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1.4.1

Principle ozone production and destruction reac-

tions

In the stratosphere the principle chemical reaction that leads to ozone production is the photolysis of molecular oxygen by UV solar radiation with wavelengths less than 250 nm:

O2 +hν → O + O (λ < 242 nm),

where hν represents the energy of a photon (with h being Planck’s Constant and ν being the frequency) and λis the wavelength. The oxygen atoms combine

rapidly with molecular oxygen in the presence of a third reactant molecule, M (usually N2,Staehelin et al., 2001), needed to balance energy and momentum in

the reaction, to produce ozone:

O + O2 + M → O3 + M.

However, sunlight with wavelength<850 nm also dissociates ozone molecules by the equation:

O3 +hν → O2 + O (λ <850 nm).

Since the oxygen atoms have very short lifetimes, they either quickly combine with O2 form ozone again, or react with an ozone atom to form two oxygen

molecules again:

O + O3 → 2 O2.

This cycle of ozone production and destruction was proposed in 1930 by Sidney Chapman and is called theChapman Cycle. If not perturbed, it produces and destroys ozone in the stratosphere in equal amounts (Reid, 2000), such that an equilibrium of O, O2, O3 and O4 is maintained.

1.4.2

Catalytic cycles

Calculating the ozone content of the stratosphere according to the Chapman Cy- cle, however, overestimates the ozone maximum by a factor of two compared to observations. One of the main reasons for this is the neglect of other ozone- destroying reactions, thecatalytic cycles (Andrews, 2000). The two main reac- tions in a catalytic destruction cycle are:

X + O3 → XO + O2

XO + O → X + O2.

The net effect of a catalytic cycle therefore is: O + O3 → 2 O2,

which has the same effect as the ozone destruction path suggested by Chapman. X in the above equations is a catalyst atom, molecule or radical which takes part in a reaction but is not itself consumed by it. One single catalyst molecule can therefore destroy many ozone molecules before it breaks the catalytic cycle by reacting with some other atom or molecule. The most common species for X belong to the families of odd-nitrogen (NOx), odd-hydrogen (HOx), odd-chlorine (Clx) (Reid, 2000), and odd-bromine (Brx) (McConnell and Jin, 2008), and are NO, OH, H, Cl and Br. Catalytic cycles are most important in the stratosphere at mid-latitudes (McConnell and Jin, 2008), although their importance varies at different atmospheric levels: in the lower stratosphere (15-25 km) all four catalytic families can contribute to ozone destruction, whereas at higher altitudes (25- 40 km) the destruction is dominated by the odd-nitrogen family (Dessler, 2000;

Reid, 2000). Although the total chlorine amount in the stratosphere is about 160 times larger than total bromine, bromine plays an important role in ozone depletion since it is much more effective than chlorine on a per-atom basis (World

Meteorological Organization, 2007, Chapter 4). All reactions discussed above

happen between gases only, therefore they are all classified as gas-phase chemistry.

1.4.3

Heterogeneous chemistry

Gas-phase chemistry alone contributes to ozone depletion only in the upper stratosphere (Staehelin et al., 2001). At lower stratospheric levels, and espe- cially at high latitudes, reactions of gases with liquid or solid particles play an important role, the so-called heterogeneous reactions. The main processes of those reactions are adsorption (accumulation of atoms or molecules on the sur- face) and absorption (atoms or molecules are taken up by the volume) of the gaseous molecules by the condensed particles. Chemical reactions then follow, forming different products. Heterogeneous reactions are extremely important for understanding the stratospheric ozone budget/balance, since they free chlorine and bromine from reservoir species (long-lived compounds which are relatively inert) into reactive forms.

There are two common species of particles on which heterogeneous chemistry can occur: stratospheric sulfate aerosols and PSCs. The stratospheric sulfate aerosols are typically composed of sulfuric acid (H2SO4) and water (McConnell

and Jin, 2008), at least at the mid-latitudes where temperatures are warm enough

to maintain the particles in a liquid state. The sulfuric acid originates from tropospheric carbonyl sulfide (COS) and sulfur dioxide (SO2) which was carried

along two main pathways into the stratosphere: either by lifting along the tropical branch of the Brewer-Dobson circulation or by the direct injection of SO2 from

large volcanic eruptions. Polar stratospheric clouds form under very cold conditions which are found only within the polar vortex regions during winter and early spring, from the tropopause up to ∼25 km (Dessler, 2000). PSCs are fundamental to an understanding of polar stratospheric ozone depletion: on their surface unreactive chlorine reservoir species (like hydrochloric acid, HCl, and chlorine nitrate, ClONO2) can be converted into reactive chlorine radicals. They are a condensation product of water vapor, sulfate aerosols and nitirc acid. The exact formation temperature depends on pressure and on the abundance of trace gases necessary to the formation. Two classes of PSCs can be differentiated: Type I and Type II. Type I PSCs form at temperatures around 195 K and are composed of nitric acid trihydrate (NAT) which consists of frozen nitric acid and water, as well as supercooled ternary solutions (STS) which combine water, sulfuric acid and nitric acid (McConnell and Jin, 2008). Type I PSCs are further divided into Type Ia and Type Ib. The PSCs of the former type consist of crystalline particles whereas the particles of the latter type are liquid. Type II PSCs form only at temperatures below 188 K and consist of frozen water particles. With typical radii of 5-20µm they are considerably larger than PSCs of Type I (0.5-1µm;Dessler, 2000).

