Chapter 8. Ozone

Table of Contents

8.1. Stratospheric ozone
8.1.1. A short history
8.1.2. Natural balance of stratospheric ozone
8.1.3. Circulation patterns in the stratosphere
8.1.4. Gas-phase chemistry in the stratosphere
8.1.5. Polar ozone chemistry
8.2 Tropospheric ozone
8.2.1. Global tropospheric ozone budget
8.2.2. Ozone production in the troposphere
8.2.3. Sinks of the tropospheric ozone
8.2.4. Spatial and temporal variability of ozone in the near surface layer

Ozone (O3) is a tri-atomic molecule consisting of three oxygen atoms was discovered and first published by a German-Swiss chemist, Christian Friedrich Schönbein in 1840. It is a toxic, oxidative, unstable gas with pungent odour (Schönbein denominated ozone after its characteristic odour from the Greek word „ozein”, which is meaning, “to smell”. Ozone has a greater oxidation potential than oxygen. About 90% of atmospheric ozone can be found in the stratosphere, particularly between 19 and 30 km above the Earth’s surface, and only about 10% of total ozone is located in the troposphere, especially near ground level.

Ozone plays important role in radiation balance of the atmosphere and in the formation of photochemical air pollution, too. Stratospheric ozone, sometimes called as ozone layer, filters out the incoming ultraviolet radiation in the bands between 210 and 290 nm to protect the live. In the troposphere, it works as a greenhouse gas. Near ground level, ozone can affect harmfully both vegetation and human life (see Chapter 13).

8.1. Stratospheric ozone

8.1.1. A short history

After the discovery of ozone by Schönbein, Walter Hartley recognized in 1881 that a sharp cut-off in radiation spectrum between 210 and 290 nm (called Hartley band) correspond to the absorption of ultraviolet radiation by ozone. In 1993, Albert Wigand obtained solar spectra from a balloon at 9000 m, and he found essentially the same cut-off as at the sea level. John William Strutt (better known as Lord Rayleigh) showed in 1918 that ozone in the lowest layer of the atmosphere is not enough for uv absorption, and assumed that most of the atmospheric ozone can be found above 10 km over the surface. In 1921, Charles Fabry and Henry Buisson published the results of their ozone column measurement. They found that the mean thickness of atmospheric ozone is about 3 mm, but this value varies from day to day. First ozone spectrophotometer was developed by Gordon Miller Bourne Dobson in 1924, and regular measurements of ozone abundance started in 1925. This instrument measures the total ozone in a column of air overhead. The amounts of ozone are regularly reported in thickness of the total column ozone, commonly known as Dobson Units (DU). This measure is defined as the amount of ozone contained in a vertical column of base 1 cm2 at standard air pressure and temperature. The world average is about 300 DU, and it varies geographically from about 230 to 500 DU.

Just a few years after first measurements with Dobson spectrophotometer, Paul Götz discovered a method, called Umkehr-effect, for determining the vertical profile of ozone concentration. This effect is observed when the ratio of the measured intensities of scattered zenith skylight at two different wavelengths (311 and 332 nm) is plotted against the increasing zenith angle of the Sun between 60° and 90°. The ratio increases with increasing zenith angle up to about 86°, then a reversal (Umkehr) occurs. The spectrally different changes in zenith sky uv radiation around Sun rises and sets are controlled by the ozone vertical distribution.

In 1930, Sydney Chapman proposed a theory to describe the natural balance (production and destruction) of stratospheric ozone.

8.1.2. Natural balance of stratospheric ozone

Ozone (O3) can be found in the earth’s atmosphere from the ground up to about 100 km, but most of the ozone (about 90%) resides in the stratosphere (between about 10 and 50 km above the Earth's surface). This stratospheric ozone is commonly known as the „ozone layer”. The ozone concentration maximum varies with latitudes, but can be localized at an altitude around 20 km. The maximum concentration varies here between 4 and 8 ppm. Typical vertical profile of ozone can be seen in Figure 8.1.

Figure 8.1: Ozone in the atmosphere

Production of ozone

Figure 8.2: Production of ozone from an oxygen atom and an oxygen molecule in the presence of a third molecule (usually N2, or another O2).

Ozone is constantly produced and destroyed in a natural cycle. Ozone is formed from a single recombination reaction of an atomic oxygen and a molecular oxygen (Figure 8.2.) in the presence of a third body M, which is required to carry away the energy released in the reaction (M is usually O2 or N2):

O + O2 + M → O3 + M


Here, O atoms are the fundamental triplet state O(3P), which are highly reactive due to their two unpaired electrons. They combine rapidly with O2 to form ozone.

