Metals – What You Should Know
Corrosion is a chemical reaction in which the surface of a metal
changes into a compound. This causes the metal to be eaten away.
Metals like Mg, Zn and Fe corrode quickly whereas metals like Sn, Pb,
Cu, Ag and Au corrode slowly.
The corrosion of iron is known as rusting. Both water and oxygen must
be present for rusting to occur.
When iron rusts each iron atom loses two electrons to form Fe2+ ions.
Fe2+ ions are detected by ferroxyl indicator (sometimes stated as
Ferrox indicator). This indicator changes from yellow to blue in the
presence of Fe2+ ions. The depth of the blue colour can be used to
estimate the extent of rusting i.e. the darker the blue the more
rusting has taken place.
Salts and acids speed up the rate at which car bodies made of iron will
rust. Salts and acids produce solutions which contain ions i.e. they are
electrolytes which help to carry away the electrons when Fe2+ ions are
formed from Fe atoms. Acid rain acts in a similar way.
When iron is connected to the negative terminal of a D.C. power supply
it does not corrode. Using the negative terminal of a D.C. power supply
to protects a metal object is called cathodic protection. Oil rigs and
car bodies are examples that use cathodic protection. In both cases
the iron is protected by a supply of electrons flowing from the
negative terminal of the battery.
Metals can be protected against corrosion by stopping air and water
attacking the surface of the metal. This can be done by using a method
which places a coating on the surface of the metal e.g.
Protective Coating
Painting
Examples of Use
Bridges, buildings, cars, ships,
railings
Greasing
Bicycle chains, machine parts, car
door hinges
Electroplating
Computer contacts, cutlery,
jewellery
Galvanising
Dust bins, railings, crash barriers
Tin-plating
Food cans
Plastic coating
Lamp posts, wire fencing, drying
racks
Electroplating means coating metals with other metals, by using
electricity. In the example below the zinc ions in solution are changed
into zinc atoms and deposited on the surface of the iron nail. The
object to be plated is always connected to the negative terminal of the
supply.
The zinc strip connected to the positive terminal of the supply
maintains the concentration of Zn2+ ions in solution.
Galvanising means coating iron with zinc for protection. This is done by
dipping the iron object into molten zinc.
Tin-plating means coating iron objects with tin.
Different metals are often in contact with each other. If a metal
higher in the electrochemical series is connected to iron, then
electrons flow from the more reactive metal to the iron. The iron is
then protected by the more reactive metal. If a metal lower in the
electrochemical series is connected to iron, electrons flow from the
iron the less reactive metal. The iron then rusts (corrodes) to protect
the metal lower in the series.
Iron can be protected by connecting it to metals higher in the
electrochemical series, such as magnesium and zinc. Magnesium scrap is
connected to underground oil pipe lines, zinc is connected to ship’s
hulls. The iron is protected because electrons flow from the magnesium
or zinc to the iron and prevent corrosion of the iron. The same happens
if the zinc coating on galvanised iron objects is broken.
Electrolysis is the process of breaking down ionic compounds using
electricity. Chemical changes take place at the positive and negative
electrodes. The production of elements from ionic compounds using
electricity is an example of electrolysis. The metallic elements are
formed at the negative electrode and the non-metals are formed at
the positive electrode. Electrolysis can take place in molten ionic
compounds or in solutions of ionic compounds in water. Metals high in
the electrochemical series are extracted from their ores by
electrolysis of the molten ore.
Whenever two different metals are connected, a voltage is produced.
In general, the further apart the metals are in the reactivity series,
the greater the voltage produced. E.g. copper and magnesium produce a
larger voltage than copper and tin.
A listing of the metals in order of the voltage produced is called the
electrochemical series.
Electricity can be produced by connecting together two metals that
are dipped into solutions of their ions. The two beakers are then
connected using a salt bridge.
A salt bridge (or ion bridge) is a piece of filter paper soaked in an
electrolyte which allows ion conduction between beakers in a cell (i.e. it
completes the circuit).
In addition to water and oxygen another electrolyte such as dissolved
carbon dioxide or salt is required for rusting to take place.
Oxidation is a process in which electrons are lost or given away.
Corrosion is an example of oxidation because the metal atoms lose
electrons to form metal ions e.g.
Fe
Fe2+ + 2e-
The Fe2+ ions formed can be further oxidised to Fe3+ ions as shown
below:
Fe2+
Fe3+ + e-
The electrons lost by the iron are accepted by water and oxygen
molecules to form hydroxide ions.
2H2O + O2 + 4e-
4OH-
Iron is attacked by acid rain because iron is above hydrogen in the
electrochemical series. Iron can displace the hydrogen from acids,
leading to the formation of Fe2+ ions, i.e. the iron corrodes.
If a piece of tin-plated iron is scratched and the surface broken, the
iron corrodes faster than a piece of ordinary iron. This happens
because the tin is lower in the electrochemical series than iron. The
iron protects the tin by supplying electrons and forming Fe2+ ions, i.e.
the iron corrodes.
Steel is an alloy of carbon and iron. The corrosion of steel containing
carbon is faster than iron alone. This is so because a cell is set up
between the iron and the carbon. What happens is shown below:
If the carbon in the cell above is replaced by a metal lower in the
electrochemical series, electrons flow from the iron to the metal. This
results in the formation of Fe2+ ions which can be detected by the
ferroxyl indicator.
If the carbon is replaced with a metal higher in the electrochemical
series then electrons flow to the iron and no corrosion of iron takes
place. No Fe2+ ions are formed and the indicator stays yellow/green
around the nail.
Electrolysis can be used to purify metals such as copper. The process
and electrode reactions are shown below:
At the positive electrode copper atoms change into copper ions. This is
an example of OXIDATION.
Cu(s)
Cu2+(aq) + 2e
At the negative electrode copper ions in solution change into copper
atoms. This is an example of REDUCTION.
Cu2+(aq) + 2e
Cu(s)
Electricity can be produced by cells which have half-cells using nonmetals. An example is shown below:
Electrons flow from the zinc electrode to the carbon electrode and
react with hydrogen ions.
Oxidation is the loss of electrons by any reactant in a chemical
reaction. When a metal reacts to form a compound it loses electron(s).
An example of oxidation is: copper reacting with oxygen to form copper
oxide. The copper atoms are oxidised to copper ions.
Reduction is the gain of electrons by any reactant in a chemical
reaction. When a compound reacts to form a metal, the metal ions gain
electrons. An example of reduction is: copper oxide reacting with
hydrogen to produce copper metal and water. The copper ions in copper
oxide are reduced to copper atoms.
In a redox reaction both oxidation and reduction take place at the
same time. Examples of more complex redox reactions are shown
below:
Fe2+ / Br2 cell – this cell uses carbon electrodes, with oxidation
occurring in the Fe2+ solution and reduction occurring in the Br2
solution.
Fe2+
Fe3+ + e
Br2 + 2e
2Br-
(OXIDATION)
(REDUCTION)
Cells with non-metals and negative ions can also be set up, again using
carbon electrodes: