8.2. The Nature of Photochemical Smog

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8.2. The Nature of Photochemical Smog

Photochemical air pollution consists of a complex mixture of gaseous pollutants and aerosols, some of which are photochemically produced. Among the gaseous compounds are the oxidizing species ozone, nitrogen dioxide, and peroxyacyl nitrate:

The member of this series most commonly found in the atmosphere is peroxyacetyl nitrate (PAN):

The three compounds, O₃, NO₂, and PAN, are often grouped together and called photochemical oxidant.

Photochemical smog is composed of mixtures of particulate matter and noxious gases, similar to those that occurred in the typical London-type “pea soup” smog. The London smog was a mixture of particulates and oxides of sulfur, chiefly sulfur dioxide. But the overall system in the London smog was chemically reducing in nature. This difference in redox chemistry between photochemical oxidant and SO_x-particulate smog is important in several respects. Note in particular the problem of quantitatively detecting oxidant in the presence of sulfur dioxide. Because it is a reducing agent, SO_x tends to reduce the oxidizing effects of ozone and thus produces low quantities of the oxidant.

In dealing with the heterogeneous gas−liquid−solid mixture characterized as photochemical smog, it is important to realize from a chemical as well as a biological point of view that synergistic effects may occur.

8.2.1. Primary and Secondary Pollutants

Primary pollutants are those emitted directly to the atmosphere whereas secondary pollutants are those formed by chemical or photochemical reactions of primary pollutants after they have been admitted to the atmosphere and exposed to sunlight. Unburned hydrocarbons, NO, particulates, and the oxides of sulfur are examples of primary pollutants. The particulates may be lead oxide from the oxidation of tetraethyllead in automobiles, fly ash, and various types of carbonaceous solids. Peroxyacyl nitrate and ozone are examples of secondary pollutants.

Some pollutants fall into both categories. Nitrogen dioxide, which is emitted directly from auto exhaust, is also formed in the atmosphere photochemically from NO. Aldehydes, which are released in auto exhausts, are also formed in the photochemical oxidation of hydrocarbons. Carbon monoxide, which arises primarily from autos and stationary sources, is likewise a product of atmospheric hydrocarbon oxidation.

8.2.2. The Effect of NO_x

It has been well established that if a laboratory chamber containing NO, a trace of NO₂, and air is irradiated with ultraviolet light, the following reactions occur:

$N O_{2} + hv (3000 \overset{\circ}{A} \leq λ \leq 4200 \overset{\circ}{A}) \to NO + O ({}^{3}P)$ $N O_{2} + hv (3000 \overset{\circ}{A} \leq λ \leq 4200 \overset{\circ}{A}) \to NO + O ({}^{3}P)$

(8.1)

$O + O_{2} + M \to O_{3} + M$ $O + O_{2} + M \to O_{3} + M$

(8.2)

$O_{3} + NO \to O_{2} + N O_{2}$ $O_{3} + NO \to O_{2} + N O_{2}$

(8.3)

The net effect of irradiation on this inorganic system is to establish the dynamic equilibrium

$N O_{2} + O_{2} \overset{hv}{\leftrightarrow} NO + O_{3}$ $N O_{2} + O_{2} \overset{hv}{\leftrightarrow} NO + O_{3}$

(8.4)

However, if a hydrocarbon, particularly an olefin or an alkylated benzene, is added to the chamber, the equilibrium represented by reaction (8.4) is unbalanced and the following events take place:

1. The hydrocarbons are oxidized and disappear.

2. Reaction products such as aldehydes, nitrates, PAN, etc., are formed.

3. NO is converted to NO₂.

4. When all the NO is consumed, O₃ begins to appear. On the other hand, PAN and other aldehydes are formed from the beginning.

Basic rate information permits one to examine these phenomena in detail. Leighton [2], in his excellent book Photochemistry of Air Pollution, gives numerous tables of rates and products of photochemical nitrogen oxide–hydrocarbon reactions in air; this early work is followed here to give fundamental insight into the photochemical smog problem. The data in these tables show low rates of photochemical consumption of the saturated hydrocarbons, compared with the unsaturates, and the absence of aldehydes in the products of the saturated hydrocarbon reactions. These data conform to the relatively low rate of reaction of the saturated hydrocarbons with oxygen atoms and their inertness with respect to ozone.

