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Atmospheric oxidation of hydrocarbons: Formation of hydroperoxides and peroxyacids
00044981/8313.00 + 0.00
Enuironmenr Vol. 17, No. I I. pp.2259-2265.1983 Printed in GreatBritain.
Atmospheric
Pergamon Press Ltd.
ATMOSPHERIC OXIDATION OF HYDROCARBONS: FORMATION OF HYDROPEROXIDES AND PEROXYACIDS PHILIP L. HANST
and
BRUCE W.
GAY,
JR.
Environmental Sciences Research Laboratory, U.S. Environmental Protection Agency, Research Triangle
Park, NC 27711, U.S.A. (First received 6 December 1982 and infinalform 11 April 1983) Abstract-Hydrocarbons at ppm levels in air have been oxidized in the absence of nitrogen oxides. Chlorine atoms served as initiators of the oxidations. Infrared analysis showed alkyl hydroperoxides to be formed early in the oxidation sequences. Aldehydes and ketones were also formed, followed by the appearance of peroxyacids. Peroxyacetic acid was found to be an especially stable reaction product. Laboratory rate data and recent atmospheric measurements of NO, NOz and HO2 indicate that hydroperoxides and peroxyacids are also formed in the real atmosphere.
INTRODUCTION
Atmospheric oxidation of organic compounds proceeds in steps that include hydrogen abstraction from the stable molecules, oxygen addition to the resulting radicals and further reactions of the oxygenated radicals. The hydrogen abstractions are the slow steps that control the rate of the overall oxidation process. The principal abstracting species is the hydroxyl radical. RH+OH-+R+HzO. (1)
atmosphere have apparently not been carried out because of the lack of measurement methods. In this work a qualitative laboratory demonstration of the hydrogen abstractions of (5) and (6) has been carried out. It has also been shown that the degradation of many different molecules yields peroxyacetic acid which is then slow in being oxidized further. The formation of hydroperoxides and peroxy acids may be relevant to the problems of photochemical air pollution, acid rain and stratospheric 0, depletion.
The abstractions are followed quickly by the formation of peroxide radicals. R+O,
-+RO,.
EXPERIMENTAL TECHNIQUE
(2)
A peroxide radical will react in one of several ways, depending on the concentrations of other trace gases. Radical reduction by NO is very fast. ROO + NO + RO + NO,. There can also be combination peroxynitrates.
with NO,
ROO + NO, -+ ROONO*.
(3) to yield (4)
In the urban smog, where the nitrogen oxides are present in high concentrations, reactions (3) and (4) proceed rapidly. In the clean atmosphere, however, the concentrations of NO and NOz are lower than in the smog by a factor of possibly 1000 (Kley et al., 1981) and the fate of the peroxide radicals there is not so well established. The most likely reaction of peroxide radicals in clean tropospheric air appears to be abstraction of hydrogen from another molecule or radical, forming hydroperoxide or peroxy acid. ROO+RH-*ROOH+R
(5)
or ROO + HO0 + ROOH + 0,. Measurements AE17:11-I
(6)
of organic peroxides in the non-urban
In the laboratory it is convenient to use Cl atoms to initiate the oxidation of atmospheric trace gases. The course of a reaction can then be followed either in the presence or absence of the NO,. Using this method with NO, present, reaction (4) was previously studied in detail. Many types of peroxyacyl nitrates and peroxyalkyl nitrates were formed (Gay et al., 1976;Edney et al., 1979).In this work the Cl initiation method has been used with NO, absent in order to look for evidence of reactions (5) and (6). The reactions were conducted at ppm concentrations in dry air. Ultraviolet light was used to release Cl atoms from Cl gas, and long-path infrared absorption spectroscopy was used to monitor reactants and products. The apparatus and technique have been described in more detail in the earlier papers. In the present work a typical procedure was as follows: (1) Dry, hydrocarbon-free air was placed in the 700-e reaction vessel at a pressure of 0.13 atm. Using a 360-m folded light path, the blank spectrum was recorded and stored. For apparatus description see Edney et al. (1979). (2) Hydrocarbon vapor and molecular Cl gas were added to give partial pressures of 10e5 and 5 x lo-’ atm respectively. The spectrum of this starting mixture was recorded and plotted, using the cell blank as reference. (3) The ultraviolet lamps were then turned on for a short time-typically 1 min-in order to release a number of Cl atoms equal to only a fraction of the starting number of hydrocarbon (HC) molecules. The spectrum was again recorded. (4) Additional short periods of Cl atom release were then carried out, with the spectrum being recorded after each. This was carried on until the HC was all consumed, as
2259
PHILIP
2260
L. HANSTand BRUCEW. GAY, JR.
indicated by the disappearance of the C-H stretch band in the infrared spectrum. (5) From the collection of spectra, the con~ntrations of reactants and products were calculated and the progress of the reaction was plotted. The absorption band used for measurement of the hydroperoxides falls at 36OOcm-‘. The intensity of this band in a published reference spectrum of methyl hydroperoxide enabled estimates to be made of the hydroperoxide yields. (Hanst and Calvert, 1959). For m~surement of the peroxy acids the characteristic band near 33OOcm-’ was used. Spectra of peroxyformic and peroxyacetic acids were published many years ago by Giguere and Olmos (1956). Spectra of peroxypropionic and peroxybutyric acids were published by Stephens et al. (1957). More recently a spectrum of peroxyformic acid was published by Su er al. (1979). A new spectrum of peroxyacetic acid was also recorded during the present study in order to see the exact shape of the 3300-cm-’ band for di~~mination between peroxyacetic and peroxypropionic acids.
