Suggest a mechanism with two steps for this reaction. step 1 \rightarrow step 2 \rightarrow
Identify the rate-determining step for this mechanism. Explain your answer.
Complete the rate equation for this reaction, then state the overall order of the reaction. rate $=$ overall order of reaction $=$
Use your rate equation in (i) to calculate the reaction rate when the concentrations of $\text{F}_2$ and $\text{ClO}_2$ are both $2.00 \times 10^{-3}\,\text{mol dm}^{-3}$. rate $=\;\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\,\text{mol dm}^{-3}\,\text{s}^{-1}$
In a different set of conditions, with a large excess of $\text{ClO}_2$, the rate equation is shown below. rate $= k_1[\text{F}_2]$ The half-life, $t_{\frac{1}{2}}$, of the concentration of $\text{F}_2$ is $4.00\,\text{s}$ under these conditions. Calculate the numerical value of $k_1$, giving its units. Give your answer to three significant figures. $k_1 = \;\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots$ units $\ldots\ldots\ldots\ldots$
An experiment is carried out under these conditions with an initial $\text{F}_2$ concentration of $0.00200\,\text{mol dm}^{-3}$. Draw a graph on the grid in Fig. 1.1 to show how the concentration of $\text{F}_2$ changes during the first $12\,\text{s}$ of the reaction.
Use your graph in Fig. 1.1 to find the rate of the reaction when the concentration of $\text{F}_2$ is $0.00100\,\text{mol dm}^{-3}$. Show your working on the graph. rate $= \;\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\ldots\,\text{mol dm}^{-3}\,\text{s}^{-1}$