Considering the ${\mathrm{S}}_{\mathrm{N}}2$ reaction of butyl bromide with ${}^{-}\text{OH}$ ion:

${\text{CH}}_3{\text{CH}}_2{\text{CH}}_2{\text{CH}}_2\text{Br}+\overline{\text{O}}\text{H}\to {\text{CH}}_3{\text{CH}}_2{\text{CH}}_2{\text{CH}}_2\text{OH}+{\text{Br}}^{-}$.

Assuming no other change, what effect on the rate would result from simultaneously doubling the concentration of both butyl bromide and $\overline{\text{O}}\text{H}$ ion?

Consider the given plots for a reaction obeying Arrhenius equation $(0°C<T<300°C):$ ($K$and $E_a$ are rate constant and activation energy, respectively )

(I) 

(II) 

For a reaction, consider the plot of $lnk$ versus $1/T$ given in the figure. If the rate constant of this reaction at $400K$ is ${10}^{-5}{s}^{-1}$ , then the rate constant at $500K$ is:

For an elementary chemical reaction, ${\mathrm{A}}_2 \underset{{\mathrm{k}}_{-1}}{\overset{{\mathrm{k}}_1}{\rightleftarrows }}2\mathrm{A}$ , the expression for $\frac{\mathrm{d}\left[\mathrm{A}\right]}{\mathrm{dt}}$ is: