Q.Describe the preparation of potassium permanganate. How does the acidified permanganate solution react with

The key idea is the pyrolusite-based fusion to convert insoluble $MnO_2$ into soluble $K_2MnO_4$, followed by electrolytic oxidation to $KMnO_4$.

Preparation:

1. Fusion: Pyrolusite ($MnO_2$) is fused with KOH and an oxidising agent (KNO$_3$ or air) to form green potassium manganate:

2. Oxidation: The manganate solution is electrolysed. At the anode, $MnO_4^{2-}$ is oxidised to $MnO_4^-$:

The purple $KMnO_4$ crystallises out.

Reactions of acidified $KMnO_4$ (in dilute $H_2SO_4$): The $MnO_4^-$ ion is reduced to $Mn^{2+}$ (colourless) in acidic medium.

1. With iron(II) ions ($Fe^{2+}$): $Fe^{2+}$ is oxidised to $Fe^{3+}$.

   $MnO_4^- + 5Fe^{2+} + 8H^+ \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O$

2. With $SO_2$: $SO_2$ is oxidised to $SO_4^{2-}$ (sulphate ion). …

The Concept: Why This Approach Works

Potassium permanganate ($\text{KMnO}_4$) is one of the most important oxidising agents in inorganic chemistry. Its preparation is a two-step process because manganese in its +7 oxidation state is not directly available from common ores. Manganese dioxide ($\text{MnO}_2$), where Mn is in +4 state, is the starting material. The key insight is that we first convert Mn(IV) to Mn(VI) (manganate) by fusion with an alkali in the presence of an oxidising agent, and then oxidise Mn(VI) further to Mn(VII) (permanganate) — this second step is best done electrolytically because chemical oxidants would contaminate the product.

The reactions with acidified permanganate all follow the same principle: in acidic medium, the permanganate ion ($\text{MnO}_4^-$) is reduced to the colourless $\text{Mn}^{2+}$ ion, gaining 5 electrons. The deep purple colour disappears as the reaction proceeds — a useful visual indicator in titrations.

Step-by-Step Preparation

1. Conversion of MnO₂ to potassium manganate ($\text{K}_2\text{MnO}_4$)

Finely powdered pyrolusite ($\text{MnO}_2$) is fused with potassium hydroxide ($\text{KOH}$) and an oxidising agent like potassium nitrate ($\text{KNO}_3$) or potassium chlorate ($\text{KClO}_3$). The reaction is:

The green-coloured potassium manganate is formed. Air can serve as the oxidising agent if the fusion is done in a current of air, but using $\text{KNO}_3$ or $\text{KClO}_3$ makes the process faster and more reliable.

2. Oxidation of manganate to permanganate

The green $\text{K}_2\text{MnO}_4$ solution is treated to convert $\text{MnO}_4^{2-}$ (manganate ion, Mn in +6) to $\text{MnO}_4^-$ (permanganate ion, Mn in +7). This can be done in two ways:

• Electrolytic oxidation (preferred method): An alkaline solution of $\text{K}_2\text{MnO}_4$ is electrolysed using nickel electrodes. At the anode, the manganate ion loses an electron:

The purple permanganate solution is then concentrated by evaporation and crystallised.

• Chemical oxidation: Passing chlorine gas or ozone through the solution also works:

However, this introduces chloride impurities, so the electrolytic method is preferred for pure crystals.

Reactions of Acidified Permanganate

In acidic medium, the half-reaction is:

This high reduction potential makes it a strong oxidising agent. The purple colour fades as $\text{Mn}^{2+}$ (pale pink, nearly colourless in dilute solutions) forms.

(i) Reaction with Iron(II) ions

Iron(II) ions ($\text{Fe}^{2+}$) are oxidised to iron(III) ions ($\text{Fe}^{3+}$). The half-reaction for iron is:

To balance electrons, multiply the iron half-reaction by 5 and add to the permanganate half-reaction:

Adding:

This is the classic redox titration used to estimate iron in ores and alloys.

(ii) Reaction with SO₂

Sulphur dioxide ($\text{SO}_2$) is oxidised to sulphate ions ($\text{SO}_4^{2-}$). In aqueous acidic medium, $\text{SO}_2$ exists as sulphurous acid ($\text{H}_2\text{SO}_3$), and the oxidation half-reaction is:

To balance electrons, multiply the permanganate half-reaction by 2 and the SO₂ half-reaction by 5:

Adding and cancelling water and protons:

$2\text{MnO}_4^- + 5\text{SO}_2 + 2\text{H}_2\text{O} \rightarrow 2\text{Mn}^{2+} + 5\text{SO}_4^{2-} + 4\text{H}^+$ …

Method: Pyrolusite → Alkaline Fusion → Electrolytic Oxidation

This is the industrial preparation method for potassium permanganate ($\ce{KMnO4}$) from the mineral pyrolusite ($\ce{MnO2}$).

