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METALS
IMPORTANCE OF METALS
More then three quarters of elements in the periodic table are metals.
We cook food, heat the food on metals stoves.
We travel in metallic vehicals like car, airplanes, ships.
The concreate wall and roofs of houses are reinforced with metal rods.
Coins, taps, cutlary, pins, needles paper pins etc. are all made up of metals.
PHYSICAL PROPERTIES OF METALS
PROPERTY
REASON
Metals are malleable and ductile i.e. they can be
drawn into sheets and wires
layers of metals can slide over each other.
Metals conduct electricity.
they have free moving electrons in their
outer most shell.
Metals usually have high melting points, high
boiling points and high densities
Strong metallic bond
CHEMICAL PROPERTIES OF METALS
Reaction with water
Potassium reacts vigrously with cold water to form salt and hydrogen gas. The reaction is so exothermic
that the hydrogen gas produced, burn in air.
Potassium + water ---------- Potassium hydroxide + Hydrogen
2K(s) + 2H
2
O (l) ------------- 2KOH(aq) + H
2
(g)
Sodium reacts with cold water in the same way.
Sodium + water -------------- Sodium hydroxide + hydrogen gas
2Na(s) + 2H
2
O (l) ------------- 2NaOH(aq) + H
2
(g)
Calcium reacts readily with cold water and vigrously with hot water to produce salt and hydrogen gas.
Calcium + water ------------ Calcium hydroxide + hydrogen gas
Ca(s) + 2H
2
O (l) ------------- Ca(OH)
2
(aq) + H
2
(g)
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Magnesium reacts very slowly with cold water but vigrously with steam to form salt and hydrogen gas.
Magnesium + steam --------- Magnesium oxide + hydrogen gas
Mg(s) + H
2
O (g) ------------- MgO(s) + H
2
(g)
Zinc do not react with cold water but reacts slowly with steam to form zinc oxide and hydrogen gas.
Zink + steam --------------- Zink oxide + hydrogen gas
Zn(s) + H
2
O (g) ------------- ZnO(s) + H
2
(g)
Iron do not react with cold water but rusting occur very slowly in the presence of oxygen. Red hot iron
reacts very slowly with steam to produce salt and hydrogen gas.
Iron + steam --------------- Iron oxide + hydrogen
3Fe(s) + 4H
2
O (g) ------------- Fe
3
O
4
(s) + 4H
2
(g)
Copper do not react with water under any condition
Silver do not react with water in any condition.
Reaction with Hydrochloric acid
Potassium and sodium reacts explosively to form salt and hydrogen gas. The reaction is so exothermic
that the hydrogen gas produced, burn in air.
Potassium + Hydrochloric acid ------ Potassium chloride + hydrogen
2K(s) + 2HCl (aq) ------------- 2KCl(aq) + H
2
(g)
Sodium + hydrochloric acid ------ Sodium chloride + hydrogen
2Na(s) + 2HCl (aq) ------------- 2NaCl(aq) + H
2
(g)
Calcium reacts vigorously` to produce calcium chloride and hydrogen gas.
Calcium + hydrochloric acid ------- Calcium chloride + hydrogen gas
Ca(s) + 2HCl (aq) ------------- CaCl
2
(aq) + H
2
(g)
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Magnesium reacts very fastly to form magnesium chloride and hydrogen gas.
Magnesium + hydrochloric acid -------- Magnesium chloride + hydrogen gas
Mg(s) + 2HCl (aq) ------------- MgCl
2
(aq) + H
2
(g)
Zinc reacts moderately to form zinc chloride and hydrogen gas.
Zink + hydrochloric acid ------ Zink chloride + hydrogen gas
Zn(s) + 2HCl (aq) ------------- ZnCl
2
(aq) + H
2
(g)
Iron reacts slowly to produce iron chloride and hydrogen gas.
Iron + hydrochloric acid -------- Ironchloride + hydrogen gas
Fe(s) + 2HCl (aq) ------------- FeCl
2
(aq) + H
2
(g)
Copper do not react with dilute HCl
Silver do not react with dilute HCl
Reaction with oxygen
Potassium tarnishes in the presence of oxygen to form potassium oxide K
2
O
Potassium + oxygen ------- Potassium oxide
4K(s) + O
2
(g) ------------------- 2K
2
O(s)
Sodium burns with a yellow flame to produce odium oxide Na
2
O
Sodium + Oxygen --------------- Sodium Oxide
4 Na(s) + O
2
(g) ------------- 2 Na
2
O(s)
Copper powder burns with dull red glow to form copper oxide. CuO
Copper + Oxygen ----------------- Copper oxide
2Cu(s) + O
2
(g) ------------------- 2CuO(s)
Iron powder or wire burns with a bright yellow flame to form iron oxide Fe
3
O
4
Iron + Oxygen ---------------- Iron oxide
Fe(s) + O
2
(g) ------------------- 2Fe
3
O
4
(s)
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Magnesium burns with a bright white flame to produce white solid magnesium oxide. MgO
Magnesium + oxygen ----------- Magnesium oxide
2Mg(s) + O
2
(g) ------------------- 2MgO(s)
REACTIVITY SERIES
Metals can be arranged in order of their chemical reactivity. The reactivity series is based on the reaction
of metals with water or dilute hydrochloric acid. When metal recats with acid or water, the metal atom
lose electron to become ion.
