CFC's, Ozone and Free Radicals

Much of this page is only suitable for advanced chemistry students!

Summary of the ideas and issues involving ozone

Abbreviations used:

CFC/CFCs = chlorofluorocarbon; HCFC/HCFCs = hydrochlorofluorocarbon; HFC/HFCs = hydrofluorocarbon

is a typical CFC molecule CCl₂F₂, dichlorodifluoromethane, known commercially as CFC-12.

These kinds of molecules 'where' used in aerosol sprays as propellant gases and refrigerant gases.

Unfortunately, when CFCs get into the atmosphere at ground level, because they are chemically inert, they diffuse up into the upper atmosphere where they are decomposed by ultraviolet (uv) radiation producing chlorine radicals.

(chlorine atoms are free radicals with an unpaired electron)

The chlorine radicals decompose (destroy) ozone via free radical chain reactions and it is estimated that one chlorine atom can lead to the decomposition of 100, 000 ozone molecules in a catalytic cycle.

In the Antarctic winter the chlorine radicals build up on small ice crystals in the air high up in the upper atmosphere.

When these ice crystals melt in the spring sun, the chlorine radicals are released causing massive depletion of the ozone layer - in fact they 'did' cause a large hole in the (uv protecting) polar ozone layer of the southern hemisphere.

Fortunately, CFCs are now banned and new refrigerant coolants and aerosol propellant gases have been developed that contain hydrogen atoms - known as HCFCs and HFCs

Examples: a HCFC CHClF₂ chlorodifluoromethane and a HFC difluoromethane CH₂F₂

HCFCs and HFCs are more reactive and are decomposed at lower altitudes, before they can diffuse up into the ozone layer and HFCs obviously cannot generate chlorine radicals.

However, they still have the disadvantage of being powerful greenhouse gases!

• The problem!

  • The liquids and gases in refrigerators and aerosols etc. has had a very negative effect on the ozone concentration in the upper atmosphere.

  • Some of the chemistry of the upper atmosphere (stratosphere) involves the creation and destruction of ozone.

  • Unfortunately, CFCs get involved in this chemistry and destroy ozone - which is only present in concentrations of ~2-8 parts per million.

  • Ozone absorbs most of the most harmful high energy uv radiation that can cause damage to DNA in skin cells and hence the potential development of skin cancer.

  • BUT, here we have a success story, by stricter regulation and using alternative chemicals, the situation is being reversed.

  • Note:

  • In the text uv is an abbreviation of ultraviolet radiation, whose photons are higher in energy than visible electromagnetic radiation.

  • The single dot (•)represents the unpaired electron on the free radical (explained in detail later) and, wherever possible, I've put the 'electron dot' on the atom associated with the unpaired electron.

• CFC's – what is so good about them? (before we get into the problems they cause!)

  • Chlorofluorocarbons (CFCs) are covalently bonded relatively small organic molecule of carbon, chlorine and fluorine atoms e.g.

  •  dichlorodifluoromethane, (CFC–12), trichlorofluoromethane, (CFC–11).

  • both these CFCs are made by replacing all the hydrogen atoms in methane with chlorine and fluorine atoms.

  • CFCs have very useful in many applications with their advantageous properties. They are non–toxic, non–flammable and insoluble in water. They also have low boiling points – typical of small volatile covalent molecules.

  • CFCs were used as heat exchanging coolants in refrigerators, in air–conditioning systems, propellants in aerosol spray cans (now banned in Europe and US), expanded foams e.g. polystyrene, cleaning solvents e.g. cleaning agents for electrical components.

    • In other words CFCs are very useful chemicals.

• The discovery!

  • It was in the early 1970s that scientists found out that chlorine could be involved with destroying ozone in the ozone layer of the upper atmosphere.

  • In the 1980s scientists produced evidence for a decrease in ozone levels in the atmosphere above Antarctica.

  • These findings were described as a 'hole in the ozone layer'.

  • Further tests on chemicals in the upper atmosphere showed the presence of CFCs breaking down and this chemistry was facilitating the breakdown of ozone itself.

  • As the scientific evidence accumulated, we are now sure that CFCs are one cause of depletion in ozone concentration in the upper atmosphere.

• An introduction to free radical chemistry (for advanced level chemistry students)

  • Its essential to know about this if you want to understand why CFCs destroy ozone

  • If enough energy is supplied by heat or by visible/uv electromagnetic radiation (higher energy photons), and the bond is weak enough, it can break.

  • A covalent bond can break in two ways. This illustrated with the molecule chloromethane CH₃Cl.

  • The bond breaks unevenly where the electron bond pair can stick with one fragment and a positive and negative ion form.

