Chemistry

Chemistry

Problem

Consider the following reversible reaction [JEE 2018] A(g)+B(g)⇌AB(g)\text{A(g)} + \text{B(g)} \rightleftharpoons \text{AB(g)} The activation energy of the backward reaction exceeds that of the forward reaction by 2RT2RT (in J mol−1\text{J mol}^{-1}). If the pre-exponential factor of the forward reaction is 4 times that of the reverse reaction, the absolute value ΔG∘\Delta G^\circ (in J mol−1\text{J mol}^{-1}) for the reaction at 300 K is: (Given: ln⁡(2)=0.7\ln(2) = 0.7, RT=2500 J mol−1RT = 2500\ \text{J mol}^{-1} at 300 K and GG is the Gibbs energy)

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