2. Good Ozone (Stratospheric) vs. Bad Ozone (Tropospheric)

Most (about 90 %) of the ozone in the atmosphere is in the stratosphere, where it plays a beneficial role in regulating solar UV radiation. The rest is in the lower atmosphere. Some of this ozone in the troposphere is leakage downward from the stratosphere. Some of it is generated naturally by lightning. In many urban areas, however, there is a significant problem with too much ozone that is produced as a by-product of air pollution. At concentrations found in many U.S. cities during the summer, ozone is known to reduce the lung capacity of healthy adults; its effects on athsma sufferers can be severe. It is also known to damage crops at high concentrations. Note, however, that the highest ozone concentrations seen during air pollution episodes in urban areas are still far lower than the concentrations of ozone found in the stratosphere. Hence, shipping our bad ozone from down here up to the stratosphere is not a solution. Further background on ozone in the lower and upper layers of the atmosphere is given in the following web links:

3. Natural Production and Destruction of Stratospheric Ozone

Thanks to the abundance of oxygen molecules in our atmosphere, much of the solar radiation with wavelengths less than 240 nm is absorbed before it reaches the surface of Earth. In the process, the oxygen molecule is blasted apart into two oxygen atoms. Individual oxygen atoms are very chemically unstable and will react very quickly with other oxygen molecules that they run into.

The molecule O3 is called ozone. In this reaction, M is another molecule that collides with the newly formed ozone molecule to absorb some of the energy left over from the collision with the photon that broke up the original oxygen atom. This transfer of energy from incoming solar radiation to gas molecules causes the temperature of the air to rise.

Ozone also absorbs incoming solar UV radiation. It captures light with wavelengths shorter than 310 nm. In the process, an oxygen atom is knocked off of the ozone, leaving an oxygen molecule and a free oxygen atom, as indicated in chemical equation 2.

The result of reaction 2 provides the essential ingredients for reaction 1 to occur once again. In fact, in the natural atmosphere, reactions 1 and 2 proceed continuously providing an ?equilibrium concentration of ozone as a balance between the production reaction (1) and the destruction reaction (2). These chemical reactions are discussed in the Text, Edition 3 in regard to Figure 13.39.

Because of the combined effects of ozone and oxygen, most of the solar radiation with wavelengths shorter than 310 nm is absorbed before reaching the surface of the Earth. As mentioned above, this absorption of light is accompanied by a local increase in air temperatures. Most of this absorption of solar UV radiation takes place from 30 to 50 km above the surface of the Earth, and the effects can be clearly seen in measurements of increasing temperature as a function of altitude upward from the troposphere to the altitudes of maximum ozone. This layer of the atmosphere where temperature increases with altitude is called the stratosphere. It is home to the so called "ozone layer" where ozone concentrations are highest. It is also the location of the jet stream. It forms an effective lid on the lowest layer of the atmosphere, the troposphere. In the troposphere, there are weather disturbances that churn up the atmosphere. The stratosphere, in contrast, does not have storms and there is very little vertical mixing of air there.

The two key chemical reactions, and the natural equililbrium in ozone concentration, are illustrated in the web links below showing pictorial representations of the reactions with additional explanation.

4. Experimental Discovery of an Ozone Hole

Perhaps the clearest suspicion of degradations in Antarctic ozone concentration came from space observations in the early to middle 1980s. Spacecraft instrumentation can infer ozone concentration in the atmosphere using measurements of backscattered ultra violet radiation or absorption of longer wavelength infrared radiation as it moves upward from the warm earth surface and is partially absorbed by atmospheric ozone. These two ventures were illustrated in the web links of section 1. Both types of radiation measurement from spacecraft have helped identify changes in ozone concentration over time. The following three web links provide a sample of the experimental evidence of ozone depletion over the Antarctic since about 1980.

5. Theory of Human-Induced Destruction of Stratospheric Ozone - The Chemical Effects of Chloroflourocarbons (CFCs)

Chlorofluorocarbons (or CFC’s for short) are a family of chemical compounds that are not known to occur naturally. They were invented in the 1920’s and found wide use as refrigerants and propellants. One of the most widely used CFC’s is called Freon-12. It is similar in structure to a methane molecule (CH4), except that two of the hydrogen atoms have been replaced by chlorine atoms and two more by fluorine atoms. The chemical bonds between the central carbon atom and the chlorine and fluorine atoms is much stronger than a normal carbon to hydrogen bond. As a result of these strong chemical bonds, Freon-12 and the other CFC’s are almost chemically inert. This means that they last a very long time, don’t corrode equipment and are non-toxic.

