Explain the uses of indicators in the analysis of redox reactions.

(N/A) $(i)$ In one situation,the reagent itself is intensely coloured,e.g.,permanganate ion,$MnO_{4}^{-}.$ Here $MnO_{4}^{-}$ acts as the self-indicator. The visible end point in this case is achieved after the last of the reductant (e.g.,$Fe^{2+}$ or $C_{2}O_{4}^{2-}$) is oxidised and the first lasting tinge of pink colour appears at $MnO_{4}^{-}$ concentration as low as $10^{-6} \ mol \ L^{-1}$. This ensures a minimal "overshoot" in colour beyond the equivalence point,the point where the reductant and the oxidant are equal in terms of their mole stoichiometry.
$(ii)$ If there is no dramatic auto-colour change (as with $MnO_{4}^{-}$ titration),there are indicators which are oxidised immediately after the last bit of the reactant is consumed,producing a dramatic colour change. The best example is afforded by $Cr_{2}O_{7}^{2-}$,which is not a self-indicator,but oxidises the indicator substance diphenylamine just after the equivalence point to produce an intense blue colour,thus signalling the end point.
$(iii)$ There is yet another method which is interesting and quite common. Its use is restricted to those reagents which are able to oxidise $I^{-}$ ions,for example: $2Cu_{(aq)}^{2+} + 4I_{(aq)}^{-} \rightarrow Cu_{2}I_{2(s)} + I_{2(aq)}$.
This method relies on the fact that iodine itself gives an intense blue colour with starch and has a very specific reaction with thiosulphate ions $(S_{2}O_{3(aq)}^{2-})$,which is also a redox reaction: $I_{2(aq)} + 2S_{2}O_{3(aq)}^{2-} \rightarrow 2I_{(aq)}^{-} + S_{4}O_{6(aq)}^{2-}$.
$I_{2}$,though insoluble in water,remains in solution containing $KI$ as $KI_{3}$.
On addition of starch after the liberation of iodine from the reaction of $Cu^{2+}$ ions on iodide ions,an intense blue colour appears. This colour disappears as soon as the iodine is consumed by the thiosulphate ions. Thus,the end-point can easily be tracked and the rest is the stoichiometric calculation only.