An analytical chemist determines the concentration of iron(II) ions, Fe2+\text{Fe}^{2+}Fe2+, in a commercial moss-killing supplement.
A 10.00 cm3 sample of the supplement is titrated against acidified potassium dichromate(VI), K2Cr2O7(aq)\text{K}_2\text{Cr}_2\text{O}_7(aq)K2Cr2O7(aq), according to the following ionic equation:
Cr2O72−(aq)+6Fe2+(aq)+14H+(aq)⟶2Cr3+(aq)+6Fe3+(aq)+7H2O(l) \text{Cr}_2\text{O}_7^{2-}(aq) + 6\text{Fe}^{2+}(aq) + 14\text{H}^+(aq) \longrightarrow 2\text{Cr}^{3+}(aq) + 6\text{Fe}^{3+}(aq) + 7\text{H}_2\text{O}(l) Cr2O72−(aq)+6Fe2+(aq)+14H+(aq)⟶2Cr3+(aq)+6Fe3+(aq)+7H2O(l)In the titration, the 10.00 cm3 sample requires exactly 22.50 cm3 of 0.0200 mol dm-3 K2Cr2O7(aq)\text{K}_2\text{Cr}_2\text{O}_7(aq)K2Cr2O7(aq) to reach the end point.
What is the concentration of Fe2+\text{Fe}^{2+}Fe2+ in the supplement, in mol dm−3\text{mol dm}^{-3}mol dm−3?
0.007500.007500.00750
0.04500.04500.0450
0.2250.2250.225
0.2700.2700.270