Metal(s) + H
2
O (l) ------------------ Metal
+
ion + OH
-
ion +
H
2
(g)
Metal (s) + HCl (aq) ---------------- Metal
+
ion + Cl
-
ion +
H
2
(g)
The more readily a metal gives up electrons to form ions, the more reactive it is.
A metal high up in the reactivity series
Reacts vigorously with chemicals
Readily gives up electrons in reactions to form positive ions
Corrode easily
A metal low down in the reactivity series
Does not Reacts vigorously with chemicals
Does not Readily gives up electrons in reactions to form positive
ions
Does not Corrode easily
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Hydrogen is sometimes placed in the reactivity series. Metals below hydrogen in the series do not react
with acids to produce hydrogen gas.
Please -------- Potassium
Send -------- Sodium
Cats ------- Calcium
Monkeys -------Magnesium
And -------Aluminium
Zebras -------Zinc
In ------Iron
Large ----- Lead
Hired ------ Hydrogen
Cages ------Copper
Make ---- Mercury
Sure -----Silver
Good -----Gold
Padlock ----- Platinium
DISPLACEMENT OF METALS
Displacement of metals from solutions
A more reactive metal will displace the ions of any less reactive metal in the reactivity series, from
solution.
Zinc + copper (II) sulphate solution ----------- Copper + zinc sulphate solution.
Zn (s) + CuSO
4
(aq) -------------------------- Cu(s) + ZnSO
4
(aq)
Zinc displace copper from the copper sulphate solution because it is more reactive than copper and readily
give up electrons to form positive ions. The electrons are transferred from zinc atom to copper (II) ions.
Cu
2+
(aq) + Zn(s) ---------------- Cu (s) + Zn
2+
(aq)
blue solution redish-brown solid colourless
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Other examples:
Ag
+
(aq) + Cu (s) ------------------
Zn
2+
(aq) + Mg (s) ----------------
Displacement of metals from metallic oxides by more reactive metals
A metal will take oxygen from the oxide of any metal below it in the reactivity series. For example, when
magnesium powder and copper (II) oxide powder is heated there is a vigrous exothermic reaction. The
magnesium takes oxygen from copper (II) oxide to from magnesium oxide and copper metal.
Magnesium + Copper oxide ----------- Magnesium oxide + Copper
Mg(s) + CuO(s) ------heat------------- MgO(s) + Cu(l)
Thermite reaction reaction.
Aluminium + Iron oxide ---------------- Iron + Aluminium oxide
2Al (s) + Fe
2
O
3
(s) -------heat--------- 2Fe(l) + Al
2
O
3
(s)
Reaction of metallic oxides with hydrogen
Hydrogen can take oxygen from metallic oxides, producing the metal and water. For example when
hyrogen is passed over hot lead (II) oxide, lead metal and water are produced.
Lead (II) oxide + hydrogen -------------- lead + water.
PbO (s) + H
2
(g) ---------------- Pb(s) + H
2
O (l)
Copper (II) oxide + hydrogen ----heat--------- copper + water.
CuO (s) + H
2
-------heat------------ Cu (s) + H
2
O (l)
The less reactive the metal, the easier it is for hydrogen to take oxygen from its oxide. The oxides of vary
recative metals such as aluminium oxide and sodium oxide cannot be reduced to the metal by hydrogen.
Reaction of metallic oxides with carbon.
Carbon can take up oxygen from the oxide of metals which are not too high in the reactivity series. For
example a mixture of charcoal and copper (II) oxide reacts when heated together
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Copper (II) oxide + Carbon -----------heat---------- copper + carbon dioxide.
2CuO (s) + C(s) -------heat--------- 2Cu (s) + CO
2
(g)
The more reactive the metal the more harder it for carbon to take oxygen from its oxide. Iron is more
reactive than copper, iron oxide and carbon must be heated very strongly for the reaction to take place.
