J. Phys. Chem. 1982, 86, 3424-3430
3424
Mechanism of the Photooxidation of Formaldehyde Studled by Flash Photolysis of CH20-Op-NO Mixtures Bernard Veyret, Jean-Claude Rayez, and Robert Lesclaux Laboratolre de Chimis Physique A, Universlt6 de Bordeaux I, 33405 Talence Cadex, France (Received: December 30, 1981; In Final Form: April 2, 1982)
The mechanism of the chain process leading to formic acid in the photooxidation of CH20 has been studied with the flash photolysis technique. Mixtures of CH20, Oz, and NO were photolyzed and the rate of appearance and yield of NOz were monitored. Kinetic simulations of both sets of data allowed the determination of the rate constants for the main reactions H02 + CHzO OzCH20H(6), OCHzOH + Oz HOz + HCOzH (8), k8 = (3.5 f 1.6) X It9 = (4.0 f 1.9) X lo-'' cm3 OCHzOH + NO products (9) ( k , = (7.5 f 3.5) X molecule-' s-'). Quantum calculations provided estimates of the heats of formation for the radicals involved. The effect of temperature was investigated, suggesting the importance of the decomposition of the radical HOCHzO into H atom and formic acid. The validity of the global scheme is discussed along with its importance for the removal of CHzO and the production of formic acid in the atmosphere.
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Introduction The mechanism by which formaldehyde is photooxidized in the atmosphere has been the subject of many investigations. Although a basic understanding has emerged from these studies, recently, some new insights have been gained about the formation of formic acid following the suggestion made by Su et al.' that HO, radicals add to formaldehyde as the first step in the photooxidation process. This mechanism was proposed after the observation of H02CH20H by FTIS during the photolysis of formaldehyde in the presence of Oz. Su et aL2have tested the validity of the mechanism through simulation of the product formation in the photolysis of C12-02-CH20 mixtures. The steady-state methods employed by these authors allowed only the determination of ratios of several elementary rate constants. By combining these ratios with other rate constant estimates based on analogy with similar simple species, they could derive estimates for the main rate constants. In this paper we report a different method for testing the mechanism: experiments were done by flash photolysis of CHZ0-O2-NO mixtures. HO, radicals are formed shortly after the flash and react with formaldehyde and nitric oxide. The appearance of the NO2 product is monitored for various mixture compositions and kinetic simulations are carried out.
TABLE I
Experimental Section Photolysis of CH20-02-N0 mixtures was studied with a flash photolysis apparatus that has been already de~ c r i b e d . The ~ absorption of the NOz product wqs monitored by sending the beam from an argon laser directly into the multiple pass system (set to a 14-m optical pathlength). The flash lamps were always operated at 110-5 electrical energy and - 5 - ~ sduration. The preparation of formaldehyde and handling of the gases were the same as in experiments described previou~ly.~ We observed a near exponential growth of the NO2absorption. Kinetic analysis of the oscilloscope trace yielded two pieces of information for each shot: the first-order rate of NO2 growth, KNo2,and the amount of product, [NO,-],,
calculated from the extinction coefficient €NO, = 36 M-' cm-' at X = 514.5 nm. Simulations were carried out on a HP85 microcomputer with a standard Runje-Kutta procedure for solving the set of differential equations.
(1)F. Su, J. G. Calvert, J. H. Shaw, H. Niki, P. D. Maker, C. M. Savage, and L. D. Breitenbach, Chem. Phys. Lett., 65, 221 (1979). (2)F. Su,J. G. Calvert, and J. H. Shaw, J. Phys. Chem., 83, 3185 (1979). (3)B. Veyret and R. Lesclaux, J.Phys. Chem., 85, 1918 (1981). 0022-3654/82/2086-3424$0 1.2510
2 2
5 5
100 100
0.36 0.48
7.8 7.8
5
2.5
100
0.83
3.5
1.oo 1.25 1.41
1.6 2.4 2.4 3.2 2.2 3.4 3.0 3.4 3.0 3.6 3.2 3.8 4.1
10 10 10 10 10 10 10 10 10 10 10 10 10
2 5 5 5 10 20 20 45 45 45 45 45 45
50 100 100 200 50 200 200 20 50 100 100 200 200
2.08 0.74 1.11 1.67 0.12 0.41 1.00 0.80 1.65 1.60
20 20 20 20 20
5 5 20 20 20
100 100 15 20 200
1.80 1.60 0.17 0.20 1.92
1.9 1.6 1.6
30
5
100
2.70
1.3
1.5 1.6
Results and Discussion
Room Temperature Photolysis of CH20-02-N0 Mixtures. Initial concentrations of CHzO, 02,and NO were varied over a wide range within the limits imposed by the performances of the apparatus (KNo2I2 X lo4 s-l and [NO,]- L 0.7 mtorr). Results of the experiments are listed in Table I. Two points about these results are especially noteworthy: (i) the first-order growth of [NO,] is in every case slower than that predicted by considering the reaction of HO, with NO alone (this would give KN0 = 2.5 X lo4 s-l for 100 mtorr of NO); (ii) the yield of dozis always larger than the initial concentration of HO,. (Initial concentrations of H atoms formed during the flash were measured in separate experiments with i-C4H8and NO as 0 1982 American Chemical Society
Mechanism of the Photooxidation of Formaldehyde
The Journal of Physical Chemistry, Vol. 86, No. 17, 1982 3425
Reaction 7 of HOCH202(RO,) with NO was considered similar to that of CH3O2 and a starting value for k7 was 7.5 x W2cm3 molecule-' s-'. Reactions 8 and 9 were at first assumed to be similar to those of the methoxy radical: CH30 + O2 CH20 + H 0 2 (8')
+ - +
CH30
Figure 1. Mechanism used in the simulations.