The net production of oxygen atoms results almost exclusively from the photodissociation of molecular oxygen (Figure 8.3). The bond energy of the oxygen molecule (498 kJ mol–1) corresponds to the energy of a 242 nm uv photon, therefore only photons of wavelengths less than 242 nm can photolyze these molecules:

O2 + hν (λ < 242 nm) → O+ O


Photodissociation of oxygen molecule

Figure 8.3: Photodissociation of oxygen molecule into to oxygen atoms

Ozone molecule produced in reaction (R8.1) absorbs solar radiation and decomposes back to O2 and O (Figure 8.4.). Because the bonds in the O3 molecule are weaker (364 kJ mol–1) than those in the O2 molecule, photolysis is achieved with lower-energy photons (in the wavelength range of 240 to 320 nm):

O3 + hν (λ < 320 nm) → O2 + O(1D)


O(1D) + M → O+ M



O3 + hν (λ < 320 nm) → O2 + O


where O(1D) is the O atom in an excited singlet state which is rapidly stabilized to O(3P) by collision with N2 or O2.

Destruction of ozone molecule by solar ultraviolet radiation

Figure 8.4: Destruction of ozone (O3) molecule by solar ultraviolet radiation (λ < 320 nm). The result is one oxygen molecule and one excited oxygen atom.

Reactions (R8.1) and (R8.3) are very fast. This dissociation and recombination cycle taking place continuously and occur in less than 100 s throughout the stratosphere. During this cycle, solar uv radiation breaks ozone molecule and oxygen atom reacts an oxygen molecule. This process converts uv radiation to thermal energy, heating the stratosphere.

Additionally, ozone can react with atomic oxygen to regenerate two oxygen molecules:

O + O3 → O2 + O2


However, stratospheric ozone concentration calculated by Chapman scheme is much greater than the observed one. The reaction (R8.6) could only account for about 20% of odd oxygen (O and O3) loss in the stratosphere. Since the 1970s, it is well known that ozone is mainly removed in natural stratosphere through catalytic cycles. During these catalytic cycles, several ozone molecules destructed, while the catalyst that started the reaction is reformed to continue the process (Figure 8.5). Several catalytic cycles are important in ozone chemistry, involving homogeneous gas-phase reactions of active radical species from hydrogen (HOx), nitrogen (NOx) and chlorine (ClOx) families. In the lower stratosphere, cycles catalyzed by bromine (BrO) also contribute to the ozone loss. A large variability of possible catalytic cycles can be described generally by the following reactions:

X + O3 → XO+ O2


XO + O → X+ O2



O + O3 → O2 + O2


where X = H, OH, NO, Cl, Br, or F as catalysts.

A catalytic reaction chain in the stratosphere

Figure 8.5: A catalytic reaction chain in the stratosphere. One catalyst can destroy several ozone molecules.

8.1.3. Circulation patterns in the stratosphere

The spatial distribution of stratospheric ozone is controlled also by dynamical processes. Most of the ozone is produced in the tropical stratosphere, where solar radiation is strongest and reaction (R8.2) is the most effective. At the same time, higher stratospheric ozone concentration can be observed outside this tropical source region. This picture is a result of a slow atmospheric circulation, called Brewer–Dobson circulation (see e.g. in Gerber, 2012). Brewer–Dobson circulation (BDC) characterizes the transport of mass through the stratosphere. First, in the tropics, the air is rising from the troposphere, then is drifting poleward in the mid latitudes, and finally is descending in the mid and high latitudes. Descending air can be transported back into the troposphere in the mid latitudes, while it can be accumulated in the lower stratosphere at higher latitudes. Due to this circulation pattern, the total ozone column varies in a wide range, from about 220 Dobson Units over tropical region, to above 450 Dobson Units in polar region. This general distribution however has a daily, seasonal and interannual variability, especially during ozone hole events (see later).

8.1.4. Gas-phase chemistry in the stratosphere

The behaviour of the stratospheric ozone layer depends on the chemical reactions, which govern the formation, chemical transformation and removal processes of chemically active components. Several chemical species slowly diffusing from the surface through the troposphere, and they can reach the stratosphere.

