Assuming 3.0% hydrogen peroxide by weight has a density of 1.00 g/mL, demonstrate that an excess of hydrogen peroxide was used in the oxidation step of Fe(II) to Fe(III). Think about redox balancing where you used half reactions to add up to get the final, overall reaction. What is the limiting reagent?
The synthetic process will begin with the light green double salt ferrous ammonium sulfate hexahydrate. Addition of oxalic acid to this salt produces insoluble yellow ferrous oxalate by the following reaction: Fe2+ + H2C2O4 + 2H2O → FeC2O4⋅2H2O + 2H+ This precipitate is then oxidized by hydrogen peroxide in the presence of oxalate ion to form the desired bright green complex ion in solution: 2FeC2O4⋅2H2O(s) + H2O2 + 4C2O4 2- → 2[Fe(C2O4)3] 3- + 4H2O + 2OHThe complex is then made to precipitate as a potassium salt by destroying the excess hydroxide and decreasing the polarity of the solvent by adding ethyl alcohol. The precipitation reaction is: 3K+ + [Fe(C2O4)3] 3- + 3H2O → K3[Fe(C2O4)3]⋅3H2O
5g of ferrous ammonium sulfate hexahytdrate is used (Fe 2+) with 15 mL of water, acidified with 6 drops of 3M sulfuric acid. 25 mL of 1m oxalic acid is added. Decant, add 20 mL of water. 5g of potassium oxalate is mixed with 10 mL of water then added to the solution. 20 mL of hydogen peroxide is then added.