Abstract
Several gaseous chemical species may lead to adverse health effects and, therefore, several of those are regulated. Brief descriptions of those chemical species, including their major sources and atmospheric fate, are presented. Next, the focus of this chapter is on urban and regional pollution, since it corresponds to most of the population exposure to ambient air pollution. The gaseous pollutants that are currently the most relevant at the urban/regional scale in terms of adverse health effects are ozone and nitrogen dioxide. These pollutants are major components of photochemical smog, which results from chemical reactions between nitrogen oxides (NOx) and volatile organic compounds (VOC) in the presence of sunlight. The fact that photochemical smog precursors such as NOx and some VOC (alkenes) are both producers and destructors of ozone makes the development of efficient strategies to reduce photochemical smog difficult.
Several gaseous chemical species may lead to adverse health effects and, therefore, several of those are regulated. Brief descriptions of those chemical species, including their major sources and atmospheric fate, are presented. Next, the focus of this chapter is on urban and regional pollution, since it corresponds to most of the population exposure to ambient air pollution. The gaseous pollutants that are currently the most relevant at the urban/regional scale in terms of adverse health effects are ozone and nitrogen dioxide. These pollutants are major components of photochemical smog, which results from chemical reactions between nitrogen oxides (NOx) and volatile organic compounds (VOC) in the presence of sunlight. The fact that photochemical smog precursors such as NOx and some VOC (alkenes) are both producers and destructors of ozone makes the development of efficient strategies to reduce photochemical smog difficult. Therefore, this chapter addresses gaseous air pollutants with a focus on the complex processes leading to the formation of photochemical smog. Following a presentation of the main reactions that govern the formation of ozone, nitrogen dioxide, and other gaseous pollutants, the different chemical regimes are analyzed in order to understand how efficient approaches to reduce the concentrations of the main constituents of photochemical smog can be developed.
8.1 General Considerations on Gaseous Pollutants
Gaseous pollutants include primary pollutants, secondary pollutants, and precursors of secondary pollutants. Some chemical species may belong to more than one category. The main gaseous pollutants that are regulated or are precursors of regulated pollutants are briefly presented in the following sections.
8.1.1 Sulfur Dioxide
Sulfur dioxide (SO2) is a primary pollutant. SO2 is the main pollutant of air pollution episodes such as that of 1952 in London. It is regulated because it leads to respiratory problems. As a primary pollutant, its impacts occur near its sources. In North America and Europe, those impacts are now mostly limited to some industrial sites and maritime traffic, because sulfur content in fuels used by road traffic is now regulated.
SO2 is oxidized to sulfuric acid in the atmosphere (see Chapter 10). Sulfuric acid contributes to fine particle air pollution, because its low volatility implies that it is preferentially present in the particulate phase, typically as ammonium salts. In addition, sulfuric acid is an important component of acid rain. SO2 is slightly soluble in water, and it may be removed from the atmosphere by dry and wet deposition.
8.1.2 Carbon Monoxide
Carbon monoxide (CO) is an inorganic carbonaceous compound, because it does not contain any hydrogen atoms. CO is regulated because it combines with hemoglobin in the blood stream and leads to anoxia (lack of oxygen), when the blood stream cannot carry enough oxygen. CO is a primary pollutant, which is mostly emitted by combustion processes, including internal combustion engines. Historically, CO concentrations were, therefore, high near roadways. CO was the first pollutant to be regulated for road traffic in North America and Europe, using catalytic converters, which convert it to carbon dioxide, CO2 (see Chapter 2). Consequently, ambient concentrations of CO are now fairly low in these regions.
CO is oxidized slowly to CO2 in the atmosphere. Its oxidation, which is described in Section 8.3.3, leads also to the formation of ozone, the main gaseous pollutant of photochemical smog. However, the contribution of CO to photochemical air pollution is less than that of other carbonaceous compounds because of its low chemical reactivity and the fact that its emissions are currently regulated.
