Field Studies

Combustion of fossil fuels containing sulfur produces S02. Because virtually all of the sulfur is emitted in the form of S02, with only small amounts in other forms such as H2S04, sulfates, and S03, one can readily calculate S02 emissions from the sulfur content of the fuel. Once emitted, the gaseous S02 is oxidized in the plume itself or, after dilution with the surrounding air, to form H2S04 and sulfates. It is these oxidation reactions that are the major focus of this chapter.

Numerous field studies of the rate of S02 oxidation in the troposphere have shown that the oxidation rate depends on a number of parameters. These include the presence of aqueous phase in the form of clouds and fogs, the concentration of oxidants such as H202 and

03, and sunlight intensity. For example, sulfur in cloudwater in the Los Angeles area over a two-year period from 1983 to 1985 was essentially all in the form of sulfate; at the same time, excess H202 was present in the cloudwater (Richards, 1995). Similarly, fogs and clouds have been shown to be associated with significant sulfate production (e.g., see Eatough et al., 1984; Jacob et al., 1987; Pandis et al., 1992; and Laj et al., 1997). S02 oxidation rates in a plume from a coal-fired power plant have been measured to be typically less than 0.5% h~' in February but 1-3% at midday in June when the sunlight intensity and availability of oxidants are much greater (Lusis et al., 1978). The same effects of clouds and fogs, oxidants, and sunlight intensity on the rate of oxidation of S02 have been observed after the contents of the plume have been dispersed in air (e.g., see McMurry and Wilson, 1983; and Husain and Dutkiewicz, 1990). Indeed rates of oxidation as high as ~30% h"1 have been measured in ambient air in such locations as Budapest, Hungary (Mesaros et al., 1977), and St. Louis, Missouri (Breeding et al., 1976; Alkezweeny and Powell, 1977).

There are some sampling sites located on mountains that have been used to assess the relative amounts of S02 oxidation that occurs in clouds versus in the gas phase. In these cases, fixed sampling sites can be used and the concentrations of S02, sulfate, and associated species measured in the gas and particle phases as well as in the cloudwater as the cloud passes over the sampling site (e.g., see Saxena and Lin, 1990; Aneja and Kim, 1993; and Burkhard et al., 1995).

Husain and co-workers, for example (Husain, 1989; Husain et al., 1991; Husain and Dutkiewicz, 1992; Burkhard et al., 1995; Dutkiewicz et al., 1995), have developed techniques using Se, As, and Sb as tracers to follow the oxidation of S02 with time in clouds at Whiteface Mountain in New York State. The principle is based on the fact that the major sources of these metals are high-temperature combustion, e.g., of coal and oil. Thus these metals are found in particles (see Chapter 9) that act as condensation nuclei for cloud formation. These particles also contain sulfate formed from the gas-phase oxidation of S02. Because of the low vapor pressure of H2S04, it becomes associated with particles, either by homogeneous condensation or by condensing out on preexisting particles (see Chapter 9.C). As a result, the metals can be used as tracers for sulfate formed in the gas phase, as opposed to sulfate formed by the uptake of S02 into cloudwater followed by oxidation.

Figure 8.2 depicts the principle of such an experiment. In air mass A below the cloud, there are sulfate and trace metals such as selenium in suspended particles, and S02 and oxidants in the gas phase. As the air

FIGURE 8.2 Schematic of cloud S02 oxidation studies (from Burkhard et al, 1995).

rises, the temperature falls and condensation to form cloudwater depicted in B results. The water content of the cloud increases with height through C and D. From B to D, not only is there sulfate in the cloudwater from the particles in A that served as condensation nuclei for formation of the cloud, but uptake of S02 and oxidants occurs, leading to in-cloud sulfate formation. As a result, gaseous S02 decreases and sulfate in the aqueous phase increases. Using this technique, Burkhard et al. showed the percentage of in-cloud oxidation to be quite large, up to ~50% of the total measured sulfate (Burkhard et al., f994). Similar results have been obtained elsewhere using other techniques such as stable sulfur isotopes (e.g., see Tanaka et al., 1994).

