Nitrogen Oxides N0X in the Absence of Volatile Organic Compounds VOCs

As pointed out in Section 6.7, the main anthropogenic sources of nitrogen oxides in a polluted atmosphere are mobile and stationary petroleum and coal combustion chambers, where inside temperatures and pressures are high enough for the fixation of nitrogen to occur,

(NO* represents a mixture of oxides of nitrogen—mainly NO and NO2) and where quenching to low temperatures outside the chambers is rapid enough to prevent the thermodynamically favored back-reaction (dissociation) from occurring. The major oxide of nitrogen produced is nitric oxide, NO, a relatively nontoxic gas, along with some nitrogen dioxide, NO2. The concentration of NO in a polluted atmosphere is quite sensitive to automobile traffic patterns, in general being a maximum during the peak morning traffic hours (Figure 5-9). However, NO does not absorb radiation above 230 nm, and therefore it cannot be the primary initiator of photochemical reactions in a polluted lower atmosphere. On the other hand, NO2 is a brown gas with a broad, intense absorption band absorbing over most of the visible and ultraviolet regions with a maximum at 400 nm (amax = 6 x 10~19 cm2/molecule) as shown in Figure 5-10, and this gas is a major atmospheric absorber leading to photochemical smog. As we shall see later, there are other light-absorbing species present at very low concentrations in a polluted atmosphere that can also initiate the complex photochemical reactions.

37A mathematical modeling [R. A. Harley, A. G. Russell, and G. R. Cass, Mathematical modeling of the concentrations of volatile organic compounds: Model performance using a lumped chemical mechanism, Environ. Sci. Technol., 27, 1638-1649 (1993)] based on the Southern California Air Quality Study of the Los Angeles Basin on August 27-29, 1987 involved 35 species and 106 chemical reactions, with the atmospheric diffusion equation solved for each of the chemical species.

300 350 400 450 500

Wavelength (nm)

FIGURE 5-10 Absorption spectrum of NO2 at room temperature averaged over 5-nm intervals. Drawn from data of M. H. Harwood and R. L. Jones, ]. Geophys. Res., 99, 22,955-22,964 (1994).

300 350 400 450 500

Wavelength (nm)

FIGURE 5-10 Absorption spectrum of NO2 at room temperature averaged over 5-nm intervals. Drawn from data of M. H. Harwood and R. L. Jones, ]. Geophys. Res., 99, 22,955-22,964 (1994).

The dissociation energy of NO2 is 3.12 eV, which corresponds to a photon having a wavelength of 397.5 nm, and absorption below this wavelength leads to almost complete photodissociation (^ « 1)38:

In the absence of other species, molecular oxygen may be formed by a reaction of an O atom with NO2, reaction (5-43), and indeed the quantum yield of O2 formation approaches unity below 380 nm. In the presence of ground-state molecular oxygen, however, ozone is produced by the three-body reaction (5-29) between O2 and O(3P2). Combination of reactions (5-71), (5-29), and (5-42) leads to the photostationary state

and hence to a steady-state ozone concentration, [O3]ss. Since NO and NO2 appear on different sides of reaction (5-72), the ozone concentration will depend on the ratio of nitrogen oxide concentrations: that is, [O3] is

38Some dissociation even occurs at wavelengths greater than 397.5 nm—for example, the O2 quantum yield is 0.36 at 404.7 nm—and thus energetically below the dissociation energy. But this dissociation is strongly temperature dependent and presumably is due to the contributions of internal rotational and vibrational energies.

proportional to [NO2] / [NO]. However, since we have assumed a photosta-tionary state, [O3]ss will be a function of the light intensity.

Thus, NO2 is the primary light-absorbing species in the atmosphere leading to photochemical smog and the ozone concentration is enhanced by increasing NO2, but NO is the major oxide of nitrogen produced in high-temperature combustion chambers. Since the small amount of NO2 produced directly would be completely depleted in a very short time in direct sunlight, at least one additional process leading to the oxidation of NO to NO2 is required to produce the quantities of O3 and other products generated in a polluted atmosphere. It is now agreed that the hydroxyl (.OH) and hydroperoxyl (HO2.) radicals play key roles in the rapid oxidation of NO to NO2 in the absence of VOCs.39 The HO2- radical is directly involved:

However, we will see shortly that HO2. radicals are formed in free-radical chain mechanisms in which .OH radicals serve as chain carriers, so that .OH and HO2. are interconverted provided sufficient NO is present for reaction (5-73) to be dominant.

The major primary source of .OH in the troposphere is the reaction of excited O(XD2) atoms with water, reaction (5-35), following photolysis of ozone below 310nm [see reaction (5-32)]. The global-averaged noontime concentration of the hydroxyl radical in the lower troposphere is very low—of the order of 106 molecules/cm3—with a lifetime of approximately one second. It does not react with the major tropospheric constituents (N2, O2, CO2, H2O), but it is the major initiator of oxidation in photochemical smog generation and in the removal of most trace gases in the atmosphere. The HO2. radical can react with ozone, regenerating the .OH radical

It can also self-combine to form hydrogen peroxide

which dissolves in fog and clouds and is washed out of the troposphere by rain.

