Statement-I : For an endothermic reaction, the activation energy is less than the change in
enthalpy.
Statement-II: Activation energy of a reaction, at a given temperature can never be greater than
its change in enthalpy
are these statements true? give reason
To determine whether these statements are true, we need to understand the concepts of activation energy and change in enthalpy in relation to endothermic reactions.
Statement-I: For an endothermic reaction, the activation energy is less than the change in enthalpy.
This statement is generally true for endothermic reactions. Activation energy refers to the energy required to initiate a chemical reaction by breaking the bonds of the reactants. Change in enthalpy (ΔH) represents the difference in energy between the products and reactants. In an endothermic reaction, ΔH is positive, indicating that energy is absorbed from the surroundings.
The activation energy required to break the bonds and initiate the reaction needs to be less than the overall energy absorbed during the reaction for it to occur. Therefore, statement-I is accurate.
Statement-II: Activation energy of a reaction, at a given temperature, can never be greater than its change in enthalpy.
This statement is not universally true. The relationship between activation energy and change in enthalpy depends on various factors, including the specific reaction and temperature.
While it is common for activation energy to be less than the change in enthalpy, it is possible for the activation energy to be greater in certain cases. This could occur if there are additional barriers or complex reaction mechanisms involved that require higher energy input to start the reaction.
Therefore, statement-II is not always accurate.
In summary, statement-I is generally true, while statement-II is not universally true.