Urban Ozone The Photochemical Smog Process The Origin and Occurrence of Smog

Many urban centers in the world undergo episodes of air pollution during which relatively high levels of ground-level ozone—an undesirable constituent of air if present in appreciable concentrations at low altitudes in the air that we breathe—are produced as a result of the light-induced chemical reaction of pollutants. This phenomenon is called photochemical smog and is sometimes characterized as "an ozone layer in the wrong place," to contrast it with the beneficial stratospheric ozone discussed in Chapter 1. The word smog is a combination of smoke and fog. The process of smog formation involves hundreds of different reactions, involving dozens of chemicals, occurring simultaneously. Indeed, urban atmospheres have been referred to as giant chemical reactors. The most important reactions that occur in such air masses will be discussed in detail in Chapter 5, In the material below, we investigate the nature and origin of the pollutants—especially nitrogen oxides—that combine to produce photochemical smog.

The chief original reactants in an episode of photochemical smog are molecules of nitric oxide, NO, and of unburned and partially oxidized hydrocarbons that are emitted into the air as pollutants from internal combustion engines; nitric oxide is also released from electric power plants. The concentrations of these chemicals are orders of magnitude greater than are found in clean air. Gaseous hydrocarbons and partially oxidized hydrocarbons are also present in urban air as a result of the evaporation of solvents, liquid fuels, and other organic compounds. Collectively, the substances, including hydrocarbons and their derivatives, that readily vaporize into the air are called volatile organic compounds, or VOCs. (Formally, VOCs are defined as organic compounds having boiling points that lie between 50°C and 260°C.) For example, vapor is released into the air when a gasoline tank is filled unless the hoses nozzle is specially designed to minimize this loss. Evaporated, unburned gasoline is also emitted from the tailpipe of a vehicle before its catalytic converter has been warmed sufficiently to operate. Two-cycle engines such as those in outboard motor boats are particularly notorious for emitting significant proportions of their gasoline unburned into the air. Personal watercraft manufactured in the 1990s, before pollution controls came into effect, emitted more smog-producing emissions in a day's operation than an automobile of the same era driven for several years! Regulations proposed recently in California would require new lawn mowers to be outfitted with a catalytic converter, though this issue is controversial since some mower manufacturers claim that a hot converter could pose a fire hazard to the engine.

Another vital ingredient in photochemical smog is sunshine, which increases the concentration of free radicals that participate in the chemical processes of smog formation. Although the reactants— NO and VOCs—are relatively innocuous, the final products of the smog reaction— ozone, nitric acid, HN03, and partially oxidized (and in some cases nitrated) organic compounds—are much more toxic.

VOCs + NO + 02 + sunlight-»-► mixture of 03, HN03, organics

Substances such as NO, hydrocarbons, and other VOCs that are emitted directly into air are called primary pollutants; the substances into which they are transformed, such as 03 and HNO3, are called secondary pollutants. A summary of the relative importance of various economic sectors in emissions of the primary pollutants sulfur dioxide, nitrogen oxides, and VOCs in the United States and Canada is given in Figure 3-1.

Other than those that absorb sunlight and subsequently decompose, most atmospheric molecules that are transformed in air begin by reacting with the hydroxyl free radical, OH, which consequently is the key reactive species in the troposphere. The most reactive VOCs in urban air are

1 no

Percentage of the emission, by sector





Nitrogen Sulfur Volatile organic oxides dioxide compounds

£ Transportation 1_ Fuel combustion

L_1 Industrial processes HH Other

FIGURE 3-1 North American emissions of primary gaseous air pollutants from various sectors. [Source: U.S. EPA 1999 National Air Quality Trends Report.)

FIGURE 3-1 North American emissions of primary gaseous air pollutants from various sectors. [Source: U.S. EPA 1999 National Air Quality Trends Report.)

hydrocarbons that contain a carbon-carbon double bond, C=C, and aldehydes, since their reactions with OH—and also with sunlight in the latter case—are very fast. Other hydrocarbons such as methane are also present in air, but due to the higher activation energy required, their reaction with OH is sluggish. However, their reaction can become important in late stages of photochemical smog episodes.

Nitrogen Oxide Production During Fuel Combustion

Nitrogen oxide gases are produced by two different reactions whenever a fuel is burned in air with a hot flame. Some nitric oxide is produced from the oxidation of nitrogen atoms contained in the fuel itself; it is called fuel NO. About 30-60% of a fuel's nitrogen is converted to NO during combustion. However, most fuels do not contain much nitrogen, so this process accounts for only a small fraction of NO emissions.

