Encyclopedia of Environmental Science and Engineering, Volume I and II

130 ATMOSPHERIC CHEMISTRY

Calvert and Stockwell (1983) have shown that the gas-

phase oxidation of sulfur dioxide is primarily by the reaction

of the hydroxyl radical with SO 2 :

HO  SO 2  M → HOSO 2  M (33)

HOSO 2  O 2 → HO 2  SO 3 (34)

SO 3  H 2 O → H 2 SO 4 (35)

In this sequence of reactions, the OH radical initiates the oxi-

dation of SO 2. The bisulfite radical (HOSO 2 ) product reacts

rapidly with oxygen to form sulfur trioxide (SO 3 ) and HO 2.

The HO 2 radical can be converted back to OH by reaction (9),

and the SO 3 can react with water to form sulfuric acid. The

details of the kinetics of these processes have been presented

by Anderson et al. (1989). This sequence of reactions can be

simplified for modeling purposes to the reaction

OH  SO 2 ( O 2 , H 2 O) → H 2 SO 4  HO 2 (36)

The modeling suggests that for moderately polluted and

mildly polluted cases described above, the maximum SO 2

oxidation rates were 3.4%/hour and 5.4%/hour. These maxi-

mum conversions occurred near noon, when the OH concen-

tration was a maximum. The conversion of SO 2 to H 2 SO 4 for

a clear summertime 24-hour period was 16% and 24% for the

moderately and mildly polluted conditions. The gas-phase

oxidation of both NO 2 and SO 2 vary considerably, depending

on the concentrations of other species in the atmosphere. But

the gas-phase oxidation of SO 2 is always going to be much

slower than that for NO 2.

Aqueous-Phase Chemistry

Aqueous-phase oxidations of nitrogen oxides are not

believed to be very important in the atmosphere. On the

other hand, the aqueous-phase oxidations of sulfur dioxide

appear to be quite important. Sulfur dioxide may dissolve in

atmospheric water droplets, to form mainly the bisulfite ion

(HSO 3 − ):

SO 2  Cloud → SO 2 ·H 2 O → HSO 3 −  H^ ^ (37)

The concentration of the bisulfite ion in the droplet is depen-

dent on the Henry’s law constant (H), which determines

the solubility of SO 2 in water, the equilibrium constant (K)

for the first dissociation of the hydrated SO 2 , the gas-phase

SO 2 concentration, and the acidity of the solution.

[HSO 3 − ] = KH [SO 2 ] gas /[H^ ^ ]

SO 2 ·H 2 O, HSO 3 − , and SO 3 2− are all forms of sulfur (IV)

(S(IV)). At normal pH levels, the bisulfite ion is the predomi-

nate form of sulfur (IV) in aqueous systems, and the form

that needs to be oxidized to the sulfate ion (SO 4 2− ), sulfur (VI).

HSO 3 − can be oxidized by oxygen, but this process is very

slow. The reaction may be catalyzed by transition metal ions,

such as manganese (Mn^2 ^ ) and iron (Fe^3 ^ ). The importance

of these metal-catalyzed oxidations depends strongly on the

concentration of metal ions present. This may be enhanced

by passing through heavily industrialized areas, where there

might be sources of these metals for the atmosphere.

Ozone and hydrogen peroxide are likely to be more

important catalysts for the oxidation of S(IV). The rate of

ozone-catalyzed oxidation of S(IV) decreases as the pH of

the solution decreases (or as the solution becomes more

acidic). Since the HSO 3 − concentration depends inversely on

[H^ ^ ], the rate of oxidation of S(IV) slows down considerably

as the pH decreases ([H^ ^ ] increases). This reaction is likely

to be of importance at pH 4.5.

Hydrogen peroxide, on the other hand, is much more solu-

ble than ozone. Hence, even though the gas-phase concentra-

tions are much lower than ozone, the aqueous concentrations

can be high. The rate constant for the hydrogen-peroxide-

catalyzed reaction increases as the pH decreases, down to

a pH of about 2.0. At a pH of 4.5, the oxidation catalyzed

by 1 ppb of gaseous H 2 O 2 in equilibrium with the aqueous

phase is about 100 times faster than the ozone-catalyzed oxi-

dation by 50 ppb of gaseous O 3 in equilibrium with the aque-

ous phase. Figure 8 shows a comparison of aqueous-phase

0 1 2 34 5^6

pH

10 –18

10 –16

10 –14

10 –12

10 –10

10 –8

10 –6

O 3

H 2 O 2

• d

[S(IV)]/

dt

, M s

–1

Fe (III)

NO 2

Mn2+

FIGURE 8 Comparison of aqueous-phase oxidation paths; the

rate of conversion of S(IV) to S(VI) as a function of pH. Conditions

assumed are: [SO 2 (g)] = 5 ppb; [NO 2 (g)] = 1 ppb; [H 2 O 2 (g)] = 1 ppb;

[O 3 (g)] = 50 ppb; [Fe(III)] = 0.3 μM; and [Mn(II)] = 0.03 μM. From

Seinfeld and Pandis (1998). With permission.

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