1. How are metals extracted from their ores?
How are metals prevented from reverting to their oxides?
Reactivity series
Electrolysis
background
How to extract metals
IRON
Preventing iron reverting
BRINGS YOU
BACK HERE
ALUMINIUM
Preventing aluminium reverting
Copper purification
END

2. Some metals are more reactive than others.
A list of metals in order of their reactivity is called the REACTIVITY SERIES.
Reactions with air / oxygen
HIGH
REACTIVITY
LOW
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Tin (Sn)
Lead (Pb)
Copper (Cu)
Silver (Ag)
Gold (Au)
Platinum (Pt)
All these metals burn vigorously with a bright flame.
copper + oxygen 
copper
oxide
The metal burns with a bright, white flame, leaving a white ash – magnesium oxide. 2
Cu + O2  CuO 
balanced 2
NOT balanced
The burning of magnesium ….. 2
Mg + O2  MgO
The metal slowly gets a black covering of copper oxide when heated. 
balanced
An unreactive element, found in the earth’s crust as the metal itself. Little reaction with oxygen.

3. IT ALL FITS IN WITH THE REACTIVITY SERIES
SUMMARY OF REACTIONS WITH WATER
The less reactive metals do nothing.
Magnesium reacts only when heated with steam.
Calcium reacts gently when cold
Sodium & potassium react increasingly violently when cold.
IT ALL FITS IN WITH THE REACTIVITY SERIES
Reactions of metals with water
HIGH REACTIVITY
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Copper (Cu)
Gold (Au)
Calcium slowly dissolves in cold water and fizzes.
The solution is an ALKALI
The gas is hydrogen.
Sodium and potassium run around on the surface of cold water, produce heat & dissolve, with lots of fizzing.
The solution is a strong ALKALI
The gas is hydrogen.
The hot metal burns with a bright, white flame when steam is passed over it.
It produces magnesium oxide Hydrogen can be detected..
No reactions.
magnesium (s) + water (g)  hydrogen(g)
magnesium
oxide (s)
calcium(s) + water(l)  hydrogen(g)
calcium
hydroxide(aq)
sodium (s) + water (l)  sodium hydroxide(aq) + hydrogen(g)
This state symbol, (g) for “gas” tells us it’s STEAM 
balanced 2
Na + H2O  NaOH H2
Mg + H2O  MgO + H2 2 2 
balanced
Note: metals that react with STEAM produce the OXIDE
metals that react with WATER produce the HYDROXIDE
liquid water
alkalis contain the
hydroxide ion, OH- 
balanced
Ca + H2O  Ca(OH)2 + H2 2
NOT balanced
NOT balanced

4. Put some zinc metal into blue copper sulphate solution.
A MORE REACTIVE METAL CAN DISPLACE A LESS REACTIVE ONE FROM ITS COMPOUNDS
MORE REACTIVE
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Copper (Cu)
Gold (Au)
Put some zinc metal into
blue copper sulphate solution.
The zinc displaces the copper and forces the copper to become an element.
Because zinc is more reactive, it has a greater tendency to become a compound than copper has.
NOTE : Lead, for example would not displace iron from iron chloride solution because lead is LESS reactive.
Pb(s) + FeCl3 (aq) x
LESS REACTIVE
Zinc(s) + copper sulphate(aq)  zinc sulphate(aq) + copper(s)
In the test tube we see the zinc dissolving, the copper sulphate solution getting less blue and copper metal appearing at the bottom.
copper displaced from its compound
zinc is more reactive & becomes a compound 
balanced
Zn +
CuSO 
ZnSO4 + Cu
REMEMBER
The MORE reactive a metal is, the more it likes being in a COMPOUND
The LESS reactive a metal is, the more it likes being an ELEMENT
The zinc has displaced the copper from its compound.