8.1.3 Ozone and Gaseous Photochemical Oxidants
Photochemical pollution is generally called “photochemical smog.” Smog is the contraction of smoke and fog, because its appearance falls between these two phenomena. Photochemical pollution was initially identified in the Los Angeles Basin in the 1950s. Heavy pollution was present in that region of southern California, because of a large number of anthropogenic pollution sources, such as road traffic, fossil-fuel fired power plants, refineries, and other industrial sources, as well as meteorological conditions that were conducive to air pollution (strong sunlight and low atmospheric dispersion). Arie Haagen-Smit, a biochemistry professor at the California Institute of Technology (Caltech), was the first to identify the processes that lead to the formation of photochemical air pollution, and in particular ozone (O3), which is its main gaseous pollutant. He showed that photochemical air pollution results from atmospheric chemical reactions occurring among precursor gases that include nitrogen oxides (NOx) and volatile organic compounds (VOC), in the presence of sunlight. Hence, the adjective “photochemical” was attributed to that form of air pollution, since photolytic reactions induced by sunlight (see Chapter 7) initiate the set of photochemical reactions that lead to ozone formation.
Ozone is not emitted in the atmosphere and is, therefore, a secondary air pollutant. Its precursors are NOx, VOC, and, to a lesser extent, CO. Although O3 is not the only photochemical oxidant, it is the main gaseous constituent of photochemical smog in terms of ambient concentrations. Therefore, O3 is targeted for regulations pertaining to the gaseous fraction of photochemical smog. Since O3 is produced by photochemical reactions, its concentrations are highest during spring and summer.
8.1.4 Nitrogen Oxides
NOx include by definition nitric oxide (NO) and nitrogen dioxide (NO2). Their emissions result mostly from combustion processes. These two compounds represent in terms of ambient concentrations the majority of nitrogen oxides present in the urban atmosphere. However, there are other nitrogen oxides, such as nitrogen protoxide (N2O), which is a greenhouse gas, nitric acid (HNO3), nitrogen pentoxide (N2O5), the nitrate radical (NO3), and a large number of organic nitrates. By convention, NOy represent all nitrogen oxides, with the exception of N2O, and NOz represent the difference between NOy and NOx. Thus, in summary:
NO2 is the only nitrogen oxide that is regulated in terms of its ambient concentrations. It leads to respiratory problems. NO2 is both a primary and a secondary pollutant, because (1) it constitutes a fraction of NOx emissions and (2) it is a product of the atmospheric oxidation of NO.
NOx are rapidly oxidized in the atmosphere. They are precursors of a large number of secondary pollutants, such as ozone, nitric acid, and organic nitrates. Nitric acid may contribute significantly to secondary particulate matter formation. It is also a major constituent of acid rain. Some organic nitrates contribute to the secondary fraction of organic particulate matter. In addition, inorganic and organic nitrates play a major role in the eutrophication of ecosystems.
8.1.5 Volatile and Semi-volatile Organic Compounds
Volatile organic compounds (VOC) and semi-volatile organic compounds (SVOC) are mostly emitted from combustion processes and also from the evaporation of liquid fuels and some organic-containing products (e.g., some paints, solvents, and cleaning products). They play a major role in the formation of photochemical smog. VOC are precursors of ozone. SVOC and some VOC are also important precursors of the secondary fraction of fine particulate matter. In addition, some VOC, such as benzene, 1,3-butadiene, and formaldehyde, are carcinogenic. These carcinogenic VOC may be regulated individually as such (as it is the case in Europe for benzene) or may be regulated indirectly via regulatory approaches that target carcinogenic compounds as a whole (as it is the case in the United States).
VOC include mostly alkanes (hydrocarbons with only single bonds, also called paraffins), alkenes (hydrocarbons with one or more double bonds, also called olefins), aromatic compounds (compounds with one or more phenyl rings), and aldehydes (compounds with a carbonyl group, HC=O). Alcohols (compounds with a C-OH group), alkynes (hydrocarbons with a triple bond), ethers (compounds with a C-O-C group), and other organic compounds are typically present in the atmosphere to a lesser extent. However, alcohols are becoming more prominent due to the use of biofuels and their increased use in gasoline.
