III. Sources of Ozone Depletion

A. General Findings

Based on data collected since the 1950s, scientists have determined that ozone levels were relatively stable until the late 1970s. Severe depletion over the Antarctic has been occurring since 1979 and a general downturn in global ozone levels has been observed since the late 1970s.

Concerns about the ozone layer were first raised in the late 1960s when it was argued that water vapor and oxides of nitrogen from proposed subsonic and supersonic aircraft might deplete stratospheric ozone. Although the effect of aircraft, and later, space shuttles, was determined to be negligible, a new concern soon arose.

In 1974, two scientists, Drs. F. Sherwood Rowland and Mario Molina hypothesized that CFCs were able to persist in the atmosphere long enough to diffuse upward into the stratosphere. Once there, intense solar radiation would break them up, releasing reactive chlorine atoms which would then destroy ozone. Their theory initially met with skepticism but mounting evidence and the discovery of the Antarctic ozone hole in 1985 galvanized the interest of scientists and policy makers.

Subsequent research has indicated that chemical contamination from our industrialized society is the root cause of ozone layer depletion.

Several catalytic chemical reactions have been identified as ozone destruction mechanisms. The chemicals that start these reactions are called catalysts because they are not used up by the reaction. Rather, they are regenerated by the reaction and therefore are capable of reacting with ozone over and over again. Each of them can destroy thousands or even hundred’s of thousands of ozone molecules before being destroyed itself by some other process. The chemicals involved in these catalytic reactions include chlorine oxide, hydrogen oxide, and nitrogen oxide. Relatively recently, human activities have introduced large quantities of these catalysts into the atmosphere.

For over 50 years, chlorofluorocarbons (CFCs) were thought of as miracle substances. They are stable, nonflammable, low in toxicity, and inexpensive to produce. Over time, CFCs found uses as refrigerants, solvents, spray cans, foam packaging and in other smaller applications. Other chlorine-containing compounds include methyl chloroform, a solvent, and carbon tetrachloride, an industrial chemical. Halons, extremely effective fire extinguishing agents, and methyl bromide, an effective produce and soil fumigant, contain bromine. All of these compounds have atmospheric lifetimes long enough to allow them to be transported by winds into the stratosphere. Because they release chlorine or bromine when they break down, they damage the protective ozone layer. The discussion of the ozone loss chemistry in the subsection below focuses on CFCs, but the basic concepts apply to all of the ozone-depleting substances (ODS).

Researchers have also examined the potential impacts of other chlorine sources, such as swimming pools, industrial plants, sea salt, and volcanoes. However, chlorine compounds from these sources readily combine with water and repeated measurements show that they rain out of the troposphere very quickly before they have a chance to reach the stratosphere. In contrast, CFCs are very stable and do not dissolve in rain. There are no natural processes that remove the CFCs from the lower atmosphere. Over time, winds drive the CFCs into the stratosphere.

The CFCs are so stable that only one process breaks them down: exposure to strong UV radiation. When a CFC molecule breaks down, it releases atomic chlorine. One chlorine atom can destroy over 100,000 ozone molecules. The net effect is to destroy ozone faster than it is naturally created.

Large fires and certain types of marine life produce one stable form of chlorine that does reach the stratosphere. However, numerous experiments have shown that CFCs and other widely-used chemicals produce roughly 82% of the chlorine in the stratosphere, while natural sources contribute only 18%. Figure 14 shows the relative contribution of the various man-made and natural sources of chlorine in the stratosphere (WMO, 1994) .

Figure 14
Primary Sources of Chlorine Entering the Stratosphere

Other much smaller sources of ozone depleting chemicals include the direct injection of nitrogen oxide from the exhausts of supersonic aircraft and other aircraft flying at high altitudes. The widespread use of artificial fertilizers may also release nitrogen oxides into the stratosphere, although this potential effect is not yet fully defined.

Large volcanic eruptions can have an indirect effect on ozone levels. Although Mt. Pinatubo’s 1991 eruption had no effect on stratospheric chlorine concentrations, it did produce large amounts of particles called aerosols. These aerosols increase chlorine’s effectiveness at destroying ozone. The aerosols only increased depletion because of the presence of CFC-based chlorine. In effect, the aerosols increased the efficiency of the CFC ozone destruction mechanism, lowering ozone levels even more than would have otherwise occurred. Unlike long-term ozone depletion, however, this effect is short-lived. The aerosols from Mt. Pinatubo have already disappeared, but satellite, ground-based, and balloon data show ozone depletion continuing.

