This document outlines learning outcomes for a chapter on electrochemistry. It describes key concepts such as classifying substances as conductors or non-conductors, different types of electrical conduction, electrolysis, oxidation and reduction reactions, and industrial applications of electrolysis. The document also defines important terms like electrolyte, cathode, anode, cation, anion, and Faraday's constant. It explains how to calculate masses and volumes of substances liberated during electrolysis.
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C12 electrochemistry
1. LEARNING OUTCOMES
Electrochemistry
Chapter 12
Describe investigations leading to the classification of substances as
conductors or non-conductors
Distinguish between metallic and electrolytic conduction
Define electrolysis, cathode, anode, cation and anion
Define electrolytes as strong and weak based on their conductivity
Predict the electrode to which an ion will drift
Define oxidation and reduction reactions including reactions at
electrodes
Predict chemical reactions making use of electrochemical series
Identify ions present in electrolytes
Discuss the electrolysis of certain substances
Define Faraday constant
Calculate the masses and volumes of substances liberated during
electrolyses
Describe industrial applications of electrolysis
2. Electrochemistry
Chapter 12
Conductors and electrolytes
The movement of electrically charged particles through a medium
constitutes an electric current. The medium of conduction is known
as a conductor.
Electrolytes can conduct electricity in their molten and aqueous
states. E.g. Metals
Non-conductors like plastic, ceramic and rubber can be used as
insulation materials.
Eg. The plastic coating on copper wires
3. Electrochemistry
Chapter 12
Electrical conductions in conductors
Electrical conductions in conductors can be explained through how the atoms are
bonded.
In metals, the electrons become delocalised due to overlapping orbitals of the
tightly packed atoms in the crystal lattice.
The valence electrons of each atom are loosely held as they are relatively distant
from the nucleus. Thus they can be easily separated and move about randomly
throughout, conducting electricity.
The only non-metal that conducts electricity is graphite.
In graphite, not all electrons are used in bonding. Thus, the free electrons can
move along the layers and carry the electric current.
4. The ionic theory proposes that it is the presence of freely-moving
charged particles called ions in the electrolyte that allows electrolysis
to occur.
These ions are only mobile when the electrolytes are in molten or
aqueous state.
These free moving ions arise only when an ionic solid melts or
dissolves in water or when certain polar covalent compounds dissolve
in water and their molecules ionise (dissociate into freely-moving
ions).
Electrochemistry
Chapter 12
The ionic theory
5. What is Electrolysis ?
Electrolysis is a process by which a substance is
broken up into its components by the passage of
electricity through it.
The substance must be an ionic compound and must
be molten or dissolved in water in order for the ions to
be mobile.
A direct current must be used for electrolysis.
During this process, electrical energy is changed into
chemical energy.
Electrochemistry
Chapter 12
6. One electrode is connected to the positive terminal of the battery. It is
called the anode. Oxidation occurs here.
The other electrode is connected to the negative terminal of the
battery. It is called the cathode. Reduction occurs here.
An electrolyte is a substance that is being electrolysed.
An electrolyte is able to conduct electricity due to the presence of
mobile ions.
Positive ions are known as cations.
Negative ions are known as anions.
Electrochemistry
Chapter 12
Electrolysis
7. At the anode,, Cl -
ions give away electrons to become Cl2 gas. We say
that Cl-
ions are discharged.
Na+
(l) + e-
Na(l)
Molten sodium chloride contains Na+
and Cl-
ions.
The Na+
ions are attracted to the cathode, while the Cl-
ions are attracted
to the anode.
At the cathode, Na+
ions take in electrons to become Na atoms. We say
that the Na+
ions are discharged.
2Cl-
(l) Cl2(g) + 2e-
Overall reaction:
2NaCl(l) 2Na(l) + Cl2(g)
Electrolysis of molten sodium chloride
Electrochemistry
Chapter 12
8. When electricity passes through the molten ionic compound, the positive ions will
migrate to the negative electrode, also known as cathode, while the negative ions
will migrate to the positive electrode, known as anode.
At the cathode, the positive ions will gain electrons and become metal atoms while
the negative ions will lose electrons at the anode and become non-metallic atoms.
Electrochemistry
Chapter 12
Electrolysis of Molten Ionic Compounds
At the Cathode
Sodium ions gain electrons and become sodium atoms, sodium ions are discharged.