EXPERIMENTAL RESULTS
The following compounds were oxidized: methane, ethane, propane, n-butane, n-hexane, toluene, acetaldehyde, propionaldehyde, acetone and diethyl ketone. All except methane produced peroxyacids. The non-methane hydrocarbons produced detectable amounts of hydroperoxides, but the carbonyl compounds did not. Methane
Methane oxidation did not yield detectable amounts of organic peroxides, even though it is expected that methyl hydroperoxide was formed during the oxidation. 1.5 x lo- 5 atm of methane, 5 x lo-’ atm of Cl2 and 0.13 atm of dry air were placed in the reaction
i
~-
r---
7 -7.---
vessel and irradiated for a total of 18min. This irradiation period was long enough to release about 95 % of the Cl from the molecular state. The Cl atom attack reduced the methane concentration by about two-thirds, with about 90% of the reacted methane being converted to CO and 10% to CO,. Small amounts of formic acid and H,O, appeared, but at the end of the period of irradiation there was no detectable methyl hydro~roxide, formaldehyde or peroxyformic acid. This is a plausible result because methane reacts only very slowly with Cl atoms, while the products react fast. For example, the published rate constants show that at 300 K the reaction between H&O and Cl is 700 times faster than the reaction between CH, and Cl (H~pson, 1980). Thus the intermediate products were destroyed shortly after being formed. Peroxyformic acid was not expected as a product because if methane or formaldehyde is oxidized in the presence of NO,, peroxyformyl nitrate is never detected (Edney et al., 1979). It appears that peroxyformyi radicals either cannot be formed or, if they are formed, they decompose extremely fast to CO and HO,. Ethane
Ethane proved to be much better than methane as a test compound to exhibit various reaction products. Hydroperoxide, acetaldehyde, peroxyacetic acid, CO and CO, were followed as the oxidation progressed. Figure 1 shows a set of spectra obtained during ethane oxidation. The infrared frequencies chosen for monitoring reactants and products and the methods used for making concentration estimates are listed in Table 1. The calculated concentrations are listed in Table 2, ,-
I
-! j j I
loo-
/
100-j
OLLIL 1600
.-_L 2000
_/ 2400
-I
L-1 2800
3200
Fig. 1. Oxidation of ethane. 1 x 10ms atm C,H,; 6 x foe5 atm C1,$.13 atm air; path = 360 m.
3600
Formation of hydroperoxides and peroxyacids
2261
Table 1. Ethane study Frequency of band (cm-‘)
Method of estimating concentration
CA
2970
C,H, metered into cell volumetrically. The starting spectrum established the absorption coefficient on the side of the
CH,CH,OOH
3600
an absorption coefficient was arbitrarily assigned to account for all the starting material at the 2 min mark. The value was 5 + lScm_’ atm-‘. This is close to the value estimated for CH,OOH from the spectrum published by Hanst and Calvert (1959)
CH,CHO
2700
an absorption coefficient value of 3.0 * OScm-’ atm-’ was calculated from the spectrum published by Hanst (1971)
CH,C(O)OOH
3300
an absorption coefficient value of 16 &- Scm-‘atm-’ was measured on a reference sample of the acid
CO2
2350
an absorption coefficient was calculated by assuming that 457; of the carbon atoms were oxidized to CO,, as reported by Hanst et al. (1980)
co
2120
an absorption coefficient was calculated by assuming that 40% of the carbon atoms were oxidized to CO, as reported by Hanst et al. (1980)
Compound
band at 3040cm-’
Table 2. Ethane
Time (min) 0 1 2 4 8 16
CH,CH,OOH
C,H,
12 6.7 5.0 2.2 0.4
&l & 0.7 + 0.5 * 0.4 f 0.2 0
and based on these numbers oxidation is plotted in Fig. 2.
0.8 1.7 2.4 1.2 0.4
0 k f k + +
0.3 0.5 0.8 0.4 0.2
the progress
results
Concentrations CH,CHO
0 3.1 * 0.5 3.5 + 0.6 2.0 * 0.4 0 0
of the
n-Butane The oxidation of n-butane yielded hydroperoxides, peroxyacid, other carbonyl compounds, CO and COz, as plotted in Fig. 3. The shape of the peroxyacid band at 33OOcm- ’ identifies the acid as peroxyacetic acid. The yield of observed products was much smaller than the yield in the ethane case. Only about half of the C atoms appeared as CO and CO, at the end of the oxidation, even though the gaseous organic matter was nearly all gone. In the ethane case, about 80 % of the C appeared as CO and C02. These yields are similar to those reported in an earlier paper (Hanst et al., 1980). It is presumed that the C not seen as gaseous oxidation products was condensed into fine particles or de-
(in units of 10m6 atm) CH,C(O)OOH CO,
0 0 0.08 f 0.3 f 0.6 + 0.5 *
0.03 0.1 0.2 0.2
1.0 2.0 7 10 11
0 * 0.2 + 0.4 k 1.5 + 2.0 k 2.0
CO
0.7 1.5 6 9 8
0 + f f * +
0.2 0.3 1.0 1.5 1.5
posited on the cell walls. The previous paper pointed out that the higher the molecular weight of the starting HC, the larger the fraction of the C atoms that did not appear in gaseous products. Propane, acetone and propionaldehyde Oxidation of each of these 3-C molecules yielded peroxyacetic acid equivalent to a few per cent of the starting material. Hydroperoxide was produced from propane but not from the two carbonyl compounds. The shape of the 3300-cm-’ band showed that in the case of acetone, peroxyacetic acid was formed from the beginning of the reaction period. Propionaldehyde initially yielded peroxypropionic acid, but as the reaction progressed this acid was transformed to peroxyacetic acid. In the propane case, the acid was not clearly defined at the beginning. No doubt it was a
2262
PHILIP L. HANST and BRUCE
0
a
2
6
8
MINUTES
OF IRRADIATION
W.
GAY,
JR.
12
10
16
Fig. 2. Oxidation of ethane. 6 x lo- 5 atm Cl,; 0.13atm air; U.V.light.
8
BENZYL HYDROPEROXIDE
a
1
2
3
4
5
6
7
6 3
MINUTES OF IRRADIATION 4
Fig. 3. Oxidation of n-butane. 1 x 10e4 atm Cl,; 0.13 atm air; U.V.light.
6 MINUTES
OF IRRADIATION
Fig. 4. Oxidation of toluene. 5 x 10T5 atm Cl,; 0.13atm air; U.V.light.
mixture, changing to peroxyacetic acid as the reaction progressed. Toluene The oxidation of 1 x 10m5atm of toluene by 5 x 10T5atm Cl, in 0.13atm of air produced benzyl
hydroperoxide, benzaldehyde, peroxybenzoic acid, CO and CO,. Other carbonyl compounds were formed but not identified. Only about 50% of the C atoms of the toiuene were seen as gaseous reaction
products. The benzyl hydroperoxide yield was high but the peroxybenzoic acid yield was low, as shown in Fig. 4. Other compounds Acetaldehyde oxidation readily yielded peroxyacetic acid, CO and COZ. As in the other cases, continued Cl atom attack on the mixture slowly removed the acid and slowly converted the CO to COz.
