Redox Chemistry and
Corrosion
Chapter 16
Oxidation and Reduction
► So
far we have looked at precipitation
reactions and acid-base reactions.
► Now we shall look at a third group of
chemical reactions.
► They are called oxidation-reduction
reactions.
► These reactions are commonly referred to
as redox reactions.
Redox Reactions
► Many
of the chemical reactions that play a
significant role in maintaining our environment are
redox reactions.
► Corrosion and the deterioration of metals are
redox reactions.
► Iron which is used as a structural base for
buildings and bridges is particularly prone to
corrosion.
► Australia spends about 3 billion dollars a year in an
effort to prevent corrosion and replacing
structures that have corroded.
Redox Reactions
► These
reactions are also
used in the processing of
mineral ores to extract from
then the metals our society
requires.
► One of Australia’s biggest
exports is the mining of
these mineral ores.
Redox Reactions
► Other
redox reactions include:
 The respiration reaction that is the source of energy in
almost all living things.
 Photosynthesis in green plants
 Burning of fuels to propel our cars.
 Combustion of coal in electricity power stations.
 Use of chemicals such as chlorine to disinfect swimming
pools.
 Manufacture and use of explosives.
 Use of electrolysis to produce many chemicals.
 Production and use of fertilisers.
Redox Reactions
► Many
chemicals react with oxygen.
► Reactions such as these were described as
oxidation reactions.
► In air, the combustion of carbon, sulfur, iron
or even octane always produced at least
one oxide:




C(s) + O2(g) ―› CO2(g)
S(s) + O2(g) ―› SO2(g)
4Fe(s) + 3O2(g) ―› 2Fe2O3(s)
2C8H18(l) + 25O2(g) ―›16CO2(g) + 18H2O(l)
Oxidation
► Oxidation
means the addition of oxygen.
► When oxygen reacts with an element, the
element is said to be oxidised.
► Because elemental iron reacts with oxygen,
there are no large deposits of elemental iron
found on earth.
► Iron is generally found as a compound of
mineral oxide ores (haematite (Fe2O3) and
magnetite (Fe3O4))
Reduction
► Iron
used in modern society has been
extracted from iron ores.
► This extraction process involves reduction of
the iron oxide to iron.
► It involves the removal or oxygen.
► When oxygen is removed from a substance,
that material has been reduced.
Reduction
►
The production of iron from haematite can be represented
by the reduction equation:
Reduction – loss of oxygen
Fe2O3(s) + 3CO(g) ―› 2Fe(s) + 3CO2(g)
Oxidation – gain of oxygen
The iron(III) oxide has lost an oxygen – it has been reduced.
Reduction cannot occur without oxidation occurring at the
same time.
In this reaction the carbon monoxide has gained an oxygen –
it has been oxidised.
A Better Definition
► There
are many oxidation and reduction
reactions that don’t involve oxygen.
► Instead we define oxidation as the loss of
electrons.
► Similarly, reduction is the gain of electrons
rather than the loss of oxygen.
OIL RIG
►Oxidation
is the loss of electrons
►Reduction is the gain of electrons.
Magnesium Oxide
► You
have used magnesium in class before,
remember how it has a coating on it that
sometimes we have had to scrape off.
► That is magnesium oxide which results in
corrosion of magnesium in air.
Magnesium Oxide
► The
magnesium has reacted with
atmospheric oxygen to form magnesium
oxide.
► The magnesium has been oxidised.
2Mg(s) + O2(g) ―> 2MgO(s)
2Mg(s) + O2(g) ―> 2MgO(s)
► Magnesium
oxide is an ionic compound and
consists of Mg2+ ions and O2- ions.
► Each magnesium ion, therefore must have
lost two electrons to form an Mg2+ ion. Each
oxygen atom in the oxygen molecule must
have gained two electrons to form an oxide
ion O2-.
► The reaction can now be represented by
two half equations.
2Mg(s) + O2(g) ―> 2MgO(s)
► The
first half equation show the gain of two
electrons by each oxygen atom in the
oxygen molecule:
Mg(s) ―> Mg2+(s) + 2e► The second show the gain of two electrons
by each oxygen atom in the oxygen
molecule.
O2(g) + 4e- ―> 2O2-(s)
Mg(s) ―> Mg2+(s) + 2eO2(g) + 4e- ―> 2O2-(s)
► So
the oxidation of magnesium involves the
transfer of electrons from magnesium atoms to
oxygen atoms.
► Note that there is no real ‘loss of electrons’ but
rather a transfer of electrons from the magnesium
to the oxygen.
► If an atom loses electrons, there must be another
atom that can gain electrons.
► Therefore oxidation and reduction occur
simultaneously.
Writing Redox Half Equations
► Worked
► 16.2b
Example 16.2a page 275
Your Turn
► Page
278
► Question 1
► Question 2
Writing an Overall Redox Equation
► When
we write equation for redox reactions,
we normally write the two half equations
first.
► We then follow this with the overall
equation.
► In the overall equation we do not show any
electrons transferred as:
 The electrons lost in the oxidation reaction are
gained in the reduction reaction.
Copper and the solution of silver ions
► In
the previous example:
 Each copper atom that is oxidised loses two
electrons
 Each Ag+ ion that is reduced gains one electron.
► When
writing full equations we must
balance the electrons first.
