Environmental Issues: CFCs and Ozone Depletion

Chlorofluorocarbons (CFCs) are a family of haloalkanes that were once celebrated as wonder-chemicals: non-toxic, non-flammable, chemically inert, cheap, and ideal as refrigerants, aerosol propellants and foam-blowing agents. The same chemical inertness that made them so useful, however, meant that when released they survived long enough to reach the stratosphere, where ultraviolet light broke them apart and set in motion a radical chain reaction that destroyed the ozone layer. CFCs are a textbook case of how the chemistry taught in A-Level can have direct, global environmental consequences.

This lesson covers the OCR A-Level Chemistry A (H432) specification point 4.2.2 (f): the role of chlorofluorocarbons (CFCs) in the depletion of the ozone layer; the use of radical chain mechanisms; alternatives such as HCFCs and HFCs.

1. The Ozone Layer — What and Why

Ozone (O₃) is an allotrope of oxygen. In the stratosphere (altitude ~15–35 km) it forms a thin layer that absorbs most of the Sun's harmful UV-B and UV-C radiation:

$O_3 + UV \rightarrow O_2 + O\bullet$ $O_2 + O\bullet + \text{collision partner} \rightarrow O_3$

The net effect is the conversion of UV photon energy into heat, shielding the surface. Without the ozone layer, UV-B would cause massively increased rates of skin cancer, cataracts, and damage to ecosystems (especially phytoplankton at the base of the food chain).

Key Fact: Ozone in the stratosphere is "good" — it shields us from UV. Ozone at ground level is "bad" — it is a pollutant from vehicle exhausts and causes respiratory problems.

1.1 Natural Steady State

Without interference, ozone is continuously formed and destroyed in a natural cycle. Oxygen molecules are split by UV-C to give atomic oxygen, and the atoms combine with O₂ to give O₃. Ozone molecules also absorb UV-B and split to give O₂ + O. The overall concentration reaches a dynamic equilibrium of a few parts per million in the stratosphere.

2. CFCs — Structure and Uses

CFCs are alkanes in which all the hydrogens have been replaced by chlorine and fluorine. Common examples:

Formula	Name	Common Use
CCl₃F	Trichlorofluoromethane (CFC-11)	Foam blowing
CCl₂F₂	Dichlorodifluoromethane (CFC-12)	Refrigerant, aerosol propellant
CClF₃	Chlorotrifluoromethane	Specialty refrigerant
C₂Cl₃F₃	1,1,2-Trichloro-1,2,2-trifluoroethane (CFC-113)	Solvent for cleaning electronics

They were adopted enthusiastically in the mid-20th century because:

Non-toxic (unlike ammonia, the previous refrigerant).
Non-flammable (unlike hydrocarbons or SO₂).
Chemically inert — don't react with food, packaging, or people.
Ideal boiling points — turn into gas and back to liquid with small temperature changes, perfect for refrigeration.
Easy to synthesise and cheap.

Unfortunately, their chemical inertness is their downfall.

3. The Road to the Stratosphere

When a CFC leaks from an old refrigerator or aerosol, it drifts through the troposphere (the lower atmosphere, up to ~15 km) unaffected. Water doesn't dissolve it, OH radicals don't react with it (no C–H bonds to attack), and UV-A is too weak to break its bonds.

Over years, atmospheric mixing carries CFCs up into the stratosphere. There, at altitudes above the ozone layer, high-energy UV-C radiation (wavelengths shorter than 242 nm) finally hits them and breaks a C–Cl bond homolytically — each atom takes one electron — producing a chlorine radical:

$CCl_2F_2 \xrightarrow{UV} \bullet CClF_2 + \bullet Cl$

Notice that C–Cl breaks rather than C–F, because the C–Cl bond enthalpy (338 kJ mol⁻¹) is lower than C–F (484 kJ mol⁻¹).

Key Fact: Photolysis of CFCs releases a chlorine radical (Cl•). This single radical is the catalyst that destroys thousands of ozone molecules.

4. The Radical Chain Mechanism for Ozone Destruction

The chlorine radical is the propagator of a destructive chain reaction. The two key propagation steps are:

Propagation step 1:

$\bullet Cl + O_3 \rightarrow ClO\bullet + O_2$

A chlorine radical abstracts an O atom from ozone, giving a chlorine monoxide radical and molecular oxygen.

Propagation step 2:

$ClO\bullet + O_3 \rightarrow \bullet Cl + 2\,O_2$

(Or more precisely, ClO• reacts with a free O atom: ClO• + O• → •Cl + O₂.)

The chlorine radical is regenerated. It can react with another ozone molecule, and another, and another. One Cl• can destroy ~100,000 ozone molecules before it is finally terminated by combining with another radical or being scavenged by other stratospheric species.

Overall (catalytic) effect: 2 O₃ → 3 O₂, with Cl• acting as a catalyst.

graph TD
  A[CFC e.g. CCl2F2] -->|UV-C stratosphere<br/>homolytic fission of C-Cl| B[Cl radical + CClF2 radical]
  B --> C[Propagation 1:<br/>Cl + O3 -> ClO + O2]
  C --> D[Propagation 2:<br/>ClO + O3 -> Cl + 2 O2]
  D --> C
  D --> E[Net: 2 O3 -> 3 O2<br/>Cl regenerated as catalyst]

4.1 Termination

Eventually the chain is broken by termination steps — two radicals combining to form a non-radical product:

$\bullet Cl + ClO\bullet \rightarrow Cl_2 + O$ $\bullet Cl + \bullet Cl \rightarrow Cl_2$

But because radicals are so dilute in the stratosphere, termination is rare, and each propagation cycle can repeat thousands of times before termination occurs.

Lesson 7: Environmental Issues — CFCs and Ozone Depletion