5. x  2 2 Al + Fe2O3  Al2O3 + Fe Cu(s) + ZnO (s)  Sodium (Na)
A MORE REACTIVE METAL CAN DISPLACE A LESS REACTIVE ONE FROM ITS OXIDE
MORE REACTIVE
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Copper (Cu)
Gold (Au)
Heat aluminium powder with powdered iron(III) oxide.
The mixture flares up.
The aluminium is displacing the iron from its oxide.
The aluminium is taking the oxygen from the iron, leaving iron metal.
LESS REACTIVE
NOTE : There would be no reaction between copper and zinc oxide because copper is LESS reactive.
Cu(s) + ZnO (s)  x
Aluminium(s) + iron oxide(s)  iron(s)
Aluminium
oxide(s) 
balanced 2
Al +
Fe2O 
Al2O 2 Fe
NOT balanced

6. REACTIONS OF METALS WITH ACIDS
MORE REACTIVE
REACTIONS OF METALS WITH ACIDS
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Hydrogen (H)
Copper (Cu)
Gold (Au)
React with cold water : violent with acids
Less reactive than hydrogen. Cannot displace it.
These metals have no reaction with acid.
A more vigorous reaction.
Hot solution & lots of bubbles.
A reaction of moderate speed.
Steady production of bubbles.
Acids are compounds containing hydrogen as the H+ ion.
We must include hydrogen in the reactivity series.
LESS REACTIVE
Metal + acid  salt + hydrogen
Metals that are more reactive than hydrogen, displace it from the acid and set the hydrogen free as the gaseous element.
Zinc + sulphuric acid 
zinc sulphate + hydrogen
(s)
(aq)
(g)

7.    metal(s) + acid(aq) salt(aq) +hydrogen(g) 2 2 Examples iron(II)
sulphate
iron + sulphuric acid 
+ hydrogen 
balanced
H2SO4 
Fe +
FeSO4 + H2
magnesium
chloride
hydrochloric
acid
magnesium 
+ hydrogen 
balanced
Mg + 2
HCl 
MgCl2 + H2 
balanced
lead
nitrate
NOT balanced
lead + nitric acid 
+ hydrogen H2
Pb + 2
HNO3 
Pb(NO3)
NOT balanced

8. EXTRACTING METALS FROM THEIR ORES.
The earth’s crust contains metals and metal compounds, always found mixed with other substances.
Gold is an unreactive metal, at the bottom of the reactivity series.
It is found in the earth as the metal itself and chemical separation is not needed.
In ORES, the metal or its compound is concentrated enough to make it economic to extract the metal.
Often an ore contains a metal oxide or something that can easily be changed into a metal oxide.
The ore haematite is iron(III) oxide, Fe2O3
The ore bauxite is aluminium oxide, Al2O3
To extract the metal, the oxygen must be removed from the metal oxide.
ThIs process is called REDUCTION.

9. HOW A METAL IS EXTRACTED FROM ITS ORE DEPENDS ON HOW REACTIVE IT IS.
MORE REACTIVE
Hydrogen will displace less reactive metals from oxides of those metals.
Because its more reactive, it is able to remove the oxygen and combine with it to make water.
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Hydrogen (H)
Copper (Cu)
Silver (Ag)
Gold (Au)
So hydrogen can reduce ;
copper oxide
silver oxide
Pass a stream of hydrogen gas over heated, black copper oxide.
It will gradually go brown as copper metal is made.
Hydrogen cannot reduce ;
LESS REACTIVE
lead oxide
iron oxide
zinc oxide, etc
copper oxide(s) + hydrogen(g) = copper(s) + water(g)
H2 
CuO +
Cu +
H2O

10. Electricity generation is a modern development.
HOW CARBON CAN BE USED TO EXTRACT A METAL FROM ITS ORE. (this is an important, central point)
Using first wood and then coal, this has been known about for thousands of years. (since the Iron Age)
The difficulty of extracting aluminium from its oxide meant a stronger method, electrolysis, was needed & this is expensive.
Electricity generation is a modern development.
So we have only been able to extract the more reactive metals in recent times.
MORE REACTIVE
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Carbon (C)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
Copper (Cu)
Gold (Au)
Aluminium is MORE reactive than carbon and so carbon does NOT have the ability to remove oxygen from aluminium oxide.
This is the position of the non-metal, carbon in the reactivity series.
or zinc
A metal such as iron, which is less reactive than carbon CAN be extracted from its ore using carbon.
Because it is more reactive, carbon can take the oxygen from the iron oxide.
As a result, iron is obtained from its ore using carbon in the Blast Furnace.
Aluminium has to be extracted from Bauxite using ELECTROLYSIS.
These two methods are considered in the following slides.
LESS REACTIVE

