Help please! Oxidation/Reduction question on iron sulfate :(

(c) Aqueous iron(II) sulfate takes part in redox reactions.
Using oxidation numbers, show that both reduction and oxidation have taken place in the
redox reaction of aqueous iron(II) sulfate shown below.
6FeSO4 + 7H2SO4 + Na2Cr2O7 3Fe2(SO4)3 + Cr2(SO4)3 + Na2SO4 + 7H2O
............................................................................................................................................................................................................................................................................. [2]

I can do simple redox questions most of this time, but with this i get a little confused.
Could someone tell me how i would do it?

Work out the oxidation state for iron on the left and on the right.

It has changed by either adding or losing electrons. You have to decide which.

Then do the same for chromium.

I understand that at the start Fe = +2 and Cr = +6
but then..

So what is the oxidation state of iron in Fe₂(SO₄)₃ ?

and chromium in Cr₂(SO₄)₃ ?

Could you explain how i would do that please?

Well you know that the SO4 ion has a 2- charge. So if there are 3 sulphate ions there is an overall negative charge of 6. There are also 2 iron ions in that compound. The charge of the 2 ions must add up to +6 to balance out the -6. That means the charge of each Fe ion is +3. So their oxidation state is +3. Which is an increase from +2, so they have been oxidised.

Now do the same with Chromium, you should find that the oxidation state has decreased which means they have been reduced.