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Answer and Explanation:
- In spontaneous reactions, activation energy (Ea) of the forward reaction is less than that for the same reaction running in the reverse direction.
- Firstly, The activation energy <em>is the minimum energy requirement for the reacting molecules to initiate the reaction.</em> <em>It is the potential energy difference between the activated complex and the reactants on the energy profile diagram.</em>
- Spontaneous reactions are mainly exothermic reactions and are represented on the energy profile diagram as shown in <em>the attached image.</em>
- <u><em>The forward reaction looks like: </em></u>Reactants (A) → products (B) + heat
As the spontaneous reaction (exothermic), the energy of the products is lower than that of the reactants.
The difference in energy between reactants and products is the enthalpy change of the reaction (ΔH).
The activation energy (Ea) of the forward direction is the difference between the energy of the the activated complex state and the reactants state.
- <u><em>In the case of the reverse direction:</em></u> B + heat → A
In the reverse direction we need a supply of energy because the reaction is converting from a more stable state (low energy) to a lower stable state (high energy).
The supply in the reverse direction is ΔH, the amount of heat released in the forward direction.
Here, the activation energy of the reverse direction will be <em>(Ea of the forward direction + ΔH)</em>.
The difference between the Ea of the reverse and forward direction is ΔH.
Ea reverse = Ea forward + ΔH.
<em>So, in spontaneous reactions, the value of the Ea of forward direction is lower than the same reaction running in the reverse direction.</em>
