Some methods of iron determination in the oxidation states II and III are discussed in Problem 12. Y — Physical Chemistry — Thermodynamics Chemistry Question
Titrimetric determination of Fe in different oxidation states
Some methods of iron determination in the oxidation states II and III are discussed in Problem 12. You are invited to test one more approach to solving that problem in practice.
Reagents and solutions required
* KIO3 reagent grade, solid
* Ascorbic acid, solid
* KI aqueous solution, 5%
* HCl conc. aqueous solution and solution with c = 2 mol dm-3
* HNO3, conc. aqueous solution,
* Sulfosalicylic acid, 25% aqueous solution,
* NH3, 10% aqueous solution,
* EDTA, standard aqueous solution, concentration about 0.05 mol dm-3 (the exact concentration will be given)
1. Preparation of a primary standard solution of KIO3
1.1. Calculate with the accuracy of 0.0001 g the mass of KIO3 necessary for the preparation of 200.0 cm3 of KIO3 solution (c = 0.01000 mol dm-3)
1.2. Using analytical balance weigh out accurately a portion of KIO3. The mass of the portion may differ from the calculated one no more than by 0.05 g and it should be measured with an accuracy of 0.0001 g.
1.3. Transfer the portion into 200.0 cm3 volumetric flask, dissolve it in water, dilute to the mark and mix.
1.4. Calculate the exact concentration (in mol dm-3 ) of the solution prepared.
2. Preparation of the titrant solution (ascorbic acid)
2.1. Calculate with the accuracy of 0.01 g the mass of ascorbic acid necessary for preparation of 200 cm3 of its aqueous solution with c = 0.1 mol dm–3.
2.2. Using technical balance weigh out a portion of ascorbic acid. Its mass may differ from the calculated one no more than by 0.05 g.
2.3 Dissolve the portion in ~200 cm3 of water, mix well, transfer the solution into a vessel and close it tightly with a stopper.
3. Standardization of the ascorbic acid solution
3.1. Fill in a burette with the ascorbic acid solution.
3.2. With a pipette transfer 10.00 cm3 of standard KIO3 solution into a 100 cm3 Erlenmeyer flask, add 20 cm3 of 5% KI solution and 5 cm3 HCl aqueous solution with c = 2 mol dm–3
3.3. Titrate the mixture with the ascorbic acid solution until the iodine color disappears.
Note. When titrating iodine with solutions of reducing agents, starch is usually added as an indicator. Here it is not recommended to do so because the reaction rate decreases significantly in presence of starch.
3.4. Repeat the titration until three titrant volumes differ no more than by 0.10 cm3.
3.5. Calculate the average titrant volume.
3.6. Calculate the concentration of the ascorbic acid in the solution in mol dm-3.
4. Determination of Fe(III) by ascorbimetric titration
4.1. From your instructor you obtain a sample solution containing Fe(II) and Fe(III) in a 100.0 cm3 volumetric flask. Dilute the solution to the mark with water and mix.
4.2. Fill in the burette with the standardized ascorbic acid solution.
4.3. Measure with a pipette 10.00 cm3 of the sample solution into a 100 cm3 Erlenmeyer flask, add 40 cm3 of water and heat nearly to boiling.
4.4. Into the hot solution add 4 – 5 drops of 25% sulfosalicylic acid solution as an indicator.
4.5. Titrate the solution with the ascorbic acid solution until the violet color disappears. During the titration and especially near the end point the solution must be hot. You may need to heat it additionally, if necessary. Near the end point the ascorbic acid solution should be added slowly.
4.6. Repeat the titrations until three titrant volumes differ no more than by 0.10 cm3.
4.7. Calculate the average titrant volume.
4.8. Calculate the mass of Fe(III) in the sample solution under investigation.
Note. Ascorbic acid, especially in aqueous solutions, is instable and oxidizes with oxygen from the air. Therefore the standardization of ascorbic acid solution and ascorbimetric determination of Fe(III) must be carried out during one workday.
5. Determination of total iron by complexometric titration
5.1. Fill in the burette with an EDTA standard solution.
5.2. With a pipette transfer 10.00 cm3 of the sample solution into a 100 cm3 Erlenmeyer flask. Add 5 cm3 of conc. HCl and 2 cm3 of conc. HNO3 to oxidize Fe(II) present in the sample to Fe(III). Cover the flask with a watch glass, heat until boiling and continue heating for 3 – 5 min avoiding splashing.
5.3. Cool down the solution and neutralize it carefully adding 10% NH3 aqueous solution dropwise until color changes from lemon yellow to yellowish brown and slight turbidity persists.
5.4. Add 1 – 2 drops of HCl solution (2 mol dm-3) to dissolve the precipitate, then add 0.5 cm3 of the HCl solution more, dilute up to 50 cm3 with distilled water and heat nearly to boiling.
5.5. Into the hot solution add 4 – 5 drops of 25% sulfosalicylic acid solution as an indicator.
5.6. Titrate the solution until color changes from violet to clear yellow. During the titration and especially near the end point the solution must be hot. You may need to heat it additionally, if necessary. Near the end point the EDTA solution should be added slowly.
5.7. Repeat the titrations until three titrant volumes differ no more than by 0.10 cm3.
5.8. Calculate the average titrant volume.
5.9. Calculate the total mass of iron in the sample solution given to you.
5.10. Calculate the mass of Fe(II) as a difference between the results obtained in 5.9 and 4.8.
Write down the balanced equations of all the reactions taking place during standardization of ascorbic acid solution. Ascorbic acid C6H8O6 is being oxidized to dehydroascorbic acid C6H6O6.
Model Answer
IO3– + 5 I– + 6 H+ → 3 I2 + 3 H2O
I2 + C6H8O6 → 2 I– + C6H6O6 + 2 H+
KIO3 in presence of excess of KI can be used as a primary standard for HCl standardization as well. The method is similar to that described above with the exception that no HCl is added to the titrated solution in this case. Which compound(s) can be used as an indicator(s) for that titration:
□ - starch,
□ - sulfosalicylic acid,
□ - methyl orange,
□ - methyl orange + Na2S2O3 (in excess).
Model Answer
Methyl orange + Na2S2O3 (in excess)
Write down the balanced equations of all the reactions taking place during Fe(III) determination. Ascorbic acid C6H8O6 is being oxidized to dehydroascorbic acid C6H6O6.
Model Answer
2 Fe3+ + C6H8O6 → 2 Fe2+ + C6H6O6 + 2 H+
In what media does ascorbic acid exhibit its reducing properties most markedly?
□ - in acidic,
□ - in neutral.
□ - in alkaline,
□ - reducing properties of ascorbic acid do not depend on the pH.
Model Answer
In alkaline media
Write down the balanced equations of all the reactions taking place during total Fe determination.
Model Answer
3 Fe2+ + NO3– + 4 H+ → 3 Fe3+ + NO + H2O
Fe3+ + Y4– (EDTA anion) → FeY–
One of the crucial items in the Fe(III) determination by complexometric titration is strict maintenance of solution acidity. What are the reasons for that?
□ - If the acidity is too low, Fe(OH)3 precipitates.
- If the acidity is too high, complex of Fe(III) with sulfosalicylic acid does not form.
- If the acidity is too high, complex of Fe(III) with EDTA acid does not form.
- If the acidity is too low and/or too high, the titrant decomposes.
Model Answer
- at too low acidity Fe(OH)3 precipitates,
- at too high acidity complex of Fe(III) with sulfosalicylic acid does not form,
- at too high acidity complex of Fe(III) with EDTA acid does not form.