Permanganometry

Permanganometry is one of the techniques used in quantitative analysis in chemistry. It is a redox titration and involves the use of permanganates and is used to measure the amount of analyte present in unknown chemical samples.^[1] It involves two steps, namely the titration of the analyte with potassium permanganate solution and then the standardization of potassium permanganate solution with standard sodium oxalate solution. The titration involves volumetric manipulations to prepare the analyte solutions.^[2]

Depending on how the titration is performed, the permanganate ion can be reduced to Mn^x, where x is +2, +3, +4 and +6. Using permanganometry we can estimate the quantitative presence of Fe⁺², Mn⁺², Fe⁺² and Mn⁺² when they are both present in a mixture, C₂O₄²⁻, NO₂⁻, H₂O₂ etc.

In the most cases permanganometry is performed in a very acidic solution in which the following reaction occurs:^[3]

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn⁺² + 4H₂O

The standard potential of this electrochemical reaction is:^[4]

E^o=+1.51 V

which shows that KMnO₄ (in an acidic medium) is a very strong oxidizing agent. With this method we can oxidize:

• Fe⁺² (E^o_{Fe⁺³/Fe⁺²}=+0.77 V)
• Sn⁺² (E^o_{Sn⁺⁴/Sn⁺²}=+0.2 V)

and even

• Cl⁻ ( E^o_Cl2/Cl⁻=+1.36 V) etc.

In weak acidic medium MnO₄⁻ can not accept 5 electrons to form Mn⁺², this time it accepts only 3 electrons and forms MnO₂(s) by the following electrochemical reaction:

MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O

The standard potential is E^o=+1.69 V.

And if the solution has a concentration c_(NaOH)>1 mol dm⁻³ the following reaction occurs:

MnO₄⁻ + e⁻ → MnO₄²⁻ E^o=+0.56 V.^[5]

References

1. Redox titrations: Permanganometry. In: University Chemistry, Vol. 1. C. Parameshwara Murthy. New Age International, 2008. ISBN 81-224-0742-0. p.632
2. Louis Rosenfeld. Four Centuries of Clinical Chemistry. CRC Press, 1999, p. 130-175.
3. https://books.google.com/books?id=XQIIAQAAIAAJ Volumetric analysis, Vol 2. Izaak Maurits Kolthoff, Heinrich Menzel, Nathaniel Howell Furman. J. Wiley & Sons, inc., 1929. page 297
4. Table of standard reduction potentials. In: Chemistry and chemical reactivity. John C. Kotz, Paul Treichel, John R. Townsend. Cengage Learning, 2008. ISBN 0-495-38703-7. p. 920
5. Louis Rosenfeld. Four Centuries of Clinical Chemistry. CRC Press, 1999, p. 72-75.