What is the activation energy for a reaction if its rate doubles when the temperature is raised from 20°C to 35°C?

$\left(\mathrm{R}=8.314{\mathrm{Jmol}}^{-1}{\mathrm{K}}^{-1}\right)$

Given that $k_2=2k_1{T}_1={20}^{\circ }C+273=293K$
$\begin{array}{l}{T}_2={35}^{\circ }\mathrm{C}+273=308\mathrm{K}\\ \log \frac{k_2}{k_1}=\frac{E_a}{2.303R}\left[\frac{1}{T_1}-\frac{1}{T_2}\right]\\ \log 2=\frac{E_a}{2.303\times 8.314}\left[\frac{1}{293}-\frac{1}{308}\right]\end{array}$
$\begin{array}{l}{\mathrm{E}}_{\mathrm{a}}=34676.9\mathrm{J}/\mathrm{mol}\\ =34.68\mathrm{KJ}/\mathrm{mol}=34.7\mathrm{KJ}/\mathrm{mol}\end{array}$