Page 97 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

EQUlllBRlUM CONSTANTS OF OXIDATION-REDUCTION REACTIONS 2.33

The term 4H20 is omitted, since the reaction is carried out in dilute solution,
and the water concentration may be assumed constant. The hydrogen ion
concentration is taken as molar. The complete reaction may be divided into
two half-ce11 reactions corresponding to the partial equations:

and Fe2+eFe3++e (36)
For (35) as an oxidation-reduction electrode, we have:

The partial equation (36) may be multiplied by 5 in order to balance (35)
electrically :

For (37) as an oxidation-reduction electrode:

Combining the two electrodes into a cell, the e.m.f. will be zero when equilibrium
is attained, i.e.

[Mn2+] x [Fe3+]' - 5(1.52 - 0.77)
or log - = 63.5
[Mn041 x [Fe2+]' x [H+I8 0.0591

This result clearly indicates that the reaction proceeds virtually to completion.
It is a simple matter to calculate the residual Fe(I1) concentration in any
particular case. Thus consider the titration of 10 mL of a 0.1 M solution of
iron(I1) ions with 0.02M potassium permanganate in the presence of hydrogen
ions, concentration 1 M. Let the volume of the solution at the equivalence point
be 100mL. Then [Fe3'] =O.OlM, since it is known that the reaction is
practically complete, [Mn2+] = f x [Fe3+] = 0.002 M, and [Fe2+] = x. Let
the excess of permanganate solution at the end-point be one drop or 0.05 mL;
its concentration will be 0.05 x 0.1/100 = 5 x IO-' M = [MnOi]. Substituting
these values in the equation:

It is clearfrom what has already been stated that standard reduction potentials
may be employed to determine whether redox reactions are sufficiently complete