2263
Formation of hydroperoxides and peroxyacids Oxidation of 3-pentanone also yielded peroxyacetic acid as a stable end product.
became acetaldehyde molecules:
Stability of peroxyacetic acid
As the reaction progressed, the HO, radicals and acetaldehyde molecules built up as hydrogen donors, promoting the generation of ethyl hydroperoxide:
Whenever peroxyacetic acid was formed, it proved to be stable under attack by Cl atoms. The loss rate for the acid under Cl attack seemed no greater than the rate at which the cell walls absorb other acids such as formic, nitric and hydrochloric acids. To demonstrate the slowness of the peroxyacidC1 reaction, an experiment was carried out in which ethane was oxidized in the presence of molecular hydrogen. The mixture tested was lo-‘atm of ethane, low4 atm of Cl,, lo- 3 atm of H, and 0.13 atm of dry air. Published data show the reaction between Cl and C2H, at room temperature to be more than 3000 times faster than the reaction between Cl and H, (Hampson, 1980). Thus, even though H, was in lOO-fold excess over CzH6, ethane consumption was almost as fast as it would have been if H, were absent. Subjecting the gas mixture to 30s of U.V.radiation reduced the ethane concentration to about one-third of its starting value. After 90 s of irradiation the C,H, appeared to be about 95% reacted. At that point the gaseous mixture contained about 3 x lo-’ atm ofperoxyacetic acid. An additional 90s of irradiation then reduced the peroxyacetic concentration by about 10% to 2.7 x lo- ’ atm. The increase of HCl during this second irradiation period showed that the chlorine atoms were being produced at almost as great a rate as at the beginning of the reaction. Thus the Cl atoms were reacting mainly with Hz while the rate of Cl attack on the peroxyacetic acid was at least 10 times slower than the rate of Cl attack on the C,H, that had been present at the beginning.
OXIDATION
MECHANISMS
The oxidation of C2H, is of special interest because a relatively large amount of this 2-C molecule is present in the clean tropospheric air. After methane, ethane appears to be the next most abundant atmospheric HC, with a world mixing ratio averaging near 1O-9 (Singh and Salas, 1982). The C,H, degradation exemplifies the types of reactions that occur for all HCs. The C2H6 results may be accounted for in the following sequence: Cl atoms abstracted hydrogen: Cl + C2H, + HCl + C2H,. The ethyl radical added 0,: C2H, + 0, + CH,CH,OO. At the beginning, these radicals reacted with themselves, yielding alkoxy radicals and 0,: 2 CH3CH200 The alkoxy
radicals
+ 2 CH3CH,0 yielded
+ OZ.
up hydrogen
and
CH,CH,O
+ 0, + CH3CH0 + HO,.
HO* + CH,CH,OO
+ 0, + CH,CH,OOH
CH3CH0 + CH,CH,OO
+ CH,CO + CH,CH,OOH.
Acetaldehyde was directly attacked by the Cl atoms: CH,CHO + Cl + CH,CO + HCl. The acetyl radicals added 0,: CH,CO + 0, + CH,C(O)OO. Some of the resulting peroxyacetyl radicals gained a hydrogen to become peroxyacetic acid: CH,C(O)OO
+ R-H + CH,C(O)OOH
+ R.
Some lost oxygen to become acetoxy radicals which then split out CO,: 2 CH,C(O)OO
+ 2CH,C(O)O + 0,
CH,C(O)O -+ CH, + COZ. The methyl groups were oxidized to CO: CH, + 0, 2 CH,OO CH,O + O2 H&O + Cl HCO + 0,
+ CH,OO + 2CH,O + O2 -+ H&O + HO, + HCl + HCO + HO, + CO.
Meanwhile the hydroperoxide molecules were also being attacked by Cl and oxidized to CO, and CO. The reaction sequence in that case probably involves reactions similar to those described above. The peroxyacetic acid molecules apparently were not further oxidized in these laboratory studies but eventually were taken up on the vessel walls. For the heavier HC molecules the oxidation processes are longer and more complex. The same types of hydrogen abstraction and oxygen addition processes take place in all cases, but when the starting molecules are larger, complex intermediate molecules of low vapor pressure can form and condense into fine particles. In this work it was seen that n-butane and toluene only had about half of their C atoms appear as gaseous CO and CO,. In an earlier study it was seen that when alpha-pinene (C10H16) was oxidized, only 30% of the C atoms appeared as CO and COZ. In all these cases the spectra indicated the presence of carbonyl compounds that were resistant to further oxidation. It is suspected that these compounds existed in fine particle form. In the oxidation of the carbonyl compounds, the results were entirely in line with the mechanism steps given above. Acetaldehyde, for example, yielded peroxyacetic acid, which remained stable under Cl attack.
PHILIPL. HANST and BRUCE W.
2264
Table 3. Rate constants
GAY, JR.
for reactions of peroxide radicals at normal atmospheric
temperature and pressure Reaction
Rate constant (cm3mall 1set-‘)
CH,OO + NO + CH30 + NO,
8 x 10-r* 7.5 x lo- I2
HO0 + NO, 5 HOONO,
6 x10-”
CH,OO + NO, y CH,00N02 HO0 + HO0 -+ H,O, + 0,
1.6 x lo-l2 2.5 x IO-‘* 6 x lo-l2
HOO+NO+HO+NO,
CH,OO + HO0 + CH,OOH + O2
Acetone also yielded peroxyacetic acid. Presumably the removal of hydrogen from acetone leads first to a di-carbonyl compound which then either photolyzes or is attacked by Cl atoms, producing acetyl radicals. These radicals then are oxidized to peroxyacetic acid. There could have been transitory formation of a hydroperoxide from acetone, but this was not detected. In the propionaldehyde case, peroxypropionic acid was formed initially, but under further oxidation it gave way to peroxyacetic acid. Again some transitory hydroperoxide seems possible, but was not detected. 3Pentanone also yielded peroxyacetic acid, which undoubtedly was preceded by various other oxygenated molecules, including di-carbonyls and hydroperoxides.
