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Oxidation of methanol by hydroxyl radicals in aqueous solution under simulated cloud droplet conditions
Atmospheric Environment 34 (2000) 5283 } 5294
Oxidation of methanol by hydroxyl radicals in aqueous solution under simulated cloud droplet conditions Anne Monod *, Abderaouf Chebbi, ReH gine Durand-Jolibois, Patrick Carlier Laboratoire de Chimie et Environnement, EA2678, Universite& de Provence, case 29, 3 Place Victor Hugo, 13331 Marseille, Cedex 03, France Laboratoire Interuniversitaire des Syste% mes Atmosphe& riques, Universite& s Paris 7 et Paris 12, URA CNRS 7583, 61 avenue du Ge& ne& ral de Gaulle, 94010 Cre& teil Cedex, France Received 30 November 1999; received in revised form 6 March 2000; accepted 24 March 2000
Abstract The results of a detailed mechanistic study of aqueous-phase OH-oxidation of methanol are presented. Analysis of reaction products by speci"c chromatographic methods revealed that hydrated formaldehyde is not the only stable primary reaction product. Formic acid and/or formate ion are also stable primary molecular reaction products of methanol OH-oxidation. The branching ratios for their formation are highly pH dependent. At pH"7, hydrated formaldehyde is the dominant molecular reaction product (ratio 4.5 : 1 for hydrated formaldehyde : formate ion), whereas at pH"2, formic acid is the dominant product (ratio 3.7 : 1 for formic acid : hydrated formaldehyde). At all pH studied, the sum of the primary stable products represents 49 ($11)% of methanol removal, in agreement with the amount of OOCH OH radicals formed relative to methanol removal 48 ($2)%. The formation of primary formic acid at pH"2 is attributed to OOCH OH self-reaction, and the strong pH e!ect is attributed to the base-catalyzed decomposition of OOCH OH leading to the formation of hydrated formaldehyde. Evaporation and/or an addition reaction between CH OH and HO radicals leading to the formation of hydroxymethyl hydroperoxide is proposed to explain the missing yields. The implications of this mechanism to atmospheric chemistry are discussed. 2000 Elsevier Science Ltd. All rights reserved. Keywords: Oxygenated volatile organic compounds; Hydroxymethyl peroxyl radicals; Aaqueous-phase photochemistry; Hydrated formaldehyde; Formate and formic acid
1. Introduction Cloud chemistry models (Jacob, 1986; Jacob et al., 1989; Pandis and Seinfeld, 1989; Lelieveld and Crutzen, 1991; Walcek et al., 1997; Monod and Carlier, 1999) have shown that, at a local scale and during the daytime, aqueous-phase photochemistry can have a drastic e!ect on gas-phase concentrations of important tropospheric species such as OH, HO and O . This e!ect depends on aqueous-phase mechanisms. OH to HO conversion by dissolved organic compounds can lead to
* Corresponding author. Tel.: #33-491-106227; fax: #33491-106377. E-mail address: [email protected] (A. Monod).
signi"cant HO build-up. Thus, it is important to understand the detailed mechanisms of aqueous-phase OH-oxidation of dissolved organic compounds under atmospheric conditions (Jacob, 2000). In some cloud chemistry models (Lelieveld and Crutzen, 1991), only simple aldehydes and acids are considered, neglecting the in#uence of other soluble species such as alcohols. When the latter are considered, methanol is taken into account (Jacob et al., 1989; Pandis and Seinfeld, 1989; Walcek et al., 1997), but it is assumed that its aqueous-phase OHinitiated oxidation leads solely to the production of formaldehyde. The experimental evidence on which this assumption is based is scant. Aqueous OH-oxidation of methanol has been previously studied, but not under atmospheric conditions (Asmus et al., 1973; Bothe et al., 1977, 1978; Bothe and Schulte-Frohlinde, 1978; Elliot
1352-2310/00/$ - see front matter 2000 Elsevier Science Ltd. All rights reserved. PII: S 1 3 5 2 - 2 3 1 0 ( 0 0 ) 0 0 1 9 1 - 6
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and McCracken, 1989; Neta et al., 1990; Huie and Clifton, 1993). However, the "rst steps of this oxidation were clearly identi"ed, and should not be modi"ed under atmospheric conditions. Aqueous-phase OH-attack on one of the H-atoms linked to the carbon atom of methanol is dominant (93%, Asmus et al., 1973): CH OH#OHPCH OH#H O,
(1)
k "1.0;10exp(!685/T) l mol\ s\ (Elliot and 2 McCracken, 1989). The organic radical formed is rapidly oxidized by dissolved O to produce hydroxymethyl peroxy radical: CH OH#O POOCH OH,
(2)
k "4.5;10 l mol\ s\ (Buxton et al., 1988). ) The behavior of the formed peroxy radical is less certain. A slow unimolecular decomposition leading to formaldehyde has been observed (Bothe et al., 1977, 1978): OOCH OHPHCHO#HO ,
(3)
k (10 s\ (Bothe et al., 1977, 1978) and this reac ) tion can be base catalyzed through an electron transfer mechanism: OOCH OH#OH\PHCHO#OH\#HO ,
(4)
k "18;10 l mol\ s\ (Neta et al., 1990). ) In aqueous solutions where hydroxymethyl peroxy radicals are highly concentrated (higher than 10\ M, such as in #ash photolysis experiments), a self-reaction has been previously observed (Bothe and Schulte-Frohlinde, 1978; Huie and Clifton, 1993): OOCH OH#OOCH OHPproducts,
(5)
k "1.0;10exp(!1400/T ) l mol\ s\ (Huie and 2 Clifton, 1993). However, the mechanism and speci"c reaction products of this reaction are not yet clear. The current study was undertaken to clarify the fate of the hydroxymethyl peroxy radical (OOCH OH) in atmospheric water droplet conditions. We present a laboratory study of the products formed during the aqueous-phase OH-initiated oxidation of methanol. All of the experimental conditions have been taken as close as possible to atmospherically relevant conditions, except for the initial concentrations of reactants, especially those of methanol. The latter were taken at 0.22}0.23 mM in order to obtain detectable amounts of any reaction product. Although we estimate such concentrations to be 10}100 times higher than in typical atmospheric droplets, these are much more dilute (by a factor '1000) than in previous #ash photolysis experiments (which sometimes
have been carried out in pure methanol). Therefore, in our experiments, which were as close as possible to typical atmospheric conditions, the organic radical self-reaction (5) was less likely than in previous #ash photolysis studies. Because OOCH OH radical decomposition is slow, especially at low pH (Bothe et al., 1977, 1978) we assumed the existence of an addition reaction between HO (or O\) and hydroxymethyl peroxy radical, leading to the formation of hydroxymethyl hydroperoxide (HMHP) OOCH OH#HO PHOOCH OH#O ,
(6)
which may be faster at pH'pK (HO /O\) in an elec tron transfer mechanism OOCH OH#O\P\OOCH OH#O &- HOOCH OH#OH\#O .(7) &
2. Experimental 2.1. Continuous photolysis apparatus The apparatus consisted of a Pyrex vessel of capacity 500 cm. The irradiation source was a metal halide lamp, HMI 575 W, OSRAM. Fig. 1 shows its spectrum, along with a standard solar spectrum for comparison. The lamp was mounted inside the vessel, above the solution, at a constant and reproducible distance from the surface. The volume of the aqueous solution was 220 cm, it was continuously stirred, and maintained at 279.0$0.2 K. The surface of the solution was swept by a constant #ow of ultrapure N in order to avoid any contamination by ambient air. This procedure also allowed us to avoid dissolved O saturation, thus enabling the monitoring of O concentrations as a function of time. The in#uence of pH was tested by studying two very di!erent constant values (pH"2 and 7, using phosphoric acid and a phosphate bu!er). Furthermore, a `freea pH (no bu!er) was also used, starting at pH+7 (and decreasing to pH+5 during the experiment because of acid production). This range of pH values adequately covers the range typically encountered in atmospheric water droplets. 2.2. OH radical yield OH radicals were generated from the continuous photolysis of H O , thus enabling the following mecha nism: H O #hlP2OH,
(8)
H O #OHPHO #H O,
(9)
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
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Fig. 1. Photoreactor's lamp spectrum (OSRAM HMI 575 W) compared to solar spectrum (zenith angle"03) (Demerjian et al., 1980).