    • e.g. CH₃Cl  ===>  CH₃⁺  +  Cl^–   (at Advanced Level this is called heterolytic bond fission)

    • shows what happens to the molecule.

  • The bond can also break evenly, where the bonding pair of electrons are equally divided between two highly reactive fragments called free radicals.

    • Free radicals are characterised by having an unpaired electron not involved in a chemical bond.

    • The ^. means the 'lone' electron on the free radical, which is not part of a bond anymore, wants to pair up with another electron to form a stable bond – that's why free radicals are so very reactive!

    • e.g.  CH₃Cl  ===>  CH₃•  +  •Cl    (at Advanced Level this is called homolytic bond fission)

    •   shows what happens to the molecule

    • Free radicals can be fragments of molecules or single atoms.

    • The single dot (•)represents the unpaired electron on the free radical.

• In the stratosphere small amounts of unstable ozone O₃ (trioxygen) are formed by free radical reactions.

• Why CFCs are used and are now banned

  • The chemistry of free radicals is important in the current environmental issue of ozone layer depletion.

  • Chlorofluorocarbons (CFC's for shorthand) are organic molecules containing carbon, fluorine and chlorine

    • e.g.  dichlorodifluoromethane has the formula CCl₂F₂ 

    • and chlorotrifluoromethane  (shown in right diagrams).

  • They are very useful low boiling organic liquids or gases, until recently, extensively used in refrigerators and aerosol sprays e.g. repellents.

  • They are relatively unreactive, non–toxic and have low flammability, so in many ways they are 'ideal' for the job they do.

  • The relatively strong C-F and C-Cl bonds gives CFC molecules their chemical stability.

  • However it is their chemical stability in the environment that eventually causes the ozone problem but first we need to look at how ozone is formed and destroyed in a 'natural cycle'.

  • This presumably has been in balance for millions of years and explains the uv ozone protection in the upper atmosphere.

• How is ozone formed? Why is the ozone layer so important to life on Earth?

  • Ozone is formed and destroyed in the stratosphere by free radical reactions.

  • 'ordinary' stable oxygen O₂ (dioxygen) is split (dissociates) into two by high energy ultraviolet radiation (uv photon energy 'wave packets) into two oxygen atoms (which are themselves radicals) and then one of these 'free' oxygen atom radicals combines with an oxygen molecule to form the molecule ozone (O₃ trioxygen).

    • (1)   O₂  == uv photon ==> 2O•

    • (2)  O•  +  O₂ ===>  •O₃     (formation of ozone)

  • Ozone molecules are found in the 'ozone layer' high up in the stratosphere, part of the upper atmosphere.

  • The ozone is a highly reactive and unstable molecule and decomposes into dioxygen when hit by other uv light photons.

    • (3)  •O₃  +  == uv photon ===> O₂  +  O•    (removal of ozone by uv)

    • The oxygen molecule and oxygen atom can then rejoin to make ozone (2), so you have a natural 'recycling system' of ozone decomposition and ozone formation.

    • This last reaction is the main uv screening effect of the upper atmosphere and the ozone absorbs a lot of the harmful incoming uv radiation from the Sun. It is the higher energy uv photons that are most likely to be absorbed by the ozone, and this is the most harmful part of the ultraviolet radiation spectrum.

  • Oxygen atoms can recombine to form the stable oxygen molecule: (4) O•  +  O• ===> O₂  (termination step)

  • Ozone is 'naturally' destroyed by reaction (5)   O•  +  •O₃ ===> 2O₂   (termination, removal of ozone by oxygen atoms)

  • So, even before the introduction of CFCs we have five free radical reaction just involving oxygen in which ozone is created and destroyed!

  • BUT, until CFCs entered the atmosphere there was balanced situation of a fairly constant and protective concentration of ozone in the upper atmosphere.

• The unplanned intervention of human activity!

  • However until the intervention of human activity, these five reactions, along with other free radical chemistry allowed a fairly constant level of ozone to exist and protect our planet from too much harmful ultra-violet radiation.

  • If the ozone levels are reduced more harmful uv radiation reaches the Earth's surface and can lead to medical problems such as increased risk of sunburn and skin cancer and it also accelerates skin aging processes.

  • There is strong evidence to show there are 'holes' in the ozone layer with potentially harmful effects, so back to the CFC problem for some explanations and solutions!

• The mechanism and chemistry of ozone depletion

  • The chemistry of ozone formation and its depletion/destruction is very complex.

  • There are dozens of free radical reactions going on in the upper atmosphere and only some of them are described below to give you an idea of how using 'manmade' chemicals led to the serious situation describe above.

  • However, the problem of ozone levels in the atmosphere has, and is being monitored and alternative chemicals introduced to improve the situation.

