Redox titration with KMnO4

Introduction

This is a particularly good experiment to use as part of the Scaffolding to prepare students for their individual scientific investigation. It introduces students to the practical technique of redox titration and also provides much scope for good evaluation. Students may already have met a simple form of redox titration (when they used dropping pipettes to determine the amount of chlorine in swimming pool water) but this practical uses traditional burettes, pipettes and volumetric flasks. In fact it could be used to cover one of the mandatory laboratory components (Topic 1.3. Use of the experimental method of titration to calculate the concentration of a solution by reference to a standard solution).

A standard solution of acidified manganate(VII) ions is used to determine the percentage by mass of Fe2+ in a sample of iron(II) ammonium sulfate. Iron(II) ammonium sulfate is used in preference to iron(II) sulfate as it is less susceptible to oxidation by air to Fe3+ ions. The experiment can easily be modified to test for the amount of Fe2+ in an ‘iron’ tablet or to determine the amount of oxalic acid, (COOH)2, in a sample. If it is used for oxalic acid then it will be necessary to warm it first to about 60 oC as this is an example of an autocatalytic reaction. Warming provides the necessary activation energy until some of the Mn2+(aq) produced can act as a catalyst.

Teacher’s notes

Provided students carry it out competently this experiment gives excellent results. Problems will arise if they heat the solution of iron(II) ammonium sulfate to speed up its dissolution as this will result in some of it being oxidized by the air to Fe3+. One practical difficulty they will find will be actually reading the burette. Even though the concentration of the potassium permanganate is quite dilute (0.0200 mol dm-3) it is still a very intense colour. Placing a small piece of white paper behind the burette gradations can help here. Make sure they rinse the burettes out thoroughly after the experiment otherwise the residual brown stain of manganese(IV) oxide, MnO2, will take some shifting. The students do not need to be given the molar mass of the salt as its formula is on the bottle – note that it contains water of crystallisation. The uncertainties can be calculated in the usual way and compared to the percentage error. Because of the possibility of oxidation by air to Fe3+ and the difficulty with reading the burette there is plenty to discuss if you wish students to evaluate the experiment.

Standard Level Higher Level Student worksheet

REDOX TITRATION WITH POTASSIUM PERMANGANATE

The amount of Fe2+ in a sample can be determined by titrating with a standard solution of potassium manganate(VII), KMnO4 in the presence of an acid. An acidified solution of manganate(VII) ions, (or permanganate ions as they are commonly called) is a strong oxidizing agent. It is particularly useful for titrations as it acts as its own indicator. The end-point can be taken when one drop just causes the purple colour due to the MnO4−(aq) ions to disappear. It will be used in this practical to determine the percentage of iron(II) ions in a sample of ammonium iron(II) sulfate. The two half-equations are:

MnO4−(aq) (purple) + 8H+(aq) + 5e− → Mn2+(aq) (colourless) + 4H2O(l)

Fe2+(aq) (pale green) → Fe3+(aq) (pale yellow) + e−

Mn2+(aq) ions are in fact a very pale pink but in the concentrations used here they appear colourless. Similarly Fe3+(aq) can be a yellow-brown colour in higher concentrations but under these conditions the final solution will be pale yellow.

ENVIRONMENTAL CARE:

To save on chemicals this analysis has been scaled down and 10 cm3 pipettes and 100 cm3 volumetric flasks are used instead of the more normal 25 cm3 pipettes and 250 cm3 flasks. Manganese and iron, although essential minerals, are heavy metals and all residues containing them should be placed in the container in the fume cupboard marked 'Heavy Metal Waste'.

SAFETY:

It is not advisable to get manganate(VII) ions on the skin as they stain, however apart from the presence of dilute sulfuric acid there are no particular hazards associated with this practical.

PROCEDURE:

Weigh out accurately about 2.5 grams of AR (analytical reagent) ammonium iron(II) sulfate crystals and dissolve them in about 40 cm3 of approximately 1 mol dm-3 sulfuric acid solution. It is important that you do not heat the solution to assist dissolving. Make up to 100 cm3 with distilled water in a volumetric flask and thoroughly mix the solution. Pipette 10 cm3 of this solution into a conical flask, add about an equal volume of distilled water and titrate with 0.0200 mol dm-3 potassium manganate(VII) solution to a faint pink colour. Repeat the titration at least twice for accuracy, recording all your results.

CALCULATION AND QUESTIONS

1. Use the two half-equations to arrive at the overall equation for the reaction.

2. What amount (in mol) of Fe2+ is required to react with one mole of MnO4− ?

3. What amount (in mol) of MnO4− is present in your average titre value?

4. What amount (in mol) of Fe2+ is present in 10 cm3 of the Fe2+solution?

5. What amount (in mol) of Fe2+ is present in 100 cm3 of the Fe2+ solution?

6. What mass of Fe2+ is present in the ammonium iron(II) sulfate you weighed out?

7. What is the percentage of Fe2+ in ammonium iron(II) sulfate ?

8. Find the correct chemical formula for ammonium iron(II) sulfate crystals from the bottle
and calculate what the correct answer should be.

9. Compare your answer with the correct value and comment on your result.

10. Why do you think you should not heat the iron(II) salt to assist dissolving?

This worksheet can also be downloaded from:

  Permanganate redox titration

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