22CT103 - ENGINEERING CHEMISTRY
Unit-2: Electrochemical Devices
Prof. Koya Prabhakara Rao
HoD, Department of Chemistry
E-mail:drkpr_sh@vignan.ac.in
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Contents
Introduction to Electro-chemistry:
Redox reactions, Electrode potential,
Galvanic Cell
EMF of an electrochemical cell,
Electrochemical series
Batteries: Classification of batteries and characteristics
Construction, working and applications of: Lead-acid
storage cell, Lithium-ion battery.
Fuel Cells- Classification, Construction, working and
applications of H2-O2 fuel cell.
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Introduction
Electro Chemistry: The branch of chemistry that deals with the relations between electrical
and chemical phenomena.
Electro motive force (EMF): Electro motive force (EMF) is the difference of potential produced
by sources of electrical energy, which can be used to drive current through electrical circuit.
Redox reactions: Oxidation is a process, which involves loss of electrons by a substance;
while reduction is a process which involves gain of electrons by a substance. Oxidation and
reduction reaction occurs side by side is called Redox reactions.
Electrochemistry involves two main types of processes ELECTRODE POTENTIAL
A. Voltaic(galvanic) cells – which are spontaneous chemical
reactions (battery)
B. Electrolytic cells – which are non-spontaneous and require
external e− source (DC power source)
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electrode potential of a metal is the 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 unit molar concentration at
25oC.
The tendency of an electrode to lose electrons is direct measure of its tendency to get
oxidized; and this tendency is called oxidation potential.
The tendency of an electrode to gain electrons is a direct measure of its tendency to get
reduced; and this tendency is known as reduction potential.
EMF OF AN ELECTRO CHEMICAL CELL
Ecell = Eright - Eleft
where Ecell = e.m.f. of the cell;
Eright = reduction potential of right hand side electrode;
and Eleft = reduction potential of left hand side electrode.
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Voltaic Cells (Galvanic cell )
We can use that energy to do work if we make the electrons flow through an external device,
such setup is called a voltaic cell. A typical cell looks like this.
The oxidation occurs at the anode.
The reduction occurs at the cathode.
The flow of electrons is always from the anode to the
cathode through the wire
Once even one electron flows from the anode to the
cathode, the charges in each beaker would not be
balanced and the flow of electrons would stop.
Therefore, we use a salt bridge, usually a U-shaped
tube that contains a salt solution, to keep the
charges balanced.
Cations move toward the cathode.
Anions move toward the anode.
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Electrochemical series:
The electrode potentials of different electrodes can be
finding using standard hydrogen electrode. The potential
of hydrogen electrode is assumed as zero volts. So the
measured EMF. Itself is the standard electrode potential
of that electrode.
The arrangement of different electrode potentials of
different electrodes from highest -ve to highest +ve are
called electrochemical series.
Standard Reduction Potentials
• The larger the difference between Ered
values, the larger Ecell.
• In a voltaic (galvanic) cell (spontaneous)
Ered(cathode) is more positive than
Ered(anode).
Ecell = Ered(cathode) - Ered(anode)
Note Cell Diagram: Anode || Cathode
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Electro chemical cells or batteries
An electrochemical cell (Battery) is a device capable of either deriving electrical energy
from chemical reactions, or facilitating chemical reactions through the introduction of electrical
energy. A common example of an electrochemical cell is a standard 1.5-volt "battery“.
Types of Electro chemical cells:
1. Primary cells (or batteries) -----In which the cell reaction is not reversible.
Eg: Dry or Laclanche cell 1. Nickel-Cadmium (Ni-Cd) Battery
2. Lead-Acid Battery
2. Secondary cells (or batteries)-------- 3. Lithium-Ion (Li-ion) Battery
In which the reaction can be reversed by passing 4. Sodium-Ion (Na-ion) Battery
direct electric current in opposite direction. 5. Lithium-Sulfur (Li-S) Battery
6. Sodium-Sulfur (Li-S) Battery
Eg: Lead-acid batteries 7. Zinc-oxygen (Zn-O2) Battery
Lithium Ion battery 8. Zinc Organic Battery
9. Redox Flow Battery
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3. Flow battery and fuel cells------- In which, materials (reactants, products, electrolytes) pass
through the battery, which converts chemical to electrical energy.
