Question

Complete and balance the following equations: (1) \[ 2MnO_4^- + 6OH^- \rightarrow \; ? \]

Hint:

In alkaline medium, MnO₄⁻ reduces to MnO₄²⁻ (permanganate → manganate).

Answer 1

Complete and Balance the Following Equation:
(1) \[ 2 \text{MnO}_4^- + 6 \text{OH}^- \rightarrow \; ? \]

Step 1: Identify the Reaction Type
This is a redox reaction where manganese undergoes a change in oxidation state. The permanganate ion (MnO₄⁻) is reduced to Mn²⁺ or MnO₂, and hydroxide ions (OH⁻) are involved in the reaction.

Step 2: Write the Half-Reactions
- The manganese in MnO₄⁻ is reduced from an oxidation state of +7 to +2 in Mn²⁺ (a common product in basic medium).
- The hydroxide ions are oxidized to produce water and oxygen.

Step 3: Write the Balanced Redox Half-Reactions
- Reduction half-reaction: \[ \text{MnO}_4^- + 8 \text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4 \text{H}_2O \] However, since the reaction occurs in a basic medium, we add OH⁻ ions to both sides to balance the equation in the presence of hydroxide ions. The final reduced product in basic medium is MnO₂. - Oxidation half-reaction: \[ 4 \text{OH}^- \rightarrow 2 \text{H}_2O + O_2 + 4e^- \] Now, balance the two half-reactions by multiplying the oxidation half-reaction by 5 and the reduction half-reaction by 2.

Step 4: Combine the Half-Reactions and Balance
\[ 2 \text{MnO}_4^- + 6 \text{OH}^- \rightarrow 2 \text{MnO}_2 + 3 \text{H}_2O + O_2 \] This is the balanced equation. The reaction occurs in a basic medium, and the manganese is reduced to MnO₂.

Conclusion:
The balanced equation is: \[ 2 \text{MnO}_4^- + 6 \text{OH}^- \rightarrow 2 \text{MnO}_2 + 3 \text{H}_2O + O_2 \]