1-8 Man-made Chemicals, CFCs

Ozone Depletion

Ozone (O₃) occurs naturally in the stratosphere because of intense ultra-violet (UV) radiation on oxygen (O₂) molecules. UV radiation also destroys ozone by the Chapman process, but the balance of production and destruction leads to a constant level of ozone (in the absence of CFCs and chlorine) that is much higher than in the troposphere.

Figure 1 is the first published evidence of major ozone depletion over Antarctica.

This graph gives spring (in the Southern Hemisphere) ozone concentrations, in Dobson units, over a 30-year period. In the 1950's the ozone concentration was about 320 Dobson units but has dropped dramatically in recent years to about 200 as shown on this graph. In fact in more recent years it has dropped to 100 averaged over a fairly large area, and in some places it has dropped completely to zero - no ozone at all.

This peculiar behavior in ozone concentrations was not limited to Antarctica, but the effect was most pronounced in that region. Figure 2 gives observations from Switzerland plotted as the difference in average ozone levels between the period 1931-1969 and 1970-1986.

The graph shows that the reduction in ozone over this location in recent years is most pronounced in the winter months.

Scientists became concerned about the trend of these observations, but thought that some unaccounted-for drift or other problem had developed with their instruments. It wasn't until surface measurements (see photo of instrument) were compared with satellite measurements and observations were compared from different locations that it was realized that ozone concentrations were significantly depleted at certain times and in certain places.

Figure 3 shows data taken more recently from a vertical profile of partial pressure of ozone, in nanobars, as a function of height.

Recall that the troposphere extends from the surface to about 10 kilometers and the region beyond it is the stratosphere, which continues up to about 50 kilometers. These ozone measurements were taken in the lower stratosphere over the South Pole during the Southern Hemisphere spring in 1992. It is apparent from these plots that ozone concentrations drop dramatically over a very short period near the end of September to readings of essentially zero in the region around 15 to 20 kilometers above sea level. Complete destruction of the ozone layer occurred over a fairly narrow vertical range. Measurements of total ozone in vertical columns of the atmosphere are made by the Total Ozone Mapping Spectrometer (TOMS).

Reactions and Compounds that Destroy Ozone

Reactions that destroy ozone are shown in Figure 4.

Chlorine monoxide (ClO) reacts with an ozone molecule (O₃) to produce a chlorine molecule (Cl₂) and diatomic oxygen (O₂), so that, in effect, we've taken an O₃ and an O and created 2O₂ molecules. The chlorine has not been neutralized in this process, but is free to repeat the reaction and does so as many as 100,000 times before being captured. Chlorine is not the only constituent that can do this: NO and the OH radical can do the same thing, as shown in the figure.

So if chlorine (along with NO and OH, which will be considered later) causes ozone destruction in the stratosphere, where does it come from, how does it get there, and why is there more of a problem now than 50 years ago? The major source of free chlorine at the earth's surface is sea salt, NaCl, which can dissociate and provide an abundant supply of chlorine atoms over the global oceans. While this represents a very large natural source of free chlorine, only a tiny fraction survives to reach the stratosphere (Figure 5). The amount that does escape from the troposphere contributes to the natural destruction of ozone, which balances natural creation by solar radiation.

Anthropogenic sources of chlorine include industrial and household processes using chlorine for cleaning, disinfecting, and bleaches. Chlorine also is widely used in water supplies and swimming pools. These represent very large anthropogenic releases of free chlorine, but these sources are not considered to be the culprits leading to ozone destruction as revealed by the previously shown graphs.

The reason for this, and the reason only a very tiny fraction of Cl from the ocean reaches the stratosphere is that chlorine is very reactive in the lower atmosphere (Figure 6) and quickly combines with some other atoms or molecules to become neutralized near the surface. To get free chlorine to the stratosphere without destruction in the lower atmosphere, we would have to design a compound that would protect the chlorine from chemical reactions in the troposphere and carry it sufficiently high to where the density of potential reacting molecules is small, and then release it. Enter chlorinated fluorocarbons (CFCs).