When CFC’s were invented, no one thought about the possible fate of these chemicals once they were released into the atmosphere. It wasn’t until the 1970’s that atmospheric chemists deduced that if there was no mechanism for destroying CFC’s in the troposphere, they would eventually diffuse into the stratosphere. (Evolution of this line of chemistry research led to the award of the Nobel Prize in 1995 to Crutzen, Molina, and Rowland). Laboratory measurements showed that if CFC’s were exposed to UV radiation with wavelengths less than 230 nm, the interaction would result in a chlorine atom being knocked off of the CFC molecule, with the possibility of additional ozone-affecting reactions to follow. Because of the natural UV shielding of oxygen and ozone, UV radiation less than 230 nm is only found in the upper stratosphere and above. Hence, the eventual fate of the CFC’s would be their breakup by UV radiation in the stratosphere. It was thought at that time that it would take 50 to 100 years for the CFC’s to slowly make their way up to this stratospheric region high in the atmosphere. However, we have subsequently learned that the relatively heavy CFC molecules can be effectively carried upward into the stratosphere by naturally occuring weather systems (convective thunderstorms, etc.) operating as an effective pipeline between earth's surface, through the troposphere, and into the stratosphere. The slow 50-100 year scenario quickly gave away to a more eminent threat of stratospheric ozone depletion due to human-manufactured agents, such as CFCs, carried upward from ground-based sources through naturally occuring weather processes.

6. The Catalytic Destruction of Stratospheric Ozone

Theory and experiment described in the previous section leads to an expectation that realtively inert chloroflourocarbons created by humans at ground level will work their way into the stratosphere. At such high altitudes, the CFCs are no nearly so chemically stable owing to the action of ultraviolet radiation from the sun which can break the chemical bonds and free up a chlorine atom. Unlike the CFC molecules themselves, the chlorine atoms released in this process are very chemically reactive. Because of the relatively high concentrations of ozone at the stratospheric altitudes where the CFCs are being broken down, it seemed very likely that the chlorine atoms would react with ozone. An hypothesis about this can be developed from chemical experiments carried out in traditional, earth-bound laboratories. When CFC molecules are bombarded with ultraviolet light (radiation), a chlorine atom splits off which then captures one of the oxygen atoms from an ozone molecule leaving chlorine monoxide (ClO) plus molecular oxygen (O2):

Because there are also relatively high concentrations of oxygen atoms at these altitudes, the ClO can react with an oxygen atom to return the chlorine to its original atomic state, leaving molecular oxygen as a final remnant:

In the process of reactions 3a and 3b, the chlorine atom is left unaffected, but an ozone molecule has been converted to an oxygen molecule in reaction 3a. Then, in reaction 3b an oxygen atom that might otherwise have gone on to produce an ozone molecule, as in Reaction 1, has instead been converted into an oxygen molecule. The chlorine atom is then able to repeat this cycle many times. The net result of the chemical reactions 3a+3b is that ozone is converted to molecular oxygen through reactions stimulated by the presence of a chlorine atom, and those reactions leave a chlorine atom free to continue the process. The chlorine atom is present at both the beginning and end of these reactions so that a single chlorine atom can convert very many ozone molecules to molecular oxygen before it is finally captured by some other exotic reation which binds it into a more stable molecule. The chlorine atom here is called a catalyst, and can cycle through this reaction breaking down a lot of O3 into O2 without any substantial change to itself. Like Yentl, the "matchmaker" catalyst of the stage musical "Fiddler on the Roof", the chlorine atom wanders through the population of stratospheric ozone selecting "eligible bachelor" oxygen atoms to marry-up with "eligible maiden" oxygen atoms producing a lot of oxygen "couples" without ever getting attached herself. Similar ozone-destroying catalytic reactions involve flourine and bromine in replacemnt of chlorine. Another catalytic reaction cycle involving nitric oxide is Reaction 4, cited in the Text, Edition 3, Figure 13.39; Edition 2, Figure 11.33:

Chlorine, flourine, bromine, and nitric oxide are referred to as catalysts for ozone destruction in these situations because they accelerate the removal of ozone from the atmosphere without themselves being consumed in the process. The chlorine-based catalytic reaction is shown pictorially in the following web links. In the second of these, the reaction is animated to dramatize the role of the chlorine atom catalyst at it moves through the population of ozone molecules leaving molecular oxygen gas in its wake. To help you appreciate the characteristics of catalytic chemical cycles in the stratospheric ozone problem, the third link below illustrates the analogous action of the matchmaker "Yentl" in "Fiddler on the Roof".

