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For two reactions : R-I and R-II the rate constants at 300 K and 320 K are 0.1 sec1 and 0.3 sec1 respectively (for R-I) and 0.2 sec1 and 0.8 sec1 (for R-II}. Which of the options regarding activation energy is correct?

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Activation energies for both the reactions is same.

Activation energy for second reaction is $\frac{4}{3}$ times as that of first reaction.

Activation energy for first reaction is more.

Activation energy for second reaction is more.

answer is C.

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Detailed Solution

$\ln (\frac{k_{2}}{k_{1}}) = \frac{E_{a}}{R} [\frac{1}{T_{1}} - \frac{1}{T_{2}}]$

$\frac{\ln (\frac{k_{2}}{k_{1}})}{\ln (\frac{k_{2}^{'}}{k_{1}^{'}})} = \frac{E_{a}}{E_{a}^{'}}$

$\Rightarrow E_{a}^{'} > E_{a}$

So, we can says that Activation energy for second reaction is more.

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