Write an equation showing the formation of $\text{NO}_2$ in these situations.
By what method is $\text{NO}_2$ taken out of the exhaust gases from motor vehicles?
Write a balanced equation for this process.
Suggest whether the amount of pollutant $\text{NO}_2$ produced would fall if fossil fuels were replaced with hydrogen as the fuel for combustion. Explain your answer.
In the atmosphere, $\text{NO}_2$ acts as a catalyst in the oxidation of $\text{SO}_2$ to $\text{SO}_3$.\n\n$\text{SO}_2(g) + \tfrac{1}{2}\text{O}_2(g) \xrightarrow{\text{NO}_2} \text{SO}_3(g)$\n\nWhat is the environmental importance of this reaction?
The oxidation happens in two stages. The first reaction is the one between $\text{NO}_2$ and $\text{SO}_2$.\n\nreaction 1:\n\n$\text{NO}_2(g) + \text{SO}_2(g) \rightleftharpoons \text{NO}(g) + \text{SO}_3(g) \qquad \Delta H = -168\,\text{kJ mol}^{-1}$\n\nWrite an equation to show how the $\text{NO}_2$ is regenerated in the second stage of the oxidation.
Write an expression for the equilibrium constant, $K_p$, for reaction 1, and state its units.
If equal amounts of $\text{NO}_2(g)$ and $\text{SO}_2(g)$ are left to react at room temperature, it is found that $99.8\%$ of the gases have changed into products at equilibrium. Calculate a value for $K_p$.
The temperature of the atmosphere falls with height. How will this change the position of the equilibrium in reaction 1? Explain your answer.