The Activation Energy

• The activation energy is the minimum amount of energy required for a chemical reaction to occur.

• It acted as a barrier to the reaction and is represented on energy profile diagrams as the energy hump.

• Without enough energy to surpass this threshold, reactants will not transform into products.

• Reaction rate increases with an increase in activation energy; in other words, higher the activation energy, more is the reaction rate.

• During a chemical reaction, bonds in the reactants need to be broken before new bonds can form to make the products. The energy needed to break these bonds is provided by the activation energy.

• Catalysts are substances that lower the activation energy, making it easier for the reaction to occur. They do so by providing an alternative reaction pathway with a lower activation energy.

• Exothermic reactions release energy into the environment. Here, the energy required to activate the reaction is less than the total energy released.

• Conversely, in an endothermic reaction, energy is absorbed from the environment. These reactions have higher activation energy than the total energy absorbed.

• Remember that the magnitude of activation energy does not indicate the speed of a reaction. A reaction with a high activation energy may still occur quickly if the reactant molecules have enough energy.

• Activation energy can be calculated from the Arrhenius equation. It is used in the calculation of rate constants, which determine the speed of reactions.

• Activation energy is commonly measured in kilojoules per mole (kJ/mol). The lower the activation energy for a reaction, the greater the proportion of colliding particles will have sufficient energy to react, making the reaction likely to occur faster.