Question

The hydroxyl radical has been referred to as the “chief clean-up agent in the troposphere.” Its concentration is approximately zero at night and becomes as high as$\mathbf{1}\mathbf{\times }{\mathbf{10}}^{\mathbf{7}}$molecules per${\mathbf{\text{cm}}}^{\mathbf{3}}$ in highly polluted air. Write an equation for the reaction of HO with${\mathbf{NO}}_{\mathbf{2}}$ in the atmosphere. How does the oxidation state of nitrogen change in this reaction? What is the ultimate fate of the product of this reaction?

Short answer

When the HO radicle is reacted with ${\text{NO}}_2$in the atmosphere, the chemical equation is as follows.

role="math" localid="1663566006438" ${\mathrm{N}}_2+{\mathrm{O}}_2\to 2\mathrm{NO}$

The oxidation state of nitrogen changes in this reaction as follows.

role="math" localid="1663566047166" $\mathrm{NO}+\frac{1}{2}{\mathrm{O}}_2\to {\mathrm{NO}}_2$

The ultimate fate of the product of this reaction as follows.

$\text{NO}+\stackrel{·}{\text{OH}}\to {\text{HNO}}_3$

Step-by-step solution

Definition of the troposphere

The troposphere and the stratosphere are separated by a boundary called tropopause. The troposphere is an oxidizing medium. The dominant oxidizing species in the troposphere is the hydroxyl radical OH.

Given information

The concentration becomes approximately zero at night and becomes high as $1\times {10}^7$ molecules per${\text{cm}}^3$.

Calculation of equation of reactions

When the HO radicleis reactedwith ${\text{NO}}_2$in the atmosphere, the chemical equation is as follows:

role="math" ${\mathrm{N}}_2+{\mathrm{O}}_2\to 2\mathrm{NO}$

The oxidation state of nitrogen changes in this reaction as follows:

$\mathrm{NO}+\frac{1}{2}{\mathrm{O}}_2\to {\mathrm{NO}}_2$

The ultimate fate of the product of this reaction is as follows:

$\text{NO}+\stackrel{·}{\text{OH}}\to {\text{HNO}}_3$