Permanganate titrations: oxalic acid and Mohr's salt

A titration that shows its own end point

In an acid and base titration, we need an indicator. In a redox titration with potassium permanganate, we do not. The permanganate ion, $\text{MnO}_{4}^{-}$, is deep purple. In acid solution it changes to the almost colourless manganese(II) ion, $\text{Mn}^{2+}$. As long as there is something left to oxidise, each drop of purple permanganate loses its colour at once. When everything is oxidised, the next drop gives a faint, permanent pink colour. That is the end point. Permanganate is a self indicator.

The two equations

In acid solution, permanganate takes five electrons: $\text{MnO}_{4}^{-}\ +\ 8\text{H}^{+}\ +\ 5\text{e}^{-}\ \rightarrow \ \text{Mn}^{2+}\ +\ 4\text{H}_{2}\text{O}$. The equations for the two titrations are:

• Oxalic acid. $2\text{MnO}_{4}^{-}\ +\ 5\text{C}_{2}\text{O}_{4}^{2-}\ +\ 16\text{H}^{+}\ \rightarrow \ 2\text{Mn}^{2+}\ +\ 10\text{CO}_{2}\ +\ 8\text{H}_{2}\text{O}$
• Iron(II) in Mohr's salt. $\text{MnO}_{4}^{-}\ +\ 5\text{Fe}^{2+}\ +\ 8\text{H}^{+}\ \rightarrow \ \text{Mn}^{2+}\ +\ 5\text{Fe}^{3+}\ +\ 4\text{H}_{2}\text{O}$

So 2 permanganate ions react with 5 oxalate ions, and 1 permanganate ion with 5 iron(II) ions.

Rules for the permanganate titration

• Permanganate goes in the burette. It is so dark that the bottom of the meniscus cannot be seen. Read the top of the meniscus.
• Acidify with dilute sulphuric acid. Not hydrochloric acid, because permanganate oxidises chloride ions to chlorine and the titre becomes too large. Not nitric acid, because it is an oxidising agent itself. Too little acid gives a brown precipitate of manganese(IV) oxide.
• Warm the oxalic acid to about 333 to 343 K (60 to 70 degrees Celsius). The reaction is slow at room temperature. The first drops lose their colour slowly. Once some $\text{Mn}^{2+}$ forms, it acts as a catalyst, and the colour then disappears at once. This is autocatalysis.
• Do not warm the Mohr's salt. The reaction with iron(II) is fast at room temperature. Warming would only let air oxidise some of the iron(II).
• Stop at the first permanent pink that lasts about half a minute.

Mohr's salt solution is made with a little dilute sulphuric acid. The acid stops iron(II) from hydrolysing and being oxidised by air.

Calculations

The molarity of permanganate. A known mass of oxalic acid dihydrate is made up to 250.0 mL. A 25.0 mL portion is titrated with permanganate.

1. Molarity of the oxalic acid: $c = \dfrac{m}{126 \times 0.2500}$.
2. Amount of oxalate in the flask: $n = c \times 0.0250$.
3. Amount of permanganate: $\dfrac{2}{5} \times n$.
4. Molarity of permanganate: amount divided by the titre in litres.

The factor $\dfrac{2}{5}$ comes from the equation: 2 permanganate ions for 5 oxalate ions.

Set 1.575 g and 25.0 mL. The molarity is 0.0200 mol/L. Now double the titre. The molarity falls to half, because the same amount of oxalate is spread over twice the volume of permanganate.

Percentage purity of Mohr's salt. Use the permanganate whose molarity you know. The amount of iron(II) is five times the amount of permanganate. Each formula unit of Mohr's salt gives one iron(II) ion, so the mass of pure salt is $n \times 392\ \text{g/mol}$.

The water of crystallisation. For $\text{H}_{2}\text{C}_{2}\text{O}_{4}\cdot \text{xH}_{2}\text{O}$, find the molar mass from the titration. Then $x = \dfrac{M - 90}{18}$.

Try it in Python

The program works out the molarity of potassium permanganate from the mass of oxalic acid and the titre.

Common mistakes

• Using hydrochloric acid to acidify. Chloride is oxidised, and the titre is too large.
• Reading the bottom of the meniscus. The solution is too dark. Read the top.
• Not warming the oxalic acid. The first drops lose their colour very slowly.
• Warming the Mohr's salt. It is not needed, and the iron(II) may oxidise.
• Using the wrong ratio. 2 : 5 for oxalate. 1 : 5 for iron(II).
• Going past the end point. A deep pink colour means too much permanganate was added.

Remember these

• Permanganate in the burette, dilute sulphuric acid in the flask, a faint permanent pink at the end.
• $2\text{MnO}_{4}^{-}\ +\ 5\text{C}_{2}\text{O}_{4}^{2-}\ +\ 16\text{H}^{+}\ \rightarrow \ 2\text{Mn}^{2+}\ +\ 10\text{CO}_{2}\ +\ 8\text{H}_{2}\text{O}$
• $\text{MnO}_{4}^{-}\ +\ 5\text{Fe}^{2+}\ +\ 8\text{H}^{+}\ \rightarrow \ \text{Mn}^{2+}\ +\ 5\text{Fe}^{3+}\ +\ 4\text{H}_{2}\text{O}$
• Warm oxalic acid to about 335 K. Do not warm Mohr's salt.
• Mohr's salt has $M = 392\ \text{g/mol}$ and one iron(II) in each formula unit.

References

• NCERT Chemistry Lab Manual for Class 12, the experiments on the permanganate titrations of oxalic acid and Mohr's salt.
• OpenStax Chemistry 2e, Chapter 17, section 17.4 on balancing redox equations and the use of permanganate.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Why is permanganate called a self indicator?

2. Why is dilute sulphuric acid used and not hydrochloric acid?

3. Why do the first drops of permanganate lose their colour slowly when titrating oxalic acid?

4. How many moles of permanganate react with 0.0100 mol of Mohr's salt?