green manganate is paramagnetic because of one unpaired electron but the permanganate is diamagnetic due to the absence of unpaired electron. Acidified permanganate solution oxidises oxalates to carbon dioxide, iron(II) to iron(III), nitrites to nitrates and iodides to free iodine. The half-reactions of reductants are: COO COO 10CO + 10e Fe + ® Fe + + 5e – 5NO – + 5H O ® 5NO – + 10H + + l0e – 10I – ® 5I + 10e – The full reaction can be written by adding the half-reaction for KMnO to the half-reaction of the reducing agent, balancing wherever necessary. If we represent the reduction of permanganate to manganate, manganese dioxide and manganese(II) salt by half-reactions, MnO – + e – ® MnO – ( E o = + .
V) MnO – + 4H + + 3e – ® MnO + 2H O ( E o = + . V) MnO – + 8H + + 5e – ® Mn + + 4H O ( E o = + . V) We can very well see that the hydrogen ion concentration of the solution plays an important part in influencing the reaction. Although many reactions can be understood by consideration of redox potential, kinetics of the reaction is also an important factor.
Permanganate at [H + ] = should oxidise water but in practice the reaction is extremely slow unless either manganese(ll) ions are present or the temperature is raised. A few important oxidising reactions of KMnO are given below: . In acid solutions: (a) Iodine is liberated from potassium iodide : 10I – + 2MnO – + 16H + ® 2Mn + + 8H O + 5I (b) Fe + ion (green) is converted to Fe + (yellow): 5Fe + + MnO – + 8H + ® Mn + + 4H O + 5Fe (c) Oxalate ion or oxalic acid is oxidised at K: 5C O – + 2MnO –