Steps of Preparation

1. Alkaline fusion Pyrolusite is fused with potassium hydroxide ($\ce{KOH}$) and an oxidising agent (air or $\ce{KNO3}$) to form potassium manganate ($\ce{K2MnO4}$).

2. Leaching

   The fused mass is dissolved in water, giving a green solution of $\ce{K2MnO4}$.

3. Electrolytic oxidation

   The manganate solution is electrolysed. At the anode, manganate ions are oxidised to permanganate:

The solution turns purple, and $\ce{KMnO4}$ crystallises on concentration.

Reactions of Acidified $\ce{KMnO4}$ (in $\ce{H2SO4}$ medium)

Acidified permanganate is a strong oxidising agent. In acidic medium, $\ce{MnO4-}$ is reduced to $\ce{Mn^{2+}}$ (colourless).

(i) With iron(II) ions ($\ce{Fe^{2+}}$)

• Observation: Purple colour of $\ce{MnO4-}$ disappears; solution turns pale yellow/green (due to $\ce{Fe^{3+}}$).
• Ionic equation:

(ii) With $\ce{SO2}$ (sulphur dioxide) …

Common Mistakes in Potassium Permanganate Preparation & Reactions

Mistake 1: Confusing the Raw Material (Pyrolusite vs. Manganese Dioxide)

The Mistake: Students often write "manganese dioxide" as the starting material but forget that it is the mineral pyrolusite ($MnO_2$). They also confuse the oxidation state of Mn in the ore.

How to Avoid:

• Remember: Pyrolusite = $MnO_2$ (Mn is in +4 state)
• The preparation starts from this ore, not from pure Mn metal or $MnCl_2$
• Write clearly: "Finely powdered pyrolusite ($MnO_2$) is fused with KOH in the presence of air or an oxidising agent like $KNO_3$."

Mistake 2: Forgetting the Fusion Step Conditions

The Mistake: Students skip the exact conditions — temperature, presence of air, or the role of $KNO_3$.

How to Avoid:

• The fusion is done at high temperature (around 500–600°C)
• Air ($O_2$) or $KNO_3$ provides oxygen to oxidise $MnO_2$ to $K_2MnO_4$
• Write the fusion equation:

• If $KNO_3$ is used:

Mistake 3: Confusing the Green Manganate ($K_2MnO_4$) with Permanganate ($KMnO_4$)

The Mistake: Students think the green solution obtained after fusion is already $KMnO_4$.

How to Avoid:

• $K_2MnO_4$ is green (Mn in +6 state)
• $KMnO_4$ is purple (Mn in +7 state)
• The green manganate must be oxidised further to permanganate
• The oxidation is done by:
  • Electrolytic oxidation (commercial method)
  • Or by passing $CO_2$ / $Cl_2$ / $O_3$ through the solution

Electrolytic equation:

Oxidation by $CO_2$:

Mistake 4: Writing Incorrect Ionic Equations for Acidified $KMnO_4$ Reactions

(i) Reaction with Iron(II) ions ($Fe^{2+}$)

Common Mistake: Students forget to balance electrons or write $Fe^{3+}$ incorrectly. They also miss that the medium is acidic ($H^+$ present).

Correct Ionic Equation:

How to Avoid:

• Half-reaction method:
  • Reduction: $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$
  • Oxidation: $Fe^{2+} \rightarrow Fe^{3+} + e^-$
• Multiply oxidation half by 5, then add
• Key check: Electrons cancel (5e⁻ on both sides)

(ii) Reaction with $SO_2$

Common Mistake: Students write $SO_4^{2-}$ as the product but forget to balance oxygen with water and hydrogen with $H^+$.

Correct Ionic Equation:

How to Avoid:

• $SO_2$ is oxidised to $SO_4^{2-}$ (sulphate)
• $MnO_4^-$ is reduced to $Mn^{2+}$
• Half-reactions:
  • Reduction: $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$ (×2)
  • Oxidation: $SO_2 + 2H_2O \rightarrow SO_4^{2-} + 4H^+ + 2e^-$ (×5)
• Add and simplify

(iii) Reaction with Oxalic Acid ($H_2C_2O_4$)

Common Mistake: Students forget that oxalic acid is a diprotic acid and write $C_2O_4^{2-}$ directly without considering the acidic medium.

Correct Ionic Equation:

How to Avoid:

• Oxalic acid ($H_2C_2O_4$) is oxidised to $CO_2$
• Half-reactions:
  • Reduction: $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$ (×2)
  • Oxidation: $H_2C_2O_4 \rightarrow 2CO_2 + 2H^+ + 2e^-$ (×5)
• Important: This reaction is slow at room temperature — it is catalysed by $Mn^{2+}$ (autocatalysis)

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