OCR A2 Assessment - Determination of the Formula of Hydrated Iron(II) Sulphate

Hey there,

I would appreciate some help with my coursework that has to be in tomorrow (o dear!). I basically have to find x in FeSO4.xH2O.

Method 1:
1. Using a balance that weighs to two decimal places, weigh a crucible. Add between 1.30 – 1.50 g of hydrated iron II sulphate crystals. Record the masses.
2. Place the crucible containing the hydrated iron(II) sulphate crystals on the pipe-clay triangle and gently heat for about two minutes.

3. Allow to cool and weigh the crucible and the iron(II) sulphate.
4. Repeat steps 2 and 3.
5. Record all of the masses.

Method 2: Titration
1. Weigh as accurately as possible between 2.85 and 3.10 g of hydrated iron(II) sulphate crystals, FeSO4.xH2O. Record the mass and then dissolve the crystals in 50.0 cm3 of 1 mol dm-3 H2SO4(aq) and make up to 250 cm3 in a volumetric flask with distilled water. Invert the volumetric flask several times to ensure that the solution is evenly mixed.
2. Using a pipette filler, pipette 25.0 cm3 of the solution of iron(II) sulphate into a conical flask and add approximately 20.0 cm3 of the 1 mol dm-3 H2SO4(aq) solution provided.
3. Titrate the acidified Fe2+(aq) solution with 0.0100 mol dm-3 potassium permanganate, KMnO4(aq), and continue the titration to the normal end-point.
4. Repeat the titration until you have obtained concordant results

Right so that's basically what we had to do and all went fine. However, coming to analysis I realized that the results I'm getting are nowhere near right. Our teacher told us that the ratio we should be getting for the formula (x) should be 7 as that is the correct formula. However, I only got 4.

Here are the masses I used:
mass of iron(II) sulphate [after drying off the oxygen) = 0.944g
mass of water dried off = 0.481g
this gave me a ratio of 0.0062:0.0267 of moles and therefore the following ration in terms of molecules:
1:4 - this would lead to the FeSO4.4H2O formula which is not correct

I am wondering whether I have done the right calculations to arrive to this answer as I have a feeling I may not have dried all the water off and that is why I am getting such a high ratio.

Oh and I almost forgot; our teacher told us that for a level 8, the mark scheme says:

"The accuracy of the formula is derived and justified, with regards to all measurements and data used"

What on earth does that mean?

I have one more question - how would I be able to calculate the formula from method 2 because I don't know the concentration of iron (II) sulphate only in combination with the 1 mol H2SO4 ?

Okay I hope someone can help me out here!

Cheers

Reply 1

vVPiKaVv

Your calculations are correct. Perhaps you didn't completely remove all the water of crystallisation from the hydrated salt.

As for method 2, if you work at it backwards,
you'll first have the volume of potassium permanganate required to react with the FeSO4 in solution.
You can then balance the redox equation between potassium permanganate and Fe (II) ions.
Work out the no. of moles of potassium permanganate used, and then the no. of moles of Fe (II) ions in solution [also the no. of moles of FeSO4].
Calculate the mass of FeSO4 present and then subtract it from the mass of hydrated FeSO4 used. You'll then know how to work it out via the calculations you did for method 1.