Potentiometry : Background
If a strip of silver is dipped into a silver nitrate solution, the silver ions cling on to the silver
metal and attract the electrons out of the strip. This leaves the silver strip positively
charged and the solution near the silver metal slightly negatively charged. After sometime
the following equilibrium will be established.
Ag+(aq) + e-

Ag(s)
Fig 1: An electrical double layer between silver and silver nitrate solution
A metal strip in equilibrium with its ions in solution constitutes an electrode (Fig. 1). The
separation of charges between electrode and electrolyte gives rise to the electrode potential.
Now how to measure this electrode potential? One can join a metal wire to the silver strip
but how to connect the electrolyte solution? If we dip a metal wire into the solution, a new
electric double layer is created which gives rise to a new electrode potential. In fact it is
impossible to measure a single electrode potential. To measure the electrode potential, the
hydrogen electrode is coupled with the electrode whose electrode potential is to be
measured as shown in the diagram (Figure 2)
Figure 2: An electrochemical cell for measurement of electrode potential
The electrode potential of standard hydrogen electrode is arbitrarily assigned to be equal to
zero. The two electrodes are electrically connected through a salt bridge. The emf of the cell
may be measured using digital potentiometer. The electrochemical cell diagrammed above
(fig 2) may be represented by the following cell notation.
Pt, H2(l atml / H+ (a = 1) // Ag+ (x M) / Ag
Here Pt, H2(l atm) / H+ (a = l) represents the standard hydrogen reference half cell which
acts as an anode, and Ag/Ag+( x M) is another half cell which acts as cathode and whose
potential we want to measure. The // represents a salt bridge which connects the two half
cells electrically. The emf of this cell is given by

Ecell = EAg+/Ag
-
EH+/H2
The electrode potential of standard hydrogen electrode is, by convention taken as equal to
zero, therefore the E1 itself becomes the electrode potential of Ag/Ag + electrode. The
electrode potential of Ag/Ag+ electrode depends upon the concentration of silver ion in the
solution and temperature of the cell. The exact dependence of the electrode potential on
these factors is given by the Nernst's equation.
EAg+/Ag = EoAg+/Ag€ + 2.3O3 RT/nF log [Ag+l
Where,
Eo = standard electrode potential (reduction)
R = universal gas constant equal to 8.315J mol-l deg-l
T = temperature in Kelvin
n = number of electron involved in the electrode reaction
F = Faraday, 96500 coulombs
At 25oC, 2.303 RT/F works out to be equal to 0.0591 substituting this value in the above
equation we get
EAg+/Ag = EoAg+/Ag + O.O59l log [Ag+]
The standard hydrogen electrode is difficult to construct and inconvenient to use, so a
number of secondary standard reference electrodes have been devised. One of the most
frequently used reference electrodes which is commercially available is saturated calomel
electrode. The reduction potential of SCE has been accurately determined which is equal to
+0.242 volt at 25oC. When the Ag+/Ag electrode is coupled with the calomel electrode, the
cell may be represented as follows:
Hg/Hg2Cl2 (satd.), KCI (satd.)// Ag+(x M) / Ag
and the emf of the cell is
Ecell
=EAg+/Ag - Ecal
= EAg+/Ag – 0.242
= EoAg+/Ag + 0.059l log [Ag+] – 0.242
= Econst + 0.059l log [Ag+]
Thus by measuring the emf of a cell with SCE as reference electrode, electrode potential of
Ag/Ag+ can be readily calculated from the above equation.
In potentiometric titration of silver nitrate with sodium chloride, the E cell may be plotted
against volume of sodium chloride added to locate the end point since the chage in Ecell is
due to the change in concentration of silver ion.
Ecell
= EoAg+/Ag + 0.059l log [Ag+] – 0.242
= Econst + 0.059l log [Ag+]
where Econst = EoAg+/Ag – 0.242