CURRENT RESEARCH Quenching of Singlet Molecular Oxygen by Some Atmospheric Pollutants Ralf-Dieter Penzhorn" and Hans Gusten Kernforschungszentruim Karlsruhe, lnstitut fur Radiochemie, 75 Karlsruhe, Germany
Ulrich Schurath and Karl Heinz Becker lnstitut fur Physikalische Chemie der Universitat Bonn, 53 Bonn, Germany
w Room temperature rate constants for the deactivation by SOz, H2S, of singlet molecular oxygen 02(a11,) CH3SH, CHaCOCH3, CH20, ambient air, and dry synthetic air were measured in a 220-m3 low-pressure reactor under static conditions. Emission a t 7620 8, from 0 2 ( b 1 S g + ) ,which is continuously produced from the energy pooling reaction 202(11,) O Z ( ~ C ~ + ) 02(3?;,-), was taken as a relative measure of the 0 2 ( l A g ) concentration in the system. Thle compounds deactivate singlet oxygen at rates comparable to or lower than ground state molecular oxygen. The rate constants for deactivation by CH3SH, dry air, and ambient air are critically compared with literature values. It is concluded that singlet oxygen does not contribute to the atmospheric conversion of the trace gases investigated.
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+
Since the role of singlet molecular oxygen 0 2 ( I A g ) in air pollution is still a matter of discussion ( I ) , it is evident that more data on the physical and chemical quenching of this excited molecule by atmospheric gases and air pollutants are needed. We wish to report some 0 2 ( l A g ) quenching rate constants obtained from experiments in a 22O-m3 stainless steel sphere. The reactor, as well as the experimental method, has been described in previous publications (2, 3). 'Briefly, 0~(~1,) was generated by passing a stream of high-purity oxygen (Air Liquide A 48) through a microwave discharge. The effluent gas was fed into the previously evacuated reactor up to a pressure of 7 mtorr. Atomic oxygen was eliminated by passing the discharged gas over a HgO film (4) and by reaction with 0.4 mtorr NO2 which was introduced into the reactor before each experiment. A quenching gas was subsequently added and the first order decay of 0~(~1,) in the reactor measured a t various pressures. The decay can either be observed directly by measuring the infrared atmospheric band 0~(~1, 38,-) a t 1.27 p using a lead sulfide detector, or indirectly by measuring the intensity of the atmospheric band 02(1Sg+ 32,-) a t 7620 A, which decays twice as fast as the infrared atmospheric band, with an infrared sensitive photomultiplier. Both methods have been used in this laboratory and found compatible ( 2 , 3 ) . The second method was chosen in this work because of its higher sensitivity. The atmospheric band was isolated by a Schott interference filter centered a t 7600 8, and the
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intensity measured with a liquid nitrogen-cooled EM1 9558 B photomultiplier. 0&2,+) is continuously produced in the reactor from OZ(ll,), by the well-known energy pooling Reaction 1:
This reaction, as well as others which are second order in 0 2 ( l A g ) , are slow and do not interfere measurably with the first order decay of 0 2 ( l A g ) under our experimental conditions ( 2 ) . Removal of OZ(~Z,+) by deactivation a t the surface of the reaction chamber, in the gas phase, and by radiation is so fast compared with the decay of its precursor 02(11,), that steady state kinetics may be applied to calculate the equilibrium concentration of 02(1S,') which is proportional to the square of the OZ(l1,) concentration. Thus, the atmospheric band a t 7620 8, gives a relative measure of the Oz(l1,) concentration. It has been shown that the decay constant l / T p , measured at 7620 A, is related to the Oz(l1,) quenching rate constant k.wi of various quenching gases M, by the relationship
