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The Ozone Hole Tour!

Centre for
Atmospheric Science

University of Cambridge

Tour Home Page     Part I     Part II     Part III     Part IV


Part III. The Science of the Ozone Hole


Evidence that human activities affect the ozone layer has been building up over the last 20 years, ever since scientists first suggested that the release of chlorofluorocarbons (CFCs) into the atmosphere could reduce the amount of ozone over our heads.

The breakdown products (chlorine compounds) of these gases were detected in the stratosphere. When the ozone hole was detected, it was soon linked to this increase in these chlorine compounds. The loss of ozone was not restricted to the Antarctic - at around the same time the first firm evidence was produced that there had been an ozone decrease over the heavily populated northern mid-latitudes (30-60N). However, unlike the sudden and near total loss of ozone over Antarctica at certain altitudes, the loss of ozone in mid-latitudes is much less and much slower - only a few percentage per year. However, it is a very worrying trend and one which is the subject of intense scientific research at present. More on this in Part IV of the tour.

Many of these findings have since been reinforced by a variety of internationally supported scientific investigations involving satellites, aircraft, balloons and ground stations, and the implications are still being quantified and assessed. More about these international investigations in Part IV.

The Recipe For Ozone Loss

In trying to understand how the ozone loss occurs and the things that need to happen to destroy so much ozone, it helps to think of it as a 'recipe'. We need several ingredients to make the ozone loss occur. We'll now look at these 'ingredients' one at a time.

The Special Features of Polar Meteorology

We start by looking at the way the atmosphere behaves over the poles - the features of the meteorology in the stratosphere.

Schematic figure showing the vortex over Antarctica The figure to the right shows schematically what happens over Antarctica during winter. During the winter polar night, sunlight does not reach the south pole. A strong circumpolar wind develops in the middle to lower stratosphere. These strong winds are known as the 'polar vortex'. This has the effect of isolating the air over the polar region.

Since there is no sunlight, the air within the polar vortex can get very cold. So cold that special clouds can form once the air temperature gets to below about -80C. These clouds are called Polar Stratospheric Clouds (or PSCs for short) but they are not the clouds that you are used to seeing in the sky which are composed of water droplets. PSCs first form as nitric acid trihydrate. As the temperature gets colder however, larger droplets of water-ice with nitric acid dissolved in them can form. However, their exact composition is still the subject of intense scientific scrutiny. These PSCs are crucial for ozone loss to occur.

So, we have the first few ingredients for our 'ozone loss recipe'. We must have:

  1. Polar winter leading to the formation of the polar vortex which isolates the air within it.
  2. Cold temperatures; cold enough for the formation of Polar Stratospheric Clouds. As the vortex air is isolated, the cold temperatures persist.

Chemical Processes Leading To Polar Ozone Depletion

It is now accepted that chlorine and bromine compounds in the atmosphere cause the ozone depletion observed in the `ozone hole' over Antarctica and over the North Pole. However, the relative importance of chlorine and bromine for ozone destruction in different regions of the atmosphere has not yet been clearly explained. Nearly all of the chlorine, and half of the bromine in the stratosphere, where most of the depletion has been observed, comes from human activities.

The figure above shows a schematic illustrating the life cycle of the CFCs; how they are transported up into the upper stratosphere/lower mesosphere, how sunlight breaks down the compounds and then how their breakdown products descend into the polar vortex.

The main long-lived inorganic carriers (reservoirs) of chlorine are hydrochloric acid (HCl) and chlorine nitrate (ClONO2). These form from the breakdown products of the CFCs. Dinitrogen pentoxide (N2O5) is a reservoir of oxides of nitrogen and also plays an important role in the chemistry. Nitric acid (HNO3) is significant in that it sustains high levels of active chlorine (as explained soon).

Production of Chlorine Radicals

One of the most important points to realise about the chemistry of the ozone hole is that the key chemical reactions are unusual. They cannot take place in the atmosphere unless certain conditions are present: our first two ingredients in our recipe for ozone loss.

The central feature of this unusual chemistry is that the chlorine reservoir species HCl and ClONO2 (and their bromine counterparts) are converted into more active forms of chlorine on the surface of the polar stratospheric clouds. The most important reactions in the destruction of ozone are:

HCl + ClONO2 -> HNO3 + Cl2 (1)
ClONO2 + H2O -> HNO3 + HOCl (2)
HCl + HOCl -> H2O + Cl2 (3)
N2O5 + HCl -> HNO3 + ClONO (4)
N2O5 + H2O -> 2 HNO3 (5)
It's important to appreciate that these reactions can only take place on the surface of polar stratospheric clouds, and they are very fast. This is why the ozone hole was such as surprise. Heterogeneous reactions (those that occur on surfaces) were neglected in atmospheric chemistry (at least in the stratosphere) before the ozone hole was discovered. Another ingredient then, is these heterogeneous reactions which allow reservoir species of chlorine and bromine to be rapidly converted to more active forms.

