Ozone has a Dr. Jeckyl and Mr. Hyde personality.  Tropospheric ozone (Mr. Hyde) is an air pollutant and a key ingredient in photochemical smog.  Stratospheric ozone (the ozone layer) is beneficial because it absorbs dangerous high-energy ultraviolet light.  Here's a review of how ozone is produced and destroyed in the stratosphere, a list of some of the harmful effects of ultraviolet (UV) light, and information about anthropogenic (man-caused) destruction of stratospheric ozone and the ozone hole.  You'll find most of the figures below on pps. 17-22 in the photocopied ClassNotes.


The top two equations show how ozone is produced in the stratosphere.  Ultraviolet (UV) light splits an O2 molecule into two O atoms (photodissociation).  Each of the O atoms can react with unsplit O2 to make O3 (ozone).

Ozone is destroyed when it absorbs UV light and is split into O and O2 (the two pieces move away from each other and don't recombine and remake ozone).  This is how the ozone protects us.  O3 is also destroyed when it reacts with an oxygen atom (thereby removing one of the "raw ingredients" used to make ozone).  Two molecules of ozone can also react with each other to make 3 molecules of O2 (probably the least likely of the three possibilities just because there aren't many ozone molecules around to react with each other).

The ozone concentration in the stratosphere will shift up and down until the natural rates of production and destruction balance each other (analogous to your bank account not changing once the amounts of money being deposited and withdrawn are equal).  The black box represents the O3 layer concentration once equilibrium is achieved.  If an additional man-caused destruction process is added (dotted red arrow) the ozone layer concentration will decrease (sort of like someone else coming along and starting to spend some of the money in your bank account, your balance will decrease).

Knowing that you need O2 and UV light to make ozone, you can begin to understand why the ozone layer is found in the middle of the atmosphere.


There is plenty of UV light high in the atmosphere but not much oxygen (air gets thinner at higher and higher altitude).  Near the ground there is plenty of oxygen but not as much UV light (it is absorbed by gases above the ground).  You find the optimal amounts of UV light and oxygen somewhere in between, near 25 km altitude.

This next figure lists some of the problems associated with exposure to UV light.  Thinning of the ozone layer will result in increased amounts of UV light reaching the ground.

Skin cancer and cataracts are probably the best known hazards associated with UV light. 

You may have heard of UVA and UVB; there is also UVC.  UVA is a relatively long wavelength (0.315 to 0.4 micrometers), low energy form of UV light; it is the light emitted by a "black light".  UVB has shorter wavelength (0.28 to 0.315 micrometers).  UVC has even shorter wavelength (0.1 to 0.28 micrometers) and is the most dangerous of the three types of UV light.  Germicidal bulbs emit UVC light and are used to sterilize and purify air, food, and water.  Fortunately All of the UVC in sunlight is absorbed by the upper atmosphere. 

There is some question about whether tanning booths, which emit mostly UVA, are safe.  You can find information online about this question.  Here is an example from the US Food and Drug Administration.

It is worth mentioning that thinning of the ozone layer and increased amounts of UV light reaching the ground is not the cause of global warming.  The worry is that increasing concentrations of greenhouse gases and strengthening of the greenhouse effect might cause global warming.
 

Human activities add substances to the atmosphere that can potentially reduce ozone concentration in the ozone layer (which would result in increased exposure to UV light at the ground). 

The first set of reactions above involve nitric oxide, NO.  First, NO reacts with O3 to form NO2 and O2. Then notice the NO2 reacts with an oxygen atom (one of the raw ingredients needed to make O3) to form NO again and O2. The NO is available again to react with and destroy another ozone molecule.

At one time many countries were considering building fleets of supersonic aircraft that would fly in the stratosphere.  The plans were scrapped partly due to concern that the NO emissions from jet engines would damage the ozone layer.

The main threat now comes from chlorofluorocarbons (CFCs).  CFCs were at one time thought to be an ideal industrial chemical and had a variety of uses.  CFCs are unreactive, non toxic, and stable.  Once they get into the atmosphere they remain there a long time, as much as 100 years.   The reactions involving CFCs are shown on the next figure (p. 19 in the ClassNotes).



CFCs released at ground level [lower left corner in the figure above] remain in the atmosphere long enough that they can eventually make their way up into the stratophere.  UV light can then break chlorine atoms off the CFC molecule Point (a)).  The resulting "free chlorine" can react with and destroy ozone.  This is shown in (b) above.  Note how the chlorine atom reappears at the end of the two step reaction.  A single chlorine atom can destroy 100,000 ozone molecules before undergoing a different reaction and being removed from the atmosphere.

There are ways of keeping chlorine from reacting with ozone.  A couple of these so called "interference reactions" are shown in (c) above.    The reaction products, reservoir molecules (because they store chlorine),  might serve as condensation nuclei for cloud droplets (the small water drops that clouds are composed of) or might dissolve in the water in clouds.  In either event the chlorine containing chemical is removed from the atmosphere by falling precipitation.  Clouds are probably the most effective way of cleaning the atmosphere.



The ozone hole that forms above the S. Pole every year in late September-early October was one of the first real indications that CFCs could react with and destroy stratospheric ozone.  The hole is not really a hole in the ozone layer, just a temporary thinning of the ozone layer above the S. Pole and the continent of Antarctica.  The ozone concentration decreases to perhaps 30% of its normal value.



It is unusual to find clouds in the stratosphere, most of the water vapor stays in the troposphere.  However, because it gets very cold above the S. Pole in the winter, polar stratospheric clouds do sometimes form (they are made from water and other materials). This together with an unusual wind pattern above the S. Pole in the winter (the polar vortex) are thought to create the ozone hole when the sun returns in the spring.

The ozone destruction reactions are circled in red above.  Cl reacts with O3 to make ClO.  This reacts with O to produce Cl and O2.  The Cl is now available to react again with other ozone molecules.

In blue is one of the "interference" reactions.  ClO reacts with NO2 to make ClNO3.  The Cl in this "reservoir" molecule is stored in a form that can't react with any more ozone.

Now what happens above the S. Pole in the winter is that the reservoir molecules react on the surfaces of the polar stratospheric cloud particles to make a new and different kind of compound.  This reaction is shown in green above.  The Cl that would be stored away in ClNO3  reacts and produces a new compound, HOCl, that accumulates in the air during the winter.  When the sun reappears in the spring, the UV light splits off all the Cl molecules which react with ozone.  A lot of chlorine suddenly becomes available and the ozone concentration takes a nosedive.