Activation energy is the minimum energy that is needed to break the bonds in reactant molecules so that a chemical reaction can proceed.
Earlier it was mentioned that it is the collision of particles that causes reactions to occur and that only some of these collisions are ’successful’. This is because the reactant particles have a wide range of kinetic energy, and only a small fraction of the particles will have enough energy to actually break bonds so that a chemical reaction can take place. The minimum energy that is needed for a reaction to take place is called the activation energy.
Even at a ﬁxed temperature, the energy of the particles varies, meaning that only some of them will have enough energy to be part of the chemical reaction, depending on the activation energy for that reaction. Increasing the reaction temperature has the eﬀect of increasing the number of particles with enough energy to take part in the reaction, and so the reaction rate increases.
A catalyst functions slightly diﬀerently. The function of a catalyst is to lower the activation energy so that more particles now have enough energy to react. The catalyst itself is not changed during the reaction, but simply provides an alternative pathway for the reaction, so that it needs less energy. Some metals e.g. platinum, copper and iron can act as catalysts in certain reactions.
In our own human bodies, enzymes are catalysts that help to speed up biological reactions. Catalysts generally react with one or more of the reactants to form a chemical intermediate which then reacts to form the ﬁnal product. The chemical intermediate is sometimes called the activated complex.
Deﬁnition: A catalyst speeds up a chemical reaction, without being altered in any way. It increases the reaction rate by lowering the activation energy for a reaction.
Energy diagrams are useful to illustrate the eﬀect of a catalyst on reaction rates. Catalysts decrease the activation energy required for a reaction to and therefore increase the reaction rate.