A redox titration can accurately determine the concentration of an analyte by measuring it… 1 answer below »

A redox titration can accurately determine the concentration of an analyte by

measuring it against a standardised titrant. A well-known example is the redox titration

of a standardised solution of potassium permanganate (KMnO4) against an analyte

containing an unknown concentration of iron(II) ions (Fe2+).

The use of KMnO4 as a titrant is particularly useful because it can act as its own

indicator; due to KMnO4 solution being bright purple, while the Fe2+ solution is

colourless. It is therefore possible to see when the titration reaches its endpoint,

because the solution will remain slightly purple from unreacted KMnO4.

During this titration;

• Fe2+ ions are oxidised to Fe3+ ions

• MnO4

– ions are reduced to Mn2+ ions

The KMnO4 solution is placed in the burette (see Figure 1 below). Before the

equivalence point is reached, any KMnO4 added to the conical flask will not remain in

the solution but will be consumed by the other reactant in the flask. Once the

equivalence point has been reached there is no more reactant in the flask to consume

the KMnO4. If one extra drop of KMnO4 is added from the burette it will remain in the

solution. This additional drop will give a purple-pink colour to the solution. The first sign

of this colour is the end-point. It approximates very closely to the equivalence point.

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