KMnO4 - Potassium Permanganate - Practice Questions & MCQ

Potassium permanganate is prepared by fusion of MnO₂ with an alkali metal hydroxide and an oxidising agent like KNO₃. This produces the dark green K₂MnO₄ which disproportionates in a neutral or acidic solution to give permanganate.

$2 \mathrm{MnO}_2+4 \mathrm{KOH}+\mathrm{O}_2 \rightarrow 2 \mathrm{~K}_2 \mathrm{MnO}_4+2 \mathrm{H}_2 \mathrm{O}$

$3 \mathrm{MnO}_4^{2-}+4 \mathrm{H}^{+}+\mathrm{O}_2 \rightarrow 2 \mathrm{MnO}_4^{-}+\mathrm{MnO}_2+2 \mathrm{H}_2 \mathrm{O}$

Commercially it is prepared by the alkaline oxidative fusion of MnO₂ followed by the electrolytic oxidation of manganate (Vl).

In the laboratory, a manganese (II) ion salt is oxidised by peroxodisulphate to permanganate.

$2 \mathrm{Mn}^{2+}+5 \mathrm{~S}_2 \mathrm{O}_8^{2-}+8 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{MnO}_4+10 \mathrm{SO}_4^{2-}+16 \mathrm{H}^{+}$

Potassium permanganate forms dark purple (almost black) crystals which are isostructural with those of KClO₄. The salt is not very soluble in water (6.4 g/100 g of water at 293 K), but when heated it decomposes at 513 K.

$2 \mathrm{KMnO}_4 \rightarrow \mathrm{~K}_2 \mathrm{MnO}_4+\mathrm{MnO}_2+\mathrm{O}_2$

It has two physical properties of considerable interest: its intense colour and its diamagnetism along with temperature-dependent weak paramagnetism. These can be explained by the use of molecular orbital theory which is beyond the present scope.
The manganate and permanganate ions are tetrahedral; the π-bonding takes place by overlap of p orbitals of oxygen with d orbitals of manganese. The green manganate is paramagnetic because of one unpaired electron but the permanganate is diamagnetic due to the absence of unpaired electron.

A few important oxidising reactions of KMnO₄ are given below:
(1) In acidic solutions:

• Iodine is liberated from potassium iodide:$10 \mathrm{I}^{-}+2 \mathrm{MnO}_4^{-}+16 \mathrm{H}^{+} \rightarrow 2 \mathrm{Mn}^{2+}+8 \mathrm{H}_2 \mathrm{O}+5 \mathrm{I}_2$
• Fe²⁺ ion (green) is converted to Fe³⁺ (yellow):$5 \mathrm{Fe}^{2+}+\mathrm{MnO}_4^{-}+8 \mathrm{H}^{+} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_2 \mathrm{O}+5 \mathrm{Fe}^{3+}$

(2) In neutral or faintly alkaline solutions:

• A notable reaction is the oxidation of iodide to iodate:$2 \mathrm{MnO}_4^{-}+\mathrm{H}_2 \mathrm{O}+\mathrm{I}^{-} \longrightarrow 2 \mathrm{MnO}_2+2 \mathrm{OH}^{-}+\mathrm{IO}_3^{-}$
• Thiosulphate is oxidised almost quantitatively to sulphate:$8 \mathrm{MnO}_4^{-}+3 \mathrm{~S}_2 \mathrm{O}_3^{2-}+\mathrm{H}_2 \mathrm{O} \longrightarrow 8 \mathrm{MnO}_2+6 \mathrm{SO}_4^{2-}+2 \mathrm{OH}^{-}$