|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|||||||
|
|
|
|
|
|
|
|
|
||
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).
NEXT: Chlorinated Fluorocarbons (CFCs)
