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Ozone and the Stratosphere

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Because this section uses the chemical reactions in the earth’s atmosphere for examples, we start with some background information about our atmosphere. Atmospheric scientists view the atmosphere as divided into layers, each with its own characteristics. The lowest layer, which extends from the surface of the earth to about 10 km (about 6 miles), is called the troposphere. For our discussion, we are most interested in the stratosphere, which extends from about 10 km to about 50 km (about 31 miles). Above the stratosphere is the mesosphere (from about 50 km to 80 km) and above that is the thermosphere (from about 80 km to 150 km).

The stratosphere contains a mixture of gases, including oxygen molecules, O2, and ozone molecules, O3. These gases play a very important role in protecting the earth from high-energy ultraviolet radiation from the sun. The ultraviolet portion of the radiant energy spectrum can be divided into three parts: UV-A, UV-B, and UV-C. UV-A, which includes radiant energy of wavelengths from about 320-400 nm, passes through the stratosphere and reaches us on the surface of the earth. We are glad it does, because UV-A radiation provides energy for our production of vitamin D.

The shorter wavelength UV-B radiation (from about 290 nm to 320 nm) has higher energy than the UV-A radiation. Some of the UV-B radiation is removed by the gases in the stratosphere, but some of it reaches the surface of the earth. Radiation in this portion of the spectrum has high enough energy so that excessive exposure can cause sunburn, premature skin aging, and skin cancer.

The highest energy ultraviolet radiation is UV-C with wavelengths from about 40-290 nm. We are very fortunate that this radiant energy is almost completely removed by the gases in the atmosphere. This radiation has high enough energy to cause serious damage not only to ourselves but to all life on earth. DNA, which carries genetic information, absorbs UV radiation at about 260 nm. Proteins, which are important for both the structure and function of living systems, absorb radiation with wavelengths of about 280 nm. If these wavelengths were to reach the earth in significant quantity, the changes caused by altering DNA and protein molecules would lead to massive crop damage and general ecological disaster.

Removal of UV Radiation by Oxygen and Ozone Molecules

The oxygen molecules, O2, in the stratosphere play an important role in preventing high-energy UV radiation from reaching the earth’s surface. This radiant energy can supply the energy necessary to break the covalent bonds between oxygen atoms in oxygen molecules. Remember that the shorter the wavelength of light is the higher the energy. Radiant energy of wavelengths less than 242 nm has enough energy to break the O-O bond, but radiant energy with wavelengths longer than 242 nm does not supply enough energy to separate the atoms.

             < 242 nm
O2(g)                  2O(g)

Ozone, O3, is a pale blue gas with a strong odor. The concentration of ozone in the air you breathe is too low to see the color, but the odor you might have noticed around electric motors. Ozone is formed when an electric spark passes through oxygen gas. Because ozone is a very powerful oxidizing agent, it is considered a pollutant in the lower atmosphere, but in the stratosphere, it performs the important function of absorbing the potentially harmful UV radiation with wavelengths from 240-320 nm.

As in the case of O2, UV radiation can provide the energy to break the bond between oxygen atoms in ozone molecules. Because it does not take as much energy to break a bond in the O3 molecule as it does to break the bond in O2, the UV photons that break the bond in O3 are associated with longer wavelengths. The O3 molecules will absorb UV radiation of wavelengths from 240 nm to 320 nm.

            240-320 nm
O3(g)                      O(g) + O2(g)

We are glad to have the ozone layer in the stratosphere removing these wavelengths of radiant energy, because these wavelengths can cause problems, including skin aging, skin cancer, and crop failure. Because O2 does not remove this radiation, it is extremely important that the ozone layer be preserved.

Oxygen molecules, O2, and ozone molecules, O3, work together to absorb high-energy UV radiation. O2 molecules absorb UV radiation with wavelengths less than 242 nm, and O3 molecules absorb radiant energy with wavelengths from 240 nm to 320 nm.

There are several ways that ozone in the stratosphere is destroyed naturally other than in the reaction initiated by UV radiation described above. Perhaps the most important of these other reactions are described by the following equations.

