The activation energy of one of the reactions in a biochemical...

$\ln \left(\dfrac{{K}_{2}}{{K}_{1}}\right) = \dfrac{{E}_{a}}{R}\left(\dfrac{1}{{T}_{1}} - \dfrac{1}{{T}_{2}}\right)$

$\ln \left(\dfrac{{K}_{2}}{{K}_{1}}\right) = \dfrac{532611}{8 . 3} \times \left(\dfrac{10}{310 \times 300}\right)$

where ${K}_{2}$ is at $310K \text{ and } {K}_{1}$ is at $300K$

$\ln \left(\dfrac{{K}_{2}}{{K}_{1}}\right) = 6 . 9$

$= 3 \times \ln 10$

$ln\dfrac{{K}_{2}}{{K}_{1}} = ln{10}^{3}$

${K}_{2} = {K}_{1} \times {10}^{3}$

${K}_{1} = {K}_{2} \times {10}^{3}$

So