Question

A sample of 1.55g of iron ore is dissolved in a acid solution in which the iron is converted into Fe2+. The solution formed is then titrated with KMnO4 which oxidises Fe2+ to Fe3+ while the MnO4- ions are reduced to Mn2+ ions. 92.95 mL of 0.020M KMnO4 is required for titration to reach the equivalent point.

a) Write the balanced equation for the titration.

b) Calculate the percentage of iron in the sample.
1. Here is the part of the equation you need to work the problem. You can fill in all the rest of it.
MnO4^- + 5Fe^2+ ==> 5Fe^3+ + Mn^2+

2. mols KMnO4 = M x L = ?
mols Fe = 5 x mols KMnO4
grams Fe = mols Fe x atomic mass Fe = ?
%Fe in sample = (grams Fe/grams sample)*100 = ?

Post yuour work if you get stuck.
Where is the KMnO4 in the chemical equation?
Why Fe2+ and Fe3+ become 5Fe2+ and 5Fe3+
5.6
The answer is 33.49%

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