Ozone in the Stratosphere

In 1930, the English geophysicist Sydney Chapman (1888-1970) worked out the cycle of ozone formation and destruction in the stratosphere. There are four chemical reactions in the cycle  

Short-wavelength ultraviolet radiation (light with a wavelength less than 210 run) is the pump that drives the process. The single-oxygen atom species (O) that form are extremely reactive. They can collide either with an molecule and form ozone (as shown in the second chemical reaction in the cycle), or they can colUde with an O3 molecule and destroy ozone (as shown in the fourth chemical reaction in the cycle). Because the number of molecules is significantly greater than the number of O3 molecules, ozone formation is favored over ozone destruction. 

Although it is true that the absorption of ultraviolet radiation with a wavelength of about 220 to 320 nm destroys ozone molecules, it is precisely this ability of ozone that makes ozone in the stratosphere so important. Although some UV radiation still reaches Earth s surface (and causes sunburns, for example), most of the radiation is absorbed before reaching the surface, protecting living organisms on the surface from the harmful effects of UV radiation. 

By the time radiation from the Sun reaches an altitude of 90 km above Earth s surface, most of the short-wavelength radiation capable of photoionization has been absorbed. At this altitude, however, radiation capable of dissociating the O2 molecule is sufficiently intense for photodissociation of O2 (Equation 18.1) to remain important down to an altitude of 30 km. In the region between 30 and 90 km, however, the concentration of 

TABLE 18.3 Photoionization Reactions for Four Components of the Atmosphere 

O2 is much greater than the concentration of atomic oxygen. From this finding, we conclude that the oxygen atoms formed by photodissociation of O2 in this region frequently collide with O2 molecules and form ozone  

The asterisk on O3 denotes that the molecule contains an excess of energy. This reaction releases 105 kj/mol. This energy must be transferred away from the O3 molecule quickly or else the molecule will fly apart into O2 and atomic O —a decomposition that is the reverse of the reaction by which O3 is formed. 

An energy-rich O3 molecule can release its excess energy by colliding with another atom or molecule and transferring some of the excess energy to it. Let s use M to represent the atom or molecule with which O3 collides. (Usually M is N2 or O2 because these are the most abundant molecules in the atmosphere.) The formation of O3 and the transfer of excess energy to M are summarized by the equations 

Why do O2 and N2 molecules fall to filter out ultraviolet light with wavelengths between 240 and 310 nm  

The photodissociation of ozone reverses the reaction that forms it. We thus have a cycle of ozone formation and decomposition, summarized as follows  

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The importance of ozone in the stratosphere has been stressed in Section 9.3.8. The fact that ozone can be decomposed by the halogen monoxides CIO, BrO and 10 means that their presence in the stratosphere contributes to the depletion of the ozone layer. For example, iodine, in the form of methyl iodide, is released into the atmosphere by marine algae and is readily photolysed, by radiation from the sun, to produce iodine atoms which can react with ozone to produce 10 ... 

In the last decade, the refrigerant issue is extensively discussed due to the accepted hypothesis that the chlorine and bromine atoms from halocarbons released to the environment were using up ozone in the stratosphere, depleting it specially above the polar regions. Montreal Protocol and later agreements ban use of certain CFCs and halon compounds. It seems that all CFCs and most of the HCFCs will be out of produc tion by the time this text will be pubhshed. 

During the mid-1980s, each September scientists began to observe a decrease in ozone in the stratosphere over Antarctica. These observations are referred to as "ozone holes." In order to understand ozone holes, one needs to know how and why ozone is present in the earth s stratosphere. 

F. S. Rowland and M. Molina showed that man-made chlorofluorocarbons, CFCs, could catalytically destroy ozone in the stratosphere (Nobel Prize for Chemistry, with P. Crutzen, 1995). 

Despite their instability (or perhaps because of it) the oxides of chlorine have been much studied and some (such as CI2O and particularly CIO2) find extensive industrial use. They have also assumed considerable importance in studies of the upper atmosphere because of the vulnerability of ozone in the stratosphere to destruction by the photolysis products of chlorofluorocarbons (p. 848). The compounds to be discussed are ... 

