For a reversible reaction A rightleftharpoons | vedclass

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(C) From the given energy profile diagram:$E_a$ = activation energy of forward reaction$E_a^{\prime}$ = activation energy of backward reaction$E_t$ =

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The threshold energy is less than that of activation energy

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(C) From the given energy profile diagram:$E_a$ = activation energy of forward reaction$E_a^{\prime}$ = activation energy of backward reaction$E_t$ = threshold energy$1$. Since the energy level of product $B$ is higher than reactant $A$,the reaction is endothermic.$2$. The activation energy of the forward reaction $(E_a)$ is the difference between threshold energy $(E_t)$ and energy of reactant $(E_R)$.$3$. The activation energy of the backward reaction $(E_a^{\prime})$ is the difference between threshold energy $(E_t)$ and energy of product $(E_p)$.$4$. From the diagram,$E_a > E_a^{\prime}$.$5$. The relationship is $E_a = E_a^{\prime} + \Delta E$,where $\Delta E$ is the heat of reaction.$6$. The threshold energy $(E_t)$ is always greater than or equal to the activation energy ($E_a$ or $E_a^{\prime}$),as it represents the minimum energy required for the reaction to occur. Therefore,the statement that 'threshold energy is less than that of activation energy' is incorrect.

For a reversible reaction $A \rightleftharpoons B$,which one of the following statements is wrong from the given energy profile diagram?

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