DISCUSSION
A peroxide radical in the atmosphere will most likely react with NO, NO2 or HO*. Rate constants published for the reactions of hydroperoxide and methyl peroxide radicals are listed in Table 3 (Hampson, 1980). These reaction rates are all of similar magnitude (fast), indicating that the product yields in various regions of the atmosphere will depend mainly on the reactant concentrations. In urban areas, in the morning NO is the dominant reactant, while in the afternoon the NO2 reactions take over. This has been known for many years, but was re-affirmed recently through spectroscopic measurements made in Los Angeles (Hanst et al., 1982). The NO, reactions will occur in the stratosphere, where the mixing ratios of NO and NOz are on the order of l&100 times higher than the mixing ratio of HO*. Stratospheric HO, measurements are reported by Anderson et al. (1981) while NO and NO, measurements have been summarized by the National Academy of Sciences (1976). In the troposphere, outside of urban areas, the reactions with HO, radicals will be more important than in the stratosphere because there the mixing ratios of NO and NO, are much smaller than in the stratosphere. McFarland et nl. (1979) have reported NO mixing ratios in near-surface air over the Pacific ocean to range from about 1 x lo- “to4x lo-“.Thisisabout IO4 times smaller than a typical NO mixing ratio in urban smog. Kley et al. (1981) have reported mixing ratios for NO, (NO + NO,) in the clean lower tropos-
phere on the order of 5 x 10-i’. This is some lo3 times lower than in urban smog. Mixing ratios of HO, in the troposphere have apparently not been measured, but they have been estimated from model calculations of the tropospheric chemistry. Logan et al. (1981) calculated noon-time tropospheric HO, mixing ratios in the range 10-io -lo- ‘i. Stedman and Cantrell (1982) also mention tropospheric HO, mixing ratios on the order of 10-16. McFarland et al. (1979) have already pointed out that when the NO mixing ratio is in the 10-“-10-‘2 range a substantial portion of the peroxide radicals will react by abstracting a hydrogen from HO, rather than by oxidizing the NO. If the NO, mixing ratio is also as low as or lower than the HO, mixing ratio, then it appears that HOz becomes the principal reaction partner for peroxy radicals formed in the troposphere. Thus there appears to be good reason to believe that there are high yields of hydroperoxides and peroxyacids in the chemistry of the real atmosphere, just as there were high yields in the laboratory tests reported here. Peroxyacetic acid ought to be the principal peroxide reaction product that accumulates. When the peroxy radicals react with NO,, peroxyacetyl nitrate is the principal product (Singh and Hanst, 1981; Singh and Salas, 1983). This combination, however, is a reversible reaction (Hendry and Kenley, 1977). HO, radicals will therefore slowly convert peroxyacetyl nitrate to peroxyacetic acid. The nitrates and acids will also be subject to degradation in the atmosphere by photolysis and by attack of OH radicals. The rates of these degradations should be measured in the laboratory. It is hoped that ways will be found to actually detect and measure the hydroperoxides and peroxyacids in the air.
REFERENCES
Anderson J. G., Grass1 H. J., Shetter R. E. and Margitan J. J. (1981) HO, in the stratosphere: three in-situ observations. &ophys. l&s. Left. 8, 289-292. Ednev E. 0.. Snence J. W. and Hanst P. L. (1979) Peroxv nitrate air’pdllutants: synthesis and thermal stability. In Nitrogeneous Air Pollutants (edited by Grosjean D.), pp, 111-135. Ann Arbor Science Publishers, Ann Arbor, Ml. Gay B. W., Jr., Noonan R. C., Bufahni J. J. and Hanst P. L.
Formation of hydroperoxides and peroxyacids (1976) Photochemical synthesis of peroxy acyl nitrates in the gas phase via chlorine-aldehyde reactions. Envir. Sci. Technol. 10, 82-85. Giguere P. A. and Olmos A. W. (1956) A spectroscopic study of hydrogen bonding in performic and peracetic acids. Can. J. Chem. 30, 821-830. Hampson R. F. (1980) Chemical Kinetics and Photochemical Data Sheets for Atmospheric Reactions. Report NO. FAAEE-80-17, U.S. Department of Transportation, available through the National Technical Information Service, Springfield, VA. Hanst P. L. (1971) Spectroscopic methods for air pollution measurement. In Advances in Environmental Science and Technology (edited by Pitts J. N., Jr. and Metcalf R. L.), Vol. 2, pp. 91-213. Wiley-Interscience, New York. Hanst P. L. and Calvert J. G. (1959) The oxidation of methyl radicals at room temperature. J. phys. Chem. 63, 71-77. Hanst P. L., Spence J. W. and Edney E. 0. (1980) Carbon monoxide production in photooxidation of organic molecules in air. Atmospheric Enoironment 14, 1077-1088. Hanst P. L., Wong N. G. and Bragin J. (1982) A long-path infrared study of Los Angeles smog. Atmospheric Environment 16, 969-981.
Hendry D. G. and Kenley R. A. (1977) Generation of peroxy radicals from peroxy nitrates (RO,NO,): decomposition of peroxyacyl nitrates. J. Am. Chem. Sot. 99, 3198-3199. Kley D., Drummond J. W., McFarland M. and Liu S. C. (198 I) Tropospheric profiles of NO,. J. geophys. Res. 86, 315333161.
2265
Logan J. A., Prather M. J., Wofsey S. C. and McElroy M. B. (1981) Tropospheric chemistry: a global perspective. J. geophys. Res. 86, 721s-7254. McFarland D., Kley D., Drummond J. W., Schmellekopf A. L. and Winkler R. H. (1979) Nitric oxide measurements in the equatorial Pacific region. Geophys. Res. Left. 6, 605608. National Academy of Sciences (1976) Halocarbons: Eficts on Stratospheric Ozone. Printing and Publishing Office, National Academy of Sciences, Washington, DC. Singh H. B. and Hanst P. L. (1981) Peroxyacetyl nitrate (PAN) in the unpolluted atmosphere: an important reservoir for nitrogen oxides. Geophys. Res. Left. 8,941-944. Singh H. B. and Salas L. J. (1982) Measurement of selected light hydrocarbons over the Pacific ocean. Latitudinal and seasonal variations. Geophys. Res. Left. 9, 842-845. Singh H. B. and Salas L. J. (1983) Peroxyacctyl nitrate in the free troposphere. Nature 302, 326328. Stedman D. H. and Cantrell C. A. (1982) Laboratory studies of an HOr/RO, detector. Proceedings of 2nd Symposium on the Composition of the Nonurban Troposphere. American Meteorological Society, Boston, MA. Stephens E. R., Hanst P. L. and Doerr R. C. (1957) Infrared spectra of aliphatic peroxyacids. Analyt. Chem. 29, 776777. Su F., Calvert J. G., Shaw J. H., Niki H., Maker P. D., Savage C. M. and Breitenbach L. D. (1979) Spectroscopic and kinetic studies of a new metastable species in the photooxidation of gaseous formaldehyde. Chem. Phys. Lett. 65, 221-225.
