Which of the following is an incorrect statement?

The correct answer is Option (1) → Activation energies of reactions are negative

For reactant molecules to change into products, they must first reach a high-energy activated complex (transition state). The extra energy they need, above their own average energy, is the activation energy:

    \( \displaystyle E_a=E_{\text{activated complex}}-E_{\text{reactants}} \)

• Option (1) Activation energies of reactions are negative — The incorrect statement. The activated complex always lies above the reactants, so E_a is always positive; a negative energy barrier has no meaning.

• Option (2) More molecules in a gas have energies above some threshold value as the temperature is increased — A correct statement. On raising the temperature the energy distribution curve broadens and shifts to higher energies, so a larger fraction of molecules has energy at or above E_a.

• Option (3) Reaction coordinate represents the profile of energy change in a reaction. — A correct statement. A plot of potential energy against the reaction coordinate shows the energy change from reactants through the activated complex to products.

• Option (4) The lower the activation energy, the faster the rate of reaction. — A correct statement. By the Arrhenius equation, k = A e^−E_a/RT; a smaller E_a gives a larger k.

NCERT (2026-27 edition):
Temperature Dependence of the Rate of a Reaction – “… It can be understood clearly using the following simple reaction H₂(g) + I₂(g) → 2HI(g) … The energy required to form this intermediate, called activated complex (C), is known as activation energy (E_a). Fig. 3.7 is obtained by plotting potential energy vs reaction coordinate. Reaction coordinate represents the profile of energy change when reactants change into products.”