Among the major products in the olefin reactions are aldehydes and ketones. Such results correspond to the splitting of the double bond and the addition of an oxygen atom to one end of the olefin.

Irradiation of mixtures of an olefin with nitric oxide and nitrogen dioxide in air shows that the nitrogen dioxide rises in concentration before it is eventually consumed by reaction. Since the photodissociation of the nitrogen dioxide initiates the reaction, it would appear that a negative quantum yield results. More likely, the nitrogen dioxide is being formed by secondary reactions more rapidly than it is being photodissociated.

The important point is that this negative quantum yield is realized only when an olefin (hydrocarbon) is present. Thus, adding the overall step

$\begin{matrix} O \\ O_{3} \end{matrix}} + olefin \to products$ $\begin{matrix} O \\ O_{3} \end{matrix}} + olefin \to products$

(8.5)

to reactions (8.1)–(8.3) would not be an adequate representation of the atmosphere photochemical reactions. However, if one assumes that O₃ attains a steady-state concentration in the atmosphere, one can perform a steady-state analysis (see Chapter 2, Section 2.4) with respect to O₃. Furthermore, if one assumes that O₃ is largely destroyed by reaction (8.3), one obtains a useful approximate relationship:

$(O_{3}) = - (j_{1} / k_{3}) (N O_{2}) / (NO)$ $(O_{3}) = - (j_{1} / k_{3}) (N O_{2}) / (NO)$

where j is the rate constant for the photochemical reaction. Thus, the O₃ steady-state concentration in a polluted atmosphere is seen to increase with decreasing concentration of nitric oxide, and vice versa. The ratio of j₁/k₃ approximately equals 1.2 ppm for the Los Angeles noonday condition [2]. Reactions such as

$O + N O_{2} \to NO + O_{2}$ $O + N O_{2} \to NO + O_{2}$

$O + N O_{2} + M \to N O_{3} + M$ $O + N O_{2} + M \to N O_{3} + M$

$N O_{3} + NO \to 2 N O_{2}$ $N O_{3} + NO \to 2 N O_{2}$

$O + NO + M \to N O_{2} + M$ $O + NO + M \to N O_{2} + M$

$2 NO + O_{2} \to 2 N O_{2}$ $2 NO + O_{2} \to 2 N O_{2}$

$N O_{3} + N O_{2} \to N_{2} O_{5}$ $N O_{3} + N O_{2} \to N_{2} O_{5}$

$N_{2} O_{5} \to N O_{3} + N O_{2}$ $N_{2} O_{5} \to N O_{3} + N O_{2}$

do not play a part. They are generally too slow to be important.

Furthermore, it has been noted that when the rate of the oxygen atom–olefin reaction and the rate of the ozone–olefin reaction are totaled, they do not give the complete hydrocarbon consumption. This anomaly is also an indication of an additional process.

An induction period with respect to olefin consumption is also observed in the photochemical laboratory experiments, thus indicating the buildup of an intermediate. When illumination is terminated in these experiments, the excess rate over the total of the O and O₃ reactions disappears. These and other results suggest that the intermediate formed is photolyzed and contributes to the concentration of the major species of concern.

Possible intermediates that fulfill the requirements of the laboratory experiments are alkyl and acyl nitrites and pernitrites. The second photolysis effect eliminates the possibility that aldehydes serve as the intermediate.

