Electrochemical Cell, its components and types.

1. PRESENTATION TOPIC
“ELECTROCHEMICAL CELLS’’
Saman Tanoli
Department: Chemistry

2. “ELECTROCHEMICAL CELLS’’
CONTENTS:
 DEFINITION
 COMPONENTS
 TYPES
 VOLTAIC or GALVANIC CELLS
 ELECTROLYTIC CELLS
 DIFFERENCE BETWEEN GALVANIC & ELECTROLYTIC CELLS
 APPLICATIONS OF ELECTROCHEMICAL CELLS

3. DEFINITION;
An electrochemical cell is a device that can
generate electrical energy from the chemical reactions
occurring in it, or use the electrical energy supplied to it to
facilitate chemical reactions in it. These devices are capable of
converting chemical energy into electrical energy, or vice
versa..
EXAMPLE;
A common example of an electrochemical cell is a
standard 1.5-volt cell which is used to power many electrical
appliances such as TV remotes and clocks.

4. COMPONENTS;
Here is the list of the all the components.
Two half cells
Two metal electrodes
One voltmeter
One salt bridge
Two aqueous solutions for each half cell
All of these components create the Electrochemical
Cell.

5. HALF CELLS;
A Voltaic Cell (also
known as a Galvanic Cell) is an
electrochemical cell that uses
spontaneous redox reactions
to generate electricity. It
consists of two separate half-
cells. A half-cell is composed of
an electrode (a strip of metal,
M) within a solution containing
Mn+ ions in which M is any
arbitrary metal. The two half
cells are linked together by a
wire running from one
electrode to the other. A salt
bridge also connects to the half
cells.

6. TWO METAL ELECTRODES;
Electrochemical cells have two conductive electrodes, called the anode and the
cathode. The anode is defined as the electrode where oxidation occurs. The cathode is the
electrode where reduction takes place.
The key features of the cathode and the anode are tabulated below.
CATHODE
 Denoted by a positive sign
since electrons are consumed
here
 A reduction reaction occurs in
the cathode of an
electrochemical cell
 Electrons move into the
cathode
ANODE
 Denoted by a negative sign since
electrons are liberated here
 An oxidation reaction occurs
here
 Electrons move out of the anode

7. VOLTMETER;
DEFINITION;
A voltmeter is an instrument
used for measuring electric potential
difference between two points in an
electric circuit.
Electric potential is
the potential energy per charge. The
concept of electric potential is used
to express the effect of
an electric field of a source in terms
of the location within
the electric field.

8. SALT BRIDGE;
DEFINITION;
A salt bridge is a device
used in an electrochemical cell
for connecting its oxidation and
reduction half cells wherein a
weak electrolyte is used.
In other words, a salt bridge is a
junction that connects the
anodic and cathodic
compartments in a cell or
electrolytic solution.

9. TYPES OF ELECTROCHEMICAL CELLS;
Electrochemical Cells are of two types.
1. Voltaic or Galvanic Cells
2. Electrolytic Cells

10. VOLTAIC CELL
 A voltaic cell, often known as a
galvanic cell, provides
electrical energy. The source
of this energy is a
spontaneous chemical
reaction, more specifically a
spontaneous redox reaction.
 The energy conversion is
achieved by spontaneous (ΔG
< 0) redox reactions producing
a flow of electrons.
ELECTROLYTIC CELL
 In an electrolytic cell,
electrical energy is used to
drive a non-spontaneous
chemical reaction.
 A flow of electrons drives non-
spontaneous (ΔG ≥ 0) redox
reactions.

11. CONTINUED……………..
VOLTAIC CELL
 For example, all batteries are
made of one or more voltaic
cells; batteries go flat when
most or all of their reactants
have been converted to
products, transforming their
chemical potential energy to
electrical energy.
 A battery powering something
is an example of a galvanic
cell.
ELECTROLYTIC CELL
 For example, water can be
split into hydrogen and oxygen
in an electrolytic cell. Also,
when a rechargeable battery is
recharged, it operates as an
electrolytic cell.
 Examples of electrolytic cells
also include those that convert
aluminum ore to aluminum
metal.

12. Basic Voltaic Cell Systematic
 Spontaneous redox reactions at the
electrodes produce a voltage.
Correctly set up, this voltage can
drive electrons through electric
devices, such as the light bulb
shown here. In this diagram,
species transfer electrons to the
anode from where they flow
through the light bulb to the
cathode, where they bring about
reduction.

13.  Non-spontaneous redox
reactions are driven by an
external voltage. The
electrolytic cell's processes
are the opposite of the voltaic
cell's. The current from the
power source pushes
electrons on to the cathode,
where they cause reduction of
species to take place.
Basic Electrolytic Cell Systematic

14. Galvanic Cell (aka Voltaic Cells)
A galvanic cell produces
an electrical charge from
the flow of electrons. The
electrons move due to the
Redox reaction. As we can
see, Zn oxidizes to Zn2+ ,
while Cu2+ reduces to Cu.
In order to understand
the redox reaction, Solve
the Redox equation.

15. First, split the reaction into two half reactions,
with the same elements paired with one
another.
Zn(S) → Zn+2
(aq) Oxidation Reaction: takes place at the Anode
Cu+2
(aq) → Cu(s) Reduction Reaction: Takes place at the Cathode

16. Next, we balance the two equations.
Oxidation: Zn(S) →Zn2+
(aq) + 2e- (Anode)
Reduction: 2e- + Cu2+
(aq) → Cu(s) (Cathode)
Spontaneous redox reaction releases energy; The system does work on the
surroundings.

17. Electrolytic Cell;
An electrolytic cell is a cell which requires
an outside electrical source to initiate the
redox reaction. The process of how
electric energy drives the non-
spontaneous reaction is
called electrolysis. Whereas the galvanic
cell used a redox reaction to make
electrons flow, the electrolytic cell uses
electron movement (in the source of
electricity) to cause the redox reaction. In
an electrolytic cell, electrons are forced to
flow in the opposite direction. Since the
direction is reversed of the voltaic cell,
the E0
cell for electrolytic cell is negative.
Also, in order to force the electrons to
flow in the opposite direction, the
electromotive force that connects the two
electrode-the battery must be larger than
the magnitude of E0
cell. This additional
requirement of voltage is
called overpotential.

18. Electrolytic cell for the example above:
Oxidation: Cu(s) → Cu2+ (aq)+2e- (anode)
Reduction: Zn2+ (aq)+2e- → Zn(s) (cathode)
(Nonspontaneous redox reaction absorbs energy to
drive it; The surroundings do work on the system. )

19. The most common form of Electrolytic cell is
the rechargeable battery (cell phones, mp3's,
etc) or electroplating. While the battery is
being used in the device it is a galvanic cell
function (using the redox energy to produce
electricity). While the battery is charging it is
an electrolytic cell function (using outside
electricity to reverse the completed redox
reaction).

20. Differences Between Galvanic Cell &
Electrolytic Cell;
Galvanic Cell
 In galvanic cell, electrical energy is
produced.
 In galvanic cell, reaction taking place is
spontaneous.
 The two half cells are set up in different
containers and are connected through
salt bridge or porous partition.
 In galvanic cell, anode is negative and
cathode is positive.
 The electrons move from anode to
cathode in external circuit.
Electrolytic Cell
 In electrolytic cell, electrical energy is
consumed.
 In electrolytic cell, reaction taking place
is non-spontaneous.
 Both the electrodes are placed in the
solution or molten electrolyte in the
same container.
 In electrolytic cell, the anode is positive
and cathode is negative.
 The electrons are supplied by the
external source. They enter through
cathode and come out through anode.

21. Applications of Electrochemical Cells;
 Electrolytic cells are used in the electrorefining of many non-ferrous
metals. They are also used in the electro winning of these metals.
 The production of high-purity lead, zinc, aluminium, and copper
involves the use of electrolytic cells.
 Metallic sodium can be extracted from molten sodium chloride by
placing it in an electrolytic cell and passing an electric current
through it.
 Many commercially important batteries (such as the lead-acid
battery) are made up of Galvanic cells.
 Fuel cells are an important class of electrochemical cells that serve
as a source of clean energy in several remote locations.

22. THANK YOU