Engineering Chemistry S Datta

ELECTROCHEMICAL CELLS 243
The electrode potential is given by,
∈ = ∈° + RT
F
1/2
[H 2 ]
log
+
[H ]
∈ = ∈° + 0.059 log
∈ = 0 – 0.059 log [H + ]
= 0.059 pH at 25°C.
1
+ (n = 1)
[H ]
A full cell is constructed by coupling the above half cell with a standard hydrogen
electrode. The e.m.f. E of the cell is then determined potentiometrically. Since electrode potential
of standard hydrogen electrode is zero, the observed e.m.f. of the cell gives directly the electrode
potential (∈) of the half cell. So,
the cell is: (Pt) H 2
| H + (a = 1) || H + (a = unknown) | H 2
(Pt)
E Obs
= ∈ = 0.059 pH
∴ pH = E Obs
0.
059
+ HgCl
paste
–
+
Sat.
KCl
HgCl
&
KCl
H 2
H 2
Calomel
electrode
Salt bridge
Hydrogen
electrode
H 2
Trap
H 2
Solution
Fig. 10.7 Determination of pH by Hydrogen Electrode.
Reference Electrodes
For determination of standard electrode potentials (∈°), standard hydrogen electrode is
used as reference. But there are some difficulties with hydrogen electrode. Maintaining 1 atm.
pressure of hydrogen and to make a H
+ = 1 is very difficult. To overcome this difficulty another
electrode which is known as secondary reference electrode is used for determining standard
electrode potential. One of such electrodes is calomel electrode. To determine pH of a solution
the cell constructed with the calomel electrode will be: (Pt) H 2
| H + (a) || Normal calomel
electrode (Fig. 10.7).