Various mechanisms have been proposed to explain these laboratory results. The following low-temperature (atmospheric) sequence based on isobutene as the initial fuel was first proposed by Leighton [2] and appears to account for most of what has been observed:

$O + C_{4} H_{8} \to C H_{3} + C_{3} H_{5} O$ $O + C_{4} H_{8} \to C H_{3} + C_{3} H_{5} O$

(8.6)

$C H_{3} + O_{2} \to C H_{3} OO$ $C H_{3} + O_{2} \to C H_{3} OO$

(8.7)

$C H_{3} OO + O_{2} \to C H_{3} O + O_{3}$ $C H_{3} OO + O_{2} \to C H_{3} O + O_{3}$

(8.8)

$O_{3} + NO \to N O_{2} + O_{2}$ $O_{3} + NO \to N O_{2} + O_{2}$

(8.9)

$C H_{3} O + NO \to C H_{3} ONO$ $C H_{3} O + NO \to C H_{3} ONO$

(8.10)

$C H_{3} ONO + h v \to C H_{3} O^{*} + NO$ $C H_{3} ONO + h v \to C H_{3} O^{*} + NO$

(8.11)

$C H_{3} O^{*} + O_{2} \to H_{2} CO + HOO$ $C H_{3} O^{*} + O_{2} \to H_{2} CO + HOO$

(8.12)

$HOO + C_{4} H_{8} \to H_{2} CO + {(C H_{3})}_{2} CO + H$ $HOO + C_{4} H_{8} \to H_{2} CO + {(C H_{3})}_{2} CO + H$

(8.13)

$H + O_{2} + M \to HOO + M$ $H + O_{2} + M \to HOO + M$

(8.14)

$HOO + NO \to OH + N O_{2}$ $HOO + NO \to OH + N O_{2}$

(8.15)

$OH + C_{4} H_{8} \to {(C H_{3})}_{2} CO + C H_{3}$ $OH + C_{4} H_{8} \to {(C H_{3})}_{2} CO + C H_{3}$

(8.16)

$C H_{3} + O_{2} \to C H_{3} OO (as above)$ $C H_{3} + O_{2} \to C H_{3} OO (as above)$

(8.17)

$2 HOO \to H_{2} O_{2} + O_{2}$ $2 HOO \to H_{2} O_{2} + O_{2}$

(8.18)

$2 OH \to H_{2} + O_{2}$ $2 OH \to H_{2} + O_{2}$

(8.19)

$HOO + H_{2} \to H_{2} O + OH$ $HOO + H_{2} \to H_{2} O + OH$

(8.20)

$HOO + H_{2} \to H_{2} O_{2} + H$ $HOO + H_{2} \to H_{2} O_{2} + H$

(8.21)

There are two chain-propagating sequences (reactions (8.13) and (8.14) and reactions (8.15)–(8.17)) and one chain-breaking sequence (reactions (8.18) and (8.19)). The intermediate is the nitrite, as shown in reaction (8.10). Reaction (8.11) is the required additional photochemical step. For every NO₂ used to create the O atom of reaction (8.6), one is formed by reaction (8.9). However, reactions (8.10), (8.11), and (8.15) reveal that for every two NO molecules consumed, one NO and one NO₂ form—hence, the negative quantum yield of NO₂.

With other olefins, other appropriate reactions may be substituted. Ethylene would give

$O + C_{2} H_{4} \to C H_{3} + HCO$ $O + C_{2} H_{4} \to C H_{3} + HCO$

(8.22)

$HOO + C_{2} H_{4} \to 2 H_{2} CO + H$ $HOO + C_{2} H_{4} \to 2 H_{2} CO + H$

(8.23)

$OH + C_{2} H_{4} \to H_{2} CO + C H_{3}$ $OH + C_{2} H_{4} \to H_{2} CO + C H_{3}$

(8.24)

Propylene would add

$OH + C_{3} H_{6} \to C H_{3} CHO + C H_{3}$ $OH + C_{3} H_{6} \to C H_{3} CHO + C H_{3}$

(8.25)

Thus, PAN would form from

$C H_{3} CHO + O_{2} \overset{h v}{\to} C H_{3} CO + HOO$ $C H_{3} CHO + O_{2} \overset{h v}{\to} C H_{3} CO + HOO$

(8.26)

$C H_{3} CO + O_{2} \to C H_{3} (CO) OO$ $C H_{3} CO + O_{2} \to C H_{3} (CO) OO$

(8.27)

