# How does CFC and NO2 deplete ozone layer? Explain it wih balanced chemical equation?

Jun 16, 2018

Warning! Long Answer. Here's what I get.

#### Explanation:

CFCs and ${\text{NO}}_{2}$ enable oxygen atoms to react with ozone and destroy it.

Sunlight reacts with oxygen an breaks it into oxygen atoms.

$\boldsymbol{\left(1.\right)} \textcolor{w h i t e}{m} \text{O"_2 stackrelcolor(blue)(hnucolor(white)(mm))(→) "O·" + "·O}$

CFCs and ${\text{NO}}_{2}$ provide the catalysts that enable oxygen atoms to react quickly with ozone with ozone and destroy it.

CFCs and ozone depletion

The process is a free radical chain reaction.

In the initiation step (2), a photon of ultraviolet light hits a CFC molecule, say, ${\text{CFCl}}_{3}$.

A $\text{C-Cl}$ bond breaks, forming a $\text{Cl}$ atom.

$\boldsymbol{\left(2.\right)} \textcolor{w h i t e}{m} {\text{Cl-CFCl"_2 stackrelcolor(blue)(hnucolor(white)(mm))(→) "Cl·" + "·CFCl}}_{2}$

In the propagation steps (3 and 4), the chlorine atom reacts with an ozone molecule, breaking it apart and destroying the ozone.

$\boldsymbol{\left(3.\right)} \textcolor{w h i t e}{m} {\text{Cl·" + "O"_3 → "ClO" + "O}}_{2}$

The reaction forms a chlorine monoxide molecule and an oxygen molecule.

Then a free oxygen atom reacts with the chlorine monoxide to form oxygen and a chlorine atom.

$\boldsymbol{\left(4.\right)} \textcolor{w h i t e}{m} {\text{ClO" + "O·" → "Cl·" + "O}}_{2}$

The chlorine atom is free to repeat the process of destroying more ozone molecules.

A single CFC molecule can destroy 100 000 ozone molecules.

If we add Equations 3 and 4, we get the overall equation 5 for the reaction.

bb((3.))color(white)(m)color(red)(cancel(color(black)("Cl·"))) + "O"_3 → color(red)(cancel(color(black)("ClO"))) + "O"_2
bb((4.))color(white)(m)ul(color(red)(cancel(color(black)("ClO"))) + "O·" → color(red)(cancel(color(black)("Cl·"))) + "O"_2)
$\boldsymbol{\left(5.\right)} \textcolor{w h i t e}{m m} {\text{O·" + "O"_3 → 2"O}}_{2}$

${\text{NO}}_{2}$ and ozone depletion

The process is a free radical chain reaction.

In the initiation step (6), a photon of ultraviolet light hits an ${\text{NO}}_{2}$ molecule, breaking an $\text{N-O}$ bond and forming an $\text{·NO}$ molecule and an $\text{O}$ atom.

$\boldsymbol{\left(6.\right)} \textcolor{w h i t e}{m} \text{NO"_2 stackrelcolor(blue)(hnucolor(white)(mm))(→) "·NO" + "O·}$

In the propagation steps (7 and 8), the nitrogen monoxide reacts with an ozone molecule, breaking it apart and destroying the ozone.

$\boldsymbol{\left(7.\right)} \textcolor{w h i t e}{m} {\text{·NO" + "O"_3 → "NO"_2 + "O}}_{2}$

The reaction forms a nitrogen dioxide molecule and an oxygen molecule.

Then a free oxygen atom reacts with the nitrogen dioxide to form oxygen and nitrogen monoxide.

$\boldsymbol{\left(8.\right)} \textcolor{w h i t e}{m} {\text{NO"_2 + "O·" → "·NO" + "O}}_{2}$

The nitrogen monoxide is free to repeat the process of destroying more ozone molecules.

The nitrogen monoxide is a catalyst because it reacts at the beginning and is regenerated at the end.

If we add Equations 7 and 8, we get the overall equation 9 for the reaction.

bb((7.))color(white)(m)color(red)(cancel(color(black)("·NO"))) + "O"_3 → color(red)(cancel(color(black)("·NO"_2))) + "O"_2
bb((8.))color(white)(m)ul(color(red)(cancel(color(black)("·NO"_2))) + "O·" → color(red)(cancel(color(black)("·NO"))) + "O"_2)
$\boldsymbol{\left(9.\right)} \textcolor{w h i t e}{m m l l} {\text{O·" + "O"_3 → 2"O}}_{2}$

Note that the overall reaction is the same whether a CFC or nitrogen dioxide is the culprit.