  • The chemically very stable CFCs diffuse up into the stratosphere and eventually decompose when hit by ultraviolet light (uv) high energy photons to produce free radicals, including free chlorine atoms, which themselves are highly reactive free radicals.

  • e.g. for the CFC dichlorodifluoromethane   and  chlorotrifluoromethane

  • (6a)   CCl₂F₂  == uv ==>  Cl•  +  •CClF₂  or  (6b)  CCl₃F  == uv ==>  Cl•  +  •CCl₂F (initiations)

  • Note the C–Cl bond is weaker than the C–F bond and breaks more easily to give the very reactive chlorine atom free radical. Bond enthalpies/kJmol-1  C-Cl = 338,  C-F = 484

  • Reactions such as (6a/b) introduce extra free radicals into the upper atmosphere that would not normally be there!

    • Not all the chlorine free radicals destroy ozone, there are very low concentrations of methane in the upper atmosphere and they react with chlorine atoms to hydrogen chloride and a methyl radical.

    • (6c)   Cl•  +  CH₄  ===> HCl  +  •CH₃ (propagation)

• The formation of chlorine atom radicals

  •  The formation of chlorine atoms/radicals from CFCs is the root of the problem because they readily react with ozone and change it back to much more stable ordinary oxygen, equation (7).

    • (7)   O₃  +  Cl• ===>  O₂  +  ClO• (propagation)

    • and this is competing with reaction (5) O•  +  • O₃ ===>  2O₂ in removing ozone from the atmosphere (termination)

  • Removal of ozone means less uv light removed in the upper atmosphere, so more potentially harmful uv light hits the Earth's surface.

  • AND the chlorine oxide radical, ClO• reacts with an oxygen atom to regenerate the chlorine atom radical ...

    • (8)   ClO•  +  O ===>  Cl•  +  O₂   (propagation)

  • So, the 'destructive' Cl radical is still around to destroy even more ozone!

  • The two reactions above, (7) and (8) involving chlorine atoms, are examples of chain propagation reactions and make a free radical catalytic cycle of ozone destruction, because the chlorine atoms from CFC's etc. go through the cycle many times acting as a catalyst in the destruction of ozone. 

    • If you add up reactions (7) and (8) you get equation (9) O•  +  •O₃  ===>  2O₂

      • (termination, but ozone lost and very exothermic)

    • Due  to the very minute concentration of radicals, the probability of chain termination step like

      • (10) 2Cl•  ===>  Cl₂    is very low probability of termination.

    • Therefore because of the catalytic cycle just one chlorine atom can destroy hundreds/thousands? of ozone molecules before it joins up with another radical giving a molecule that isn't a reactive free radical.

    • You can construct a simple catalytic cycle for ozone destruction via chlorine atoms/radicals.

    • (i)     Cl•  +  O3  ===>  ClO•  +  O2

      (ii)    ClO•  +  O  ===>  Cl•  +  O2

      (iii)     O•  +  •O3  ===>  2O2   (i) + (ii)

  • Extra note on free radical chemistry for A level chemistry students on ozone depletion

    • There are lots of other free radical reactions going on and some involve other reactive species e.g. the nitrogen(II) oxide molecule NO ('nitric oxide', nitrogen monoxide) which also has unpaired electron and so can act as a free radical.

    • (11)  •NO  +  O₃  ===>  •NO₂  +  O₂  (∆H = -100 kJmol^-1, propagation, followed by 12)

    • (12) •NO₂  +  •O₃  ===>  •NO₃  +  O₂

    • so, both reactions remove ozone.

    • However NO is naturally produced in the atmosphere, but fossil fuel burning does increase the concentration of nitrogen oxides in the atmosphere.

    • You can also have a catalytic cycle with NO involving reactions (11)  and  (13)

    • (13)  •NO₂  +  O  ===>  •NO  +  O₂  (∆H = -192 kJmol^-1, propagation)

    • If you add up reactions  (11) and (13) you get reaction (9) O•  +  •O₃  ===>  2O₂

      • (∆H = -292 kJmol^-1, termination and very exothermic)

      • The general catalytic cycle is shown below and then quoted for a 3rd time when the OH radical from water is the catalyst X to remove ozone from the upper atmosphere.

