For experiments 1 and 2, the value of $k$ is $26.4$ under the stated conditions. The reaction rate is measured in $\text{mol dm}^{-3}\text{ s}^{-1}$. State the units of $k$.
In experiment 1, the starting concentrations of $\text{NO}$ and $\text{Cl}_2$ are equal. The initial rate in experiment 1 is $2.57 \times 10^{-6}\ \text{mol dm}^{-3}\text{ s}^{-1}$. Calculate the initial concentration of $\text{NO}$. Show your working.
In experiment 2, the starting concentrations of $\text{NO}$ and $\text{Cl}_2$ are each ten times the starting concentrations used in experiment 1. Calculate the initial rate of the reaction in experiment 2.
The graph of [$\text{Cl}_2$] against time indicates that the reaction has a constant half-life, $t_{1/2}$. Explain this observation.
Under the conditions used in experiment 3, the rate constant is $105.6$. Show that $t_{1/2}$ of [$\text{Cl}_2$] is $6.56 \times 10^{-3}\,\text{s}$ in these conditions.
Calculate the time taken, in $\text{s}$, for [$\text{Cl}_2$] to fall to $1.25 \times 10^{-5}\,\text{mol dm}^{-3}$ in experiment 3.
Sulfur dioxide, $\text{SO}_2$, reacts very slowly with oxygen in the atmosphere to form sulfur trioxide, $\text{SO}_3$. This reaction is much faster when NO is present. Explain the role of NO in this process. Include chemical equations in your answer.