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Forrester High School: Chemistry Summary notes
Topic 10 / 11 – Metals
1) Introduction
a) Properties of Metals
Most elements in the periodic table are metals. The use of these metals depends on their
properties.
Density
Page 2 in the data book gives values in g/cm3.
The values for most metals are higher than for non-metals.
Density is high because the atoms are packed closely together.
Example:
Dense metals like lead are used in a diver's belts.
Less dense metals, like aluminium, are used to make aircraft.
Thermal Conductivity
Metals all conduct heat well because of the close contact of the atoms.
Example:
Metals, like copper and iron, are used in cooking utensils and radiators.
Electrical Conductivity
Metals all conduct electricity when solid and when molten because electrons can travel
easily through the structure.
Example:
Copper is used for electrical wiring.
Malleability
Metals can be beaten into shape
Example:
The metals used in car bodies are pressed into shape.
Strength
Most metals are strong because of the metallic bond which holds the atoms together.
Example:
The Forth Rail Bridge is made from steel.
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Forrester High School: Chemistry Summary notes
b) Alloys
The properties of metals can be improved by mixing them with other metals or with
non-metals.
Iron can be changed into stainless steel by mixing it with small amounts of chromium.
This stops the metal rusting.
The table below
Alloy
Stainless
steeel
Mild steel
Solder
Brass
Main
Metal
Iron
Iron
Lead
Other elements
present
Chromium,
Nickel
Carbon
Tin
Copper
Zinc
Uses
Reason
Sinks Cutlery
Non-rusting Strong
Girders, Cars
Joining metals
Electrical contacts
Machine bearings,
ornmanets
Strong, rust resistant
Low M.P.
good conductor
Hard wearing,
attractive
c) Recycling
Metals are finite and so will not last for ever.
Recycling metals will help the reserves of them to last longer.
Recycling also saves energy as it takes less energy to recycle a metal than it does to get it
from the ground.
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2) Reactions of Metals
a)
Metals reacting with oxygen
Oxygen can be made by heating potassium permanganate in a test tube and
allowing the gas to pass through rocksil wool to the preheated metal as shown.
Rocksil wool
metal
Potassium permanganate
Heat 2
Heat 1
Different metals react with different intensities. Some like magnesium react very
violently giving out a lot of heat and light while others like copper just give a dull
glow.
The order of metals reacting with oxygen is:
Most
reactive
Least
reactive
magnesium > aluminium > zinc > iron > tin > lead > copper > mercury
Silver and gold do not react.
Potassium, sodium and calcium are too reactive to react with oxygen in this
way.
The equation for a metal reacting with oxygen is:
i.e.
metal
+
oxygen
metal oxide
magnesium
+
oxygen
magnesium oxide
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Forrester High School: Chemistry Summary notes
b) With water
Only a few metals react quickly with water.
Potassium reacts vigorously, sodium very quickly, calcium quickly and
magnesium slowly.
Some other metals will react with steam.
Copper does not react with water
The order of metals reacting with water is:
Most
reactive
Least
reactive
potassium > sodium > calcium >magnesium
The equation for a metal reacting with water is:
metal
i.e.
+ water
potassium + water
2K
+ 2H2O
metal hydroxide
+ hydrogen
potassium hydroxide + hydrogen
2KOH
+
H2
c) Metal reacting with Acid
Potassium, sodium and calcium are too reactive to add to acid.
Copper, mercury, silver and gold do not react,
Most
reactive
Least
reactive
magnesium > aluminium > zinc > iron > tin > Lead
The equation for a metal reacting with acid is:
metal
+ acid
i.e. magnesium + nitric acid
Mg
salt
+ hydrogen
magnesium nitrate + hydrogen
+ 2HNO3
Mg(NO3)2
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+
H2
Forrester High School: Chemistry Summary notes
These reactions give an indication of the reactivity of the metal and are summarised
below. This is called the reactivity series.
Metal
Reaction with
Oxygen
Potassium
Sodium
React
React
Lithium
Calcium
Water
react
Aluminium
with
form
forming
and hydrogen
Magnesium
Zinc
React
forming
metal hydroxide
to
Acid
steam
salt
and
Iron
Tin
metal oxide
hydrogen
Lead
Copper
No
Mercury
No
Silver
No
Gold
Reaction
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Reaction
Reaction
Forrester High School: Chemistry Summary notes
d) Why do metals react?
