volumetric analysis
Objective: Find the percentage of iron in a given sample through titration.
EDTA CONCENTRATION
0.02060 M
Mass of Ferric Salt
1.8193 g
Burette Readings ( in mL)
Trial 1
Trial 2
Trial 3
Final
42.19
21.29
39.69
Initial
23.35
2.61
21.21
Endpoint Volume
18.85
18.6
18.47
Average of the endpoint volumes
V= 18.64 mL
% Fe ± Sv
11.78% ± 0.19313
Calculations: sV= {[(v1 – V)^2 + (v2 – V)^2 + (v3 – V)^2]/2}^(1/2) {[(18.85 – 18.64)^2 + (18.6 – 18.64)^2 + (18.47 – 18.64)^2]/2}^(1/2) Sv = 0.19313
Percentage of Fe(III):
X mol/ (0.01864 L of EDTA) = 0.02060 M Mols of EDTA = 3.834 x 10^-4 mol Mols of Fe = 3.834 x 10^-4 mol (Fe (III) titrates with EDTA at 1:1 ratio)
3.834 x 10^-4 mol x 58.845g/mol = 0.02144 grams of Fe(III)
(0.2144 g Fe(III) / 1.8193 g Ferric salt sample) x 100 = 11.78%
1) The experiment was performed acceptably well, the %RSD was not below 1%, but it only exceeded 1% slightly. % RSD = (0.1913/ 18.64) x 100 = 1.026
2) The indicator is necessary to determine if the solution is fully titrated because it indicates the presence of untitrated Fe(III) ions and if additional EDTA to the solution is necessary. Without it, the solution can easily be “over” or “under” titrated; because the equivalence point is not clear. This would lead to inaccurate readings of EDTA that would cause errors in the data and calculations.
3) If the indicator chosen forms a stronger complex with the Fe (III) than the FeEDTA- complex, the solution will not titrate at all. The weak bond between the Fe(III) and the indicator allows the Fe(III) ions to break away and bond with the EDTA. If it is stronger, the Fe(III) ions will remain bonded with the indicator and the solution will not titrate.
Conclusion: Through finding the amount of EDTA necessary to titrate with the solution and knowing that Fe titrates with EDTA on an 1:1 ratio, the percentage amount of Fe present in the sample was determined to be 11.78% ± 0.19313.
EDTA CONCENTRATION
0.02060 M
Mass of Ferric Salt
1.8193 g
Burette Readings ( in mL)
Trial 1
Trial 2
Trial 3
Final
42.19
21.29
39.69
Initial
23.35
2.61
21.21
Endpoint Volume
18.85
18.6
18.47
Average of the endpoint volumes
V= 18.64 mL
% Fe ± Sv
11.78% ± 0.19313
Calculations: sV= {[(v1 – V)^2 + (v2 – V)^2 + (v3 – V)^2]/2}^(1/2) {[(18.85 – 18.64)^2 + (18.6 – 18.64)^2 + (18.47 – 18.64)^2]/2}^(1/2) Sv = 0.19313
Percentage of Fe(III):
X mol/ (0.01864 L of EDTA) = 0.02060 M Mols of EDTA = 3.834 x 10^-4 mol Mols of Fe = 3.834 x 10^-4 mol (Fe (III) titrates with EDTA at 1:1 ratio)
3.834 x 10^-4 mol x 58.845g/mol = 0.02144 grams of Fe(III)
(0.2144 g Fe(III) / 1.8193 g Ferric salt sample) x 100 = 11.78%
1) The experiment was performed acceptably well, the %RSD was not below 1%, but it only exceeded 1% slightly. % RSD = (0.1913/ 18.64) x 100 = 1.026
2) The indicator is necessary to determine if the solution is fully titrated because it indicates the presence of untitrated Fe(III) ions and if additional EDTA to the solution is necessary. Without it, the solution can easily be “over” or “under” titrated; because the equivalence point is not clear. This would lead to inaccurate readings of EDTA that would cause errors in the data and calculations.
3) If the indicator chosen forms a stronger complex with the Fe (III) than the FeEDTA- complex, the solution will not titrate at all. The weak bond between the Fe(III) and the indicator allows the Fe(III) ions to break away and bond with the EDTA. If it is stronger, the Fe(III) ions will remain bonded with the indicator and the solution will not titrate.
Conclusion: Through finding the amount of EDTA necessary to titrate with the solution and knowing that Fe titrates with EDTA on an 1:1 ratio, the percentage amount of Fe present in the sample was determined to be 11.78% ± 0.19313.