The term hydrocarbon (HC) refers to organic compounds that contain only carbon and hydrogen atoms (therefore, alkanes, alkenes, alkynes, and some aromatic compounds). We will use hereafter the term VOC to refer to all organic compounds, thus including aldehydes and alcohols. Among the alkanes, methane has a very low atmospheric reactivity (atmospheric lifetime on the order of 10 years) and, therefore, it is typically not included among the precursors of photochemical air pollution. Thus, when referring to VOC that are precursors of photochemical pollution, one should use the term non-methane VOC. However, for the sake of simplicity, we will use the term VOC to mean non-methane VOC hereafter.
SVOC are organic compounds with a saturation vapor pressure that is such that they can be present in the atmosphere in both the gas phase and the particulate phase. They play a major role in the formation of particulate organic compounds and may also be involved in the formation of gaseous air pollutants.
The ultimate chemical fate of VOC and SVOC is CO2, via atmospheric oxidation. However, organic compounds may be removed from the atmosphere before being converted to CO or CO2. Their removal may occur either as gaseous or particulate compounds, via dry and wet deposition.
8.1.6 Ammonia
Ammonia (NH3) is the reduced form of nitrogen in the atmosphere. It is mostly emitted by agricultural activities. NH3 does not present any adverse health effects at the ambient concentrations usually observed in the atmosphere. However, it contributes to the formation of sulfate and nitrate ammonium salts, which may constitute a significant fraction of fine particulate matter. In addition, NH3 contributes to nitrogen deposition and, therefore, may lead to the eutrophication of ecosystems.
8.2 Oxidizing Power of the Atmosphere and Chemical Reactivity
The atmosphere is an oxidizing environment because of the presence of 21 % of oxygen. However, oxygen is not the main oxidizing species in the atmosphere and atmospheric oxidation processes are due mostly to other chemical species that contain oxygen atoms and are formed photochemically in the atmosphere. The main atmospheric oxidants are:
Their formation in the atmosphere is described in Section 8.3, along with the chemistry of photochemical air pollution.
The oxidation of chemical species, such as VOC, NOx, CO, and SO2 by these oxidants can occur more or less rapidly depending on the oxidant concentrations and the reactivity of the chemical species toward those oxidants. For a given chemical reaction, two terms are generally used to characterize this chemical reactivity:
– The half-life
– The lifetime (also called residence time)
In the case of a reaction with a constant oxidant concentration, the half-life corresponds to the median of the times needed for all the individual molecules of the chemical species initially present to react. The lifetime corresponds to the mean of those reaction times.
The half-life, t½, is the time needed for half of the molecules initially present to react. Let [X]0 be the initial concentration of chemical species X, which here will be oxidized by OH (with a constant concentration), as an example:
The change with time of the concentration of X is given by the following equation (see Chapter 7):
(8.1) where k is the rate constant of the chemical reaction. Integrating this equation between times 0 and t leads to the following solution:
where [X]0 is the initial concentration, i.e., at t = 0. Therefore, when:
(8.3) Thus, the half-life of X is:
(8.4) where ln is the natural logarithm. The lifetime, tl, is the characteristic time of the chemical reaction and is, therefore, defined simply via a dimensional analysis as:
(8.5) It can be shown that tl corresponds to the mean of the reaction times of all X molecules initially present. Let p(t) be the normalized distribution of the individual reaction times of all X molecules (some will react right away or almost right away, whereas others will react after a long, or very long, time). Since all molecules have the same probability of reacting, the number of molecules that will react is proportional to the concentration of these molecules and, by definition, this function, p(t), is proportional to the concentration of X molecules:
This function is normalized and its integration over time must be equal to 1. Therefore, the pre-exponential factor is calculated to be equal to (k [OH]):
The mean of the reaction times of all X molecules can be obtained by integrating the reaction time weighted by the distribution of those reaction times:
(8.8) Integrating by parts:
(8.9) The half-life is related to the lifetime (i.e., mean reaction time) as follows:
If the chemical species X undergoes several oxidation reactions, its overall half-life and overall lifetime can be calculated by considering all the oxidation reactions. For example, if X is oxidized by OH, NO3, and O3:
(8.11) where kOH, kNO3, and kO3 are the rate constants of the different oxidation reactions. Integrating this equation leads to the following solution:
(8.12) The half-life and the lifetime are, respectively, as follows:
(8.13) The overall half-life and lifetime can be expressed in terms of the half-lives and lifetimes of the individual reactions:
(8.14) The lifetimes and half-lives of the main atmospheric pollutants vary greatly, ranging from a few hours for species such as NOx and propane (C3H8) to about 10 years for methane. The chemical reactivity of VOC is related to their ozone formation potential as discussed in Section 8.5.3. Table 8.1 lists typical atmospheric chemical lifetimes of selected chemical species undergoing oxidation by OH, NO3, and O3, as well as photolysis.
