The Behaviour of Metals
It is the free electrons (delocalised electrons) which give metals their
properties. You should remember that metals are strong, malleable, ductile,
sonorous and generally high density. There are also thermal and electrical
conductors and mostly have high melting points.
The free electrons hold the structure together as if the metal ions were in a
"glue" (called a "sea of electrons"). Delocalised electrons hold the ions
together by strong electrostatic forces. This gives metals high melting
points and boiling points. They are all solid at room temperature except
mercury, which is liquid. The free electrons also cause metals to conduct
heat and electricity.
The free electrons allow the metal ions to slide over each other. This makes
metals malleable which means that they are soft, easily bent and shaped, and
can be pressed or beaten into thin sheets. Metals are ductile which means
that they can be drawn down into wires.
The reaction of metals with air (oxygen)
Potassium, sodium, lithium, calcium and magnesium react with oxygen and will
burn in air. Metals in the reactivity series from aluminium to copper do not
burn but may react slowly with oxygen in the air to form the metal oxide.
Metal oxides are bases and can be used to neutralise acids forming salts and
water. Aluminium is the fastest and copper is the slowest of the six.
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Silver, gold and platinum do not react at all with oxygen and so remain shiny
and attractive.
Zinc reacts reasonably quickly with oxygen to form zinc oxide.
Zinc + oxygen
2Zn(s)
zinc oxide.
+ O2(g)
2ZnO(s)
The reaction of metals with water
Potassium, sodium, lithium and calcium react vigorously with cold water to
form metal hydroxides. Hydroxides are basic/alkaline in nature and group 1
are known as the alkali metals.
Metals in the reactivity series from magnesium to iron react with steam
(H2O(g)) but not water (H2O(l)). The reaction of magnesium or iron with steam
forms the metal oxide and hydrogen gas. You should know the test for
hydrogen gas. This gas ‘ignites with a squeaky pop’.
Sodium +
2Na(s)
Magnesium +
Mg(s)
water
+ 2H2O(l)
sodium hydroxide + hydrogen.
2NaOH(aq) + H2(g)
steam
+ H2O(g)
magnesium oxide + hydrogen.
MgO(s) +
H2(g)
Tin, lead, copper, silver, gold and platinum do not react with water or steam.
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The reaction of metals with dilute acid
Potassium, sodium, lithium and calcium all react violently with dilute
hydrochloric acid. It is dangerous to put these metals into an acid.
The reaction is similar to the reaction with water, forming the metal
chloride and H2(g). The reaction of zinc with acid is often used to
make a small amount of hydrogen in the laboratory. Iron, tin and lead all
react slowly with acids whilst metals below hydrogen in the reactivity series
(copper, silver, gold and platinum) will not react with dilute acids.
For example
Magnesium + hydrochloric acid
Mg(s) +
2HCl(aq)
magnesium chloride + hydrogen.
MgCl2(aq) +
H2(g)
Why is aluminium different?
Aluminium is a reactive metal but does not appear so. It reacts quickly with
the air to produce a very thin oxide layer on its surface which then acts to
protect itself. You can drop the metal in acid and you won’t see any reaction
as acid does not actually touch any aluminium.
Displacement Reactions
A metal will displace (take the place of) a less reactive metal in a metal salt
(solid or solution). The element carbon can do this also.
For example,
magnesium + copper sulfate
Mg(s) + CuSO4(aq)
magnesium sulfate +copper.
MgSO4(aq) + Cu(s)
Copper sulfate is blue, magnesium sulfate is colourless. During the reaction
the blue solution loses its colour and the magnesium metal is seen to turn
brown as the displaced copper metal becomes deposited on its surface.
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In these displacement reactions the metals are competing for the non-metal
anion (negative ion). In the above examples the non-metal anion is sulfate SO42-. Since magnesium is more reactive, it wins the SO42-. The order of the
metals in the reactivity series can be worked out by using these type of
reactions.