11. EXTRACTION OF IRON IN THE BLAST FURNACE
The raw materials are :
coke (C)
haematite (Fe2O3)
limestone (CaCO3)
EXTRACTION OF IRON IN THE BLAST FURNACE
1. The coke burns in an exothermic reaction. Energy is released, raw material gets hot and carbon dioxide is formed.
waste
gases
2. At the high temperatures, the carbon dioxide reacts with more coke to make carbon monoxide.
3. The carbon monoxide REDUCES the iron oxide in the ore, into molten iron which flows to the bottom of the furnace.
carbon + oxygen  carbon dioxide
C + O2  CO2
carbon
dioxide
monoxide
+ carbon 
4. Limestone is added to remove acidic impurities forming a molten slag which floats on top of the iron.
Hot air is blown in
CO2 + C  CO 2 
balanced
NOT balanced
SLAG
MOLTEN
IRON

12. The process of removing oxygen from the ore is called REDUCTION.
Hot air
MOLTEN
IRON
SLAG
The carbon monoxide REDUCES the iron oxide in the ore, into molten iron which flows to the bottom of the furnace.
This is the central, important reaction in the Blast Furnace.
The process of removing oxygen from the ore is called REDUCTION.
The carbon monoxide combines with the oxygen from the iron ore to produce carbon dioxide.
This is called OXIDATION.
Iron(III) oxide  iron +
carbon
monoxide
dioxide
This is OXIDATION
This is REDUCTION
Fe2O CO  Fe CO2 3 
balanced 2 3
NOT balanced

13. HOW CAN IRON BE PREVENTED FROM REVERTING
TO ITS OXIDE ?
Iron & steel corrode with air and water more quickly than most transition metals. It can be prevented by :
connecting iron with a more reactive metal “Galvanised iron” is the name given to iron plated with zinc.
The more reactive metal corrodes first, (zinc) making sure
the iron does NOT corrode.
That is why the zinc would be called a “sacrificial metal”.
magnesium
wire
underground
iron pipe
Here, the more reactive magnesium is being sacrificed,
saving the iron pipe.
Steel can be mixed with other metals such as chromium (Cr) to make an alloy which will not rust.
Lots of cutlery made in Sheffield is manufactured from this “stainless steel”.
Theory of
Electrolysis
ALUMINIUM

14. - MAKING ALUMINIUM METAL INDUSTRIALLY USING ELECTROLYSIS. +
This large block of carbon is connected to the positive of the voltage supply.
The ions are attracted as shown
Oxygen forms at the positive electrode
MOLTEN ALUMINIUM OXIDE IN CRYOLITE
O2-
Al3+ -
The lining of this large tank is made of carbon and is connected to the negative of the voltage supply.
The ions must be able to move but melting the aluminium oxide is difficult because of its high melting point.
So the aluminium oxide is dissolved in molten aluminium compound called CRYOLITE at a much lower temperature.
The raw material from which aluminium is produced is aluminium oxide, Al2O3 which contains the ions Al3+ and O2-
The ore is called BAUXITE which needs purifying to get the aluminium oxide.
The aluminium forms at the negative electrode
The oxide ion loses electrons and is oxidised : 2O2- - 4e  O2
This oxygen reacts with the carbon (+) electrode
and makes carbon dioxide gas.
So the electrode burns away quickly and often has to be replaced.
The aluminium ion gains electrons and is reduced
Al3+ + 3e  Al

15. HOW CAN ALUMINIUM BE PREVENTED FROM REVERTING TO ITS OXIDE ?
Aluminium is almost as reactive as magnesium & we might expect it to corrode quickly over time, or to burn up (oxidise) completely if it got hot.
But it does not do this.
Gets covered in a thin layer of aluminium oxide as soon as it is made.
Aluminium metal
This surface layer forms a barrier to oxygen and water & so prevents further corrosion.
Aluminium is a useful structural metal, low in density.
It can be made harder, stronger & stiffer by mixing it will small amounts of other metals such as magnesium, to make an alloy.