► Therefore two Ag+ ion must be reduced to
take up the electrons lost by each copper
atom that is oxidised.
Copper and the solution of silver ions
Cu(s) ―> Cu2+(aq) + 2e( Ag+(aq) + e- ―> Ag(s) ) x 2
So we need to times the silver ions by 2
The overall equation is:
Cu(s) + 2Ag+(aq) ―> Cu2+(aq) + 2Ag(s)
Remember
► In
both half and overall equations.
 The number of atoms of each element present
in the products is equal to the number present
in the reactants.
 Atoms are conserved in all chemical equations.
 The total charge on the product side of the
equation is equal to the total charge on the
reactant side of the equation.
 Charge is conserved in chemical reactions.
Worked Example 16.2c
When sodium is oxidised by atmospheric
oxygen, the reaction can be represented by
the following half equations:
Na(s) ―> Na+(s) + eO2(g) + 4e- ―> 2O2-(s)
Identify the half equation representing the
oxidation reaction and write the balanced
overall equation.
Oxidants and Reductants
► An
oxidant (or oxidising agent) is a species
that causes another to be oxidised.
► A reductant (or reducing agent) is a species
that causes another to be reduced.
► The oxidant itself is reduced.
► The reductant is oxidised.
Your Turn
► Page
278
► Question 3 and 4
Predicting electron transfer
► Read
pages 283 – 285
► What is a galvanic cell?
Galvanic Cell
► All
galvanic cells are composed of two half
cells.
► Oxidation occurs in one half cell.
► Reduction occurs in the other.
► A half cell must contain an electrode and an
electrolyte.
► An electrode is an electronic conductor – a
material that has delocalised electrons that
can move through the circuit.
Galvanic Cells
► The
electrode at which oxidation takes place
is called the anode.
► The
electrode at which reduction takes
place is called the cathode.
Galvanic Cells
► Zinc
is the anode.
► Copper is the cathode.
► In galvanic cells the anode is negatively charged
and the cathode is positively charged.
Galvanic Cells
► Cu2+
ions are reduced to Cu atoms at the
cathode.
► Cations will migrate from the salt bridge into
the beaker containing that cathode to
compensate for the loss of the Cu2+ ions.
► At the anode, zinc metal is oxidised and so
more Zn2+ ions are added to the solution in
that beaker.
The salt bridge
► To
avoid the build up of a positive charge,
anions (negatively charged ions) will
migrate from the salt bridge into the beaker
and so maintain electrical neutrality.
Electrolyte
► An
electrolyte contains ions that are free to
move through the solution.
► In the example the electrolyte in beaker A
was the zinc chloride.
► The electrolyte in beaker B was the copper
sulfate solution.
Galvanic Cells Comprise Of:
► Two
half cells, which are separate and do
not mix.
► A length of wire connecting the electrodes
of the half cells. This is the external current.
► A salt bridge to connect the solutions in the
half cells. This is the electrical conductor.
► The salt bridge balances the overall charge
during the circuit.
The electrochemical series
► Sodium,
magnesium and iron are all metals
that corrode easily because they are easily
oxidised.
► Sodium is oxidised so easily that it is stored
under paraffin oil.
► Other metals, however, do not corrode
readily. Platinum and gold are sufficiently
inert to be found free in nature.
The Electrochemical Series
► Table
16.2 on page 287 represents the
electrochemical series.
► What
can you tell me about the
electrochemical series?
The electrochemical series
► Each
half equation represents the reduction
reactions.
► The top equation is the strongest oxidant so it is
most easily reduced.
► The strongest reductants are at the bottom and
are oxidised quite easily. What kind of metals do
these mainly consist of?
► In general the smaller amount of energy required
to remove a valance electron the more readily the
metal will act as a reductant and itself be oxidised.
Electrochemical Series
► Non-metals
tend to gain electrons and
therefore act as oxidants.
► Reactive metals tend to be stronger
reductants.
► Transition metals are less readily oxidised.
Predicting Redox Reactions
► We
use the electrochemical series to predict
redox reactions.
► More reactive metals tend to be found on
the lower right of the electrochemical series.
► A more reactive metal will be oxidised by,
and donate its electrons to the cation of a
less reactive metal.
► The cation receives the electrons and is
reduced.
Predicting Redox Reactions
►A
spontaneous redox reaction can be
expected to occur when a relatively strong
oxidant is mixed with a relatively strong
reductant.
► The oxidant is reduced and the half
equation occurs in the forward direction.
► The reductant is reduced and the half
equation occurs in the reverse direction of
the that on the electrochemical series.
Predicting Redox Reactions
► We
can predict that zinc metal with react
with Cu2+ ions because zinc is more reactive
than copper.
Is reduced
Cu2+(aq)
What is the overall Equation????
+ 2e- ―> Cu(s)
Reacts with
Is oxidised
Zn2+ + 2e- <―
Zn(s)
AN OIL RIG CAT
► Anode
+ Oxidation is loss of electrons:
► Reduction is gain of electrons + Cathode
►A
way to remember oxidation occurs at the
anode. Reduction occurs at the cathode.
Predicting Reactions
► For
reactions to occur spontaneously, the
aqueous cation in the solution must be a
stronger oxidant than the cation of the
metal added.
► Your
Turn
► Try Question 13 on page 291
► Try Question 15 as well
Your Turn
► Finish
reading this chapter yourself about
corrosion.