8.6. Stratospheric Ozone

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8.6. Stratospheric Ozone

The ozone balance in the stratosphere is determined through complex interactions of solar radiation, meteorological movements within the stratosphere, transport to and from the troposphere, and the concentration of species based on elements other than oxygen that enter the stratosphere by natural or artificial means (such as flight of aircraft).

It is not difficult to see that ozone initially forms from the oxygen present in the air. Chapman [121] introduced the photochemical model of stratospheric ozone and suggested that the ozone mechanism depended on two photochemical and two chemical reactions:

$O_{2} + hv \to O + O$ $O_{2} + hv \to O + O$

(8.148)

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

(8.149)

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

(8.150)

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

(8.151)

Reactions (8.148) and (8.149) are the reactions by which the ozone is generated. Reactions (8.150) and (8.151) establish the balance, which is the ozone concentration in the troposphere. If one adds reactions (8.150) and (8.151), one obtains the overall rate of destruction of the ozone: namely:

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

(8.150)

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

(8.151)

$\bar{Net : 2 O_{3} + hv \to 3 O_{2}}$ $\bar{Net : 2 O_{3} + hv \to 3 O_{2}}$

(8.152)

The rates of reactions (8.148)–(8.151) vary with altitude. The rate constants of reactions (8.148) and (8.150) are determined by the solar flux at a given altitude, and the rate constants of the other reactions are determined by the temperature at that altitude. However, precise solar data obtained from rocket experiments and better kinetic data for reactions (8.149)–(8.151), coupled with recent meteorological analysis have shown that the Chapman model was seriously flawed. The concentrations predicted by the model were essentially too high. Something else was affecting the ozone.

8.6.1. The HO_x Catalytic Cycle

Hunt [122,123] suggested that excited electronic states of O atoms and O₂ might account for the discrepancy between the Chapman model and the measured meteorological ozone distributions. But he showed that reactions based on these excited species were too slow to account for the differences sought. Realizing that water could enter the stratosphere, Hunt considered the reactions of free radicals (H, HO, and HOO) derived from water. Consistent with the shorthand developed for the oxides of nitrogen, these radicals are specified by the chemical formula HO_x. The mechanism that Hunt postulated was predicated on the formation of hydroxyl radicals. The photolysis of ozone by ultraviolet radiation below 310 nm produces excited singlet oxygen atoms which react rapidly with water to form hydroxyl radicals:

$O_{3} + hv \to O_{2} + O ({}^{1}D)$ $O_{3} + hv \to O_{2} + O ({}^{1}D)$

(8.153)

$O ({}^{1}D) + H_{2} O \to 2 OH$ $O ({}^{1}D) + H_{2} O \to 2 OH$

(8.154)

Only an excited singlet oxygen atom could react with water at stratospheric temperatures to form hydroxyl radicals.

At these temperatures, singlet oxygen atoms could also react with hydrogen or methane to form OH. The OH reacts with O₃ to produce hydroperoxy radicals HO₂. Both HO and HO₂ destroy ozone by an indirect reaction which sometimes involves O atoms:

$HO + O_{3} \to H O_{2} + O_{2}$ $HO + O_{3} \to H O_{2} + O_{2}$

(8.155)

$H O_{2} + O_{3} \to HO + O_{2} + O_{2}$ $H O_{2} + O_{3} \to HO + O_{2} + O_{2}$

(8.156)

$HO + O \to H + O_{2}$ $HO + O \to H + O_{2}$

(8.157)

$H O_{2} + O \to HO + O_{2}$ $H O_{2} + O \to HO + O_{2}$

(8.158)

$H + O_{3} \to HO + O_{2}$ $H + O_{3} \to HO + O_{2}$

(8.159)

Numerous reactions of HO₂ radicals are possible in the stratosphere. The essential reactions for the discussion of the ozone balance are:

$HO + O_{3} \to H O_{2} + O_{2}$ $HO + O_{3} \to H O_{2} + O_{2}$

(8.155)

$H O_{2} + O_{3} \to HO + O_{2} + O_{2}$ $H O_{2} + O_{3} \to HO + O_{2} + O_{2}$

(8.156)

$\bar{Net : 2 O_{3} \to 3 O_{2}}$ $\bar{Net : 2 O_{3} \to 3 O_{2}}$

The reaction sequence (8.155) and (8.156) is a catalytic chain for ozone destruction and contributes to the net destruction. However, even given the uncertainty possible in the rates of these reactions and the uncertainty of the air motions, this system could not explain the imbalance in the ozone throughout the stratosphere.