The above is a REDOX reaction since magnesium is going from magnesium
atoms to magnesium ions by losing electrons. At the same time copper ions
are gaining the electrons and turning to atoms. The overall ionic equation is
Mg(s) + Cu2+(aq)
Mg2+ (aq) + Cu(s)
If a less reactive metal is added to a metal salt solution there will be no
reaction – nothing will happen! Iron is less reactive than zinc so
iron + zinc chloride
no reaction.
If carbon is used instead
lead oxide + carbon
2PbO(s)
carbon dioxide + lead
+ C (s)
CO2(aq)
+
2Pb(s)
You should be familiar with the reactions between metals and acids. These
too are displacement reactions. A reactive metal can displace hydrogen has
from an acid.
zinc + hydrochloric acid
Zn + 2HCl
zinc chloride + hydrogen
ZnCl2 + H2
Look at the reactivity series below to see that zinc is above hydrogen.
Since copper is below hydrogen in the reactivity series, copper will not
displace hydrogen from an acid so there is no reaction.
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The reactivity series of metals
Reactivity
The reactivity of a metal is a measure of how readily the metal becomes a
compound so that its atoms gain a full outer shell. In the case of metals this
means how readily a metal loses electrons to form positive ions. Once a
reactive metal has formed a compound these compounds are extremely
stable and it is very difficult to obtain the metal on its own once again. This
means that extracting a reactive metal from its ore is very difficult.
Whereas the less reactive the metal, the less stable its compounds and so
the easier it is to obtain from its ore.
This can be illustrated further by considering the following thermal
decomposition reactions.
Many carbonates thermally decompose to form metal oxides and CO2.
CuCO3(s)
CuO (s) + CO2 (s)
Sodium and potassium carbonates do not decompose because they are
particularly stable.
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Most hydroxides thermally decompose to form metal oxides and H2O.
Cu(OH)2(s)
CuO (s) + H2O (l)
Sodium and potassium hydroxides do not decompose again because there are
particularly stable.
Nitrates thermally decompose to form metal oxides, O2
2Cu(NO3)2(s)
2CuO (s) + O2
(g)
(g)
and NO2.
+ 4 NO2 (g)
Sodium and potassium hydroxides do decompose slightly to make nitrites and
O2. Be Careful here - Nitrites (whose formulae end NO2) are not the same as
nitrates (whose formulae end NO3)!
2Na(NO3)(s)
2NaNO2 (s) + O2
(g)
Using the difference in reactivity of two metals
Thermite reaction
We can make use of the difference in the reactivity of two metals in a
number of different ways. Railway lines can be joined together using the
‘thermite reaction’ where aluminium displaces iron from its oxide in a very
exothermic reaction. The liquid iron produces flows into the gap between the
rails joining them when it cools.
Fe2O3(s) +
2Al(s)
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Al2O3 (s) + 2Fe (l)
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Cells
A simple cell consists of two different metals in a
solution called an electrolyte. Under these conditions a
voltage is produced. Electrons flow from the more
reactive metal (making it the negative electrode) to the
less reactive one (which is positive). The larger the
difference in reactivity, the higher the voltage will be.
Sacrificial protection
Iron is a very useful metal with one
major drawback; it rusts. Rusting is
the reaction of iron with water and
oxygen. To prevent rusting a more
reactive metal, such as blocks of zinc
or magnesium, is welded to the
structure made of iron and the zinc
reacts in preference to the iron. The
image shows a block of zinc welded
to the hull of a ship.
Galvanising
Galvanising is another way of using
zinc to protect iron by completely
coating iron or steel (an alloy
containing mostly iron) with a thin
layer of zinc. The zinc coating
means that the iron or steel is not
exposed to water and oxygen in
the air and so can not rust. Most
food cans are made of steel which
has been galvanised.
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Topic 4:The Behaviour of Metals
Summary questions
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