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.

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  1. 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.

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  2. Where is the KMnO4 in the chemical equation?

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  3. Why Fe2+ and Fe3+ become 5Fe2+ and 5Fe3+

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