3.6: Biogeochemical reactions in the troposphere

The planetary boundary layer, PBL, is the initial recipient of substances from ground sources. Because of the dynamic character of the atmosphere, material from the PBL will mix into the free troposphere and can be transported horizontally over great distances. Vertical transport across the PBL is significantly enhanced in the presence of deep convective cloud systems. These dynamic processes provide a transition from high mixing ratios of trace gases near source regions to low mixing ratios in the remote troposphere. These changes in mixing ratios result in changes in chemical reaction paths, thus justifying a chemical classification of the troposphere into the background troposphere, the source region, and a "transition" regime between the two. The troposphere is governed by heterogeneous chemistry far more so than the stratosphere. Heterogeneous processes of interest here involve scavenging of trace gases by aerosols, cloud and precipitation elements leading to aqueous phase chemical reactions and to temporary and permanent removal of material from the gas phase. They also include exchange of material with the earth's surface, such as dry deposition. Dry deposition is a major removal process for ozone, as well as for other gases of importance in tropospheric photochemistry (e.g., \(\mathrm{NO}_2\) ).

The Chemistry of Ozone

Ozone is formed by the association reaction of ground state \(\mathrm{O}\) atoms with \(\mathrm{O}_2\) :
\[
\mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{O}_2+\mathrm{M} \rightarrow \mathrm{O}_3+\mathrm{M}
\]
and is removed by photolysis
\[
\begin{array}{l}
\mathrm{O}_3+\mathrm{h} v \rightarrow \mathrm{O}\left({ }^1 \mathrm{D}\right)+\mathrm{O}_2 \\
\mathrm{O}_3+\mathrm{h} v \rightarrow \mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{O}_2 .
\end{array}
\]
The metastable \(\mathrm{O}\left({ }^{\prime} \mathrm{D}\right)\) is quenched by \(\mathrm{O}_2\) and \(\mathrm{N}_2\)
\[
\mathrm{O}\left({ }^1 \mathrm{D}\right)+\mathrm{M} \rightarrow \mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{M}
\]
though it can also react with \(\mathrm{H}_2 \mathrm{O}\) to form the hydroxyl radical \(\mathrm{HO}\)
\[
\mathrm{O}\left({ }^1 \mathrm{D}\right)+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{HO}+\mathrm{HO}
\]
Nitric oxide reacts with ozone to form \(\mathrm{NO}_2\)
\[
\mathrm{NO}+\mathrm{O}_3 \rightarrow \mathrm{NO}_2+\mathrm{O}_2
\]
with \(\mathrm{NO}_2\) removed by photolysis
\[
\mathrm{NO}_2+\mathrm{h} v \rightarrow \mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{NO}
\]
Ozone removed by (6) is reconstituted by (7) followed by (1). It is convenient, because of the cyclic nature of the chemistry, to define a family of species undergoing rapid reactions leading to formation or removal of ozone. We define this family, odd oxygen, as the sum of \(\mathrm{O}_3, \mathrm{O}\left({ }^1 \mathrm{D}\right), \mathrm{O}\left({ }^3 \mathrm{P}\right)\) and \(\mathrm{NO}_2\).

