Environmental Chemistry II: Atmospheric Chemistry — Particles, Ions and Radicals, Chemical Speciation, the Formation of Particulate Matter, Thermochemical and Photochemical Reactions, Oxygen and Ozone Chemistry, Photochemical Smog, and CO, O₃, PAN, VOCs, POPs and Carcinogens in the Air

Secondary inorganic particles form from gases. Sulphur dioxide is oxidised to sulphuric acid in the gas phase, beginning with SO₂ + OH, and faster still in cloud droplets, where hydrogen peroxide and ozone oxidise dissolved SO₂. Nitrogen dioxide is converted to nitric acid by OH by day (NO₂ + OH → HNO₃) and through N₂O₅ at night. Both acids are neutralised by ammonia, mostly from agriculture, forming ammonium sulphate and ammonium nitrate particles, which is why ammonia control helps reduce PM₂.₅. Secondary organic aerosol forms when volatile organic compounds, from vehicles and solvents and from trees (isoprene and terpenes), are oxidised to products of low volatility that condense. Primary particles, by contrast, are emitted as particles: soot, fly ash, dust and sea spray.

Thermochemical reactions are driven by heat. In flames, nitrogen oxides form in three ways. Thermal NOx comes from atmospheric nitrogen by the Zeldovich mechanism, N₂ + O → NO + N and N + O₂ → NO + O, which needs very high temperatures (significant above about 1,500 °C) and rises steeply with flame temperature and residence time; fuel NOx comes from nitrogen bound in the fuel, important for coal; and prompt NOx forms fast in fuel-rich flame fronts. This is why low-NOx burners, staged combustion and flue-gas recirculation work by lowering peak flame temperature or oxygen. Photochemical reactions are driven by the absorption of a photon. The key tropospheric one is the photolysis of nitrogen dioxide by light below about 420 nm: NO₂ + hν → NO + O, followed by O + O₂ + M → O₃ + M, where M is any third molecule that carries off excess energy; ozone is then removed by NO + O₃ → NO₂ + O₂. Left alone these three reactions make a null cycle, a photostationary state in which ozone does not accumulate.

Sydney Chapman explained stratospheric ozone in 1930 with four reactions. Ultraviolet light below about 242 nm splits oxygen molecules, O₂ + hν → 2O; each atom combines with O₂, O + O₂ + M → O₃ + M, forming ozone; ozone absorbs ultraviolet between about 200 and 320 nm and is split back, O₃ + hν → O₂ + O, which is the reaction that shields the surface and heats the stratosphere; and an oxygen atom destroys ozone, O + O₃ → 2O₂. Because the Chapman cycle alone predicts more ozone than is observed, catalytic cycles must also remove it. In each, a catalyst X is regenerated: X + O₃ → XO + O₂, then XO + O → X + O₂, net O + O₃ → 2O₂. X can be OH, NO or, most destructively, chlorine and bromine atoms released from chlorofluorocarbons and halons, which reach the stratosphere intact because they are inert in the troposphere; a single chlorine atom can destroy on the order of a hundred thousand ozone molecules before it is locked into a reservoir such as HCl or ClONO₂.

The Antarctic ozone hole, reported by Farman, Gardiner and Shanklin in 1985, forms every southern spring. In the polar winter vortex, temperatures below about −78 °C form polar stratospheric clouds, on whose surfaces the reservoirs HCl and ClONO₂ are converted into Cl₂; when sunlight returns in September, Cl₂ is photolysed and chlorine destroys ozone rapidly. Total ozone is measured in Dobson units: one Dobson unit is a layer of pure ozone 0.01 mm thick at standard temperature and pressure, so the global average of about 300 DU would be a layer just 3 mm thick, and the ozone hole is conventionally the area below 220 DU. Stratospheric ozone is protective; tropospheric ozone, formed in smog, is a pollutant and a greenhouse gas. The two are chemically identical, and the difference is where they are.

Photochemical smog forms when nitrogen oxides and volatile organic compounds react in strong sunlight, typically in warm, sunny, still weather under a subsidence inversion, the conditions first studied in Los Angeles in the 1950s by Haagen-Smit. The key step breaks the photostationary null cycle. OH attacks a hydrocarbon to form an organic peroxy radical, RO₂, which converts NO to NO₂ without using up ozone: RO₂ + NO → RO + NO₂. Each NO₂ then photolyses to make more ozone, while the NO that would have destroyed it has been consumed, so ozone accumulates. The products include ozone, aldehydes, nitric acid and peroxyacetyl nitrate (PAN, CH₃C(O)OONO₂), a strong eye irritant and phytotoxin that also carries NOx over long distances and releases it on warming.

The daily cycle on a smoggy day follows the chemistry: NO peaks with the morning traffic, NO₂ peaks an hour or two later as NO is oxidised, and ozone peaks in the early afternoon, after hours of sunlight, often downwind of the city. Whether ozone falls more when NOx or when VOCs are cut depends on the regime: in NOx-rich city centres it is VOC-limited, and cutting NOx alone can briefly raise ozone because less NO is left to remove it; in suburbs and rural areas it is NOx-limited.

Carbon monoxide, from incomplete combustion, binds haemoglobin roughly 200 to 250 times more strongly than oxygen, forming carboxyhaemoglobin and starving tissues of oxygen; the heart and brain are affected first, and exposure in closed rooms with charcoal stoves or running engines can be fatal. Ground-level ozone is a strong oxidant that inflames the airways, reduces lung function and damages leaves, reducing crop yields; PAN adds eye irritation and plant injury. Volatile organic compounds are organic chemicals with high vapour pressure at room temperature, among them benzene, toluene, xylenes and formaldehyde; they matter both as smog precursors and directly, benzene being a cause of leukaemia and formaldehyde of nasopharyngeal cancer.

Persistent organic pollutants are organic chemicals that resist degradation, accumulate in fatty tissue, travel long distances through the air and are toxic. Their long-range transport works by repeated evaporation in warm places and condensation in cold ones, the "grasshopper effect" or global distillation, which is why they are found in Arctic mammals and in the breast milk of Inuit women far from any source. The Stockholm Convention (2001, in force 2004) began with twelve, the "dirty dozen": aldrin, chlordane, DDT, dieldrin, endrin, heptachlor, hexachlorobenzene, mirex, toxaphene, polychlorinated biphenyls, and the dioxins and furans; others such as lindane and endosulfan have been added since.