Final answer:
By determining the moles of iron needed and the molar masses of FeSO4 and Fe2(SO4)3, the required mass of each iron salt to prepare two solutions, each containing 50.0 mg/L of iron, can be calculated.
Step-by-step explanation:
Calculating Molar Mass and Required Mass for Iron Salts
To determine the mass of two iron salts needed to prepare two solutions of 500.0 mL each containing 50.0 mg/L of iron (Fe2+ and Fe3+), we first need to look up their molar masses. Assuming the two iron salts are FeSO4 (iron(II) sulfate) and Fe2(SO4)3 (iron(III) sulfate), we can use the molar masses of these salts to perform the calculation.
Firstly, the quantity of iron per liter desired is 50.0 mg, which is equivalent to 0.05 g. Knowing that the molar mass of iron is 55.8 g/mol, we can calculate the moles of iron needed per liter and subsequently for 500.0 mL of each solution.
Next, using the molar masses of FeSO4 and Fe2(SO4)3, along with the desired concentration of iron, we can calculate the required mass of each compound for the 500.0 mL solutions. The molar mass of FeSO4 is approximately 152 g/mol and that of Fe2(SO4)3 is approximately 400 g/mol.
After calculating the moles of Fe needed, we account for the fact that FeSO4 contains one atom of Fe per formula unit, while Fe2(SO4)3 contains two. This step is crucial to ensure that the calculated mass of each iron salt will indeed provide the 50.0 mg/L concentration of iron in each solution.