Iron tablets contain iron (II) sulfate which is a soluble inexpensive form of ‘iron supplement’. The experiment is to determine the percentage by mass of iron (II) sulfate in each tablet. Iron (II) ions can be oxidised to iron (III) ions by potassium manganate (VII) in acidic solution. In acidic conditions the deep purple solution of manganate (VII) ions is reduced to a very pale pink solution of manganese (II) ions. This solution is so pale as to appear colourless when dilute and, in practice, the marked difference in colour between these two oxidation states is useful as an end-point for this redox reaction.
The manganate (VII) ion accepts electrons and is reduced to colourless Mn2+ ions according to the following half-equation.
MnO4 – (aq) + 8H+ (aq) + 5e– Mn2+(aq) + 4H2O(l) (purple to colourless)
The electrons are provided by the iron (II) ions which act as the reducing agent. Fe2+(aq) à Fe3+(aq) + e– The potassium manganate (VII) solution is added from the burette to the solution of the reducing agent and is immediately decolourised. As soon as the reducing agent is used up, the next drop of potassium manganate (VII) solution is not decolourised and colours the solution in the conical flask a pale purple colour (often described as a permanent pink colour). The end-point is the first appearance of this pale purple colour. Potassium manganate (VII) is therefore self-indicating and no other indicator is required. The acid used to provide H+ (aq) is dilute sulfuric acid, which should always be in excess otherwise insoluble brown manganese (IV) oxide (MnO2) will form. The method described in this experiment ensures that there is a large excess of sulfuric acid in each titration.
To analyse iron tablets by titration using potassium manganate (VII) in acidic solution. Whenever possible, students should work individually. If it is essential to work in a pair or in a small group, because of the availability of apparatus, supervisors must be satisfied that they are able to assess the contribution from each student to the practical activity
- Weighing bottle or boat
- Five iron tablets
- Approximately 1 mol dm-3 sulfuric acid (50 cm3)
- 100 cm3 conical flask with stopper
- Filter funnel and paper
- Deionised or distilled water in a wash bottle
- 100 cm3 graduated (or volumetric) flask
- Burette stand and clamp
- 250 cm3 beaker
- 25 cm3 pipette and pipette filler
- 25 cm3 measuring cylinder
- Two 250 cm3 conical flasks
- 0.0200 mol dm-3 potassium manganate (VIi) solution (150 cm3)
- Approximately 1 mol dm-3 sulfuric acid (100 cm3)
- Using a weighing bottle, weigh accurately five iron tablets.
- Place the iron tablets into a 100 cm3 conical flask and add approximately 50 cm3 of the 1 mol dm-3 sulfuric acid provided.
- Stopper the conical flask, shake its contents well and then leave the tablets to dissolve. This is a slow process and should be carried out at least one day before the titration is to be attempted. The outer coating of each tablet is insoluble in water, but slowly breaks down in the acidic solution. The solution will need filtering before carrying out the titration.
- Without disturbing the residue, which will have settled to the bottom of the flask, carefully filter the solution directly into a 100 cm3 graduated (volumetric) flask.
- Rinse the residue in the filter paper into the graduated flask using a small volume of de-ionised or distilled water.
- Add dilute sulfuric acid to make the solution in the graduated flask up to the mark.
- Ensure that the contents of the graduated flask are fully mixed. You now have an acidified solution of iron (II) sulfate.
- Fill a burette with the 0.0200 mol dm-3 potassium manganate (VII) solution provided.
- Pour some of the contents of the graduated flask into a clean 250 cm3 beaker and, using a 25 cm3 pipette and a pipette filler, measure out a 25.0 cm3 sample of the iron (II) sulfate solution into a clean 250 cm3 conical flask.
- Using a 25 cm3 measuring cylinder, measure out 25 cm3 of the 1 mol dm-3 sulfuric acid provided and add this to the contents of the conical flask.
- Titrate this acidified sample of iron (II) sulfate solution by adding potassium manganate (VII) from the burette until the first permanent pink colour is seen.
- You will only be able to carry out three titrations and if you are careful, you should be able to obtain at least two results that are concordant. Record the three results that you obtain.
- Calculate and record the mean volume of potassium manganate (VII) solution used in the titration (the average titre). Show your working.
You should be able to:
- Combine the two half-equations (given in the introduction) to give the overall redox equation for the reaction that has taken place during the titration.
- Use your overall equation to determine the ratio of moles of manganate (VII) ions that react with iron (II) ions.
- Use the average titre to calculate the moles of manganate (VII) ions which have been used in the titration.
- Calculate the amount, in moles, of iron (II) ions in the 25 cm3 sample of iron (II) sulfate.
- Calculate the amount, in moles, of iron (II) ions in the 100 cm3 graduated flask at the start of the experiment.
- Calculate the mass of Fe in the original five iron tablets and hence the mass of Fe in one iron tablet.
- Compare your value for the mass of Fe with the information from the supplier about the composition of each iron tablet.