Explaining the colour change in the potassium permanganate titration of iron(II) ions In the redox titration:

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

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?
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.