9.2
ELECTROCHEMICAL CELLS
ESSENTIAL IDEA
Voltaic cells convert chemical energy to
electrical energy and electrolytic cells
convert electrical energy to chemical
energy.
NATURE OF SCIENCE (4.5)
Ethical implications of research – the desire to
produce energy can be driven by social needs or
profit.
INTERNATIONAL-MINDEDNESS
Research in space exploration often
centers on energy factors. The basic
hydrogen-oxygen fuel cell can be used
as an energy source in spacecraft, such
as those first engineered by NASA in the
USA. The International Space Station is a
good example of a multinational project
involving the international scientific
community.
THEORY OF KNOWLEDGE
Is energy just an abstract concept used
to justify why certain types of changes
are always associated with each other?
Are concepts such as energy real?
UNDERSTANDING/KEY IDEA
9.2.A
Voltaic cells convert energy from
spontaneous, exothermic chemical
processes to electrical energy.
APPLICATION/SKILLS
Be able to construct and annotate
a voltaic cell.
APPLICATION/SKILLS
Be able to distinguish between
electron and ion flow in a voltaic
cell.
The redox reaction involves a transfer of
electrons from the reducing agent
(species oxidized) to the oxidizing agent
(species reduced).
 When redox reactions occur in the same
container, the chemical energy is
released as heat so no useful work is
obtained.
 Chemists wish to harness this chemical
energy.

Voltaic Cells
The key is to separate the oxidizing
agent from the reducing agent, thus
requiring the electron transfer to occur
through a wire.
 The current from the electrons in the
wire can then be harnessed to provide
useful work.
 A voltaic (galvanic) cell uses a
spontaneous redox reaction to produce
a current that can be used to do work.

APPLICATION/SKILLS
Be able to explain how a redox
reaction is used to produce
electricity in a voltaic cell.
 The
reaction in a voltaic cell is always
a redox reaction which can be broken
into two half-reactions.
 Oxidation occurs at the anode.
 Reduction occurs at the cathode.
 (My way to remember this is that
reduction involves the gain of electrons
and electrons were discovered by the
cathode-ray experiments.)
If we want to harness the energy, we
need to separate the ½ reactions into
half-cells.
 Then the electrons are only allowed to
flow through an external circuit.
 This creates the voltaic cell.

The easiest way to make a half cell is to put
a strip of metal into a solution of its own ions.
 If you put a strip of zinc into a solution of zinc
sulfate, the zinc atoms form ions by releasing
electrons which make the surface of the
metal negatively charged with respect to the
solution. An equilibrium is set up between
the metal and its solution of ions.
 This charge separation is the electrode
potential.
 The more reactive a metal, the more
negative its electrode potential in its half-cell.

If you connect the two half cells by an
external wire, electrons will flow
spontaneously from the electrode of higher
potential to the one with lower potential.
 The half cells connected in this way are
called electrodes.
 The electrode where reduction occurs is
called the cathode.
 The electrode where oxidation occurs is
called the anode.

To make a voltaic cell

An electrical wire must connect to the metal
electrodes in each half-cell.
 A voltmeter can be attached to the external circuit to
record the voltage generated.
 Electrons always flow from anode to cathode through
the wire.

A salt bridge completes the circuit.
 Without a bridge, no voltage will be generated.
 The bridge is a glass tube or such that contains an
aqueous solution of ions such as KNO3.
 This allows ion flow in the opposite direction so that
potential difference can be maintained.
○ Cations flow towards the cathode.
○ Anions flow towards the anode.





Any two metal half-cells can be connected to
make a voltaic cell.
The direction of electron flow and voltage
generated will be determined by the difference
in reducing strength of the 2 metals.
In most cases this can be judged by the relative
position of the metals in the reactivity series.
Electrons will flow from the higher element to
the lower element in the series.
The greater reducing agent is always the
anode.
UNDERSTANDING/KEY IDEA
9.2.B
Oxidation occurs at the anode
(negative electrode) and reduction
occurs at the cathode (positive
electrode) in a voltaic cell.
In a voltaic cell, oxidation always occurs
at the anode.
 Reduction always occurs at the cathode.
 In the voltaic cell, the anode has a
negative charge.
 The cathode has a positive charge.

