Standard Grade Chemistry
Summary Notes
Topic 12. Corrosion
General
Learning Outcomes
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Corrosion is a chemical reaction which involves the surface of a metal changing from an
element to a compound.
Give examples to show that different metals corrode at different rates.
Rusting is the corrosion of iron.
Water and oxygen (from the air) are required for rusting.
When iron rusts, initially the iron atoms lose two electrons to form Fe2+ ions.
Ferroxyl indicator for Fe2+(aq) can be used to show the extent of the rusting process.
Salt spread on roads increases the rate of corrosion on car body work.
Salt acts as an electrolyte.
Iron does not rust when attached to the negative terminal of a battery.
The electrons flowing to the iron prevents rusting.
Metals higher in the electrochemical series donate electrons to the iron. Iron donates
electrons to metals lower in the electrochemical series.
You must be able to explain :
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A surface barrier to air and water can provide physical protection against corrosion.
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Painting
Greasing
Electroplating
Galvanising
Tin-plating
Coating with plastic
Sacrificial protection involves a metal higher in the electrochemical series supplying
electrons to iron.
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Electroplating
Galvanising
Tin-plating
Galvanising
Scrap magnesium
You must be able to give everyday examples of anti-corrosion methods.
Credit
Learning Outcomes
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Corrosion is an example of oxidation.
Water, oxygen and dissolved carbon dioxide, or another electrolyte, are required for
rusting.
Iron (II) ions can be further oxidised to give iron (III) ions.
Electrons lost by the iron during rusting are accepted by the water and oxygen to form
hydroxide ions.
You must be able to explain the reaction at the iron electrode in an iron/ carbon cell
using the colour change of the ferroxyl indicator and direction of electron flow.
You must be able to explain the reactions in a cell using the colour change of ferroxyl
indicator and direction of electron flow.
You must be able to explain the effect of scratching tinplate in increasing the rate of
rusting of iron.
Corrosion
General
Corrosion is a chemical reaction which involves the surface of the metal changing
from an element to a compound.
Corrosion of metals is very costly. Money is spent to replace corroded objects with
new ones, and limited world resources of metals are used up faster than necessary.
Iron, in the form of its alloy steel, is the most commonly used metal in the world. The
corrosion of iron is given a special name because it is so important to us. Rusting is the
special name given to the corrosion of iron.
Different metals corrode at different rates. The shiny surface of a piece of freshly
cut sodium will begin to corrode within seconds; a piece of cleaned magnesium ribbon
will show signs of surface corrosion within days; iron will corrode slowly over a period
of months or years and some metals such as gold do not corrode at all.
Metals corrode when they are left exposed to the atmosphere. There must be
substances in the atmosphere which are causing the corrosion.
Rusting can only happen if oxygen and water can both get to the iron together. As air
contains both, any iron object will rust if left outside unprotected. The brown solid
which forms (rust) is a compound of iron, oxygen and water.
The layer of rust is very weak and falls off the surface of the iron easily. The iron
underneath then starts to rust, until eventually the whole piece of metals corrodes.
Detecting Rust
General
Rusting of iron is a slow process. We need a quick method of detecting rusting so that
we can investigate it.
The corroded iron solution gives a blue colour with ferroxyl indicator because it
contains Fe2+ ions.
This should give us a clearer picture of what happens to iron atoms when rusting takes
place. In rusting the iron atoms change into iron (2+) ions by losing electrons.
Fe  Fe2+ + 2e
Indicators Again
Credit
Iron can form two different ions, Fe2+ and Fe3+,. The Fe2+ ion is formed when iron rusts.
But what about the other ion? Is Fe3+ also formed when iron rusts? To find out we must
find an indicator that tests for this ion.
Ammonium thiocyanate solution is an indicator which gives a blood red colour in the
presence of Fe3+ ions.
In rusting, the iron atoms initially change to Fe3+ ions by losing electrons. The iron
atoms are oxidised to iron (II) ions. Adding ammonium thiocyanate solution shows that
further oxidation must take place, in which Fe2+ ions are converted into Fe3+ ions.
Fe2+  Fe3+ + e
This is the second stage in the rusting process.
Ferroxyl indicator turns blue in the presence of iron (II) ions. This makes it a good
indicator for investigating rusting. It can tell use even more.