$C H_{3} (CO) OO + N O_{2} \to C H_{3} (CO) OON O_{2}$ $C H_{3} (CO) OO + N O_{2} \to C H_{3} (CO) OON O_{2}$

(8.28)

and an acid could form from the overall reaction

$C H_{3} (CO) OO + 2 C H_{3} CHO \to C H_{3} (CO) OH + 2 C H_{3} CO + OH$ $C H_{3} (CO) OO + 2 C H_{3} CHO \to C H_{3} (CO) OH + 2 C H_{3} CO + OH$

(8.29)

Because pollutant concentrations are generally in the parts-per-million range, it is not difficult to postulate many types of reactions and possible products.

8.2.3. The Effect of SO_x

Historically, the sulfur oxides have long been known to have a deleterious effect on the atmosphere, and sulfuric acid mist and other sulfate particulate matter are well established as important sources of atmospheric contamination. However, the atmospheric chemistry is probably not as well understood as the gas-phase photoxidation reactions of the nitrogen oxides−hydrocarbon system. The pollutants form originally from the SO₂ emitted to the air. Just as mobile and stationary combustion sources emit some small quantities of NO₂ as well as NO, so do they emit some small quantities of SO₃ when they burn sulfur-containing fuels. Leighton [2] also discussed the oxidation of SO₂ in polluted atmospheres, and an excellent review by Bulfalini [3] has appeared. This section draws heavily from these sources.

The chemical problem here involves the photochemical and catalytic oxidation of SO₂ and its mixtures with the hydrocarbons and NO; however the primary concern is the photochemical reactions, both gas-phase and aerosol-forming.

The photodissociation of SO₂ into SO and O atoms is markedly different from the photodissociation of NO₂. The bond to be broken in the sulfur compound requires about 560 kJ/mol. Thus, wavelengths greater than 2180 Å do not have sufficient energy to initiate dissociation. This fact is significant in that only solar radiation greater than 2900 Å reaches the lower atmosphere. If a photochemical effect is to occur in the SO₂−O₂ atmospheric system, it must be that the radiation electronically excites the SO₂ molecule but does not dissociate it.

There are two absorption bands of SO₂ within the range 3000–4000 Å. The first is a weak absorption band and corresponds to the transition to the first excited state (a triplet). This band originates at 3880 Å and has a maximum around 3840 Å. The second is a strong absorption band and corresponds to the excitation to the second excited state (a triplet). This band originates at 3376 Å and has a maximum around 2940 Å.

Blacet [4], who carried out experiments in high O₂ concentrations, reported that ozone and SO₃ appear to be the only products of the photochemically induced reaction. The following essential steps were postulated:

$S O_{2} + h v \to S O_{2}^{*}$ $S O_{2} + h v \to S O_{2}^{*}$

(8.30)

$S O_{2}^{*} + O_{2} \to S O_{4}$ $S O_{2}^{*} + O_{2} \to S O_{4}$

(8.31)

$S O_{4} + O_{2} \to S O_{3} + O_{3}$ $S O_{4} + O_{2} \to S O_{3} + O_{3}$

(8.32)

The radiation used was at 3130 Å, and it would appear that the excited

S O_{2}^{*}

$S O_{2}^{*}$

in reaction (8.30) is a singlet. The precise roles of the excited singlet and triplet states in the photochemistry of SO₂ are still unclear [3]. Nevertheless, this point need not be of great concern since it is possible to write the reaction sequence

$S O_{2} + h v \to {}^{1}S O_{2}^{*}$ $S O_{2} + h v \to {}^{1}S O_{2}^{*}$

(8.33)

${}^{1}S O_{2}^{*} + S O_{2} \to {}^{3}S O_{2}^{*} + S O_{2}$ ${}^{1}S O_{2}^{*} + S O_{2} \to {}^{3}S O_{2}^{*} + S O_{2}$

(8.34)

Thus, reaction (8.30) could specify either an excited singlet or triplet

S O_{2}^{*}

$S O_{2}^{*}$

. The excited state may, of course, degrade by internal transfer to a vibrationally excited ground state that is later deactivated by collision, or it may be degraded directly by collisions. Fluorescence of SO₂ has not been observed above 2100 Å. The collisional deactivation steps known to exist in laboratory experiments are not listed here, to minimize the writing of reaction steps.

