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Appendix A The Chemistry of Acid Formation It is well established that the oxides of sulfur--sulfur dioxide (SO2) and sulfur trioxide (SO3)--and nitrogen-- nitric oxide (NO) and nitrogen dioxide (NO2)--are converted (oxidized) in the troposphere to sulfuric acid (H2SO4) and nitric acid (HNC3), respectively. However, the most prevalent end products of these reactions--H2SO4, HNO3, ammonium bisulfate (NH4HSO4), ammonium nitrate (NH4NO3), etc.--give few clues about which of several oxidizing pathways are important. Yet for the develop- ment of scientifically sound, predictive models of acid deposition that define theoretical source-receptor rela- tionships r knowledge of the elementary chemical steps that are involved (among many other quantities such as deposition rates and transport rates) is required. The rates of these reactions follow well-defined mathematical laws (rate expressions) that relate reaction rates to concentrations of the reactants, temperature, pressure, and other variables. Considerable progress has been made in developing an understanding of the nature of this chemistry, but a large number of uncertainties still remain in key areas. As our knowledge of the atmospheric chemistry of SO2, NO, and NO2 continues to grow, it has become increasingly clear that many different pathways exist for generating H2SO4 and HNO3 in the troposphere. Reactions can occur in the gas phase; in the solution phase in cloud water and rainwater, for example; and in reactions on surfaces of solid particles in the atmosphere. Thus we are concerned with the rates of exchange of gaseous reactants and their reaction products between liquids and the surfaces of solids as well as the rates of interactions among gaseous molecules, aqueous phase molecules and ions, and species adsorbed on solid 155
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156 surfaces. For current purposes, it is not necessary to review and evaluate every detail of these chemical processes, but we should note the nature of the many chemical processes, the variety of interactions among the many species involved, and the current state of knowledge related to the chemistry of acid deposition. We consider first the homogeneous gas-phase chemistry that results in oxidation of SO2, NO, and NO2 in the troposphere (Calvert and Stockwell 1983, Calvert et al. 1978). GAS-PHASE REACTIONS LEADING TO GENERATION OF ACID IN THE TROPOSPHERE Oxidation by Stable Atmospheric Molecules The thermodynamic properties of the oxides of sulfur indicate that sulfur dioxide has a strong tendency to react with oxygen in the air under normal tropospheric conditions: 2SO2 + O2 ~ 2SO3. (1) Thus the ratio of [SO3]/tSO2] is about 8 x 1011 at equi- librium in air at 1 atm and 25°C. (Square brackets indicate the concentration of the species inside the brackets.) Thermodynamic arguments also tell us that at humidities normally encountered in the lower troposphere, the SO3 produced by reaction (1) will be converted efficiently to sulfuric acid, H2SO4(aq), according to SO3 + H2O ~ H2SO4(aq). (2) The reaction of SO3 with H2O is so fast that any process in which SO3 is formed in the moist troposphere can be considered equivalent to the formation of H2SO4. Certain metal ions (Mn2+, Fe3+, etc.) in aqueous solutions of SO2(HSO5) can catalyze the overall sequence of reactions (1) and (2). However, thermo- dynamics tells us nothing about the rates of chemical reactions, and the rate of reaction (1) is so slow under catalyst-free conditions in the gas phase that it can be neglected as a source of sulfuric acid in the atmosphere. The thermal oxidation of NO and NO2 in the gas phase is also slow. Pathways that are thermodynamically favored are