H atom quenchers. As shown below H 0 2 radicals are produced in the oxidation of HCO and [HO2l0= 2[HCOIo = 2[Hjo. Hence [HO?lo 1.25 X 10-4[CH20].) These two facta indicate clearly that a chain process is occurring. Formyl radicals and hydrogen atoms are formed in the photolysis of formaldehyde: N
CH20 H + HCO (1) H atoms react with either formaldehyde or oxygen: H + CH20 -* H2 + HCO (1) H + 0 2 + M - HOP + M (2) HCO radicals react with oxygen by the fast reaction HCO + 0 2 HO2 + CO (3) Thus because of the initial conditions ([CH20]1 2 torr, [O,] 1 2 torr) and since reactions 1 and 3 are fast, H 0 2 radicals are formed shortly after the pulse in a known amount. H 0 2radicals can react with NO and enter a chain process forming NO2 without consuming H02: chain process A HO2 + NO -* NO2 + OH (4) OH + CH2O -* H20 + HCO (54 HCO 02 -* CO + HO2 (3) The addition of H 0 2to formaldehyde suggested by Su et al.' can initiate a second chain process: chain process B HO2 + CH2O --c H02CH20 HOCH202 (6) HOCH202 + NO -* HOCHzO + NO2 (7) HOCH20 + 0 2 -* HCOOH + HO2 (8) The chain termination reaction for process B is HOCHZO + NO -.* HOCHzONO (94 HCOOH HNO (9b) This basic scheme shown in Figure 1 was included in computer simulations with the rate constants found in the cm3 molecule-' literature for reaction 4: k4 = 7.9 x s-l .4 Rate constants for reactions 6-9 were treated as parameters for the simulations. Reactions 1, 2, 3, and 5a were considered as instantaneous owing to the experimental conditions. (k3[02]> 1.8 X lo5 s-' with k3 = 5.6 x cm3 molecule-' s-'.~) The range for the variation of k6 was 1x 1O-lL32 x cm3 molecule-' s-l based on two earlier estimation^.^^^ +
+
+
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+
(4) C. J. Howard, J . Chem. Phys., 71,2352 (1979).
NO
CH30N0
@a')
CH2O HNO (9b') Simulations were carried out with accepted values for k8, and ky (6.1 X and 2.4 X lo-'' cm3 molecule-' s-', respectively6.'). The starting value for k6 was 1 x cm3 molecule-' s-'. The agreement with the experimental data was poor; the simulated yields of NO2were much too high because of the weak coupling between processes A and B. Increasing k6 to 10 X cm3 molecule-l s-' did improve the fit but large discrepencies remained in the case of slow NO2 buildup and the effect of O2 was not simulated at all. It had been suggested8 that the products of reaction 5 could be H 2 0 and an excited HCO* radical which could decompose rapidly: OH
-- + +
+ CH20 HCO*
HzO
H
HCO*
CO
(5b) (10)
This process was added to the simulations thereby slowing down process A. The agreement between simulation and experiment remained poor; simulations gave too small values for KNo2and a nonexponential shape for the growth of the NO2 absorbance. The existence of reaction 10 is discussed further below. The effect of [O,] variations was the clue to improvements in the scheme: [N02],/[H0z]o increases with added O2 (cf. Table I). In the absence of reaction 10, process A is independent of [O,] since reaction 3 is always much faster than reaction 4. Thus any effect of [O,] can only come from process B. With the value of k8 (= k,)that we initially used, process B is independent of [02]because of competition between reactions 8 and 9. Therefore we examined more closely the similarities between HOCH20(RO)and CH30. The presence of an extra oxygen atom in RO could increase the lability of the hydrogen atoms, thereby implying a higher value of ke. This hypothesis was confiied by means of thermochemical and quantum calculations (MNDO program) on RO which gave a weaker C-H bond strength in RO as compared to CH30 (see Appendix). The weak C-H bond in RO could be the source of H atoms when the temperature is increased. This result also implies that reaction 9b could be faster than reaction 9b' suggesting that k, could be significantly greater than k y. Letting k8 increase to 3.5 X cm3 molecule-l s-' considerably improved the agreement between the simulation and the experiment. Results of the simulations are shown in Figures 2-4 where KNoz and [NO,], are plotted vs. [CH20],[NO], and [O,], respectively. Figure 2 shows that KNo2goes to zero with [CH20] whereas [NO,], is still large at low formaldehyde con2 mtorr when centration ([CH,O] = 2 torr, [NO,], [HO,], is only equal to -0.25 mtorr). This behavior is easily explained by considering that, at low [CH20],the
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(5)B. M. Morrison,Jr., and J. Heicklen,J . Photochem., 15,131(1981). (6)J. R. Barker, S. W. Benson, and D. M. Golden, Int. J. Chem. Kinet., 9,31 (1977). (7) L. Batt, R. T. Milne, and R. D. McCulloch, Int. J. Chem. Kinet., 9,567 (1977). (8) A. Horowitz,F. Su, and J. G. Calvert,Int. J . Chem. Kinet., 9,1099 (1978).