B. Key Research Findings

In 1985, scientists from the British Antarctic Survey reported that the ozone layer over Antarctica had shrunk each September and October since the late 1970s, which corresponds to the start of the Southern Hemisphere's spring season. In 1989, the first comprehensive research expedition of the Northern Hemisphere's polar region showed that the Arctic stratosphere is loaded in winter with the same destructive chlorine species which contribute to ozone depletion (Zurer, 1993}.

Subsequently, three years of data from NASA's Upper Atmosphere Research Satellite (UARS) provided conclusive evidence that human-made chlorine in the stratosphere is the cause of the Antarctic ozone hole. UARS instruments found chlorofluorocarbons (CFCs)--human-made products used in electronics and refrigeration systems—in the stratosphere. The satellite’s global data set also has traced worldwide buildup of stratospheric fluorine gases corresponding to the breakdown of CFCs, according to NASA scientists. (NASA, 1994)

In the past few years, some debate has occurred over the origin of ozone-destroying chlorine. Sea spray and volcanic gases have been put forth as possible sources for chlorine reaching the stratosphere. The UARS data have ended that debate. "These new results confirm our theories about CFCs," said Dr. Mark Schoeberl, UARS Project Scientist. "The detection of stratospheric fluorine gases, which are not natural, eliminates the possibility that chlorine from volcanic eruptions or some other natural source is responsible for the ozone hole." (NASA, 1994)

In addition to CFCs, UARS has detected hydrogen fluoride, a product of the chemical breakdown of CFCs, in the stratosphere. "Hydrogen fluoride has no natural source, it is not produced by volcanic eruptions or salt spray," said Dr. Anne Douglass, UARS Deputy Project Scientist. "Furthermore, scientists can calculate how much chlorine in the stratosphere is man-made using the hydrogen fluoride data." This calculation shows that almost all of the chlorine in the stratosphere comes from human-made chlorofluorocarbons. (NASA, 1994)

Two global networks of monitoring stations, now funded by the National Aeronautics & Space Administration and by the National Oceanic & Atmospheric Administration, have been measuring CFCs in the atmosphere since 1978. That research has confirmed that fully halogenated compounds are essentially inert in the troposphere, gradually floating unchanged into the stratosphere.

In the past 50 or so years, chlorine concentrations in the stratosphere hare increased from their background level of about 0.5 ppb to around 3.5 ppb today. Because of their long lifetimes, it will take several centuries, even after production stops, before the atmosphere is free of CFCs and Halons. The rise in atmospheric concentrations of selected ozone depleting chemicals between 1979 and 1994 is shown in Figure 15.

Figure 15
Atmospheric Concentration of
Selected Ozone-Depleting Chemicals

 

C. Ozone Chemistry

Ozone is produced continually in the upper stratosphere where solar UV radiation (hv) dissociates molecular oxygen (O2) to form atomic oxygen (O):

O2 + hv --> O + O

O + O2 --> O3

These reactions occur very rapidly in the stratosphere over the tropics, where solar radiation is most intense. But even though most ozone is produced at low latitudes, it is not as abundant there as it is at higher latitudes, because circulation in the stratosphere constantly moves ozone away from the equator toward both poles.

Ozone is destroyed when it absorbs UV light that otherwise would reach Earth's surface:

O3 + hv --> O2 + O

There is no net ozone depletion, however, because the process produces atomic oxygen that reacts with molecular oxygen to form another ozone molecule. Refer to Figure 5.

Ozone also is continually being destroyed through reactions with families of naturally occurring radicals that contain chlorine, nitrogen, hydrogen, or oxygen atoms. Atmospheric scientists like to use an analogy of a bathtub to describe how natural production and destruction processes are roughly in balance: Even with the drain unplugged, the level of water in the tub will remain constant so long as water from the tap is entering as fast as water is flowing out.

If more holes are punched in the bathtub, however, the level of water will drop. That's exactly the trouble caused by chlorine and bromine carried into the stratosphere by chlorofluorocarbons (CFCs) and halons (fluorocarbons that contain bromine).

Once CFCs and halons rise above the bulk of the ozone layer (which is most dense between 15 and 30 km altitude, depending on latitude), they are photolyzed by ultraviolet light, producing halogen atoms. For example

CCl2F2 (CFC-12) + hv -> CCIF2 + Cl

These chlorine atoms can abstract a hydrogen from methane to form hydrogen chloride (HCl). Or they can participate in catalytic cycles that destroy ozone, such as

Cl + O3 --> ClO + O2

ClO + O --> Cl + O2

O3 + O --> 2O2 (net)

Because the cycle regenerates the radicals that actually attack ozone, one chlorine atom can destroy hundreds of thousands of ozone molecules. This process of catalytic ozone destruction brought about by the presence of CFC molecules in the stratosphere is shown in Figure 16.