When metals react they lose their outer electrons forming ions and obtaining the stable
outer electron arrangement of their nearest noble gas.
Loss of electrons is called OXIDATION.
potassium has electron arrangement 2,8,8,1.
potassium loses its 1 outer electron to become stable
K
K+
2,8,8,1
2,8,8
+
e-
oxidation
The name oxidation comes from the reaction of metals with oxygen.
The metal gained oxygen and so the reaction was called oxidation.
For example:
magnesium
+ oxygen
magnesium oxide
oxidation
It is when we look at what is happening to the magnesium that we see that it is losing
electrons.
Mg
Mg2+ O2-
+ O2
2,8,2
2,8
The group 1 metals are the most reactive as they only need to lose 1 outer electron to
become stable.
Remember
Loss of Electrons:
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Oxidation
Forrester High School: Chemistry Summary notes
When the metals react they form ionic compounds. The electrons that they lose are given
to the non metal.
Gaining electrons is called REDUCTION.
For example when magnesium reacts with sulphuric acid the reaction is:
Mg + (H+)2SO42-
Mg2+SO42- + H2
If we remove the spectator ions we get the following equation:
Mg + (H+)2
Mg2+
+ H2
Mg2+
+ H2
This can be written as:
Mg + 2H+
The magnesium is losing its electrons
Mg
Mg2+ +
2,8,2
2,8,
2e-
oxidation
The hydrogen ions are gaining the electrons
2H+
Remember
+
2e-
H2
Gain of Electrons:
Reduction
LEO: Loss Electrons Oxidation
GER: Gain Electrons Reduction
So remember LEO the lion says GER
Reduction reactions are found on page 7 of the data booklet.
For Oxidation reactions turn the reduction equations round.
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reduction
Forrester High School: Chemistry Summary notes
3) Elements on earth
Metals such as gold and silver are found uncombined on earth because they are
unreactive. As a result of this, these elements were among the first to be discovered.
Other metals are found in compounds called ORES and have to be extracted.
a) Metal Ores
Ores are naturally-occuring compounds of metals from which metals can be extracted.
The three main types of ore are metal carbonates, metal oxide and metal sulphides.
Common name
Haematite
Bauxite
Galena
Malachite
Chemical name
iron oxide
aluminium oxide
lead sulphide
copper (II) carbonate
Metal present
lron
aluminium
lead
copper
b) Extraction of metals from ores
When the metals react to form compound the more reactive metals form more stable
compounds.
This means that it is more difficult to obtain metals from ores of reactive metals than
from ores of unreactive metals.
The more reactive the metal the harder it is to break up the compound (ore).
As we are trying to turn metal ions back into metals these reactions are called Reduction
because they are losing oxygen, and in the process gaining electrons.
For example:
Iron ore
iron
Fe2O3
2Fe
(Fe3+)2(O2-)3
2Fe
Fe3+ + 3e-
Fe
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Reduction
Forrester High School: Chemistry Summary notes
Methods of extraction
a) Heating metal oxides
The least reactive metals can be obtained from their ores simply by heating.
The main reaction is:
i.e.
Metal oxide
metal
+
oxygen
mercury oxide
mercury
+
oxygen
2HgO
2Hg
+
O2
This method is used to extract metals below mercury in the reactivity series.
b) Heating metal oxides with carbon
More reactive metals are extracted using carbon or carbon monoxide to remove the
oxygen.
The main reaction is:
i.e.
Metal oxide
+
carbon
metal +
carbon dioxide
lead oxide
+
carbon
lead
+
carbon dioxide
PbO2
+
C
Pb
+
CO2
This method is used to extract metals below aluminium in the reactivity series.
c) Electrolysis
The most reactive metals form the most stable compounds and so need the most
energy to extract the metal.
Electricity is needed to split ionic compounds into their elements in a process
called electrolysis.
