For experiments 1 and 2, the value of $k$ is $26.4$. The reaction rate is measured in $\text{mol dm}^{-3}\text{ s}^{-1}$. State the units of $k$.
In experiment 1, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are the same. The initial rate for experiment 1 is $2.57 \times 10^{-6}\, \text{mol dm}^{-3}\text{ s}^{-1}$. Calculate the initial concentration of $\text{NO}$. Show all working.
In experiment 2, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are both 10 times larger than the initial concentrations in experiment 1. Calculate the initial rate for experiment 2.
The graph of $[\text{Cl}_2]$ against time shows that the reaction has a constant half-life, $t_{\frac{1}{2}}$. Explain this observation.
Under the conditions used in experiment 3, the value of the rate constant is 105.6. Show that $t_{\frac{1}{2}}$ of $[\text{Cl}_2]$ is $6.56 \times 10^{3}\ \text{s}$ for these conditions.
Calculate the time taken, in s, for $[\text{Cl}_2]$ to drop to $1.25 \times 10^{-5}\ \text{mol dm}^{-3}$ in experiment 3.
Sulfur dioxide, $\text{SO}_2$, reacts extremely 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.