Nitrogen monoxide, $\mathrm{NO}$, reacts with hydrogen, as shown in reaction 3: $2\mathrm{NO} + 2\mathrm{H_2} \rightarrow \mathrm{N_2} + 2\mathrm{H_2O}$. For reaction 3, the rate equation is $\text{rate} = k[\mathrm{H_2}][\mathrm{NO}]^2$. Complete Table 4.1 by giving the order with respect to $\mathrm{[H_2]}$, the order with respect to $\mathrm{[NO]}$, and the overall order of the reaction.
Predict how the initial rate for reaction 3 changes when the concentration of NO is halved.
Predict how the initial rate for reaction 3 changes when the concentrations of NO and $\mathrm{H_2}$ are both increased three times.
Suggest why reaction 3 is unlikely to proceed by a mechanism involving only a single step.
Suggest equations for the three steps of the reaction mechanism for reaction 3. Each step involves a reaction between two molecules.
Suggest the role of $\mathrm{N_2O}$ in this mechanism. Explain your reasoning.
Iodine, $\mathrm{I_2}$, reacts with thiosulfate ions, $\mathrm{S_2O_3^{2-}}$, as shown in reaction 4: $2\mathrm{S_2O_3^{2-}} + \mathrm{I_2} \rightarrow \mathrm{S_4O_6^{2-}} + 2\mathrm{I^-}$. Reaction 4 is done with a large excess of $\mathrm{I_2}$. Under these conditions, the reaction is first order with respect to $\mathrm{[S_2O_3^{2-}]}$ and zero order with respect to $\mathrm{[I_2]}$. The half-life, $t_{1/2}$, for reaction 4 is 720 s under certain conditions. Calculate the value of the rate constant, $k$, for reaction 4. Include the units of $k$.
The reaction between iodide ions, $\mathrm{I^-}(aq)$, and peroxydisulfate ions, $\mathrm{S_2O_8^{2-}}(aq)$, is catalysed by $\mathrm{Co^{3+}}(aq)$. The mechanism is similar to the mechanism of this reaction when $\mathrm{Fe^{3+}}(aq)$ is used as the catalyst. State the type of catalysis that occurs in this reaction. Explain your reasoning.
Write two equations to show how $\mathrm{Co^{3+}}(aq)$ catalyses this reaction.
Suggest why this reaction is slow in the absence of $\mathrm{Co^{3+}}(aq)$.