Week 8, Lesson 1
Chapter 16 – Redox Chemistry &
Corrosion
Oxidation & Reduction
• Redox reactions are reactions that are called
oxidation-reduction reactions.
• Many of the chemical reactions that play a
significant role in maintaining our
environment are redox reactions.
• For example, the reactions of water and the
gases of the atmosphere that cause corrosion
and so the deterioration of metals are redox
reactions.
Redox Reactions
• 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 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 electrolytes to produce many chemicals
– Production and us of fertilisers.
Redox Reactions cont…
• Once the reactive component of air was first
produced by Joseph Priestley identified by
Antoine Lavoisier as the element 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.
• This type of reaction was classified as oxidation,
which here means the addition of oxygen.
Redox Reactions cont…
• When oxygen reacts with an element, the
element is said to be oxidised.
• The iron that society uses has been extracted
from iron ores.
• This process involves reduction of the iron
oxide to iron.
• When oxygen has been removed from a
substance the material has been reduced.
Redox Reactions Example…
• The iron (III) oxide has lost oxygen – it has
been reduced.
• Reduction can not occur without oxidation
occurring at the same time.
• In this reaction it is the carbon monoxide that
has gained oxygen – and it is said to have been
oxidised.
• Oxidation and reduction always occur
simultaneously.
An Improved Definition – Electron
Transfer
• While the definition of oxidation as the addition of oxygen
was adequate for combustion reactions occuring in air, many
similar reactions not involving oxygen as a reactant have since
been investigated.
• Magnesium ribbon is often dull and covered with a coating
that needs to be scraped of. This is magnesium oxide which
results from the corrosion of Mg in air.
• The magnesium reacted with atmospheric oxygen to form
magnesium oxide – the magnesium has been oxidised.
• This reaction is represented by the equation:
2Mg(s) + O2  2MgO(s)
Electron Transfer cont…
• Magnesium oxide is an ionic compound and consists
of Mg2+ and O2- ions.
• Each magnesium atom has lost 2 electrons to form
this ion.
• Each oxygen molecule has gained 2 electrons to form
this ion.
• The reaction can be represented by two half
equations.
• The first shows the loss of electrons:
Mg(s)  Mg2+(s) + 2e• The second shows the gain of electrons:
O2(g) + 4e-  2O2-(s)
Electron Transfer cont…
• Therefore oxidation is defined as the loss of
electrons.
• Reduction is defined as the gain of electrons.
• Reactions that involve oxidation and reduction
are called redox reactions.
Writing Half Equations example…
When a strip of copper wire is suspended in a solution of silver
nitrate, long crystals of silver metal can be observed. The
solution changes to a pale blue colour, indicating the presence
of Cu2+ ions. Write half equations for this reaction and identify
the substance oxidised and reduced.
STEP 1: The first equation is the copper metal forming copper
ions.
Cu(s) Cu2+(aq) + 2eSTEP 2: The second equation involves the silver ions form silver
metal.
Ag+(aq) + e-  Ag(s)
Since the copper has lost electrons is has undergone oxidation
and the silver has undergone reduction.
Writing an Overall Redox Reaction
• When writing equations for redox reactions,
normally the two half equation are written
first and then they are added together to
produce an overall equation.
• In an overall equation, no electron transfer is
shown.
• That is, the electrons that are lost in oxidation,
are gained in the reduction reaction.
Writing an Overall Redox Reaction
Example…
In the example of copper and the solution of silver ions:
- Each copper atom that is oxidised loses two electrons.
- Each silver ion that is reduced gains on electron.
Two silver ions must be reduced to take up the electrons lost by each
copper atom oxidised.
To write an overall equation, the half equation involving the reduction
of silver is multiplied by 2, before it is combined with the half
equation for the oxidation of copper.
Cu(s)Cu2+(aq) + 2e(Ag+(aq) + e-  Ag(s)) x 2
The overall equation is
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
Overall Redox Reactions cont…
• Remember in both half and full redox
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 reactions.
– The total charge on the product side of the
equation is equal to the total charge on the
reactants side of the equation. Charge is
conserved in all reactions.
Oxidants and Reductants
• An oxidant (or oxidising agent) is a species
that causes another to be oxidised.
• A reductant is a species that causes another to
be reduced.
• An oxidant causes oxidation and is reduced.
• A reductant causes reduction and is oxidised.
Oxidants and Reductants cont…
Week 8, Lesson 2
Oxidation Numbers
• To determine whether or not an equation is a
redox reaction and to help us identify oxidants
and reactants, oxidation numbers are assigned
to elements involved in a reaction.