About 51 percent of solar energy incident at the top of the atmosphere reaches Earth s surface. Energetic solar ultraviolet radiation affects the chemistry of the atmosphere, especially the stratosphere where, through a series of photochemical reactions, it is responsible for the creation of ozone (O,). Ozone in the stratosphere absorbs most of the short-wave solar ultraviolet (UV) radiation, and some long-wave infrared radiation. Water vapor and carbon dioxide in the troposphere also absorb infrared radiation. 

Ever) year our planet is bombarded with enough energy from the Sun to destroy all life. Only the ozone in the stratosphere protects us from that onslaught. The ozone, though, is threatened by modern life styles. Chemicals used as coolants and propellants, such as chlorofluorocarbons (CFCs), and the nitrogen oxides in jet exhausts, have been found to create holes in Earth s protective ozone layer. Because they act as catalysts, even small amounts of these chemicals can cause large changes in the vast reaches of the stratosphere. 

Figure 4. Processes that Control Ozone in the Stratosphere.

The concentration of ozone in the stratosphere is lower than predicted from reactions 1-4. This is due to the presence of trace amounts of some reactive species known as free radicals. These species have an odd number of electrons and they can speed up reaction 4 by means of catalytic chain reactions. Nitrogen oxides, NO and NO2, which are naturally present in the stratosphere at levels of a few parts per billion (ppb), are the most important catalysts in this respect. The reactions, first suggested by Paul Crutzen (2) and by Harold Johnston (3) in the early 1970 s, are as follows ... 

Events that take place on a grand scale often can be traced to the molecular level. An excellent example is the depletion of the ozone layer in the Earth s stratosphere. The so-called ozone hole was first observed above the Antarctic in the 1980s and is now being observed above both the Arctic and Antarctic poles. The destruction of ozone in the stratosphere is caused primarily by reactions between chlorine atoms and ozone molecules, as depicted in our molecular inset view. 

Catalysts are immensely beneficial in industry, but accidental catalysis in the atmosphere can be disastrous. Recall from Box that the chemishy of ozone in the stratosphere involves a delicate balance of reactions that maintain a stable concentration of ozone. Chlorofiuorocarbons (CFCs) shift that balance by acting as catalysts for the destruction of O3 molecules. 

We see that chiorine atoms provide an aitemative mechanism for the reaction of ozone with oxygen atoms. The iower-energy pathway breaks down ozone in the stratosphere at a significantiy faster rate than in the absence of the cataiyst. This disturbs the deiicate baiance among ozone, oxygen atoms, and oxygen molecules in a way that poses a serious threat to the iife-protecting ozone iayer. 

C15-0094. NO is an atmospheric poiiutant that destroys ozone in the stratosphere. Here is the accepted O3 + NO NO2 + O2 (slow)... 

C15-0117. Write chemical equations that show how CF2 CI2 contributes to the destruction of ozone in the stratosphere. 

Concern has been expressed over the destruction of ozone in the stratosphere brought about by its reactions with chlorine atoms produced from chlorofluoroalkanes that are persistent in the troposphere, and that may contribute to radiatively active gases other than COj. 

Chapman was the first to provide a clear picture of the formation of ozone in the stratosphere.9 Figure 1 summarizes the principal production... 

In order to calculate the steady-state concentration of ozone in the stratosphere, we need to balance the rate of production of odd oxygen with its rate of destruction. Chapman originally thought that the destruction was due to the reaction O + 03 —> 2O2, but we now know that this pathway is a minor sink compared to the catalytic destruction of 03 by the trace species OH, NO, and Cl. The former two of these are natural constituents of the atmosphere, formed primarily in the photodissociation of water or nitric oxide, respectively. The Cl atoms are produced as the result of manmade chlorofluorocarbons, which are photodissociated by sunlight in the stratosphere to produce free chlorine atoms. It was Rowland and Molina who proposed in 1974 that the reactions Cl + 03 —> CIO + O2 followed by CIO + O —> Cl + O2 could act to reduce the concentration of stratospheric ozone.10 The net result of ah of these catalytic reactions is 2O3 — 3O2. 

