO and CH3.
The reaction scheme relating to the decay of the HCO radical would be HCO + HCO -+ products CH, + HCO products -+
CH,
+ CH,
-+
products
(4) (5) (6)
Hence the following equations are obtained: -d[HCO]/dt = 2k4[HC0I2 + k,[HCO][CH3] (I) (11) -d[CH,]/dt= k,[CH3][HCO] + 2k6[CH3l2 The simulation was made to obtain the observed decay profile, using values of 3.7 X 4.3 X and 3.6 X cm3 molecule-' s-l for k4, k5, and kg, re~pectively,'~ and assuming that the initial concentrations of HCO and CH3 radicals were equal. The solid line in Figure l b shows the results of the simulation. The probable values of the initial concentrations of HCO and CH3 were determined to be (6 f 2) X 1013molecule cmM3. When a suitable amount of NO or O2 is added to the samples, reactions of the HCO radical with added scavengers predominate over the second-order radical reactions. Under these conditions, the decay will be governed by the reaction HCO t Q + products
(7)
where Q represents the radical scavengers, that is, NO or 02.
Typical decay plots of HCO radical are shown in Figure 2a ([NO] = 0.036 Torr) and 2b ( [ O , ] = 0.071 Torr). Apparently these plots display the pseudo-first-order decay. Therefore, most of HCO radicals react with NO and 02. In order to evaluate the rate constants, k7Q,numerical simulations were performed by considering the effects of the second-order radical reactions. Since the CH, radical also reacts with the scavengers CH,
+Q
-+
products
(8)
then the decay profiles are expressed as follows: -d[HCO]/dt= k7Q[HCO][Q] + 2k4 [ HCO]' + k5 [ HCO] [ CH3] -d[CHs]/dt = k,Q[CH,][Q] + kS[HCO][CHj] + 2k6[CH3I2
50
0
100
TIME ( p s e c ) Figure 2. Time variation of In ([HCO],/[HCO]o) for the reactions of
HCO with the scavengers NO and 02:(a) NO (0.036 Torr)/CH&HO (20 Torr)/He (510 Torr), k,No = 7, 8,9 X lo-'* om3 molecule-' s-', respectively, from the top: (b) O2(0.071 Torr)/CH CHO (20 Torr)/He (510 Torr), kToP= 4, 5, 6 X crn3 rnolecuIe-'s-', respectively, from the top.
Using the initial concentrations of HCO and CH3 radicals estimated above and the rate constants k4, k5,k6,and k*Q, the decay curves of the HCO radical were calculated at given pressures of scavengers (solid lines in Figure 2a and and 2b). The values of k8Qused were (4.8 f 0.4) X cm3 molecule-l s-l for NO and (3.0 f (3.2 f 0.3) X cm3 molecule-l s-l for 0.3) X and (5.2 f 0.5) X O2at totalpressures of 20 and 530 Torr, respectively, where the standard deviations are quoted.14 The values of k7Q or k7Q[Q]were obtained from the curves fit for the experimental plots similar to Figure 2a and 2b at various pressures of NO and 02.In the flash photolysis of the mixture of acetaldehyde and 02,the CH3C0 radical is produced and the following chain reactions are known to occur:15
+ 0, CH,CO, CH,CO, + CH,CHO CH,CO,H + CH,CO CH,CO
-+
-+
(9) (10)
However, after one flash was irradiated to the mixture, the concentration of O2was decreased only a few percent and the effect of the chain reaction was negligible within experimental error. Besides, in the experiment with NO, we could not observe the decrease of the concentration of NO due to the chain reaction. In order to avoid the pressure changes of the scavengers,we performed one flash for each mixture. The values of k7Q[Q]vs. the concentrations of NO and O2are plotted in Figures 3 and 4 , respectively. The linear relations obtained indicate that the rate constants, k7Q,are independent of the concentration of scavengers. The bimolecular rate constants k7Q,were evaluated from the slopes of these plots. The rate constants obtained are listed in Table I. The bimolecular reaction of the HCO radical with NO is proposed16
+ NO
+ CO
(111)
HCO
(IV)
but the information on the third body effect has not yet been reported. In our experiments, the value for k7N0
-+
HNO
(11)
The Journal of Physlcal Chemistry, Vol. 8 1, No. 24, 1977
2294
Obi et ai.
2 h
c
V
* 1
m
v
:‘o - 0
,
lk$( , , ,, 0
10
5
0
(NO1
x
0
15
lo2(Torr)
10
5
0
(02)
,
x
,
,
1 15
102(Torr)
Figure 3. Plots of the HCO radical decay rate against NO concentration in the photolysis of (a) NO/CH3CH0 (20 Torr) and (b) NO/CH3CH0 (20 Torr)/He (510 Torr).
