UNIT 1
ELECTRODE SYSTEM AND CORROSION SCIENCE
Electrode and Cells
Electrochemistry is the branch of chemistry which deals with the study of interconversion of
electrical energy and chemical energy and study of behaviour of electrolytes.
Chemical reactions that involved in the interconversion of electrical energy and chemical energy are
called as electrochemical reactions.
Electrochemical cell is a device which is used to interconvert chemical energy and electrical energy.
It consists of two electrodes and an electrolyte solution.
A metal rod which is dipped in a solution containing its own ions is called an electrode. The electrode
at which oxidation takes place (loss of electrons) is called anode and the electrode at which reduction
takes place (gain of electrons) is called cathode.
Ex : Zn in contact with ZnSO4 solution.
Cu in contact with CuSO4 solution.
Each electrode of the cell is called half cell.
A spontaneous chemical process is the one which can take place on its own and in such a
process Gibbs free energy of a system decreases.
Types of Electrochemical Cell
1. Galvanic cells
2. Electrolytic cells
Galvanic Cell
The galvanic cell converts chemical energy into electrical energy i.e, electricity can be obtained with
the help of redox reaction. The oxidation and reduction take place in two separate compartments.
Each compartment consists of an electrolyte solution and metallic conductor which acts as an
electrode. The compartment containing the electrode and the solution of the electrolyte is called half
cells. Ex: Daniel cell.
Electrolytic Cell
The electrolytic cell converts electrical energy to chemical energy. Here the electrodes are dipped in
an electrolytic solution containing cations and anions. On supplying current the ions move towards
electrodes of opposite polarity and simultaneous reduction and oxidation take place.
Ex: - Electroplating of Cu,Ni,Zn etc.
1
,Differences between Galvanic cell and Electrolytic cell:
Galvanic cell Electrolytic cell
1 Chemical energy is converted into Electrical energy is converted into chemical
electrical energy energy
2 Anode is -ve electrode & Anode is + ve electrode &
Cathode is +ve electrode Cathode is -ve electrode
3 It can be used as a source of energy It cannot be used as a source of energy
Concentration cells:
Concentration cells can be defined as electrochemical cells that consist of two half-cells
wherein the electrodes are the same, but they vary in concentration. As the cell as a whole strives to
reach equilibrium, the more concentrated half cell is diluted and the half cell of lower concentration
has its concentration increased via the transfer of electrons between these two half cells. Therefore,
as the cell moves towards chemical equilibrium, a potential difference is created.
The cell potential depends on the concentration of electrolyte. A concentration cell is
defined as a galvanic cell obtained by dipping the same electrode in the same electrolyte of
different concentrations.
Consider 2 silver rods dipped in AgNO3 solution of
different concentration say C1 and C2 such that C2 ˃ C1 , The two
electrodes are connected externally by a metallic wire and
internally through a salt bridge.
The cell is represented as Ag/Ag+ (C1)// Ag+/Ag (C2)/Ag
Electrode reaction:
At anode : Ag → Ag + (C1) + e-
At cathode : Ag + (C2) + e- → Ag
EMF of the cell:
Ecell = Ecathode – Eanode
0.0591 0.0591
= 𝐸 0𝑐𝑒𝑙𝑙 + log[C2] − 𝐸 0 𝑐𝑒𝑙𝑙 + log[C1]
𝑛 𝑛
= 0.0591 0.0591
𝐸 0 𝑐𝑒𝑙𝑙 + log[C2] − 𝐸 0 𝑐𝑒𝑙𝑙 − log[C1]
𝑛 𝑛
0.0591 [C2]
Ecell = 𝑛
log
[C1]
2
, When C2=C1, Ecell=0 Hence no net current flows through the cell.
When C2˃C1 Ecell is positive. Hence net cell reaction is spontaneous
Higher the C2/C1 ratio , higher is the cell potential
REFERENCE ELECTRODES may be defined as an electrode whose electrode potential value is
either arbitrarily fixed or exactly known, using which it is possible to determine the potential of the
electrodes.
Classification of reference electrodes
There are two types of reference electrodes viz.,
(1) Primary reference electrode Ex: Standard hydrogen gas electrode whose potential is fixed as
zero and
(2) Secondary reference electrode Ex: Calomel electrode and Ag-Agcl electrode.
Limitations of standard hydrogen electrode:
1. It is difficult to maintain the pressure of hydrogen gas uniformly at 1 atmosphere
2. It is difficult to maintain the H+ ions concentration at 1M throughout.
`CONSTRUCTION AND WORKING OF CALOMEL ELECTRODE:
In order to overcome the practical difficulties in setting
primary reference electrode, SHE, calomel electrode is devised
and since potential of calomel electrode is determined with
respect to SHE, it is called as secondary reference electrode.
Calomel electrode consists of a glass vessel containing
mercury at its bottom and is covered with a paste of Mercurous
chloride (Hg2Cl2) called calomel. The vessel is then filled with a
saturated solution of KCl and it is provided with a platinum wire
which dips in mercury for establishing electrical connection. It has a side U-tube which acts as salt
bridge to combine this electrode with other electrode to form an electrochemical cell.
It is represented as Hg, Hg2Cl2(s) / KCl(aq)
Electrode reaction: Hg2Cl2 + 2e 2 Hg+ + 2Cl-
According to Nernst equation, electrode potential is given by the expression
E = E0 – 0.0591 log [Cl-]
The potential of calomel electrode depends on the concentration of KCl. The potential of calomel
electrode w.r.t. SHE at 250C for various concentrations of KCl are
0.1 N KCl solution E = 0.334 V
1 N KCl solution E = 0.280 V
3
ELECTRODE SYSTEM AND CORROSION SCIENCE
Electrode and Cells
Electrochemistry is the branch of chemistry which deals with the study of interconversion of
electrical energy and chemical energy and study of behaviour of electrolytes.
