OK...in order to answer your question fully we need to consider two points.
a) What would the products be if the reaction had to occur?
Fe3+ is reduced to Fe2+ according to the half-equation:
Fe3+ + e
Fe2+
Sn2+ is oxidised to Sn4+ according to the half-equation:
Sn2+
Sn4+ + 2e
However, since the number of electrons involved in both cases has to be the same, then we have to multiply the first equation by 2, to obtain:
2Fe3+ + 2e
2Fe2+
Summing up the two half-equations:
2Fe3+ + Sn2+
2Fe2+ + Sn4+
b) Should the reaction occur?
Just because you have determined the redox products does not necessarily guarantee that the reaction will occur. To determine this you need electrode potentials. In this particular case they are:
Fe3+ + e
> Fe2+ E = +0.77V
Sn4+ + 2e
> Sn2+ E = +0.154V
Now, a positive electrode potential implies that the equilibrium prefers to go to the right, so Fe2+ and Sn2+ are the preferred states for both half-equations. However, Fe3+ prefers to be reduced (0.77) much more than Sn2+ does not prefer to be oxidised (0.154), so the Fe3+ will win and the reaction should occur.
Please notice that I say SHOULD occur because this is merely a thermodynamic prediction. It is no guarantee that the reaction will, from a kinetic perspective, be fast enough or even proceed at all.
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