The document discusses electrode potential, which describes how metals behave in solutions of their own salts, detailing oxidation and reduction processes. It explains the standard hydrogen electrode as a reference for measuring electrode potentials and introduces the saturated calomel electrode as a secondary reference. Additionally, it covers electrochemical cells, particularly galvanic cells, illustrating how different metals interact in terms of oxidation and reduction, along with the use of the electromotive series to predict chemical reactions.

1. Electrochemistry
(Electrode Potential)
Dr.S.SURESH
Assistant Professor
Email:avitsureshindia@gmail.com

2. Electrode potential
• The tendency of a metal to get oxidised or reduced when it
is placed in a solution of its own salt is called electrode
potential.
• When a metal [M] is placed in a solution containing its own
ions [Mn+
], then the metal may undergo either oxidation or
reduction. If the metal undergoes oxidation, then the
positive metal ions may pass into the solution
M Mn+
+ ne‒
• If the metal undergoes reduction, then the negative ions
may get deposited over the metal.
Mn+
+ ne‒
M

3. Zn in ZnSO4
When Zn is placed in a solution of
its own salt, zinc undergoes
oxidation with the release of
electrons. The electrons liberated
in the process, accumulate over the
surface of the metal and hence, the
metal is negatively charged. Now,
the negatively charged metal
attracts the positive ions from the
solution, and hence formation of a
double layer takes place near the
surface of the metal.

4. Cu in CuSO4
When Cu is placed in the
solution of CuSO4, the copper
ions in the solution gets
deposited over the metal and
hence the metal becomes
positively charged. The
positively charged metal attracts
the negatively charged sulphate
ions in the solution and hence a
doubly charged layer (Helmholtz
electrical double layer) is formed
near the metal.

5. Single Electrode Potential:
It is a measure of tendency of a metallic
electrode to lose or gain electrons when it is in
contact with a solution of its own salt.
Standard Electrode Potential:
It is a measure of tendency of a metallic
electrode to lose or gain electrons, when it is
in contact with a solution of its own salt of 1
Molar concentration at 25°C.

6. Primary Reference
Electrodes
The electrode potential is found out by coupling
the electrode with a primary reference electrode,
the potential of which is arbitrarily fixed as zero.
The important primary reference electrode used
is a standard hydrogen electrode,

7. Standard hydrogen electrode(SHE)
It consists of a platinum wire
in a inverted glass tube.
Hydrogen gas is passed
through the tube at 1 atm. A
platinum foil is attached at
the end of the wire. The
electrode is immersed in 1M
H+
ion solution at 25°C. The
electrode potential of SHE is
zero at all temperatures.

8. Standard hydrogen electrode(SHE)
It is represented as
Pt, H2(1atm)/H+
(1M)
In a cell when the standard hydrogen electrode acts as
anode, the electrode reaction can be written as
H2(g) 2 H+
+ 2 e‒
When the standard hydrogen electrode acts as cathode,
the electrode reaction can be written as
2H+
+ 2e‒
H2(g)
Based on the electrode potential obtained with reference
to hydrogen, electrochemical series is obtained.

9. Secondary Reference Electrode:
(Saturated calomel electrode)

10. Need for Secondary Reference Electrode
• The use of SHE is difficult, because it is difficult to
maintain 1M H+
ion concentration and the pressure
of the gas at one atmosphere. Also, the electrode
will easily get poisoned in case of traces of
impurities in the gas and hence, other reference
electrodes are used.
Example: Saturated calomel electrode (Saturated
KCl)

11. Saturated calomel electrode
It is a commonly used reference
electrode, it consists of a glass tube,
that contains Hg at the bottom
covered with solid Hg2Cl2 and above
this the tube is filled with KCl
solution. A platinum wire is in touch
with Hg and it is used for electrical
contact. The KCl solution inside the
tube can have ionic contact with
solution outside and acts as a salt
bridge.
The electrode potential of the
calomel electrode is +0.2422V.

12. Determination of standard electrode potential of Silver using
Saturated Calomel Electrode
• To determine the electrode potential of Silver (electrode)
immersed in 1M solution of AgCl, the Ag half cell is
connected with the calomel half cell, through a salt bridge of
potassium chloride. Since the reduction potential of the
coupled Ag electrode is more than E° of calomel electrode
(+0.2422V), the calomel electrode behaves as anode and Ag
acts as cathode

13. Determination of standard electrode
potential of Zinc
The cell may be represented as
Hg, Hg2Cl2(s) KCl (Sat.Sol.) // AgCl (1M), Ag
The EMF of this cell, is measured potentiometrically.
At 25°C, it is found to be 0.56 V.
E°Cell =
0.56 = ER - 0.2422 EAg
= 0.56 + 0.24
EAg = 0.8022 V
ο
L
ο
R E-E

14. Electromotive series
Definition:
When the metals (electrodes) are arranged in
the order of their increasing values of
standard reduction potential on the
hydrogen scale, then the arrangement is
called electromotive series.

15. Electromotive series
• Electrode Electrode Reaction E°

16. Applications of Electromotive series
• The standard EMF of the cell can be calculated if the
standard electrode potential values are known.
E°cell = E°R E°‒ L
• The relative tendency of metals to go into the solution
can be noted with the help of electrochemical series.
Metals on the top are more easily ionised into solution.
• The anode or more active metal with high negative
electrode potential in the series are more prone to
corrosion. The cathode or more noble metals with less
negative electrode potential are less prone to corrosion

17. Applications of Electromotive series
• Using electrochemical series we can predict whether a metal
will displace another metal from its salt solution or not.
Example: Zinc metal having low reduction potential in the
series is easily oxidised to Zn2+
, while copper having higher
reduction potential in the series is easily reduced to copper.
• Metals hydrogen displacement behaviour can be predicted.
Any metal that like above hydrogen in the electrochemical
series can liberate hydrogen from an acid solution.
Example: Zn lying above hydrogen in the electrochemical
series reacts with dilute H2SO4 to liberate hydrogen.
↑+→+ ++
2
2
HZnHZn

18. Electrochemical cell (or) Galvanic cell
A galvanic cell is an electrochemical cell in which
the electrons are transferred due to redox reaction
to get electrical energy. In a galvanic cell, two
different electrodes are kept immersed in their
respective salt solutions and connected by means
of a salt bridge
Example: Daniel cell

20. Daniel Cell
• When a zinc rod in contact with 1M ZnSO4 and a Cu rod
in contact with 1M CuSO4 are connected, Zn goes into
the solution as Zn2+
ions and the electrons released
flows through the external wire reaches the copper
electrode where copper gets reduced. A salt bridge is
used to maintain the electrical continuity between the
two half cells, also it eliminates the liquid junction
potential.
Zn Zn2+
+ 2e-
E° = - 0.76 V
Cu2+
+ 2e-
Cu E° = + 0.34 V
Zn + Cu2+
Zn2+
+ Cu -
E°Cell = 1.10 V

21. REPRESENTATION OF A GALVANIC CELL
The following conventions are used in representing an
electrochemical cell:
1. A galvanic cell is represented by writing the anode
(where oxidation occurs) on the left hand side and
cathode (where reduction occurs) on the right hand
side.
Anode // Cathode
2. The anode of the cell is represented by writing metal
first and then the electrolyte (or the cation of the
electrolyte)
Zn/Zn2+

22. Representation of a Galvanic cell
3.The cathode is represented by writing the
electrolyte first and then metal.
Cu2+
/Cu
4.The two half cells are separated by a salt bridge,
which is indicated by two vertical lines.
Zn/ZnSO4 // CuSO4/Cu
or
Zn/Zn2+
// Cu2+
/Cu