NO(g) + O3(g)    NO2(g) + O2(g)

NO2(g) + O(g)    NO(g) + O2(g)

The first reaction destroys one ozone molecule directly. The second reaction destroys an oxygen atom that might have become an ozone molecule. (Because oxygen atoms can collide with oxygen molecules to form ozone molecules, the ozone concentration is depleted indirectly through the removal of oxygen atoms.) The main reason that this series of reactions is so efficient at destroying ozone molecules is that the NO(g) that is destroyed in the first reaction is regenerated in the second reaction. The net reaction involves the conversion of an ozone molecule and an oxygen atom into two oxygen molecules with no change in the number of NO molecules. This makes NO(g) a catalyst for this reaction. A catalyst is a substance that speeds a chemical reaction without being permanently altered itself. The equation for the net reaction is below.

                                                NO catalyst
net reaction: O3(g) + O(g)                     2O2(g)


Chlorofluorocarbons: A Chemical Success Story Gone Bad

In 1972 the chemical industry was producing about 700,000 metric tons (about 1.5 billion pounds) of chlorofluorocarbons, CFCs, per year. A chlorofluorocarbon is a compound composed of just carbon, chlorine, and fluorine. Most of the CFCs produced in the early 70’s were either CFC-11, which is CFCl3, or CFC-12, which is CF2Cl2. The development of these chemicals was considered a major success for the industry because they seemed to be perfect for certain uses. By the 1970’s, they were used as propellants in aerosol cans, as solvents, as expansion gases in the production of foams, as the heat-exchanging fluid in air conditioners, and as the working fluid in refrigerators.

One of the reasons that they were so useful is that they are very stable compounds, so there are very few substances that will react with them. This means very low toxicity and very low flammability. Another important characteristic is that they are gases at normal room temperatures and pressures, but they can be liquefied by putting them under pressures just slightly above normal pressures. These characteristics made them appear to be perfect chemicals for the uses listed above.

CFCs and the Ozone Layer

"The announcement today, suggesting a worse situation than we thought, affirms the warning I issued last Spring: Upper atmosphere ozone depletion remains one of the world’s most pressing environmental threats."

William K. Reilly,

Administrator of the U.S. Environmental Protection Agency    1991

Gases are removed from the lower atmosphere in two general ways. They either dissolve in the clouds and are rained out, or they react chemically to be converted into other substances. Neither of these mechanisms are important for CFCs. Chlorofluorocarbons are insoluble in water, and they are so stable that they can exist in the lower atmosphere for years. For instance, CFC-11 molecules have an average lifetime of 50 years in the atmosphere, and CFC-12 molecules have an average lifetime of about 102 years. During this time, the CFC molecules wander around in the atmosphere, moving wherever the air currents take them. They can eventually make their way up into the stratosphere where for the first time they come into contact with radiant energy of high enough energy to cause them to break down. For instance, radiant energy of wavelength less than 215 nm will break the covalent bond between one of the chlorine atoms and the carbon atom in CF2Cl2.

CF2Cl2(g)    CF2Cl(g) + Cl(g)

Click here to see molecular models of CFCs and other related molecules.

The chlorine atoms released in this sort of reaction can destroy ozone molecules. They react with ozone molecules and oxygen atoms in a similar way to the catalytic reactions between NO, O3, and O described in the last section.

Cl(g) + O3(g)    ClO(g) + O2(g)

ClO(g) + O(g)    Cl(g) + O2(g)

Each chlorine atom destroys one ozone molecule directly. The oxygen atom that reacts with ClO might have collided with an oxygen molecule to form an ozone molecule, so a second ozone molecule is prevented from forming. The chlorine atom is regenerated in the second reaction, so it is a catalyst for the reaction. The equation for the net reaction is below.

                                              Cl catalyst
net reaction: O3(g) + O(g)                 2O2(g)

Each chlorine atom is thought to destroy an average of 1000 ozone molecules before it is converted into an inactive form. Two important inactive chlorine compounds are HCl and ClONO2. They are formed in reactions like those below.

CH4(g) + Cl(g)    CH3(g) + HCl(g)

ClO(g) + NO2(g) + M    ClONO2(g) + M

M = molecule to carry off excess energy

In 1985 it was discovered that in the atmosphere over Antarctica there was a large decrease in the concentration of the ozone compared to what was expected. This "ozone hole" could not be explained with the model of the time, but it has since been explained in terms of the more rapid reformation of chlorine atoms from inactive chlorine compounds like HCl and ClONO2. The new model suggests that reactions like the ones described by the equations below take place on the surface of ice crystals which form in the cold air in the atmosphere above Antartica.

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

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

HOCl(g) + HCl(s)    Cl2(g) + H2O(s)

                 radiant energy
HOCl(g)                         Cl(g) + OH(g)

            radiant energy
Cl2(g)                        2Cl(g)

The chlorine atoms freed in these reactions can once again react with ozone molecules and oxygen atoms. It has been suggested that each chlorine atom generated in the stratosphere might be able to destroy tens of thousands of ozone molecules before it is permanently removed from the stratosphere.