Odd oxygen is conserved in reactions (1)-(4), (6) and (7). It is formed by reactions of \(\mathrm{NO}\) with \(\mathrm{HO}_2\), \(\mathrm{CH}_3 \mathrm{O}_2\), and \(\mathrm{RO}_2\).
\[
\mathrm{HO}_2+\mathrm{NO} \rightarrow \mathrm{HO}+\mathrm{NO}_2
\]
\[
\begin{array}{c}
\mathrm{CH}_3 \mathrm{O}_2+\mathrm{NO} \rightarrow \mathrm{CH}_3 \mathrm{O}+\mathrm{NO}_2 \\
\mathrm{RO}_2+\mathrm{NO}-\mathrm{RO}+\mathrm{NO}_2
\end{array}
\]
The species \(\mathrm{RO}_2\) represents a variety of complex organic peroxy radicals. These reactions occur during the photooxidation of \(\mathrm{CO}, \mathrm{CH}_4\) and hydrocarbons, for example by
\[
\begin{array}{c}
\mathrm{HO}+\mathrm{CO}+\mathrm{O}_2 \rightarrow \mathrm{HO}_2+\mathrm{CO}_2 \nonumber \\
\mathrm{HO}_2+\mathrm{NO} \rightarrow \mathrm{NO}_2+\mathrm{HO} \nonumber \\
\mathrm{NO}_2+\mathrm{h} v \rightarrow \mathrm{NO}+\mathrm{O} \nonumber \\
\mathrm{O}+\mathrm{O}_2+\mathrm{M} \rightarrow \mathrm{O}_3+\mathrm{M} \nonumber \\
\mathrm{NET}: \mathrm{CO}+2 \mathrm{O}_2 \rightarrow \mathrm{CO}_2+\mathrm{O}_3
\nonumber \end{array} \nonumber
\]
and by
\[
\begin{aligned}
\mathrm{HO}+\mathrm{CH}_4 & \rightarrow \mathrm{CH}_3+\mathrm{H}_2 \mathrm{O} \nonumber \\
\mathrm{CH}_3+\mathrm{O}_2+\mathrm{M} & \rightarrow \mathrm{CH}_3 \mathrm{O}_2+\mathrm{M} \nonumber \\
\mathrm{CH}_3 \mathrm{O}_2+\mathrm{NO} & \rightarrow \mathrm{CH}_3 \mathrm{O}+\mathrm{NO}_2 \nonumber \\
\mathrm{CH}_3 \mathrm{O}+\mathrm{O}_2 & \rightarrow \mathrm{CH}_2 \mathrm{O}+\mathrm{HO}_2 \nonumber \\
\mathrm{HO}_2+\mathrm{NO} & \rightarrow \mathrm{NO}_2+\mathrm{HO} \nonumber \\
2 \times\left(\mathrm{NO}_2+\mathrm{h} v\right. & \left.\rightarrow \mathrm{NO}+\mathrm{O}\left({ }^3 \mathrm{P}\right)\right) \quad \lambda<420 \mathrm{~nm} \nonumber \\
2 \times\left(\mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{O}_2+\mathrm{M}\right. & \left.\rightarrow \mathrm{O}_3+\mathrm{M}\right) \nonumber \\
\mathrm{NET}: \mathrm{CH}_4+4 \mathrm{O}_2 & \rightarrow \mathrm{CH}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O}+2 \mathrm{O}_3
\nonumber \end{aligned} \nonumber
\]
Odd oxygen is removed by reaction of \(\mathrm{O}\left({ }^1 \mathrm{D}\right)\) with \(\mathrm{H}_2 \mathrm{O}\), by reaction of \(\mathrm{O}_3\) with \(\mathrm{HO}_2\) and \(\mathrm{HO}\),
\[
\begin{array}{c}
\mathrm{HO}_2+\mathrm{O}_3 \rightarrow \mathrm{OH}+2 \mathrm{O}_2 \\
\mathrm{HO}+\mathrm{O}_3 \rightarrow \mathrm{HO}_2+\mathrm{O}_2
\end{array}
\]
by formation of nitric acid from \(\mathrm{NO}_2\),
\[
\mathrm{HO}+\mathrm{NO}_2+\mathrm{M} \rightarrow \mathrm{HNO}_3
\]
and by heterogeneous reactions of \(\mathrm{O}_3\) and \(\mathrm{NO}_2\) at the earth's surface.

Oxidation of \(\mathrm{CO}, \mathrm{CH}_4\) and hydrocarbons leads to net production of odd oxygen in the presence of adequate \(\mathrm{NO}_{\mathrm{x}}\). The rate for production of ozone is roughly proportional to the concentration of \(\mathrm{NO}\), while the rate of loss is almost independent of \(\mathrm{NO}_{\mathrm{x}}\left(\mathrm{NO}+\mathrm{NO}_2\right)\) for concentrations below 200 ppt (Fishman et al., 1979). Loss of odd oxygen, primarily by reactions (5) and (11) is balanced by production in reactions (8) and (9) for concentrations of NO near 30 ppt (Fishman et al., 1979; Logan et al., 1981; Crutzen, 1983). Hence, regions of the globe characterized by extremely low concentrations of NO, such as the remote Pacific, are likely to provide a net photochemical sink for odd oxygen (Liu et al., 1983), while the continental boundary layer at midlatitudes, characterized by higher concentrations of NO, is likely to provide a net source. Measurements of \(\mathrm{NO}_{\mathrm{x}}\) in the troposphere are few, and the lack of data for \(\mathrm{NO}_{\mathrm{x}}\) contributes significantly to uncertainties in global estimates for the photochemical source of ozone.
Production of ozone in the troposphere is limited ultimately by supply of \(\mathrm{CO}, \mathrm{CH}_4\), and hydrocarbons, if \(\mathrm{NO}_{\mathrm{x}}\) is available. One molecule of ozone may be formed for each molecule of CO (Crutzen, 1973b), while the yield of ozone from oxidation of \(\mathrm{CH}_4\) could be as large as 3.5 (e.g., Logan et al., 1981). Production of ozone from nonmethane hydrocarbons is discussed in more detail later in this chapter.

