From what I understand about concentration cells - which I never learned about when I took general chemistry - they rely on the fact that if we have differing concentrations of some solute separated by a permeable membrane, they will try to exchange solute and/or solvent to achieve the same concentration or as close to it as possible. Correct?
Now, in a concentration cell, there is no physical exchange of ions. Instead, a salt bridge facilitates the flow of electrons, which either:
a) helps convert solids into ions
or
b) helps convert ions into solids
So in the above problem, we obviously have unequal concentrations of iron(III) ions. We can describe the problem using cell notation:
${|Fe|Fe^{3+} (4.8*10^-7 M)||Fe^3+ (1.3*10^-3 M)|Fe|}$
The flow of electrons will be from half-cell A to half-cell B. These are the reactions:
Half-cell A: $\ce{Fe -> Fe^3+ + 3e^-}$
Half-cell B: $\ce{Fe^3+ + 3e^- ->Fe}$
This increases the $\ce{Fe^3+}$ concentration in half-cell A while decreasing the $\ce{Fe^3+}$ concentration in half-cell B. This will continue until both half-cells have the same $\ce{[Fe^3+]}$.
So, we can also describe half-cell A as the anode, because it is where oxidation occurs, and half-cell B as the cathode, as it is where reduction occurs.
So far so good?
And if we want to calculate the electrochemical potential of this particular half-cell (wired up to make a spontaneous reaction), we can apply the Nernst equation as follows:
