"Chemists poke holes in ozone theory: Reaction data of crucial chloride compounds called into question." - "As the world marks 20 years since the introduction of the Montreal Protocol to protect the ozone layer, Nature has learned of experimental data that threaten to shatter established theories of ozone chemistry. If the data are right, scientists will have to rethink their understanding of how ozone holes are formed and how that relates to climate change.

Long-lived chloride compounds from anthropogenic emissions of chlorofluorocarbons (CFCs) are the main cause of worrying seasonal ozone losses in both hemispheres. In 1985, researchers discovered a hole in the ozone layer above the Antarctic, after atmospheric chloride levels built up. The Montreal Protocol, agreed in 1987 and ratified two years later, stopped the production and consumption of most ozone-destroying chemicals. But many will linger on in the atmosphere for decades to come. How and on what timescales they will break down depend on the molecules' ultraviolet absorption spectrum (the wavelength of light a molecule can absorb), as the energy for the process comes from sunlight. Molecules break down and react at different speeds according to the wavelength available and the temperature, both of which are factored into the protocol.

So Markus Rex, an atmosphere scientist at the Alfred Wegener Institute of Polar and Marine Research in Potsdam, Germany, did a double-take when he saw new data for the break-down rate of a crucial molecule, dichlorine peroxide (Cl₂O₂). The rate of photolysis (light-activated splitting) of this molecule reported by chemists at NASA's Jet Propulsion Laboratory in Pasadena, California¹, was extremely low in the wavelengths available in the stratosphere - almost an order of magnitude lower than the currently accepted rate.

"This must have far-reaching consequences," Rex says. "If the measurements are correct we can basically no longer say we understand how ozone holes come into being." What effect the results have on projections of the speed or extent of ozone depletion remains unclear.

The rapid photolysis of Cl₂O₂ is a key reaction in the chemical model of ozone destruction developed 20 years ago² (see graphic). If the rate is substantially lower than previously thought, then it would not be possible to create enough aggressive chlorine radicals to explain the observed ozone losses at high latitudes, says Rex. The extent of the discrepancy became apparent only when he incorporated the new photolysis rate into a chemical model of ozone depletion. The result was a shock: at least 60% of ozone destruction at the poles seems to be due to an unknown mechanism, Rex told a meeting of stratosphere researchers in Bremen, Germany, last week.

Other groups have yet to confirm the new photolysis rate, but the conundrum is already causing much debate and uncertainty in the ozone research community. "Our understanding of chloride chemistry has really been blown apart," says John Crowley, an ozone researcher at the Max Planck Institute of Chemistry in Mainz, Germany.

"Until recently everything looked like it fitted nicely," agrees Neil Harris, an atmosphere scientist who heads the European Ozone Research Coordinating Unit at the University of Cambridge, UK. "Now suddenly it's like a plank has been pulled out of a bridge."

The measurements at the Jet Propulsion Laboratory were overseen by Stanley Sander, a chemist who chairs a NASA panel for data evaluation. Every couple of years, the panel recommends chemical kinetics and photochemical data for use in atmosphere studies. Until the revised photolysis rate has been evaluated, which won't be before the end of next year, "modellers must make up their minds about what to do," says Sander. One of the problems with checking the data is that the absorption spectra of chloride compounds are technically challenging to determine. Sander's group used a new technique to synthesize and purify Cl₂O₂. To avoid impurities and exclude secondary reactions, the team trapped the molecule at low temperatures, then slowly warmed it up.

"Reactions in experimental chambers are one thing - the free atmosphere is something else," says Joe Farman, one of the scientists who first quantified the ozone hole over Antarctica³. "There's no doubt that ozone disappears at up to 3% a day - whether or not we completely understand the chemistry." But he adds that insufficient control of substances such as halon 1301, used as a flame suppressor, and HCFC22, a refrigerant, is a bigger threat to the success of the Montreal Protocol than are models that don't match the observed losses.

Hot topic

Meanwhile, atmosphere researchers have started to think about how to reconcile observations of ozone depletion with the new chemical models. Several thermal reactions, or combinations of reactions, could fill the gap. Sander's group has started to study possible candidates one by one - but so far without success.

Rex thinks that a chemical pathway involving a Cl₂O₂ isomer - a molecule with the same atoms but a different structure - might be at play. But even if the basic chemical model of ozone destruction is upheld, the temperature dependency of key reactions in the process could be very different - or even opposite - from thought. This could have dramatic consequences for the understanding of links between climate change and ozone loss, Rex says.

The new measurements raise "intriguing questions", but don't compromise the Montreal Protocol as such, says John Pyle, an atmosphere researcher at the University of Cambridge. "We're starting to see the benefits of the protocol, but we need to keep the pressure on." He says that he finds it "extremely hard to believe" that an unknown mechanism accounts for the bulk of observed ozone losses.

Nothing currently suggests that the role of CFCs must be called into question, Rex stresses. "Overwhelming evidence still suggests that anthropogenic emissions of CFCs and halons are the reason for the ozone loss. But we would be on much firmer ground if we could write down the correct chemical reactions."

Quirin Schiermeier

1. Pope, F. D., Hansen, J. C., Bayes, K. D., Friedl, R. R. & Sander, S. P. J. Phys. Chem. A 111, 4322-4332 (2007).
2. Molina, L. T. & Molina, M. J. J. Phys. Chem. 91, 433-436 (1987).
3. Farman, J. C., Gardiner, B. G. & Shanklin, J. D. Nature 315, 207-210 (1985).

Copyright 2007, Nature