The Chemistry of HO

Reaction with \(\mathrm{HO}\) in the troposphere is the primary removal mechanism for many trace gases which influence the chemical composition and radiative balance of the stratosphere. Consequently, tropospheric and stratospheric chemistry are inextricably linked through the \(\mathrm{HO}\) radical.

The hydroxyl radical is formed by reaction (2) followed by (5). It is removed by reaction with carbon monoxide and methane,
\[
\begin{array}{c}
\mathrm{HO}+\mathrm{CO}+\mathrm{O}_2 \rightarrow \mathrm{HO}_2+\mathrm{CO}_2 \\
\mathrm{HO}+\mathrm{CH}_4+\mathrm{O}_2 \rightarrow \mathrm{CH}_3 \mathrm{O}_2+\mathrm{H}_2 \mathrm{O}
\end{array}
\]
with reaction (14) dominant over (15). \(\mathrm{HO}_2\) is removed by reaction (8), and (16)
\[
\mathrm{HO}_2+\mathrm{O}_3 \rightarrow \mathrm{HO}+2 \mathrm{O}_2
\]
leading to regeneration of \(\mathrm{HO}\), or by
\[
\mathrm{HO}_2+\mathrm{HO}_2 \rightarrow \mathrm{H}_2 \mathrm{O}_2+\mathrm{O}_2
\]
Hydrogen peroxide is photolysed,
\[
\mathrm{H}_2 \mathrm{O}_2+\mathrm{h} v \rightarrow \mathrm{HO}+\mathrm{HO}
\]
it can react with \(\mathrm{HO}\),
\[
\mathrm{HO}+\mathrm{H}_2 \mathrm{O}_2 \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{HO}_2
\]
or may be removed by heterogeneous processes such as precipitation scavenging, as discussed below.
\[
\mathrm{H}_2 \mathrm{O}_2+\text { rain } \rightarrow \text { removal. }
\]

It is convenient, because of the fairly rapid interconversion of \(\mathrm{H}, \mathrm{HO}, \mathrm{HO}_2\) and \(\mathrm{H}_2 \mathrm{O}_2\) to consider these species as a family, odd hydrogen \(\left(=\mathrm{H}+\mathrm{HO}+\mathrm{HO}_2+2 \times \mathrm{H}_2 \mathrm{O}_2\right)\). Reaction (5) provides the dominant source for odd hydrogen, while reactions (19) and (20) provide important sinks for odd hydrogen. Reactions (8), (14) and (16)-(18) do not influence the concentration of odd hydrogen.

Atmospheric oxidation of methane (Figure 2-1) may also provide sources and sinks for odd hydrogen, in addition to providing a source for odd oxygen. \(\mathrm{CH}_3 \mathrm{O}_2\), formed in (15), is removed primarily by reaction with \(\mathrm{NO}\) or \(\mathrm{HO}_2\).
\[
\begin{array}{c}
\mathrm{CH}_3 \mathrm{O}_2+\mathrm{NO} \rightarrow \mathrm{CH}_3 \mathrm{O}+\mathrm{NO}_2 \\
\mathrm{CH}_3 \mathrm{O}_2+\mathrm{HO}_2 \rightarrow \mathrm{CH}_3 \mathrm{OOH}+\mathrm{O}_2
\end{array}
\]
Subsequent reaction of methyl hydroperoxide with HO leads to loss of odd hydrogen.
\[
\mathrm{CH}_3 \mathrm{OOH}+\mathrm{HO} \rightarrow \mathrm{CH}_3 \mathrm{O}_2+\mathrm{H}_2 \mathrm{O}
\]
\[
\mathrm{HO}_2+\mathrm{HO} \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2 .
\]
Recent measurements of a relatively fast rate for reaction (22) (Niki et al., 1983) indicate that this reaction may provide an important sink for odd hydrogen at low concentrations of \(\mathrm{NO}_{\mathrm{x}}\) (Logan, 1985, private communication).

