May I ask a small question that I doubt would ever come up but has been bugging me all week - why do you get a sudden colour change? I realise it's the equivalence point, but if I add reactant A to reactant B in deficit, why doesn't the deficit amount of reactant A react to form product A - which gives off some colour?
I think I kind of understand what you mean.
Say you have a reaction involving MnO4- and Fe2+. What you have is MnO4- (strong purple) + Fe2+ (essentially colorless) + (other reactants) --> Mn2+ (essentially colorless) + Fe3+ (essentially colorless) + (other products)
What you already know: In the initial reaction vessel, you have MnO4-, the solution is purple. As you add some clear Fe2+, the color fades. Some of the purple is converted to the colorless product. Since Fe2+ is added in 'deficit' for now, the color of the remaining MnO4- persists, the solution is still purple. As the reaction continues, the MnO4- eventually is used up, and the color disappears (the final stage of fading is quite sudden). There is a sharp transition from coloured --> clear
What you are asking about: If, however, the reaction is actually reversed. That is Mn2+ + Fe3+ + (other things) --> MnO4- + Fe2+ + (other things), so the reaction
1 is going from clear to purple. We will get a tinge of purple straight away as we add the reactants, and this color will deepen as the reaction proceeds. It is impossible to tell when we have reach the equivalence point, because this method's end point (point of sharp color change) is at the very beginning of the reaction. Thus this method simply does not work, and will not be used for a titration. Another method must be sought.
Did that answer your question?
1. This reaction won't actually happen spontaneously since it is the back reaction of a spontaneous reaction. It is used in this scenario as a hypothetical case of a clear --> coloured redox addition.
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