Iron (II) oxide + Carbon -----strong heating------------- Iron + carbon dioxide.
2FeO (s) + C (s) --------strong heating---------- 2Fe + CO
2
Carbon is unable to take oxygen from the oxides of very reactive such as calcium and sodium.
THE EXTRACTION OF METALS
Most of the metals are found as compounds called minerals. Minerals are usually found mixed with large
amounts of impurities. These impure minerals are called ores.
A ROCK is a mixture of minerals from which useful substances can be made.
A MINERAL is a solid element or compound found naturally in the Earths crust.
A METAL ORE is a mineral or mixture of minerals from which economically viable amounts of metal
can be obtained. Two important ores to know:
Haematite for Iron [contains iron(III) oxide, Fe
2
O
3
]
Bauxite for Aluminium [contains aluminium oxide, Al
2
O
3
]
Some important minerals
Name of Mineral
Chemical Name
Formula
Metal extracted
Usual method of Extraction
Bauxite Aluminium oxide Al
2
O
3
Aluminium Electrolysis of oxide dissolved
in molten cryolite.
Galena Lead sulphide PbS Lead Sulphide is roasted in air and
the oxide produced is
Haematite Iron (III) oxide Fe
2
O
3
Iron Heat oxide with carbon
Sphalerite Zinc Sulphide ZnS Zinc Sulphide is roasted in air and
the oxide produced is heated
with carbon.
Copper pyrite Copper iron
sulphide
CuFeS
2
Copper Sulphide ore is roasted in air
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Introduction
The Earth's crust contains many different rocks. Rocks are a mixture of minerals and from some we
can make useful substances.
A mineral can be a solid metallic or non-metallic element or a compound found naturally in the
Earth's crust.
A metal ore is a mineral or mixture of minerals from which economically viable amounts of
metal can be extracted, i.e. its got to have enough of the metal, or one of its compounds, in it to be
worth digging out! Ores are often oxides, carbonates or sulphides. They are all finite resources so
we should use them wisely!
In order to extract a metal, the ore or compound of the metal must undergo a process called
reduction to free the metal (i.e. the positive metal ion gains negative electrons to form the
neutral metal atom, or the oxide loses oxygen, to form the free metallic atoms).
Generally speaking the method of extraction depends on the metals position in the reactivity
series.
The reactivity series of metals can be presented to include two non-metals, carbon and
hydrogen, to help predict which method could be used to extract the metal.
o lower Pt Au Ag Cu (H) Pb Sn Fe Zn (C) Al Mg Ca Na K higher in series
o RULE: Any element higher in the series can displace any other lower element
Metals above zinc and carbon in the reactivity series cannot usually be extracted with carbon or
carbon monoxide. They are usually extracted by electrolysis of the purified molten ore or other
suitable compound
o eg aluminium from molten aluminium oxide or sodium from molten sodium chloride.
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o The ore or compound must be molten or dissolved in a solution in an electrolysis cell to
allow free movement of ions (electrical current).
Metals below carbon can be extracted by heating the oxide with carbon or carbon monoxide.
The non-metallic elements carbon will displace the less reactive metals in a smelter or blast
furnace e.g. iron or zinc and metals lower in the series.
o Metals below hydrogen will not displace hydrogen from acids. Their oxides are easily
reduced to the metal by heating in a stream of hydrogen, though this is an extraction method
rarely used in industry. In fact most metal oxides below carbon can be reduced when heated in
hydrogen, even if the metal reacts with acid.
Some metals are so unreactive that they do not readily combine with oxygen in the air or any other
element present in the Earth's crust, and so can be found as the metal itself. For example gold (and
sometimes copper and silver) and no chemical separation or extraction is needed. In fact all the metals
below hydrogen can be found as the 'free' or 'native' element.
Other methods are used in special cases using the displacement rule. A more reactive metal can be
used to displace and extract a less reactive metal but these are costly processes since the more
reactive metal also has to be produced in the first place! See Titanium or see at the end of the section
on copper extraction
Sometimes electrolysis is used to purify less reactive metals which have previously been extracted
using carbon or hydrogen (eg copper and zinc). Electrolysis is also used to plate one metal with
another.
The demand for raw materials does have social, economic and environmental implications eg
conservation of mineral resources by recycling metals, minimising pollution etc.
Historically as technology and science have developed the methods of extraction have improved to
the point were all metals can be produced. The reactivity is a measure of the ease of compound
formation and stability (ie more reactive, more readily formed stable compound, more difficult to reduce
to the metal).
o The least reactive metals such as gold, silver and copper have been used for the past 10000
years because the pure metal was found naturally.
o Moderately reactive metals like iron and tin have been extracted using carbon based
smelting for the past 2000-3000 years.
o BUT it is only in the last 200 years that very reactive metals like sodium or aluminium have
been extracted by electrolysis.
