Ozone depletion — AQA A-Level Chemistry

Chlorofluoroalkanes were chosen as refrigerants and solvents because they are unreactive, non-flammable and non-toxic. That stability is exactly the problem: nothing destroys them at ground level, so over years they drift up into the stratosphere. There the ultraviolet radiation is energetic enough to break a bond homolytically, one electron going to each atom. It is the C–Cl bond that breaks, not C–F, because C–Cl has the lower bond enthalpy; if a question supplies the two values, quote both, but it is the comparison that carries the mark. For dichlorodifluoromethane the step is CCl₂F₂ → •CClF₂ + Cl•. The species released is a chlorine atom with an unpaired electron, written Cl•, not a chloride ion; the whole catalytic cycle that follows depends on that.

A single chlorine radical destroys ozone over and over. The two steps are Cl• + O₃ → ClO• + O₂ and ClO• + O₃ → 2O₂ + Cl•. Add them and the radicals cancel, leaving 2O₃ → 3O₂. The chlorine radical is regenerated in the second step, so it is a catalyst, not a reactant, and one atom can remove of the order of 10⁵ ozone molecules before it is finally removed from the cycle. ClO• is the intermediate: it is made and used up inside the cycle, so it appears in neither the overall equation nor the catalyst role. Because catalyst and reactants are all gases in the same phase, this is homogeneous catalysis; AQA has set ozone decomposition as the odd one out in a list of heterogeneous processes.

This statement is about how science becomes policy, so the credit is for the process, not the chemistry. A laboratory prediction in the 1970s that chlorine radicals could catalyse ozone loss was only a hypothesis. It gained weight when independent teams measured the effect in the real atmosphere: ground-based measurements over Antarctica showed a seasonal collapse in ozone, satellite data confirmed the same hole, and laboratory work reproduced the radical chemistry. Because separate groups using different methods agreed, and the work was published and peer reviewed, governments treated the conclusion as reliable and signed an international agreement restricting chlorofluorocarbons as refrigerants and solvents. Replication by independent groups, not the authority of any one scientist, is what turned evidence into legislation.

The replacements work by removing the source of the problem: with no carbon–chlorine bond, ultraviolet light cannot release a chlorine radical, so no catalytic cycle can start. Hydrofluorocarbons such as CH₂FCF₃ keep the useful properties of the old refrigerants, low boiling point, low flammability and low toxicity. While they contain C–F, C–H, and C–C bonds, they crucially lack C–Cl bonds; the C–F bonds are too strong to break in the stratosphere. Simple alkanes such as pentane are used in domestic fridges for the same reason, since they contain no C–Cl bonds at all. Neither answer is perfect: hydrofluorocarbons are potent greenhouse gases and hydrocarbons are flammable, so the choice trades one environmental or safety cost against another.

Here you must work with the equations themselves. Write the two steps, Cl• + O₃ → ClO• + O₂ and ClO• + O₃ → 2O₂ + Cl•, then add them. Cl• appears on the left of the first and the right of the second, and ClO• appears on the right of the first and the left of the second, so both cancel, leaving 2O₃ → 3O₂. That cancellation is the proof of catalysis: the chlorine radical is consumed and then regenerated, so it does not appear in the overall equation and a tiny amount keeps working. The species that is made and then used up inside the cycle, ClO•, is the intermediate. The same pattern is used for other radicals, such as nitrogen monoxide, so recognise the shape of the argument rather than memorising one pair of equations.