Table 8.1. Lifetimes of selected chemical species in the atmosphere at 1 atm and 25 °C for various oxidation reactions and photolysis. Sources of the rate constants: Calvert et al., 2000, 2002, 2008, 2011; Finlayson-Pitts and Pitts, 2000; Mollner et al., 2010.
| Chemical species | Photolysisa | OHa | NO3a | O3a |
|---|---|---|---|---|
| NO2 | (b) | 30 h | (b) | (b) |
| SO2 | – | 12 d | – | – |
| CO | – | 48 d | – | – |
| Methanec | – | 5 a | >300 a | – |
| Propane | – | 11 d | >4 a | – |
| n-Butane | – | 5 d | 7 a | – |
| Hexane | – | 2 d | 3 a | – |
| Octane | – | 36 h | 20 mo | – |
| Ethylene | – | 33 h | 19 mo | 7 d |
| Propylene | – | 11 h | 12 d | 28 h |
| trans-2-Butene | – | 4 h | 7 h | 90 min |
| 1,3-Butadiene | – | 4 h | 28 h | 44 h |
| 1-Hexene | – | 8 h | (d) | 25 h |
| trans-3-Hexene | – | (d) | (d) | 100 min |
| trans-4-Octene | – | 4 h | (d) | 2 h |
| Benzene | – | 8 d | 11 a | – |
| Toluene | – | 2 d | 5 a | – |
| o-Xylene | – | 20 h | 9 mo | – |
| Formaldehyde | 18 h | 33 h | 7 mo | – |
| Acetaldehyde | 9 d | 18 h | 45 d | – |
| Isoprene | – | 3 h | 4 h | 22 h |
| MBOe | – | 4 h | 10 d | 31 h |
| α-Pinene | – | 5 h | 27 min | 3 h |
| Δ3-Carene | – | 3 h | 18 min | 8 h |
| Humulene | – | 1 h | 5 min | 1 min |
| Longifolene | – | 6 h | 4 h | >23 d |
(a) Concentrations: [OH] = 2 × 106 cm−3 over 12 h per day (daytime); [NO3] = 2 × 108 cm−3 over 12 h per day (nighttime); [O3] = 40 ppb over 24 h per day. Photolysis for the spring equinox (March 20) in Paris calculated over 24 h; cos(θs) = 0.39 on average over 12 h during daytime (see Equations 7.7 to 7.10). Lifetimes are calculated with these values (concentrations or sunlight) averaged over 24 h; therefore, for species with lifetimes less than 24 h, the half-lives are shorter during daytime for photolytic reactions and reactions with OH and O3; they are shorter during nighttime for reactions with NO3.
(b) These reactions were not taken into account because they produce NO, which can subsequently be converted back to NO2 (see Section 8.3); the reaction with OH is the only terminal reaction (see R8.44).
(c) The atmospheric lifetime of methane is actually longer because the kinetics depends on temperature (<25 °C on average) and [OH] decreases with altitude.
(d) No data available on the rates of these reactions.
(e) 2-Methyl-3-buten-2-ol.
8.3 Gas-phase Chemistry of Photochemical Air Pollution
8.3.1 Oxidants
The reactions leading to the formation of the three main oxidant species of photochemical air pollution, OH, NO3, and O3, are described in this section.