21.2 Metallurgy
Metallurgy is the combination of science and technology used to extract metals from their ores. Ores
are complex mixtures of metal-containing material and useless impurities called gangue. The steps
involved in extracting a metal include the following:
concentrating the ore, and chemically treating it if necessary
reducing the mineral to free metal
refining and purifying the metal.
The metal may be mixed with other elements to modify its properties or to form an alloy, a metallic
solution of two or more elements
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Concentration and Chemical Treatment of Ores
The different physical properties of the mineral and the gangue, such as density and magnetic charge,
can be used to concentrate the mineral and remove the gangue. Metal sulfide ores are concentrated by
flotation, a process that exploits differences in the ability of water and oil to wet the surfaces of the
mineral and the gangue. Mineral particles float to the top of the tank along with soapy air bubbles, while
the gangue sinks to the bottom.
Ores can also be concentrated by chemical means. In the Bayer process, the Al
2
O
3
in bauxite is
separated from Fe
2
O
3
impurities by treating the ore with NaOH.
Roasting, or heating in air, is another chemical treatment used to convert minerals to compounds that
are more easily reduced to the metal.
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Reduction
The more active metals are obtained by reducing their ores with a chemical reducing agent. Zinc is
obtained by reducing ZnO with coke, a form of carbon.
Carbon cannot be used to reduce metals that form stable carbides, such as tungsten. Tungsten(VI) oxide
is reduced with hydrogen gas.
The most active metals cannot be reduced with chemical reducing agents, so these metals are produced
by electrolytic reduction,.
Refining
The metals obtained from reducing ores generally require purification. Some metals, including zinc, can
be purified by distillation. Nickel is purified using the Mond process, a chemical method in which
Ni(CO)
4
is formed and then decomposed at a higher temperature. The equilibrium shift at the higher
temperature favors pure nickel.
Extraction
of Metal
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The Extraction of Iron
The solid mixture of haematite ore, coke and limestone is continuously
fed into the top of the blast furnace.
The coke is ignited at the base and hot air blown in to burn the coke
(carbon) to form carbon dioxide in an oxidation reaction (C gains O).
The heat energy is needed from this very exothermic reaction to raise
the temperature of the blast furnace to over 1000
o
C to effect the ore
reduction. The furnace contents must be heated.
o carbon + oxygen ==> carbon dioxide
o C
(s)
+ O
2(g)
==> CO
2(g)
at high temperature the carbon dioxide formed, reacts with more coke
(carbon) to form carbon monoxide
o carbon dioxide + carbon ==> carbon monoxide
o CO
2(g)
+ C
(s)
==> 2CO
(g)
o (note: CO
2
reduced by O loss, C is oxidised by O gain)
The carbon monoxide is the molecule that actually removes the oxygen
from the iron oxide ore. This a reduction reaction (Fe
2
O
3
loses its O, or
Fe
3+
gains three electrons to form Fe) and the CO is known as the reducing
agent (the O remover and gets oxidised in the process).
This frees the iron, which is molten at the high blast furnace temperature,
and trickles down to the base of the blast furnace. The main reduction
reaction is ...
o iron(III) oxide + carbon monoxide ==> iron + carbon
dioxide
o Fe
2
O
3(s)
+ 3CO
(g)
==> 2Fe
(l)
+ 3CO
2(g)
o note, as in the two reactions above, oxidation and reduction
always go together!
Other possible ore reduction reactions are ...
Fe
2
O
3(s)
+ 3C
(g)
==> 2Fe
(l)
+ 3CO
(g)
2Fe
2
O
3(s)
+ 3C
(g)
==> 4Fe
(l)
+ 3CO
2(g)
The original ore contains acidic mineral impurities such as silica (SiO
2
,
silicon dioxide). These react with the calcium carbonate (limestone) to
form a molten slag of e.g. calcium silicate.
o calcium carbonate + silica ==> calcium silicate + carbon dioxide
o CaCO
3
+ SiO
2
==> CaSiO
3
+ CO
2
o this is sometimes shown in two stages:
CaCO
3
==> CaO + CO
2
CaO + SiO
2
==> CaSiO
3
The molten slag forms a layer above the more dense molten iron and they
can be both separately, and regularly, drained away. The iron is cooled and
cast into pig iron ingots OR transferred directly to a steel producing furnace.
Iron from a blast furnace is ok for very hard cast iron objects BUT is too
brittle for many applications due to too high a carbon content from the coke.
