Propose a two-step mechanism for this reaction. step 1 $\rightarrow$ step 2 $\rightarrow$
State which step is the rate-determining step in this mechanism. Justify your answer.
Complete the rate equation for this reaction, and state the overall order of the reaction. rate $=$ overall order of reaction $=$
Apply your rate equation in (i) to find the rate of the reaction when the concentrations of $\text{F}_2$ and $\text{ClO}_2$ are each $2.00 \times 10^{-3}\,\text{mol dm}^{-3}$. rate $=$ $\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots$ $\text{mol dm}^{-3}\,\text{s}^{-1}$
In another set of conditions, with a large excess of $\text{ClO}_2$, the rate equation is shown here. rate $= k_1[\text{F}_2]$ The half-life, $t_{\frac{1}{2}}$, for the concentration of $\text{F}_2$ is $4.00\,\text{s}$ under these conditions. Calculate the numerical value of $k_1$, stating its units. Give your answer to three significant figures.
An experiment is carried out under these conditions with an initial concentration of $\text{F}_2$ of $0.00200\,\text{mol dm}^{-3}$. Draw a graph on the grid in Fig. 1.1 to show how the concentration of $\text{F}_2$ varies over the first $12\,\text{s}$ of the reaction.
Use your graph in Fig. 1.1 to determine the rate of the reaction when the concentration of $\text{F}_2$ is $0.00100\,\text{mol dm}^{-3}$. Show the method you use on the graph. rate $=$ $\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots$ $\text{mol dm}^{-3}\,\text{s}^{-1}$