The reaction between potassium permanganate (\( \text{KMnO}_4 \)) and ammonium iron(II) sulfate hexahydrate (\( (\text{NH}_4)_2\text{Fe(SO}_4)_2 \cdot 6\text{H}_2\text{O} \)) is a classical redox reaction, often performed in a laboratory setting to demonstrate potassium permanganate's role as an oxidizing agent.
In acidic conditions, \( \text{KMnO}_4 \) oxidizes the ferrous ion (\( \text{Fe}^{2+} \)) from ammonium iron(II) sulfate to ferric ion (\( \text{Fe}^{3+} \)), while itself is reduced from permanganate (\( \text{MnO}_4^- \)) to manganese(II) ions (\( \text{Mn}^{2+} \)). The overall reaction can be represented as follows:
\[ \text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O} \]
Balanced Chemical Equation
In acidic conditions, the balanced equation for this redox reaction is:
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Oxidation half-reaction: \[ 5\text{Fe}^{2+} \rightarrow 5\text{Fe}^{3+} + 5e^- \]
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Reduction half-reaction: \[ \text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} \]
Full Reaction
Combining these half-reactions gives you the full reaction: \[ \text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O} \]
Conditions
- The reaction typically takes place in an acidic medium (usually sulfuric acid, \( \text{H}_2\text{SO}_4 \)).
- The purple color of the permanganate ion disappears as it is reduced to colorless manganese(II) ions.
Observations
- The initial purple color of the KMnO4 solution fades as it is consumed in the reaction.
- The solution will turn yellow/orange due to the formation of \( \text{Fe}^{3+} \) ions, which impart a distinct color.
This reaction showcases the powerful oxidizing nature of potassium permanganate and is frequently used in titrations and analytical chemistry for determining the concentration of iron(II) ions in a solution.