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157 2NO + O2 + 2NO2, 2NO2 + H2O ~ HNO3 + MONO. (3) (4) Reaction (3) requires literally days for the conversion of a significant fraction of the nitric oxide present [at concentrations of the order of parts per billion (ppb)] in the troposphere, and the homogeneous gas-phase reaction (4) is immeasurably slow (Schwartz and White 1982). Obviously, other, seemingly less direct reactions must be invoked to account for the observed rates of SO2, NO, and NO2 oxidation, which are between 1 and 100 percent/in. A number of the more complex pathways involve photo- chemistry. In one, sulfur dioxide absorbs light in the ultraviolet region of the solar radiation incident in the troposphere, and, in principle, excited states of SO2 generated in this fashion could lead to SO2 oxidation in the troposphere. Figure A.1 indicates that there is a significant overlap between the actinic flux incident in the lower troposphere and two distinct absorption regions of SO2. Excitation in the "forbidden," long-wavelength band forms the excited SCt(3B1) species, while excitation at the wavelengths below 330 nm generates higher excited states, presumably the 1A2 and 1B1 species. SO2(XlAl) + hv(340 < A < 400 nm) ~ So2(3Bl), SO2(XlAl) + hv(240 < ~ < 330 nm) ~ SO2(1A2, 1Bl). (6) These excited states of SO2 are nondissociative; only quanta of light at wavelengths below 218 nm (which do not penetrate to the troposphere) provide sufficient energy to allow photodissociation: ~1 SO2(X*A1) + hv(A < 218 nm) ~ S02(1B2) ~ 0(3P) + So(3£ ) e (7) The lower excited singlet states of SO2(1A1, 1Bl) appear to be very short lived in air at 1 atm, and they are rapidly converted by collisional perturbations to S02(3B1) molecules, possib~ SO2(3A2) and SO2(3Bk) molecules, and ground-state S02(X A1) molecules. The rate of excitation of SC2 through absorption of sunlight can be very sig- nificant. If this excitation were the rate-determining step in the photooxidation of SO2, that is, if every molecule of SO2 that is photoexcited were oxidized
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158 400 _ 300 c, . _ cat - c' to 0 200 - c~ - x 111 100 _ O 240 1 1 1 " ~.~ 1 - 320 Wavelenoth (nary) 0.10 0.08 A, c . _ . _ - o 0.04 _ C1 0.02 400 FIGURE A.1 Comparison of the extinction coefficients (liters mole~1 cm~i, base 10) of SO2 within the first allowed band (left), the "forbidden" band (right), and a typical distribution of the flux of solar quanta (relative) at ground level (dashed curve). SOURCE: Calvert et al. (l978~. through subsequent reaction with O2 or other reactants, then the lifetime of SO2 in the lower troposphere with overhead sun should be as low as 52 minutes (Sidebottom et al. 1972). Of course, this is not the case. The SO2(3B1) species appears to be one of the most favored states, which is ultimately populated through absorption of sunlight and collisional processes in the lower atmosphere. The reactions of this species with various atmospheric gases and many atmospheric impurities have been studied extensively (Calvert et al. 1978). Quenching of SO2(3B1) by atmospheric gases is expected to be the dominant process. In air at 1 atm, 25°C, and 50 percent relative humidity, quenching by N2, 02, H2O, and Ar will occur 45.7, 41.7, 12.2, and 0.3 percent of the time, respectively. Quenching by impurity gases is highly improbable. Even when impurities are present at concentrations of the order of ppm, the rates of SO2 conversion by such species are very slow (Calvert et al. 1978). All available evidence suggests that the only . . . . . . significant chemical result of the major quenching reactions of SO2(3B1) occurs with O2, and this does not lead efficiently to any overall chemical change in the SO2, but low-lying, excited electronic states of molecular oxygen, O2(1Ag) and O2(1£+g), are formed by energy transfer.