Fuel cells are different from batteries in that they require a constant source of fuel and oxygen/air
to sustain the chemical reaction, they can however produce electricity continually for as long as
these inputs are supplied.
Eg: Hydrogen-oxygen fuel cell
Methanol-oxygen fuel cell
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1. Primary cells (or batteries)
(Dry or Laclanche cell): The anode of the cell is zinc can (or container) containing an electrolyte
consisting of NH4Cl, ZnCl2, and MnO2 to which starch is added to make it thick paste-like so that it
is less likely to leak.
A carbon (graphite) rod serves as the cathode, which is immersed in the electrolyte in the center
of the cell.
Anode: Zn (S) → Zn2 (Aq) + 2e
Cathode: 2MnO2 (S) + 2NH4 (aq) + 2e → Zn(NH3)2Cl2 (s)
Net: : Zn (S) 2NH4 (aq) + 2Cl (aq) 2MnO2 (S) → Mn2O3 (S) + Zn(NH3)2Cl2 (s) + 2H2O (l)
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1859 Gaston Plante Lead-acid batteries:
At anode: Pb (S) → Pb2 + 2e
The Pb2 ions then combines with sulphate (SO42) ions.
Pb2 + SO42- → PbSO4
The electrons released from the anode (Lead plate) flows to the dioxide electrode. Here PbO2 gains
electrons to form Pb2 ions. In other words, lead undergoes reduction at cathode from oxidation
state of 4 to 2. The Pb2 ions then combine with (SO42) ions.
At cathode: PbO2 + 4H 2e- → Pb2 2H2O
Pb2 + SO42- → PbSO4
So the net reaction during use (or discharging) is:
Pb PbO2 + 4H +2SO42- → 2PbSO4 + 2H2O + Energy
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Charging: When both anode and cathode become covered with PbSO4, the cell ceases to function as a
voltaic cell.
To re-charge a lead storage cell the reactions taking place during discharging are reversed by passing an
external EMF greater than 2 volts from a generator.
The ve pole of the generator is attached to the ve pole of the cell (or battery) and the following reactions
take place:
Reactions at the ve terminal (cathode):
PbSO4 + 2e- → Pb SO42
Reactions at the ve terminal (anode):
PbSO4 + 2H2O 2e- → PbO2 4H SO42
Hence the net reaction during charging is:
2PbSO4 + 2H2O Energy → Pb PbO2 4H 2SO42
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Li-ion battery system
Electrochemical Reactions
• Cathode
c
LiCoO2 Li1-xCoO2 + xLi+ + x e-
d
• Anode
c
Cn + xLi+ + x e- CnLix
d
• Overall
c
LiCoO2 + Cn Li1-xCoO2 + CnLix
d
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Lithium-Ion Battery Charge
Electrolyte
Cu AL
Current Current
Collector Collector
Graphite LiMO2
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SEI
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Lithium-Ion Battery Discharge
Electrolyte
Cu AL
Current Current
Collector Collector
Graphite LiMO2
SEI SEI
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Lithium Batteries market
Lithium batteries: high energy density (3 times lead-
acid).
Power sources of choice for the consumer
electronics market
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Fuel cells
Fuel Oxygen Oxidation products Electricity
Hydrogen-oxygen fuel cell:
One of the simplest and most successful fuels is hydrogen-oxygen fuel cell.
It consists ---an electrolytic solution such as 25% KOH solution, and two inert porous
electrodes.
Anode: 2H2 (g) 4 OH (aq) → 4H2O (l) 4 e E0 -0.83
Cathode: O2 (g) 2H2O (l) 4 e → 4OH (Aq) E0 0.40
Net: 2H2 (g) O2 → 2H2O (l)
The standard emf of the cell, E0 E0red E0oxi 0.40V (-0.83V) 1.23 V.
Sir William Grove invented the first Fuel
Cell in 1838.
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The fuel (direct H2 or
reformed H2) undergoes
oxidation at
anode and releases
electrons.
These electrons flow
through the external circuit
to the cathode.
At cathode, oxidant (O2
from air) gets reduced.
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