Define the term rate of reaction.
Identify one change to the reaction mixture that would make it possible to investigate the rate of this reaction.
Use the data and the rate law $\text{rate} = k[\text{NO}]^2[\text{H}_2]$ to find the value of the rate constant $k$. State the units of $k$.
A second experiment is carried out at the same temperature. The starting concentration of $\text{H}_2\text{(g)}$ is $4.60 \times 10^{-3}\,\text{mol dm}^{-3}$. The initial reaction rate is $2.31 \times 10^{-3}\,\text{mol dm}^{-3}\,\text{s}^{-1}$. Calculate the initial concentration of $\text{NO(g)}$.
State the order of the reaction with respect to $\text{NO(g)}$ and with respect to $\text{H}_2(\text{g})$, together with the overall order of the reaction.
The reaction is thought to occur in three stages: 1 $\ \text{2NO} \rightarrow \text{N}_2\text{O}_2$ 2 $\ \text{N}_2\text{O}_2 + \text{H}_2 \rightarrow \text{N}_2\text{O} + \text{H}_2\text{O}$ 3 $\ \text{N}_2\text{O} + \text{H}_2 \rightarrow \text{N}_2 + \text{H}_2\text{O}$ Deduce which of the three stages is the rate-determining step.
Explain your answer to (i).
A third experiment is carried out under different conditions. A small amount of $\text{H}_2(\text{g})$ with concentration $0.0200\ \text{mol dm}^{-3}$ is mixed with a large excess of $\text{NO(g)}$. Under the conditions used, the concentration of $\text{H}_2(\text{g})$ is found to have a constant half-life of $2.00$ seconds. Define the term half-life.
Using the axes below, draw a graph to show how the concentration of $\text{H}_2(\text{g})$ changes during the first $6$ seconds under the conditions used.
$\text{NO(g)}$ acts as a catalyst in the oxidation of atmospheric sulfur dioxide. Give two equations to show how $\text{NO(g)}$ acts as a catalyst in this process.
Explain why $\text{NO(g)}$ may be described as a catalyst in this reaction.
Describe, with the aid of an equation, one environmental consequence of the oxidation of atmospheric sulfur dioxide.