Na+
(l) + e-
Na (l)
At the Anode
Chloride ions lose electrons and become chlorine gas, chloride ions are discharged.
2Cl-
(l) Cl2 (g) + 2e-
9. Electrolysis of other molten compounds
When a molten binary ionic compound is electrolysed,
the metal is always produced at the cathode and non-
metal is produced at the anode.
Electrochemistry
Chapter 12
10. Quick check 1
1. What is meant by (i) cathode, (ii) anode ?
2. What is meant by an electrolyte? What type of compounds must
electrolytes be?
3. During electrolysis, to which electrode do:
(a) the positive ions of the electrolyte move to;
(b) the negative ions of the electrolyte move to?
4. Predict the products formed when the following substances (in the
molten state) are electrolysed.
Compound Product at Anode Product at Cathode
Potassium chloride, KCl
Calcium fluoride, CaF2
Solution
Electrochemistry
Chapter 12
11. 1. (i) Cathode: the electrode connected to the negative terminal of the
battery.
(ii) Anode: the electrode connected to the positive terminal of the
battery.
2. An electrolyte is a substance which conducts electricity when molten or
dissolved in water. Electrolytes must be ionic compounds.
3. (a) The positive ions move to the cathode.
(b) The negative ions move to the anode.
4.
Compound Product at Anode Product at Cathode
Potassium chloride, KCl Chlorine, Cl2 Potassium, K
Calcium fluoride, CaF2 Fluorine, F2 Calcium, Ca
Solution to Quick check 1
Return
Electrochemistry
Chapter 12
12. When solutions are electrolysed, gases are usually produced.
The gases produced can be collected in test tubes to be identified later
by simple tests.
The electrolysis of solutions is
more complicated than
electrolysis of molten
compounds, because the
products at the electrodes can
come from the electrolyte as well
as from water.
Electrolysis of solutions
Electrochemistry
Chapter 12
13. At the cathode:
At the anode:
Selective discharge of ions
Positive ions from the electrolyte are discharged if they are H+
ions or ions of less reactive metals
such as Cu2+
, Pb2+
or Ag+
.
Positive ions of reactive metals such as Na+
, K+
and Ca2+
are not discharged in the presence of
water. Instead, H+
ions from water are discharged and H2 gas is produced.
Negative ions from the electrolyte are discharged if they are halide
ions such as I-
, Br -
, and Cl-
.
SO4
2-
and NO3
-
ions are not discharged. Instead, OH-
ions from water
are discharged and O2 gas is produced.
Electrochemistry
Chapter 12
14. The table shows the electrode products from solutions of ions, assuming the solutions
are fairly concentrated.
Cation Product at
Cathode
Ease of
discharge
Anion Product at
Anode
K+
Na+
Ca2+
Mg2+
Al3+
Ni2+
Pb2+
(H+
)
Cu2+
Ag+
Hydrogen from
water
Lead
Copper
Silver
Difficult
Easy
SO4
2-
NO3
2-
Cl-
Br-
I-
(OH-
)
Oxygen from
water
Chlorine
Bromine
Iodine
Oxygen
Selective discharge of ions
Electrochemistry
Chapter 12
15. Pure water will not conduct electricity, so some dilute sulphuric
acid is added to make it conduct electricity.
The electrolysis of dilute sulphuric acid is therefore essentially
the same as the electrolysis of water.
Electrolysis of dilute H2SO4
Electrochemistry
Chapter 12
16. OH-
ions are discharged in preference over SO4
2-
. Oxygen gas is formed.
2H+
(aq) + 2e-
H2(g)
Dilute sulphuric acid contains H+
, SO4
2-
and OH-
ions.
H+
ions take in electrons to become H2 molecules; H+
ions are discharged:
4OH-
(aq) O2(g) + 2H2O(l) + 4e-
Overall reaction:
2H2O(l) 2H2(g) + O2(g)
At the cathode: 2 volumes of hydrogen are produced.
At the anode: 1 volume of oxygen is produced.
Electrolysis of dilute H2SO4
At the cathode:
At the anode:
Electrochemistry
Chapter 12
The electrolysis of dilute solutions is essentially the same as the electrolysis of
water.
17. The ions present in the solution are: Na+
, Cl-
(from sodium
chloride) and H+
, OH-
(from water).