As mentioned in Chapter 7, hydroxyl radicals can be formed via the photolysis of ozone:
A chemical kinetic mechanism will also need to take into account the fact that only a fraction of the O(1D) excited oxygen atoms reacts with water vapor, because most of them lose their excess energy by collision with air molecules (N2 or O2) and produce O3 back by reacting next with O2. (At 100 % relative humidity and 25 °C, the reaction with air molecules is about five times faster than that with water vapor.)
In addition to this OH formation pathway, there are two other important photolytic reactions that lead to OH formation in the troposphere: (1) the photolysis of hydrogen peroxide (H2O2) and (2) the photolysis of nitrous acid (HNO2):
In addition, as discussed in Section 8.3.3, OH is also formed by reaction of hydroperoxyl radicals, HO2. They originate mostly from the photolysis of aldehydes. Aldehyde photolysis leads to the production of a hydrogen atom, H, which is then oxidized rapidly by O2 to form HO2. Since the OH radical is formed by photolytic reactions, it is present mostly during daytime. There are, however, some formation pathways that do not require photolysis, such as the decomposition of peroxyacetylnitrate (PAN) in presence of NOx (see the chemistry of PAN in Section 8.3.4) and the oxidation of alkenes by O3 (see Section 8.3.6). However, these reactions are very limited sources of OH and nighttime OH concentrations are negligible.
The nitrate radical (not to be confused with the nitrate ion, NO3−, which is present in the aqueous phase, see Chapter 10) is formed via the reaction of nitrogen dioxide with ozone:
This radical is rapidly photolyzed:
Its formation does not require any photochemical reaction; therefore, it can be formed either at night or during the day. However, its rapid photolysis diminishes significantly its concentration during daytime. Thus, this oxidant plays a role mostly at night. Since the kinetics of the reaction of O3 with NO is about 1,000 times faster than that with NO2, NO concentrations are negligible when NO3 is present, because NO will have almost entirely been oxidized into NO2.
Ozone is formed in the stratosphere by photolysis of oxygen molecules (see Chapter 7). Solar radiation that leads to oxygen photolysis is filtered in the stratosphere and is, therefore, unavailable in the troposphere to lead to ozone formation there. However, the photolysis of nitrogen dioxide, which takes place in the visible and the near ultraviolet (UV) range, takes place in the troposphere:
Therefore, ozone formation takes place in the presence of sunlight, i.e., during daytime. However, the lifetime of ozone ranges from several hours to a few days. Thus, its oxidizing power can also occur at night.
In summary, the atmospheric gaseous oxidants are the following:
– During daytime: OH and O3
– During nighttime: NO3 and O3
8.3.2 The Photostationary State of Leighton
Ozone formation is balanced by its destruction by nitric oxide:
This reaction is very fast and can, therefore, be called a titration reaction when considered in isolation, i.e., it stops when one of the two reactants (NO or O3) has been entirely depleted. In the atmosphere, in the presence of sunlight, O3 can be continuously regenerated by NO2 photolysis. Thus, a system of three reactions that are at equilibrium occurs, i.e., the rates of these three reactions are identical. This set of three reactions, which are at steady state, is called the photostationary state of Leighton, named after the Stanford chemistry professor, Philip A. Leighton.
These three reactions are as follows:
(R8.11) 
(R8.9) 
(R8.10) The rate constants are given here at 1 atm and 25 °C. A typical average daytime photolysis rate is used. At steady state, the rates of the three reactions are equal:
The first equation leads to:
(8.16) In addition, the sum of the concentrations of nitrogen oxides must remain constant:
where the subscript 0 indicates the initial concentration. Each ozone molecule reacting with NO leads to a molecule of NO2, and each photolyzed NO2 molecule leads to a molecule of ozone; therefore, the sum of the concentrations of O3 and NO2 remains constant:
Thus, the concentrations of nitrogen oxides can be calculated as a function of the initial concentrations and the ozone concentration:
(8.19) Replacing [NO2] and [NO] in Equation 8.16:
(8.20) The ozone concentration is then the solution of a quadratic equation:
(8.22) Example: Calculation of the ozone concentration produced from nitrogen oxides
The initial nitrogen oxide concentrations are as follows: [NO]0 = 100 ppb, [NO2]0 = 5 ppb. There is no ozone present initially: [O3]0 = 0 ppb.