The reactions described here are just a small fraction of the reactions in the stratosphere that directly or indirectly involve ozone. It has been estimated that modern calculations to predict ozone concentrations may involve 50 or more chemical species and 500 or more chemical reactions.

Green Chemistry - Substitutes for Chlorofluorocarbons

If you have purchased a television or a computer in recent years, it has probably come surrounded by polystyrene foam (Styrofoam) to protect it. This same foam is often used as an insulation material for coolers that you might take on a picnic. This foam was produced by blowing gas into the polystyrene as it solidifies to produce a very low density, stiff solid which is a good insulator. Over 700 million pounds of this foam were produced in 1995.

Chlorofluorocarbons, CFCs, have been used as the blowing agents in the past, but because of the damage that chlorofluorocarbons can do to the ozone layer, chemists are actively seeking alternatives for all of the jobs that CFCs have done in the past. The Dow Chemical Company received the 1996 Alternative Solvents/Reaction Conditions Award for their development of an alternative process for making polystyrene foam using 100% carbon dioxide, CO2, as the blowing agent. This award is part of the Presidential Green Chemistry Challenge Awards Program. Carbon dioxide is nonflammable, nontoxic, and does not deplete the ozone layer. The process does not even increase the level of CO2 in the atmosphere, because the carbon dioxide used comes from other commercial or natural sources, like ammonia plants or natural gas wells. This new technology eliminates the need for the use of 3.5 million pounds of CFCs per year.

Important Scientists Involved in the Study of the Ozone Layer

In 1930, an English physicist named Sidney Chapman suggested the following four reactions to explain the formation and destruction of the ozone in the atmosphere.

              UV Radiant Energy
O2(g)                                  2O(g)

O(g) + O2(g) + M    O3(g) + M           M = molecules to carry off excess energy

            UV Radiant Energy
O3(g)                                 O(g) + O2(g)

O(g) + O3(g)    2O2(g)

The model described by Sidney Chapman led to a prediction of much greater concentrations of ozone in the atmosphere than were found, so it was assumed that another mechanism must exist for its destruction. In 1970, Paul Crutzen, now at the Max-Planck-Institute for Chemistry, suggested the following reactions that involve nitrogen monoxide gas as a catalyst.

NO(g) + O3(g)    NO2(g) + O2(g)

NO2(g) + O(g)    NO(g) + O2(g)

                                             NO catalyst
net reaction O3(g) + O(g)                   2O2(g)

In 1971, Harold Johnston raised the concern about depletion of the ozone layer from human sources by suggesting that flying a fleet of supersonic transport planes in the stratosphere might produce enough NO(g) in the stratosphere to do significant damage to the ozone layer from the reactions suggested by Paul Crutzen. This fleet of supersonic aircraft was never built, but Johnston’s suggestion raised the awareness that it was possible for humans to have a serious effect on the ozone layer.

In 1974, Mario Molina, now at the Massachusetts Institute of Technology, and F. Sherwood Rowland, at the University of California, Irvine, showed that there was a threat to the ozone layer from chlorofluorocarbons, CFCs. They showed that the CFC molecules would almost all reach the stratosphere where they would liberate chlorine atoms that would then destroy ozone in a catalytic reaction similar to that involving NO(g).

                   UV Radiant Energy
CCl3F(g)                                CCl2F(g) + Cl(g)

Cl(g) + O3(g)    ClO(g) + O2(g)

ClO(g) + O(g)    Cl(g) + O2(g)

                                              Cl catalyst
net reaction O3(g) + O(g)                  2O2(g)

The Englishman Joseph Farman and his colleagues raised the issue of ozone depletion again in 1985 when they found a serious depletion of the ozone layer over the Antarctic. This is now called the "ozone hole". Molina, Crutzen, and others helped to identify the mechanism that explained this drastic decrease in ozone concentration in terms of reactions that take place on the surface of cloud particles in the atmosphere.

Professors Paul Crutzen, Mario Molina, and F. Sherwood Rowland were rewarded with the 1995 Nobel Prize for Chemistry for their outstanding contributions to the explanation of how ozone is formed and destroyed and for their contributions to the effort to save us from a potentially catastrophic environmental problem.