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159 SO2 ( B1) + O2 ( Eg) + SO2 (X A1) + O2 ( fig) ~ SO2(2 Al) + O2( Ag) (8) . (9) These reactions are not unique sources of excited oxygen, species that are also formed in other atmospheric reactions at much higher rates (Calvert et al. 1978). Thus we conclude that photooxidation of SCt is also not an important source of acids in the atmosphere. Similar conclusions are reached from a study of the photochemistry of NO and NO2. NO does not absorb solar radiation in the wavelength range available near the Earth's surface, so its photochemistry is not important in the troposphere. NO2 does absorb radiation over a wide range, and absorption of the wavelengths below 430 nm leads to dissociation (NO2 + he ~ O + NO). However, formation of acid as a direct result of the photochemistry . of NO and NO2 is unimportant. Reactive Transient Species in the Troposphere Most of the gas-phase tropospheric chemistry of SO2, NO, NO2, and other impurity molecules involves reactions with a variety of reactive excited molecules, atoms, and free radicals (neutral fragments of stable molecules) formed by absorption of sunlight by trace gases in the atmosphere. As a background to discussions to come we review briefly here some of this important chemistry since it enters directly or indirectly into many of the reaction pathways that lead to the formation of acids in the troposphere. In the polluted troposphere, NO2 is dissociated by sunlight absorption (A < 430 nm) to form reactive, ground state oxygen atoms, O(3P), and NO, while the oxygen atom reacts rapidly to form ozone (O3): NC: + hv(A < 430 nm) + O(3P) + NO, O(3P) + O2(+M) ~ O3(+M). Ozone can reoxidize NO to NC2 in (12) or react with alkenes to give highly reactive ozonides in (13) and Criegee intermediates in (14): (10) Ill)
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160 0 3 + NO ~ O2 + NO2 0 - 0 - 0 O3 + RHC=CHR ~ RHC- CHR 0 - 0 - 0 1 1 RHC - CHR ~ RCHO2 + RCHO Ozone can also oxidize NO2 to the reactive transient NO3 in (15), and this can lead to N2O5 in (16): O3 + NO2 + O2 + NO3, Net + Net (+M) By N2O5 (+M) . (12) (13) (14) (15) (16) The photodecomposition of ozone may generate electroni- cally excited oxygen atoms, O(1D), and excited molecular oxygen with absorption in the short-wavelength region of the spectrum: O3 + hv(290-306 nm) ~ O(1D) + O2( fig), O3 + hv(290-350 nm) ~ O(1D) + O2 or O + 02(1Ag' 3£g)' O3 + hv(450-700 nm) ~ O + O2. (17) ( 18) (19) The O(1D) species formed in (17) is much more reactive than the ground-state oxygen atoms [O(3P), often simply symbolized by O]. O(1D) reacts efficiently when it collides with a water molecule to form a highly important transient in atmospheric chemistry, the hydroxy radical, HO: O (ID) + H2O ~ 2HO. (20) This radical, unlike many molecular fragments formed from carbon-containing molecules, is unreactive toward oxygen, and it survives to react with most atmospheric impurities such as the hydrocarbons, aldehydes, NO, NO2, SO2, and CO. Its reactions with carbon monoxide and the hydro- carbons (RH) lead to another important class of reactive transients, the peroxy radicals: HO + CO ~ H + CO2, (21)