k "8.25;10exp(!1680/T ) l mol\ s\ (Christen2 sen et al., 1982); HO #HO PH O #O ,
(10)
k "3.39;10exp(!2480/T ) l mol\ s\ (Christen2 sen and Sehested, 1988); &> H O #O , HO #O\ P
(11)
k "2.17;10exp(!914/T ) l mol\ s\ (Christen2 sen and Sehested, 1988). Because the solution was continuously stirred, it is assumed that H O was homogeneously irradiated. However, spectroradiometric measurements indicated that the actinic #ux of the lamp could vary (by as much as 25%) from one run to another. The variations were lower (less than 6%) during one continuous run of 9 h. Thus, the H O photolysis rate had to be evaluated for each experiment, which, therefore, was divided into three steps. In step 0, the aqueous solution (with no reactant) was irradiated in order to reach constant actinic #ux and temperature (O concentration was monitored). In step 1, a single aliquot of H O (2.3}11 mM) was introduced and H O photolysis was observed for 1 h (H O and O concentrations were monitored). In
step 2, a single aliquot of CH OH (0.23}0.31 mM) was introduced, and OH-oxidation of methanol was observed for 4 h (reactants and reaction products were monitored). Initial reactant concentrations were chosen so that k [CH OH]+k [H O ]. This allowed reaction (9) to be competitive, and thus, relatively high production of HO and O\ radicals, enabling reactions (6) and (7). 2.3. Analysis of reaction products The expected reaction products were formaldehyde, formic acid (formate ions) and HMHP. Formaldehyde in aqueous solution is in equilibrium with its hydrated form, methanediol, HCHO#H O&H C(OH)
(12)
with K "[H C(OH) ]/[HCHO]"4.9;10 (Bell, 1966). Therefore, we only consider hydrated formaldehyde in the following. Product analyses were performed using gas-phase chromatography (GC-FID), two highpressure liquid chromatography systems (HPLC -UV and HPLC -#uorimetry) and ionic chromatography (IC-conductimetry). The GC was equipped with a polar capillary column HP-INNOWAX. Solutions were analyzed for methanol by adding 300 ll from the vessel to an
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internal standard (1-butanol), prior to direct injection into the GC. HPLC was equipped with a reversed phase C column and UV detection at 360 nm. Solutions were analyzed for formaldehyde by removing 300 ll from the vessel and diluting it to 1.20 ml of 2,4-dinitrophenylhydrazine (2,4-DNPH, 0.16 g/l). Ionic chromatography was performed with a Dionex AS5 column, an AMMS suppressor column and a conductivity detector. Solutions were analyzed for formic acid (and formate ions) by removing 300 ll from the vessel and adding it to 4 ml of frozen ultrapure water. HPLC was equipped with a re versed phase C column. Following Hellpointner and GaK b (1989), parahydroxyphenyl acetic acid (and peroxidase) derivatization was used for detection of peroxides by #uorimetry (j "320 nm, and j "400 nm). Solutions were analyzed for H O and HMHP by re moving 20 ll from the vessel and diluting it to 2 ml of refrigerated (73C) ultrapure water. Dissolved O was continuously monitored by direct potentiometric measurements in the vessel. Except for the ionic chromatography (operated 12 h after each experiment, the samples being stored at !203C), all analyses were performed less than 2 h after sampling from the vessel. Each analytical technique gave rise to sharp peaks, and calibration of each compound (in the range covering the concentrations encountered in the experiments) gave retention times, statistical error limits and detection limits that are summarized in Table 1. 2.4. Reagents The reagents used were H O (not stabilized, Fluka, 30 wt%); CH OH and 1-butanol (Prolabo, normapur grade, 99.8%); N (I grade, Air Liquide); H PO (Pro labo, normapur grade); NaOH (Prolabo, normatom grade). Solutions were prepared using water puri"ed by an ELGASTAT maxima-HPLC system, including reverse osmosis, micro-"ltration, nuclear-grade deionization, activated carbon modules and an irradiation module at 254 nm.
2.5. Glass cleaning procedure Because we operated with diluted solutions, and in order to avoid any contamination, the reactor and all the sampling vessels were carefully cleaned using the following procedure: three rinses with reverse osmosed water; 3 h soaking in 2% of DECON diluted in reverse osmosed water; 10 rinses with ultrapure EGASTAT water; 3 h soaking in 2% of HCl diluted in ultrapure EGASTAT water; 10 rinses with ultrapure EGASTAT water; 12 h soaking in ultrapure EGASTAT water; three rinses with ultrapure EGASTAT water; drying in an ultraclean laminar #ow vent (class 10). 2.6. Data analysis Some tests by curve "tting were carried out using the FACSIMILE chemical modelling package (FACSIMILE, 1994).
3. Results Monitored reactants and products are presented in Fig. 2 as a function of time, for an experiment performed at pH"2. O pro"les clearly mimic the di!erent steps: during step 0, O decreases by evaporation because the surface of the solution is swept by pure N . At the beginning of step 1, O production by reaction (10) starts and dominates over evaporation. Just before step 2, because of the [H O ] decrease, the e$ciency of reaction (10) decreases, and O evaporation becomes competitive with reaction (10), leading to stabilization of [O ]. At the beginning of step 2, O consumption by reaction (2) starts and dominates over production by reaction (10). When 80% of methanol is consumed, the e$ciency of reaction (2) decreases, and O production by reaction (10) becomes competitive with reaction (2) leading to stabilization of [O ]. When all organic compounds are consumed (both methanol and reaction products), O production by reaction (10) dominates
Table 1 Calibration of each compound with each analytical technique Compound
Analytical technique
Retention time (min)
Detection limit (M)
Statistical error ($2.6p) (%)
CH OH H O HMHP CH (OH) HCOOH O
GC-FID HPLC (HPLC-Fl) HPLC (HPLC-Fl) HPLC (HPLC-UV) Ionic chromatography Potentiometry
2.93 5.0 5.3 6.0 8.3 *
10\ 10\ 10\ 10\ 5;10\ 3;10\
$10 $15 $15 $9 $10 $2.6
HPLC-Fl"high-pressure liquid chromatography equipped with #uorescence detection. HPLC-UV"high-pressure liquid chromatography equipped with UV detection.
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Fig. 2. Aqueous-phase OH-oxidation of methanol: reactant and product pro"les as a function of time at pH"2.
over reaction (2). At the end of the experiment, because initial [H O ] has decreased by a factor of 2 or more, O production by reaction (10) is in competition with O evaporation, which leads to stabilization of [O ]. 3.1. Reaction products 3.1.1. Hydrated formaldehyde and formic acid or formate ion At all pH levels studied, substantial amounts of hydrated formaldehyde, formic acid, and formate ion (both are designed as `formatea in the following) were detected (Fig. 3). Furthermore, `formatea pro"les indicate that this compound was produced directly from methanol without a stable molecular intermediate. Assuming that `formatea was exclusively produced from hydrated formaldehyde, one would have observed a null initial formation #ux of `formatea, and an in#ection on its pro"le when the maximum of hydrated formaldehyde concentration occurred. Fig. 3 shows that this assumption is not valid in our experiments, especially at pH"2 where the initial `formatea formation #ux was even higher than that of hydrated formaldehyde. 3.1.2. Hydroxymethyl hydroperoxide No detectable amount of HMHP was identi"ed at any pH value. However, using HPLC , the HMHP peak is
eluted close to the H O peak (Table 1). Calibration experiments have shown that these two compounds coelute when [HMHP]/[H O ])1/10. Thus, when initial H O concentrations were 2.3 mM (respectively, 11 mM), HMHP could not be detected under 0.23 mM (respectively, 1.1 mM). We performed several separate tests, which have shown that (i) HMHP response in HPLC is 100%; (ii) in the equilibrium (reaction (13)) of peroxidation of hydrated formaldehyde, reaction >13 is too slow in our conditions (279 K) to be signi"cant in our experiments; (iii) (reaction \13) is too slow to be shifted to hydrated formaldehyde formation in the presence of 2,4-DNPH prior to injection on HPLC : CH (OH) #H O & HOOCH OH#H O,
(13)
17($2);10 k> "exp 51($8)! (M\ s\) 2 ¹ (between 279 and 298 K),
2.1($0.3);10 K "exp !2.3($0.96)# (M\) 2 ¹ (between 279 and 298 K). Therefore, if HMHP was formed in our experiments, it was produced only from the studied reaction (OHoxidation of CH OH), and its concentration did not
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Fig. 3. pH e!ect on the reaction products.
exceed one-tenth of H O concentrations (i.e. less than 2}3;10\ M for experiments starting with 2}3;10\ M of H O and less than 1}1.1;10\ M for experiments starting with 10}11;10\ M of H O ).
hydrated formaldehyde formation rate was on average 4.5 ($1.0) times higher than that of `formatea, whereas, at `freea pH (i.e. initial pH"7 decreasing to 5 as the acid appeared), this ratio was 3.7 ($1.0), and at pH"2, this ratio decreased to 0.34 ($0.10).