The nitric acid (HNO3) formed in these reactions remains in the PSC particles, so that the gas phase concentrations of oxides of nitrogen are reduced. This reduction, 'denoxification' is very important as it slows down the rate of removal of ClO that would otherwise occur by the reaction:

ClO + NO2 + M -> ClONO2 + M (6)
(where M is any air molecule)

... and so helps to maintain high levels of active chlorine. Here is some more information on Polar Stratospheric Clouds.

This movie shows a 3D model simulation of how chlorine nitrate (ClONO2) changes during a northern hemisphere winter in the lower stratosphere. Remember that ClONO2 is destroyed when the PSCs form, so for a large part of the movie, you see nothing. But as sunlight returns to the polar night region over the Arctic we see the ClONO2 start to recover. This first happens around the edge of the polar vortex, and we the the now classic doughnut shape of the so-called 'chlorine nitrate collar'.

Inline Evolution of ClONO2 over the North Pole during winter 1994
(3.4 Mb)

MPEG Evolution of ClONO2 over the North Pole during winter 1994 (small)
(554 Kb)


MPEG Evolution of ClONO2 over the North Pole during winter 1994 (large)


The Return Of Sunlight

Lastly note that we have still only formed molecular chlorine (Cl2) from reactions (1)-(5). To destroy ozone requires atomic chlorine.

Molecular chlorine is easily photodissociated (split by sunlight):

Cl2 + hv -> Cl + Cl
This is the key to the timing of the ozone hole. During the polar winter, the cold temperatures that form in the 'vortex' lead to the formation of polar stratospheric clouds. Heterogeneous reactions convert the reservoir forms of the ozone destroying species, chlorine and bromine, to their molecular forms. When the sunlight returns to the polar region in the southern hemisphere spring (northern hemisphere autumn) the Cl2 is rapidly split into chlorine atoms which lead to the sudden loss of ozone. This sequence of events has been confirmed by measurements before, during and after the ozone hole.

There is still one more ingredient for our recipe of ozone destruction. We have most of it but we have still not explained the chemical reactions that the atomic chlorine actually takes part in to destroy the ozone. We'll discuss this next.

Catalytic Destruction of Ozone

Measurements taken of the chemical species above the pole show the high levels of active forms of chlorine that we have explained above. However, we still have many more atoms of ozone than we do of the active chlorine so how it is possible to destroy nearly all of the ozone?

The answer to this question lies in what are known as 'catalytic cycles'. A catalytic cycle is one in which a molecule significantly changes or enables a reaction cycle without being altered by the cycle itself.

The production of active chlorine requires sunlight, and sunlight drives the following catalytic cycles thought to be the main cycles involving chlorine and bromine, responsible for destroying the ozone:

(I)  ClO + ClO + M -> Cl2O2 + M
   Cl2O2 + hv -> Cl + ClO2
   ClO2 + M -> Cl + O2 + M
 then: 2 x (Cl + O3) -> 2 x (ClO + O2)
 net: 2 O3 -> 3 O2
(II)  ClO + BrO -> Br + Cl + O2
   Cl + O3 -> ClO + O2
   Br + O3 -> BrO + O2
 net: 2 O3 -> 3 O2

The dimer (Cl2O2) of the chlorine monoxide radical involved in Cycle (I) is thermally unstable, and the cycle is most effective at low temperatures. Hence, again low temperatures in the polar vortex during winter are important. It is thought to be responsible for most (70%) of the ozone loss in Antarctica. In the warmer Arctic a large proportion of the loss may be driven by Cycle (II).

The Recipe For Ozone Loss

To summarise then, we have looked at the 'ingredients' or conditions necessary for the destruction of ozone that we see in Antarctica. The same applies more or less to the loss of ozone in the Arctic stratosphere during winter. Although in this case the loss is not nearly so severe.

To recap then, the requirements for ozone loss are:

Part IV: Current Research Work at Cambridge

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© 1998. Centre for Atmospheric Science, Cambridge University, UK. No text or graphics can be used or reproduced without explicit written permission. This version designed and maintained by Dr. Glenn Carver. Original concept and design Owen Garrett.