For a reversible reaction A rightleftharpoons | vedclass

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(C) From the energy profile diagram,we have:$E_a = 	ext{activation energy of forward reaction}$$E_a^{\prime} = 	ext{activation energy of backward re

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The threshold energy is less than that of activation energy

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(C) From the energy profile diagram,we have:$E_a = 	ext{activation energy of forward reaction}$$E_a^{\prime} = 	ext{activation energy of backward reaction}$$E_t = 	ext{threshold energy}$$1$. The diagram shows that the peak energy $(E_t)$ is higher than the energy of the reactants $(E_R)$ and products $(E_P)$.$2$. The activation energy for the forward reaction is $E_a = E_t - E_R$.$3$. The activation energy for the backward reaction is $E_a^{\prime} = E_t - E_P$.$4$. Since $E_P > E_R$,it follows that $E_a > E_a^{\prime}$. Thus,option $A$ is correct.$5$. Since the potential energy of the product is greater than that of the reactant $(E_P > E_R)$,the reaction is endothermic. Thus,option $B$ is correct.$6$. The threshold energy $(E_t)$ is the minimum energy required for the reaction to occur,which is always greater than or equal to the activation energy of the forward reaction $(E_t = E_a + E_R)$. Therefore,the statement that threshold energy is less than the activation energy is incorrect. Thus,option $C$ is wrong.$7$. The energy of the forward reaction is $E_a = \Delta H + E_a^{\prime}$,where $\Delta H$ is the heat of reaction. Thus,option $D$ is correct.

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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