3426
The Journal of Physical Chemistry, Vol. 86, No. 17, 1982
Veyret et al.
KN02~10-4 ( 5-11
E04, ( m Torr)
't
0 [H*CO]
10
(Torr)
Flgure 2. Plots of K, and [NO,], vs. [H2CO]. ([021 = 5 torr; [NO] = 100 mtorr.) Solid Ihe Is from the kinetic simulations.
0 0
b
100
20
[NO] (mTorr) Flgure 3. Plots of K,,,, and [NO,], vs. [NO]. Solid line and dotted line correspond to slmuiations with 10 torr of CH,O-45 and 5 torr of O2 respectively.
20
30
40
50
[02] ( T o r r )
Floure 4.
Plots of K, and [NO,], vs. [O,]. [H2CO] = 10 ton, [NO] = 50 mtorr, for closed circles and 100 mtorr for open circles. Lines
are output of the simulation.
path out of process A (reaction 6) is slow and NO2is being formed mainly through process A. [NO,], will obviously when reaction 5a beapproach zero at very low [CH20] comes inefficient. The effect of [NO] as shown in Figure 3 is twofold when [O,] is large, the plot of [NO& vs. [NO] goes through a minimum corresponding to an equivalent efficiency for processes A and B. From thereon, when [NO] increases, process A becomes dominant as a result of competition between reactions 4 and 6. As [NO] decreases, reaction 6 becomes faster than reaction 4 and process B becomes dominant, while the termination reaction for process B becomes slower. This explanation remains valid for low [O,]but then process B cannot be efficient at low [NO] because of the less rapid reaction 8. This situation is again different in Figure 4 which suggests finite intercepts for the plots of KNoland [NO,], vs. [O,]. This is easily explained by considering that process A is independent of [O,] as long as reaction 3 is fast enough ([O,] > 0.5 torr or k 3 [ 0 2 ]> 9 X lo4 s-l). A t low oxygen concentrations process A is dominant while at higher [O,], process B starts yielding NOz through reactions 8 and 4 therefore slowing down the growth of [NO,]. Both KNOl and [NO,], will go to zero with [O,] as less and less HOz will be formed and both processes A and B will cease functioning. Thus the observed agreement between the simulation and the experiment can be rationalized in terms of the roles of processes A and B and the path out of process B through reaction 9. Several possible radical-molecule reactions were considered and the effect of their inclusion into the global scheme was studied separately.