Nitrogen oxides (NO and NO2) dominate the balance between ozone production and destruction in the natural stratosphere. The main source of nitrogen species is nitrous oxide (N2O) emitted from soil:

O(1D) + N2O → NO


Additionally, significant amount of NO are emitted directly to the lower stratosphere by the exhaust gases of aircrafts. The main ozone destruction cycle in the middle stratosphere (25–35 km) is the following:

NO + O3 → NO2 + O2


NO2 + O → NO + O2


In the upper stratosphere, odd hydrogen species play important role in the ozone destruction. The main sources of the hydroxyl-radical (OH) and subsequently of hydrogen (H) and hydroperoxyl (HO2) radicals are the reactions of excited atomic oxygen with water vapour and methane:

O(1D) + H2O → 2OH


O(1D) + CH4 → OH + CH3


and the ozone destruction mechanism if hydrogen species:

OH + O3 → HO2 + O2


HO2 + O→ OH + O2


or in the lower stratosphere, where atomic oxygen is much less abundant, due to the rapid formation of ozone (R8.4):

HO2 + O3 → OH + 2O2


Additionally, in the lower stratosphere the coupling between the odd nitrogen and odd hydrogen chemistry can lead to a net production of ozone.

Chlorine species have also a key role in the stratospheric chemistry. However, the major chlorine compounds at the surface (Cl2, HCl, NaCl) emitted from oceans and continuous volcanic activity are water soluble and they have a very short lifetime. They removed from the lower atmosphere to the surface by wet deposition (see. Chaper 12) during a few days and they cannot transferred to the stratosphere. The only known natural source of stratospheric chorine is methylchloride (CH3Cl) produced by micro-organisms in the ocean. However, methylchloride is only produced about 20% of stratospheric chlorine.

In the 1970’s, it became clear that the most important cause for the depletion of stratospheric ozone is the growth in the anthropogenic chlorofluorocarbon gases. The chemically inert chlorofluorocarbons may remain in the atmosphere for 40 – 150 years, and their photodissociation in stratosphere produce significant amounts of chlorine radicals (Molina and Rowland, 1974). The most significant chlorofluorocarbons are CFCl3 (CFC11), CF2Cl2 (CFC12), carbon-tetrachloride (CCL4), methylchloroform (CH3CCl3) while chlorine radicals (ClOx) are chlorine (Cl) and chlorine monoxide (ClO). Chlorine formation can be described generally as:

XCl + hν (λ < 215 nm) → X + Cl


The catalytic chain reaction leading to the net destruction of O3 and O related to chlorine compounds:

Cl + O3 → ClO + O2


ClO + O→ Cl + O2


This ozone depletion by chlorine radicals occurs especially about 30 km above the surface. In the lower part of stratosphere, the effects of chlorine radicals on ozone depletion are moderated due to the strong chemical interactions between chlorine radicals and other compounds. The presence of nitric-oxide (NO) can hinder the ozone depletion, as it reacts with ClO:

ClO + NO → Cl + NO2


Resulting nitrogen dioxide (NO2) takes part in the formation of ozone, which moderates the ozone depletion mechanisms. Additionally, further chemical reactions between ClOx, and NOx–HOx can lead to convert the catalysts into less reactive reservoir species which do not react with ozone.

ClO + NO2 + M → ClONO2 + M


ClO + HO2 → HOCl + O2


ClO + OH → HCl + CH3


CH4 + Cl → HCl + CH3


These reservoir species, such as chlorine nitrate (ClONO2), hypochloic acid (HOCl) or chlorhydric acid (HCl) can slowly transported downward to the troposphere, where they can dissolved and removed into the surface by wet deposition. However, HCl is only a temporary reservoir in the stratosphere, as chlorine can be formed again through the reaction with hydroxyl radical:

OH + HCl → Cl + H2O


Brominated hydrocarbons (halons) from anthropogenic sources can also reach the stratosphere, where bromine catalyst can be formed during photodissociation processes. These bromine catalysts (active bromine – Br, or hypobromite – BrO) can also take part in ozone depletion, similarly, as chlorine compound.

8.1.5. Polar ozone chemistry

Significant seasonal ozone depletion was observed in the lower stratosphere (between 15 and 22 km) over Antarctica during the Southern Hemisphere spring (Farman et al., 1985). Each year for the past few decades during the period between September and November, the stratospheric ozone layer in the southern polar region was destroyed rapidly causing a large area, where total column ozone decreased below 220 Dobson Unit. This region is known as the „ozone hole”.