A large electric current is passed through the molten compound, and metal
appears at the negative electrode.
i.e.
aluminium oxide
2Al2O3
aluminium
+
oxygen
4Al
+
3O2
(Al3+)2(O2-)3
4Al
Al3+ + 3e-
Al
Reduction
This method is used to extract reactive metals above zinc in the reactivity series.
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Forrester High School: Chemistry Summary notes
d) Heating with carbon monoxide
Iron is extracted from its ore in the blast furnace by heating with carbon (coke) in
the presence of air.
In zones 1 and 2 carbon is turned into carbon dioxide and then carbon monoxide.
In zone 3 the carbon monoxide turns the iron ore into iron.
Zone 1: At the bottom of the furnace the reaction makes carbon dioxide
C
+
O2
CO2
Zone 2: The carbon dioxide reacts with carbon to make carbon monoxide.
CO2
+
C
2CO
Zone 3: The carbon monoxide reacts with iron oxide to make iron and carbon dioxide.
Fe2O3 +
CO
2Fe
+
2CO2
The formation of a metal from a compound is a REDUCTION reaction.
This is because the metal ion is being given electrons to turn it into the metal.
(Fe3+)2(O2-)3 +CO
2Fe
Fe3+ + 3e-
Fe
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+
2CO2
Reduction
Forrester High School: Chemistry Summary notes
4) Displacement
When metals react they lose electrons forming metal ions: Oxidation
The more reactive metals readily lose their electrons forming stable compounds.
The less reactive metals form less stable compounds in which the metal ions can be easily
turned back into the metal atoms.
These statements can be seen happening in displacement reactions:
Displacement:
A metal can displace ions of a less reactive metal.
For example:
Adding zinc metal to a solution of copper (II) sulphate
produces copper metal and a solution of zinc (II) sulphate.
Zn(s) + Cu2+SO42-
Cu(s) + Zn2+SO42-
Omitting the sulphate ions (SO42-) and the state symbols makes the process clearer.
Zn
+ Cu2+
Cu
+ Zn2+
Being the more reactive metal the zinc metal is losing its electrons forming zinc ions.
Copper is a less reactive metal so the electrons are pushed onto the copper ions turning
them into copper metal.
Zn
Zn2+ +
Cu2+ + 2e-
Cu
2 e-
Oxidation
Reduction
Adding the oxidation and reduction equations together gives us the REDOX equation
Zn
+ Cu2+
Cu
+ Zn2+
REDOX
5) Hydrogen
Metals reacting with acids is an example of a displacement reaction.
Zn(s) + (H+)2SO42-
Zn2+SO42- + H2
The zinc metal is displacing the hydrogen ions.
The metals which do not react with acids, can not displace the hydrogen ions and so
these metals must be less reactive than hydrogen.
Therefore hydrogen is put into the reactivity series between lead and copper.
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Forrester High School: Chemistry Summary notes
Topic 10 Making Electricity
1) Introduction
Electricity is a flow of charged particles:
flow of electrons through metals,
flow of ions through solutions or melts.
When placed in an electrolyte metals tend to lose electrons and form positively charged
ions. The ions go into the solution and the electrons are left on the metal giving the
metal a negative charge. However metals react at different speeds resulting in different
metals having a different build up of negative charge.
Metal A
-ve
-ve
-ve
+ ve
-ve
+ve
-ve
+ ve
+ve +ve
Metal B
-ve
-ve
-ve
+ve
-ve
+ve
-ve
+ve
+ve +ve
-ve
-ve
+ ve
+ ve
-ve
-ve
-ve
+ve
+ve
+ve
+ve
Metal A is more reactive than metal B and so it has a greater potential to form ions and
leave electrons on the surface of the metal.
Connecting the metals, electrons will flow along the wire from metal A to metal B.
Flow of electrons
V
Ion Bridge
-ve
-ve
-ve
+ ve
-ve
+ve -ve
+ ve
+ve +ve
-ve
-ve
-ve
-ve +ve
+ve
-ve
+ve
+ve +ve
-ve
-ve
+ ve
+ ve
-ve
-ve
-ve
+ve
+ve
+ve
+ve
With an ion bridge in place to complete the circuit, a voltmeter can measure the
difference in potential of metals to form ions. This is known as the voltage. The
metals can be placed in an order which compares the energy involved in metal atoms
forming ions. This is called the electrochemical series. It is found on page 7 of the data
booklet and is very similar to the reactivity series.
