# Explaining the colour change in the potassium permanganate titration of iron(II) ions

In the redox titration:

$$\ce{MnO4- + 8 H+ + 5 Fe^2+ -> Mn^2+ + 4 H2O + 5 Fe^3+}$$

the colour change which occurs is purple to colourless, because of the decreased concentration of permanganate ions. But doesn't this fail to account for the visible iron ions? $$\ce{Fe^2+(aq)}$$ is green and $$\ce{Fe^3+(aq)}$$ is brown.

Anyone have an explanation for why the end result is colourless and not brown solution?

• It's permanganate, not manganate. As for your question, consider the intensity of color. Compared to permanganate, both Fe2+ and Fe3+ are almost colorless. – Ivan Neretin Apr 26 '18 at 7:53
• Fe(3+) is very pale yellow, hardly brown – orthocresol May 27 '18 at 0:19
• Furthermore, note the common addition of $\ce{H3PO4}$ whose dissociated anions give colourless complexes with $\ce{Fe^{3+}}$. – Linear Christmas Jun 25 '18 at 20:46

I think this is because $\ce{MnO4-}$ ions have a very distinct colour. At the end of the reaction, the solution is colourless. I think that the $\ce{Fe^3+}$ ions give the solution a brown colour, but it is in a very low concentration. The brown $\ce{Fe^3+}$ ions are too few to give the whole solution a brown colour.