    • (i)     X•  +  O3  ===>  XO•  +  O2

      (ii)    XO•  +  O  ===>  X•  +  O2

      (iii)     O  +  O3  ===>  2O2   (i) + (ii)

    • By adding up (i) + (ii) to give (iii) describes the catalytic cycle promoted by many active radicals denoted by catalyst X (e.g. Cl (from CFCs), NO (from fossil fuel combustion) and even H₂O (see below) which is ever present naturally.
    • The catalytic cycle for ozone depletion via water:
      • (i) H₂O  == uv photon ==> H•  +  HO•  (initiation, photolysis, homolytic bond fission)
      • (ii) HO•  +  O₃  ===>  HO₂•  +  O₂  (propagation)
      • (iii) HO₂•  +  O₃  ===>  HO•  +  2O₂  (propagation)
      • (iv) 2O₃  ===>  3O₂   {(ii) +  (iii)}
      • (iv) sums up the catalytic cycle by which water removes ozone from the atmosphere.
      • (v) HO•  +  HO₂•  ===>  H₂O  +  O₂  (termination)

  • -

• Does BANNING CFCs have any effect? are there alternatives to CFCs?

  • Obviously the ban on using CFCs is needed, the case for ozone depletion due to CFCs has been made.

  • But why did the problem persist? and why, even now, the ozone layer has not fully recovered?

    • The problem is that CFCs are not very reactive and are quite stable in the lower atmosphere where the CFC molecules don't get hit by the high energy uv photons to give chlorine atoms.

    • BUT, they will still drift up into the stratosphere and contribute to ozone destruction.

    • In other words, its going to take a long time for all the CFC bans to have complete effect because

      • (i) the 'long-life' of CFCs and

      • (ii) so a few CFC molecules, allowing chlorine atom formation, can be responsible for destroying so many more ozone molecules.

  • The conclusive evidence of ozone destruction by CFCs did cause much concern, spreading from the scientific community to the wider public at large, and most importantly, for any action to be taken, politicians also realised something needed to be done.

    • things didn't happen fast as governments, quite rightly, demanded a good body of proven evidence e.g. fully evaluated peer reviewed research papers.

    • Even by the late 1970s Canada, Norway, Sweden and the USA had banned the use of CFCs as aerosol propellants.

    • Once the 'ozone hole' was discovered other countries (including other European countries like the UK) reduced chlorofluorocarbon production and have now banned the use of CFCs completely.

  • Therefore many countries are banning the use of CFCs, but not all despite the fact that scientists predict it will take many years for the depleted ozone layer to return to its 'original' O₃ concentration and alternatives to CFC's are already being marketed.

    • In fact from 1987 an international ban on the use of CFC molecules called the Montreal protocol is in place and organic chemists have been working hard to create new 'ozone-friendly' alternatives to CFCs.

  • BUT at least the ozone layer is recovering thanks to some world-wide co-operation and the work of chemists in developing less environmentally harmful alternatives.

  • Alternatives to CFCs

    • The idea is to use replacement compounds that are less harmful to the ozone layer e.g. in aerosol products.

    • The HCFC and HFC molecules listed below contain strong C–H and C-F bonds.

      • However, they are susceptible to attack by the hydroxyl radical (•OH) are more easily broken down in the lower atmosphere (troposphere) before reaching the ozone layer in the upper atmosphere.

        • The hydroxyl radical abstracts a hydrogen atom to begin the degradation process

        • e.g  CH₂F₂  +  •OH  ===>  H₂O  +  •CHF₂

      •  Any HCFC and HFC molecules that reach the upper atmosphere are more stable than CFCs because the C-H and C-F bonds in them are stronger than the C-Cl bond in CFCs.

      • This is an excellent example of the practical and theoretical use of bond enthalpies.

        • Bond enthalpies: C-F (484 kJ/mol)  >  C-H (412 kJ/mol)  >  C-Cl (338 kJ/mol)

    • Hydrochlorofluorohydrocarbons (HCFCs)

      • e.g. CH₃CFCl₂ named 1,1–dichloro–1–fluoroethane

      • and named dichlorofluoromethane (a HCFC)

      • They are less destructive than CFCs, but still contain chlorine.

    • Hydrofluorocarbons (HFCs)

      • e.g. CH₂FCF₃ named 1,1,1,4–tetrafluoroethane  (HFC 134a) is used as refrigerant gas and solvent.

      •  CH₃CHF₂ named 1,1-difluoroethane  (HFC 152a)

      • and named difluoromethane

      • Pentafluoroethane CF₃CHF₂ is used in fire extinguishing systems.

      • HFCs are similar to CFCs but they don't contain chlorine in the molecule, so even if they reached the stratosphere, they can't give rise to the catalytic cycle of destruction caused by the chlorine atom free radicals.

      • HFCs are now considered to be safe to use.

    • Alkanes

      • e.g. butane CH₃CH₂CH₂CH₃

      • These hydrocarbons do not contain chlorine, but they are very flammable and strongly infrared absorbing 'greenhouse' effect molecules.!

    • However, note that all of these molecules are greenhouse gases and will contribute to global warming!

    • Its sometimes very difficult to win 100% on these complex environmental issues, its great to be an idealist, but a compromise has to be accepted sometimes!

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