Electrons always flow from metals high in the electrochemical series through the wires to
metals lower in the electrochemical series.
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2) Making electricity using two different metals
The further apart two metals are in the electrochemical series the larger the voltage
obtained.
For example when magnesium is connected to copper a larger voltage is
produced than when zinc is connected to copper.
2.7V
1.1V
Mg
Cu
Zn
Cu
electrolyte
Cells (Batteries)
A cell contains chemicals which react to make electricity.
The energy change is:
Chemical Energy
Dry Cells
Electrical Energy
In dry cells, the chemicals are used up and the cells then have to be
replaced.
In a simple dry cell, the chemicals are shown in the diagram below.
metal top (+)
carbon rod
ammonium chloride
paste
zinc case (-)
The paste containing ammonium chloride is the electrolyte needed to complete the
circuit .
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Forrester High School: Chemistry Summary notes
3) Rechargeable cells.
In rechargeable cells (like car and mobile phone batteries) the chemicals are not used up
and can be regenerated by recharging the cell.
Energy Changes in Cells
Using the cell (Discharging): Reaction A
Chemical Energy
Electrical Energy
Charging the cell: Reaction B
Electrical Energy
Chemical Energy
This can be seen in a lead acid Cell (car battery).
If the battery is used when the car is not running the chemicals get used up quickly
(reaction A) and the battery stops working. If the battery is recharged by connecting it to
the mains electricity through a battery charger the chemicals are remade (reaction B) and
the battery can be used again.
Battery or Mains?
There are advantages and disadvantages with all sources of electricity.


Batteries are expensive to buy but make the appliance portable.
Battery-operated appliances are safer but use finite resources such as metals in
the battery.


Rechargeable cells use less of the finite resources, but still use fossil fuels to make
the electricity they need.
Rechargeable cells contain toxic chemicals and should be disposed of carefully.


Mains appliances are cheaper to run, but can be dangerous.
Mains appliances are not as portable, and still use up fossil fuels.
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Forrester High School: Chemistry Summary notes
5) Oxidation and Reduction in cells.
Electron flow
A
Zn
Zn2+SO42-
Ion Bridge
Cu
Cu2+SO42-
The ammeter shows flow of electrons through the wire from the more reactive zinc to
the less reactive copper.
The copper ions in the solution accept these electrons and turn into copper metal.
Zn
Cu2+ + 2e-
Zn2+
Cu
+
2e-
Oxidation
Reduction
Note:
 These two equations are on page 7 of the data booklet. The equation higher in
the list has been reversed.
 The ion bridge completes the circuit by allowing ions to move through it.
 It is important that the ions in the ion bridge do not react with any of the
electrolyte to form a solid as they will then not be able to move and the cell will
stop.
 When the chemicals are used up (i.e. the Zn metal or the Cu2+ ions) the cell will
stop producing electricity.
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Forrester High School: Chemistry Summary notes
6) Electricity from cells involving non-metals
The electrochemical series in the data book has some reactions that involve non-metals
e.g.
I2 + 2e-
2I-
A cell can be set up by using a carbon rod dipping into a solution of iodine dissolved in
potassium iodide solution (which contains iodide ions) as one half cell and connecting
this to another half cell as shown below:
Electron flow
A
Zn
Ion Bridge
Zn2+SO42-
C
K+I-
The diagram shows that electrons flow from zinc metal to the carbon rod and the
reactions are shown below
I2
+
Zn
2e-
Zn2+
2I-
+
2e-
Reduction
Oxidation
2I-
+
Zn2+
REDOX
The redox equation will be:
I2
+
Zn
7) Identifying Redox equations
Wherever there is an oxidation there must be a reduction and vice versa. In other words
if one chemical is losing electrons another chemical must be gaining electrons.
Whenever metal atoms react they lose electrons.
In a cell where a non metal molecule is turning into its ions, it is gaining electrons.
Therefore we must have a REDOX reaction when an equation contains either:
 A metal (i.e. Zn)
 A non-metal molecule and its ions (i.e. Br2 / Br-)
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Forrester High School: Chemistry Summary notes
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