The net source of odd hydrogen is determined first by competition between reaction (9) and reaction (21), and second by competition between reactions (22)-(23) which remove odd hydrogen,
\[
\mathrm{CH}_3 \mathrm{OOH}+\text { rain } \rightarrow \text { removal }
\]
and reaction (24)
\[
\mathrm{CH}_3 \mathrm{OOH}+\mathrm{h} v \rightarrow \mathrm{CH}_3 \mathrm{O}+\mathrm{HO}
\]
which recycles odd hydrogen. Reaction of the methoxy radical with oxygen yields formaldehyde.
\[
\mathrm{CH}_3 \mathrm{O}+\mathrm{O}_2 \rightarrow \mathrm{CH}_2 \mathrm{O}+\mathrm{HO}_2
\]
Photolysis of \(\mathrm{CH}_2 \mathrm{O}\)
\[
\begin{array}{l}
\mathrm{CH}_2 \mathrm{O}+\mathrm{h} v \rightarrow \mathrm{H}+\mathrm{HCO} \\
\mathrm{CH}_2 \mathrm{O}+\mathrm{h} v \rightarrow \mathrm{H}_2+\mathrm{CO}
\end{array}
\]
via path (26a) provides an important source for odd hydrogen in the upper troposphere, while path (26b) and reaction of \(\mathrm{HO}\) with \(\mathrm{CH}_2 \mathrm{O}\) have no net effect on odd hydrogen.

The major pathways in the methane oxidation chain are thought to be understood fairly well. Recent kinetic data for some of the intermediate species have changed the net effect of the cycle on the budget of odd hydrogen. It should be emphasized that the influence of \(\mathrm{CH}_4\) chemistry on the budgets of both odd hydrogen and odd oxygen depends critically on ambient concentrations of \(\mathrm{NO}_{\mathrm{x}}\), because of competition between reactions (9) and (21). A discussion of the chemistry of \(\mathrm{HO}\) and associated uncertainties may be found in Chameides and Tan (1981), Logan et al. (1981) and NRC (1984). Logan et al. (1981) present a detailed discussion of the potential influence of \(\mathrm{CH}_4\) chemistry on odd hydrogen.

The Chemistry of Oxides of Nitrogen

Nitrogen oxides act as catalysts in the photochemical production of ozone. Any reaction which converts \(\mathrm{NO}\) to \(\mathrm{NO}_2\), other than the reaction of \(\mathrm{NO}\) with ozone, provides a photochemical source of ozone. Present measurements of \(\mathrm{NO}_{\mathrm{x}}\) are inadequate for definition of its distribution. Preliminary data indicate that ambient concentrations are highly variable in space and time, as discussed in Chapter 3. The lack of data for \(\mathrm{NO}_{\mathrm{x}}\) precludes accurate quantification of the net global chemical source for ozone.

Nitrogen oxides are produced in the troposphere primarily in the form of NO. Nitric oxide and nitrogen dioxide are rapidly interconverted by reactions (6), (7), (8) and (9) on a time scale of minutes. Nitrogen oxides are removed from the atmosphere by conversion to nitric acid,
\[
\mathrm{HO}+\mathrm{NO}_2+\mathrm{M} \rightarrow \mathrm{HNO}_3
\]
followed by heterogeneous processes, i.e., rainout or surface deposition of \(\mathrm{HNO}_3\). Surface deposition of \(\mathrm{NO}_2\) could provide another important tropospheric sink. There have been few field studies of the rate of uptake of \(\mathrm{NO}_{\mathrm{x}}\) by surfaces, as discussed later in this chapter, and removal rates of \(\mathrm{NO}_{\mathrm{x}}\) are not well defined at present.

Current models suggest that \(\mathrm{NO}_{\mathrm{x}}\) is converted to \(\mathrm{HNO}_3\) by reaction (13) within 1-2 days in summer. Nitric acid may be converted to aerosol nitrate, for example by reaction with sea-salt aerosol (Savoie and Prospero, 1982) or by reaction with ammonia (Tang, 1980). Nitric acid and aerosol nitrate are likely to be removed from the atmosphere by precipitation and surface deposition, with mean lifetimes of a few days (Junge, 1963; Levine and Schwartz, 1982). Nitric acid is converted back to \(\mathrm{NO}_{\mathrm{x}}\) by reactions (20) and (21) more slowly, with a time scale of 2-4 weeks.
\[
\begin{array}{c}
\mathrm{HNO}_3+\mathrm{HO} \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{NO}_3 \\
\mathrm{HNO}_3+\mathrm{h} v \rightarrow \mathrm{HO}+\mathrm{NO}_2
\end{array}
\]
These processes are significant primarily in the upper troposphere.
Nitric oxide and \(\mathrm{NO}_2\) may be converted to \(\mathrm{HNO}_2, \mathrm{HO}_2 \mathrm{NO}_2, \mathrm{NO}_3, \mathrm{~N}_2 \mathrm{O}_5\) and organic nitrates in addition to \(\mathrm{HNO}_3\). Most of these molecules decompose thermally or photolytically and therefore provide temporary reservoirs for \(\mathrm{NO}_{\mathrm{x}}\). Peroxynitric acid is formed by reaction of \(\mathrm{HO}_2\) with \(\mathrm{NO}_2\)
\[
\mathrm{HO}_2+\mathrm{NO}_2+\mathrm{M} \rightarrow \mathrm{HO}_2 \mathrm{NO}_2+\mathrm{M} .
\]
It decomposes rapidly in the lower troposphere,
\[
\mathrm{HO}_2 \mathrm{NO}_2 \rightarrow \mathrm{HO}_2+\mathrm{NO}_2
\]