3.2. pH ewect on reaction products
3.3. First step carbon balance
pH had a drastic e!ect on the initial #uxes of the identi"ed reaction products (Fig. 3). At pH"7, the
Because measurements of "nal products such as CO and CO could not be made, "rst step carbon balance
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Table 2 Experimental results: initial reactant concentrations, percentages of reaction products formed relative to methanol consumption, and "rst step carbon balance pH
[H O ] (;10\ M)
[CH OH] (;10\ M)
% Hydrated formaldehyde formed
% `formatea formed
Carbon balance (%)
2 2 2 2 `freea `freea 7 7
3.0$0.4 2.7$0.4 11$2 3.2$0.5 2.3$0.3 2.6$0.4 2.8$0.4 10$2
2.4$0.2 2.3$0.5 3.1$0.4 3.0$0.3 2.5$0.2 2.2$0.5 2.3$0.2 2.4$0.5
11$2 8$2 15$3 14$2 43$7 35$11 31$5 45$13
49$9 31$10 37$8 36$7 10$2 12$4 9$2 8$2
60$11 39$12 52$11 50$9 53$9 47$15 40$7 53$15
calculations were performed. This method consisted in a comparison between methanol consumption and the primary stable products identi"ed (i.e. `formatea and hydrated formaldehyde) within the "rst 10}20 min of each experiment (Table 2). Only the very "rst few minutes of each experiment were taken into account, so `formatea and hydrated formaldehyde OH-oxidation was not important. The induced underestimate of the "rst step carbon balance was evaluated as follows: CH (OH) #OHPCH(OH) #H O, (14) k "2.4;10exp(!1020/T) l mol\ s\ (Chin and 2 Wine, 1994); HCOOH#OHPCOOH#H O, (15) k "3.0;10exp(!991/T ) l mol\ s\ (Chin and 2 Wine, 1994); HCOO\#OHPCOO\#H O, (16) k "2.0;10exp(!1240/T ) l mol\ s\ (Chin and 2 Wine, 1994). The "rst step carbon balance underestimate was calculated taking into account k /k , k /k and k /k ratios for our temperature conditions. For hydrated formaldehyde, the underestimate was 10% of the "rst step carbon balance percentage. For formic acid and formate ion, the analogous underestimates were 5 and 20%, respectively. All these underestimates are within the error bars indicated in Table 2. The average "rst step carbon balance was "nally 49$11%, indicating that half of CH OH removal was not identi"ed. 4. Discussion 4.1. CH3 OH conversion into OOCH2 OH radicals Was CH OH entirely converted into OOCH OH radicals? The answer to this question is essential before
discussing the fate of OOCH OH radicals. We evaluated the amount of RO radicals formed relative to methanol consumption. Stoechiometric coe$cients are equal to 1 in reactions (1) and (2). Thus, assuming steady state for CH OH radicals, one molecule of dissolved O should be consumed per consumed CH OH molecule, produ cing one OOCH OH radical. Therefore, comparing the amount of methanol consumed to the amount of dissolved O consumed, one can evaluate the amount of OOCH OH radicals formed relative to methanol con sumption. The evaluation of dissolved O consumed just after methanol injection (beginning of step 2) was carried out using the following equation: [O ] "[O ] ![O ] , (I) where [O ] is the measured concentrations in our experiments, and [O ] is the extrapolated concentra tions from the simulation of the `base linea of dissolved O concentrations during methanol reactivity (Fig. 4). [O ] was obtained by curve "tting both O and H O pro"les during step 1 (i.e. H O photolysis), while (reaction (8)) and also optimizing the optimizing J & - >FJ rate constant of O evaporation. An accurate curve-"t of O pro"les was obtained (Fig. 4) because of the low uncertainties in the measurements of its concentrations ($2.6%). We applied this analysis to the experiments performed at pH"2, 7 and `freea pH. Comparing the obtained [O ] to the consumed methanol (Fig. 5), we found that the ratio [consumed methanol]/[consumed O ] was 2.1$0.1. Therefore, the amount of OOCH OH radicals formed was only 48 ($2)% of the amount of CH OH consumed at all pH values. This result is in very good agreement with the "rst step carbon balance of 49 ($11)% at all pH values. We have checked in separate tests that methanol consumption by evaporation and by direct reaction with H O represented less than 5% of methanol consumption by OH-oxidation. Thus, 52 ($2)% of the CH OH radicals formed in reaction (1) must have undergone another pathway than
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Fig. 4. Experimental (exp) and simulated (simu) dissolved O pro"les at `freea pH. Simulation of step 1 (i.e. H O continuous photolysis) was performed by curve "tting of O and H O pro"les (see text). At t ("6307 s), methanol was injected into the vessel. At t ("14494 s), all the organic species (reactants and products) were consumed.
reaction (2). This is explained by the fact that dissolved O did not saturate the solution because of the continu ous N #ow at the surface of the solution. CH OH radicals were in their acid form in all our experimental conditions, because pK (CH OH/CH O\)"10.7 (Rabani et al., 1974). Following the "ndings by McElroy and Waygood (1991) on CH(OH) radicals, it is possible that the following reactions occurred: CH OH#H O PCH OH#HO (17) CH OH#H O PCH (OH) #OH. (18) However, these reactions would have been limited by reaction (9). CH OH self-reaction may also have occur red:
Fig. 5. Comparison between the amount of CH OH consumed and the amount of dissolved O consumed just after CH OH injection, at `freea pH (same experiment as shown in Fig. 3). Concentrations versus time (a); consumed methanol versus consumed dissolved O plot (b).
CH OH#CH OHPCH OH#HCHO. (19) Reformation of methanol by reaction (17) would not a!ect the "rst step carbon balance, thus, reaction (17) may have occurred, but it would not explain the elimination of CH OH radicals. Formation of hydrated formal dehyde by reacions (18) and (19) would a!ect the "rst step carbon balance. However, no in#uence of initial [H O ] was observed on the "rst step carbon balance (Table 2). If reaction (18) or (19) had occurred signi"cantly, the "rst step carbon balance would have been closer to 100%. It is probable that CH OH radicals have been eliminated from the solution by evaporation, and/or by addition of
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
HO radicals, leading to the formation of HMHP: CH OH#HO PHOOCH OH.
(20)
However, HMHP was not formed at detectable levels in our experiments (i.e. [HMHP](0.23 mM or [HMHP](1.1 mM, depending on initial concentrations of H O ). No further experiments were carried out to investigate the elimination of CH OH radicals because our goal was to examine the fate of OOCH OH radicals (atmospheric conditions). 4.2. The fate of OOCH2 OH radicals For all the reasons mentioned above, all the identi"ed products were attributed to OOCH OH radical reactivity. OOCH OH radicals were in their acid form in all our experimental conditions because pK (OOCH OH/OOCH O\) '8 (Rabani et al., 1974). 4.2.1. HMHP Compared to previous #ash photolysis studies, our experiments were carried out under conditions closer to the atmosphere, and under conditions more favorable to HMHP formation. However, insigni"cant amounts of HMHP were formed from OOCH OH radical reactiv ity. Therefore, reactions (6) and (7) were of negligible importance at all pH tested. 4.2.2. Primary `formatea formation Direct formation of formic acid from OOCH OH self-rection (reaction (5)) has been previously suggested (Bothe and Schulte-Frohlinde, 1980; Rabani et al., 1974; Huie and Clifton, 1993), and three di!erent mechanisms have been proposed for this reaction:
`formatea. This is clearly not the case, especially at pH"2 where `formatea formation rate is even higher than that of hydrated formaldehyde (Fig. 3). Thus, reaction (5) does not proceed by mechanism A. Assuming reaction (5) proceeds by mechanism B, the budget of reactions (1), (2) and mechanism B would result in the net consumption of two molecules of dissolved O per molecule of methanol. However, this is in contradiction with our experiments, where two molecules of methanol were consumed per molecule of O (Fig. 5). Finally, mechanism C does not contradict our experiments if one supposes that tetroxide (HOCH OO} OOCH OH) is so unstable that it could not be detected in HPLC . However, further experiments are needed to clarify the details of mechanism C; for example by generating OH radicals with another source (than H O ) and determining the amount of H O produced by C. 4.2.3. pH ewect The observed drastic pH e!ect on the amounts of reaction products was attributed to the catalytic e!ect of hydroxyl ions (reaction (4)). When the pH increased from 2 to 7, [OH\] increased from 10\ to 10\ M, thus the rate of decomposition of OOCH OH radicals via reac tion (4) increased to the detriment of the rate of OOCH OH self-reaction (5). This is in agreement with our experiments: comparing the pH"2 and 7 experiments, the proportion of hydrated formaldehyde formed relative to that of `formatea was inverse (Fig. 3 and Table 2). The ratio of hydrated formaldehyde formed at pH"7 to that formed at pH"2 was 3.2$1.3, which is equal to the ratio: +rate of reaction (4),/+rate of reaction (3),, thus enabling an estimate of k at pH"7 and at 279 K: k [OH\] "3.2$1.3. k
2OOCH OHP2OCH OH#O followed by 2OCH OHPHCOOH#CH (OH) or followed by
(A)
OCH OH#O PHCOOH#HO , or: 2OOCH OH[PHOCH OO}OOCH OH]
(B)
2HCOOH#H O . (C) Assuming reaction (5) proceeds by mechanism A, equal amounts of `formatea and hydrated formaldehyde would be formed from this sequence and an additional amount of hydrated formaldehyde would be generated from reactions (3) and (4). Therefore, in our experiments, one would have detected the formation of hydrated formaldehyde in equal or in higher amounts than that of
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(II)
Considering that k (298 K)(10 s\ (Bothe et al., 1977, 1978), thus, at 279 K, and pH"7, k (279 K)(3.2 ($1.3);10 l mol\ s\. Comparing this value with k (295 K)"18;10 l mol\ s\ (Neta et al., 1990), it is clear that k is highly exothermic. These considerations give an estimate of the lower value of the activation energy of reaction (4) of the order of 172 kJ mol\. Thus, temperature should have an important in#uence on the fate of OOCH OH radicals, especially at high pH levels. 4.3. Implications for atmospheric chemistry The results of our experiments have shown that the presence of relatively high levels of HMHP in some atmospheric water droplets (up to 1 lM in cloudwater,