Nitric oxide, produced by the oxidation at high combustion temperatures of atmospheric nitrogen, is called thermal NO. At high flame temperatures, some of the nitrogen and oxygen gases in the air passing through the flame combine to form NO:

hor flame

The higher the flame temperature, the more NO is produced. Since this reaction is very endothermic, its equilibrium constant is very small at normal temperatures but increases rapidly as the temperature rises. One might expect that the relatively high concentrations of NO that are produced under combustion conditions would revert back to molecular nitrogen and oxygen as the exhaust gases cool, since the equilibrium constant for the above reaction is much smaller at lower temperatures. However, the activation energy for the reverse reaction is also quite high, so the process cannot occur to an appreciable extent except at high temperatures. Thus the relatively high concentrations of nitric oxide produced during combustion are maintained in the cooled exhaust gases; equilibrium cannot be quickly re-established, and the nitrogen is "frozen" as NO.

Because the reaction between N2 and 02 has a high activation energy, it is negligibly slow except at very high temperatures, such as occur in the modern combustion engines of vehicles—particularly when they are traveling at high speeds—and in power plants. Very little NO is produced by the burning of wood and other natural materials since the flame temperatures involved in such combustion processes are relatively low.

Two distinct mechanisms are involved in the initiation of the reaction of molecular nitrogen and oxygen to produce thermal nitric oxide; in one it is atomic oxygen that attacks intact N2 molecules, whereas in the other it is free radicals, such as CH, that are derived from the decomposition of the fuel. The initial reaction steps of the first mechanism are

The rate of the second, slower step is proportional to [O] [N2]. However, since, from the equilibrium in the first step, [O] is proportional to the square root of [02], it follows that the rate of NO formation will be proportional to [N21 [Oz]1/2.

The nitric oxide released into air is gradually oxidized to nitrogen dioxide, N02, over a period of minutes to hours, the rate depending upon the concentration of the pollutant gases present. Collectively, NO and NOz in air is referred to as NOx, pronounced "nox." The yellow-brown color in the atmosphere of a smog-ridden city is due in part to the nitrogen dioxide present, since this gas absorbs visible light, especially near 400 nm (see its spectrum in Figure 3-2), removing sunlight's purple component while allowing most yellow light to be transmitted. The small levels of NOx in clean air result in part from the operation of the above reaction in the very energetic environment of lightning flashes and in part from the release of NOx and of ammonia, NH3, from biological sources. Recently it has been discovered that NOx is emitted from coniferous trees when sunlight shines on them and when the ambient concentrations of these gases are low.

FIGURE 3-2 Absorption spectrum of N02 in the solar ultraviolet region. [Source: Reprinted from A. .M. Bass et al., /. Res. U.S. Natl. Bur. Stand, km (1976): 143.]

Mechanism For Photochemical Smog

290 300 310 320 330 340 350 360 370 380 390 400 410 420

Wavelength (nm)

290 300 310 320 330 340 350 360 370 380 390 400 410 420

Wavelength (nm)

Ground-Level Ozone in Smog

Photochemical smog is a widespread phenomenon in the modern world. If we are to prevent or limit its formation, we must understand the main chemical reactions that occur in it. Although the detailed reaction mechanism in smog is quite complicated (discussed further in Chapter 5), its most important aspects are discussed below.

In order for a city to generate photochemical smog, several conditions must be fulfilled. First, there must be substantial vehicular traffic to emit sufficient NO, reactive hydrocarbons, and other VOCs into the air. Second, there must be warmth and ample sunlight in order for the crucial reactions, some of them photochemical, to proceed at a rapid rate. Finally, there must be relatively little movement of the air mass so that the reactants are not quickly diluted. For reasons of geography (e.g., the presence of mountains) and dense population, cities such as Los Angeles, Denver, Mexico City, Tokyo, Athens, Sao Paulo, and Rome all fit the bill splendidly and consequently are subject to frequent smog episodes. Indeed, the photochemical smog phenomenon was first observed in Los Angeles in the 1940s and has generally been associated with that city ever since, although pollution controls have partially alleviated the smog problem in recent decades.

As in the stratosphere, ground-level ozone is produced by the reaction of oxygen atoms with diatomic oxygen. The main source of the oxygen atoms in the troposphere, however, is the photochemical dissociation by sunlight of nitrogen dioxide molecules:


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