So it is converted into steel alloys for a wide range of uses.
The waste slag is used for road construction or filling in quarries which
can then be landscaped or making cement.
Raw Materials:
Iron Ore eg
haematite ore
[iron(III) oxide,
Fe
2
O
3
]
coke (carbon, C)
hot air (for the O
2
in it)
limestone (calcium
carbonate, CaCO
3
)
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21.3 Iron and Steel
The cast iron or pig iron produced in a blast furnace must be purified. In the basic oxygen process,
molten iron from the blast furnace is mixed with pure oxygen gas in a furnace lined with basic
oxides. The impurities in the iron are oxidized and the acidic oxides react with CaO to yield a
molten slag. Phosphorus impurities react in this process to form a calcium phosphate slag.
The Extraction of Aluminium
The purified bauxite ore of aluminium oxide
is continuously fed in. Cryolite is added to
lower the melting point and dissolve the ore.
Ions must be free to move to the
electrode connections called the cathode (-
), attracting positive ions eg Al
3+
, and the
anode (+) attracting negative ions eg O
2-
.
When the d.c. current is passed through
aluminium forms at the positive cathode
(metal*) and sinks to the bottom of the tank.
At the negative anode, oxygen gas is
formed (non-metal). This is quite a problem.
At the high temperature of the electrolysis cell it
burns and oxidises away the carbon electrodes
to form toxic carbon monoxide or carbon
dioxide. So the electrode is regularly
replaced and the waste gases dealt with!
It is a costly process (6x more than Fe!) due to
the large quantities of expensive electrical
energy needed for the process.
Raw materials for the electrolysis process:
Bauxite ore of impure aluminium oxide
The redox details of the electrode processes:
At the negative (-) cathode, reduction
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[Al
2
O
3
made up of Al
3+
and O
2-
ions]
Carbon (graphite) for the electrodes.
Cryolite reduces the melting point of the ore
and saves energy, because the ions must be
free to move to carry the current
Electrolysis means using d.c. electrical
energy to bring about chemical changes
eg decomposition of a compound to
form metal deposits or release gases. The
electrical energy splits the compound!
At the electrolyte connections called the
anode electrode (+, attracts - ions) and
the cathode electrode (-, attracts + ions).
An electrolyte is a conducting melt or
solution of freely moving ions which
carry the charge of the electric current.
occurs (electron gain) when the positive
aluminium ions are attracted to it. They gain
three electrons to change to neutral Al atoms.
Al
3+
+ 3e
-
==> Al
At the positive (+) anode, oxidation takes
place (electron loss) when the negative
oxide ions are attracted to it. They lose two
electrons forming neutral oxygen molecules.
2O
2-
==> O
2
+ 4e
-
Note: Reduction and Oxidation always go
together!
The overall electrolytic decomposition is ...
o aluminium oxide => aluminium +
oxygen
o 2Al
2
O
3
==> 4Al + 3O
2
o and is a very endothermic process,
lots of electrical energy input!
The original extraction of copper from copper ores
from copper carbonate ores* ...
o The ore can be roasted to concentrate the copper as its oxide.
o Water is driven off and the carbonate thermally decomposed.
o copper(II) carbonate ==> copper oxide + carbon dioxide
o CuCO
3(s)
==> CuO
(s)
+ CO
2(g)
o The oxide can be smelted by heating with carbon (coke, charcoal) to reduce the
oxide to impure copper, though this method isn't really used much these days (the 'bronze
age' method archaeologically!).
o copper(II) oxide + carbon ==> copper + carbon dioxide
o 2CuO
(s)
+ C
(s)
==> 2Cu
(s)
+ CO
2(g)
from copper sulphide ores ...
o copper sulphide ores can roasted in air to form impure copper
o nasty sulphur dioxide gas is formed, this must be collected to avoid pollution and can be used
to make sulphuric acid to help the economy of the process
o copper(I) sulphide + oxygen ==> copper + sulphur dioxide
o Cu
2
S
(s)
+ O
2(g)
==> 2Cu
(s)
+ SO
2(g)
sulphur dioxide is a nasty toxic acidic gas, it is collected and used to make sulphuric
acid, helps pay for the extraction process.
o or *CuS
(s)
+ O
2(g)
==> Cu
(s)
+ SO
2(g)
* the CuS might be part of an ore like chalcopyrite CuFeS
2
which is the principle ore
copper is extracted from.
