STAGE 2 CHEMISTRY
Using and Controlling Reactions
READING
Unit 3.3
The Essentials book pg 166-170
Q 3.29, 3.30
Electrochemistry
Part 2: Electrochemical Cells: Electrolytic Cells
describe the main energy difference between the two types of electrochemical cells,
write balanced redox equations for the reactions occurring in an electrolytic cell,
label diagrams of a electrolytic cells,
explain the function of each part of an electrolytic cell,
describe, with equations, how electrolytic cells are used in the production of active
metals.
The Syllabus statement says:
Key Ideas
Intended Student Learning
Electrochemical cells are conveniently divided into
galvanic cells, which produce electrical energy from
spontaneous redox reactions, and electrolytic cells,
which use electrical energy from an external source
to cause a non-spontaneous chemical reaction.
Identify a cell as galvanic or electrolytic, given
sufficient information.
Redox reactions can be considered as two
half-reactions, one involving oxidation and the
other reduction.
Write half-equations for half-reactions,
including those in acidic solution, given
information about the reactants and the
products.
Galvanic and electrolytic cells involve oxidation
at the anode and reduction at the cathode, with
electrons being transferred from one electrode
to the other through an external circuit.
Identify the anode and the cathode in a
galvanic cell or an electrolytic cell, given
information about the reactants and the
products.
Electrolytic cells are used in the production of
active metals.
Describe, with the aid of equations, the
electrolytic production of active metals.
Page 1
Electrochemical Cells: A Review
The KEY IDEAS are:

Electrochemical cells are divided into galvanic cells that produce electricity from spontaneous chemical
reactions and electrolytic cells that use electricity from an external source to cause a non-spontaneous
chemical reaction.
In a galvanic cell:
In an electrolytic cell:
Electricity is produced from a spontaneous redox reaction.
Oxidation and reduction take place in separate compartments that are
linked by an external circuit or wire.
The current flow is due to the movement of electrons through the circuit
from the oxidation (loss of electrons) compartment to the reduction (gain of
electrons) compartment.
Electrical energy is used to force a chemical reaction to take place.
The process is called electrolysis.
The electrical energy source pushes electrons around a circuit to one
part of the cell where reduction (gain of electrons) occurs, and
electrons are pulled from another part of the cell where oxidation (loss
of electrons) occurs.

The chemical reactions in these cases are usually redox reactions involving two half-equations, one for
oxidation and the other reduction. You need to be able to balance redox half equations and then write
overall redox equations.

Both types of cells involve oxidation at the anode and reduction at the cathode, with electrons being
transferred from the anode to the cathode through an external wire.
Comparing galvanic cells and electrolytic cells

You will need to be able to compare the characteristic features of both types of cells and when given
information about an electrochemical cell, you will need to be able to:
identify it as a galvanic or an electrolytic cell
write half-equations for the redox reactions that occur and
identify all processes, eg movement of electrons, stating where they occur in the cell.
ELECTROCHEMICAL CELLS
The following metal Activity Series will be useful when working with electrochemical cells.
most reactive
least reactive
Page 2
ELECTROLYTIC CELLS
In an electrolytic cell, electrical energy
is converted into chemical energy.
The process is called electrolysis and
results in a chemical reaction caused
by the passage of a current through an
electrolyte.
The chemical reaction that occurs when electricity passes through a molten ionic compound or through an
electrolyte solution is called electrolysis.
Solutes that form solutions that can conduct electricity
are called electrolytes.
An electrolyte solution conducts electricity. Positive
ions gain electrons at the cathode and negative ions
lose electrons at the anode.
This transfer of electrons has the same effect as a
flow of electrons, and the solution conducts electricity.
The apparatus in which electrolysis occurs is called
an electrolytic cell.
An electrolytic cell has three essential features:
1.
an electrolyte solution that contains free-moving ions. These ions can donate or accept electrons, allowing
electrons to flow through the external circuit.
2.
two electrodes at which the electrolysis reactions occur
3.
an external source of electrons, such as a battery or power pack. This electron flow is in one direction only
and is termed DC (direct current).
