The Nernst equation

Set-up and procedure
Set up the experiment as shown in Fig. 1.
Prepare the solutions required for the experiment as follows:
– 0.01 molar K4 [Fe(CN)6] solution: Weigh 4.2239 g of potassium
hexacyanoferrate(II) (yellow prussiate of potash:
K4[Fe(CN)6] • 3H2O) into a 1000 ml volumetric flask, dissolve
it in distilled water, and make up to the mark with distilled
water.
– 0.001 molar K4[Fe(CN)6] solution: Pipette 100 ml of
0.01 molar potassium hexacyanoferrate(II) solution into a
1000 ml volumetric flask and make up to the mark with distilled
water.
– 0.01 molar K3[Fe(CN)6] solution: Weigh 3.2925 g of potassium
hexacyanoferrate(III) (red prussiate of potash):
K3[Fe(CN)6]) into a 1000 ml volumetric flask, dissolve it in distilled
water, and make up to the mark with distilled water.
– 0.001 molar K3[Fe(CN)6] solution: Pipette 100 ml of
0.01 molar potassium hexacyanoferrate(III) solution into a
1000 ml volumetric flask and make up to the mark with distilled
water.
Attach the two burettes, one for the Fe(II) solution and the other
for the Fe(III) solution, to the retort stand, rinse twice with the
respective 0.001 molar solution and fill them. Prepare the sample
solutions to be investigated as listed in Table 1.
Place the beaker containing the first sample solution on the
magnetic stirrer and put in a magnetic stirrer bar. Connect the
platinum electrode, the reference electrode and the temperature
probe to the pH-meter, and dip them into the solution. Record
the temperature and the e.m.f. of the cell. Rinse the electrodes
thoroughly in distilled water, dry them and place them in the next
sample solution. Solutions must be changed quickly. Do not
allow the reference electrode to remain out of solution for too
long. Continue until all sample solutions have been measured.
Repeat the procedure with the 0.01 molar solutions.
Plot the cell e.m.f. (E) as a function of
Table 1: Preparation of the sample solutions
Theory and evaluation
In the electrochemical cell used here, the Ag(s)
AgCl(s)
Cl- electrode
used as a reference electrode supplies a constant potential
against which we measure the potential of the redox electrode.
The silver chloride electrode consists of a silver wire covered
with silver chloride which is immersed into a potassium chloride
solution of defined concentration.
The redox system is an iron(III) / iron(II) couple
[Fe(CN)6]3- + e- S [Fe(CN)6]4-
In general, a redox reaction in which Xz+ ions are reduced by n
electrons (supplied by an inert metal electrode) to Y ions of
charge (z – n)+ can be expressed as follows
Xz+
(aq) + n e- S Y(z – n)+
(aq)
In this system, equilibrium is attained when the sums of the electochemical
potentials on each side of the reaction are equal:
?X (soln.) + n ?e (metal) = ?Y (soln.) (1)
From the definition of the electrochemical potential, it follows
that
?X (soln.) = ?X + z F
soln. (2a)
?e (metal) = ?e + F
soln. (2b)
?Y (soln.) = ?Y + (z – n) F
soln. (2c)
with

soln. Electric potential of the solution

metal Electric potential of the inert metal electrode
?i Chemical potential of species i
Combining equations (1) and (2) we obtain
?X
z+ - ?Y
(z-n)+ + n ?e- = n F
soln. - n F
metal (3)
which allows the electric potential difference ?
between the
solution and the metal to be expressed