* It is also possible to dissolve the carbonate ore or the oxide from roasted ore in dilute
sulphuric acid and extracting copper by ....
o (1) using electrolysis see purification by electrolysis above. or
o (2) by adding a more reactive metal to displace it eg scrap iron or steel is used by
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adding it to the resulting copper(II) sulphate solution.
iron + copper(II) sulphate ==> iron(II) sulphate + copper
Fe
(s)
+ CuSO
4(aq)
=> FeSO
4(aq)
+ Cu
(s)
The Extraction of Titanium by Displacement
Titanium ore is mainly the oxide TiO
2
, which is converted into titanium tetrachloride TiCl
4
The chloride is then reacted with sodium or magnesium to form titanium metal and sodium
chloride or magnesium Chloride.
This reaction is carried out in an atmosphere of inert argon gas so non of the metals
involved becomes oxidised by atmospheric oxygen.
TiCl
4
+ 2Mg ==> Ti + 2MgCl
2
or TiCl
4
+ 4Na ==> Ti + 4NaCl
These are examples of metal displacement reactions eg the less reactive titanium is
displaced by the more reactive sodium or magnesium.
Overall the titanium oxide ore is reduced to titanium metal (overall O loss, oxide => metal)
Environmental Impact and Economics of Metal and other Mineral Extraction
One of the problems of metal or mineral extraction is balancing ecological, environmental,
economic, social advantages.
It doesn't matter whether you are mining and processing iron ore or limestone, many of
the advantages and disadvantages are common to these operations.
Examples of advantages of a country exploiting it's own mineral resources:
o Valuable revenue if the mineral or its products are exported.
o Jobs for people, especially in poor countries or areas of high unemployment in
developed countries.
o Wages earned go into the local economy.
o Increase in local facilities promoted eg
transport systems, roads and recreational and health social facilities.
o ?
Examples of disadvantages of a country exploiting it's own mineral resources and
reduction of its social and environmental impact:
o Dust from mining or processing can be reduced by air filter and precipitation
systems.
o Noise from process operation or transport of raw materials and products.
Difficult to deal with, sound-proofing often not practical, but operations
can be reduced for unsociable hours eg evening and night!
o Pollution can be reduced by cleaning the 'waste' or 'used' air or water of toxic or
acidic materials eg
carbon monoxide from the blast furnace extraction of iron
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sulphur dioxide gas from copper extraction of its sulphide ore
o Mining operations will disfigure the landscape BUT it can be re-claimed and
'landscaped' in an attempt to restore the original flora and fauna.
o ?
4. How can metals be made more useful?
Extraction details Aluminium can be made more resistant to corrosion by a process called
anodising. Iron can be made more useful by mixing it with other substances to make various types of
steel. Many metals can be given a coating of a different metal to protect them or to improve their
appearance.
Aluminium is a reactive metal but it is resistant to corrosion. This is because aluminium reacts
in air to form a layer of aluminium oxide which then protects the aluminium from further
attack.
o This is why it appears to be less reactive than its position in the reactivity series of
metals would predict.
For some uses of aluminium it is desirable to increase artificially the thickness of the
protective oxide layer in a process is called anodising.
o This involves removing the oxide layer by treating the aluminium sheet with sodium
hydroxide solution.
o The aluminium is then placed in dilute sulphuric acid and is made the positive electrode
(anode) used in the electrolysis of the acid.
o Oxygen forms on the surface of the aluminium and reacts with the aluminium metal to
form a thicker protective oxide layer.
Aluminium can be alloyed to make 'Duralumin' by adding copper (and smaller amounts
of magnesium, silicon and iron), to make a stronger alloy used in aircraft components (low
density = 'lighter'!), greenhouse and window frames (good anti-corrosion properties), overhead
power lines (quite a good conductor and 'light'), but steel strands are included to make the
'line' stronger and poorly electrical conducting ceramic materials are used to insulate the wires
from the pylons and the ground.
The properties of iron can be altered by adding small quantities of other metals or carbon to
make steel. Steels are alloys since they are mixtures of iron with other metals or with non-
metals like carbon or silicon.
Making Steel:
o (1) Molten iron from the blast furnace is mixed with recycled scrap iron
o (2) Then pure oxygen is passed into the mixture and the non-metal impurities such as
silicon or phosphorus are then converted into acidic oxides (oxidation process) ..
eg Si + O
2
==> SiO
2
, or 4P + 5O
2
==> P
4
O
10
o (3) Calcium carbonate (a base) is then added to remove the acidic oxide impurities (in
an acid-base reaction). The salts produced by this reaction form a slag which can be
tapped off separately.
eg CaCO
3
+ SiO
2
==> CaSiO
3
+ CO
2
(calcium silicate slag)
o Reactions (1)-(3) produce pure iron.
o Calculated quantities of carbon and/or other metallic elements such as titanium,
manganese or chromium are then added to make a wide range of steels with particular
properties.
o Because of the high temperatures the mixture is stirred by bubbling in
unreactive argon gas!
o Economics of recycling scrap steel or ion: Most steel consists of >25% recycled
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17
iron/steel and you do have the 'scrap' collection costs and problems with varying steel
composition* BUT you save enormously because there is no mining cost or overseas
transport costs AND less junk lying around! (NOTE: * some companies send their own
scrap to be mixed with the next batch of 'specialised' steel they order, this saves both
companies money!)