• These numbers are determined by applying a
set of rules that consider unequal sharing of
electrons between two atoms.
Oxidation Number Rules
• Oxidation numbers are determine according to the following rules in
ascending order:
1. The oxidation number of a free element is 0.
2. The oxidation number of a simple ions is equal to the charge on that ion.
3. In compounds, some elements have oxidation numbers that are regarded
as fixed, except in a few exceptional circumstances.
a) Main group elements have an oxidation number equal to their valency.
b) Fluorine has an oxidation number of -1.
c) Hydrogen has an oxidation number of +1 when is forms compounds with
non metals and -1 in metal hydrides.
d) Oxygen almost always has an oxidation number of -2. In F2O it has an
oxidation number of +2 and in peroxides such as H2O2 and BaO2 its
oxidation number is -1.
4. The sum of oxidation numbers in a neutral compound is zero.
5. The sum of oxidation number in a polyatomic ion is equal to the charge on
the ion.
Using Oxidation Numbers
• A change in oxidation numbers is an indication
that a redox reaction have taken place.
• Oxidation involves an increase in oxidation
number.
• Reduction involves a decrease in oxidation
number.
Week 8, Lesson 3
Half Equations for Complex Redox
Reactions
• Look at the half equation for the oxidation of iron (II) ions to iron (III) ions:
Fe2+(aq)  Fe3+(aq) + e• The conversion of MnO4- to Mn2+ involves reduction.
• The following steps are used to balance this half equation:
1. Balance all atoms in the half equation except hydrogen and oxygen.
2. Balance the oxygen by adding water.
3. Balance the hydrogen atoms by adding H+
4. Balance the charges on both sides by adding electrons to the more
positive side.
Then add the oxidation half equation to the reduction half equation to get the
overall equation.
Make sure the number of electrons used in reduction equals the number of
electrons released during oxidation.
Week 9, Lesson 1
Evidence for Electron Transfer
• When zinc metal is placed in a solution of
copper (II) sulfate, the zinc is oxidised and a
brown deposit of copper metal forms.
• Electrons flow from the zinc atoms to the
copper(II) ions as these collide.
• Evidence for this flow can be obtained by
setting up an electrochemical or galvanic cell.
• A galvanic cell is an energy converter that
separates the oxidant from the reductant.
Galvanic Cells
• The zinc metal and solution containing copper(II) ions are in
separate beakers.
• A copper strip, dipping in the solution of copper (II) ions is
connected by a piece of wire to a galvanometer.
• A galvanometer is a meter for detecting a flow of electrons.
• A zinc strip, dipping into a solution of zinc chloride is also
connected to this by a piece of wire.
• The solutions in the two beakers are connected by a salt
bridge.
• The salt bridge contains an ionic compound.
• The ions flow through the salt bridge between the solutions
to complete the circuit.
• The salt bridge can be a piece of filter paper soaked in
potassium nitrate.
Galvanic Cells cont…
• When this particular galvanic cell operates,
the galvanometer indicates that electrons are
flowing from the zinc strip in beaker B to the
copper strip in beaker A.
• This observation provides evidence that the
oxidation and reduction reactions between
zinc atoms and copper ions also occur what
reactants are separated in a galvanic cell.
• It also supports the idea that there is a
transfer of electrons.
Zinc and Copper Galvanic Cell
• At the copper strip, in beaker A, Cu2+ ions accept electrons to
form copper atoms:
Cu2+(aq) + 2e-  Cu(s)
• At the zinc strip in beaker B, zinc atoms lose electrons to form
zinc ions:
Zn(s)  Zn2+(aq) + 2e• Because of these reactions, the electrons flow through the
zinc strip and connecting wire towards the copper strip.
• The zinc is the source of negatively charged electrons so it is
assigned a negative charge.
• The electrons are attracted to the copper strip so this is
assigned a positive charge.
Salt Bridge
• To maintain the balance of positive and negative
charge in each beaker, ions flow from the salt bridge
into the two solutions.
• In the beaker containing the copper sulfate solution,
Cu2+ ions are reduced to Cu atoms, leaving SO42- ions
in the solution.
• The K+ ions flow from the salt bridge into this beaker
to replace the Cu2+ ions.
• At the same time, NO3- ions flow into the beaker
containing ZnCl2 solution to balance the positive
charge of the Zn2+ ions being produced at the zinc
electrode.
