To prepare a standard solution of ammonium iron(II) sulfate and use this to standardise a solution of KMnO4 by titration.

Describe the appearance of the ammonium iron(II) sulfate crystals.

a green crystalline material.

Why do you use ammonium iron(II) sulfate rather than just use iron(II) sulfate?

1. It can be obtained in a high degree of purity.

2. It is not affected by the air, i.e. the Fe2+ ions in ammonium iron(II) sulfate are not oxidised to Fe3+ ions by the oxygen in the air.

FeSO4 is not a primary standard because (a) the crystals are oxidised slightly by the air and (b) the crystals lose their water of crystallisation when exposed to the air (efflorescence).

What colour change is observed during this titration?

colourless to pink.

Why was it not necessary to weigh out exactly 9.8 g of ammonium iron(II) sulfate?

In order to calculate the concentration of the ammonium iron(II) sulfate solution we need to know the exact amount that has dissolved in the dilute acid. This amount need not be exactly 9.8 g. It can be any amount that has been accurately measured. When calculating molarity all we need to know is the precise amount of salt in the solution.

Describe how you would take the reading in the burette when using KMnO4?

from the top rather than the bottom of the meniscus.

Dilute sulfuric acid is added on two occasions during this experiment. Give the reason for its addition on each occasion.

1. when making up the solution of ammonium iron(II) sulfate. It is used to prevent the Fe2+ ions being oxidised to Fe3+ ions by oxygen in the air or oxygen dissolved in the water;

2. to the conical flask at the beginning of the titration to supply the H+ ions in order for the reaction to occur

When carrying out the titration another student observed that the first few drops of KMnO4added to the conical flask were decolourised slowly, but subsequent drops were decolourised rapidly. Explain this observation.

This observation is due to autocatalysis, i.e. the reaction is catalysed by the Mn2+ ions formed in the reaction: MnO4- + 8H+ + 5e-Æ Mn2+ + 4H2O

As soon as the Mn2+ ions are formed the rate of the reaction increases.

When performing this experiment a student noticed a dark brown colour being formed in the conical flask. What conclusion would you draw from this observation?

The student has probably forgotten to add the dilute sulfuric acid at the beginning of the titration. This omission causes a dark brown precipitate of MnO2 to be formed.