Higher Atmosphere

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Chlorine chemistry

Primarily chlorine chemistry is driving the destruction of the ozone layer. With the industrial production of chlorofluorocarbons (CFCs) man brought a new chlorine source into the atmosphere. Now chlorine has six times the level it had from natural sources with fatal consequences for the ozone layer. However, conditions for the formation of the ozone hole are very special. Therefore a such drastic development has not been predicted.

CFCs - gases without natural source

The ozone-depleting gases with the largest potential to influence climate are CFC-11 (CFCl₃), CFC-12 (CF₂Cl₂), and CFC-113 (CF₂ClCFCl₂). It is now clear from measurements in polar firn air that there are no natural sources of these compounds. The only significant natural source of chlorine is methylchloride CH₃Cl and this compound has a comparable short lifetime of 1.3 years.
Due to their stability versus the OH radical and photolysis in the troposphere, CFCs have a very long tropospheric lifetime in the range of 50-100 years. They reach the stratosphere and are photolysed, which leads to a formation of chlorine radicals. The availability of these radicals does not necessarily lead to a significant ozone depletion, because other sink reactions compete. Special conditions are needed, as shown in the following sections.

Stratospheric chlorine chemistry
 - the basics

Generally, as many other radicals X, chlorine (Cl) is oxidised by ozone in the stratosphere and forms XO (ClO)

           X + O₃              -> XO + O2
           O₃ + sunlight        -> O + O2
           O + XO             -> X + O2
net:      2 O₃                       -> 3 O₂

This chain reaction drives the ozone depletion.

But the initiating radical X (here Cl) is not necessarily recycled. Cl or ClO can also be removed in other reactions. Nitrogen oxides act as sinks for ClO radicals, which are transferred into the so called reservoir species ClONO₂ and HCl, shown in the blue frame.

 ClO + NO₂ + M^*   -> ClONO₂ + M^*
and
 ClO + NO            ->  Cl + NO₂
 Cl + CH₄             -> HCl + CH₃

HCl and ClONO₂ are so called 'reservoir species', because here chlorine is not active. They do not react with ozone. Normally they stay in the gas phase and can be slowly removed again from the stratosphere. Therefore in normal stratospheric gas phase chemistry, only slight ozone depletion is expected. But these species are transported with the mean circulation to the lower stratosphere in the polar winter area ...

The special conditions of the Antarctic ozone hole

During the polar night with temperatures of about -80°C, the very small amount of water available in the stratosphere is able to freeze and to form polar stratospheric ice clouds together with nitric acid (so called nitric acid trihydrate NAT). Now five key conditions can come together:

First: The nitrogen oxide catalysts (nitrogen oxide NO and nitrogen dioxide NO₂), which help to convert ClO into HCl as shown above, are removed from the stratospheric gas phase through the reactions

NO + O₃              -> NO₂ + O₂
NO₂ + NO₃ + M^*   -> N₂O₅ + M
N₂O₅ + H₂O          -> 2 HNO₃  

thereby producing nitric acid HNO₃ which is incorporated in the particles of the polar stratospheric clouds (PSC).

Second: On the surface of the PSC ice particles the 'reservoir species' for unreactive chlorine, HCl and ClONO₂, react with each other to produce Cl₂ and HNO₃; the latter is immediately incorporated in the particles.

Third: After the return of daylight at the end of the polar night, Cl₂ is photolysed to produce 2 Cl radicals. The chlorine becomes reactivated.

Fourth: The chlorine atoms start a catalytic chain of reactions, leading to the destruction of ozone as long as no nitrogen oxides are available to remove them. Note, that the speed of ozone destruction is quadratic in the chlorine (or ClO) concentration.

    Cl + O₃                       -> ClO + O₂
    Cl + O₃                       -> ClO + O₂
    ClO + ClO + M           -> Cl₂O₂ + M
    Cl₂O₂ + sunlight -> Cl +ClO₂ -> 2 Cl + O₂
Net: 2 O₃  -> 3 O₂ 

Fifth: Normally chlorine species as Cl, ClO, and Cl₂O₂ are formed and concentrated more in the upper stratosphere and ozone more in the lower stratosphere. Therefore several decades ago experts did not expect ozone to decrease significantly. Ozone and ozone killers should only come together in border zones. At this point, the polar vortex comes into play: Chlorine species are advected to the lower stratosphere by downwind transport from the middle and upper stratosphere within this meteorologically stable vortex (circumpolar wind) with the pole more or less at the center. This is how it comes that the ozone-destroying chlorine species are transported down to lower altitudes, to this region where most of the ozone is accumulated.

All five conditions have to come together, to form the ozone hole. This is why the major ozone depletion occurs over Antarctica and only in the Antarctic spring months September / October as soon as the sun rises after the polar night. In some years, we have comparable conditions over the Arctic in March and a little ozone hole forms over Northern Europe as well. Later in the year, the polar clouds dissolve, nitrogen oxides become available again, the vortex breaks together, chlorine radicals are removed and the ozone layer recovers.

M*: In any sort of reaction A + B -> C a third partner is needed, which takes away excess energy. Otherwise the product C would have the same energy than the sum of the educts A + B and directly react back. In most cases M is the nitrogen from the air N₂.

About this page:

author: Dr. Elmar Uherek - Max Panck Institute for Chemistry, Mainz
scientific reviewer: Dr. Christoph Brühl, Max Planck Institute for Chemistry, Mainz
educational proofreading: Michael Seesing - Uni Duisburg - 2003-08-07
last published: 2004-05-11