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.