Explaining the colour change in the potassium permanganate titration of iron(II) ions [duplicate]








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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?




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.


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