Galvanic Cells
•
•
•
•
All galvanic cells are composed of two half cells.
Oxidation occurs in one and reduction in the other.
A half cell must contain al electrode and electrolyte.
An electrode is an electronic conductor – a material
that can move through the circuit.
• The electrode at which oxidation takes place is called
the anode.
• The electrode at which reduction takes place is call
the cathode.
Galvanic Cells cont…
• In galvanic cells the anode is the negatively charged electrode.
• The cathode is the positively charged electrode.
• In the previous example, the copper ions are reduced to copper
atoms at the cathode.
• Cations will migrate from the salt bridge in the beaker containing
the cathode to compensate for the loss of copper ions.
• At the anode, zinc metal is oxidised and so more zinc ions are
added to the solution in the beaker.
• To avoid a build up of positive charge, anions will migrate from
the salt bridge into the beaker and so maintain electrical
neutrality.
• An electrolyte contains ions that are free to move through the
solution.
Galvanic Cell
Electrochemical Series
• Sodium, magnesium and iron are metals that
easily corrode because they are easily
oxidised.
• Chemists have combined such information
with experimental data from constructing
galvanic cells from various combinations of
half cells to produce the electrochemical
series.
• The electrochemical series lists half equations
in order of their tendency to occur as
reduction reactions.
Electrochemical Series cont…
Electrochemical Series cont…
• The half cell represented by the equation at the top
of the list attracts electrons most strongly.
• It will form the positive half cell when it is connected
to any other half cell represented in the series.
• An electrochemical series shows the relative
reactivities of a number of metals and their
corresponding.
• The most reactive metals are those that are oxidised
most easily; the most reactive metal cations are
those that are reduced most readily.
Electrochemical Series cont…
As you go down the series, the:
- Metals, which are on the right hand side, become
increasingly easy to oxidise and therefore are more
reactive. Metals lower in the series lose electrons
more readily and so are stronger reductants.
- Metal cations which are on the left hand side,
become increasingly hard to reduce and therefore
are less reactive. Cations higher in the series have a
greater attraction for electrons and so are more
readily reduced, and therefore are strong oxidants.
Predicting Redox Reactions
• The electrochemical series is useful for predicting redox
reactions.
• More reactive metals tend to be found on the lower right of
the 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.
• A spontaneous 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 reaction occurs in the
forwards direction.
• The reductant is oxidised and its half equation is the reverse
of that on the electrochemical series.
Predicting Redox Reactions cont…
Is reduced
Oxidant
Reacts with
Reductant
Is oxidised
Predicting Reactions in Galvanic
Cells
• An electrochemical series can be used to
predict the processes that occur when two
half cells are joined in an electrochemical cell.
• The galvanic cell on the next slide is made by
joining Ag+/Ag half cell with the Cu2+/Cu half
cell and linking them with a salt bridge.
• In this galvanic cell the two possible oxidants
are Ag+ and Cu2+ and the two possible
reductants are Cu and Ag.
Predicting Reactions in Galvanic
Cells cont…
• Ag+(aq) is higher on the left-hand
side of the electrochemical series
than Cu2+ and so is the stronger of
the two oxidants.
• This means that Ag+ attracts
electrons more strongly than Cu2+,
so electrons in the wire will flow
away from the copper electrode
towards the silver electrode.
• Since oxidation occurs at the anode
and the reduction at the cathode,
the half cell reactions are:
At the cathode: Ag+(aq) + e-  Ag(s)
At the anode: Cu(s)  Cu2+(aq) + 2e-
AN OIL RIG CAT
• Use the following phrase to help you
remember the process:
ANode: Oxidation Is Loss of electrons Reduction
Is Gain of electrons CAThode.
Week 9, Lesson 2
The Corrosion Process
• Direct reaction with oxygen in the air to form a
metal ocise is known as dry corrosion.
• In cases such as aluminium, dry corrosion forms a
tough, impervious oxide coating that protects the
metal underneath from further contact from
oxygen.
• When moisture is present, the corrosion of iron
quickens.
• This process is known as wet corrosion.
• The most well known form of wet corrosion is the
formation of rust.
Wet Corrosion
• Wet corrosion can occur in moist air or by direct
immersion in water.
• Several factors affect the rate at which wet corrosion
occurs.
• In general, corrosion is accelerated in the presence
of:
– Water – corrosion in dry air is very slow
– Impurities such as salt and acidic pollutants that dissolve in
the water
– Impurities such as carbon in the iron
• Corrosion is reduced however when the iron:
– Is alloyed with certain other materials
– Has a protective coating
The Wet Corrosion Process
• Analysis of rust shows that it is a hydrated
oxide of iron with the formula Fe2O3.xH2O.