Both Na+
and H+
are attracted here, but due to their
relative positions in the reactivity series, H+
ions are preferentially
discharged:
2H+
(aq)+ 2e-
H2(g)
Both Cl-
and OH-
are attracted here, but due to the lower
position of the hydroxide ions in the reactivity series, they are
preferentially discharged:
4OH-
(aq) O2(g) + 2H2O(l) + 4e-
Overall reaction:
2H2O(l) 2H2(g) + O2(g) [ Electrolysis of water ]
Electrolysis of dilute sodium chloride solution
At the cathode:
At the anode:
Electrochemistry
Chapter 12
18. The ions present in the solution are:
Na+
, Cl-
and H+
, OH-
Both Na+
and H+
are attracted here, but due to their
relative positions in the reactivity series, H+
ions
are preferentially discharged:
2H+
(aq) + 2e-
H2(g)
Both Cl-
and OH-
are attracted here, but due to the high
concentration of the chloride ions, chloride ions are
preferentially discharged:
2Cl-
(aq) Cl2(g) + 2e-
Overall reaction:
Hydrogen gas is produced at the cathode and chlorine gas is
produced at the anode.
The Na+
and OH-
ions left in the solution combine to form sodium
hydroxide, thus making the solution alkaline.
Electrolysis of concentrated sodium chloride solution
At the cathode:
At the anode:
Electrochemistry
Chapter 12
19. Examples of electrolysis of different solutions using inert electrodes,
assuming the solutions are fairly concentrated
Electrolyte Ions in Solution
Product at
Cathode
Product at
Anode
Aq. sodium chloride
Na+
(aq), Cl-
(aq) ,
H+
(aq), OH-
(aq)
hydrogen gas chlorine gas
Aq. hydrochloric
acid
H+
(aq), Cl-
(aq),
H+
(aq), OH-
(aq)
hydrogen gas chlorine gas
Aq. copper(II)
sulphate
Cu2+
(aq), SO4
2-
(aq),
H+
(aq), OH-
(aq)
copper metal oxygen gas
Electrolysis of Solutions
Electrochemistry
Chapter 12
20. Inert and reactive electrodes
Inert electrodes do not react with the product
produced or dissolved in the electrolyte.
Carbon and platinum are examples of inert electrodes.
Reactive electrodes can react or dissolve in the electrolyte.
Copper, silver and mercury are examples of reactive
electrodes.
Electrochemistry
Chapter 12
21. Neither SO4
2-
nor OH-
ions are discharged.
Instead the copper anode dissolves in the
solution and produces electrons:
Cu Cu2+
+ 2e-
During the electrolysis the total concentration of the
CuSO4 solution remains unchanged. The cathode
increases in mass while the anode decreases in mass
proportionately.
Electrolysis of CuSO4 solution
Copper
anodeCopper
cathode
Aq. copper(II) sulphate
_ +
Cu2+
H+
SO4
2
-
OH-
Using copper (reactive) electrodes
The ions present in the solution are: Cu2+
, SO4
2-
, H+
and OH-
.
Cu2+
ions are discharged in preference over the H+
ions:
Cu2+
+ 2e-
Cu
At the cathode:
At the anode:
Electrochemistry
Chapter 12
22. Electrolysis of CuSO4 solution
Using copper (reactive) electrodes
During the electrolysis the concentration of the H+
and
SO4
2-
ions increases, and hence the solution becomes
more acidic.
This process essentially transfers
copper metal from the anode to
the cathode.
This process is used in the
industry for the purification of
impure copper to obtain pure copper.
Electrochemistry
Chapter 12
23. Quick check 2
1. Place in order the ease of discharge of the following cations (starting from the easiest
first): Ca2+
, Na+
, H+
, Al3+
, Mg2+
, K+
, Cu2+
, Pb2+
, Ag+
.
2. Place in order the ease of discharge of the following anions (starting from the easiest
first): Br-
, Cl-
, NO3
-
, OH-
, I-
, SO4
2-
,.
3. (a) State the products obtained when a solution of dilute sulphuric acid is electrolysed
using platinum electrodes.
(b) Write the ionic equations for the reactions taking place at the cathode and anode.
4. (a) State the products obtained when a concentrated solution of sodium chloride is
electrolysed using inert electrodes.
(b) Write the ionic equations for the reactions taking place at the cathode and anode.
5. (a) State the products obtained when a solution of sodium hydroxide is electrolysed
using platinum electrodes.
(b) Write the ionic equations for the reactions taking place at the cathode and anode.