Figure 4. Plots of the HCO radical decay rate against O2 concentration in the photolysis of (a) 0,/CH3CH0 (20 Torr) and (b) 02/CH3CH0 (20 Torr)/He (510 Torr).
TABLE I: R a t e Constants k,Qfor t h e Reactions of HCO Radicals with NO a n d 0,
was about 4 Torr and in our experiment the total pressures were 20 and 530 Torr. From their value and our results the rate constant for the reaction of the HCO radical and O2 did not depend on the total pressure within experimental error limits. These results suggest that the reaction of the HCO radical and O2 takes place through a bimolecular reaction in the pressure range studied.
Total press, Scavenger
Torr
NO
20 530
0,
20 530 4
R a t e constants,” c m 3 m o l e c u l e - ’ s-’
(8.6 * 0.9) x lo-’’ (8.4 i 0.9) x lo-” (14 i 2 ) x l o - ” * (6.0 i 0.9) X lo-” (5.3 t 0.7) x (5.7 i 1.2) x (3.8 i 0.6) x
References and Notes
a T h e errors c i t e d are estimates of t h e standard deviation. Reference 21. Reference 3.
*
remained constant when 510 Torr of helium was added to a mixture of acetaldehyde and NO (total pressure 20 Torr). Accordingly the termolecular reaction might be negligible in this system. The flash photolysis of mixtures of acetaldehyde and O2 was first carried out by McKellar and Norrish,17and the HCO absorption bands were not detected because of their experiments with too large a partial pressure of O2 P 2 . 5 Torr). Reactive collisions of HCO and O2 are thought to occur as f o l l o ~ s : ~ ~ ~ J ~ - ~ ~ HCO t 0, HO, t CO (12)
-
k;:y2
(13)
(14)
The HOzforming process was directly confirmed by LMR experiments,l8 and ab initio Hartree-Fock c a l c ~ l a t i o n s ~ ~ suggested that the association of HCO and 02 yielded the excited state of HCOB which decomposed readily. The total pressure at which Washida et al. measured the rate constants of the reaction of the HCO radical and 02
The Journal of Physical Chemktty, Vol. 81, No. 24, 1977
P. A. Leighton, “Photochemistry of Air Pollution”, Academic Press, New York, N.Y., 1961. A. G. Gaydon, “Spectroscopy of Flames”, Chapman & Hall, London, 1957. N. Washida, R. I. Martinez, and K. D. Bayes, Z. Naturforsch. A, 29, 251 (1974). J. U. White, J . Opt. SOC. Am., 32, 285 (1942). K. W. Welge, J. Wanner, F. Stuhl, and A. Heindrichs, Rev. S d . Instrum., 38, 1728 (1967). D. A. Ramsay, J. Chem. fhys., 21, 960 (1953); G. Herzberg and D. A. Ramsay, froc. R. SOC. London, Ser. A , 233, 34 (1955); J. W. C. Johns, S. H. Pfddle, and D. A. Ramsay, Discuss. Fara&y Soc., 35, 90 (1963). F. E. Blacet and J. D. Heldman, J. Am. Chem. Soc.,64, 889 (1942). G. Herzberg, R o c . Chem. Soc., 116 (1959). R. B. Cundall and A. S. Davies, frogr. React. Kinet., 4, 149 (1967). T. BBrces, Comp. Chem. Kinet., 5, 277 (1972). C. S. Parmenter and W. A. Noyes, Jr., J. Am. Chem. Soc.,85, 416 (1963). A. S. Archer, R. B. Cundali, G. B. Evans, and T. F. Palmer, Roc. R. SOC. London, Ser. A, 333, 385 (1973). V. N. Kondratiev, “Rate Constants of Gas Phase Reactions”, National Bureau of Standards, Washington, D.C., 1972. N. Basco, D. G. L. James, and R. D. Suart, Inf. J. Chem. Kinet., 2, 215 (1970); N. Basco, D. G. L. James, and F. C. James, !bid., 4, 129 (1972). C. A. McDowell and L. K. Sharples, Can. J. Chem., 36, 251 (1958). J. Heicklen and N. Cohen, Adv. fhotochem., 5, 157 (1968). J. F. McKellar and R. G. W. Norrish, Roc. R. SOC.London, Ser. A , 254, 147 (1960). H. E. Radford, K. M. Evenson, and C. J. Howard, J . Chem. fhys., 60, 3178 (1974). N. W. Winter and W. A. Goddard, 111, Chem. fhys. Lett., 33, 25 (1975). T. L. Osif and J. Heicklen, J . fhys. Chem., 80, 1526 (1976). C. B. Moore, private communication.