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161 H + O2(+M) + HO2(+M), HO + RH ~ H2O + R. R + O2(+M) ~ RO2(+M). (22) (23) (24) Here R represents the alkyl groups such as methyl (CH3), ethyl (C2H5), or another larger group derived from the parent hydrocarbons, methane (CH4), ethane (C2H6), or larger hydrocarbon (RH), respectively. The reaction of the HO radical with aldehydes (RCHO) forms the acyl (RCO) and acylperoxy (RCO02) radicals in similar reactions: RCHO ~ HO ~ RCO + H2O, RCO + O2(+M) ~ RC002(+M). (25) (26) The peroxy radicals react rapidly with NO to form NO2 and other classes of reactive species. In the case of the HC2-No reaction, HO is regenerated, while with the RO2 and SCOW radicals, alkoxy (RO) and acyloxy (RCO2) radicals, respectively, are formed: HO2 + NO ~ HO + NO2, RO2 + NO ~ RO + NO2, RCOO2 + NO ~ RCO2 + NO2 (27) (28) (29) The most common fate of the smaller alkoxy radicals in the lower atmosphere is reaction with oxygen, leading to HO2 radicals and a carbonyl compound. For example, with the simplest alkoxy radical, methoxy (CH3O), the following reaction occurs: CH3O + O2 + HO2 + CH2O (30) The RCC: radicals are of short lifetime, decomposing to form an alkyl radical (R) and CO2, with the subsequent generation of another peroxyalkyl radical: RCO2 ~ R + CO2, R + O2(+M) ~ RO2(+M) (31) (32) Reactions (14)-(26) combined with reactions (27)-(32) form a chain reaction. That is, a single initial HO radical
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162 may oxidize CO, hydrocarbon, or aldehyde, and additional HO radicals will be regenerated as NO is oxidized to N02; the subsequent steps occur again and again in a repeating cycle of events. The peroxy radicals and ozone are the principal oxidizing agents for NO in the lower atmosphere [reactions (IS) and (27)-(29)] (Demerjian et al. 1974). When peroxy radicals react with NO2, an additional class of highly reactive compounds is generated--the peroxynitrates: HO2 + NC2 ~ HC2NC2, RO2 ~ NO2 ~ RO2NO2 ~ RCOO2 + NO2 ~ RCOO2 NO2 Peroxynitric acid formed in (33), and the alkyperoxy- nitrates formed in (34) are relatively unstable in the lower troposphere at temperatures common in summer months; they dissociate readily to reform peroxy radicals and NC2. However, during the cold winter months or in the stratosphere they can act as temporary sinks for HO2 and RO2 radicals and NO2. Peroxyacylnitrates (RC ~ NO2), of which peroxyacetylnitrate (CH3COCtNC2) is the most common, have longer lifetimes and can be the source of radical generation even during the nighttime hours. The excited singlet delta molecular oxygen, O2(1Ag), product of reactions (17) and (18), and excited singlet sigma oxygen, O2(1~+)' the product of reaction (8), can also be created by ~irect absorption of sunlight by atmospheric O2 and by energy transfer reactions from other photoexcited species such as NO2(\ > 430 nm) and excited triplet aromatic hydrocarbons. (33) (34) (35) For current purposes the complex array of interactions that occur among the reactive species outlined here and with the various atmospheric impurities need not be considered. It suffices to say that many aspects of tropospheric chemistry, including photochemical "smog" and ozone generation, depend on these happenings (Demerjian et al. 1974). It is important, however, to evaluate the potential significance of the many highly reactive transient species of the atmosphere for reactions that oxidize SO2 and NO2 to acids.
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163 Atmospheric Oxidation of SO2 by Reactive Transient Species There are a large number of potentially significant gas-phase reactions of the reactive transients leading to oxidation of SC: in the troposphere. The potential candidate reactions, summarized in Table A.1, have the thermodynamic potential to occur as measured qualitatively by the sign of the change in enthalpy (aH°) for the overall reaction. Rate constants for most of these elementary reactions have been determined. These data, coupled with estimates of the concentrations of the transients in the