The simple chlorine-chlorine monoxide cycle can only occur at high altitudes where there are enough free oxygen atoms to allow the second step in the chain to run. A wealth of evidence has accumulated to show that these reactions are indeed destroying small amounts of ozone at about 40 km above Earth. Other families of radicals also participate in catalytic cycles that destroy ozone—nitrogen oxides, for example

NO + O3 --> NO2 + O2

NO2 + O --> NO + O2

-----------------------------------

O3 + O --> 2O2 (net)

The nitrogen oxide cycle lies behind concern over the effects that supersonic aircraft flying in the stratosphere may have on ozone. The planes' exhaust injects additional nitrogen oxide into the stratosphere, perhaps accelerating ozone depletion.

Until recently, nitrogen oxide cycles were thought to be the predominant natural loss process for ozone in the lower stratosphere. Recent research, however, suggests that the hydrogen radical family my be even more important

OH + O3 --> HO2 + O2

HO2 + O3 --> HO + 2O2

-----------------------------------

2O3 --> 3O2 (net)

 

Figure 16
Catalytic Ozone Destruction Caused by Man-made Compounds (e.g. CFCs)

 

Reactions can also occur on sulfate aerosols that encircle Earth year-round to release chlorine radicals. Heterogeneous chemistry, the importance of which has only begun to be understood since the discovery of the Antarctic ozone hole, shifts the balance between active and inactive chlorine from that which would exist if only go phase reactions took place. See discussion of the effects of polar stratospheric clouds below.

Members of the various families of radicals can react together, with varying effects on the fate of ozone. For example, when chlorine monoxide reacts with nitrogen dioxide or the chlorine radical reacts with methane, chlorine nitrate or hydrogen chloride are formed, both of which are relatively inert:

ClO + NO2 --> ClONO2

Cl + CH4 --> HCl + CH3

Hydrogen chloride and chlorine nitrate are examples of (chlorine) reservoir species that do not react with ozone, but that can later break down into destructive free radicals.

D. Polar Ozone Hole Chemistry

Different destruction mechanisms can predominate under special circumstances. The dramatic seasonal depletion of ozone known as the Antarctic ozone hole, for example, takes place at a time of year and in a region where essentially no oxygen atoms are present. The predominant destruction cycle there is one where chlorine atoms are regenerated by the reaction of chlorine monoxide with itself.

2(Cl + O3 --> ClO + O2)

ClO + ClO --> Cl2O2

Cl2O2 + hv --> Cl + ClOO

ClO2 --> Cl + O2

-----------------------------------

2O3 + hv --> 3O2 (net)

In Antarctica, the build up of active forms of chlorine is increased by the presence of polar stratospheric clouds. Normally, clouds do not form in the stratosphere due to the temperature inversion that occurs there. But, due to the extremely low prevailing temperatures, water and nitric acid condense to form ice clouds, known as polar stratospheric clouds. In addition, during the long dark Antarctic winter, stratospheric winds move in a circular pattern over the polar region creating a polar vortex that isolates the air above the Antarctic land mass.

Because nitric acid is tied up in the ice particles, the concentrations of oxides of nitrogen in the gaseous phase are significantly reduced. This in turns slows down the rate of conversion of chlorine oxide to the relatively inert chlorine nitrate in the reaction:

ClO + NO2 --> ClONO2

This reduction in the availability of gaseous NO2 helps to maintain high levels of active chlorine. Chlorine nitrate (ClONO2) along with (HCl) are called reservoir species because they bind up chlorine in a form which does not participate in catalytic ozone destruction.

In addition, the reservoir species, chlorine nitrate and hydrogen chloride, react on the surface of these polar stratospheric clouds to form molecules that dissociate into active radicals as soon as the Sun hits them. For example:

ClONO2 + HCl--> Cl2 + HNO3

Cl2 + hv --> 2Cl

As a result of all of these factors, when the sun returns to the Antarctic in the early Spring, large amounts of free chlorine become available to destroy the ozone over this area and create the now infamous Antarctic ozone hole. Figure 17 illustrates the conversion of chlorine from active to inactive states and the subsequent effect of polar stratospheric clouds on the inactive states.

Figure 17

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