Key heterogeneous reactions, either on sulfate aerosols or PSCs, as given by

Solomon (1999) and McConnell and Jin (2008), are:

ClONO2 (g) + HCl (s) → Cl2 (g) + HNO3 (s) (HC1)

ClONO2 (g) + H2O (s) → HOCl (g) + HNO3 (s) (HC2)

HOCl (g) + HCl (s) → Cl2 (g) + H2O (s) (HC3)

BrONO2 (g) + HCl (s) → BrCl (g) + HNO3 (s) (HC4)

BrONO2 (g) + H2O (s) → HOBr (g) + HNO3 (s) (HC5)

HOBr (g) + HCl (s) → BrCl (g) + H2O (s) (HC6)

N2O5 (g) + HCl (s) → ClNO2 (g) + HNO3 (s) (HC7)

N2O5 (g) + H2O (s) → 2 HNO3 (s) (HC8)

Letters in brackets denote if the reaction participant is gaseous (g) or solid (s). Reaction HC5 and HC8 are not very dependent on temperature (Dessler, 2000) unlike the other reactions. For the latter, reaction rates increase drastically for low temperatures, therefore they are most important for understanding the po- lar ozone budget. HC5 and HC8 are additionally important for heterogeneous reactions at mid-latitudes (on sulfate aerosols).

Molecules such as hydrogen chloride (HCl) and chlorine nitrate (ClONO2) are

reservoir species since they are relatively unreactive in their gaseous form. The conversion of those reservoirs on PSCs into Clx(for example by reaction HC1 and HC3) is called chlorine activation (Dessler, 2000). Since PSCs are composed of HNO3 and H2O these species are removed from the stratosphere when suffi-

ciently large PSCs undergo sedimentation. These processes are known as deni- trification and dehydration, respectively. The photochemically labile chlorine molecules (Cl2 as well as HOCl) are short-lived and can be rapidly photolyzed

when the sun comes back to high latitudes in spring. The Cl2 molecules are split

in two chlorine atoms which can participate in the odd-chlorine catalytic cycle which is described in Section 1.4.2 to destroy ozone. The photolysis of HOCl releases ClO which then can destroy ozone through the following reactions:

ClO + ClO + M → Cl2O2 + M

Cl2O2 + hν → OClO + Cl

OClO + M → Cl + O2 + M

2(Cl + O3 → ClO + O2).

The net reaction then is:

2 O3 + hν → 3 O2

A reaction set similar to the above mentioned ClO-ClO reaction exists for the reaction of BrO with ClO. Ozone loss due to the BrO-ClO reactions is about half of the loss caused by the ClO-ClO reactions, although BrO concentrations in the stratosphere are much smaller than ClO concentrations.

It is obvious from the reactions described above that Cl2O2 (also called theClO

dimer, Molina and Molina, 1987) plays a major role in high-latitude ozone de- pletion as soon as the sun comes back to the polar regions. However, with the returning sun the deactivation of chlorine also starts and, therefore, the depletion of ozone decreases: Temperatures within the polar vortex steadily increase again with the return of the sun which leads to suppressed formation of PSCs and, therefore, to less chlorine activation. HNO3 molecules still in the stratosphere

are destroyed by photolysis producing NO2 which will react with ClO to reform

the reservoir species ClONO2 (Dessler, 2000). However, with the denitrification

due to PSC sedimentation a possible pathway of chlorine deactivation is shut off and more ozone can be destroyed before Clx can be deactivated. In such deni- trified air, chlorine is, therefore, preferentially deactivated by reactions with CH4

In theory, these mechanisms are effective in the polar regions of both hemispheres. Nevertheless, ozone depletion over the Arctic is less severe than over the Antarctic. This is due to the less stable and warmer Arctic vortex (see Section 1.3.2) which reduces the frequency of temperatures low enough to form PSCs and, therefore, results in less chlorine activation (Dessler, 2000). Furthermore, in the Arctic, denitrification is reduced as a result of fewer PSCs and, therefore, deactivation of Clx is faster and the time period during which ozone can be destroyed is shorter. For a detailed description of the most common and important chemical reactions involved in stratospheric ozone depletion seeSolomon (1999),Dessler (2000) and

McConnell and Jin (2008).