Some time ago, Eli discussed the atmospheric oxidation of methane a topic to which he now returns in the context of ozone pollution and the effects on air quality and human health. Tropospheric ozone is a result of NOx chemistry, where NOx, either referred to as NO ex, or nocks, is the sum of NO (nitric oxide) and NO2 (nitrogen dioxide) in the atmosphere.
Tropospheric ozone is formed by the photodissociation of NO2 between 400 and 300 nm, with the long wavelength limit set by the bond strength of O - NO, and the lower one by the absorption cut off of stratospheric ozone which absorbs all of the solar light below 300 nm (to be fussy 306 nm or so and as long as the bunnies are being fussy somewhere like 420 nm on the high end because of thermal population of excited vibrational levels)
NO2 + hν --> NO + O(3P) 
If Reaction 2 is assumed to be fast (it is because there is so much oxygen and nitrogen around) the rate at which ozone is produced is j [NO2] where j is the flux of UV light and the rate at which it is consumed is k3[NO][O3] where k3 is the rate constant for reaction 3[X] means the concentration of X. Under steady state conditions when production via Reaction 1 equals destruction via Reaction 3
which can be solved for an estimate of the ozone concentration
in terms of the steady state concentrations of NO and NO2
While reaction  appears to limit the amount of ozone that can form, other reactions which cycle NO back to NO2 without consuming ozone would lead to much higher ozone concentrations. Volatile organic compounds (VOCs) provide such a path, and methane is the VOC with the highest atmospheric concentration. Moreover, as a non-condensible gas not very susceptible to rain out, it spreads worldwide in ways that heavier VOCs do not. Looking at the reaction scheme for methane oxidation