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Sauer et al., 1996) is not the consequence of aqueousphase photochemistry of OOCH OH radicals. HMHP is more probably formed at the droplet interface from Criegee biradicals (GaK b et al., 1995), which are produced in the gas phase by ozonolysis of alkenes. Further, the present investigation indicated that aqueous-phase OHinitiated oxidation of methanol does not solely lead to hydrated formaldehyde, it also forms formic acid, with no stable molecular intermediary, by self-reaction of OOCH OH radicals. Therefore, the importance of this reaction is sensitive to OOCH OH concentrations, and thus to initial methanol concentration. However, compared to atmospheric conditions, we have operated with relatively high concentrations of methanol. To explore whether reaction (5) is still signi"cant under atmospheric conditions, the mechanism of OOCH OH fate elucidated here was included in a cloud box model (Monod and Carlier, 1999). This box model was initialized with gas-phase methanol concentrations similar to those found in "eld measurements, and it was run at 279 K, similar to our experimental conditions. The rate of formation of `formatea relative to that of the sum of hydrated formaldehyde and `formatea from aqueousphase OH-oxidation of methanol (ratio R) has been calculated:
base-catalyzed unimolecular decomposition of OOCH OH radicals. However, compared to realistic atmospheric conditions, only a limited number of parameters have been investigated here. It is known that transition metals play an important role in aqueous radical chemistry (Walcek et al., 1997). Moreover, dissolved organic RO radicals may interact with transition metals (Graedel et al., 1986). To our knowledge, however, no experimental evidence has yet been carried out on these interactions under atmospheric conditions. Therefore, further experimental and modelling studies are needed to assess the in#uence of temperature, ionic strength, and the presence of transition metals on the behavior of RO radicals within the aqueous phase of atmospheric water droplets. 5. Conclusions Analysis of reaction products of aqueous-phase OHoxidation of methanol revealed that hydrated formaldehyde is not the only stable primary reaction product. Formic acid and/or formate ion were also identi"ed as stable primary molecular reaction products. The branching ratios for the formation of the products were highly
2k [OOCH OH] R (%)" ;100. 2k [OOCH OH]#k [OOCH OH]#k [OOCH OH][OH\] The in#uence of pH and initial gas-phase concentrations of methanol on R was tested (Table 3). Table 3 shows that R is much more sensitive to pH values than to initial gas-phase concentrations of methanol. Moreover, except at pH 6, `formatea formation is signi"cant. Therefore, the formation of primary `formatea from OH-oxidation of methanol has to be taken into account in cloud models. The present study has shown that the aqueous-phase chemistry of OOCH OH radicals is strongly in#uenced by pH. The in#uence of changing concentrations of the parent compound is less signi"cant (Table 3). This study also suggested that temperature has a strong in#uence on
pH dependent. At pH"7, hydrated formaldehyde was the dominant molecular reaction product (ratio 4.5 : 1 for hydrated formaldehyde : formate ion), whereas at pH"2, formic acid was the dominant product (ratio 3.7 : 1 for formic acid : hydrated formaldehyde). At all pH values studied, the sum of the primary stable products represented 49 ($11)% of methanol removal, in agreement with the amount of OOCH OH radicals formed relative to methanol removal (48 ($2)%). The formation of primary formic acid at pH"2 was attributed to OOCH OH self-reaction, and the strong pH e!ect
Table 3 Cloud box model results (including both gas and aqueous phases): in#uence of pH and initial gas-phase concentration of methanol on R pH
R (%) (initial [CH OH ]"4.3 ppbV)
R (%) (initial [CH OH ]"10 ppbV)
3 4.16 5.2 6
39 21 5 0.43
48 29 7.5 0.67
Rate of formation of `formatea relative to that of the sum of hydrated formaldehyde and `formatea from aqueous-phase OHoxidation of methanol: 2k [OOCH OH] R (%)" ;100. 2k [OOCH OH]#k [OOCH OH]#k [OOCH OH][OH\]
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
was attributed to the base-catalyzed decomposition of OOCH OH leading to the formation of hydrated formaldehyde. Evaporation and/or an addition reaction between CH OH and HO radicals leading to the formation of hydroxymethyl hydroperoxide was proposed to explain the missing yields. However, hydroxymethyl hydroperoxide could not be detected under 0.23 mM or under 1.1 mM, depending on H O concen trations. All the experimental conditions were taken as close as possible to atmospheric conditions, except for the initial concentrations of reactants, especially those of methanol. The latter were taken at 0.22}0.23 mM in order to obtain detectable amounts of reaction products. Although such concentrations are supposed to be 10}100 times higher than in atmospheric droplets, these are much more dilute conditions (by a factor '1000) than in previous experimental studies on OH-oxidation of methanol. To explore whether the two identi"ed OOCH OH pathways (de composition and self-reaction) were signi"cant under atmospheric conditions, the mechanism elucidated was included in a cloud box model comprising more realistic concentration and pH conditions than in our experiments. At typical pH values (pH"4}5) encountered in clouds, both pathways were signi"cant. It was concluded that under atmospheric conditions, aqueous-phase OHinitiated oxidation of methanol does not lead solely to formaldehyde, but also to primary formation of formic acid, and this should be included in cloud models. However, only a limited number of parameters have been investigated, and further experimental and modelling studies are needed to assess the in#uence of temperature, ionic strength, and the presence of transition metals species on the behavior of RO radicals within the aque ous phase. Because these parameters may have di!erent in#uences depending on the chemical nature of RO , di!erent parent compounds also need to be studied (hydroxycarbonyls, ethers, esters, peroxides, etc.). Acknowledgements The authors thank `la feH deH ration des industries de la parfumeriea for the "nancial support of this study, and Dr. J.F. Doussin, Dr. J. Colman and Dr. M. McGrath for helpful comments on the manuscript. References Asmus, K.D., MoK ckel, H., Henglein, A., 1973. Pulse radiolytic of site of OH radical attack on aliphatic alcohols in aqueous solution. Journal of Physical Chemistry 77 (10), 1218}1221. Bell, R.P., 1966. The reversible hydration of carbonyl compounds. In: Gold, V. (Ed.), Advances in Physical Organic Chemistry, Vol. 4. Academic Press, London, pp. 1}27.
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Bothe, E., Behrens, G., Schulte-Frohlinde, D., 1977. Mechanism of the "rst order decay of 2-hydroxy-propyl-2-peroxyl radicals and of O\ formation in aqueous solution. Zeitschrift fur Naturforschung 32b, 886}889. Bothe, E., Schuchmann, M.N., Schulte-Frohlinde, D., von Sonntag, C., 1978. HO elimination from a-hydroxyalkylperoxyl radicals in aqueous solution. Photochemistry and Photobiology 28, 639}644. Bothe, E., Schulte-Frohlinde, D., 1978. The bimolecular decay of a-hydroxymethylperoxyl radicals in aqueous solution. Zeitschrift fur Naturforschung 33b, 786}788. Bothe, E., Schulte-Frohlinde, D., 1980. Reaction of dihydroxymethyl radical with molecular oxygen in aqueous solution. Zeitschrift fur Naturforschung 35b, 1035}1039. Buxton, G., Greenstock, C.L., Helman, W.P., Ross, A., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atom and hydroxyl radicals (OH/O\) in aqueous solution. Journal of Physical Chemistry Reference Data 17 (2), 513}883. Chin, M., Wine, P.H., 1994. A temperature dependent competitive kinetics study of the aqueous-phase reaction of OH radicals with formate, formic acid, acetate, acetic acid, and hydrated formaldehyde. In: Helz, G.R., Zepp, R.G., Crosby, D.G. (Eds.), Aquatic and Surface Photochemistry. Boca Raton, pp. 85}98. Christensen, H., Sehested, K., 1988. HO and O\ radicals at elevated temperatures. Journal of Physical Chemistry 92, 3007}3011. Christensen, H., Sehested, K., Cor"tzen, H., 1982. Reactions of hydroxyl radicals with hydrogen peroxide at ambient and elevated temperatures. Journal of Physical Chemistry 86, 1588}1590. Demerjian, K.L., Schere, K.L., Peterson, J.T., 1980. Theoretical estimation of actinic #ux and photolytic rate constants of atmospheric species in the lower atmosphere. Advances in Environmental Science and Technology 10, 369}459. Elliot, A.J., McCracken, D.R., 1989. E!ect of temperature on O\ reactions and equilibria: a pulse radiolysis study. International Journal of Applied Radiation and Instrument C 33 (1), 69}74. Facsimile, 1994. User Guide v.4.0, copy 4013. AEA Technology, Harwell, Didcot, Oxford. GaK b, S., Turner, W.V., Wol!, S., Becker, K.H., Ruppert, L., Brockmann, K.J., 1995. Formation of alkyl and hydroxyalkyl hydroperoxides on ozonolysis in water and in air. Atmospheric Environment 29 (18), 2401}2407. Graedel, T.E., Mandich, M.L., Weschler, C.J., 1986. Kinetic model studies of atmospheric droplet chemistry 2. Homogeneous transition metal chemistry in raindrops. Journal of Geophysical Research 91 (D4), 5205}5221. Hellpointner, E., GaK b, S., 1989. Detection of methyl, hydroxymethyl and hydroxyethyl hydroperoxides in air and precipitation. Nature 337, 631}634. Huie, R.E., Clifton, C.L., 1993. Kinetics of the self-reaction of hydroxymethylperoxyl radicals. Chemical Physics Letters 205 (2,3), 163}167. Jacob, D.J., 1986. Chemistry of OH in remote clouds and its role in the production of formic acid and peroxymonosulfate. Journal of Geophysical Research 91 (D9), 9807}9826. Jacob, D.J., 2000. Heterogeneous chemistry and tropospheric ozone. Atmospheric Environment 34, 2131}2159.
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Jacob, D.J., Gottlieb, E.W., Prather, M.J., 1989. Chemistry of a polluted cloudy boundary layer. Journal of Geophysical Research 94, 12975}13002. Lelieveld, J., Crutzen, P.J., 1991. The role of clouds in tropospheric photochemistry. Journal of Atmospheric Chemistry 12, 229}267. McElroy, W.J., Waygood, S.J., 1991. Oxidation of formaldehyde by the hydroxyl radical in aqueous solution. Journal of the Chemical Society, Faraday Transactions 87 (10), 1513}1521. Monod, A., Carlier, P., 1999. Impact of clouds on the tropospheric ozone budget: direct e!ect of multiphase photochemistry of soluble organic compounds. Atmospheric Environment 33 (27), 4431}4446. Neta, P., Huie, R.E., Ross, A.B., 1990. Rate constants for reactions of peroxyl radicals in #uid solutions. Journal
of Physical and Chemical Reference Data 19 (N2), 413}498. Pandis, S.N., Seinfeld, J.H., 1989. Sensitivity analysis of a chemical mechanism for aqueous-phase atmospheric chemistry. Journal of Geophysical Research 94, 1105}1126. Rabani, J., Klugh-Roth, D., Henglein, A., 1974. Pulse radiolytic investigations of OHCH O radicals. Journal of Physical Chemistry 78 (21), 2089}2093. Sauer, F., Schuster, G., SchaK fer, C., Moortgat, G.K., 1996. Determination of H O and organic peroxides in cloud-and-rain water on the Kleiner Feldberg during Feldex. Geophysical Research Letters 23 (19), 2605}2608. Walcek, C.J., Yuan, H.H., Stockwell, W.R., 1997. The in#uence of aqueous-phase chemical reaction on ozone formation in polluted and nonpolluted clouds. Atmospheric Environment 31 (8), 1221}1237.