The electrode to which electrons flow from the external power source is the negative electrode and is called the
cathode, since reduction will occur there.
The electrode from which electrons are withdrawn by the power source is the positive electrode and is called the
anode, since oxidation will occur there.
Cations are attracted to the cathode, while anions are attracted to the anode.
The cations gain electrons from the cathode and the anions lose electrons to the anode.
The redox reactions in an electrolytic cell can involve the electrode, or ions in the electrolyte or water molecules if
the electrolyte is an aqueous solution.
Page 3
ELECTROLYTIC CELLS – SOME EXAMPLES
1
Electrolysis of molten ionic compounds
Solid sodium chloride does not conduct electricity.
However, when an electric current is passed through
molten sodium chloride, a chemical reaction may be
clearly observed. A shiny bead of sodium is produced at
the cathode and chlorine gas is evolved at the anode.
In solid sodium chloride, the oppositely charged sodium ions, Na +, and chloride ions, Cl-, are held tightly together.
Heating the solid causes the ions in the crystal to separate so they are then free to move. In an electrolytic cell, the
sodium ions are attracted to the cathode where they are reduced:
Na+(l)
+
e-
Na(l)
The chloride ions are attracted to the anode where they undergo oxidation:
2Cl-(l)
Cl2(g)
+
2e-
Since, in a redox reaction, the same number of electrons are consumed as are produced, the overall equation is:
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QUESTION
Magnesium may be obtained commercially from seawater. During the last stage of the process, molten magnesium
chloride undergoes electrolysis in a cell that contains an iron cathode and a graphite anode.
1
Draw a fully labelled diagram of an electrolytic cell that could be used to produce magnesium.
Include equations at each electrode and the overall cell reaction.
2
Predict the product formed at the anode, if iron was used to form the anode instead of graphite in this cell.
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3
Why would a molten rather than aqueous solution of magnesium chloride be used?
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2
Electrolysis of water
When a current is applied via two electrodes in pure distilled water, no visible reaction occurs.
There is no current flow and no electrolysis. If, however, an electrolyte such as H 2SO4 or KNO3 is added in low
concentration, the resulting solution conducts electricity and electrolysis will occur. The products of the electrolysis
of water, in this case, are hydrogen and oxygen.
At the cathode, water is reduced to form hydrogen:
H2O(l)
+
2e-
H2(g)
2OH-(aq)
+
At the anode, water is oxidised to form oxygen:
H2O(l)
O2(g)
+
4e-
+
4H+(aq)
The region around the cathode becomes basic owing to an increase in OH -ions, whereas the region around the
anode becomes acidic, owing to an increase in H+ ions. The overall cell reaction may be obtained by adding the
half-equations:
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3
Electrolysis of aqueous solutions of ionic compounds
When a molten salt is electrolysed, the products are predictable.
However, when an aqueous solution of an ionic compound is electrolysed, water may react at one or both of the
electrodes in preference to the ions from the salt.
Hydrogen sometimes appears at the cathode, rather than a metal, and oxygen sometimes appears at the anode
rather than a metal cation or a halogen.
Hence, a range of redox reactions can occur with aqueous solutions, depending on the ions present and and the
nature of the electrodes.
NOTE: The following information is given to help you understand the electrolysis of aqueous solutions but is not
required learning.
Reduction reactions at the cathode
 If there are zinc ions or ions of metals less reactive than zinc in the aqueous solution, they will be reduced to
the metal.
For example, this will happen if an aqueous solution of copper (II) sulfate is electrolysed.
Copper ions will be reduced to copper metal and the half-equation is:
Cu2+(aq)
2e-
+
Cu(s)
If the solution is acidified, hydrogen ions, H+(aq), will be reduced to hydrogen gas and the half-equation is:
2H+(aq)
+
2e-
H2(g)
 If there are no ions of less active metals or there are no hydrogen ions, water molecules will be reduced to
hydrogen gas and the half-equation is:
H2O(l)
+
2e-
H2(g)
2OH-(aq)
+
Oxidation reactions at the anode
 If the electrode is not an inert metal (platinum, carbon, stainless steel), but a metal (eg copper), then the
electrode will be oxidised and go into solution as metal ions.
