given that the first-order photolysis rate constant is ~ 1.3 X 1CT3 s"1 at a solar zenith angle of 40° (Calvert et al., 1994), even at a peak OH concentration of f X 107 radicals cm"3, this reaction will be too slow to compete with the loss by photolysis.
There are some intriguing observations that suggest that HONO can undergo some as yet unrecognized reactions, at least in laboratory systems, and it seems likely in air as well. The self-reaction of gaseous HONO to form NO + N02 + H20 (i.e., the reverse of reaction (15)) has been observed in laboratory systems (e.g., Ten Brink and Spoelstra, 1998) and treated theoretically (e.g., Mebel et al., 1998). This gaseous reaction is too slow to be important in the atmosphere. However, Kleffmann et al. (1994, 1998) have observed the formation of nitrous oxide, N20, during the decay of HONO in a laboratory system (Fig. 7.10), in a reaction that appears to occur on the reactor surface. While they suggest the overall reaction can be represented by (32),
this clearly cannot be an elementary reaction, and the mechanism remains unclear.
Nitrous acid/nitrite can also be oxidized in the aqueous solutions found in the atmosphere in the form of fogs, clouds, and particles. Nitrite is well known to be slowly oxidized in the dark to nitrate by dissolved oxygen in the liquid phase. However, it has been reported that the rate of this oxidation increases remarkably during freezing of the solution containing the nitrite (Takenaka et al., 1992, 1996). Figure 7.11, for example, shows the rate of nitrate formation in a nitrite solution at 25°C and in one with the cooling bath at — 21°C (Takenaka et al., 1992). This unusual phenomenon has also been observed with respect to the o o
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