Sunlight cause photochemical decomposition of $\mathrm{NO}_2$ into NO and O.

$$\mathrm{NO}_2 \xrightarrow{h v} \mathrm{NO}+[\mathrm{O}]$$

Atomic oxygen is a highly reactive species. It combines with diatomic oxygen and forms ozone.

$$\mathrm{O}_2+\mathrm{O}+\mathrm{M} \longrightarrow \mathrm{O}_3+M$$

where, $M$ is inert gas such as nitrogen. This, $\mathrm{O}_3$ is formed during the formation of smog.

Ozone in stratosphere is a product of action of UV-radiations on dioxygen $\left(\mathrm{O}_2\right)$ molecules. The UV radiation split apart molecular oxygen into free oxygen atoms. These oxygen atoms combine with the molecular oxygen to form ozone.

$\mathrm{O}_2(\mathrm{~g}) \xrightarrow[\text { UV-radiation }]{h v} \mathrm{O}+\mathrm{O}$

$\mathrm{O}_2(g)+\mathrm{O}(g) \underset{\text { radiations }}{\stackrel{\mathrm{UV}}{\rightleftharpoons}} \mathrm{O}_3$

In stratosphere, the formation of $\mathrm{O}_3$ goes on continuously but $\mathrm{O}_3$ is also decomposed by UV- radiation between 240 nm to 360 nm .

$\mathrm{O}_3+\mathrm{H}_2 \xrightarrow{(240-360 \mathrm{~nm})} \mathrm{O}_2+\mathrm{O}$

The O -atom reacts will sand $\mathrm{O}_3$ molecule

$\begin{aligned} \mathrm{O}_3+\mathrm{O} & \longrightarrow 2 \mathrm{O}_2 \\ \text{Net reaction}\quad 2 \mathrm{O}_3 & \longrightarrow 3 \mathrm{O}_2\end{aligned}$

Thus, the reaction form a delicate balance in which the rate of $\mathrm{O}_3$ decomposition match the rate of $\mathrm{O}_3$ formation is a dynamic equilibrium exists and maintains a constant concentration of $\mathrm{O}_3$.

In summer season, nitrogen dioxide and methane react with chlorine monoxide and chlorine atoms forming chlorine sinks, preventing much ozone depletion, whereas in winter, special type of clouds called polar stratospheric clouds are formed over Antarctica. These polar stratospheric clouds provide surface on which chlorine nitrate gets hydrolysed to form hypochlorous acid. It also reacts with hydrogen chloride to give molecular chlorine.

$\mathrm{ClO}^{\infty}(g)+\mathrm{NO}_2(g) \longrightarrow \underset{\text { Chlorine nitrate }}{\mathrm{ClONO}_2(g)}$

$\mathrm{Cl}^{\infty}(\mathrm{g})+\mathrm{CH}_4(\mathrm{~g}) \longrightarrow{ }^{\infty} \mathrm{CH}_3(\mathrm{~g})+\mathrm{HCl}(\mathrm{g})$

$\mathrm{ClONO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{g}) \xrightarrow{\text { Hydrolysis }} \mathrm{HOCl}(\mathrm{g})+\mathrm{HNO}_3(\mathrm{~g})$

$\mathrm{ClONO}_2(g)+\mathrm{HCl}(g) \longrightarrow \mathrm{Cl}_2(g)+\mathrm{HNO}_3(g)$

When sunlight returns to the Antarctica in the spring, the sun's warmth breaks up the clouds and $\mathrm{HOCl}, \mathrm{Cl}_2$ are photolysed by sunlight.

$\mathrm{HOCl}(g) \xrightarrow{h v} \mathrm{O}^{\infty} \mathrm{H}(g)+\mathrm{Cl}^{\infty}(g)$

$\mathrm{Cl}_2(\mathrm{~g}) \xrightarrow{h v} 2 \mathrm{Cl}^{\infty}(\mathrm{g})$

The chlorine radicals thus formed, initiate the chain reaction for ozone depletion.

The laxative effect is observed only when the sulphates present in water have concentration greater than 500 ppm . Otherwise at moderate levels it is charmless.