where k , corresponds to the radiation transition probability of 0~(~1,) and k,. is its rate constant of wall deactivation. When all pressures except that of the quencher under consideration, are kept constant, a linear plot between 1 / T p and M i s expected. Data from a series of experiments with 0.4 mtorr NOz, 7 mtorr 0 2 , and varying pressures of SOz, HzS, CHJSH, CH3COCH3, CH20, ambient air, and synthetic air a t room temperature are plotted in Figures 1-3. In each case the expected linear relationship is observed above a certain pressure. The curvature a t low pressures of SOz, ambient air, and other gases is due to adsorption phenomena which lower ku,. The effect is reversible after prolonged pumping of the reactor. The rate constants calculated from the slopes of the straight lines are summarized in Table I. Some of these may actually represent the sum of two rate constants, kz and k3, for physical quenching and chemical reaction:
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products (3) There are some interesting points concerning the quenching efficiencies of these primary and secondary air pollutants. The O&A,) has been discussed as a possible Volume 8, Number IO, October 1974
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Table I. Rate Constants for Deactivation of O,(la,) by Some Primary and Secondary Air Pollutants at Room Temperature 1018
Ackerman et a l . ( 9 )
This work
Compound
so2 Hz CHaSH C H 3COC H3 CH20
"Ambient air" Dry synthetic air
2GI
X k (cm3 molecule-' sec-1)
"
0.39 f 0.09 2.1 zt 0.2 10 =t3 16 i 4 23 i 3 3.4 i 0.4 3.0 =!= 0.1
"
I
jC
Clark and Wayne ( 1 3 )
38
4.4 i 1.3
'
'
"
M :mtorrl
1
133
Figure 1. Quenching of 0 2 ( ' A g ) by 0 HzS, 0 CH3SH, CH3COCH3, and 0 CH2O; p ( 0 2 ) = 7 mtorr and p(N02) = 0.4 mtorr I
.'t 0
i 100
200
300
40 0
M [mtorr]
+
Figure 2. Quenching of 0 2 ( ' A g ) by 0 SO2 and d r y synthetic air; p ( 0 2 ) = 7 mtorr and 4 mtorr, respectively, and p(N02) =
0.4 mtorr
Figure 3.
Quenching
of 0 2 ( ' A g ) by
ambient air;
mtorr andp(NO2) = 0.3 mtorr 908
Environmental Science & Technology
p(O2) = 4
oxidant of SO2 in the lower atmospheTe ( 5 ) . Thermodynamically, only Reaction 4 need be considered, since the reaction yielding so3 and O(3P) is 13.6 kcal/mol endothermic:
SOz + O 2 ( ' A p )
+M
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So4 + M
(4 )
The SO4 radical has been suggested as an intermediate in the homogeneous atmospheric photooxidation of SO2 (6). Our data indicate that the quenching of 0 2 ( 1 1 , ) by SO2 is first order in SO2 pressure and is extremely slow which excludes the three-body Reaction 4 as an important step of the modification of SO2 in the atmosphere. It can be shown from our measurements with ambient air and from our previously published rate constants for the quenching of 0 2 ( I A g ) by pure atmospheric gases (2), that in an atmosphere containing 0.1 ppm of S02, only one out of 108 singlet oxygen molecules will be quenched by S02. A comparison of this figure with the estimated singlet oxygen production rate of 0.5 ppm hr-1 in polluted atmospheres (7, 8) shows clearly that the SO2 photooxidation must occur via other routes. Ackerman et al. (9) determined the quenching rate of O2(lIg) with various sulfur compounds in a flow system by measuring the intensity of the OZ(~A,)emission a t 1.27 along the flow tube. They state that methyl mercaptan is essentially quenched physically. Their rate constant for the quenching of singlet oxygen by CHsSH, Table I, is higher than our value by nearly a factor of 4. The authors used mercury distillation through the discharge to remove oxygen atoms. We know, however, from previous experiments that HgO alone does not completely remove oxygen atoms from discharged oxygen, since measurable amounts of NO were evolved when oxygen from a Hg/O2 discharge was introduced into the 22O-m3 reactor containing 0.5 mtorr NO2 as a scavenger. Furukawa and Ogryzlo (10) observed that traces of oxygen atoms interfere seriously with the determination of rate constants for the quenching of singlet oxygen by organic amines. This interference is due to radicals from reactions of O(3P) atoms with organic compounds, which react with singlet oxygen much more rapidly than stable molecules (Becker et al., to be published). Since no oxygen atoms were present during our decay measurements in a static system, we conclude that our value of the CH3SH rate constant is correct. The fact that H2S is a poorer quencher of O ~ ( 1 1 than ~) CH3SH may be attributed to the higher ionization potential of H2S. A linear