Enuironmenr Vol. 17, No. I I. pp.2259-2265.1983 Printed in GreatBritain.
Atmospheric
Pergamon Press Ltd.
ATMOSPHERIC OXIDATION OF HYDROCARBONS: FORMATION OF HYDROPEROXIDES AND PEROXYACIDS PHILIP L. HANST
and
BRUCE W.
GAY,
JR.
Environmental Sciences Research Laboratory, U.S. Environmental Protection Agency, Research Triangle
Park, NC 27711, U.S.A. (First received 6 December 1982 and infinalform 11 April 1983) Abstract-Hydrocarbons at ppm levels in air have been oxidized in the absence of nitrogen oxides. Chlorine atoms served as initiators of the oxidations. Infrared analysis showed alkyl hydroperoxides to be formed early in the oxidation sequences. Aldehydes and ketones were also formed, followed by the appearance of peroxyacids. Peroxyacetic acid was found to be an especially stable reaction product. Laboratory rate data and recent atmospheric measurements of NO, NOz and HO2 indicate that hydroperoxides and peroxyacids are also formed in the real atmosphere.
INTRODUCTION
Atmospheric oxidation of organic compounds proceeds in steps that include hydrogen abstraction from the stable molecules, oxygen addition to the resulting radicals and further reactions of the oxygenated radicals. The hydrogen abstractions are the slow steps that control the rate of the overall oxidation process. The principal abstracting species is the hydroxyl radical. RH+OH-+R+HzO. (1)
atmosphere have apparently not been carried out because of the lack of measurement methods. In this work a qualitative laboratory demonstration of the hydrogen abstractions of (5) and (6) has been carried out. It has also been shown that the degradation of many different molecules yields peroxyacetic acid which is then slow in being oxidized further. The formation of hydroperoxides and peroxy acids may be relevant to the problems of photochemical air pollution, acid rain and stratospheric 0, depletion.
The abstractions are followed quickly by the formation of peroxide radicals. R+O,
-+RO,.
EXPERIMENTAL TECHNIQUE
(2)
A peroxide radical will react in one of several ways, depending on the concentrations of other trace gases. Radical reduction by NO is very fast. ROO + NO + RO + NO,. There can also be combination peroxynitrates.
with NO,
ROO + NO, -+ ROONO*.
(3) to yield (4)
In the urban smog, where the nitrogen oxides are present in high concentrations, reactions (3) and (4) proceed rapidly. In the clean atmosphere, however, the concentrations of NO and NOz are lower than in the smog by a factor of possibly 1000 (Kley et al., 1981) and the fate of the peroxide radicals there is not so well established. The most likely reaction of peroxide radicals in clean tropospheric air appears to be abstraction of hydrogen from another molecule or radical, forming hydroperoxide or peroxy acid. ROO+RH-*ROOH+R
(5)
or ROO + HO0 + ROOH + 0,. Measurements AE17:11-I
(6)
of organic peroxides in the non-urban
In the laboratory it is convenient to use Cl atoms to initiate the oxidation of atmospheric trace gases. The course of a reaction can then be followed either in the presence or absence of the NO,. Using this method with NO, present, reaction (4) was previously studied in detail. Many types of peroxyacyl nitrates and peroxyalkyl nitrates were formed (Gay et al., 1976;Edney et al., 1979).In this work the Cl initiation method has been used with NO, absent in order to look for evidence of reactions (5) and (6). The reactions were conducted at ppm concentrations in dry air. Ultraviolet light was used to release Cl atoms from Cl gas, and long-path infrared absorption spectroscopy was used to monitor reactants and products. The apparatus and technique have been described in more detail in the earlier papers. In the present work a typical procedure was as follows: (1) Dry, hydrocarbon-free air was placed in the 700-e reaction vessel at a pressure of 0.13 atm. Using a 360-m folded light path, the blank spectrum was recorded and stored. For apparatus description see Edney et al. (1979). (2) Hydrocarbon vapor and molecular Cl gas were added to give partial pressures of 10e5 and 5 x lo-’ atm respectively. The spectrum of this starting mixture was recorded and plotted, using the cell blank as reference. (3) The ultraviolet lamps were then turned on for a short time-typically 1 min-in order to release a number of Cl atoms equal to only a fraction of the starting number of hydrocarbon (HC) molecules. The spectrum was again recorded. (4) Additional short periods of Cl atom release were then carried out, with the spectrum being recorded after each. This was carried on until the HC was all consumed, as
2259
PHILIP
2260
L. HANSTand BRUCEW. GAY, JR.
indicated by the disappearance of the C-H stretch band in the infrared spectrum. (5) From the collection of spectra, the con~ntrations of reactants and products were calculated and the progress of the reaction was plotted. The absorption band used for measurement of the hydroperoxides falls at 36OOcm-‘. The intensity of this band in a published reference spectrum of methyl hydroperoxide enabled estimates to be made of the hydroperoxide yields. (Hanst and Calvert, 1959). For m~surement of the peroxy acids the characteristic band near 33OOcm-’ was used. Spectra of peroxyformic and peroxyacetic acids were published many years ago by Giguere and Olmos (1956). Spectra of peroxypropionic and peroxybutyric acids were published by Stephens et al. (1957). More recently a spectrum of peroxyformic acid was published by Su er al. (1979). A new spectrum of peroxyacetic acid was also recorded during the present study in order to see the exact shape of the 3300-cm-’ band for di~~mination between peroxyacetic and peroxypropionic acids.