In the absence of VOCs, two species with major natural and anthropogenic sources that lead to HO2. (hence to oxidation of NO to NO2) are carbon

39Thermal oxidation of NO to NO2 by O2 is thermodynamically feasible but is not significant at normal NO pressures in the atmosphere. Similarly, the oxidation of NO by ground-state O atoms is unimportant at the low O-atom concentrations in the lower troposphere.

monoxide, CO, and methane, CH4 (see Section 2.2). Carbon monoxide reacts rapidly (with very low activation energy) with the hydroxyl radical

CO + .OH ^ CO2 + H; &298 « 2 x 10-13 cm3 molecule-1 s-1 (5-76)

followed by the very fast three-body (essentially triple-collision frequency—see Section 4.3) combination reaction

Coupling these two reactions with reaction (5-73) gives the CO chain mechanism, with .OH the primary chain carrier and HO2 the secondary chain carrier:

However, this CO chain does not contribute to the overall NO oxidation to the extent that methane does in a similar chain. The .OH radical is the major tropospheric "sink" for methane via the reaction

Reaction (5-79) results in a tropospheric lifetime for methane of about 12 years. The major fate of the methyl (.CH3) radical in the troposphere is the very fast (also of the order of a triple collision) three-body combination reaction with O2 to form CH3O2.40:

The methylperoxy radical, CH3O2., can react in the troposphere with a variety of species.

1. If the NO concentration is high enough, NO is oxidized to NO2 in a fast reaction analogous to reaction (5-73), producing the methoxy (CH3O.) radical

40Actually, under tropospheric conditions the concentrations of ' CH3 and O2 and the lifetime of CH3O2. are such that the reaction order varies between 2 and 3 depending on the overall pressure. That is, at very low pressures the reaction order approaches 3, whereas at high pressures it approaches 2.

followed by

The .OH radical is regenerated by reaction (5-73). Combination of these reactions gives the .OH-initiated methane chain reaction for oxidation of NO:

.CH3 + O2 M > CH3O2. CH3O2. + NO ^ CH3O. + NO2 CH3O. + O2 ^ H2CO + HO2. HO2. + NO ^ NO2 + .OH

net: CH4 + 2O2 + 2NO —'°H'M > H2CO + 2NO2 + H2O (5-83)

[Reaction (5-75) is a possible chain-terminating step in this mechanism.] This gives an increase in ozone by the photostationary state (5-72).

2. On the other hand, if the tropospheric NO concentration is low (as in some very clean, unpolluted areas), instead of reacting with NO, the methylperoxy radical CH3O2. reacts with HO2% forming methyl hydroperoxide:

In effect reaction (5-84) serves as a "sink" for both .OH and HO2% since the methane chain reaction (5-83) is now terminated and the interconversion of .OH and HO2. is restricted, and this leads to a net decrease in ozone. The methylperoxy radical also rapidly combines with NO2 in the three-body reaction

but this reaction is generally not important in the troposphere because of the fast decomposition of methylperoxynitrate (CH3O2NO2) back to the react-ants.

Formaldehyde, H2CO, is an overall product of the methane chain reaction (5-83), formed in the chain by the reaction of the methoxy radical with molecular oxygen (step 4).41 It absorbs weakly in the near-ultraviolet spectral

41Formaldehyde is also a primary emission product of hydrocarbon combustion, and is produced in the photooxidation of a wide variety of higher hydrocarbons in VOC-containing atmospheres (next section); its average concentration in urban areas is approximately 3-16 parts per billion by volume (ppbv).

region. Its gas-phase absorption spectrum is strongly banded, with \max = 305 nm (ct305 « 7 x 10-20 cm2/molecule), and only very weak absorption near the visible region (e.g., ct370 « 2 x 10-21 cm2/molecule). This absorption forms an excited formaldehyde molecule, H2CO*, which dissociates either to H and the formyl radical HCO,

Reaction (5-86) dominates below roughly 330 nm, corresponding to the dissociation energy of H2CO, while (5-87) is the major dissociation above 330 nm. Thus, H2CO is a source of HO2. radicals by reaction (5-77) and by

HCO + O2 ) HO2. + CO (5-88) Formaldehyde also reacts with .OH

.OH + H2CO ) H2O + HCO (5-89) and combines with HO2. to form initially an alkoxy radical

HO2. + H2CO ) HO2CH2O. (5-90) which rapidly isomerizes to a peroxy radical

However, excitation and subsequent dissociation steps, reactions (5-86) and (5-87), dominate in an atmosphere free from VOCs, leading to a noonday overall atmospheric lifetime of approximately 3 hours for formaldehyde.

A reaction not yet considered is the reaction between NO2 and O3:

The nitrate radical (NO3.) plays an important role in photochemical smog, particularly in nighttime reactions involving aldehydes and simple olefins. It reacts at close to collisional frequency with NO to produce NO2:

It combines with NO2 to form dinitrogen pentoxide, N2O5, by the reversible reaction

with a formation constant K29s = 4.5 x 10~n cm3/molecule. (Equilibrium is reached in about one minute under tropospheric conditions.) N2O5 also reacts with water to form nitric acid,

thus providing a sink for both NO3" and NO2. In the daytime NO3" is rapidly photolyzed to NO and O2

or to NO2 and O

so that its daytime concentration is very low. However, at night it may be one of the most important tropospheric oxidizing agents. Additional reactions of NO3" in the presence of methane and VOCs are covered next.

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