Oxidation of methanol by hydroxyl radicals in aqueous solution under simulated cloud droplet conditions Anne Monod *, Abderaouf Chebbi, ReH gine Durand-Jolibois, Patrick Carlier Laboratoire de Chimie et Environnement, EA2678, Universite& de Provence, case 29, 3 Place Victor Hugo, 13331 Marseille, Cedex 03, France Laboratoire Interuniversitaire des Syste% mes Atmosphe& riques, Universite& s Paris 7 et Paris 12, URA CNRS 7583, 61 avenue du Ge& ne& ral de Gaulle, 94010 Cre& teil Cedex, France Received 30 November 1999; received in revised form 6 March 2000; accepted 24 March 2000
Abstract The results of a detailed mechanistic study of aqueous-phase OH-oxidation of methanol are presented. Analysis of reaction products by speci"c chromatographic methods revealed that hydrated formaldehyde is not the only stable primary reaction product. Formic acid and/or formate ion are also stable primary molecular reaction products of methanol OH-oxidation. The branching ratios for their formation are highly pH dependent. At pH"7, hydrated formaldehyde is the dominant molecular reaction product (ratio 4.5 : 1 for hydrated formaldehyde : formate ion), whereas at pH"2, formic acid is the dominant product (ratio 3.7 : 1 for formic acid : hydrated formaldehyde). At all pH studied, the sum of the primary stable products represents 49 ($11)% of methanol removal, in agreement with the amount of OOCH OH radicals formed relative to methanol removal 48 ($2)%. The formation of primary formic acid at pH"2 is attributed to OOCH OH self-reaction, and the strong pH e!ect is attributed to the base-catalyzed decomposition of OOCH OH leading to the formation of hydrated formaldehyde. Evaporation and/or an addition reaction between CH OH and HO radicals leading to the formation of hydroxymethyl hydroperoxide is proposed to explain the missing yields. The implications of this mechanism to atmospheric chemistry are discussed. 2000 Elsevier Science Ltd. All rights reserved. Keywords: Oxygenated volatile organic compounds; Hydroxymethyl peroxyl radicals; Aaqueous-phase photochemistry; Hydrated formaldehyde; Formate and formic acid
1. Introduction Cloud chemistry models (Jacob, 1986; Jacob et al., 1989; Pandis and Seinfeld, 1989; Lelieveld and Crutzen, 1991; Walcek et al., 1997; Monod and Carlier, 1999) have shown that, at a local scale and during the daytime, aqueous-phase photochemistry can have a drastic e!ect on gas-phase concentrations of important tropospheric species such as OH, HO and O . This e!ect depends on aqueous-phase mechanisms. OH to HO conversion by dissolved organic compounds can lead to
* Corresponding author. Tel.: #33-491-106227; fax: #33491-106377. E-mail address: [email protected] (A. Monod).
signi"cant HO build-up. Thus, it is important to understand the detailed mechanisms of aqueous-phase OH-oxidation of dissolved organic compounds under atmospheric conditions (Jacob, 2000). In some cloud chemistry models (Lelieveld and Crutzen, 1991), only simple aldehydes and acids are considered, neglecting the in#uence of other soluble species such as alcohols. When the latter are considered, methanol is taken into account (Jacob et al., 1989; Pandis and Seinfeld, 1989; Walcek et al., 1997), but it is assumed that its aqueous-phase OHinitiated oxidation leads solely to the production of formaldehyde. The experimental evidence on which this assumption is based is scant. Aqueous OH-oxidation of methanol has been previously studied, but not under atmospheric conditions (Asmus et al., 1973; Bothe et al., 1977, 1978; Bothe and Schulte-Frohlinde, 1978; Elliot
1352-2310/00/$ - see front matter 2000 Elsevier Science Ltd. All rights reserved. PII: S 1 3 5 2 - 2 3 1 0 ( 0 0 ) 0 0 1 9 1 - 6
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and McCracken, 1989; Neta et al., 1990; Huie and Clifton, 1993). However, the "rst steps of this oxidation were clearly identi"ed, and should not be modi"ed under atmospheric conditions. Aqueous-phase OH-attack on one of the H-atoms linked to the carbon atom of methanol is dominant (93%, Asmus et al., 1973): CH OH#OHPCH OH#H O,
(1)
k "1.0;10exp(!685/T) l mol\ s\ (Elliot and 2 McCracken, 1989). The organic radical formed is rapidly oxidized by dissolved O to produce hydroxymethyl peroxy radical: CH OH#O POOCH OH,
(2)
k "4.5;10 l mol\ s\ (Buxton et al., 1988). ) The behavior of the formed peroxy radical is less certain. A slow unimolecular decomposition leading to formaldehyde has been observed (Bothe et al., 1977, 1978): OOCH OHPHCHO#HO ,
(3)
k (10 s\ (Bothe et al., 1977, 1978) and this reac ) tion can be base catalyzed through an electron transfer mechanism: OOCH OH#OH\PHCHO#OH\#HO ,
(4)
k "18;10 l mol\ s\ (Neta et al., 1990). ) In aqueous solutions where hydroxymethyl peroxy radicals are highly concentrated (higher than 10\ M, such as in #ash photolysis experiments), a self-reaction has been previously observed (Bothe and Schulte-Frohlinde, 1978; Huie and Clifton, 1993): OOCH OH#OOCH OHPproducts,
(5)
k "1.0;10exp(!1400/T ) l mol\ s\ (Huie and 2 Clifton, 1993). However, the mechanism and speci"c reaction products of this reaction are not yet clear. The current study was undertaken to clarify the fate of the hydroxymethyl peroxy radical (OOCH OH) in atmospheric water droplet conditions. We present a laboratory study of the products formed during the aqueous-phase OH-initiated oxidation of methanol. All of the experimental conditions have been taken as close as possible to atmospherically relevant conditions, except for the initial concentrations of reactants, especially those of methanol. The latter were taken at 0.22}0.23 mM in order to obtain detectable amounts of any reaction product. Although we estimate such concentrations to be 10}100 times higher than in typical atmospheric droplets, these are much more dilute (by a factor '1000) than in previous #ash photolysis experiments (which sometimes
have been carried out in pure methanol). Therefore, in our experiments, which were as close as possible to typical atmospheric conditions, the organic radical self-reaction (5) was less likely than in previous #ash photolysis studies. Because OOCH OH radical decomposition is slow, especially at low pH (Bothe et al., 1977, 1978) we assumed the existence of an addition reaction between HO (or O\) and hydroxymethyl peroxy radical, leading to the formation of hydroxymethyl hydroperoxide (HMHP) OOCH OH#HO PHOOCH OH#O ,
(6)
which may be faster at pH'pK (HO /O\) in an elec tron transfer mechanism OOCH OH#O\P\OOCH OH#O &- HOOCH OH#OH\#O .(7) &
2. Experimental 2.1. Continuous photolysis apparatus The apparatus consisted of a Pyrex vessel of capacity 500 cm. The irradiation source was a metal halide lamp, HMI 575 W, OSRAM. Fig. 1 shows its spectrum, along with a standard solar spectrum for comparison. The lamp was mounted inside the vessel, above the solution, at a constant and reproducible distance from the surface. The volume of the aqueous solution was 220 cm, it was continuously stirred, and maintained at 279.0$0.2 K. The surface of the solution was swept by a constant #ow of ultrapure N in order to avoid any contamination by ambient air. This procedure also allowed us to avoid dissolved O saturation, thus enabling the monitoring of O concentrations as a function of time. The in#uence of pH was tested by studying two very di!erent constant values (pH"2 and 7, using phosphoric acid and a phosphate bu!er). Furthermore, a `freea pH (no bu!er) was also used, starting at pH+7 (and decreasing to pH+5 during the experiment because of acid production). This range of pH values adequately covers the range typically encountered in atmospheric water droplets. 2.2. OH radical yield OH radicals were generated from the continuous photolysis of H O , thus enabling the following mecha nism: H O #hlP2OH,
(8)
H O #OHPHO #H O,
(9)
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
5285
Fig. 1. Photoreactor's lamp spectrum (OSRAM HMI 575 W) compared to solar spectrum (zenith angle"03) (Demerjian et al., 1980).