EXPERIMENTAL RESULTS
The following compounds were oxidized: methane, ethane, propane, n-butane, n-hexane, toluene, acetaldehyde, propionaldehyde, acetone and diethyl ketone. All except methane produced peroxyacids. The non-methane hydrocarbons produced detectable amounts of hydroperoxides, but the carbonyl compounds did not. Methane
Methane oxidation did not yield detectable amounts of organic peroxides, even though it is expected that methyl hydroperoxide was formed during the oxidation. 1.5 x lo- 5 atm of methane, 5 x lo-’ atm of Cl2 and 0.13 atm of dry air were placed in the reaction
i
~-
r---
7 -7.---
vessel and irradiated for a total of 18min. This irradiation period was long enough to release about 95 % of the Cl from the molecular state. The Cl atom attack reduced the methane concentration by about two-thirds, with about 90% of the reacted methane being converted to CO and 10% to CO,. Small amounts of formic acid and H,O, appeared, but at the end of the period of irradiation there was no detectable methyl hydro~roxide, formaldehyde or peroxyformic acid. This is a plausible result because methane reacts only very slowly with Cl atoms, while the products react fast. For example, the published rate constants show that at 300 K the reaction between H&O and Cl is 700 times faster than the reaction between CH, and Cl (H~pson, 1980). Thus the intermediate products were destroyed shortly after being formed. Peroxyformic acid was not expected as a product because if methane or formaldehyde is oxidized in the presence of NO,, peroxyformyl nitrate is never detected (Edney et al., 1979). It appears that peroxyformyi radicals either cannot be formed or, if they are formed, they decompose extremely fast to CO and HO,. Ethane
Ethane proved to be much better than methane as a test compound to exhibit various reaction products. Hydroperoxide, acetaldehyde, peroxyacetic acid, CO and CO, were followed as the oxidation progressed. Figure 1 shows a set of spectra obtained during ethane oxidation. The infrared frequencies chosen for monitoring reactants and products and the methods used for making concentration estimates are listed in Table 1. The calculated concentrations are listed in Table 2, ,-
I
-! j j I
loo-
/
100-j
OLLIL 1600
.-_L 2000
_/ 2400
-I
L-1 2800
3200
Fig. 1. Oxidation of ethane. 1 x 10ms atm C,H,; 6 x foe5 atm C1,$.13 atm air; path = 360 m.
3600
Formation of hydroperoxides and peroxyacids
2261
Table 1. Ethane study Frequency of band (cm-‘)
Method of estimating concentration
CA
2970
C,H, metered into cell volumetrically. The starting spectrum established the absorption coefficient on the side of the
CH,CH,OOH
3600
an absorption coefficient was arbitrarily assigned to account for all the starting material at the 2 min mark. The value was 5 + lScm_’ atm-‘. This is close to the value estimated for CH,OOH from the spectrum published by Hanst and Calvert (1959)
CH,CHO
2700
an absorption coefficient value of 3.0 * OScm-’ atm-’ was calculated from the spectrum published by Hanst (1971)
CH,C(O)OOH
3300
an absorption coefficient value of 16 &- Scm-‘atm-’ was measured on a reference sample of the acid
CO2
2350
an absorption coefficient was calculated by assuming that 457; of the carbon atoms were oxidized to CO,, as reported by Hanst et al. (1980)
co
2120
an absorption coefficient was calculated by assuming that 40% of the carbon atoms were oxidized to CO, as reported by Hanst et al. (1980)
Compound
band at 3040cm-’
Table 2. Ethane
Time (min) 0 1 2 4 8 16
CH,CH,OOH
C,H,
12 6.7 5.0 2.2 0.4
&l & 0.7 + 0.5 * 0.4 f 0.2 0
and based on these numbers oxidation is plotted in Fig. 2.
0.8 1.7 2.4 1.2 0.4
0 k f k + +
0.3 0.5 0.8 0.4 0.2
the progress
results
Concentrations CH,CHO
0 3.1 * 0.5 3.5 + 0.6 2.0 * 0.4 0 0
of the
n-Butane The oxidation of n-butane yielded hydroperoxides, peroxyacid, other carbonyl compounds, CO and COz, as plotted in Fig. 3. The shape of the peroxyacid band at 33OOcm- ’ identifies the acid as peroxyacetic acid. The yield of observed products was much smaller than the yield in the ethane case. Only about half of the C atoms appeared as CO and CO, at the end of the oxidation, even though the gaseous organic matter was nearly all gone. In the ethane case, about 80 % of the C appeared as CO and C02. These yields are similar to those reported in an earlier paper (Hanst et al., 1980). It is presumed that the C not seen as gaseous oxidation products was condensed into fine particles or de-
(in units of 10m6 atm) CH,C(O)OOH CO,
0 0 0.08 f 0.3 f 0.6 + 0.5 *
0.03 0.1 0.2 0.2
1.0 2.0 7 10 11
0 * 0.2 + 0.4 k 1.5 + 2.0 k 2.0
CO
0.7 1.5 6 9 8
0 + f f * +
0.2 0.3 1.0 1.5 1.5
posited on the cell walls. The previous paper pointed out that the higher the molecular weight of the starting HC, the larger the fraction of the C atoms that did not appear in gaseous products. Propane, acetone and propionaldehyde Oxidation of each of these 3-C molecules yielded peroxyacetic acid equivalent to a few per cent of the starting material. Hydroperoxide was produced from propane but not from the two carbonyl compounds. The shape of the 3300-cm-’ band showed that in the case of acetone, peroxyacetic acid was formed from the beginning of the reaction period. Propionaldehyde initially yielded peroxypropionic acid, but as the reaction progressed this acid was transformed to peroxyacetic acid. In the propane case, the acid was not clearly defined at the beginning. No doubt it was a
2262
PHILIP L. HANST and BRUCE
0
a
2
6
8
MINUTES
OF IRRADIATION
W.
GAY,
JR.
12
10
16
Fig. 2. Oxidation of ethane. 6 x lo- 5 atm Cl,; 0.13atm air; U.V.light.
8
BENZYL HYDROPEROXIDE
a
1
2
3
4
5
6
7
6 3
MINUTES OF IRRADIATION 4
Fig. 3. Oxidation of n-butane. 1 x 10e4 atm Cl,; 0.13 atm air; U.V.light.
6 MINUTES
OF IRRADIATION
Fig. 4. Oxidation of toluene. 5 x 10T5 atm Cl,; 0.13atm air; U.V.light.
mixture, changing to peroxyacetic acid as the reaction progressed. Toluene The oxidation of 1 x 10m5atm of toluene by 5 x 10T5atm Cl, in 0.13atm of air produced benzyl
hydroperoxide, benzaldehyde, peroxybenzoic acid, CO and CO,. Other carbonyl compounds were formed but not identified. Only about 50% of the C atoms of the toiuene were seen as gaseous reaction
products. The benzyl hydroperoxide yield was high but the peroxybenzoic acid yield was low, as shown in Fig. 4. Other compounds Acetaldehyde oxidation readily yielded peroxyacetic acid, CO and COZ. As in the other cases, continued Cl atom attack on the mixture slowly removed the acid and slowly converted the CO to COz.