but it is thermally stable in the colder upper troposphere, where it is removed by photolysis and by reaction with HO. Current models suggest that a significant fraction of acidic nitrate may be present in the form of \(\mathrm{HO}_2 \mathrm{NO}_2\) in the upper troposphere, but there are no observational data at present.

Peroxyacetyl nitrate (PAN) is formed during the degradation of hydrocarbons (see below). PAN is more stable than \(\mathrm{HO}_2 \mathrm{NO}_2\) with respect to thermal decomposition, with a lifetime of about a day at \(275^{\circ} \mathrm{K}\) and several years at temperatures characteristic of the upper troposphere (Cox and Coffey, 1977; Hendry and Kenley, 1979). Recent kinetic data indicate that the lifetime of PAN towards photolysis is about four months (Senum et al., 1984), with a similar value for attack by HO (Wallington \(e t\) al., 1984). Recent measurements of PAN show that it is an important reservoir for \(\mathrm{NO}_{\mathrm{x}}\) in clean marine air (Singh and Salas, 1983a, b) and in rural air in North America (Bottenheim et al., 1984; Spicer et al., 1983) and in Europe (Brice et al., 1984).
The nitrate radical \(\left(\mathrm{NO}_3\right)\) is formed by
\[
\mathrm{NO}_2+\mathrm{O}_3 \rightarrow \mathrm{NO}_3+\mathrm{O}_2
\]
and is photolysed rapidly during the day. At night a steady state should be established between \(\mathrm{NO}_3\) and \(\mathrm{N}_2 \mathrm{O}_5\).
\[
\mathrm{NO}_3+\mathrm{NO}_2+\mathrm{M} \leftrightarrows \mathrm{NO}_2 \mathrm{O}_5+\mathrm{M} .
\]
The nitrate radical has been observed in a variety of "clean-air" locations including Hawaii and rural Colorado. In most cases the concentration of \(\mathrm{NO}_3\) at a given level of \(\mathrm{NO}_2\) was lower than predicted by model calculations (Noxon et al., 1980; Platt et al., 1980b, 1981, 1984; Platt and Perner, 1980; Noxon, 1983). Some processes, as yet unidentified, appear to be removing \(\mathrm{NO}_3\) and/or \(\mathrm{N}_2 \mathrm{O}_5\) from the atmosphere at night. If these processes lead to removal of \(\mathrm{NO}_{\mathrm{x}}\), then nighttime removal of \(\mathrm{NO}_{\mathrm{x}}\) by reactions involving \(\mathrm{NO}_3\) could be comparable to daytime by reaction (13). Platt \(e t\) al. \((1981,1984)\) speculate that reactions of \(\mathrm{NO}_3\), or more likely \(\mathrm{N}_2 \mathrm{O}_5\), on wet aerosols may account for rapid removal of \(\mathrm{NO}_3\) under conditions of high humidity, and may provide another sink for \(\mathrm{NO}_{\mathrm{x}}\) in clouds and fog. A detailed discussion of \(\mathrm{NO}_3\) scavenging by cloud droplets, and its possible conversion to \(\mathrm{NO}_3^{-}\)by chemical reactions in the liquid phase is given in the section on heterogeneous chemistry. The rate of the homogeneous gas-phase reaction between \(\mathrm{N}_2 \mathrm{O}_5\) and water vapor,
\[
\mathrm{N}_2 \mathrm{O}_5+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{HNO}_3+\mathrm{HNO}_3
\]
is extremely slow, but this reaction could provide an important source for \(\mathrm{HNO}_3\) if it proceeds at the upper limit \(\left(\mathrm{k}<1.3 \times 10^{-21} \mathrm{~cm}^3\right.\) molecule \(\left.{ }^{-1} \mathrm{sec}^{-1}\right)\) given by Tuazon et al. (1983). The role of (33) in nighttime chemistry will remain uncertain until its rate constant is better defined.