Since they involve one species in large concentrations, reactions (8.30)–(8.32) are the primary ones for the photochemical oxidation of SO₂ to SO₃. A secondary reaction route to SO₃ could be

$S O_{4} + S O_{2} \to 2 S O_{3}$ $S O_{4} + S O_{2} \to 2 S O_{3}$

(8.35)

In the presence of water a sulfuric acid mist forms according to

$H_{2} O + S O_{3} \to H_{2} S O_{4}$ $H_{2} O + S O_{3} \to H_{2} S O_{4}$

(8.36)

The SO₄ molecule formed by reaction (8.31) would probably have a peroxy structure, and if

S O_{2}^{*}

$S O_{2}^{*}$

were a triplet, it might be a biradical.

There is conflicting evidence with respect to the results of the photolysis of mixtures of SO₂, NO_x, and O₂. However, many believe that the following should be considered with the NO_x photolysis reactions:

$S O_{2} + NO \to SO + N O_{2}$ $S O_{2} + NO \to SO + N O_{2}$

(8.37)

$S O_{2} + N O_{2} \to S O_{3} + NO$ $S O_{2} + N O_{2} \to S O_{3} + NO$

(8.38)

$S O_{2} + O + M \to S O_{3} + M$ $S O_{2} + O + M \to S O_{3} + M$

(8.39)

$S O_{2} + O_{3} \to S O_{3} + O_{2}$ $S O_{2} + O_{3} \to S O_{3} + O_{2}$

(8.40)

$S O_{3} + O \to S O_{2} + O_{2}$ $S O_{3} + O \to S O_{2} + O_{2}$

(8.41)

$S O_{4} + NO \to S O_{3} + N O_{2}$ $S O_{4} + NO \to S O_{3} + N O_{2}$

(8.42)

$S O_{4} + N O_{2} \to S O_{3} + N O_{3}$ $S O_{4} + N O_{2} \to S O_{3} + N O_{3}$

(8.43)

$S O_{4} + O \to S O_{3} + O_{2}$ $S O_{4} + O \to S O_{3} + O_{2}$

(8.44)

$SO + O + M \to S O_{2} + M$ $SO + O + M \to S O_{2} + M$

(8.45)

$SO + O_{3} \to S O_{2} + O_{2}$ $SO + O_{3} \to S O_{2} + O_{2}$

(8.46)

$SO + N O_{2} \to S O_{2} + NO$ $SO + N O_{2} \to S O_{2} + NO$

(8.47)

The important reducing effect of the SO₂ with respect to different polluted atmospheres mentioned in the introduction of this section becomes evident from these reactions.

Some work [5] has been performed on the photochemical reaction between sulfur dioxide and hydrocarbons, both paraffins and olefins. In all cases, mists were found, and these mists settled out in the reaction vessels as oils with the characteristics of sulfuric acids. Because of the small amounts of materials formed, great problems arise in elucidating particular steps. When NO_x and O₂ are added to this system, the situation is most complex. Bulfalini [3] summed up the status in this way: “The aerosol formed from mixtures of the lower hydrocarbons with NO_x and SO₂ is predominantly sulfuric acid, whereas the higher olefin hydrocarbons appear to produce carbonaceous aerosols also, possibly organic acids, sulfonic or sulfuric acids, nitrate-esters, etc.”

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Table of Contents for 8.2. The Nature of Photochemical Smog

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8.2. The Nature of Photochemical Smog

8.2.1. Primary and Secondary Pollutants

8.2.2. The Effect of NOx

8.2.3. The Effect of SOx

Table of Contents for
8.2. The Nature of Photochemical Smog

8.2.2. The Effect of NO_x

8.2.3. The Effect of SO_x