Reaction profiles: explain that the activation energy is the energy needed for a reaction to occur.

The idea

Activation energy is the minimum energy that colliding particles must possess for a reaction to occur. On a reaction profile, it is the vertical distance from the reactants' energy level to the top of the curve, the peak representing the transition state. Even exothermic reactions need this input because existing bonds must be stretched and broken before new bonds form. For example, hydrogen and oxygen can coexist without reacting until a spark supplies enough energy; the spark is not the fuel but the trigger. A larger activation energy means fewer collisions are successful, so the reaction is slower at a given temperature. Explaining the concept therefore means linking the energy barrier to successful collisions, bond breaking and the rate of reaction, not simply naming the peak on a graph.

What earns the marks

• Activation energy is the minimum energy that particles need for a collision to result in reaction.
• Particles must collide with at least this energy so that bonds in the reactants can begin to break.
• A reaction with a higher activation energy needs more energetic collisions, so fewer collisions are successful and the rate is lower at the same temperature.
• Activation energy is not the overall energy change of the reaction; the overall change is the difference between reactants and products.
• Catalysts lower the activation energy by providing an alternative pathway, which increases the proportion of successful collisions.

Common mistakes and how to put them right

• Confusing activation energy with the overall energy change.

  Put it right: The overall change is measured between reactants and products, while activation energy is measured from reactants to the peak.

• Thinking that exothermic reactions need no energy input.

  Put it right: Bonds must still be broken first, so an activation energy is always required.

• Believing that all collisions cause reaction.

  Put it right: Only collisions with energy greater than or equal to the activation energy are successful.

Exam tips

• Label the activation energy on a reaction profile with a vertical arrow from the reactants' level to the peak, and label the overall energy change separately.
• When explaining rate effects, state that a greater proportion of collisions has energy greater than or equal to the activation energy.
• Use the phrase minimum energy needed for a reaction to occur, and avoid saying energy needed to start the reaction without qualification.