Permanganometry is a redox titration technique that uses potassium permanganate as a strong oxidizing agent. It is commonly used to determine the concentration of reducing agents such as iron(II) ions or oxalic acid in acidic solution.
Potassium permanganate acts as a strong oxidizing agent in acidic medium. The permanganate ion is reduced from manganese(VII) to manganese(II), while the reducing agent is oxidized.
In acidic solution, the half-reaction for permanganate is:
\[\ce{MnO4^- + 8H^+ + 5e^- -> Mn^{2+} + 4H2O}\]
For example, when iron(II) ions are used as the reducing agent, the oxidation half-reaction is:
\[\ce{Fe^{2+} -> Fe^{3+} + e^-}\]
Combining the half-reactions gives the overall equation:
\[\ce{MnO4^- + 5Fe^{2+} + 8H^+ -> Mn^{2+} + 5Fe^{3+} + 4H2O}\]
Potassium permanganate acts as its own indicator, because it is purple, while the reduced form Mn (II) is colorless. The endpoint is reached when a faint pink color persists for about 30 seconds. In the case of organic compounds, such as oxalic acid, the oxidation might be slow, so the soution needs to be heated up, in order to speed up the reaction. Aditionally, in some cases, you need to wait a few seconds between drops, to allow the reaction to proceed.
The solution remains colorless during most of the titration. At the endpoint, a light pink color appears and persists, indicating excess permanganate ions.
The concentration of the reducing agent solution is determined using the volume of potassium permanganate required to reach the endpoint.
Video courtesy of Xylem Class 12 CBSE (YouTube). Video used for educational purposes under YouTube’s embedding policy.