Ozone: the Shield We Almost Lost and How We Saved It

Two ozones, two opposite stories

Ozone (O₃) leads a perfect environmental double life. In the stratosphere (15–35 km up) it is the shield that protects us from solar ultraviolet radiation; in the troposphere (the lower layer of the atmosphere where we live) it is a photochemical pollutant that damages lungs and plants.

Tropospheric ozone forms when NOₓ and volatile organic compounds react in sunlight; it is the main component of photochemical smog. Stratospheric ozone forms when UV radiation splits O₂, and it is destroyed naturally by catalytic cycles that keep it in balance. The problem of recent decades is that we introduced artificial catalysts, the CFCs, which speed up stratospheric destruction far beyond that natural balance.

The structure of ozone: a resonant shield

Ozone (O₃) is a bent molecule of three oxygen atoms with a bond angle of 116.8°. Its electronic structure is best described with an sp²-hybridised central oxygen and resonance structures that spread the charge:

• The central O carries a formal charge of +1; the terminal O atoms share a charge of −1 (−½ each on average)
• Both O–O bonds have a bond order of 1.5 (between single and double)
• The central oxygen holds one non-bonding lone pair

This resonance has important spectroscopic consequences. Ozone absorbs UV radiation between 200 and 310 nm (peaking around 260 nm), which covers the UV-B range (280–315 nm) that damages DNA and causes melanoma. The absorption happens because a UV-B photon has enough energy to excite ozone's π system and break one of the O–O bonds, giving back O₂.

The chemistry of the ozone layer: the Chapman cycle

Stratospheric ozone is created and destroyed naturally by the Chapman cycle (1930):

Formation:

O₂ + hν (UV-C, λ < 240 nm) → 2 O· O· + O₂ + M → O₃ + M (M = a third molecule that carries off the excess energy)

Natural destruction:

O₃ + hν (UV-B, λ = 200–310 nm) → O₂ + O· O· + O₃ → 2 O₂

The Chapman cycle produces an equilibrium ozone layer at about 25 km altitude. However, the destruction observed in the real atmosphere is faster than Chapman predicts, because natural catalytic cycles speed it up:

The NO cycle:

NO + O₃ → NO₂ + O₂ NO₂ + O· → NO + O₂ Net: O₃ + O· → 2 O₂ (NO is regenerated, so it is catalytic)

The HO· cycle:

HO· + O₃ → HO₂· + O₂ HO₂· + O· → HO· + O₂ Net: O₃ + O· → 2 O₂ (HO· is regenerated, so it is catalytic)

These natural catalytic cycles are part of the normal balance of the ozone layer. The trouble began when we added another one, the chlorine cycle, fed by CFCs.

CFCs: artificial catalysts of destruction

Chlorofluorocarbons (CFCs), specifically CFC-11 (CCl₃F), CFC-12 (CCl₂F₂) and CFC-113, were developed by Thomas Midgley Jr. in 1928 as non-toxic, non-flammable refrigerants, ideal for air conditioners and refrigerators. For decades they were considered completely safe.

The problem is that they are extremely stable in the troposphere (lifetimes of 45–100 years), so they rise intact into the stratosphere. There, UV-C radiation (which never reaches the troposphere because ozone higher up absorbs it) breaks the C–Cl bond and releases chlorine radicals:

CCl₂F₂ + hν → ·CClF₂ + Cl·

The chlorine radical attacks ozone catalytically:

Cl· + O₃ → ClO· + O₂ ClO· + O· → Cl· + O₂ Net: O₃ + O· → 2 O₂ (Cl· is regenerated, so it is catalytic)

Each chlorine atom can destroy up to 100,000 ozone molecules before it is deactivated (by reacting with methane or NO₂ to form HCl or ClONO₂, the temporary chlorine "reservoirs"). The scale of this catalytic destruction is why small amounts of CFCs have such a disproportionate effect.

The Antarctic hole: the chemistry of polar clouds

The Antarctic ozone hole, discovered in 1985 by Joe Farman and colleagues at the British Antarctic Survey, is more severe than anywhere else for a specific chemical reason: polar stratospheric clouds (PSCs).

During the Antarctic winter (May–August), stratospheric temperatures drop below −78 °C, cold enough for nitric acid and water to freeze into PSC particles. These particles catalyse heterogeneous reactions that turn the inert chlorine reservoirs (HCl, ClONO₂) into reactive forms (Cl₂, HOCl). When the Antarctic spring arrives (September–October) and sunlight returns to the stratosphere, photolysis of Cl₂ and HOCl releases Cl· radicals en masse, producing the accelerated ozone loss we see as the "hole".

The ozone hole is not a literal hole but a region with less than 220 DU of total ozone (the Dobson unit: 1 DU = a 0.01 mm thick layer of ozone at standard pressure). In the worst years (1998–2006) it covered more than 28 million km², twice the area of Antarctica.

The Montreal Protocol: the treaty that worked

The Montreal Protocol, signed in 1987, set out to phase out the production of CFCs and other ozone-depleting substances. In 2009 it became the first UN treaty to achieve universal ratification (today it has 198 parties: all 197 states plus the European Union), and it is also the most successful environmental treaty ever:

• CFC production has fallen by 99.7 % since 1987
• The stratospheric concentration of active chlorine has started to decline
• Models predict full recovery of the ozone layer by 2060–2080

Antarctic expeditions led by Susan Solomon in 1986–87 confirmed the role of chlorine and PSCs in the hole, the scientific basis of the Protocol. Paul Crutzen, Mario Molina and F. Sherwood Rowland received the 1995 Nobel Prize in Chemistry for working out the chemistry of stratospheric chlorine.

The Montreal Protocol is the counter-example that shows coordinated global environmental action is possible, although the ozone problem had the advantage of clear technological substitutes. Those substitutes, the HFCs, are themselves potent greenhouse gases, a problem tackled by the 2016 Kigali Amendment.

The next article comes down from the stratosphere to ground level to examine the other major greenhouse gas: methane, with a warming power about 80 times that of CO₂ over 20 years, and sources ranging from livestock to Arctic permafrost.