8.6.2. The NO_x Catalytic Cycle

In the late 1960s, direct observations of substantial amounts (3 ppb) of nitric acid vapor in the stratosphere were reported. Crutzen [123] reasoned that if HNO₃ vapor is present in the stratosphere, it could be broken down to a degree to the active oxides of nitrogen NO_x (NO and NO₂) and that these oxides could form a catalytic cycle (or the destruction of the ozone). Johnston and Whitten [124] first realized that if this were so, then supersonic aircraft flying in the stratosphere could wreak harm to the ozone balance in the stratosphere. Much of what appears in this section is drawn from an excellent review by Johnston and Whitten [125]. The most pertinent of the possible NO_x reactions in the atmosphere are:

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

(8.3)

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

(8.160)

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

(8.1)

whose rate constants are now well known. The reactions combine in two different competing cycles. The first is catalytic destructive:

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

(8.3)

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

(8.160)

$\bar{Net : O_{3} + O \to O_{2} + O_{2}}$ $\bar{Net : O_{3} + O \to O_{2} + O_{2}}$

and the second, parallel cycle is essentially a “do nothing” one:

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

(8.3)

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

(8.1)

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

(8.149)

$\bar{Net : no chemical change}$ $\bar{Net : no chemical change}$

The rate of destruction of ozone with the oxides of nitrogen relative to the rate in pure air (Chapman model) is defined as the catalytic ratio, which may be expressed either in terms of the variables (NO₂) and (O₃) or (NO) and (O). These ratio expressions are:

$β = \frac{rate of ozone destruction with N O_{x}}{rate of ozone destruction in pure air}$ $β = \frac{rate of ozone destruction with N O_{x}}{rate of ozone destruction in pure air}$

(8.161)

$β = 1 + {k_{160} (N O_{2}) / k_{151} (O_{3})}$ $β = 1 + {k_{160} (N O_{2}) / k_{151} (O_{3})}$

(8.162)

$β = 1 + {(k_{3} k_{160} / k_{151} j_{1}) (NO)} / {1 + k_{160} (O) / j_{1}}$ $β = 1 + {(k_{3} k_{160} / k_{151} j_{1}) (NO)} / {1 + k_{160} (O) / j_{1}}$

(8.163)

Here, as throughout this book, the k's are the specific rate constants of the chemical reactions, while the j's are the specific rate constants of the photochemical reactions.

At low elevations, where the oxygen atom concentration is low and the NO₂ cycle is slow, another catalytic cycle derived from the oxides of nitrogen may be important:

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

(8.164)

$N O_{3} + hv (visible, day) \to NO + O_{2}$ $N O_{3} + hv (visible, day) \to NO + O_{2}$

(8.165)

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

(8.3)

$\bar{Net : 2 O_{3} + hv \to 3 O_{2} (day)}$ $\bar{Net : 2 O_{3} + hv \to 3 O_{2} (day)}$

The radiation involved here is red light, which is abundant at all elevations. Reaction (8.164) permits another route at night (including the polar night), which converts a few percent of NO₂ to N₂O₅:

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

(8.164)

$N O_{2} + N O_{3} + M \to N_{2} O_{5} + M (night)$ $N O_{2} + N O_{3} + M \to N_{2} O_{5} + M (night)$

(8.166)

The rate of reaction (8.164) is known accurately only at room temperature, and extrapolation to stratospheric temperature is uncertain; nevertheless, the extrapolated values indicate that the NO₃ catalytic cycle [reactions (8.164) and (8.165)] destroys ozone faster than the NO₂ cycle below 22 km and in the region where the temperature is at least 220 K.

The nitric acid cycle is established by the reactions

$HO + N O_{2} + M \to HN O_{3} + M$ $HO + N O_{2} + M \to HN O_{3} + M$

(8.167)

$HN O_{3} + hv \to OH + N O_{2}$ $HN O_{3} + hv \to OH + N O_{2}$

(8.168)

$HO + HN O_{3} \to H_{2} O + N O_{3}$ $HO + HN O_{3} \to H_{2} O + N O_{3}$

(8.169)

The steady-state ratio of nitrogen dioxide concentration to nitric acid can be readily found to be

${[(N O_{2}) / (HN O_{3})]}_{ss} = [k_{169} / k_{167}] + [j_{168} / k_{167} (OH)]$ ${[(N O_{2}) / (HN O_{3})]}_{ss} = [k_{169} / k_{167}] + [j_{168} / k_{167} (OH)]$

(8.170)

For the best data available for the hydroxyl radical concentration and the rate constants, the ratio has the values

$0.1 at 15 km, 1 at 25 km, > 1 at 35 km$ $0.1 at 15 km, 1 at 25 km, > 1 at 35 km$

Thus, it can be seen that nitric acid is a significant reservoir or sink for the oxides of nitrogen. In the lowest stratosphere, the nitric acid predominates over the NO₂, and a major loss of NO_x from the stratosphere occurs by diffusion of the acid into the troposphere where it is rained out.

By using the HO_x and NO_x cycles just discussed and by assuming an NO_x concentration of 4.2 × 10⁹ molecules/cm³ distributed uniformly through the stratosphere, Johnston and Whitten [124] were able to make the most reasonable prediction of the ozone balance in the stratosphere. Measurements of the concentration of NO_x in the stratosphere show a range of 2 − 8 × 10⁹ molecules/cm³.