• x represents a number between 1 and 3 and
xH2O indicates that this is the number of
water molecules associated with the iron
oxide.
The Wet Corrosion Process cont…
• Experimental results indicate that the first step in the
formation of rust is the oxidation of iron to form Fe2+
ions.
• These ions can be detected in solution on the surface
of the iron.
• An Fe2+ ion has two less electrons than and iron
atom.
• This step therefore involves the loss of electrons and
can be represented by the half equation:
Fe(s)  Fe2+(aq) + 2e-
The Wet Corrosion Process cont…
• This only represents have the equation. The
electrons lost by the iron must be accepted by
another reactant.
• Oxygen is also involved in corrosion – it gains
the electrons lost by the iron.
• The oxygen is reduced in the presence of
water to hydroxide ions, represented by:
O2(aq) + 2H2O(l) + 4e-  4OH-(aq)
The Wet Corrosion Process cont…
• The overall equation can be written as:
2Fe(s) + O2(aq) + 2H2O(l)  2Fe2+(aq) + 4OH-(aq)
The Wet Corrosion Process cont…
• The iron (II) ions then react with the hydroxide
ions to form a green precipitate of iron(II)
hydroxide.
• On exposure to more moisture and oxygen,
the iron(II) hydroxide is further oxidised to
iron(III) hydroxide, a red brown precipitate.
The Wet Corrosion Process cont…
• The wet corrosion of iron is a galvanic process.
• Galvanic cells are set up on the metal surface
where different regions act as anodes and
cathodes:
– The cathode region, where oxygen gains electrons
and is reduced, occurs at sites of high oxygen
concentration.
– The anode region, where iron loses electrons and
is oxidised, oftern occurs in an area of stress in the
metal where the metal lattice has been deformed,
for example at the tip of a nail.
The Wet Corrosion Process cont…
• The anode region is one where there is a lower
concentration of oxygen.
• Corrosion is therefore greatest at the centre of a
water drop or under a layer of chipped paint.
• The cathode region is at the edge of the water
drop or where paint is missing as the oxygen
concentration is the greatest.
• Wet corrosion is accelerated by the presence of
dissolved salts, if the iron contains impurities or if
it is in less contact with a less reactive metal.
Protecting Against Corrosion
• There are several ways to protect iron and steel
structures against corrosion.
• These depend on the environment in which the
structure is located, the design and intended use of
the structure, as well as the relative costs of various
protection measures.
• The simplest of these methods involves creating a
barrier between the iron and oxygen and water to
provide surface protection.
• Other methods are based on an understanding of the
electrochemical nature of corrosion.
Surface Protection
• Involves covering the surface of the iron to prevent contact
with oxygen and moisture.
• Materials such as paint and plastic can be used for this
purpose.
• Iron can also be protected by alloying with small quantities of
metals such as chromium, nickel, manganese and
molybdenum to produce stainless steel.
• The atoms of these metals used to make the allot are all
bonded into the metallic lattice.
• These metals oxidise slightly in air, but the oxide coating
produced is continuous and unreactive.
• This oxide layer forms a dense, impervious barrier between
the metal and oxygen and water, protecting the iron from
further oxidation.
Surface Protection cont…
• Metals such as iron and steel can also be
coated with thin layers of other metals in a
process known as electroplating.
• These other metals, which include chromium
and nickel, again form protective oxide
coatings.
Electrochemical Protection – Impressed
Current Cathodic Protection
• Impressed current protection involved reversing
the direction of electron flow by using a low
voltage direct current power supply to give the
iron being protected a negative charge.
• Because the iron is receiving electrons, oxidation
is inhibited (oxidation would produce more
electrons).
• The iron becomes the site of the reduction
reaction, that is the cathode.
• This method is used to protect large scale steel
structures such as wharves and pipelines.
Electrochemical Protection –
Sacrificial Protection
• This also involves the iron acting as the cathode.
• A more easily oxidised metal, such as zinc, forms the
anode and is sacrificed.
• The more reactive metal loses electrons and forms metal
cations in preference to the iron.
• The iron is not oxidised in the presence of zinc because
the zinc has protected the iron from corrosion.
• This type of protection is known as electrochemical or
sacrificial protection.
• This form of protection is used in roofing, galvanised iron
is iron with a zinc coating.