CHEMISTRY OF OZONE FORMATION IN THE POLLUTED TROPOSPHERE

The discussion on clean tropospheric chemistry indicated that the oxidation of \(\mathrm{CO}, \mathrm{CH}_4\), and possibly other biogenic hydrocarbons leads to a net production of odd oxygen, if sufficient \(\mathrm{NO}_{\mathrm{x}}\) is present (i.e., \(\mathrm{NO}_{\mathrm{x}}>30 \mathrm{ppt}\) ). In the polluted troposphere, unlike the clean troposphere, the rate of ozone formation is not necessarily proportional to the concentration of \(\mathrm{NO}_{\mathrm{x}}\), but varies in a complex way that is dependent upon the ratios and concentrations of the hydrocarbons and \(\mathrm{NO}_{\mathrm{x}}\) as well as the chemical composition of the hydrocarbons themselves. The chemistry of ozone in the polluted troposphere has been studied for over thirty years. The research emphasis has been predominantly in elucidating the processes for oxidant/ozone formation in urban areas. The focus of interest in urban areas is due mainly to the high emission densities of ozone precursors (hydrocarbons and \(\mathrm{NO}_{\mathrm{x}}\) ) which result typically in very high ozone concentrations under summertime meteorological conditions. In many cases the concentration levels of ozone in urban areas exceed ambient air quality standards which have been set to protect public health and welfare.
Several reviews of the chemistry of polluted atmospheres are available (Leighton, 1961; Stern, 1977; Seinfeld, 1975; Heiklen, 1976), as are detailed discussions of reaction mechanisms (Demerjian et al., 1974; Carter et al., 1979; Baldwin et al., 1977; Falls and Seinfeld, 1978; Whitten et al., 1980) and reaction rate constant reviews (Baulch et al., 1980; Demerjian et al., 1980; Baulch et al., 1982). The discussions which follows provides an overview of the chemistry of the polluted troposphere.

The chemistry that occurs in sunlight-irradiated polluted atmospheres involves the interaction of a host of chemical species. These include: hydrocarbons such as alkanes, alkenes, and aromatics; other organics such as aldehydes and ketones; nitric oxide \((\mathrm{NO})\); nitrogen dioxide \(\left(\mathrm{NO}_2\right)\); ozone \(\left(\mathrm{O}_3\right)\); peroxyacetyl nitrate (PAN); nitric acid \(\left(\mathrm{HNO}_3\right)\); atomic oxygen \(\left(\mathrm{O}^3 \mathrm{P}\right)\) and its first electronic excited state \(\mathrm{O}\left({ }^1 \mathrm{D}\right)\); hydroxy radical (HO); hydroperoxyl radical \(\left(\mathrm{HO}_2\right)\); alkylperoxyl radicals \(\left(\mathrm{RO}_2\right)\); acylperoxyl radicals \(\left(\mathrm{R}(\mathrm{O}) \mathrm{O}_2\right)\); nitrogen trioxide \(\left(\mathrm{NO}_3\right)\); and dinitrogen pentoxide \(\left(\mathrm{N}_2 \mathrm{O}_5\right)\).

This chemistry explains the rapid conversion of \(\mathrm{NO}\) to \(\mathrm{NO}_2\) observed in the ambient polluted atmosphere (Leighton, 1961; Altshuller and Bufalini, 1965; Demerjian et al., 1974). The key lies in a sequence of reactions involving the same free radical species thought to be important in the chemistry of the clean troposphere and a number of organic radicals derived from the host of more complex hydrocarbon and organic molecules present additionally in the polluted atmosphere.