Other Ozone Depleting Chemicals

Although CFCs have gotten the most attention as a threat to the ozone layer, there are other chemical that have been found to be potentially damaging to ozone.

One category of these ozone depleting chemicals is bromine-containing compounds, including halons. Halons are compounds that are similar to CFCs, but they contain at least one bromine atom. Halon-1301 (CF3Br) and halon-1211 (CF2ClBr) have been used as fire extinguishing agents, but because they have been linked with the destruction of ozone in the stratosphere, alternatives are being actively sought. These halons have very long lifetimes in the atmosphere (over 60 years for halon-1301 and about 20 years for halon-1211), so they can exist long enough in the atmosphere to make their way to the stratosphere. Once they get there, in reactions that are very similar to the reactions involving CFCs, bromine atoms are liberated and react to destroy ozone. One of the possible reaction sequences is below.

                  Radiant Energy
CF3Br(g)                           CF3(g) + Br(g)

Br(g) + O3(g)    BrO(g) + O2(g)

Cl(g) + O3(g)    ClO(g) + O2(g)

                                Radiant Energy
ClO(g) + BrO(g)                               Br(g) + Cl(g) + O2(g)

Chlorine atoms spend a relatively large amount of time in an inactive form as HCl or ClONO2, but because HBr forms much more slowly than HCl and BrONO2 breaks down more rapidly than ClONO2, reactive bromine makes up a much larger portion of the bromine substances in the stratosphere. For this reason, each bromine atom has been estimated to be 50 times more efficient at destroying ozone than each chlorine atom.

There has been general agreement that halons should be banned, but another bromine containing compound has been more controversial. This compound is methyl bromide, CH3Br. The controversy revolves around whether or not it has a significant effect on the ozone layer and whether the benefits of using methyl bromide outweigh the potential hazards. This ozone depleting compound is different from CFCs and halons because it is not only produced by humans, but there are also large natural sources. For instance, the ocean is thought to both release and absorb significant amounts of CH3Br, and wildfires generate methyl bromide as well. The actual contribution of the various sources of methyl bromide in the atmosphere is still uncertain, but it is agreed that significant amounts come from several human sources, including biomass burning (like the burning of the rain forests), agriculture (where it is used is as an insecticide, herbicide, and fungicide), the fumigation of structures, and exhaust from cars using leaded gasoline that contains ethylene dibromide. These human sources are called anthropogenic sources. One of the reasons that methyl bromide has been considered less threatening to the ozone layer than CFCs or halons is that it has a much shorter lifetime. The best estimates predict its average lifetime to be from 1-2 years compared to over 50 years for the shortest lived common CFC and over 20 years for the most common halons. Although research on the possible effect of methyl bromide continues, it is now considered damaging enough to begin to phase it out.

With the discovery of the damaging effects of CFCs, alternatives were developed that as much as possible have the characteristics of CFCs but which are less stable in the lower atmosphere and less likely to reach the stratosphere. These chemicals called hydrochlorofluorocarbons, HCFCs, have a very similar structure to CFCs, but they contain at least one hydrogen atom. For instance, HCFC-22 is CF2HCl, and HCFC-123 is CF3CHCl2. Although these chemicals are thought to be less damaging to the ozone layer, they too can reach the stratosphere and lead to some depletion of O3. Thus, they are viewed as transitional compounds that will be used until better compounds are developed. They are being phased out over time.

The Copenhagen Amendment to the Montreal Protocol

In 1987, a worldwide agreement was reached to phase out the production and use of ozone depleting chemicals. This agreement called the Montreal Protocol was amended twice, in 1990 in London and again in 1992 in Copenhagen. The Copenhagen Amendment called for the following reductions of emissions of ozone depleting chemicals. (This is a partial list.) As more information is supplied by the scientific community, these guidelines are likely to change.

Ozone Depleting Chemicals

Industrialized Countries

Developing Nations


100% reduction by 1996 100% reduction by 2006


100% reduction by 1994 100% reduction by 2006


35% reduction by 2004 35% reduction by 2014
  65% reduction by 2010 65% reduction by 2020
  90% reduction by 2015 90% reduction by 2015
  99.5% reduction by 2020 99.5% reduction by 2030
  100% reduction by 2030 100% reduction by 2040


Fumigant uses Freeze in 1995 at 1991 levels Freeze in 1995 at 1991 levels
From auto exhaust Eliminate between 2000 and 2005 Eliminate between 2000 and 2005
From biomass burning 50% reduction between 2010 and 2040 50% reduction between 2010 and 2040

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