7. Experiments to Test the Hypothesis - Ozone Depletion at the South (and North?) Poles

The chemistry of ozone destruction through catalytic reactions based on human manufactured chemicals was first established by carefully controlled laboratory experiments. This research, honored by the 1995 Nobel Prize in chemistry to Crutzen, Molina, and Rowland, set the stage for the next step in the scientific method, that of testing the hypothesis against observations in the atmosphere. The theory of ozone depletion by CFC'c caused alarm in the 1970's, but it wasn't until the early 1980's that scientists discovered significant ozone depletion over the continent of Antarctica. Initially, there were many who believed that this ozone depletion might be unrelated to CFC's. Scientists needed a good test of the CFC hypothesis. Following the deductive part of the scientific method, they assumed that the ozone depletion was caused by CFC's as described above. From this assumption, they then deduced that areas of the stratosphere with depleted ozone should also have higher than normal concentrations of ClO, a product of reaction 3a. This is exactly what was measured by high-altitude aircraft that flew into the ozone hole. Additional observational support for this hypothesis came through the use of special research experiments designed for the platform of earth-orbiting satellites. A variety of satellite-based experiments have since been launched which are summarized in the next web link. Also see Text, Edition 2, Figures 13.41, 13.42, 13.43, 13.44.

For example, appropriate to the basic hypothesis that chlorine-based catalytic chemical reactions might destroy ozone in the stratosphere, scientists tested this hypothesis against satellite-based experimental observations of the concentrations of both chlorine monoxide and ozone in the stratosphere, with rather dramatic results. Sample observations are shown in the next web link.

A single cycle of the scientific method does not establish "absolute truth" about the ozone hole. A single experiment, such as the UARS experiment, does not prove that CFC's were responsible for the ozone depletion. But it is very difficult to explain the high concentrations of ClO and lowered conmcentrations of ozone without invoking the CFC-related ozone depletion scheme described above. The growing collection of laboratory experiments, chemical theories, field studies of meteorological processes, computer simulation models, and space-based observations seem to be leading to a very durable hypothesis. Human-manufactured chemicals, such as the CFCs, are working their way from earthbound sources upward into the stratosphere. There they participate in photochemical reactions that reduce the concentration of ozone far below that which would be occuring naturally without the human influence.

While the early phases of the scientific method of study of the Antarctic ozone hole demonstrates that CFC's are likely to have an impact on ozone, it created more questions than it answered, which is the outcome good scientific methodology should produce. For example, why did the ozone hole open up during the polar springtime? And why was the ozone depletion concentrated over Antarctica? Part of the springtime question relates simply to the fact that the catalytic reactions described above require sunlight (the source of the ultraviolet radiation), and therefore proceed only after sunlight returns to the polar region after the dark winter period. Subsequent research has also shown that certain types of ice crystals play a key role in accelerating ozone depletion. These ice crystals only form when it is very cold (-80 Celsius), as occurs over Antarctica during the winter months. At this time, however, there is no sunlight, and hence there is nothing to break the chlorine atoms off of the CFC molecules. When the sun rises in the polar spring, chlorine is released and the ice particles speed up the process of ozone depletion for a month or more until warming temperatures melt the ice particles.

The apparent prominence of an ozone hole at the South Pole as compared with the North Pole is due to differences in meteorology. It doesn't get as cold at the North Pole as it does over Antarctica, and thus it is less common to have the necessary clouds of ice particles. The link below shows maps of wintertime temperature over the North and South Poles, illustrating the much colder temperatures over the South Polar region.

Additionally, atmospheric circulations around the South Pole act to trap the chemical morass for a longer period of chemical "cooking" than the more wavy, meandering flow around the North Pole. The different wind flow regimes over the opposite poles during their respective winters are shown in the next link.

In recent years, however, there has been increasing evidence that ozone is also being depleted over the North Pole, as shown in the web links given in Section 1 of this essay. Global measurements show that ozone has been depleted by 5% or more over much of the world. Thus, while the problem is most severe over the South Pole, ozone depletion is an issue that the whole world must consider.