Even with our much-advanced understanding of the chemistry of the stratosphere, it appears that there are still some discrepancies between the calculated amount of ozone in the stratosphere and the amount measured. Toumi and Kerridge have summarized data showing that the range of calculated concentrations is some 10-15% below the range of measured... 

Because it is the UV-B radiation (280-320 nm) that causes the degradation, the absorption spectra of the UV-absorber must coincide with these wavelengths. UV-A (320-400 nm) does not cause damage (it is not energetic enough) and UV-C (wavelength less than 280 nm) does not reach the troposphere (it is filtered out by ozone in the stratosphere). The problem is to find an additive that absorbs UV-B but does not have an absorption "tail" in the UV-A and visible wavelengths, and therefore would have a yellow appearance. 

Ban on chlorofluorocarbons (CFCs) as aerosol propellants react with ozone in the stratosphere, causing an increase in the penetration of ultraviolet sunhght and increase the risk of skin cancer. 

You probably know that compounds called chlorofluorocarbons (CFCs) are responsible for depleting the ozone layer in Earth s stratosphere. Did you know, however, that CFCs do their destructive work by acting as homogeneous catalysts Use the Internet to find out how CFCs catalyze the decomposition of ozone in the stratosphere. To start your research, go to the web site above and click on Web Links. Communicate your findings as a two-page press release. 

Sticksel discussed vertical profile measurements of ozone in the stratosphere and the troposphere over the last several years. Transient ozone maximums in the troposphere are illustrated and explained by three possible mechanisms a channel-like r on conducted ozone from the stratosphere into the troposphere ozone-laden air descended from the stratosphere and was compressed as it subsided and ozone-rich layers leaked through the break between the polar and middle tropopauses by differential advection. Surface variations of ozone soundings were mostly attributed to anthropogenic pollution however, relatively thick high-... 

A number of experimental and theoretical studies have focused on the causes of mass-independent fractionation effects, but as summarized by Thiemens (1999), the mechanism for mass-independent fractionations remains uncertain. The best studied reaction is the formation of ozone in the stratosphere. Mauersberger et al. (1999) demonstrated experimentally that it is not the symmetry of a molecule that determines the magnitude of enrichment, but it is the difference in the geometry of the molecule. Gao and Marcus (2001) presented an advanced model, which has led to a better understanding of nonmass-dependent isotope effects. 

Mauersberger K (1981) Measurement of heavy ozone in the stratosphere. Geophys Res Lett 8 ... 

Long-range effects of having less ozone in the stratosphere involve greater ultraviolet sunlight transmission, alteration of weather, and an increased risk of skin cancer. The ozone depletion potential for CFCs and other fluorocarbons have been measured and are given below relative to CFC-11 and -12. Notice that the HCFCs with lower chlorine content have lower depletion potentials than the CFCs, and the one HFC studied shows no depletion potential because it contains no chlorine. 

The role of CFCs in the destruction of ozone in the stratosphere was something of a surprise to some researchers because those compounds are normally quite stable. In fact, their stability is one of their most desirable properties for many industrial and commercial applications. But, when CFCs escape into the atmosphere and drift upward, they are exposed to ultraviolet radiation in sunlight and, as is oxygen itself, are dissociated by that radiation. In the case of Freon-12 (CCI2F2), photodissociation results in the formation of free chlorine atoms ... 

CFCs are the most important, hut by no means the only, chemicals capable of destroying ozone molecules. For many years, researchers have recognized that oxides of nitrogen have the capacity both to increase and to decrease the concentration of ozone in the stratosphere. They can increase ozone concentrations in the presence of ultraviolet (uv) radiation by undergoing uv-mediated reactions similar to those that occur in the lower troposphere. For example ... 

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