For example:
Cu(s)
Cu2+(aq)
+
2e-
 If the electrode is inert, and the solution is dilute, water molecules will be oxidised to oxygen gas.
H2O(l)
O2(g)
+
4e-
+
4H+(aq)
 When the electrolyte contains halide ions (Group 7 ions like Cl-), and is reasonably concentrated, the halide
ions get oxidised to the halogen.
For example:
2Cl-(aq)
Cl2(g)
+
2e-
Page 6
COMPARING GALVANIC AND ELECTROLYTIC CELLS
The following represents the main differences between an electrolytic cell and a galvanic cell:
 In an electrolytic cell, electrical energy is transformed into chemical energy. The electricity supplied causes
non-spontaneous chemical reactions to occur at the electrodes.
In a galvanic cell, chemical energy is transformed into electrical energy. The spontaneous chemical reactions
that occur generate a flow of electricity.
 In an electrolytic cell, the electricity supplied forces electrons around a circuit.
Reduction occurs at the electrode where electrons are gained and oxidation occurs at the electrode where
electrons are lost.
In a galvanic cell, there is a spontaneous flow of electrons around the circuit.
Oxidation occurs at the electrode where electrons are lost. The electrons move around the circuit to the other
electrode where they are gained and reduction occurs.
In both cells the flow of electrons is balanced by a movement of ions in either the electrolyte or the salt
bridge.
 In a galvanic cell, the anode is negative and the cathode positive.
In an electrolytic cell, the anode is positive and the cathode negative.
These differences can be summarized in the table below:
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SUPPORTING QUESTIONS – Galvanic/Electrolytic Cells
1
A galvanic cell was assembled where the overall cell reaction was:
Mg(s)
+
Mg2+(aq)
Br2(l)
+
2Br-(aq)
a State the energy transformation in this type of cell. _________________________________________
b Magnesium metal is used as the anode in this cell.
Write the half-equation for the reaction at the anode.
c Suggest a suitable material to use for the cathode electrode. _________________________________
d In the space below, sketch a diagram of this cell.
Identify on your diagram above:
i the charge on the electrodes,
ii the anode and cathode,
iii the direction of electron flow,
iv in which the direction the positive ions in the salt bridge would move.
2
An electrolytic cell is assembled where the overall cell reaction is:
MgBr2(l)
Mg(s)
+
Br2(l)
In the space opposite, sketch a diagram of the cell
where this reaction could occur.
On your diagram:
i
Name and give the formula for the electrolyte.
ii Name the material of which each electrode is made.
iii Show the direction of electron flow.
iv Show the movement of ions in the electrolyte.
v Label each electrode as cathode or anode.
vi Assign a charge to each electrode.
vii Write half-equations for oxidation and reduction
near each electrode.
Page 8
Practical:
ELECTROLYTIC CELLS
PURPOSE:
to assemble a variety of electrolytic cells,
to determine the products formed at each electrode,
to examine the changing products due to different,
to predict the likely products of an electrolysis.
METHOD
1
2
3
4
Construct each cell according to the instructions given in the table using a 6V power supply.
Draw a labelled diagram of each cell in the space provided.
State any observations made.
Identification tests:
phenolphthalein turns purple in the presence of OH- ions, a base,
Chlorine gas bleaches damp red litmus paper.
Write half-equations for reactions occurring at each electrode.
Cell 1: inert Norwood electrodes, tap water as the
Cell 2: copper electrodes, dilute sulfuric acid as the
electrolyte.
In the cathode, add 2-3 drops of phenolphthalein.
electrolyte.
Cell 3: inert Norwood electrodes, dilute sodium
Cell 4: inert Norwood electrodes, concentrated
chloride (0.2 g in 50 mL) as the electrolyte.
In the cathode, add 2-3 drops of phenolphthalein.
sodium chloride (15g in 50 mL) as the electrolyte.
In the cathode, add 2-3 drops of phenolphthalein.
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EXTRA NOTES:
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