k "8.25;10exp(!1680/T ) l mol\ s\ (Christen2 sen et al., 1982); HO #HO PH O #O ,
(10)
k "3.39;10exp(!2480/T ) l mol\ s\ (Christen2 sen and Sehested, 1988); &> H O #O , HO #O\ P
(11)
k "2.17;10exp(!914/T ) l mol\ s\ (Christen2 sen and Sehested, 1988). Because the solution was continuously stirred, it is assumed that H O was homogeneously irradiated. However, spectroradiometric measurements indicated that the actinic #ux of the lamp could vary (by as much as 25%) from one run to another. The variations were lower (less than 6%) during one continuous run of 9 h. Thus, the H O photolysis rate had to be evaluated for each experiment, which, therefore, was divided into three steps. In step 0, the aqueous solution (with no reactant) was irradiated in order to reach constant actinic #ux and temperature (O concentration was monitored). In step 1, a single aliquot of H O (2.3}11 mM) was introduced and H O photolysis was observed for 1 h (H O and O concentrations were monitored). In
step 2, a single aliquot of CH OH (0.23}0.31 mM) was introduced, and OH-oxidation of methanol was observed for 4 h (reactants and reaction products were monitored). Initial reactant concentrations were chosen so that k [CH OH]+k [H O ]. This allowed reaction (9) to be competitive, and thus, relatively high production of HO and O\ radicals, enabling reactions (6) and (7). 2.3. Analysis of reaction products The expected reaction products were formaldehyde, formic acid (formate ions) and HMHP. Formaldehyde in aqueous solution is in equilibrium with its hydrated form, methanediol, HCHO#H O&H C(OH)
(12)
with K "[H C(OH) ]/[HCHO]"4.9;10 (Bell, 1966). Therefore, we only consider hydrated formaldehyde in the following. Product analyses were performed using gas-phase chromatography (GC-FID), two highpressure liquid chromatography systems (HPLC -UV and HPLC -#uorimetry) and ionic chromatography (IC-conductimetry). The GC was equipped with a polar capillary column HP-INNOWAX. Solutions were analyzed for methanol by adding 300 ll from the vessel to an
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internal standard (1-butanol), prior to direct injection into the GC. HPLC was equipped with a reversed phase C column and UV detection at 360 nm. Solutions were analyzed for formaldehyde by removing 300 ll from the vessel and diluting it to 1.20 ml of 2,4-dinitrophenylhydrazine (2,4-DNPH, 0.16 g/l). Ionic chromatography was performed with a Dionex AS5 column, an AMMS suppressor column and a conductivity detector. Solutions were analyzed for formic acid (and formate ions) by removing 300 ll from the vessel and adding it to 4 ml of frozen ultrapure water. HPLC was equipped with a re versed phase C column. Following Hellpointner and GaK b (1989), parahydroxyphenyl acetic acid (and peroxidase) derivatization was used for detection of peroxides by #uorimetry (j "320 nm, and j "400 nm). Solutions were analyzed for H O and HMHP by re moving 20 ll from the vessel and diluting it to 2 ml of refrigerated (73C) ultrapure water. Dissolved O was continuously monitored by direct potentiometric measurements in the vessel. Except for the ionic chromatography (operated 12 h after each experiment, the samples being stored at !203C), all analyses were performed less than 2 h after sampling from the vessel. Each analytical technique gave rise to sharp peaks, and calibration of each compound (in the range covering the concentrations encountered in the experiments) gave retention times, statistical error limits and detection limits that are summarized in Table 1. 2.4. Reagents The reagents used were H O (not stabilized, Fluka, 30 wt%); CH OH and 1-butanol (Prolabo, normapur grade, 99.8%); N (I grade, Air Liquide); H PO (Pro labo, normapur grade); NaOH (Prolabo, normatom grade). Solutions were prepared using water puri"ed by an ELGASTAT maxima-HPLC system, including reverse osmosis, micro-"ltration, nuclear-grade deionization, activated carbon modules and an irradiation module at 254 nm.
2.5. Glass cleaning procedure Because we operated with diluted solutions, and in order to avoid any contamination, the reactor and all the sampling vessels were carefully cleaned using the following procedure: three rinses with reverse osmosed water; 3 h soaking in 2% of DECON diluted in reverse osmosed water; 10 rinses with ultrapure EGASTAT water; 3 h soaking in 2% of HCl diluted in ultrapure EGASTAT water; 10 rinses with ultrapure EGASTAT water; 12 h soaking in ultrapure EGASTAT water; three rinses with ultrapure EGASTAT water; drying in an ultraclean laminar #ow vent (class 10). 2.6. Data analysis Some tests by curve "tting were carried out using the FACSIMILE chemical modelling package (FACSIMILE, 1994).
3. Results Monitored reactants and products are presented in Fig. 2 as a function of time, for an experiment performed at pH"2. O pro"les clearly mimic the di!erent steps: during step 0, O decreases by evaporation because the surface of the solution is swept by pure N . At the beginning of step 1, O production by reaction (10) starts and dominates over evaporation. Just before step 2, because of the [H O ] decrease, the e$ciency of reaction (10) decreases, and O evaporation becomes competitive with reaction (10), leading to stabilization of [O ]. At the beginning of step 2, O consumption by reaction (2) starts and dominates over production by reaction (10). When 80% of methanol is consumed, the e$ciency of reaction (2) decreases, and O production by reaction (10) becomes competitive with reaction (2) leading to stabilization of [O ]. When all organic compounds are consumed (both methanol and reaction products), O production by reaction (10) dominates
Table 1 Calibration of each compound with each analytical technique Compound
Analytical technique
Retention time (min)
Detection limit (M)
Statistical error ($2.6p) (%)
CH OH H O HMHP CH (OH) HCOOH O
GC-FID HPLC (HPLC-Fl) HPLC (HPLC-Fl) HPLC (HPLC-UV) Ionic chromatography Potentiometry
2.93 5.0 5.3 6.0 8.3 *
10\ 10\ 10\ 10\ 5;10\ 3;10\
$10 $15 $15 $9 $10 $2.6
HPLC-Fl"high-pressure liquid chromatography equipped with #uorescence detection. HPLC-UV"high-pressure liquid chromatography equipped with UV detection.
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
5287
Fig. 2. Aqueous-phase OH-oxidation of methanol: reactant and product pro"les as a function of time at pH"2.
over reaction (2). At the end of the experiment, because initial [H O ] has decreased by a factor of 2 or more, O production by reaction (10) is in competition with O evaporation, which leads to stabilization of [O ]. 3.1. Reaction products 3.1.1. Hydrated formaldehyde and formic acid or formate ion At all pH levels studied, substantial amounts of hydrated formaldehyde, formic acid, and formate ion (both are designed as `formatea in the following) were detected (Fig. 3). Furthermore, `formatea pro"les indicate that this compound was produced directly from methanol without a stable molecular intermediate. Assuming that `formatea was exclusively produced from hydrated formaldehyde, one would have observed a null initial formation #ux of `formatea, and an in#ection on its pro"le when the maximum of hydrated formaldehyde concentration occurred. Fig. 3 shows that this assumption is not valid in our experiments, especially at pH"2 where the initial `formatea formation #ux was even higher than that of hydrated formaldehyde. 3.1.2. Hydroxymethyl hydroperoxide No detectable amount of HMHP was identi"ed at any pH value. However, using HPLC , the HMHP peak is
eluted close to the H O peak (Table 1). Calibration experiments have shown that these two compounds coelute when [HMHP]/[H O ])1/10. Thus, when initial H O concentrations were 2.3 mM (respectively, 11 mM), HMHP could not be detected under 0.23 mM (respectively, 1.1 mM). We performed several separate tests, which have shown that (i) HMHP response in HPLC is 100%; (ii) in the equilibrium (reaction (13)) of peroxidation of hydrated formaldehyde, reaction >13 is too slow in our conditions (279 K) to be signi"cant in our experiments; (iii) (reaction \13) is too slow to be shifted to hydrated formaldehyde formation in the presence of 2,4-DNPH prior to injection on HPLC : CH (OH) #H O & HOOCH OH#H O,
(13)
17($2);10 k> "exp 51($8)! (M\ s\) 2 ¹ (between 279 and 298 K),
2.1($0.3);10 K "exp !2.3($0.96)# (M\) 2 ¹ (between 279 and 298 K). Therefore, if HMHP was formed in our experiments, it was produced only from the studied reaction (OHoxidation of CH OH), and its concentration did not
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Fig. 3. pH e!ect on the reaction products.
exceed one-tenth of H O concentrations (i.e. less than 2}3;10\ M for experiments starting with 2}3;10\ M of H O and less than 1}1.1;10\ M for experiments starting with 10}11;10\ M of H O ).
hydrated formaldehyde formation rate was on average 4.5 ($1.0) times higher than that of `formatea, whereas, at `freea pH (i.e. initial pH"7 decreasing to 5 as the acid appeared), this ratio was 3.7 ($1.0), and at pH"2, this ratio decreased to 0.34 ($0.10).
3.2. pH ewect on reaction products
3.3. First step carbon balance
pH had a drastic e!ect on the initial #uxes of the identi"ed reaction products (Fig. 3). At pH"7, the
Because measurements of "nal products such as CO and CO could not be made, "rst step carbon balance
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5289
Table 2 Experimental results: initial reactant concentrations, percentages of reaction products formed relative to methanol consumption, and "rst step carbon balance pH
[H O ] (;10\ M)
[CH OH] (;10\ M)
% Hydrated formaldehyde formed
% `formatea formed
Carbon balance (%)
2 2 2 2 `freea `freea 7 7
3.0$0.4 2.7$0.4 11$2 3.2$0.5 2.3$0.3 2.6$0.4 2.8$0.4 10$2
2.4$0.2 2.3$0.5 3.1$0.4 3.0$0.3 2.5$0.2 2.2$0.5 2.3$0.2 2.4$0.5
11$2 8$2 15$3 14$2 43$7 35$11 31$5 45$13
49$9 31$10 37$8 36$7 10$2 12$4 9$2 8$2
60$11 39$12 52$11 50$9 53$9 47$15 40$7 53$15
calculations were performed. This method consisted in a comparison between methanol consumption and the primary stable products identi"ed (i.e. `formatea and hydrated formaldehyde) within the "rst 10}20 min of each experiment (Table 2). Only the very "rst few minutes of each experiment were taken into account, so `formatea and hydrated formaldehyde OH-oxidation was not important. The induced underestimate of the "rst step carbon balance was evaluated as follows: CH (OH) #OHPCH(OH) #H O, (14) k "2.4;10exp(!1020/T) l mol\ s\ (Chin and 2 Wine, 1994); HCOOH#OHPCOOH#H O, (15) k "3.0;10exp(!991/T ) l mol\ s\ (Chin and 2 Wine, 1994); HCOO\#OHPCOO\#H O, (16) k "2.0;10exp(!1240/T ) l mol\ s\ (Chin and 2 Wine, 1994). The "rst step carbon balance underestimate was calculated taking into account k /k , k /k and k /k ratios for our temperature conditions. For hydrated formaldehyde, the underestimate was 10% of the "rst step carbon balance percentage. For formic acid and formate ion, the analogous underestimates were 5 and 20%, respectively. All these underestimates are within the error bars indicated in Table 2. The average "rst step carbon balance was "nally 49$11%, indicating that half of CH OH removal was not identi"ed. 4. Discussion 4.1. CH3 OH conversion into OOCH2 OH radicals Was CH OH entirely converted into OOCH OH radicals? The answer to this question is essential before
discussing the fate of OOCH OH radicals. We evaluated the amount of RO radicals formed relative to methanol consumption. Stoechiometric coe$cients are equal to 1 in reactions (1) and (2). Thus, assuming steady state for CH OH radicals, one molecule of dissolved O should be consumed per consumed CH OH molecule, produ cing one OOCH OH radical. Therefore, comparing the amount of methanol consumed to the amount of dissolved O consumed, one can evaluate the amount of OOCH OH radicals formed relative to methanol con sumption. The evaluation of dissolved O consumed just after methanol injection (beginning of step 2) was carried out using the following equation: [O ] "[O ] ![O ] , (I) where [O ] is the measured concentrations in our experiments, and [O ] is the extrapolated concentra tions from the simulation of the `base linea of dissolved O concentrations during methanol reactivity (Fig. 4). [O ] was obtained by curve "tting both O and H O pro"les during step 1 (i.e. H O photolysis), while (reaction (8)) and also optimizing the optimizing J & - >FJ rate constant of O evaporation. An accurate curve-"t of O pro"les was obtained (Fig. 4) because of the low uncertainties in the measurements of its concentrations ($2.6%). We applied this analysis to the experiments performed at pH"2, 7 and `freea pH. Comparing the obtained [O ] to the consumed methanol (Fig. 5), we found that the ratio [consumed methanol]/[consumed O ] was 2.1$0.1. Therefore, the amount of OOCH OH radicals formed was only 48 ($2)% of the amount of CH OH consumed at all pH values. This result is in very good agreement with the "rst step carbon balance of 49 ($11)% at all pH values. We have checked in separate tests that methanol consumption by evaporation and by direct reaction with H O represented less than 5% of methanol consumption by OH-oxidation. Thus, 52 ($2)% of the CH OH radicals formed in reaction (1) must have undergone another pathway than
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Fig. 4. Experimental (exp) and simulated (simu) dissolved O pro"les at `freea pH. Simulation of step 1 (i.e. H O continuous photolysis) was performed by curve "tting of O and H O pro"les (see text). At t ("6307 s), methanol was injected into the vessel. At t ("14494 s), all the organic species (reactants and products) were consumed.
