Draw sketch graphs of reaction rate against concentration of the reactant in bold for each of the following reaction:

$\mathbf{NO}\mathbf{\left(}\mathrm{g}\right) + \mathrm{CO}\left(\mathrm{g}\right) + {\mathrm{O}}_2\left(\mathrm{g}\right) \to {\mathrm{NO}}_2\left(\mathrm{g}\right) + {\mathrm{CO}}_2\left(\mathrm{g}\right)$ for which the rate equation is: rate$= \mathrm{k}{\left[\mathrm{NO}\right]}^2$

Draw sketch graphs of reaction rate against concentration of the reactant in bold for each of the following reaction:

$2\mathbf{HI}\left(\mathrm{g}\right) \stackrel{\mathrm{gold}}{\to } {\mathrm{H}}_2\left(\mathrm{g}\right) + {\mathrm{I}}_2\left(\mathrm{g}\right)$ for which the rate equation is: $\mathrm{rate} = \mathrm{k}$

Note: the catalyst influences the order here – the order is not the same as for the uncatalysed reaction

Draw sketch graphs of reaction rate against concentration of the reactant in bold for each of the following reaction:

${\mathbf{\left(}{\mathbf{CH}}_{\mathbf{3}}\mathbf{\right)}}_{\mathbf{3}}\mathbf{CCl} + {\mathrm{OH}}^{–} \to {\left({\mathrm{CH}}_3\right)}_3\mathrm{COH} + {\mathrm{Cl}}^{–}$ for which the rate equation is: $\mathrm{k}\left[{\left({\mathrm{CH}}_3\right)}_3\mathrm{CCl}\right]$

Draw a sketch graph to show how the concentration of the bold reactant changes with time.

$\mathbf{NO}\mathbf{\left(}\mathrm{g}\right) + \mathrm{CO}\left(\mathrm{g}\right) + {\mathrm{O}}_2\left(\mathrm{g}\right) \to {\mathrm{NO}}_2\left(\mathrm{g}\right) + {\mathrm{CO}}_2\left(\mathrm{g}\right)$ for which the rate equation is: rate$= \mathrm{k}{\left[\mathrm{NO}\right]}^2$

Draw a sketch graph to show how the concentration of the bold reactant changes with time.

$2\mathbf{HI}\left(\mathrm{g}\right) \stackrel{\mathrm{gold}}{\to } {\mathrm{H}}_2\left(\mathrm{g}\right) + {\mathrm{I}}_2\left(\mathrm{g}\right)$ for which the rate equation is: $\mathrm{rate} = \mathrm{k}$

Note: the catalyst influences the order here – the order is not the same as for the uncatalysed reaction

Draw a sketch graph to show how the concentration of the bold reactant changes with time.

${\mathbf{\left(}{\mathbf{CH}}_{\mathbf{3}}\mathbf{\right)}}_{\mathbf{3}}\mathbf{CCl} + {\mathrm{OH}}^{–} \to {\left({\mathrm{CH}}_3\right)}_3\mathrm{COH} + {\mathrm{Cl}}^{–}$ for which the rate equation is: $\mathrm{k}\left[{\left({\mathrm{CH}}_3\right)}_3\mathrm{CCl}\right]$

Use the data from experiments 2 and 3 in table to calculate the rate constant for the following reaction.${\mathrm{H}}_2{\mathrm{O}}_2\left(\mathrm{aq}\right) + 2{\mathrm{I}}^{–}\left(\mathrm{aq}\right) + 2{\mathrm{H}}^{+}\left(\mathrm{aq}\right) \to 2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) + {\mathrm{I}}_2\left(\mathrm{aq}\right)$

The rate equation for this reaction is: rate of reaction $= \mathrm{k}\left[{\mathrm{H}}_2{\mathrm{O}}_2\right]\left[{\mathrm{I}}^{–}\right]$

Use the formula ${\mathrm{t}}_{\frac{1}{2}} = \frac{0.693}{\mathrm{k}}$ to calculate a value for the rate constant of a reaction which is first order and has a half-life of $480 \mathrm{s}$.

A first-order reaction has a rate constant of $9.63 \times {10}^{–5} {\mathrm{s}}^{–1}$. Calculate a value for the half-life of this reaction.