correlation between the ionization potentials and the logarithms of the quenching rate constants of amines and sulfur compounds, respectively, has been observed by Ogryzlo and Tang (11) and by Ackerman et al. (9) and was ascribed to charge transfer interaction between quencher and singlet oxygen. Recent work by Furukawa and Ogryzlo (10) indicates that the correlation is more involved than had been previously assumed. From a number of measurements using "ambient air" from outside the laboratory as a quencher of O2(lIg), a rate constant of (3.4 f 0.4) X cm3 moleculecl sec-I was obtained. The partial pressure of water in the atmosphere was calculated from the relative humidity and the temperature to be 18 f 2 torr a t the time of the experiments. To evaluate the influence of humidity on the rate constant, pure dry synthetic air was also used as a quencher, and a rate constant of (3.0 f 0.1) X cm3 molecule- 1 sec-1 was obtained. The difference between the dry and the wet air values is less than would be expected from published data on the quenching of singlet oxygen by water vapor (2, 12). The dry air value is consistent, within the error limits, with a rate constant of (4.4
f 1.3) x 10-19 cm3 molecule-1 sec-1 given by Clark and cm3 moleWayne (13) and with an estimate of 3 X cule-1 sec-1 by Evans et al. (14) to fit measured atmospheric altitude profiles of O Z ( ~ & ) .It also agrees fairly well with our published rate constant of (1.7 f 0.1) X 10-18 cm3 molecule-1 sec-1 for pure oxygen, since nitrogen does not contribute appreciably to the quenching of singlet oxygen in the atmosphere (2). A reinvestigation of the important rate constant k ( O z ) , which is not compatible with data of other investigators (12, 13), will be published separately. The reactivity toward 0 2 ( l A g ) of all pollutant molecules listed in Table I is comparable to or lower than that of pure ambient air. Reactions of these contaminant species with atmospheric Oz(lAg) are therefore unimportant. Because these reactions would probably lead to products of a highly oxidizing nature, it is not inconceivable that they may be precursors of or be themselves substances hazardous to health. Up to the present, however, it has not been investigated what products, if any, are produced from the interaction of O 2 ( l A a ) with these molecules and thus no categoric answer as to the toxicological role of 0~(~1,) in air pollution can as yet be given. N o t e A d d e d in Proof: A recent publication by Huie and Herron (15) describing results concerning reactions of 0 2 ( 1 A , ) with a wide variety of organic compounds, arrives a t essentially the same conclusion concerning the role of excited 0 2 in atmospheric processes.
Literature Cited (1) Pitts, J. N., in “Chemical Reactions in Urban Atmospheres,”
C. S. Tuesday, Ed., Elsevier, 1971. (2) Becker, K . H., Groth, W., Schurath, U., Chem. Phys. Lett., 8,259 (1971). (3) Groth, W., Becker, K . H., Comsa, G. H., Elzer, A,, Fink, E . H., J u d , W., Kley, D., Schurath, U., Thran, D., Naturwissenschaften, 59,379 (1972). (4) Elias, L., Ogryzlo, E . A., Schiff, H. I., Can. J. Chem., 37, 1680 (1959). (5) Sidebottom, H. W., Badcock, C. C., Jackson, G. E., Calvert, J . G., Reinhardt, G. W., Damon, G. K., Enuiron. Sci. Technol., 6.72 (1972). (6) ‘Leighton, P . A,, “Photochemistry of Air Pollution,” Chap. IX, Academic Press, New York, N.Y., 1961. (7) Jones, T. N., Bayes, K. D., Chen,. Phys. Lett., 11, 163 (1971). (8) Frankiewicz, T. C., Berry, R. S., J. Chem. Phys., 58, 1787 (1973). (9) Ackerman, R. A,, Rosenthal, I., Pitts, J . N., ibid., 54, 4960 (1971). (10) Furukawa, K., Ogryzlo, E. A.,J. Photochem., 1,163 (1973). (11) Ogryzlo, E . A,, Tang, C. W., J . Amer. Chem. SOC.,92, 5034 (1970). (12) Findlay, F. D., Snelling, R. D., J. Chem. Phys., 55, 545 (1971). (13) Clark, I. D., Wayne, R. P., Proc. Roy. SOC.(London), A314, 1680 (1969). (141 Evans. W.F . J . , Hunten. D. M.. Llewellvn. ” . E . J.. VallanceJones, A., J. Geophys. Res., 73,2885 (1968). (15) Huie, R. E., Herron, J . T., Int. J. Chem. Kinet., 5, 197 (1973). Received for review Nouember 12, 1973. Accepted M a y 10, 1974. Work supported by the Bundesministerium des Innern.
