# The Equilibrium Constant ```9.5
Electrolytic cells
2 hours
9.5.1
Describe, using a diagram, the essential components of an electrolytic cell.
9.5.2
State that oxidation occurs at the positive electrode (anode) and reduction occurs at the negative electrode
(cathode).
9.5.3
Describe how current is conducted in an electrolytic cell.
9.5.4
Deduce the products of the electrolysis of a molten salt.
The diagram should include the source of electric current and conductors, positive and negative electrodes, and the electrolyte.
Half-equations showing the formation of products at each electrode will be assessed.
Aim 8: This process (which required the discovery of electricity) has made it possible to obtain reactive metals such as aluminium
from their ores. This in turn has enabled subsequent steps in engineering and technology that increase our quality of life. Unlike iron,
aluminium is not prone to corrosion and is one material that is replacing iron in many of its applications.
MHR Chapter 19.3 p. 776 – 783
Electrolytic Cells
electrolyisis = splitting using electricity
Electrolytic cells
 use electric (DC) currect
 cause nonspontaneous redox reactions form products
Anode = oxidation
Cathode = reduction
(same as voltaic)
Describe how current is conducted in an electrolytic cell.
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Electrons flow from the anode to the cathode in the external circuit.
The DC cell pushes electrons to the cathode making it negatively charged.
The negative cathode donates electrons for reduction.
The DC cell attracts electrons from the anode making it positively charged.
The positive cathode accepts electrons from oxidation.
ex. electrolysis of water
reaction 2 H2O → 2 H2 + O2
Anode: water is oxidized
Cathode: water is reduced
OX
RED
2 H2O + 2e−
2 H2O
 H2 + 2OH−
 O2 + 4H+ + 4e−
E0ox = − 0.42 V
E0red = − 0.82 V
NET
2 H2O
 2 H 2 + O2
E0net = − 1.66 V
 Hydrogen gas is evolved at the anode
 Oxygen gas is evolved at the cathode
Deduce the products for the electrolysis of a molten salt.
ex. Electrolysis of molten sodium chloride (Down’s cell)
reaction 2 NaCℓ(ℓ)  2 Naℓ) + Cℓ2(g)
Anode:
Cathode:
molten chloride ions are oxidized
molten sodium ions are reduced
OX
RED
2 Cℓ−(ℓ)  Cℓ2 (g) + 2e-2 Na+(ℓ) + 2e−  2Na(ℓ)
 Chlorine gas is evolved at the anode
 Molten sodium metal collects at the cathode
Anode = oxidation
Cathode = reduction
electrode gains electrons
electrode loses electrons
negative electrode
positive electrode
positive electrode
negative electrode
accepts electrons for oxidation
provides electrons for reduction
Voltaic Cell
Electrolytic Cell
Electroplating
the application of a thin coating of one metal oven another by electrolysis
Zn(s) + CuSO4(aq) + 1.10V  ZnSO4(aq) + Cu(s)
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