reaction (2). This is explained by the fact that dissolved O did not saturate the solution because of the continu ous N #ow at the surface of the solution. CH OH radicals were in their acid form in all our experimental conditions, because pK (CH OH/CH O\)"10.7 (Rabani et al., 1974). Following the "ndings by McElroy and Waygood (1991) on CH(OH) radicals, it is possible that the following reactions occurred: CH OH#H O PCH OH#HO (17) CH OH#H O PCH (OH) #OH. (18) However, these reactions would have been limited by reaction (9). CH OH self-reaction may also have occur red:
Fig. 5. Comparison between the amount of CH OH consumed and the amount of dissolved O consumed just after CH OH injection, at `freea pH (same experiment as shown in Fig. 3). Concentrations versus time (a); consumed methanol versus consumed dissolved O plot (b).
CH OH#CH OHPCH OH#HCHO. (19) Reformation of methanol by reaction (17) would not a!ect the "rst step carbon balance, thus, reaction (17) may have occurred, but it would not explain the elimination of CH OH radicals. Formation of hydrated formal dehyde by reacions (18) and (19) would a!ect the "rst step carbon balance. However, no in#uence of initial [H O ] was observed on the "rst step carbon balance (Table 2). If reaction (18) or (19) had occurred signi"cantly, the "rst step carbon balance would have been closer to 100%. It is probable that CH OH radicals have been eliminated from the solution by evaporation, and/or by addition of
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
HO radicals, leading to the formation of HMHP: CH OH#HO PHOOCH OH.
(20)
However, HMHP was not formed at detectable levels in our experiments (i.e. [HMHP](0.23 mM or [HMHP](1.1 mM, depending on initial concentrations of H O ). No further experiments were carried out to investigate the elimination of CH OH radicals because our goal was to examine the fate of OOCH OH radicals (atmospheric conditions). 4.2. The fate of OOCH2 OH radicals For all the reasons mentioned above, all the identi"ed products were attributed to OOCH OH radical reactivity. OOCH OH radicals were in their acid form in all our experimental conditions because pK (OOCH OH/OOCH O\) '8 (Rabani et al., 1974). 4.2.1. HMHP Compared to previous #ash photolysis studies, our experiments were carried out under conditions closer to the atmosphere, and under conditions more favorable to HMHP formation. However, insigni"cant amounts of HMHP were formed from OOCH OH radical reactiv ity. Therefore, reactions (6) and (7) were of negligible importance at all pH tested. 4.2.2. Primary `formatea formation Direct formation of formic acid from OOCH OH self-rection (reaction (5)) has been previously suggested (Bothe and Schulte-Frohlinde, 1980; Rabani et al., 1974; Huie and Clifton, 1993), and three di!erent mechanisms have been proposed for this reaction:
`formatea. This is clearly not the case, especially at pH"2 where `formatea formation rate is even higher than that of hydrated formaldehyde (Fig. 3). Thus, reaction (5) does not proceed by mechanism A. Assuming reaction (5) proceeds by mechanism B, the budget of reactions (1), (2) and mechanism B would result in the net consumption of two molecules of dissolved O per molecule of methanol. However, this is in contradiction with our experiments, where two molecules of methanol were consumed per molecule of O (Fig. 5). Finally, mechanism C does not contradict our experiments if one supposes that tetroxide (HOCH OO} OOCH OH) is so unstable that it could not be detected in HPLC . However, further experiments are needed to clarify the details of mechanism C; for example by generating OH radicals with another source (than H O ) and determining the amount of H O produced by C. 4.2.3. pH ewect The observed drastic pH e!ect on the amounts of reaction products was attributed to the catalytic e!ect of hydroxyl ions (reaction (4)). When the pH increased from 2 to 7, [OH\] increased from 10\ to 10\ M, thus the rate of decomposition of OOCH OH radicals via reac tion (4) increased to the detriment of the rate of OOCH OH self-reaction (5). This is in agreement with our experiments: comparing the pH"2 and 7 experiments, the proportion of hydrated formaldehyde formed relative to that of `formatea was inverse (Fig. 3 and Table 2). The ratio of hydrated formaldehyde formed at pH"7 to that formed at pH"2 was 3.2$1.3, which is equal to the ratio: +rate of reaction (4),/+rate of reaction (3),, thus enabling an estimate of k at pH"7 and at 279 K: k [OH\] "3.2$1.3. k
2OOCH OHP2OCH OH#O followed by 2OCH OHPHCOOH#CH (OH) or followed by
(A)
OCH OH#O PHCOOH#HO , or: 2OOCH OH[PHOCH OO}OOCH OH]
(B)
2HCOOH#H O . (C) Assuming reaction (5) proceeds by mechanism A, equal amounts of `formatea and hydrated formaldehyde would be formed from this sequence and an additional amount of hydrated formaldehyde would be generated from reactions (3) and (4). Therefore, in our experiments, one would have detected the formation of hydrated formaldehyde in equal or in higher amounts than that of
5291
(II)
Considering that k (298 K)(10 s\ (Bothe et al., 1977, 1978), thus, at 279 K, and pH"7, k (279 K)(3.2 ($1.3);10 l mol\ s\. Comparing this value with k (295 K)"18;10 l mol\ s\ (Neta et al., 1990), it is clear that k is highly exothermic. These considerations give an estimate of the lower value of the activation energy of reaction (4) of the order of 172 kJ mol\. Thus, temperature should have an important in#uence on the fate of OOCH OH radicals, especially at high pH levels. 4.3. Implications for atmospheric chemistry The results of our experiments have shown that the presence of relatively high levels of HMHP in some atmospheric water droplets (up to 1 lM in cloudwater,
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Sauer et al., 1996) is not the consequence of aqueousphase photochemistry of OOCH OH radicals. HMHP is more probably formed at the droplet interface from Criegee biradicals (GaK b et al., 1995), which are produced in the gas phase by ozonolysis of alkenes. Further, the present investigation indicated that aqueous-phase OHinitiated oxidation of methanol does not solely lead to hydrated formaldehyde, it also forms formic acid, with no stable molecular intermediary, by self-reaction of OOCH OH radicals. Therefore, the importance of this reaction is sensitive to OOCH OH concentrations, and thus to initial methanol concentration. However, compared to atmospheric conditions, we have operated with relatively high concentrations of methanol. To explore whether reaction (5) is still signi"cant under atmospheric conditions, the mechanism of OOCH OH fate elucidated here was included in a cloud box model (Monod and Carlier, 1999). This box model was initialized with gas-phase methanol concentrations similar to those found in "eld measurements, and it was run at 279 K, similar to our experimental conditions. The rate of formation of `formatea relative to that of the sum of hydrated formaldehyde and `formatea from aqueousphase OH-oxidation of methanol (ratio R) has been calculated:
base-catalyzed unimolecular decomposition of OOCH OH radicals. However, compared to realistic atmospheric conditions, only a limited number of parameters have been investigated here. It is known that transition metals play an important role in aqueous radical chemistry (Walcek et al., 1997). Moreover, dissolved organic RO radicals may interact with transition metals (Graedel et al., 1986). To our knowledge, however, no experimental evidence has yet been carried out on these interactions under atmospheric conditions. Therefore, further experimental and modelling studies are needed to assess the in#uence of temperature, ionic strength, and the presence of transition metals on the behavior of RO radicals within the aqueous phase of atmospheric water droplets. 5. Conclusions Analysis of reaction products of aqueous-phase OHoxidation of methanol revealed that hydrated formaldehyde is not the only stable primary reaction product. Formic acid and/or formate ion were also identi"ed as stable primary molecular reaction products. The branching ratios for the formation of the products were highly
2k [OOCH OH] R (%)" ;100. 2k [OOCH OH]#k [OOCH OH]#k [OOCH OH][OH\] The in#uence of pH and initial gas-phase concentrations of methanol on R was tested (Table 3). Table 3 shows that R is much more sensitive to pH values than to initial gas-phase concentrations of methanol. Moreover, except at pH 6, `formatea formation is signi"cant. Therefore, the formation of primary `formatea from OH-oxidation of methanol has to be taken into account in cloud models. The present study has shown that the aqueous-phase chemistry of OOCH OH radicals is strongly in#uenced by pH. The in#uence of changing concentrations of the parent compound is less signi"cant (Table 3). This study also suggested that temperature has a strong in#uence on