Photochemical Reactivities of Aldehyde-Nitrogen Oxide Systems Stanley L. Kopczynski, Aubrey P. Altshuller,* and Francis D. Sutterfield Chemistry and Physics Laboratory, Environmental Protection Agency, Research Triangle Park, N.C.
The photooxidation of formaldehyde, acetaldehyde, and propionaldehyde in the presence of nitrogen oxides produces the same products and biological effects as do hydrocarbon photooxidations. Despite the greater consumption of formaldehyde, propionaldehyde is generally the most reactive in terms of higher product yields, eye irritation, and plant damage. Product yields from the photooxidation of formaldehyde-nitrogen oxide tends to be the most sensitive to variation in the ratio of reactants. As a result the oxidant yields from the reactions in the formaldehyde-nitrogen oxide system exceed those of the other systems a t higher aldehyde-nitrogen oxide ratios. The photochemical reactivities of the three aldehydes in terms of product yields, eye irritation, and plant damage overlap those of the olefinic and aromatic hydrocarbons. The reduction of formaldehyde and higher molecular weight aldehydes should be accomplished along with reduction of hydrocarbons in emission control strategies for photochemically reactive substances.
Aldehydes have been shown to be formed as the result of incomplete combustion in internal combustion engines ( 1 - 4 , diesel engines ( I , 5, 6), and incinerators (3, 7). Aldehydes also are major products of the laboratory photooxidation of olefins and alkylbenzenes in the presence of nitrogen oxides, all in the ppm range in air ( 8 ) . Atmo-
spheric analyses confirm that aldehydes exist a t significant concentrations in the Los Angeles atmosphere as a result of combustion and atmospheric reactions (9-1 I ) . When photooxidized in air with radiation between 2900 and 3500 A, formaldehyde and other aliphatic aldehydes in the absence of nitrogen oxides react to form products identified as hydrogen peroxide, alkylhydroperoxides, carbon monoxide, and lower-molecular-weight aldehydes, (12, 13), and the intermediates formed can react with olefins and alkylbenzenes (14). When photooxidized in air in the presence of nitrogen oxides, aliphatic aldehydes react rapidly to form oxidants including ozone, hydrogen peroxide, and peroxyacyl nitrates (15-17). These products cause damage to plants (18,19). Saturated aliphatic aldehydes, when irradiated with nitrogen oxides, also cause eye irritation (15, 18). The past studies have not involved a systematic investigation of the photooxidation of aldehydes in the presence of nitrogen oxides as a function of aldehyde or nitrogen oxide concentration. In the present work reaction products, eye irritation, and plant damage from the photooxidation of formaldehyde, acetaldehyde, or propionaldehyde with nitrogen oxide have been measured as the concentrations and ratios of reactants are varied. These results are used to achieve three objectives: develop the relationships between aldehyde and nitrogen oxide as reactants with the products or effects measured, compare the photochemical reactivity parameters for the aldehyde systems with each other and with those obtained preVolume 8, Number 10, October 1974
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