Titration Experiment to Detemine the Mass of Iron (Ii) Sulphate in an 'Iron' Tablet

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Aim: To find out the mass of Iron (II) sulphate each ‘Iron tablet’ contains Background Knowledge: ‘Iron tablets’ are prescribed to patients suffering from anaemia. They contain 200mg (0.200g) of iron (II) sulphate FeSO4 We are going to analyse iron tablets to find out how much Iron (II) sulphate they contain. The procedure we will use is a REDOX reaction, by oxidising the Fe2+ contained in the Iron (II) sulphate to Fe3+. We will use potassium manganite (VII) as the oxidising agent. The active part of this is the manganite ion, MnO4-1. The SO4-2 in the FeSO4 and the K+ in the KMnO4 do not react, we say they are ‘spectator’ ions and so we will ignore them. Apparatus:

* 2 Iron tablets
* Pestle & Mortar
* 1M Sulphuric acid
* 100cm3 Volumetric flask
* Weighing scales
* Dropping Pipette
* Potassium Manganate (VII) – KMnO4
* Burette
* 10cm3 Volumetric pipette
* Conical Flask
* White tile
Potassium Manganate (VII) solution
Potassium Manganate (VII) solution
Iron sulphate solution
Iron sulphate solution

Method:
1. Weigh the 2 iron tablets separately
2. Crush them firmly using the pestle and mortar until the tablets become crystals 3. Using a dropping pipette, add a few drops of sulphuric acid to the iron crystals and stir into a paste with the pestle 4. Transfer the iron solution into a volumetric flask carefully while using a funnel 5. Use the sulphuric acid to rinse out the mortar and pestle into the volumetric flask so all the Iron is collected 6. Rinse out the funnel as well

7. Add more sulphuric acid to a little before the mark, then use the dripping pipette to add the sulphuric acid up to the mark and cork it 8. Shake the volumetric flask
9. Rinse out the burette with the KMnO4 solution then fill it up to the zero mark 10. Rinse out the Volumetric pipette with the Iron Sulphate + Sulphuric acid solution 11. Use the volumetric pipette to measure out 10cm3 of the Iron solution and pour this solution into the conical flask 12. Titrate the KMnO4 solution against the Iron solution by adding the KMnO4 solution carefully until it turns colourless 13. Measure the volume of KMnO4 solution used to complete the reaction 14. Empty the contents of the conical flask and repeat the experiment at least twice until 2 results (volume used) between 0.1 of each other are obtained. -------------------------------------------------

Results table (Raw results): Volume of KMnO4 solution used
| Initial/cm3 (+/- 0.1cm3)| Final (cm3)(+/- 0.1cm3)| Titre (cm3)(+/- 0.1cm3)| Observations| Rough| 0.0| 10.1| 10.1| Difficult to see meniscus|
1st| 10.1| 19.7| 9.7| Pale pink |
2nd | 19.7| 29.4| 9.7| Pale pink|
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Also, mass of each Iron tablet was – 0.46g
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Observations:
* -------------------------------------------------
The iron tablets crushed quite easily
* -------------------------------------------------
The pill didn’t crush evenly
* -------------------------------------------------
When H2SO4 was added, not all the pieces dissolved.
* -------------------------------------------------
The colour change (purple to colourless) during titration was very sharp as the colours were very distinct from each other * -------------------------------------------------
The meniscus was quite difficult to see as the KMnO4 solution stained the burette a bit * -------------------------------------------------
There was no indicator needed so this means it is a self-indicating reaction Analysis:
Calculation of mass of Iron Sulphate in the tablet
Equation (without spectator ions):
5Fe+2 + MnO4-1 + 8H+ 5Fe+3 + Mn+2 + 4H2O
The Fe+2 ion is oxidised (loses an electron) to form Fe+3.
KMnO4
Volume - 9.7cm3
Concentration – 0.005moldm-3

FeSO4
Volume – 10cm3
Concentration - ?

C=n/v n (KMnO4) = C X V...
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