pH dependent. At pH"7, hydrated formaldehyde was the dominant molecular reaction product (ratio 4.5 : 1 for hydrated formaldehyde : formate ion), whereas at pH"2, formic acid was the dominant product (ratio 3.7 : 1 for formic acid : hydrated formaldehyde). At all pH values studied, the sum of the primary stable products represented 49 ($11)% of methanol removal, in agreement with the amount of OOCH OH radicals formed relative to methanol removal (48 ($2)%). The formation of primary formic acid at pH"2 was attributed to OOCH OH self-reaction, and the strong pH e!ect
Table 3 Cloud box model results (including both gas and aqueous phases): in#uence of pH and initial gas-phase concentration of methanol on R pH
R (%) (initial [CH OH ]"4.3 ppbV)
R (%) (initial [CH OH ]"10 ppbV)
3 4.16 5.2 6
39 21 5 0.43
48 29 7.5 0.67
Rate of formation of `formatea relative to that of the sum of hydrated formaldehyde and `formatea from aqueous-phase OHoxidation of methanol: 2k [OOCH OH] R (%)" ;100. 2k [OOCH OH]#k [OOCH OH]#k [OOCH OH][OH\]
A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
was attributed to the base-catalyzed decomposition of OOCH OH leading to the formation of hydrated formaldehyde. Evaporation and/or an addition reaction between CH OH and HO radicals leading to the formation of hydroxymethyl hydroperoxide was proposed to explain the missing yields. However, hydroxymethyl hydroperoxide could not be detected under 0.23 mM or under 1.1 mM, depending on H O concen trations. All the experimental conditions were taken as close as possible to atmospheric conditions, except for the initial concentrations of reactants, especially those of methanol. The latter were taken at 0.22}0.23 mM in order to obtain detectable amounts of reaction products. Although such concentrations are supposed to be 10}100 times higher than in atmospheric droplets, these are much more dilute conditions (by a factor '1000) than in previous experimental studies on OH-oxidation of methanol. To explore whether the two identi"ed OOCH OH pathways (de composition and self-reaction) were signi"cant under atmospheric conditions, the mechanism elucidated was included in a cloud box model comprising more realistic concentration and pH conditions than in our experiments. At typical pH values (pH"4}5) encountered in clouds, both pathways were signi"cant. It was concluded that under atmospheric conditions, aqueous-phase OHinitiated oxidation of methanol does not lead solely to formaldehyde, but also to primary formation of formic acid, and this should be included in cloud models. However, only a limited number of parameters have been investigated, and further experimental and modelling studies are needed to assess the in#uence of temperature, ionic strength, and the presence of transition metals species on the behavior of RO radicals within the aque ous phase. Because these parameters may have di!erent in#uences depending on the chemical nature of RO , di!erent parent compounds also need to be studied (hydroxycarbonyls, ethers, esters, peroxides, etc.). Acknowledgements The authors thank `la feH deH ration des industries de la parfumeriea for the "nancial support of this study, and Dr. J.F. Doussin, Dr. J. Colman and Dr. M. McGrath for helpful comments on the manuscript. References Asmus, K.D., MoK ckel, H., Henglein, A., 1973. Pulse radiolytic of site of OH radical attack on aliphatic alcohols in aqueous solution. Journal of Physical Chemistry 77 (10), 1218}1221. Bell, R.P., 1966. The reversible hydration of carbonyl compounds. In: Gold, V. (Ed.), Advances in Physical Organic Chemistry, Vol. 4. Academic Press, London, pp. 1}27.
5293
Bothe, E., Behrens, G., Schulte-Frohlinde, D., 1977. Mechanism of the "rst order decay of 2-hydroxy-propyl-2-peroxyl radicals and of O\ formation in aqueous solution. Zeitschrift fur Naturforschung 32b, 886}889. Bothe, E., Schuchmann, M.N., Schulte-Frohlinde, D., von Sonntag, C., 1978. HO elimination from a-hydroxyalkylperoxyl radicals in aqueous solution. Photochemistry and Photobiology 28, 639}644. Bothe, E., Schulte-Frohlinde, D., 1978. The bimolecular decay of a-hydroxymethylperoxyl radicals in aqueous solution. Zeitschrift fur Naturforschung 33b, 786}788. Bothe, E., Schulte-Frohlinde, D., 1980. Reaction of dihydroxymethyl radical with molecular oxygen in aqueous solution. Zeitschrift fur Naturforschung 35b, 1035}1039. Buxton, G., Greenstock, C.L., Helman, W.P., Ross, A., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atom and hydroxyl radicals (OH/O\) in aqueous solution. Journal of Physical Chemistry Reference Data 17 (2), 513}883. Chin, M., Wine, P.H., 1994. A temperature dependent competitive kinetics study of the aqueous-phase reaction of OH radicals with formate, formic acid, acetate, acetic acid, and hydrated formaldehyde. In: Helz, G.R., Zepp, R.G., Crosby, D.G. (Eds.), Aquatic and Surface Photochemistry. Boca Raton, pp. 85}98. Christensen, H., Sehested, K., 1988. HO and O\ radicals at elevated temperatures. Journal of Physical Chemistry 92, 3007}3011. Christensen, H., Sehested, K., Cor"tzen, H., 1982. Reactions of hydroxyl radicals with hydrogen peroxide at ambient and elevated temperatures. Journal of Physical Chemistry 86, 1588}1590. Demerjian, K.L., Schere, K.L., Peterson, J.T., 1980. Theoretical estimation of actinic #ux and photolytic rate constants of atmospheric species in the lower atmosphere. Advances in Environmental Science and Technology 10, 369}459. Elliot, A.J., McCracken, D.R., 1989. E!ect of temperature on O\ reactions and equilibria: a pulse radiolysis study. International Journal of Applied Radiation and Instrument C 33 (1), 69}74. Facsimile, 1994. User Guide v.4.0, copy 4013. AEA Technology, Harwell, Didcot, Oxford. GaK b, S., Turner, W.V., Wol!, S., Becker, K.H., Ruppert, L., Brockmann, K.J., 1995. Formation of alkyl and hydroxyalkyl hydroperoxides on ozonolysis in water and in air. Atmospheric Environment 29 (18), 2401}2407. Graedel, T.E., Mandich, M.L., Weschler, C.J., 1986. Kinetic model studies of atmospheric droplet chemistry 2. Homogeneous transition metal chemistry in raindrops. Journal of Geophysical Research 91 (D4), 5205}5221. Hellpointner, E., GaK b, S., 1989. Detection of methyl, hydroxymethyl and hydroxyethyl hydroperoxides in air and precipitation. Nature 337, 631}634. Huie, R.E., Clifton, C.L., 1993. Kinetics of the self-reaction of hydroxymethylperoxyl radicals. Chemical Physics Letters 205 (2,3), 163}167. Jacob, D.J., 1986. Chemistry of OH in remote clouds and its role in the production of formic acid and peroxymonosulfate. Journal of Geophysical Research 91 (D9), 9807}9826. Jacob, D.J., 2000. Heterogeneous chemistry and tropospheric ozone. Atmospheric Environment 34, 2131}2159.
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A. Monod et al. / Atmospheric Environment 34 (2000) 5283}5294
Jacob, D.J., Gottlieb, E.W., Prather, M.J., 1989. Chemistry of a polluted cloudy boundary layer. Journal of Geophysical Research 94, 12975}13002. Lelieveld, J., Crutzen, P.J., 1991. The role of clouds in tropospheric photochemistry. Journal of Atmospheric Chemistry 12, 229}267. McElroy, W.J., Waygood, S.J., 1991. Oxidation of formaldehyde by the hydroxyl radical in aqueous solution. Journal of the Chemical Society, Faraday Transactions 87 (10), 1513}1521. Monod, A., Carlier, P., 1999. Impact of clouds on the tropospheric ozone budget: direct e!ect of multiphase photochemistry of soluble organic compounds. Atmospheric Environment 33 (27), 4431}4446. Neta, P., Huie, R.E., Ross, A.B., 1990. Rate constants for reactions of peroxyl radicals in #uid solutions. Journal
of Physical and Chemical Reference Data 19 (N2), 413}498. Pandis, S.N., Seinfeld, J.H., 1989. Sensitivity analysis of a chemical mechanism for aqueous-phase atmospheric chemistry. Journal of Geophysical Research 94, 1105}1126. Rabani, J., Klugh-Roth, D., Henglein, A., 1974. Pulse radiolytic investigations of OHCH O radicals. Journal of Physical Chemistry 78 (21), 2089}2093. Sauer, F., Schuster, G., SchaK fer, C., Moortgat, G.K., 1996. Determination of H O and organic peroxides in cloud-and-rain water on the Kleiner Feldberg during Feldex. Geophysical Research Letters 23 (19), 2605}2608. Walcek, C.J., Yuan, H.H., Stockwell, W.R., 1997. The in#uence of aqueous-phase chemical reaction on ozone formation in polluted and nonpolluted clouds. Atmospheric Environment 31 (8), 1221}1237.