The most important atmospheric reactions governing the decomposition of alkanes, alkenes, and aromatics involve their reaction with hydroxyl radical. Reaction sequences describing the pathways for their oxidation as a result of \(\mathrm{HO}\) attack have been studied extensively (Demerjian et al., 1974; Carter et al., 1979; Niki, 1978; Falls and Seinfeld, 1978; Perry et al., 1977; Atkinson et al., 1985; Grosjean, 1984). In a very simplified form this may be represented as follows:
\[
\mathrm{HO}+\text { Hydrocarbon } \rightarrow \mathrm{R}+\mathrm{H}_2 \mathrm{O}
\]
The alkyl radical (R) produced as a result of the \(\mathrm{HO}\) attack on the hydrocarbon reacts with an oxygen molecule to form an alkylperoxyl radical \(\left(\mathrm{RO}_2\right)\)
\[
\mathrm{R}+\mathrm{O}_2 \rightarrow \mathrm{RO}_2
\]

The alkylperoxyl radical reacts with \(\mathrm{NO}\) to form \(\mathrm{NO}_2\) and an alkoxy radical (RO)
\[
\mathrm{RO}_2+\mathrm{NO} \rightarrow \mathrm{RO}+\mathrm{NO}_2
\]
Hydrogen abstraction from the alkoxy radical by molecular oxygen will produce a hydroperoxyl radical \(\left(\mathrm{HO}_2\right)\) and a carbonyl compound \((\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{H})\) :
\[
\mathrm{RO}+\mathrm{O}_2 \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{H}+\mathrm{HO}_2
\]
The aldehydes formed in the RO oxidation, or emitted as such from combustion sources, react with HO or photolyze, introducing another important source of radicals to the atmosphere.
\[
\begin{array}{r}
\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{H}+\mathrm{HO} \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O})+\mathrm{H}_2 \mathrm{O} \\
\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{H}+\mathrm{h} \nu(\lambda<400 \mathrm{~nm}) \rightarrow \mathrm{R}+\mathrm{H}(\mathrm{C}=\mathrm{O})
\end{array}
\]
The hydroperoxyl radical can then react with \(\mathrm{NO}\) to form \(\mathrm{NO}_2\) and another hydroxyl radical, which reenters the cycle.
\[
\mathrm{HO}_2+\mathrm{NO} \rightarrow \quad \mathrm{HO}+\mathrm{NO}_2
\]
Reactions (8) and (10) taken together provide an efficient mechanism for the rapid conversion of NO to \(\mathrm{NO}_2\). This increase in the concentration of "odd oxygen" (previous section) results in an increase of ozone through the perturbation of the chemical steady state relationship (39) derived from the following reactions:
\[
\begin{aligned}
\mathrm{NO}_2+\mathrm{h} \nu & \rightarrow \mathrm{NO}+\mathrm{O}\left({ }^3 \mathrm{P}\right) \\
\mathrm{O}\left({ }^3 \mathrm{P}\right)+\mathrm{O}_2+\mathrm{M} & \rightarrow \mathrm{O}_3+\mathrm{M} \\
\mathrm{O}_3+\mathrm{NO} & \rightarrow \mathrm{NO}_2+\mathrm{O}_2
\end{aligned}
\]
The chemical steady state relationship (39) has been shown to be a valid approximation over a considerable range of atmospheric pollutant conditions (Stedman and Jackson, 1975; Calvert, 1976).
\[
\left[\mathrm{O}_3\right]=\frac{\mathrm{J}_7\left[\mathrm{NO}_2\right]}{\mathrm{k}_6[\mathrm{NO}]}
\]
Ratios of \(\mathrm{NO}_2\) and \(\mathrm{NO}\) emitted into the atmosphere by man's activities are typically less than one and would lead to very low ozone steady state concentrations under typical atmospheric solar irradiation conditions. The impact of reactions (8) and (10) is to drive the ratio of \(\mathrm{NO}_2\) to \(\mathrm{NO}\) up by oxidizing the \(\mathrm{NO}\) to \(\mathrm{NO}_2\), in direct competition with the ozone-nitric oxide reaction and thereby allowing ozone to build up in concentration.

Typical observations made in a polluted atmosphere of this phenomenon are shown in Figures 4-1a and \(b\). The diurnal pattern begins with the emission of hydrocarbons, carbon monoxide and \(\mathrm{NO}_{\mathrm{x}}\) from motor vehicles in the early morning. The \(\mathrm{NO}\) is converted to \(\mathrm{NO}_2\) and ozone accumulation begins when

most of the NO has been oxidized. The ozone concentration maximizes and then declines either as a result of reaction with additional emissions of \(\mathrm{NO}\) through the day, dilution due to meteorology, or in the case of extended periods of time and transport, by interaction with the ground surface, which is a major sink for ozone in the atmosphere as will be discussed later in this chapter.

