What is the equivalence point in the context of a redox titration? [closed]








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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?
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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.




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](https://en.wikipedia.org/wiki/Reaction_quotient) reaches the reaction [equilibrium constant](https://en.wikipedia.org/wiki/Equilibrium_constant).


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