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Unit 11: Atmospheric Pollution // Section 4: Secondary Air Pollutants


Secondary pollutants form when primary pollutants react in the atmosphere. Table 1 summarizes common forms of atmospheric reactions.

Table 1. Chemical kinetics in the atmosphere.
Type of reaction Process Notation
Bimolecular Two reactants combine to produce two products. A + B → C + D
Three-body Two reactants combine to form one new product. A third, inert molecule (M) stabilizes the end product and removes excess energy. A + B + M → AB + M
Photolysis Solar radiation photon breaks a chemical bond in a molecule A + hν → B + C
Thermal decomposition A molecule decomposes by collision with an inert molecule (M) A + M → B + C

For reactions to take place, molecules have to collide. However, gases are present in the atmosphere at considerably lower concentrations than are typical for laboratory experiments or industrial processes, so molecules collide fairly infrequently. As a result, most atmospheric reactions that occur at significant rates involve at least one radical—a molecule with an odd number of electrons and hence an unpaired electron in its outer shell. The unpaired electron makes the radical unstable and highly reactive with other molecules. Radicals are formed when stable molecules are broken apart, a process that requires large amounts of energy. This can take place in combustion chambers due to high temperatures, and in the atmosphere by photolysis:

Nonradical + hν → radical + radical

Radical formation initiates reaction chains that continue until radicals combine with other radicals to produce nonradicals (atoms with an even number of electrons). Radical-assisted chain reactions in the atmosphere are often referred to as photochemical mechanisms because sunlight plays a key role in launching them.

One of the most important radicals in atmospheric chemistry is the hydroxyl radical (OH), sometimes referred to as the atmospheric cleanser. OH is produced mainly through photolysis reactions that break apart tropospheric ozone, and is very short-lived. It is consumed within about one second by oxidizing a number of trace gases like carbon monoxide, methane, and nonmethane VOCs (NMVOCs). Some of these reactions eventually regenerate OH in continuous cycles, while others deplete it.

Since OH has a short atmospheric lifetime, its concentration can vary widely. Some anthropogenic emissions, such as carbon monoxide and VOCs, deplete OH, while others such as NOx boost OH levels. Measuring atmospheric OH is difficult because its concentration is so low. Long-term trends in OH concentrations are uncertain, although the prevailing view is that trends over the past decades have been weak because of compensating influences from carbon monoxide and VOCs on the one hand and NOx on the other hand. Since OH affects the rates at which some pollutants are formed and others are destroyed, changes in OH levels over the long term would have serious implications for air quality.

Ground-level ozone (O3) is a pernicious secondary air pollutant, toxic to both humans and vegetation (Fig. 5). It is formed in surface air (and more generally in the troposphere) by oxidation of VOCs and carbon monoxide in the presence of NOx. The mechanism is complicated, involving hundreds of chemically interactive species to describe the VOC degradation pathways. A simple schematic is:

VOC + OH → HO2 + other products
HO2 + NO → OH + NO2
NO2 + hν → NO + O
O + O2 + M → O3+ M

An important aspect of this mechanism is that NOx and OH act as catalysts—that is, they speed up the rate of ozone generation without being consumed themselves. Instead they cycle rapidly between NO and NO2, and between OH and HO2.

Ozone damage to plant leaves

Figure 5. Ozone damage to plant leaves
See larger image

Source: Courtesy United States Environmental Protection Agency.

This formation mechanism for ozone at ground level is totally different from that for ozone formation in the stratosphere, where 90 percent of total atmospheric ozone resides and plays a critical role in protecting life on Earth by providing a UV shield (for details see Unit 2, "Atmosphere"). In the stratosphere ozone is produced from photolysis of oxygen (O2 + hν → O + O, followed by O + O2 + M → O3 + M). This process does not take place in the troposphere because the strong (< 240 nm) UV photons needed to dissociate molecular oxygen are depleted by the ozone overhead.

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