Chemical reactions that involved in the interconversion of electrical energy and chemical energy are
called as electrochemical reactions.
Electrochemical cell is a device which is used to interconvert chemical energy and electrical energy.
It consists of two electrodes and an electrolyte solution.
A metal rod which is dipped in a solution containing its own ions is called an electrode. The electrode
at which oxidation takes place (loss of electrons) is called anode and the electrode at which reduction
takes place (gain of electrons) is called cathode.
Ex : Zn in contact with ZnSO4 solution.
Cu in contact with CuSO4 solution.
Each electrode of the cell is called half cell.
A spontaneous chemical process is the one which can take place on its own and in such a
process Gibbs free energy of a system decreases.
Types of Electrochemical Cell
1. Galvanic cells
2. Electrolytic cells
Galvanic Cell
The galvanic cell converts chemical energy into electrical energy i.e, electricity can be obtained with
the help of redox reaction. The oxidation and reduction take place in two separate compartments.
Each compartment consists of an electrolyte solution and metallic conductor which acts as an
electrode. The compartment containing the electrode and the solution of the electrolyte is called half
cells. Ex: Daniel cell.
Electrolytic Cell
The electrolytic cell converts electrical energy to chemical energy. Here the electrodes are dipped in
an electrolytic solution containing cations and anions. On supplying current the ions move towards
electrodes of opposite polarity and simultaneous reduction and oxidation take place.
Ex: - Electroplating of Cu,Ni,Zn etc.
1
,Differences between Galvanic cell and Electrolytic cell:
Galvanic cell Electrolytic cell
1 Chemical energy is converted into Electrical energy is converted into chemical
electrical energy energy
2 Anode is -ve electrode & Anode is + ve electrode &
Cathode is +ve electrode Cathode is -ve electrode
3 It can be used as a source of energy It cannot be used as a source of energy
Concentration cells:
Concentration cells can be defined as electrochemical cells that consist of two half-cells
wherein the electrodes are the same, but they vary in concentration. As the cell as a whole strives to
reach equilibrium, the more concentrated half cell is diluted and the half cell of lower concentration
has its concentration increased via the transfer of electrons between these two half cells. Therefore,
as the cell moves towards chemical equilibrium, a potential difference is created.
The cell potential depends on the concentration of electrolyte. A concentration cell is
defined as a galvanic cell obtained by dipping the same electrode in the same electrolyte of
different concentrations.
Consider 2 silver rods dipped in AgNO3 solution of
different concentration say C1 and C2 such that C2 ˃ C1 , The two
electrodes are connected externally by a metallic wire and
internally through a salt bridge.
The cell is represented as Ag/Ag+ (C1)// Ag+/Ag (C2)/Ag
Electrode reaction:
At anode : Ag → Ag + (C1) + e-
At cathode : Ag + (C2) + e- → Ag
EMF of the cell:
Ecell = Ecathode – Eanode
0.0591 0.0591
= 𝐸 0𝑐𝑒𝑙𝑙 + log[C2] − 𝐸 0 𝑐𝑒𝑙𝑙 + log[C1]
𝑛 𝑛
= 0.0591 0.0591
𝐸 0 𝑐𝑒𝑙𝑙 + log[C2] − 𝐸 0 𝑐𝑒𝑙𝑙 − log[C1]
𝑛 𝑛
0.0591 [C2]
Ecell = 𝑛
log
[C1]
2
, When C2=C1, Ecell=0 Hence no net current flows through the cell.
When C2˃C1 Ecell is positive. Hence net cell reaction is spontaneous
Higher the C2/C1 ratio , higher is the cell potential
REFERENCE ELECTRODES may be defined as an electrode whose electrode potential value is
either arbitrarily fixed or exactly known, using which it is possible to determine the potential of the
electrodes.
Classification of reference electrodes
There are two types of reference electrodes viz.,
(1) Primary reference electrode Ex: Standard hydrogen gas electrode whose potential is fixed as
zero and
(2) Secondary reference electrode Ex: Calomel electrode and Ag-Agcl electrode.
Limitations of standard hydrogen electrode:
1. It is difficult to maintain the pressure of hydrogen gas uniformly at 1 atmosphere
2. It is difficult to maintain the H+ ions concentration at 1M throughout.
`CONSTRUCTION AND WORKING OF CALOMEL ELECTRODE:
In order to overcome the practical difficulties in setting
primary reference electrode, SHE, calomel electrode is devised
and since potential of calomel electrode is determined with
respect to SHE, it is called as secondary reference electrode.
Calomel electrode consists of a glass vessel containing
mercury at its bottom and is covered with a paste of Mercurous
chloride (Hg2Cl2) called calomel. The vessel is then filled with a
saturated solution of KCl and it is provided with a platinum wire
which dips in mercury for establishing electrical connection. It has a side U-tube which acts as salt
bridge to combine this electrode with other electrode to form an electrochemical cell.
It is represented as Hg, Hg2Cl2(s) / KCl(aq)
Electrode reaction: Hg2Cl2 + 2e 2 Hg+ + 2Cl-
According to Nernst equation, electrode potential is given by the expression
E = E0 – 0.0591 log [Cl-]
The potential of calomel electrode depends on the concentration of KCl. The potential of calomel
electrode w.r.t. SHE at 250C for various concentrations of KCl are
0.1 N KCl solution E = 0.334 V
1 N KCl solution E = 0.280 V
3