As we shall see in the following sections, these observations are readily understood in terms of the kinetics and mechanisms of oxidation of S02. The oxidation of S02 occurs in solution and on the surfaces of solids as well as in the gas phase. Indeed, under many conditions typical of the troposphere, oxidation in the aqueous phase provided by clouds and fogs predominates, consistent with the observed dependence on these factors. The presence of oxidizers to react with the S02 is, of course, also a requirement; hence the dependence on 03 (which is a useful surrogate for other oxidants as well) and sunlight, which is needed to generate significant oxidant concentrations.

Figure 8.3 summarizes the pathways that must be considered for S02 oxidation and for the deposition of sulfur compounds (Lamb et al., 1987). The focus of this chapter is on the chemistry converting S02 to sulfate in both the gas and condensed phases.

As we have seen in Chapter 7, the oxidation of NOx to HNO3 occurs to a large extent in the gas phase as well as by the hydrolysis of N2Os on surfaces, and the acid is then taken up by dissolving in clouds and fogs;



FIGURE 8.3 Summary of emission, oxidation, and deposition of S(IV) and S(VI). (Adapted from Atmos. Environ. 21, Lamb, D., Miller, D. F., Robinson, N. F., and Gertler, A. W. "The Importance of Liquid Water Concentration in the Atmospheric Oxidation of S02," pp. 2333-2344. Copyright 1987, with permission from Elsevier Science.)

FIGURE 8.3 Summary of emission, oxidation, and deposition of S(IV) and S(VI). (Adapted from Atmos. Environ. 21, Lamb, D., Miller, D. F., Robinson, N. F., and Gertler, A. W. "The Importance of Liquid Water Concentration in the Atmospheric Oxidation of S02," pp. 2333-2344. Copyright 1987, with permission from Elsevier Science.)

in contrast, a large portion of the total S02 oxidation can occur in the condensed phase (e.g., see Pandis and Seinfeld, 1989a; Wurzler et al., 1995; and Bergin et al., 1996). Another important difference between nitric and sulfuric acids is in the stable form of these two acids in the atmosphere. The vapor pressure of HNO-, is sufficiently large at tropospheric temperatures (e.g., 48 Torr at 20°C) that it exists in the gas phase at the ppt to ppb levels found in the lower atmosphere. Sulfuric acid, on the other hand, has a very low vapor pressure (9.9 X 10"6 Torr at 296 K) so that it exists in the condensed phase diluted with water to varying degrees (e.g., Roedel, 1979; Ayers et al., 1980).

The vapor pressure of H2S04 above solutions with water depends on the solution composition and the temperature. For example, the vapor pressure at 25°C varies from 2.6 X 10"9 Pa for a 54.1 wt% H2S04-H20 solution to 5.9 X 10"6 Pa for a 76.0 wt% solution (Marti et al., 1997). The vapor pressures above solutions partially neutralized with ammonia are also reported by Marti et al. (1997); as discussed in Chapter 9.B.1, the vapor pressures of the partially neutralized solutions are orders of magnitude smaller than those of the acid. As a result, ammonia may play an important role in nucleation of gaseous sulfuric acid in the atmosphere to form new particles.

2. Oxidation in the Gas Phase a. Hydroxy I Radical

The only significant oxidant for S02 in the gas phase is the OH radical:

This reaction is termolecular and is in the falloff region between second and third order at 1 atm pressure. The recommended high- and low-pressure limiting rate constants for reaction (4) at room temperature are /cx = 2 X 10~12 cm3 molecule 1 s"1 and ku = 4.0 X 10~31 [N2] cm6 molecule"2 s"1 with Fc = 0.45 (see Chapter 5.A.2) (Atkinson et al., 1997b) or, alternatively, k^ = 1.5 X 10"12 cm3 molecule"1 s"1 and klt = 3.0 X 10"31 [M] cm6 molecule"2 s"1 with Fc = 0.6 (DeMore et al., 1997). The effective bimolecular rate constant at room temperature is thus k!f = (9.7 or 8.8) X 10"13 cm3 molecule"' s"1, depending on which set of recommendations is used, and the corresponding lifetime of S02 with respect to OH at f X 10h radicals cm~3 is ~13 days.

The adduct free radical formed in reaction (4) [which has been detected directly using neutralization/reioni-zation mass spectrometry (Egsgaard et al., 1988)] subsequently reacts with 02:

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