Final answer:
The Fe(II) concentration, write down the precipitation equation, utilize the solubility product constant, and then find the new concentrations by dividing the moles by the total volume. The process may involve specific calculations for different salts or titration procedures.
Step-by-step explanation:
To calculate the concentration of Fe(II) in a stock solution, you need to understand the chemistry behind the dissolution and precipitation processes. First, write down the precipitation equation of Fe(OH)2 and state the solubility product expression:
Fe(OH)2 (s) → Fe2+ (aq) + 2OH− (aq)
Ksp = [Fe2+][OH−]2 = 4.87 × 10−17
Find the new concentrations of Fe2+ and OH− by determining the moles in the original solutions and dividing by the total volume in liters. For calculations involving different iron salts like FeCO3 and FeF2, use the same strategy except with specific Ksp values and stoichiometric coefficients for each equation.
Titrations can also determine Fe2+ concentrations. For instance:
[Fe2+] = (0.100 M) (50.0 mL) − (0.100 M) (10.0 mL) / (50.0 mL + 10.0 mL)
For calculating the ore's iron content in titration experiments, calculate the moles of titrant used and convert them to desired concentration units, as in:
The amount of Fe in a 0.4891-g sample of an ore was determined by titrating with K2Cr2O7. The iron content as %w/w Fe2O3 is reported after calculations that consider the molarity of the titrant and the volume used.