How do Catalysts work?
To recap, see Developing Fuels
In a chemical reaction, the bonds holding the reactants together must first be broken before the reaction can begin.
- Breaking bonds requires energy, and the energy needed to start a reaction is referred to as the activation energy
- Catalysts work by Providing an alternative reaction pathway for the breaking and remaking of bonds. The activation energy for this new pathway is often less than the activation energy of the normal pathway.
- So, how do catalysts work? Well, it’s quite simple really! When a homogeneous catalyst is present, one of the reactants (substrate) reacts with the catalyst forming an intermediate product. The intermediate product then reacts with the other reactant to form the final product.
- The activation energies of both these steps is lower than the activation energy without the presence of a catalyst, therefore more molecules will have the energy to react using the catalyst; hence the rate of reaction is increased.
- Take for instance the reaction between ozone and oxygen free radicals to form di-oxygen.
O3(g) + O(g) 2O2(g)
- This reaction has a high activation energy, and so the process occurs at a slow rate.
- However, the presence of CFC’s in the atmosphere catalyses this reaction by forming intermediate products.
Cl.g + O3(g)O2(g) + ClO.(g)
- This reaction has a lower activation energy than the reaction between ozone and oxygen free radicals. The intermediate product (ClO) then goes on to react with the oxygen free radical.
ClO(g) + O(g) Cl.(g)+O2(g)
- The reactants and the products remain the same; however chlorine has been used as a catalyst to form an intermediate product (requiring less activation energy).
- This increases the rate substantially, and the chlorine remains unchanged and is able to go on to catalyse another reaction.
Useful books for revision:
Revise AS Chemistry for Salters (Written by experienced examiners and teachers of Salter's chemistry)
Revise AS Chemistry for Salters (OCR) (Salters Advanced Chemistry)