The maximum of the daily ozone hole size for each year from 1979 to 2012

Figure 8.6: The maximum of the daily ozone hole size for each year from 1979 to 2012. (Source of the data:

The minimum of the daily total column ozone minimum for each year from 1979 to 2012

Figure 8.7: The minimum of the daily total column ozone minimum for each year from 1979 to 2012. (Source of the data:

Figure 8.6 shows the yearly maximum extension of ozone hole from 1979, while Figure 8.7 presents the minimum total ozone column for each year from 1979. The border of the ozone hole is represented by a line with a constant value of 220 Dobson Units. This value was chosen since total ozone values of less than 220 DU were not found over Antarctica before 1979 (measurements started in the 1950’s). This reduced ozone level is a result of the ozone loss by chlorine and bromine compounds. However, the atomic oxygen, O, has very low concentrations at this region, which limits the efficiency of the formerly described catalytic reaction chains. In that case, other type catalytic reactions become the dominant.

In the polar stratosphere, heterogeneous chemical reactions are extremely important in ozone depletion. Besides of these reactions, dynamical and radiative processes are also responsible for the emergence of the ozone hole.

Over the polar region (above about 65 °C), an isolated, quasi-cylindrical, dynamically stable persistent area, called polar vortex is formed in the mid and upper troposphere and in the lower stratosphere. This cold vortex is strongest during winter and weaker or can disappear in summer. Due to the different geographical conditions, the polar vortex is less perturbed, therefore more pronounced and persistent in the Southern Hemisphere than over the Arctic region. During the polar winter, at very low temperature, different types polar stratospheric clouds (PSCs) can be formed in the polar vortex between altitude of 15 and 25 km. Type I PSCs are formed at temperature below −78 °C and composed by supercooled solution of nitric acid, sulphuric acid and water ice (HNO3-H2SO4-H2O) or condensated nitric acid trihydrate (NAT – HNO3-3 H2O). Type II PCs are less common, and contains water ice particles that form when temperature drops below −88 °C. PSCs are mainly formed over Antarctica, in the polar vortex, but they are less frequent in the Artctic stratosphere, because the temperature in this region is on average about 10 °C warmer than in the Antarctic lower stratosphere. Therefore, ozone depletion is less significant over the Northern polar region.

Polar stratospheric clouds plays important role in ozone depletion (Crutzen and Arnold, 1986). Heterogeneous chemical reactions, which convert inactive chlorine compound to ozone destroying radicals can take place on the component of PSCc. Additionally, the sedimentation of nitric acid (HNO3) particles from polar stratospheric clouds leads indirectly to the removal active nitrogen species (this process is called denoxification). As NO2 can react with ClO to form reservoir chlorine nitrate (ClONO2) species, the removal of NO2 helps maintain large levels of ozone-destroying ClO.

The heterogeneous reactions in the polar stratosphere (Molina et al., 1987; Isaksen, 1994) involving HCl, ClONO2, HOCl, N2O5 and H2O are responsible for the conversation of reservoir species into reactive forms of chlorine, and also for removal of reactive nitrogen species into more stable forms (such as HNO3):

ClONO2 + HCl → Cl2 + HNO3


ClONO2 + H2O → HOCl + HNO3


N2O5 + HCl → ClNO2 + HNO3


N2O5 + H2O → HNO3 + HNO3


HOCl + HCl → Cl2 + H2O


Due to the heterogeneous chemistry, active species of chlorine (and also bromine) accumulated in the polar vortex. Then, after polar sunrise, the re-appearing solar radiation starts the photochemical activity, which leads to intensive ozone depletion in polar stratosphere during the polar spring. Chlorine reservoirs (such as molecular chlorine - Cl2) photodissociated rapidly by ultraviolet and visible radiation into chlorine atoms:

Cl2 + hν (visible light) → Cl + Cl


which initiates catalytic reaction chains, as chlorine atoms react with ozone to produce ClO. Due to the low solar elevation angles, atomic oxygen is not present in the polar stratosphere, therefore the ozone destruction occurs through the following catalytic chain:

Cl + O3 → ClO + O2 (2x)


ClO + CLO +M → Cl2O2 + M


Cl2O2 + hν (visible) → 2Cl + O2



2O3 → 3O2


Similar catalytic bromine cycles can occur involving Bromine compound:

ClO + BrO → Cl + Br + O2


Cl + O3 → ClO + O2


Br + O3 → BrO + O2



2O3 → 3O2