The solution is: [O3] = 0:5 ppb
The ozone concentration produced from only nitrogen oxides is, therefore, very low.
8.3.3 Oxidation of Carbon Monoxide (CO)
The chemistry of CO is simple and is, therefore, convenient to explain ozone formation when volatile carbonaceous species (CO or VOC) are present. CO is oxidized by OH radicals:
Hydrogen atoms are not stable, and they recombine rapidly with molecular oxygen to form hydroperoxyl radicals (HO2):
Then, these radicals react rapidly to oxidize NO into NO2:
The OH radical has been regenerated, and the total budget of these three reactions is then as follows:
Therefore, the oxidation of CO into CO2 leads to the conversion of NO into NO2. NO2 can be photolyzed to form NO and O3. Since NO is converted to NO2 without consumption of O3 (unlike what happens in the Leighton photostationary state), there is formation of a molecule of O3 for each molecule of CO that is oxidized. This yield is theoretical because all HO2 radicals do not react with NO and some OH radicals may react with NO2. The actual yield is, therefore, less than 1. The presence of a carbonaceous species (here CO, but VOC play a similar role, see Sections 8.3.4 to 8.3.8) leads to a perturbation of the photostationary equilibrium, thereby allowing the formation of NO2 without O3 consumption and leading, therefore, to O3 formation.
Example: Calculation of the ozone concentration produced from carbon monoxide in presence of nitrogen oxides
The initial concentration of CO is 1 ppm and its oxidation occurs over an 8-hour period. The OH radical concentration is assumed to be 106 cm−3.
The rate constant of the oxidation of CO by OH is 0.35 ppb−1 min−1 at 1 atm and 25 °C. The OH concentration in ppb is: 106 / (2.46 × 1010) = 4 × 10−5 ppb. Formation of O3 over eight hours is theoretically equivalent to the amount of CO that has reacted. Therefore:
Thus, 1 ppm of CO (which is not very reactive) has formed 6 ppb of O3 in eight hours. Therefore, it is the presence of volatile carbonaceous compounds (here CO, but also reactive VOC), which leads to ozone formation, as correctly identified originally by Haagen-Smit.
8.3.4 Photolysis and Oxidation of Aldehydes
The simplest aldehyde (i.e., the aldehyde with only one carbon atom) is formaldehyde (HCHO). Its oxidation occurs by photolysis or by reaction with OH. There are two pathways for the photolysis of formaldehyde. On one hand:
Thus, the overall budget is:
On the other hand:
The two products of this reaction are stable molecules (CO will of course be oxidized slowly as described in Section 8.3.3). These two photolysis reactions have similar kinetics. Therefore, one may write a simple overall budget as being the average of these two reactions:
Oxidation by OH leads to the following reactions:
The OH radical abstracts a hydrogen atom from the formaldehyde molecule to form a stable molecule (water vapor, H2O) and an unstable organic radical. Therefore, one obtains an overall budget that is similar to that obtained with the photolytic reactions:
Thus, the oxidation of HCHO, whether it occurs by reaction with OH or results from photolytic reactions leads to one molecule of CO and one HO2 radical. As shown in Section 8.3.3, the HO2 radical can later oxidize NO into NO2 and form an OH radical. The photolysis of NO2 leads to the formation of a molecule of O3. Since the oxidation of a molecule of CO leads also to the formation of a molecule of O3, the oxidation of HCHO can theoretically lead to the formation of two molecules of O3.
The oxidation of higher aldehydes (i.e., aldehydes with more than one carbon atom) follows the same conceptual scheme as that of formaldehyde, but leads to more complex products due to the greater number of carbon atoms. For example, the reactions of acetaldehyde (two carbon atoms, CH3CHO) are as follows:
The overall budget is as follows:
In this first oxidation step, H originating from formaldehyde has been replaced by CH3 originating from acetaldehyde. The methylperoxyl radical (also called peroxymethyl), CH3O2, behaves similarly to HO2, that is to say that it can oxidize NO into NO2 and form a methoxy radical, CH3O: