Introduction to Electrochemical Cells

1. Electrochemical
Cells

2. Energy
 Energy is the capacity to do work
 SI unit is Joule (J)
 Law of conservation of energy
- energy cannot be created or destroyed but it
can be converted from one form to another

3. Electrochemical cells
Chemical energy- electrical energy conversions take place
and can go in either direction
Voltaic/galvanic cell
Chemical energy is converted into electrical energy
From spontaneous, exothermic chemical processes to
electrical energy
Electrolytic cells
Electrical energy is converted into chemical energy
Non-spontaneous process

4. Electrodes
 In metals, electric charge are carried by
electrons
 Electrode is conductor of electricity used to
make contact with a non-metallic part in a
circuit such as the electrolyte
Voltaic cell Both Electrolytic cell
Positive electrode
CATHODE
Reduction take place
Negative electrode
Negative electrode
ANODE
Oxidation take place
Positive electrode

5. The Voltaic Cell
 Consist of two half-cells
 Oxidation occurs at one half cell (anode) and
reduction occurs at the other half-cell (cathode)
 Different types of electrode used in voltaic cell such
as metal/metal ions electrode, metal ion in two
different oxidation state and the gas ion electrode

6. Metal/metal ion electrode
 Consists of a bar of metal dipped into a
solution containing cations of the same metal
 Eg: Fe (s)l Fe2+ (aq)
 In a voltaic cell, the two half cells are
separated. If the solutions were allowed to
mix in a single container, a spontaneous
reaction would occur but there would be no
movement of electron through the external
circuit and hence no current

7. Salt bridge
 It allows physical separation of the cathode and
the anode, preventing mixing of two solutions
 It provides electrical continuity (migration of the
positive ion (cation) and the negative ion (anion))
 Reduces the liquid junction potential. This is
the voltage generated when two different
solutions come into contact with each other,
which occurs due to unequal cation and anion
migration across the junction.

8.  Salt bridge contains a concentrated solution of a
strong electrolyte.
 The high concentration of ions in the salt bridge
allows ions to diffuse out of it
 For example, the Daniell Voltaic Cell.
 Typical compounds used in the salt bridge for this
cell could be sodium sulfate or potassium chloride
 The ions in the salt bridge must be inert. They should
not react with the other ion in the solution

9. The Daniell Voltaic Cell
 At the anode (negative electrode)l oxidation
Zn (s) → Zn2+(aq) + 2e-
 Cathode (positive electrode)l reduction
Cu2+ (aq) + 2e- → Cu (s)
 Overall cell reaction
Cu2+ (aq) + Zn (s) → Zn2+(aq) + Cu (s)

10. Observation
 Blue colour of the copper(II) sulfate solution fades
 The copper bar increases in size as it become coated
in more copper
 The zinc bar gets thinner

12. Electrolytic cells
 Electrolysis is the process by which electrical
energy is used to drive a non-spontaneous
chemical reaction
 Consists of a single container, two electrodes
(cathode and anode),a solution (electrolyte) and a
battery (an electron pump)

13. Electrolysis of a molten salt (lead II bromide)
 Anode (positive) (oxidation)
2Br- → Br2(g) + 2e-
 Cathode (negative) (reduction)
Pb2+(l) + 2e- → Pb (l)
 Overall cell reaction
PbBr2(l) → Pb (l) + Br2 (g)