GUIDANCE
For voltaic cells, a cell diagram
convention should be learned.
Sometimes it is convenient to use a
shorthand notation rather than
drawing out the whole cell.
1. A single vertical line represents a phase boundary.
2. A double vertical line represents the salt bridge.
3. Anode is on the left, cathode right.
4. Spectator ions are omitted.
5. If a half cell contains 2 ions, they are separated
by a comma because they are in the same phase.
salt bridge
H2(g) H+(aq)
Cu2+(aq) Cu(s)
anode
cathode
6. Always include phases in parentheses.
UNDERSTANDING/KEY IDEA
9.2.C
Electrolytic cells convert electrical
energy to chemical energy, by
bringing about non-spontaneous
processes.
APPLICATION/SKILLS
Be able to construct and annotate
an electrolytic cell.
APPLICATION/SKILLS
Be able to distinguish between
electron and ion flow in an
electrolytic cell.
To review, a voltaic cell takes the energy
of a spontaneous redox reaction and
harnesses it to produce an electric
voltage.
 An electrolytic cell uses an external
source of voltage to bring about a redox
reaction that would otherwise be nonspontaneous.


The reactant in an electrolytic cell is the
electrolyte.
 This is usually a liquid either a molten ionic
compound or a solution of an ionic compound.
As the electric current passes through the
electrolyte, redox reactions occur at the
electrodes removing the charges on the
ions and forming products that are neutral
elements.
 The ions are said to be discharged during
this process.

Reactive metals such as aluminum, lithium,
magnesium, sodium and potassium are
found naturally in compounds such as Al2O3
and NaCl where they exist as positive ions.
 Extraction of the metal element alone
therefore involves reduction of the ions.
 But remember the E˚ of these metals are so
low or negative that there are no good
reducing agents available to do this.
 This means that electrolysis is usually the
only means by which the metals can be
extracted from their ores.

The longer lines in the power source represent the
positive terminal and the shorter lines represent
the negative terminal.
Electrolyte
Anions are
oxidized
here.
Cations are
reduced
here.
UNDERSTANDING/KEY IDEA
9.2.D
Oxidation occurs at the anode
(positive electrode) and reduction
occurs at the cathode (negative
electrode) in an electrolytic cell.
The positive ions are attracted to the
negative electrode (cathode) while the
negative ions are attracted to the
positive electrode (anode).
 At the anode (positive electrode),
negative ions lose electrons. (oxidation)
 At the cathode (negative electrode),
positive ions gain electrons. (reduction)


In both types of cells, electrons always flow
from anode to cathode.
 Oxidation always occurs at the anode.
 Reduction always occurs at the cathode.

The difference is in the charge of the
electrodes.
 Voltaic cells – anode is neg, cathode is pos
 Electrolytic cells – anode is pos, cathode is neg
APPLICATION/SKILLS
Be able to explain how current is
conducted in an electrolytic cell.
The source of electric power is a battery or DC
power source.
 This power source is indicated in the diagram with
two vertical lines, one being longer than the other.

 The longer line represents the positive terminal.
 The shorter line represents the negative terminal.
The electrodes are connected to the power supply
but are never touching.
 Electrodes are made of a conducting substance
such as metal or graphite.
 Electric wires connect the power supply to the
electrodes.






The power source pushes the electrons
towards the negative electrode where they
enter the electrolyte.
This is the cathode.
Electrons are released at the positive
terminal, the anode, and returned to the
source.
The current is not passed through the
electrolyte by electrons but by the ions that
are mobile and migrate to the electrodes.
The chemical reactions occurring at each
electrode remove the ions from the solution
and enable the process to continue.
APPLICATION/SKILLS
Be able to deduce the products of
the electrolysis of a molten salt.
When the electrolyte is a molten salt, the
only ions present are from the salt.
 There is no solvent.
 It is easy to predict the products
because the anion migrates to the
anode and the cation migrates to the
cathode.

EXAMPLE


Describe the reactions that occur during the electrolysis
of molten lead II bromide. Write an equation for the
overall reaction and comment on any observable
changes.
First: Deduce the ions and the appropriate electrodes.
PbBr2(l) → Pb2+(l) + 2Br-(l)
cathode
anode
Deduce the half equations.
Anode: Br- is oxidized to Br2
2Br- → Br2 + 2eCathode: Pb2+ is reduced to Pb
Pb2+ + 2e- → Pb
 Overall reaction: Pb2+ + 2Br- → Br2 + Pb
 Observations: A brown liquid Br2 with a strong smell at
the anode and the appearance of a grey metal Pb at
the cathode.

Most ionic compounds have very high
melting points so it is costly to supply
and maintain the energy for electrolysis.
 Sometimes another compound is added
to lower the melting point to make it
more economical.
 You may have more than one ion
attracted to the same electrode, thus E˚
data is important.

Citations
International Baccalaureate Organization. Chemistry Guide,
First assessment 2016. Updated 2015.
Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd
ed. N.p.: Pearson Baccalaureate, 2014. Print.
Most of the information found in this power point comes
directly from this textbook.
The power point has been made to directly complement the
Higher Level Chemistry textbook by Catrin and Brown and is
used for direct instructional purposes only.