atmosphere, allow us to evaluate the significance of each reactant in oxidizing SO2 in the atmosphere. Such evaluations have shown that reactions (42), (50), (56), and (59) or (61) are potentially sig- nificant sources of SO2 oxidation; some of these reactions are only important for certain peculiar atmospheric conditions. By far the most important of the gasphase reactions is the reaction of HO radicals with so2 HO + SO2(+M) ~ HOSO2(+M) . (56) The rate-constant data for reaction (56) have been reviewed recently (Calvert and Stockwell 1983) and are summarized in Figure A.2. The theoretical effect on this rate constant of the altered temperature and pressure of the atmosphere at various altitudes is shown in Figure A.3. The "best" values of kII(56) suggested from the analysis given by Calvert and Stockwell (1983) are somewhat larger than those chosen by Moortgat and Junge (1977), Zellner (1978), and the values recommended for use by the CODATA (1980) group and the NASA (1979) panel and are more consistent with currently available infor- mation on the HO-SO2 reaction. One must, however, retain considerable pessimism about the accuracy of all these data; realistic confidence limits should include +50 percent of the suggested value. Evidence is good that HOSO2 formed in reaction (56) ultimately leads to the generation of sulfuric acid aerosol. However, HOSO2 is not a stable molecule; it is a free radical that is probably highly reactive with several atmospheric compounds. It is not now clear what elementary reaction conversion to H2SO4. pathways are important in its Although there has been a great
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164 TABLE A. I Enthalpy Changes and Recommended Rate Constants for Potentially Important Reactions of Ground State SO~ and SO~ Molecules in the Lower Atmosphere - -~H ~kh Reaction (kcal/mole) (25°C) (cm3/molec-sec) (36) O. ('~g) + SO. ~ SO4 (biradical; cyclic) ~25, ~28 (37) O. ('Ax) + SO. ~ SO,3-O (3P) -13.5 3.9 x l0-2° (3X) O. (~Ag) + SO.-O. (~5g) + SO. 22.5 (39) O. (~5g) + SO2-SO4 (biradical: cyclic) ~40, ~43 (40) O. ('Ig) + SO.-SO~ + O (3P) 1.5 6.6 x 10-~6 (41) O2 (~g) + SO.-SO. + O. ('/`,~) 15.0 (42) O (3P) + SO~ (+M)-SO~ (+M) 83.0 5.7 x 10-~4 (43) O3 + SO2-O. + SO; 57.6 <8 x 10-24 (44) NO2 + SO2-NO + SO3 9.9 8.8 x 10-3° (45) NO,3 + SO2-NO2 + SO3 32.6 <7 x 10-2' (46) ONOO + SO.-NO. + SO,3 ~30 <7 x 10-2' (47) N2O~ + SO2 ~ N.O4 + SO3 24.0 <4 x 10-23 (48) HO. + SO'-HO + SO3 16.7 ) <1 x 10-'8 (49) HO. + SO. (+M)-HO.SO2 (+M) ~7 (50) CH3O2 + SO2-CH3O + SO3 ~27 <1 x 10-'8 (51) CH,;02 + SO2 (+M) _ CH3O2SO2 (+M) ~31 ~1.4 x 10-'4C (52) (CH,3),~CO2 + SO2-(CH3)3CO + SO,3 ~26 / <7.3 x 10-~9 (53) (CH~),~CO2 + SO2 _ (CH,~)3CO2SO2 ~30 (54) CH,~COO. + SO2-CH3CO. + SO,1 ~33 <7 x 10-19 (55) CH,~COO2 + SO2-CH3COO,SO2 ~37 (56) HO + SO2 (+M) ~ HOSO2 (+M) ~37 1.1 x lo-~2 (57) CH30 + SO2 (+M)_ CH3OSO2 (+M) ~24 ~5.5 x 10-13 Q~ ,0 (58) RCH-CHR + SO2 ~ 2RCHO + SO3 ~69 See text Q o,-O RCH-CHR + SO2-2RCHO + SO3 ~89 See text (59) RCHOO- + SO2_ RCHO + SO,~ ~79 k59/k60~ 6 x 10-5 (60) RCHOO- + H2O-RCOOH + H2O ~121 (R = CH3) (61) RCHO ~ SO2-RCHO + SO3 ~58 k6l/k62 ~ 4 Q Q (62) RCHO + CH2O ~ RCHOCH2O ~12 (63) SO3 + H2O ~ H2SO4 24.8 (R = H) , ,, 9.1 x 10 'Enthalpy change estimates were derived from the data of Benson (1978), Harding and Goddard (1978), and Domalski ( 1971). ~Rate constants are expressed as second-order reactions for I atm of air at 25°C; see Calvert and Stockwell (198,3) for the references to the original literature. ''The reverse reaction is so fast that the rate of oxidation of SO2 via (51) is very dependent on alternate fates of the CH3O2SO2 species.
