Quick questions on Catalysts and reaction rate: TCE Chemistry (Tasmania)

The key idea is that a catalyst lowers the activation energy. On an energy profile diagram the catalysed pathway has a lower peak (energy barrier) than the uncatalysed pathway, while the energies of the reactants and products are unchanged. Because the barrier is lower, a greater proportion of colliding particles have enough energy to react.

Catalysts are central to industrial chemistry because they let reactions run at lower temperatures and pressures, saving energy and cost. Iron catalyses the Haber process for ammonia, vanadium(V) oxide catalyses the Contact process for sulfuric acid, and platinum, palladium and rhodium in a vehicle's catalytic converter speed the conversion of toxic exhaust gases (carbon monoxide and nitrogen oxides) into less harmful carbon dioxide and nitrogen. Because a catalyst lets a process reach an acceptable rate at a lower temperature, it can also improve the equilibrium yield indirectly: for an exothermic reaction, a lower temperature gives a higher equilibrium yield, and the catalyst restores the rate that the lower temperature would otherwise cost.

A catalyst can be deactivated by a poison, a substance that binds strongly to the active sites and blocks reactant molecules. Lead poisons the platinum in catalytic converters, which is why leaded petrol had to be phased out. Poisoning explains why catalysts, though not consumed by the reaction itself, may still need periodic replacement in real plants.