OH- (hydroxide) ions give a pink colour with ferroxyl indicator.
Corrosion And Redox
General
When anything loses electrons, oxidation is said to have taken place. When iron
corrodes, the iron atoms lose electrons and form Fe2+ ions.
Fe  Fe2+ + 2e
oxidation
An oxidation reaction is always accompanied by reduction reaction. Something has to
be present which takes away the electrons.
Both oxygen and water are required for corrosion. The combination of the oxygen and
water take electrons from the metal, in this case iron, in a complicated reaction which
results in the formation of hydroxide ions.
O2 + 2H2O + 4e  4OH-
reduction
The presence of OH- ions is indicated by obtaining a pink colour with ferroxyl indicator.
Corrosion is a redox process as it involves a flow of electrons from the metal to
oxygen and water. Under normal conditions, dissolved carbon dioxide in the water
(carbonic acid), acts as an electrolyte and allows this flow of electrons to take place.
Anything which speeds up the electron flow will speed up the corrosion process.
Salt or acid are better electrolytes than dissolved carbon dioxide and moth speed up
corrosion by allowing the electron transfer to take place more easily.
Example : salt spread on the roads increases the rate of corrosion of car body work.
Physical Protection
General/ Credit
If a metal is to corrode its surface must be in contact with both oxygen and water. If
a coating is put onto the surface of the metal it will acts as a barrier against corrosion.
This is called physical protection and stops the metal coming into contact with air and
water.
Painting is the most common method used to put a coating on large steel or iron
structures. Through time paint wears and it has to be redone. The Forth Rail Bridge
has a team of painters, employed full time, continuously repainting the bridge.
Paint can’t be used to protect the moving parts in machinery. It would just flake off.
Instead they are often coated in grease or oil.
Where paint is likely to be chipped by stones or splashed by water other coatings are
used. This happens on the underside of cars. If the roads have been salted corrosion
can be very fast. Underseal, a type of plastic coating is used. This slows down rusting
but does not stop it forever.
Metals can be coated with other metals which are less likely to corrode. Tin-plating is
an example of this. If steel is dipped into liquid tin, a layer of tin forms on the surface
of the steel. Food cans (tin cans) are often made in this way.
Tin-plating works well provided the tin remains unbroken. If the tin layer is scratched,
the steel corrodes quickly. This is because the iron, in the steel, is more reactive than
the tin of the coating.
The chromium plating on some golf clubs and the silver-plating of cutlery, are done
using a process called electroplating. The metal to be protected is made the negative
electrode in an electrolysis cell containing ions of the metal which will provide the
coating.
The positive ions of the plating metal are attracted towards the metal to be protected
where they gain electrons and form a coating.
Electroplating is done to give an attractive appearance and to provide protection
against corrosion
Chemical Protection
General
When metals corrode they lose electrons and form metal ions. In rusting iron atoms
lose electrons to form Fe2+ ions.
If we can give a metal electrons, perhaps we can slow down rusting.
Magnesium is more reactive than iron. It gives electrons to the iron joined to it
preventing it rusting. The magnesium corrodes very quickly.
Iron is more reactive than copper. It gives electrons to the copper joined to it,
preventing it corroding. The iron corrodes very quickly.
A metal higher in the electrochemical series can protect a metal lower in the
electrochemical series by giving it electrons. The corrosion of the more reactive
metal is speeded up. The more reactive metal is sacrificed to protect the other metal.
This method of protection is called sacrificial protection.
Magnesium is used to protect underground steel pipelines. Bags of magnesium are hung
at different points along the pipe. These are replaced at regular intervals as the
magnesium corrodes.
Galvanising is used to protect car exhausts, dustbins, corrugated sheeting and special
types of nails. Galvanised iron is iron coated with zinc.
This is done by dipping the iron into molten zinc. Galvanising is a particularly good form
of protection because it gives both physical protection (complete coating of zinc) and
chemical protection (zinc is more reactive than iron) to the iron. Even if scratched, the
zinc layer continues to protect the iron.
Linking to a more reactive metal is not the only method of supplying electrons. A
battery or d.c power supply can also be used. The negative terminal of the battery
supplies the electrons. Cars are usually given this type of protection from corrosion by
having the negative terminal of their battery joined to the car body.