Electrolysis
• Electrolysis is the passage of an
electrical current through a cell, causing
an otherwise non-spontaneous reaction
to occur.
• An electrolytic cell is the type of cell in
which electrolysis occurs. It typically
includes a molten salt or electrolyte
solution.
An Electrolytic Cell
Electrolysis in Aqueous Solutions
• In aqueous solutions, there is the
possibility that water may be oxidized or
reduced, rather than the solute.
• Whichever half-reaction has the more
positive voltage will be the one to occur.
• Oxidation half-reactions must be written
as oxidations, not as reductions.
Electrolysis in Aqueous NaCl
• The following two oxidation reactions are
possible:
2Cl-(aq) → Cl2(g) + 2eE = -1.36 V
2H2O(l) → 4H+(aq) + O2(g) + 4e- E = -1.23 V
• Because the oxidation of water has a
more positive potential, oxidation of
water is the expected reaction.
Electrolysis in Aqueous NaCl
• The following two reduction reactions are
possible:
Na+(aq) + e- → Na(s)
E = -2.71 V
2H2O(l) + 2e- → H2(g) + 2OH-(aq) E = -0.83 V
• Because the reduction of water has a
more positive potential, reduction of water
is the expected reaction.
Quantitative Aspects of Electrolysis
• The relationship between the electric
current and amount of charge used is
given by the equation
Q = It
where:
• Q = total charge in coulombs
• I = current in amperes
• t = time in seconds
Quantitative Aspects of Electrolysis
• The Faraday constant (96,500 C/mol e-)
is used to determine the number of
moles of electrons used.
• The stoichiometry of the oxidation or
reduction half-reaction is then used to
determine the number of moles of
reactant used or product made.
Quantitative Aspects of Electrolysis
• Calculate the mass of copper deposited
through the passage of 0.400 A through
a solution of Cu2+(aq) for 25.0 min.
Cu2+(aq) + 2e- → Cu(s)
 60 s 
Q  (0.400 A)(25.0 min) 
 600. C

 1 min 
 1 mol e   1 mol Cu  63.55 g Cu 
600. C 
 0.198 g Cu


- 
 96,500 C   2 mol e  1 mol Cu 
Test Your Skill
• What mass of silver is deposited by the
passage of 0.250 A through a solution of
silver nitrate for a period of 90.0 min?
Industrial Applications of Electrolysis
• Electrolysis is used commercially to
isolate the elements sodium, fluorine,
chlorine, and aluminum, and to purify
(electrorefine) copper.
• Electrolysis is used to electroplate
metals for decorative or protective
purposes.
The Hall Process for Aluminum
• Aluminum oxide ore is mixed with
cryolite (Na3AlF6) to produce a mixture
that melts at about 980 C.
• The molten mixture is electrolyzed with
carbon electrodes at about 4.2 A.
• Molten aluminum forms at the cathode.
The Hall Process for Aluminum
Corrosion
• Corrosion is the oxidation of a metal
through interaction with the environment.
• Rust is a hydrated form of iron oxide,
formed by the corrosion of iron metal:
Fe(s) → Fe2+(aq) + 2eE = +0.44 V
O2(g) + 4H+(aq) + 4e- → 2H2O(l)
E = +1.23 V
Corrosion
• Corrosion is pH-dependent because the
reduction potential of O2 is greater at low
pH.
Protection from Corrosion
• Metal can be protected from corrosion by
isolating it from water and oxygen, either
by painting or plating.
• Anodic protection is the formation of a
thin protective layer of oxide on the
surface of the metal. Iron can be
protected by the controlled formation of
iron and chromium oxides:
2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) →
Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)
Protection from Corrosion
• In cathodic protection, a second more
reactive metal is placed in electrical
contact with the piece to be protected.
The more reactive metal serves as a
sacrificial anode.
• For example, galvanized iron has a
protective layer of zinc that corrodes
preferentially.
Cathodic Protection