Solution
Electrochemistry
Chapter 12
24. 1. Ag+
, Cu2+
, H+
, Pb2+
, Al3+
, Mg2+
, Ca2+
, Na+
, K+
2. OH-
, I-
, Br-
, Cl-
, NO3
-
, SO4
2-
3. (a) At cathode: hydrogen; At anode: oxygen
(b) At cathode: 2H+
(aq) + 2e- H2(g)
At anode: 4OH-
(aq) 2H2O(l) + O2(g)
4. (a) At cathode: hydrogen; At anode: chlorine
(b) At cathode: 2H+
(aq) + 2e- H2(g)
At anode: 2Cl-
(aq) Cl2
5. (a) At cathode: hydrogen; At anode: oxygen
(b) At cathode: 2H+
(aq) + 2e- H2(g)
At anode: 4OH-
(aq) 2H2O(l) + O2(g) Return
Solution to Quick check 2
Electrochemistry
Chapter 12
25. The Faraday constant, F, is the quantity of electricity carried by one
mole of electrons and is equivalent to 96500 Cmol-1
.
The amount of a substance deposited on each electrode of an
electrolytic cell is directly proportional to the quantity of electricity
passing through the cell.
The quantity of electricity contained in a current running for a
specified time can be calculated by Q= l x t
The quantity of electricity required to deposit an amount of metal can
be calculated by Q = n(e) x F
Electrochemistry
Chapter 12
Faraday’s Law of Electrolysis
27. Cu2+
ions are discharged as copper metal is deposited on the object:
Cu(s) Cu2+
(aq) + 2e-
The electrolyte is copper(II) sulphate (CuSO4)solution.
The object to be plated is made the cathode; copper is made the anode.
Cu2+
(aq) + 2e Cu(s)
There is a net transfer of copper from the anode to the cathode.
The concentration of the CuSO4 solution remains unchanged.
Electroplating with copper
At the cathode:
At the anode:
Copper dissolves:
Electrochemistry
Chapter 12
28. Anodising is the process of making the oxide layer on the
surface of the aluminium thicker. This will protect the aluminium
even better.
The aluminium object is made the anode while the cathode
could be copper or lead or aluminium.
When current is applied, the water in the electrolyte breaks
down and oxygen is deposited at the anode. This oxygen then
combines with the aluminium to form aluminium oxide and thus
provides a protective layer for the aluminium. This prevents
corrosion.
Anodising
Electrochemistry
Chapter 12
29. Electroplating with other metals
Many metallic objects can be electroplated in
the same way:
The object is made the cathode and the metal used for
electroplating is made the anode. The electrolyte is a
solution of ions of the metal used for the plating.
Electroplating can be used to protect iron objects from
corrosion by covering it with a less reactive metal like
chromium or copper.
Electroplating also makes the object more attractive and
increases its value e.g. plating it with gold, silver and
platinum.
Electrochemistry
Chapter 12
30. Electroplating Metal Uses
Chromium
Water taps, motorcar bumpers and bicycle
parts
Tin Tin cans
Silver
Silver-plated sports trophies, plaques,
ornaments, knives and forks
Gold Gold-plated watches, plaques, ornaments
Examples of electroplating
Electrochemistry
Chapter 12
31. 1. State two uses of electrolysis in the industry.
2. (a) What is electroplating?
(b) State two advantages of electroplating
an iron object with chromium.
3. A metal spoon is to be coated with silver. Sketch a diagram to
show how you would set up the electrolytic cell for this to be
carried out. Label the material that can be used for the
cathode and anode. Also state a suitable solution for the
electrolyte.
Quick check 3
Solution
Electrochemistry
Chapter 12
32. 1. (i) To extract reactive metals like sodium, magnesium
and aluminium from their ores;
(ii) To electroplate metallic objects with less reactive
metals for attractiveness and protection from corrosion.
2. (a) Electroplating is the coating of a more reactive metal
with a less reactive metal by electrolysis.
(b) Electroplating with chromium protects the
iron from corrosion and the silver colour of the
chromium improves its appearance.
Solution to Quick check 3
Return
Electrochemistry
Chapter 12
33. Return
Electrochemistry
Chapter 12
3. Cathode: the metal object; Anode: silver metal
Electrolyte: silver nitrate solution.
Solution to Quick check 3 (cont’d)
Electroplating with
silver