Different steels for different uses:
o High % carbon steel is strong but brittle.
o Low carbon steel or mild steel is softer and is easily shaped and pressed eg into a
motor car body.
o Stainless steel alloys contain chromium and nickel and are tougher and more
resistant to corrosion.
o Very strong steels can be made by alloying the iron with titanium or manganese
metal.
Steel can be galvanised by coating in zinc, this is physically done by dipping the object into a
bath of molten zinc. On removal and cooling a thin layer of zinc is left on. The zinc chemically
bonds to the iron via the free electrons of both metals - its all the same atoms to them! It can
also be done by electroplating.
Steel (and most metals) can be electroplated.
o The steel object to be plated is made the negative electrode (cathode) and placed in a
solution containing ions of the plating metal.
o The positive electrode (anode) is made of the pure plating metal (which dissolves and
forms the fresh deposit on the negative electrode).
o Nickel, zinc, copper, silver and gold are examples of plating metals.
o If M = Ni, Cu, Zn ....
At the positive (+) anode, the process is an oxidation, electron loss, as the
metal atoms dissolve to form metal(II) ions.
M
(s)
==> M
2+
(aq)
+ 2e
-
at the negative (-) cathode, the process is a reduction, electron gain by
the attracted metal(II) ions to form neutral metal atoms.
M
2+
(aq)
+ 2e
-
==> M
(s)
For silver plating it is Ag
+
, Ag and a single electron change
Any conducting (usually metal) object can be electroplated with copper
or silver for aesthetic reasons or steel with zinc or chromium as anti-
corrosion protective layer.
Many other metals have countless uses eg zinc
o zinc is used to make the outer casing of zinc-carbon-weak acid batteries.
o It is alloyed with copper to make the useful metal brass (electrical plug pins). Brass
alloy is stronger and more hardwearing than copper AND not as brittle as zinc.
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18
USES OF METALS
METAL USES
PROPERTIES THAT MAKE IT
SUITABLE
Aluminium
a) Structural material for ships,
planes, cars, saucepans.
b) Overhead electricity cables
a) strong but light; oxide layer prevents
corrosion.
b) light but good conductor
Zinc
a) Coating iron to give galvanized
iron
b) To make alloys e.g brass (Zn/Cu)
and bronze ( Zn/Sn/Cu).
a) Reactive- gives acrificial protection
to iron; does not corrode easily.
b) Modifies the properties of other
elements.
Iron
Structural amterial for all industries
( in the form of steel)
Strong and cheap; properties can be
made suitable by alloying.
Lead
a) Car batteries.
b) Solder (Pb/Sn) alloys
a) Design of battery makes recharging
possible.
b) low melting point.
Copper
a) Electric cables
b) Pipes
c) Alloys
d) Coins (Cu/Ni)
a) very good conductors
b)Very ductile, does not corrode easily
c)
d) A traditional metal for coins
Tin
Coating steel cans or tins.
Un reactive and non- toxic. Protevts the
steel from rusting
Nickle
Electroplating steel
Resist corrosion, shiny and attractive to
look at.
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19
Note on Alloy Structure
1. Shows the regular arrangement of
the atoms in a metal crystal and
the white spaces show where the
free electrons are (yellow circles
actually positive metal ions).
2. Shows what happens when the
metal is stressed by a strong force.
The layers of atoms can slide over
each other and the bonding is
maintained as the mobile electrons
keep in contact with atoms, so the metal remains intact BUT a different shape.
3. Shows an alloy mixture. It is NOT a compound but a physical mixing of a metal plus at least one
other material (shown by red circle, it can be another metal eg Ni, a non-metal eg C or a
compound of carbon or manganese, and it can be bigger or smaller than iron atoms). Many alloys
are produced to give a stronger metal. The presence of the other atoms (smaller or bigger)
disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The
result is a stronger harder less malleable metal.