2263
Formation of hydroperoxides and peroxyacids Oxidation of 3-pentanone also yielded peroxyacetic acid as a stable end product.
became acetaldehyde molecules:
Stability of peroxyacetic acid
As the reaction progressed, the HO, radicals and acetaldehyde molecules built up as hydrogen donors, promoting the generation of ethyl hydroperoxide:
Whenever peroxyacetic acid was formed, it proved to be stable under attack by Cl atoms. The loss rate for the acid under Cl attack seemed no greater than the rate at which the cell walls absorb other acids such as formic, nitric and hydrochloric acids. To demonstrate the slowness of the peroxyacidC1 reaction, an experiment was carried out in which ethane was oxidized in the presence of molecular hydrogen. The mixture tested was lo-‘atm of ethane, low4 atm of Cl,, lo- 3 atm of H, and 0.13 atm of dry air. Published data show the reaction between Cl and C2H, at room temperature to be more than 3000 times faster than the reaction between Cl and H, (Hampson, 1980). Thus, even though H, was in lOO-fold excess over CzH6, ethane consumption was almost as fast as it would have been if H, were absent. Subjecting the gas mixture to 30s of U.V.radiation reduced the ethane concentration to about one-third of its starting value. After 90 s of irradiation the C,H, appeared to be about 95% reacted. At that point the gaseous mixture contained about 3 x lo-’ atm ofperoxyacetic acid. An additional 90s of irradiation then reduced the peroxyacetic concentration by about 10% to 2.7 x lo- ’ atm. The increase of HCl during this second irradiation period showed that the chlorine atoms were being produced at almost as great a rate as at the beginning of the reaction. Thus the Cl atoms were reacting mainly with Hz while the rate of Cl attack on the peroxyacetic acid was at least 10 times slower than the rate of Cl attack on the C,H, that had been present at the beginning.
OXIDATION
MECHANISMS
The oxidation of C2H, is of special interest because a relatively large amount of this 2-C molecule is present in the clean tropospheric air. After methane, ethane appears to be the next most abundant atmospheric HC, with a world mixing ratio averaging near 1O-9 (Singh and Salas, 1982). The C,H, degradation exemplifies the types of reactions that occur for all HCs. The C2H6 results may be accounted for in the following sequence: Cl atoms abstracted hydrogen: Cl + C2H, + HCl + C2H,. The ethyl radical added 0,: C2H, + 0, + CH,CH,OO. At the beginning, these radicals reacted with themselves, yielding alkoxy radicals and 0,: 2 CH3CH200 The alkoxy
radicals
+ 2 CH3CH,0 yielded
+ OZ.
up hydrogen
and
CH,CH,O
+ 0, + CH3CH0 + HO,.
HO* + CH,CH,OO
+ 0, + CH,CH,OOH
CH3CH0 + CH,CH,OO
+ CH,CO + CH,CH,OOH.
Acetaldehyde was directly attacked by the Cl atoms: CH,CHO + Cl + CH,CO + HCl. The acetyl radicals added 0,: CH,CO + 0, + CH,C(O)OO. Some of the resulting peroxyacetyl radicals gained a hydrogen to become peroxyacetic acid: CH,C(O)OO
+ R-H + CH,C(O)OOH
+ R.
Some lost oxygen to become acetoxy radicals which then split out CO,: 2 CH,C(O)OO
+ 2CH,C(O)O + 0,
CH,C(O)O -+ CH, + COZ. The methyl groups were oxidized to CO: CH, + 0, 2 CH,OO CH,O + O2 H&O + Cl HCO + 0,
+ CH,OO + 2CH,O + O2 -+ H&O + HO, + HCl + HCO + HO, + CO.
Meanwhile the hydroperoxide molecules were also being attacked by Cl and oxidized to CO, and CO. The reaction sequence in that case probably involves reactions similar to those described above. The peroxyacetic acid molecules apparently were not further oxidized in these laboratory studies but eventually were taken up on the vessel walls. For the heavier HC molecules the oxidation processes are longer and more complex. The same types of hydrogen abstraction and oxygen addition processes take place in all cases, but when the starting molecules are larger, complex intermediate molecules of low vapor pressure can form and condense into fine particles. In this work it was seen that n-butane and toluene only had about half of their C atoms appear as gaseous CO and CO,. In an earlier study it was seen that when alpha-pinene (C10H16) was oxidized, only 30% of the C atoms appeared as CO and COZ. In all these cases the spectra indicated the presence of carbonyl compounds that were resistant to further oxidation. It is suspected that these compounds existed in fine particle form. In the oxidation of the carbonyl compounds, the results were entirely in line with the mechanism steps given above. Acetaldehyde, for example, yielded peroxyacetic acid, which remained stable under Cl attack.
PHILIPL. HANST and BRUCE W.
2264
Table 3. Rate constants
GAY, JR.
for reactions of peroxide radicals at normal atmospheric
temperature and pressure Reaction
Rate constant (cm3mall 1set-‘)
CH,OO + NO + CH30 + NO,
8 x 10-r* 7.5 x lo- I2
HO0 + NO, 5 HOONO,
6 x10-”
CH,OO + NO, y CH,00N02 HO0 + HO0 -+ H,O, + 0,
1.6 x lo-l2 2.5 x IO-‘* 6 x lo-l2
HOO+NO+HO+NO,
CH,OO + HO0 + CH,OOH + O2
Acetone also yielded peroxyacetic acid. Presumably the removal of hydrogen from acetone leads first to a di-carbonyl compound which then either photolyzes or is attacked by Cl atoms, producing acetyl radicals. These radicals then are oxidized to peroxyacetic acid. There could have been transitory formation of a hydroperoxide from acetone, but this was not detected. In the propionaldehyde case, peroxypropionic acid was formed initially, but under further oxidation it gave way to peroxyacetic acid. Again some transitory hydroperoxide seems possible, but was not detected. 3Pentanone also yielded peroxyacetic acid, which undoubtedly was preceded by various other oxygenated molecules, including di-carbonyls and hydroperoxides.