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165 1 000 E ~ 00 - 10 ~ 1 ~ 1 1 0.01 0.1 _ I 1--l rilill r -Al l-:lil --I 1-l 1 IrTl ~-r I I Al ITT: C Corrected data of Castleman and Tang (1976-1977) O Leu (1982) a Range of values reported by Harris and Wayne (1975) Cox and Sheppard (1980) O Davis et a I. ( 1979) D' ,/ o i I I I l_L / 1/[M] (cm3/molec x 1018) 3' new , MU _4 ~ J /^ / / / I 1 1 1 1 1 11 1 1 1 1 1 1 1 1 1 1 1 1.0 1n 100 10 FIGURE A.2 Comparison of the experimental data for the effective second-order rate constants for the reaction (56) with M = N2. SOURCE: Calvert and Stockwell (1983~. amount of speculation in this regard, there is little experimental evidence to help determine the relative importance of suggested alternative routes. Thus Cox (1974-1975, 1975), Calvert and McQuigg (1975), Calvert et al. (1978), Davis et al. (1979), Friend et al. (1980), Leu (1982), Benson (1978), and others have suggested that the HOSO2 radical may participate in a variety of radical-radical and radical-molecule reactions. These reactions are summarized in Figure A.4. Although the mechanism of H2SO4 generation following (56) is unclear today, it is probable that reaction (56) is the rate- determining step in the sequence. Recent evidence suggests that the concentration of HO in photooxidizing mixtures of MONO, NO, NO2, and CO is insensitive to even large additions of SC2 (Stockwell and Calvert 1983). Thus the following sequence of reactions seems favored:
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191 104 103 1o2 - ~ 10 a) Q - LL 1 CC A O 10 0 10 2 10-3 10-4 10-5 H 2O2 M// / / / Mn: /NO2 0 1 2 /NO2 3 4 5 6 Gas-Phase Concentrations (ppb) H2O2 1 O3 50 HNO2 2 NO2 1 Liquid-Phase Concentrations Fe3+ 3 x 10-7 M Mn2+ 3 x 10~ M C 1x10~2g/liter pH FIGURE A.13 Theoretical rates of liquid-phase oxidation of SO2 assuming 5 ppb of SO2, 1 ml/m3 of water in air, and concentrations of impurities as shown. SOURCE: Martin (1983~. For the hydrated NO2 oxidation, Martin used the Henry's law constant and rate constants of Schwartz and White (1982) and Lee and Schwartz (1981). The concen- trations of NO2, MONO, and NO are related thermo- dynamically, and so they may not be regarded as completely independent. The trend of rising conversion rates with increased pH results from either the rising equilibrium concentration of sulfur (IV) or the sensitivity of the rate constants to pH, or both. The relative conversion rate increases with rising concentrations of sulfur (IV) in this range of pH values and for this amount of liquid. H2O2, which
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192 undergoes an acid-catalyzed reaction, is the only oxidant for which the rate dependence on [H+] compensates for the decreased solubility of SO2 with increased [H+]; hence the rate of oxidation by H2O2 is nearly independent of pH. Martin emphasized that the relative positions of the curves in Figure A.13 differ for differing assumptions about the composition of the droplets. For example, an aerosol with less liquid water will, as a rule, have higher concentrations of nonvolatile species such as carbon, iron, and manganese but a smaller reaction volume, so the relative positions of the curves may differ in this situation. The relative positions will also be different at night, when the concentration of photochemically derived oxidants drops. At temperatures below 25°C, the rate constants are lower in accordance with the activation energies of the given reactions, and the Henry's law coefficients are higher. The two effects act in opposite directions. In some cases, such as that for H2O2 oxidation, the net rate rises as the temperature falls (at constant gas-phase concentration). In other cases, such as for iron catalysis, the net rate falls with temperature. The concentrations