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We know the standard reduction potential of the species: $\ce{Fe^{3+}}/\ce{Fe^{2+}}$, $\varepsilon_1$ and the standard reduction potential of the species: $\ce{MnO_4}^{-}/\ce{Mn^{2+}}$, $\varepsilon_2$. If we know the concentration of the ferrous cation in the initial solution $c_0$ in sulphuric medium $c_s$, at equivalence point in titration with potassium permanganate $c_t$ how to determine the ratio of concentrations of ferric and ferrous cations?

What I have thought on the subject is that Nernst's potential equation is useful, but I don't really know what equivalence point should mean in a redox context. I know it only from elementary acid-base equilibria.

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As the ongoing reaction is

$$\ce{5 Fe^2+(aq) + MnO4-(aq) + 8 H+(aq) -> 5 Fe^3+(aq) + Mn^2+(aq) + 4 H2O(l)}$$

the point of equivalence is when the initial molar amount of $\ce{MnO4-(aq)}$ is equal to 1/5 of the initial molar amount of $\ce{Fe^2+(aq)}$.

Another condition comes from the equilibrium requirement that (not standard) redox potentials of all redox systems must be equal

$$E = E^{\circ}_{\ce{MnO4-}/\ce{Mn^2+}} + \frac{RT}{5F} \ln{\left(\frac{[\ce{MnO4-}][\ce{H+}]^8}{[\ce{Mn^2+}]}\right)}=E^{\circ}_{\ce{Fe^3+}/\ce{Fe^2+}} + \frac{RT}{F} \ln{\left(\frac{[\ce{Fe^3+}]}{[\ce{Fe^2+}]}\right)}$$

So the reactants do not react completely, but the reaction stops when both redox potential are equal. It is similar like when a general reaction stops (its net rate) when the reaction quotient reaches the reaction equilibrium constant.

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