Topic 12 – Corrosion
All metals corrode.
Corrosion is a chemical reaction on the surface of a metal.
During corrosion, atoms of elements change into compounds.
Rusting is the name given to the corrosion of iron/steel.
Different metals corrode at different rates.
Gold may be found native (uncombined as the element) in the earth’s crust.
Gold corrodes so slowly that it virtually does not corrode at all.
Iron corrodes quite rapidly.
Iron appears rusty after a few hours, although it may take years to
completely corrode.
Sodium reacts very rapidly in air and corrosion can be complete in minutes.
When metals form compounds they lose electrons to form positive ions.
Corrosion involves the formation of metal ions:
M (s)
M+ (aq)
+
e-
During corrosion, metals undergo oxidation.
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Metal nails will corrode quicker in some conditions than in others:
Amount of corrosion:
none
none
a little
a lot
a lot
oil
calcium chloride
(dry air, oxygen
only)
boiled water
(oil prevents
air getting in)
air (oxygen)
& water
sea
water
soda water
(CO2)
Corrosion occurs when both water and oxygen are present.
An electrolyte is also required for corrosion to occur.
An example of an electrolyte is dissolved carbon dioxide
Corrosion is faster when salt is present in water.
The salt dissolves in the water forming an electrolyte.
Ions in the solution help carry the current speeding up the rate of corrosion.
Cars driven on salted winter roads should be washed as soon as possible.
If the salt is not removed the car will rust more rapidly.
Corrosion also occurs more quickly when an acid is present.
Acids solutions contain ions so are electrolytes.
Pollution, in the form of acid rain, is a major cause of increased rusting.
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When iron rusts, initially iron atoms lose two electrons to form iron (II) ions:
Fe (s)
Fe2+ (aq) + 2e-
This can be shown using ferroxyl indicator.
This indicator turns blue in the presence of Fe2+ (aq) ions.
The iron (II) ions can be further oxidised to give iron (III) ions:
Fe2+ (aq)
Fe3+ + e
The electrons iron loses are accepted by water and oxygen forming
hydroxide ions:
2H2O (l) + O2 (g) + 4e-
4OH- (aq)
This equation explains why both water and oxygen are needed for rusting.
Ferroxyl indicator turns pink in the presence of hydroxide ions.
Direct electrical protection
Iron does not rust when attached to the negative terminal of a battery.
Electrons flowing from the battery to the iron prevent rusting.
This method of preventing corrosion is used to protect car
bodies from rusting.
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This method is also used to protect piers and jetties and some ocean liners
when they are in port.
Sacrificial Protection
When 2 metals are connected together, the metal higher in the
electrochemical series loses electrons to the metal lower in the series.
Connecting iron to a metal higher in the electrochemical series prevents
rusting.
E.g. magnesium gives up electrons to iron.
Connecting iron to a metal lower in the series makes rusting occur quicker
than if iron was on its own.
E.g. copper nails will make an iron roof rust faster.
Iron pipelines, e.g. transporting gas, are protected by bags of scrap
magnesium attached to them at intervals.
The magnesium corrodes pushing electrons onto the iron, thus
preventing corrosion of the pipe.
Ships hulls, piers and oilrigs are also protected in this way.
Electroplating
Electroplating is coating a metal with a very thin film of another metal.
This technique uses electricity.
The material is placed in a solution containing ions of the metal
used for coating.
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The material to be plated is made the negative electrode.
Metal ions are positively charged so are attracted to the negative electrode.
The ions gain electrons at the negative electrode, forming solid metal.
E.g.
+
d.c. supply Ag+
metal to be coated
solution of silver ions
silver strip corrodes away
(The mass of the silver electrode will decrease.
The other metal’s mass will increase until all of the silver is used up).
Galvanising
Galvanising involves coating iron in zinc.
During galvanising the iron object is dipped into molten zinc.
Zinc is higher than iron in the electrochemical series.
Even if the zinc coating is scratched, zinc will sacrificially protect iron.
Tin-plating
Tin-plating involves coating iron with tin.
This can be done either by electroplating or dipping objects into molten tin.
Iron is higher in the electrochemical series than tin.
If tin-plating is scratched, iron will sacrificially protect tin.
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Scratched tin-plating will make iron rust faster than iron without tin-plating.
Preventing oxygen (air) and water from reaching a metal object will
prevent corrosion.
For this reason many metal objects are given a coating.
This acts as a physical barrier between the metal and water and air.
Examples of physical protection include:
painting
Forth rail bridge; cars; bikes; etc.
greasing
tools; machinery.
electroplating
car bumpers; bike wheels and handlebars.
galvanising
dustbins; corrugated iron.
tin-plating
food cans.
coating with plastic
household goods e.g. wire draining boards.
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iron/carbon cell
e-
iron/magnesium cell
e-
e-
mA
e-
iron/tin cell
e-
e-
mA
mA
iron
iron
iron
magnesium
carbon
ferroxyl
indicator
tin
ferroxyl
indicator
ferroxyl
indicator
blue colour
pink colour
Iron/carbon cell:
Iron rusts.
Blue colour around iron indicates Fe2+ (aq).
Pink colour around carbon indicates OH- (aq).
Iron/magnesium cell: Iron does not rust.
No blue colour.
Pink colour around magnesium.
Iron/tin cell:
Iron rusts.
Blue colour around iron.
Pink colour around tin.
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