of reactants used in deriving Figure A.13 are representative of those that might be anticipated in the atmosphere. Oxidation by H2O2 dominates all reactions for conditions of low pH. Oxidation rates can be greater than 100 percent/in. The rate for oxidation by ozone varies from about 10 percent/in at pH 4.5 to about 1 percent/in at pH 4.0. The contributions from the Fe3+, Mn/+, and carbon-catalyzed reactions (for the conditions specified) are below 1 percent/in for solutions of pH ~ 4.5. Oxidation by NO2 and HONO is even less significant under these conditions. An additional influence on the rates and kinetics of sulfite oxidation in clouds and rainwater can arise from the complexation of HSOi(aq) by aldehydes scavenged by the droplets. The common gaseous products of atmospheric oxidation of the impurity hydrocarbons, CH2O, CH3CHO, CH3COCH3, and possibly other carbonyl compounds, can in principle complex with HSO3. The result of such interactions could increase the solubility of SO2 in the droplet and conceivably retard sulfite oxidation by the oxidants. Although there is significant theoretical evidence that H2O2 may be the most important oxidizing agent for acid generation in cloud water and rain, unambiguous
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193 experimental measurements of H2O2 levels in air and in cloud water and rain have not been possible to date. Peroxide development in the sampling train and lack of selectivity of the luminol technique employed in previous work prevents a firm conclusion about the H2O2 levels in the atmosphere and in precipitation (Heikes et al. 1982, Zika and Saltzman 1982). Theoretical simulations of the atmospheric chemistry of a mixture of reactants in a highly diluted urban atmosphere show that H2Ok generation through reactions (71) and (73) can provide substantial levels of H2O2, CH302H, and other hydroperoxides in the gas phase. If these species are absorbed i non In aerosols, cloud water, or rainwater, for example, as is probable for H2O2 in view of its very high Henry's law constant, then oxidation by H2O2, and possibly other peroxides, can be most significant. The possible roles for peroxyacetylnitrate, peroxynitr~c acid, CH3C2H, and other peroxides in solution-phase sulfite oxidation remain to be evaluated. Solution-Phase Generation of Nitric and Nitrous Acids in the Troposphere There is some evidence of the formation of HONO2 in clouds and rainwater. Recently, both theory and experi- ment suggest that HONO2 may be formed rapidly from a combined gasphase/liquid-phase process. Through simulation, Heikes and Thompson (1981) have suggested that N205 generated by O3-NO2 reactions (15) and (16) may be scavenged effectively by H2O droplets to form HONO2 in clouds or rainwater. NO2 + O3 ~ O2 + NO3, N ~ ~ NO>(+M) ~ N2O5(+M), N2O5 + H2O(liq) ~ 2H+(aq) ~ 2NO3(aq) (15) (16) (81) A preliminary report of the experimental observation of this process has been made recently by Gertler et al. (1982). An evaluation of the contribution of this mechanism to the total [HONO2] found in the atmosphere is not now possible. Lee and Schwartz (1981) studied the rates of reactions (82) and (83):
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194 2NO2(g) + HzO(liq) ~ 2H+(aq) + NO3(aq) + NO2(aq), (82) NO(g) + NO2(g) + H2O(liq) ~ 2H+(aq) + 2NO2(aq). (83) They found that at the lower partial pressures character- istic of moderately polluted atmospheres the rates are slow (10-9 to 10-8 M/h). Unless high partial pressures (e.g., 10-7 atm) of these gases are maintained in contact with liquid water for substantial periods of time (tens of hours), reactions (82) and (83) cannot represent a substan- tial source of atmospheric acidity. Lee and Schwartz found no evidence of Fez+ ion catalysis of these reactions, but they suggested