ALLOY
COMPOSITION
%
SPECIAL PROPERTIES USES
Stain less steel
Fe = 74%
Cr = 18 %
Ni = 8 %
Resist corrosion
Car parts, kitchen sinks,
cutlery
Cupronickle
Cu = 75%
Ni = 25%
Hard wearing, attractive
silver color
Silver coins
Manganese
steel
Fe = 85 %
Mn = 13.8 %
C = 1.2
Very hard Springs
Brass
Cu = 70%
Zn = 30 %
Harder then Copper, does
not corrode
Musical instruments, taps
Bronze
Cu = 90 %
Sn = 10 %
Harder then brass, does not
corrode.
Statues, ornaments.
Magnalium
Al = 90 %
Mg = 10 %
Light but strong. Aeroplanes bodies
Solder
Pb = 50 %
Sn = 50 %
Low mwlting point but
form a strong solid
Joining wires and pipes.
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20
METAL CORROSION and the RUSTING of IRON
Iron (or steel) corrodes more quickly than most other transition
metals and readily does so in the presence of both oxygen (in air) and
water to form an iron oxide. You can do simple experiments to show
that BOTH oxygen and water are needed. Put an iron nail into (1) boiled
water in a sealed tube; (2) a tube of air and a drying agent; (3) an open
test tube with water. Rusting appears overnight with (3) only.
Rusting is speeded up in the presence of salt or acid solutions
because of an increased concentration of ions. Corrosion is a redox
process involving redox electron transfer and ion movement. The rusting
metal behaves like a simple cell and more ions enable the current, and hence the electron transfer, to
occur more readily.
Rusting is overall ... Fe
(s)
+ O
2(g)
+ H
2
O
(l)
==> Fe
2
O
3
.xH
2
O
(s)
ie rust is a hydrated iron(III) oxide (the
equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!).
o The reaction proceeds via iron(II) hydroxide Fe(OH)2 which is the oxidised further to the FeO3
o Rusting is an oxidation because it involves iron gaining oxygen (Fe ==> Fe
2
O
3
) or iron atoms
losing electrons (Fe - 3e
-
==> Fe
3+
.
o See more examples of oxidation and reduction below.
The rusting of iron is a major problem in its use as a structural material.
Iron and steel (alloy of iron) are most easily protected by paint which provides a barrier between the
metal and air/water. Moving parts on machines can be protected by a water repellent oil or grease
layer.
This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc or
magnesium). This is referred to as sacrificial protection or sacrificial corrosion, because the more
reactive protecting metal is preferentially oxidised away, leaving the protected metal intact. The picture
illustrates what might be seen after a few days.* Iron or steel can also be protected by mixing in other
metals (e.g. chromium) to make non-rusting alloys called stainless steel. The chromium, like aluminium,
forms a protective oxide layer.
* Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the
more reactive 'sacrificed' metal.
Coating iron or steel with a thin zinc layer is called 'galvanising'. The layer is produced by electrolytic
deposition by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc
(more details). The zinc preferentially corrodes or oxidises to form a zinc oxide layer that doesn't flake off
like iron oxide rust does. Also, if the surface is scratched, the exposed zinc again corrodes before the iron
and continues to protect it.
Steel tin cans are protected by relatively unreacted tin and works well as long as the thin tin layer
is complete. HOWEVER, if a less reactive metal is connected to the iron, it then the iron rusts
preferentially (try scratching a 'tin' can and leave out in the rain and note the corrosion by the
scratch!)
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21
Methods of Prevention of Rusting of Iron
Covering with Paint
Covering with Grease or Oil
Covering with Chromium ( Chrome Plating)
Covering with Tin ( Tin plating)
Covering with Zinc Metal ( Galvanising)
Using Blocks of Zinc Metal
Making Stainless Steel
Using Bocks of Magnesium Metal
\Aluminium does not oxidise (corrode) as quickly as its reactivity would suggest. Once a thin oxide
layer of Al
2
O
3
has formed on the surface, it forms a barrier to oxygen and water and so prevents
further corrosion of the aluminium.
Aluminium is a useful structural metal. It can be made harder, stronger and stiffer by mixing it with small
amounts of other metals (e.g magnesium) to make alloys.
Copper and Lead are both used in roofing situations because neither is very reactive and the compounds
formed do not flake away as easily as rust does from iron. Lead corrodes to a white lead oxide or
carbonate and copper corrodes to form a basic green carbonate (combination of the hydroxide
Cu(OH)
2
and carbonate CuCO
3
eg seen as green roof on buildings).
Both metals have been used for piping but these days lead is considered too toxic and copper is
usually used as the stronger, but equally unreactive alloy with zinc, brass. Now of course, most piping is
flowing in the plastic direction which doesn't corrode at all!
The Group 1 Alkali Metals rapidly corrode in air and need to be stored under oil.
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Apart from their structural weakness they would hardly used for any outside purpose!
DONE
