Chapter 16, Problem 74E

Nitrogen monoxide is a pollutant in the lower atmosphere that irritates the eyes and lungs and leads to the formation of acid rain. Nitrogen monoxide forms naturally in the atmosphere according to the endothermic reaction:

${\text{N}}_{\text{2}}\left(\text{g}\right)+{\text{O}}_{\text{2}}\left(\text{g}\right)\rightleftharpoons \text{2NO}\left(\text{g}\right)$

   $K_{\text{p}}=\text{ 4}.\text{1}\times \text{1}{0}^{-\text{31}}\text{at 298K}$

Use the ideal gas law to calculate the concentrations of nitrogen and oxygen present in air at a pressure of 1.0 atm and a temperature of 298 K. Assume that nitrogen composes 78% of air by volume and that oxygen composes 21% of air. Find the “natural” equilibrium concentration of NO in air in units of molecules/cm³. How would you expect this concentration to change in an automobile engine in which combustion is occurring?