Much improvement in the levels of smog and, more importantly, photochemical oxidants has been made in the United States and other countries since the late 1960s. However, room for improvement persists.
The oxidant of critical importance in the photochemical atmosphere is ozone (O3). Several miles above the Earth’s surface, in the troposphere, there is sufficient shortwave ultraviolet (UV) light to directly split molecular O2 to atomic O to combine with O2 to form O3. These UV wavelengths do not reach the Earth’s surface. In this region, nitrogen dioxide efficiently absorbs longer wavelength UV light, which leads to the following simplified series of reactions:NO2 + UV→O + NOO + O2→O3O3 + NO→NO2
This process is cyclic, with NO2 regenerated by the reaction of NO and O. In the absence of hydrocarbons, this series of reactions would approach a steady state with no excess or buildup of O3. However, near the Earth’s surface, the hydrocarbons, especially olefins and substituted aromatics, are attacked by the free atomic O, which, with NO, produces more NO2. Thus, the balance of the reactions shown in the above reactions is upset so that O3 levels build up, particularly when the Sun’s intensity is greatest at midday. The reactions with hydrocarbons are very complex and involve the formation of unstable intermediate free radicals that undergo a series of changes. Aldehydes are major products in these reactions. Formaldehyde and acrolein account for ∼50 and 5%, respectively, of the total aldehyde in urban atmospheres. Peroxyacetyl nitrate (CH3COONO2), often referred to as PAN, and its homologs, also arise in urban air, most likely from the reaction of the peroxyacyl radicals with NO2.