The Journal of phvsical Chemistry, Vol. 86, No. 17, 1982 3427
Mechanism of the Photooxidation of Formaldehyde
Two reactions of formaldehyde are RO + CHzO ROH + HCO ROz + CHzO ROzCHzO +
(11)
(12) Reaction 11 is similar to that of H atom abstraction by the methoxy radical CH30 CH20 CH30H HCO (11') cm3 molecule-' s-l. for which k4 1 X Quantum calculations on ROH gave a bond energy for the 0-H bond similar to that in CH30H (see Appendix) suggesting that reactions 11 and 11' are similar. Reaction 11, in competition with reactions 8 and 9, will become important when [O,] and [NO] are both low. In that case reactions 12 may also become important since reaction 7 becomes slower. Reaction 12 is analogous to that between HOz and CHzO and is likely to occur but probably much more slowly than reaction 6. The product of this latter reaction is likely to decompose or to react with O2to form HOz and thus reaction 12 could be part of process B but would not yield any NOz. Thus in our experimental conditions, reaction of R02with NO, is the only reaction that we need to consider, except in the case of low [NO] (120 mtorr), where we then have to take into account reaction 12 and the radical-radical reactions which are considered below. The reactions of H02, R02, and RO with NO2 are possible but very unlikely since [NO,], remains small and since the growth of [NO,] and the decays of [H02] and [RO,] are simultaneous. There are several radical-radical reactions that need to be examined although experimental conditions were chosen in order to minimize their roles in the mechanism. Formyl radicals react with hydrogen atoms and with themselves but always more slowly than with oxygen. Mutual recombination of H 0 2 radicals does occur but is much slower than reactions of H 0 2 with NO and CH20: HO2 + HO2 H202 + 02 (13) Hydroxyl radicals react very rapidly with CH20 (k,[CH20] I8.5 X lo6 s-') and their reactions with NOz, HOz, RO, and NO are negligible. However, it has been postulated that the products of reaction 5 besides HCO and H20could be H and HCOOH? OH + H2CO (RO) H + HCOzH (5~) Our simulations show clearly that few H atoms if any are formed since their presence significantly affects the results of the simulations which cannot then be readjusted in order to fit the experimental data. We believe that reaction 5a is the main path for reaction 5 and that formic acid is rather formed via process B. Reaction 5c has also been rejected by Morrison and Hei~klein.~Horowitz et al. observed an H2forming chain reaction in the photolysis of CH20-02 mixtures. They suggested the involvement of reactions 5c and 10 but did not considered reaction 6. We believe that reactions 5c and 10 do not occur to any significant extent. Reactions between peroxy radicals certainly occur in the atmosphere but not in our system because of the presence of NO which reacts rapidly with both RO, and HOP. RO2 + RO2 2 R 0 + 02 (144 HCOOH + ROH + 02 (14b) (15) HO2 + ROz -* ROOH + 02 Reaction 13 is a termination reaction and reaction 14, depending on the ratio k14+,/k14e, can also interrupt the +
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+
+
+
+
+
+
+
+
(9)B. M. Morrison, Jr., and J. Heicklen,J.Photochem., 13, 189 (1980).
2.0
t
KN02x10-4 1.21
.,'
V I I-
5.
.___---
4-
I-\
I
0
100
20d
[NO] ( m T o r r ) Flgure 5. Plots of K,, and [NOz], vs. [NO] ([HzCO] = 10 torr, [0 ] = 45 torr). SOMline comesponds to simulations wlth k, = 7.5 x 10-74 cm3 molecule-' s-' and dotted line with a 30 % increase (a) and 30 % decrease (b) of the k , value.
chain process. Formation of the peroxynitrates H02NOz and R02NOzis negligible, being much slower than reactions of the peroxy radicals with NO. Thus kinetic simulations included no radical-radical reaction but only the reactions shown in Figure 1. Final values for the rate constants were obtained after a good fit was reached between the output of the simulations and the experimental data, as shown in Figures 3-5. The effects of variations of the three main rate constants ks, k8, and kg was investigated by running the simulations with a &30% margin on one of the rate constant while keeping the two others at their final values. Two such sensitivity analysis are shown in Figure 5. Most of the experimental points are within the limits, and this indicates that the error is smaller than 30% if we assume no errors were made in the other rate constants, in the initial concentrations of reactants and radicals, or in the kinetic analysis. The total error on the rate constant values is estimated to be less than 50%. Simulations of Continuous Irradiation Experiments. Other groups have carried out photolysis of formaldehyde in the presence of oxygen with and without N0.2JoJ1 We tried to fit their experimental data using our kinetic scheme and rate constant values. In their experiments the irradiation was continuous and some of the radical-radical reactions and photolysis of the products were important. Reactions 14 and 15 were added to our scheme along with reactions 16 and -6: ROOH + hv HCOOH + H20 (16) ROz -* HO2CHzO HOz + CH20 (-6) Reaction 16 was suggested by Su et al.2 and is certainly a major path for HCOOH formation in the atmosphere. +
+
(10)B. M. Morrison, Jr., and J. Heicklen, J. Photochem., 11, 183 (1979). (11)H. Niki, P.D. Maker, C. M. Savage, and L. P. Breitenbach, Chem. Phys. Lett., 72,72 (1980).
The Journal of phvsical Chemistry, Vol. 86,No. 17, 1982
3428
Veyret et al.
TABLE I1 [CHZOl, torr
[NO], mtorr
2.0 2.0 2.0 0.4
6 10 26 10
@(NO,)
0 (HCOOH)
exptl
calcd
exptl
calcd
41 13
15.5 7.5 3.6 9.4
14 11 8 29
23.2 14.4 13.3 41
8