Chlorinated Fluorocarbons (CFCs)

CFCs were invented in 1928 by Thomas Midgley of the DuPont Corporation and have become widely used because of their many advantageous physical and chemical properties.

Chlorine-containing molecules, such as the CFC's, CF₂, CL₂, CCl₃F, that have long enough lifetimes will diffuse to the stratosphere where they are acted on by ultraviolet radiation.

A chlorine atom is ripped off, and the destruction of ozone begins. We might ask why in the stratosphere and why over Antarctica? The measurements show some decrease at other latitudes, but why is the effect so pronounced specifically over Antarctica?

To answer these questions, we need to understand a little about the meteorology and chemistry of the stratosphere in the Southern Hemisphere (Figures 7 and 8). In July the South Pole has the polar night, or polar winter, so the sun remains below the horizon all day and the whole atmosphere from the surface to the lower stratosphere gets very cold. Temperatures in the stratosphere drop to -90 degrees Celsius and clouds of ice particles begin to form. The ice crystals in these clouds provide a local surface for heterogeneous chemistry to take place.

Hydrochloric acid (HCl) combines with chlorine nitrate to form Cl₂ and NONO₂. On the surface of the cloud ice crystals, H₂O from the ice combines with the chlorine nitrate to form HOCl and NONO₂. When the sun rises over the horizon in early spring (about the end of August or first part of September), atomic chlorine is released by the first rays of solar energy. However, the nitrogen oxides, which are the only known scavengers of chlorine, are held on the ice crystals.

The free chlorine is left uninhibited to convert ozone to diatomic oxygen by the reactions in Figure 4 until additional warming releases the nitrogen oxides. The time between the release of chlorine and the release of the nitrogen oxides allows chlorine to destroy ozone without interruption until the ozone is all gone. In summary, stratospheric ice cloud particles under extremely cold conditions in the absence of sufficient sunlight to release the nitrogen compounds are what allow free chlorine to convert ozone to diatomic oxygen in very large amounts.

NASA scientists have clearly demonstrated this connection by flying through these clouds and measuring concentrations of chlorine, chlorine nitrate, and ozone. The chlorine - chlorine nitrate - ozone depletion linkage is very well established. Three stratospheric chemists have won the Nobel prize in chemistry for their work in clarifying these conditions. Despite this overwhelming scientific evidence for the linkage between increasing concentrations of CFCs and ozone depletion, articles continue to appear in otherwise very respectable business magazines questioning this connection. Such articles typically cite some scientist with no published work in this area as their authorities. This is a clear case of the need to carefully and critically review the evidence for claims that go counter to the scientific consensus.

Lifetimes of Chlorine-Containing Compounds

Figure 9 gives the lifetimes of different chlorine-containing compounds.

Note that the atmospheric lifetimes, range from about 60 to 400 years, with uncertainty levels given in parentheses. Therefore some of the CFCs we release today will be giving up chlorine leading to ozone destruction well beyond the lifetimes of people living today. The halons (fluorocarbons, chlorocarbons, carbon tetrachloride), similarly, have long lifetimes, although not as long as the CFCs. Some uses of the CFCs and halons, such as refrigeration, lubricants, foams and aerosols, are given in the table as well. Manufacture and use of some of the CFCs and halons have been dramatically reduced at least in the US, but not world wide.

It is possible to reduce the adverse effects of CFCs by adding a hydrogen atom to the molecular structure. This makes the molecule somewhat more reactive and less likely to survive long enough to diffuse to the stratosphere. The resulting molecule is called an HCFC or HFC. For example, HFC-134a is a possible replacement for CFC-12, but requires a much more complicated process for production, as is shown in Figure 10.

A recent graph from National Geographic shows trends over the last 20 years in global production and sale ("consumption") of CFCs and atmospheric concentrations of CFCs in Antarctica.

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1-8: Man-made Chemicals, CFCs