that possibly other metal ions could enhance these rates. There is no evidence related to this pos- sibility now available. Although significant uncertainty remains concerning the source of HNO3 in clouds and rainwater, the limited evidence currently available favors the probable importance of the formation of N2O5 in (15) and (16) followed by its reaction in water droplets to form HNO3. SUMMARY In summary, a wide variety of interrelated homogeneous gas-phase, solution-phase, and heterogeneous chemistry may result ultimately in oxidation of SO2 to sulfuric acid and NOX to nitric acid. The homogeneous gas-phase oxidation of SC2 by the HO radical and the solution-phase oxidation of S(IV) through H2O2, O3, and possibly other species appear to be the major sources of H2SO4. In the cloud-free, ambient, sunlight-irradiated troposphere, nitric acid is probably generated largely by the reaction of HO radicals with NO2. Both HONO2 and H2SO4 produced in the gas phase can be scavenged effectively by cloud water and precipita- tion. NO2 may be oxidized to HONO2 if sufficient O3 and NO2 are present. Following its gas phase generation, N2O5 may be scavenged effectively by water droplets to form HONO2. All the various pathways that lead to the oxidation of SO2 and NOX are coupled by common products and reactants that can directly and indirectly influence the rates of reaction by the other pathways. In general the homogene- ous gas-phase reactions can lead to maximum daylight rates of acid formation of a few percent per hour for SO2 and 20-30 percent/in for NOk. Solution-phase reactions involving H2O2 and O3 can in principle
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195 convert SO2 to H2SO4 in cloud water and precipitation at much higher rates (as high as 100 percent/in) for concen- trations Of H2O2 and O3 in the troposphere that reasonably could result from the normal homogeneous reactions characteristic of atmosphere chemistry during daylight. REFERENCES Anderson, L.G. 1980. Absolute rate constants for the reaction of OH with NO2 in N2 and He from 225 to 389K. J. Phys. Chem. 84:2152-2155. Aubuchon, C. 1976. The rate of iron catalyzed oxidation of sulfur dioxide by oxygen in water. Ph.D. thesis. The Johns Hopkins University, Baltimore, Md. Barrie, L.A., and H.W. Georgii. 1976. An experimental investigation of the absorption of sulfur dioxide by water drops containing heavy metal ions. Atmos. Environ. 10:743-749. Benson, S.W. 1978. Thermochemistry and kinetics of sulfur-containing molecules and radicals. Chem. Rev. 78:23-35. Brimblecombe, P., and D.J. Spedding. 1974. The reaction order of the metal ion catalyzed oxidation of sulfur dioxide in aqueous solution. Chemosphere 1:29-32. Brodzinski, R., S.G. Chang, S.S. Markowitz, and T. Novakov. 1980. Kinetics and mechanism for the catalytic oxidation of sulfur dioxide on carbon in aqueous suspension. J. Phys. Chem. 84:3354-3358. Calvert, J.G., and R.D. McQuigg. 1975. The computer simulation of rates and mechanisms of photochemical smog formation. Int. J. Chem. Kinet. Symp. 1:113-154. Calvert, J.G., and W.R. Stockwell. 1983. The mechanism and rates of the gas phase oxidations of sulfur dioxide and nitrogen oxides in the atmosphere. In Acid Precipitation: SO2, NO, and NO2 Oxidation Mechanisms: Atmospheric Considerations. Ann Arbor, Mich.: Ann Arbor Scientific Publications. In press. Calvert, J.G., F. Su, J.W. Bottenheim, and O.P. Strausz. 1978. Mechanism of the homogeneous oxidation of sulfur dioxide in the troposphere. Atmos. Environ. 12:197-226. Campbell, M.J., J.C. Sheppard, and B.F. Au. 1979. Measurements of hydroxyl concentration in the boundary layer air by monitoring CO oxidation. Geophys. Res. Lett. 6:175-178.
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Representative terms from entire chapter: