One mole of an equimolar mixture of ferric oxalate, , and ferrous oxalate, , is titrated with in an — Redox Reactions and Volumetric Analysis Chemistry Question
Question
One mole of an equimolar mixture of ferric oxalate, $Fe_2(C_2O_4)_3$ , and ferrous oxalate, $FeC_2O_4$ , is titrated with $KMnO_4$ in an acidic medium. How many moles of $KMnO_4$ are strictly required for the complete oxidation of this entire mixture?
💡 Solution & Explanation
One mole of an equimolar mixture contains $0.5$ moles of $Fe_2(C_2O_4)_3$ and $0.5$ moles of $FeC_2O_4$ . For $Fe_2(C_2O_4)_3$ , the $Fe^{3+}$ cannot be oxidized further; only the 3 oxalate ions oxidize to $CO_2$ (total $n=6$ ). Equivalents = $0.5 \times 6 = 3.0$ . For $FeC_2O_4$ , $Fe^{2+} \to Fe^{3+}$ ( $n=1$ ) and $C_2O_4^{2-} \to 2CO_2$ ( $n=2$ ), giving total $n=3$ . Equivalents = $0.5 \times 3 = 1.5$ . Total equivalents of mixture = $3.0 + 1.5 = 4.5$ . $KMnO_4$ in acid has an n-factor of 5. Moles of $KMnO_4$ = $4.5 / 5 = 0.9$ moles.