It is possible to similarly estimate the effect of the various cycles upon ozone destruction. The results can be summarized as follows: Between 15 and 20 km, the NO₃ catalytic cycle dominates; between 20 and 40 km, the NO₂ cycle dominates; between 40 and 45 km, the NO₂, HO_x, and O_x mechanisms are about equal; and above 45 km, the HO_x reactions are the controlling reactions.

It appears that between 15 and 35 km, the oxides of nitrogen are by far the most important agents for maintaining the natural ozone balance. Calculations show that the natural NO_x should be about 4 × 10⁹ molecules/cm³. The extent to which this concentration would be modified by anthropogenic sources such as supersonic aircraft determines the extent of the danger to the normal ozone balance. It must be stressed that this question is a complex one, since both concentration and distribution are involved (see Johnston and Whitten [124]).

8.6.3. The ClO_x Catalytic Cycle

Molina and Rowland [125] pointed out that fluorocarbons that diffuse into the stratosphere could also act as potential sinks for ozone. Cicerone et al. [126] showed that the effect of these synthetic chemicals could last for decades. Thus, possibly, a major source of atmospheric contamination arises by virtue of the widespread use of fluorocarbons as propellants and refrigerants. Approximately 80% of all fluorocarbons released to the atmosphere derive from these sources. There is no natural source.

Eighty-five percent of all fluorocarbons are F11 (CC1₃F) or F12 (CC1₂F₂). According to Molina and Rowland [125], these fluorocarbons are removed from the stratosphere by photolysis above altitudes of 25 km. The primary reactions are

$CC l_{3} F + hv \to CC l_{2} F + Cl$ $CC l_{3} F + hv \to CC l_{2} F + Cl$

(8.171)

$CC l_{2} F + hv \to CClF + Cl$ $CC l_{2} F + hv \to CClF + Cl$

(8.172)

Subsequent chemistry leads to release of additional chlorine, and for purposes of discussion, it is here assumed that all of the available chlorine is eventually liberated in the form of compounds such as HCl, C1O, C1O₂, and C1₂. The catalytic chain for ozone that develops is

$Cl + O_{3} \to ClO + O_{2}$ $Cl + O_{3} \to ClO + O_{2}$

(8.173)

$ClO + O \to Cl + O_{2}$ $ClO + O \to Cl + O_{2}$

(8.174)

$\bar{Net : O_{3} + O \to O_{2} + O_{2}}$ $\bar{Net : O_{3} + O \to O_{2} + O_{2}}$

Other reactions that are important in affecting the chain are

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

(8.175)

$Cl + H O_{2} \to HCl + O_{2}$ $Cl + H O_{2} \to HCl + O_{2}$

(8.176)

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

(8.177)

$ClO + O_{3} \to Cl O_{2} + O_{2}$ $ClO + O_{3} \to Cl O_{2} + O_{2}$

(8.178)

Table 8.8

Residence Time of Halocarbons in the Troposphere^a

Halocarbon	Average Residence Time in Years
Chloroform (CHCl₃)	0.19
Methylene chloride (CH₂Cl₂)	0.30
Methyl chloride (CH₃Cl)	0.37
1,1,1-Trichloroethane (CH₃CCl₃)	1.1
F12	330 or more
Carbon tetrachloride (CCl₄)	330 or more
F11	1000 or more

^a Based on reaction with OH radicals.

$Cl + C H_{4} \to HCl + C H_{3}$ $Cl + C H_{4} \to HCl + C H_{3}$

(8.179)

$Cl + N O_{2} + M \to ClN O_{2} + M$ $Cl + N O_{2} + M \to ClN O_{2} + M$

(8.180)

$ClN O_{2} + O \to ClNO + O_{2}$ $ClN O_{2} + O \to ClNO + O_{2}$

(8.181)

$ClN O_{2} + hv \to ClNO + O$ $ClN O_{2} + hv \to ClNO + O$

(8.182)

$ClO + N O_{2} + M \to ClON O_{2} + M$ $ClO + N O_{2} + M \to ClON O_{2} + M$

(8.183)

The unique problem that arises here is that F11 and F12 are relatively inert chemically and have no natural sources or sinks, as CC1₄ does. The lifetimes of these fluorocarbons are controlled by diffusion to the stratosphere, where photodissociation takes place as designated by reactions (8.171) and (8.172). The lifetimes of halogen species in the atmosphere are given in Ref. [127]. These values are reproduced in Table 8.8. The incredibly long lifetimes of F11 and F12 and their gradual diffusion into the stratosphere pose the problem. Even though use of these materials has virtually stopped today, their effects are likely to be felt for decades.

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Table of Contents for 8.6. Stratospheric Ozone

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8.6. Stratospheric Ozone

8.6.1. The HOx Catalytic Cycle

8.6.2. The NOx Catalytic Cycle

8.6.3. The ClOx Catalytic Cycle

Table of Contents for
8.6. Stratospheric Ozone

8.6.1. The HO_x Catalytic Cycle

8.6.2. The NO_x Catalytic Cycle

8.6.3. The ClO_x Catalytic Cycle