DISCUSSION
A peroxide radical in the atmosphere will most likely react with NO, NO2 or HO*. Rate constants published for the reactions of hydroperoxide and methyl peroxide radicals are listed in Table 3 (Hampson, 1980). These reaction rates are all of similar magnitude (fast), indicating that the product yields in various regions of the atmosphere will depend mainly on the reactant concentrations. In urban areas, in the morning NO is the dominant reactant, while in the afternoon the NO2 reactions take over. This has been known for many years, but was re-affirmed recently through spectroscopic measurements made in Los Angeles (Hanst et al., 1982). The NO, reactions will occur in the stratosphere, where the mixing ratios of NO and NOz are on the order of l&100 times higher than the mixing ratio of HO*. Stratospheric HO, measurements are reported by Anderson et al. (1981) while NO and NO, measurements have been summarized by the National Academy of Sciences (1976). In the troposphere, outside of urban areas, the reactions with HO, radicals will be more important than in the stratosphere because there the mixing ratios of NO and NO, are much smaller than in the stratosphere. McFarland et nl. (1979) have reported NO mixing ratios in near-surface air over the Pacific ocean to range from about 1 x lo- “to4x lo-“.Thisisabout IO4 times smaller than a typical NO mixing ratio in urban smog. Kley et al. (1981) have reported mixing ratios for NO, (NO + NO,) in the clean lower tropos-
phere on the order of 5 x 10-i’. This is some lo3 times lower than in urban smog. Mixing ratios of HO, in the troposphere have apparently not been measured, but they have been estimated from model calculations of the tropospheric chemistry. Logan et al. (1981) calculated noon-time tropospheric HO, mixing ratios in the range 10-io -lo- ‘i. Stedman and Cantrell (1982) also mention tropospheric HO, mixing ratios on the order of 10-16. McFarland et al. (1979) have already pointed out that when the NO mixing ratio is in the 10-“-10-‘2 range a substantial portion of the peroxide radicals will react by abstracting a hydrogen from HO, rather than by oxidizing the NO. If the NO, mixing ratio is also as low as or lower than the HO, mixing ratio, then it appears that HOz becomes the principal reaction partner for peroxy radicals formed in the troposphere. Thus there appears to be good reason to believe that there are high yields of hydroperoxides and peroxyacids in the chemistry of the real atmosphere, just as there were high yields in the laboratory tests reported here. Peroxyacetic acid ought to be the principal peroxide reaction product that accumulates. When the peroxy radicals react with NO,, peroxyacetyl nitrate is the principal product (Singh and Hanst, 1981; Singh and Salas, 1983). This combination, however, is a reversible reaction (Hendry and Kenley, 1977). HO, radicals will therefore slowly convert peroxyacetyl nitrate to peroxyacetic acid. The nitrates and acids will also be subject to degradation in the atmosphere by photolysis and by attack of OH radicals. The rates of these degradations should be measured in the laboratory. It is hoped that ways will be found to actually detect and measure the hydroperoxides and peroxyacids in the air.
REFERENCES
Anderson J. G., Grass1 H. J., Shetter R. E. and Margitan J. J. (1981) HO, in the stratosphere: three in-situ observations. &ophys. l&s. Left. 8, 289-292. Ednev E. 0.. Snence J. W. and Hanst P. L. (1979) Peroxv nitrate air’pdllutants: synthesis and thermal stability. In Nitrogeneous Air Pollutants (edited by Grosjean D.), pp, 111-135. Ann Arbor Science Publishers, Ann Arbor, Ml. Gay B. W., Jr., Noonan R. C., Bufahni J. J. and Hanst P. L.
Formation of hydroperoxides and peroxyacids (1976) Photochemical synthesis of peroxy acyl nitrates in the gas phase via chlorine-aldehyde reactions. Envir. Sci. Technol. 10, 82-85. Giguere P. A. and Olmos A. W. (1956) A spectroscopic study of hydrogen bonding in performic and peracetic acids. Can. J. Chem. 30, 821-830. Hampson R. F. (1980) Chemical Kinetics and Photochemical Data Sheets for Atmospheric Reactions. Report NO. FAAEE-80-17, U.S. Department of Transportation, available through the National Technical Information Service, Springfield, VA. Hanst P. L. (1971) Spectroscopic methods for air pollution measurement. In Advances in Environmental Science and Technology (edited by Pitts J. N., Jr. and Metcalf R. L.), Vol. 2, pp. 91-213. Wiley-Interscience, New York. Hanst P. L. and Calvert J. G. (1959) The oxidation of methyl radicals at room temperature. J. phys. Chem. 63, 71-77. Hanst P. L., Spence J. W. and Edney E. 0. (1980) Carbon monoxide production in photooxidation of organic molecules in air. Atmospheric Enoironment 14, 1077-1088. Hanst P. L., Wong N. G. and Bragin J. (1982) A long-path infrared study of Los Angeles smog. Atmospheric Environment 16, 969-981.
Hendry D. G. and Kenley R. A. (1977) Generation of peroxy radicals from peroxy nitrates (RO,NO,): decomposition of peroxyacyl nitrates. J. Am. Chem. Sot. 99, 3198-3199. Kley D., Drummond J. W., McFarland M. and Liu S. C. (198 I) Tropospheric profiles of NO,. J. geophys. Res. 86, 315333161.
2265
Logan J. A., Prather M. J., Wofsey S. C. and McElroy M. B. (1981) Tropospheric chemistry: a global perspective. J. geophys. Res. 86, 721s-7254. McFarland D., Kley D., Drummond J. W., Schmellekopf A. L. and Winkler R. H. (1979) Nitric oxide measurements in the equatorial Pacific region. Geophys. Res. Left. 6, 605608. National Academy of Sciences (1976) Halocarbons: Eficts on Stratospheric Ozone. Printing and Publishing Office, National Academy of Sciences, Washington, DC. Singh H. B. and Hanst P. L. (1981) Peroxyacetyl nitrate (PAN) in the unpolluted atmosphere: an important reservoir for nitrogen oxides. Geophys. Res. Left. 8,941-944. Singh H. B. and Salas L. J. (1982) Measurement of selected light hydrocarbons over the Pacific ocean. Latitudinal and seasonal variations. Geophys. Res. Left. 9, 842-845. Singh H. B. and Salas L. J. (1983) Peroxyacctyl nitrate in the free troposphere. Nature 302, 326328. Stedman D. H. and Cantrell C. A. (1982) Laboratory studies of an HOr/RO, detector. Proceedings of 2nd Symposium on the Composition of the Nonurban Troposphere. American Meteorological Society, Boston, MA. Stephens E. R., Hanst P. L. and Doerr R. C. (1957) Infrared spectra of aliphatic peroxyacids. Analyt. Chem. 29, 776777. Su F., Calvert J. G., Shaw J. H., Niki H., Maker P. D., Savage C. M. and Breitenbach L. D. (1979) Spectroscopic and kinetic studies of a new metastable species in the photooxidation of gaseous formaldehyde. Chem. Phys. Lett. 65, 221-225.