When a catalyst is in the same phase or state as the reactant, an intermediate species is formed
Overall reaction of peroxodisulphate ions and iodide ions:
S2O82-(aq) + 2I- à 2SO42-(aq) +I2(aq)
Peroxodisulphate ions, S2O8 ions oxidise iron (II) to iron (III):
S2O82- (aq) + 2Fe2+ (aq) à 2SO42- (aq) + 2Fe3+(aq)
Then Fe3+ oxidises the I- to I2, regenerating the Fe2+ ions so that none are used up in the reaction:
2Fe3+(aq) + 2I-(AQ) à 2Fe2+(aq) + I2(aq)
Irons job here is to give an electron to the peroxodisulphate and then later on take one back from the iodide reactions.
Notice that without the iron catalyst the reaction occurs between two negatively charged ions, which repel, so a high activation energy for the reaction is needed. However involving the catalyst means that in both reactions the reactants are oppositely charged, this explains the lower activation energy and the increase in rate of reaction.
Autocatalysis occurs when one of the products of the reaction is a catalyst for the same reaction. The reaction begins slowly, with no presence of the catalyst, then as the concentration of the catalyst product increases, the reaction speeds up to a catalysed rate.
Example: the oxidation of ethanedioic acid by manganate(VII) ions
2MnO4-(aq) + 16H+(aq) + 5C2O42- (aq) à 2Mn2+ (aq) + 8H2O (l) + 10CO2 (g)
The catalyst is the 2Mn2+ ions, which are not present at the beginning of the reaction. Once some 2Mn2+ has formed it can react with the 2MnO4- ions to form Mn3+ as an intermediate species, which then reacts with C2O42- ions to form Mn2+:
4Mn2+(aq) + MnO4-(aq) + 8H+(aq) à 5Mn3+ (aq) + 4H2O (l)
2Mn3+ (aq) + C2O42- (aq) à 2CO2 (g) + 2Mn2+ (aq)
The reaction can be followed using a colorimeter to measure the concentration of MnO4-, which is purple.