In acid-base systems we come across with a titration method for finding out the strength of one solution against the other using a pH sensitive indicator. Similarly, in redox systems, the titration method can be adopted to determine the strength of a reductant/oxidant using a redox sensitive indicator. The usage of indicators in redox titration is illustrated below:
- In one situation, the reagent itself is intensely coloured, e.g., permanganate ion. Here permanganate ion acts as the self indicator. The visible end point in this case is achieved after the last of the reductant is oxidised and the first lasting tinge of pink colour appears at permanganate ion concentration as low as 10^-6 mol/L. 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.
- If there is no dramatic auto-colour change, 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 dichromate ion, 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.
- There is yet another method which is interesting and quite common. Its use is restricted to those reagents which are able to oxidise Iodine ions, say, for example, Cu(II). This method relies on the facts that iodine itself gives an intense blue colour with starch and has a very specific reaction with thiosulphate ions, which too is a redox reaction.
Diiodine, 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 copper ions on iodine ions, an intense blue colour appears. This colour disappears as soon as the iodine is consumed by thiosulphate ions. Thus, the end-point can easily be tracked and the rest is the stoichiometric calculation only.
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