In the following reaction; $\mathrm{x}\mathrm{A}\to \mathrm{y}\mathrm{B}$

${\log }_{10}\left[-\frac{\mathrm{d}\left[\mathrm{A}\right]}{\mathrm{d}\mathrm{t}}\right]={\log }_{10}\left[-\frac{\mathrm{d}\left[\mathrm{B}\right]}{\mathrm{d}\mathrm{t}}\right]+0.3010$

‘A’ and ‘B’ respectively can be:

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:

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 first order reaction, the time required for completion of $90\%$ reaction is '$\mathrm{x}$' times the half life of the reaction. The value of '$\mathrm{x}$' is

(Given: $\ln 10=2.303$ and $\log 2=0.3010$)