16. Cu - e  Cu2+ Cu2+ + e  Cu 2 PURIFICATION OF COPPER - - - - - 2 - + +
The negative electrode is pure copper. +
is the copper ion, Cu2+
The solution must contain the Cu2+ion
Copper sulphate(aq) will do.
The positive electrode is scrap copper
containing impurities. - +
The scrap copper dissolves in the solution.
Pure copper gets plated onto the negative electrode. - - + - + - + - + + - + - -
impurities
separated
out + - +
OXIDATION AT THE POSITIVE ELECTRODE 2
Cu - e  Cu2+
REDUCTION AT THE NEGATIVE ELECTRODE
Cu2+ + e  Cu 2

17. THE END
These are the remains of some of the earliest Blast Furnaces. (Coalbrookdale Museum of Iron)
This is a painting of the
“Bedlam” Blast Furnaces at night,
hundreds of years ago.
To find out more about the Steel Industry, visit :
The Ironbridge Gorge Museums, Telford
Kelham Island Museum, Sheffield

18. Electrolysis

19. What sort of substances conduct electricity?
ALL METALS
This is because they have some FREE ELECTRONS in them that can move throughout the metal
These electrons can carry the current because:
They have an electric charge
They can move
GRAPHITE (a form of carbon)
Graphite is the ONLY non-metal element that conducts electricity
Like metals, it conducts because it has FREE ELECTRONS
IONIC SUBSTANCES when MOLTEN or in SOLUTION
When ionic substances are molten (= in liquid form) or in solution, their positively and negatively charged ions can move – so they conduct.
When they are solid, the ions can’t move – so no current flows,

20. When we pass a current through a metal or graphite, the substance is not changed.
When we pass a current through an ionic substance that’s molten or in solution, CHEMICAL CHANGES happen.
This is ELECTROLYSIS

21. What we need for electrolysis+ - e-
ELECTRODES to carry the current into the solution or molten compound
These are made of graphite or a metal
Electrical wires to make the circuit
A cell or battery e- e- e- e-
The ELECTROLYTE – that’s the solution or molten ionic compound the electrodes are dipped in.
The cell pushes electrons around the circuit
FROM its NEGATIVE terminal
TO its POSITIVE terminal
This means that
the positive electrode ends up short of electrons
the negative electrode ends up with extra electrons
YES – that IS the right way round – even though it’s opposite to the way we think of the current going…

22. - - + + Opposite charges attract
negative ion, Cl- +
positive ion, Cu2+
Opposite charges attract
The negative ions move to the positive electrode
The positive ions to the negative electrode
The ionic compound is decomposed by the passage of electric current.
copper chloride(aq)
Elements are released at the electrodes as gases or metals.
In this example, you would get copper metal made at the negative electrode and chlorine gas appearing at the positive electrode.

23. Here are Cu2+ ions moving to the negative electrode.
Positive ions in the solution are attracted to negative electrode – opposite charges attract.
When electrons are gained by a positive ion, the name of the chemical change is REDUCTION.
REDUCTION IS THE GAIN OF ELECTRONS.
THE COPPER ION HAS BEEN REDUCED
2e- + Cu2+  Cu
The electrode is negative because it has too many electrons
As they get close, the ions gain electrons from the electrode and the Cu2+ is neutralised.
e- go to ion Cu
This makes copper the element, which covers the electrode.
e- go to ion.
TWO ELECTRONS FROM THE CATHODE
A NEUTRAL ATOM OF THE ELEMENT COPPER.
ARE ADDED TO THE COPPER ION

24. When electrons are lost by a negative ion, the name of the chemical change is OXIDATION.
OXIDATION IS THE LOSS OF ELECTRONS
THE CHLORIDE ION HAS BEEN OXIDISED
e- go to cell
Cl-
This is what happens at the positive electrode when chloride ions, Cl- are present in the electrolyte
Here are negative chloride ions
attracted towards the positive
electrode. Opposite charges attract.
This electrode is positive because some electrons have been removed by the cell.
As they get close, each Cl- ion loses an electron which goes onto the electrode.
The ion becomes electrically neutral
the ion
loses an e-
Cl2
We have made chlorine the element. The neutral atoms join in pairs to make chlorine molecules, Cl2 which bubble off as a gas.
2Cl- - 2e  Cl2
TWO CHLORIDE IONS, EACH WITH AN EXTRA ELECTRON
THE TWO ELECTRONS LOST BY THE IONS GO TO THE ELECTRODE
A NEUTRAL CHLORINE MOLECULE
END