The interaction of organic free radicals produced by hydrocarbon oxidation with \(\mathrm{NO}\) and \(\mathrm{NO}_2\) represents an important aspect of the chemistry of the oxides of nitrogen in the polluted atmosphere. They represent key processes in the conversion of \(\mathrm{NO}\) to \(\mathrm{NO}_2\) and the formation of organic nitrates. It has recently been found that the more complex peroxyalkyl radicals formed during alkane photooxidation can add NO to form alkyl nitrates in non-negligible yields, in competition with reaction (10) (Atkinson et al., 1985):
\[
\mathrm{RO}_2+\mathrm{NO} \rightarrow \mathrm{RONO}_2
\]
Also the hydroperoxyl radical formed from hydrocarbon oxidation in reaction (36) supplements that formed in reactions (12) and (14) to create hydrogen peroxide via reaction (17), and aldehydes react with the hydroxy radical to produce peroxyacyl radicals and ultimately peroxyacyl nitrates by reaction with \(\mathrm{NO}_2\)
\[
\begin{array}{c}
\mathrm{RCHO}+\mathrm{OH} \rightarrow \mathrm{RCO}+\mathrm{H}_2 \mathrm{O} \\
\mathrm{RCO}+\mathrm{O}_2 \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2 \\
\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2+\mathrm{NO}_2 \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2 \mathrm{NO}_2
\end{array}
\]
A thermal equilibrium exists between the peroxyacyl nitrate and its components
\[
\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2 \mathrm{NO}_2 \leftrightarrows \mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2+\mathrm{NO}_2
\]
and the peroxyacyl radical is removed from the atmosphere by reaction with NO (Cox and Coffey 1977)
\[
\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}_2+\mathrm{NO} \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{O}+\mathrm{NO}_2
\]
The most commonly occurring peroxyacyl nitrate is peroxyacetyl nitrate or PAN and there is a great similarity between the behavior of PAN and ozone in polluted atmospheres. Their diurnal variations at ground based sites are virtually identical on many occasions (Garland and Penkett 1976) primarily because their formation and removal mechanisms are very similar. PAN is thus an excellent tracer of the photochemical reactivity of the troposphere and since it has no large stratospheric source it can be used in studies designed to investigate the seasonal behavior of tropospheric photochemical activity (Brice et al., 1984).
4.2.2 Night-time Chemistry of the Source Region
The nitrate radical \(\mathrm{NO}_3\) can be formed by reaction of \(\mathrm{NO}_2\) with ozone
\[
\mathrm{NO}_2+\mathrm{O}_3 \rightarrow \mathrm{NO}_3+\mathrm{O}_2
\]
During the day \(\mathrm{NO}_3\) is rapidly photolyzed back to its precursors but during the night an equilibrium with \(\mathrm{N}_2 \mathrm{O}_5\) can be established
\[
\mathrm{NO}_3+\mathrm{NO}_2=\mathrm{N}_2 \mathrm{O}_5
\]

Night-time concentrations of \(\mathrm{NO}_3\) as high as \(355 \mathrm{pptv}\) have been recorded in the Los Angeles basin (Platt et al., 1980b, 1984). Some \(\mathrm{NO}_3\) is converted to \(\mathrm{HNO}_3\) by reactions with hydrocarbons and aldehydes,
\[
\begin{aligned}
\mathrm{NO}_3+\mathrm{RH} & \rightarrow \mathrm{R}+\mathrm{HNO}_3 \\
\mathrm{NO}_3+\mathrm{R}(\mathrm{C}=\mathrm{O}) \mathrm{H} & \rightarrow \mathrm{R}(\mathrm{C}=\mathrm{O})+\mathrm{HNO}_3
\end{aligned}
\]
but in the absence of \(\mathrm{NO}\) no chain reactions can be initiated by the peroxy radicals, which will be rapidly lost. The \(\mathrm{N}_2 \mathrm{O}_5\) which is formed will be stable until sunrise or it can be removed either by hydrolysis in deliquescent aerosols or by uptake on ground surfaces.

It is quite possible that heterogeneous chemistry is important in the overall \(\mathrm{NO}_{\mathrm{x}}\) cycle particularly in the formation of nitrous acid, HONO, which can act as an efficient source of \(\mathrm{HO}\) radicals in conditions of low light intensity, but this is an area where much more research is required before definite statements can be made.

Excerpted from:

V.A. Mohnen, W. Chameides K.L. Demerjian D. H. Lenschow J.A. Logan R.J. McNeal, S.A. Penkett, U. Platt, U. Schurath, P. da Silva Dias. Tropospheric Chemistry, in Atmospheric Ozone 1985: Assessment of Our Understanding of the Processes Controlling Its Present Distribution and Change, World Meteorological Organization. Accessed November 2023 at https://csl.noaa.gov/assessments/ozone/1985/report.html