Question
Explain chemical properties of $KMnO _4$
✓
Answer
→ $KMnO _4$ act as strong oxidising agent in neutral, alkaline and in acidic medium.
(1) In Acidic Solution
(a) Iodine is liberated from potassium iodide:
$\begin{array}{l}2 MnO _4^{-}+16 H ^{+}+10 e ^{-} \rightarrow 2 Mn ^{2+}x+8 H _2 O \\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ 10 I ^{-} \rightarrow 5 I _2+10 e ^{-} \\ \hline 2 MnO _4^{-}+16 H ^{+}+10 I^{-} \rightarrow 2 Mn ^{2+}+5 I _2+8 H _2 O \\\end{array}$
(B) $Fe ^{2+}$ ion (green) is converted to $Fe ^{3+}$ (yellow) :
$
\begin{aligned}
MnO _4^{-}+8 H ^{+}+5 e ^{-} & \rightarrow Mn ^{2+}+4 H _2 O \\
5 Fe ^{2+} & \rightarrow 5 Fe ^{3+}+4 e ^{-} \\ \hline
5 Fe ^{2+}+ MnO _4^{-}+8 H ^{+} & \rightarrow Mn ^{2+}+4 H _2 O +5 Fe ^{3+}
\end{aligned}
$
(c) Oxalate ion or oxalic acid is oxidised at 333 K.
$\begin{array}{c}2 MnO _4^{-}+16 H ^{+}+10 e ^{-} \rightarrow 2 Mn ^{2+}+8 H _2 O \\ 5 C _2 O _4^{2-} \rightarrow 10 CO _2+10 e ^{-} \\ \hline 5 C _2 O _4^{2-}+2 MnO _4^{-}+16 H ^{+} \rightarrow 2 Mn ^{2+}+10 CO _2+8 H _4\end{array}$
(d) Hydrogen sulphide is oxidised, sulphur being precipitated:
$\begin{array}{l} H _2 S \rightarrow 2 H ^{+}+ S ^{2-} \\ 5 S^{2-}+2 MnO _4^{-}+16 H ^{+} \rightarrow 2 Mn ^{2+}+8 H _2 O +5 S\end{array}$
(e) Sulphurous acid or sulphite is oxidised to sulphate or sulphuric acid:
$\begin{array}{r}5 SO _3^{2-}+2 MnO _4^{-}+6 H ^{+} \rightarrow 2 Mn ^{2+}+3 H _2 O
+5 SO_4 ^{2-}\end{array}$
(f) Nitrite is oxidised to nitrate:
$5 NO _2^{-}+2 MnO _4^{-}+6 H ^{+} \rightarrow 2 Mn ^{2+}+5 NO _3^{-}+3 HO _2$
(2) In neutral or faintly alkaline solutions:
(a) A notable reaction is the oxidation of iodide to iodate:
$2 MnO _4^{-}+ H _2 O +\Gamma \rightarrow 2 MnO _2+2 OH ^{-}+ IO _3$
(b) Thiosulphate is oxidised almost quantitatively to sulphate:
$
8 MnO _4^{-}+3 S_2 O _3^{2-}+ H _2 O \rightarrow 8 MnO _2+6 SO _4^{2-}+2 OH^{-}
$
(c) Manganous salt is oxidised to MnO2; the presence of zinc sulphate or zinc oxide catalyses the oxidation :
$2 MnO _4^{-}+3 Mn ^{2+}+2 H _2 O \rightarrow 5 MnO _2+4 H ^{+}$
(1) In Acidic Solution
(a) Iodine is liberated from potassium iodide:
$\begin{array}{l}2 MnO _4^{-}+16 H ^{+}+10 e ^{-} \rightarrow 2 Mn ^{2+}x+8 H _2 O \\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ 10 I ^{-} \rightarrow 5 I _2+10 e ^{-} \\ \hline 2 MnO _4^{-}+16 H ^{+}+10 I^{-} \rightarrow 2 Mn ^{2+}+5 I _2+8 H _2 O \\\end{array}$
(B) $Fe ^{2+}$ ion (green) is converted to $Fe ^{3+}$ (yellow) :
$
\begin{aligned}
MnO _4^{-}+8 H ^{+}+5 e ^{-} & \rightarrow Mn ^{2+}+4 H _2 O \\
5 Fe ^{2+} & \rightarrow 5 Fe ^{3+}+4 e ^{-} \\ \hline
5 Fe ^{2+}+ MnO _4^{-}+8 H ^{+} & \rightarrow Mn ^{2+}+4 H _2 O +5 Fe ^{3+}
\end{aligned}
$
(c) Oxalate ion or oxalic acid is oxidised at 333 K.
$\begin{array}{c}2 MnO _4^{-}+16 H ^{+}+10 e ^{-} \rightarrow 2 Mn ^{2+}+8 H _2 O \\ 5 C _2 O _4^{2-} \rightarrow 10 CO _2+10 e ^{-} \\ \hline 5 C _2 O _4^{2-}+2 MnO _4^{-}+16 H ^{+} \rightarrow 2 Mn ^{2+}+10 CO _2+8 H _4\end{array}$
(d) Hydrogen sulphide is oxidised, sulphur being precipitated:
$\begin{array}{l} H _2 S \rightarrow 2 H ^{+}+ S ^{2-} \\ 5 S^{2-}+2 MnO _4^{-}+16 H ^{+} \rightarrow 2 Mn ^{2+}+8 H _2 O +5 S\end{array}$
(e) Sulphurous acid or sulphite is oxidised to sulphate or sulphuric acid:
$\begin{array}{r}5 SO _3^{2-}+2 MnO _4^{-}+6 H ^{+} \rightarrow 2 Mn ^{2+}+3 H _2 O
+5 SO_4 ^{2-}\end{array}$
(f) Nitrite is oxidised to nitrate:
$5 NO _2^{-}+2 MnO _4^{-}+6 H ^{+} \rightarrow 2 Mn ^{2+}+5 NO _3^{-}+3 HO _2$
(2) In neutral or faintly alkaline solutions:
(a) A notable reaction is the oxidation of iodide to iodate:
$2 MnO _4^{-}+ H _2 O +\Gamma \rightarrow 2 MnO _2+2 OH ^{-}+ IO _3$
(b) Thiosulphate is oxidised almost quantitatively to sulphate:
$
8 MnO _4^{-}+3 S_2 O _3^{2-}+ H _2 O \rightarrow 8 MnO _2+6 SO _4^{2-}+2 OH^{-}
$
(c) Manganous salt is oxidised to MnO2; the presence of zinc sulphate or zinc oxide catalyses the oxidation :
$2 MnO _4^{-}+3 Mn ^{2+}+2 H _2 O \rightarrow 5 MnO _2+4 H ^{+}$
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