TWEED RIVER HIGH SCHOOL
2006
PRELIMINARY CHEMISTRY
Unit 2
Metals
Part 2
Metals differ in their reactivity with other chemicals and this
influences their uses.

Describe observable changes when metals react with
dilute acid, water and oxygen.
Note: You must know the common mineral acids:
Acid
Sulfuric
Hydrochloric
Nitric
Formula
H2SO4
HCl
HNO3
1. Reactions of metals with dilute acids
Some metals react rapidly with dilute acids, others react slowly and
some have no reaction.
Metals react with dilute acids according to the general reaction
equation:
metal + acid  metal salt + hydrogen
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For example:
Magnesium reacts with sulfuric acid:
Word equation
Magnesium + sulfuric acid  magnesium sulfate + hydrogen
Chemical equation
Mg(s) + H2SO4(aq)  MgSO4(aq) + H2(g)
Observable changes:

Change in temperature

Gas evolved

Metal disappears
Homework: Copy Table 8.3,page 126, Chemistry Contexts 1
Table 7.3 p 131 new edition
2
Reactions of Metals with Water
Most metals do not react with cold water. However, Group 1 metals
do. These metals react according to the following general equation:
Metal + Water  Salt (hydroxide) + hydrogen
For example:
Word equation
Sodium + water  sodium hydroxide + hydrogen
Chemical equation:
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Observable Changes:

Gas evolved

Increase in temperature
Homework: Copy Table 8.2,page 124, Chemistry Contexts 1
Table 7.2 p130 new edition
3
Reaction of Metals with Oxygen
Metals react with oxygen at varying rates. Lithium, potassium and
sodium
react
very
quickly,
rubidium
and
caesium
react
so
vigorously that they catch fire in air. Other metals such as
aluminium and iron react slowly in oxygen. Also metals such as
silver and gold do not react with oxygen.
Metals react with oxygen according to the general reaction
equation:
Metal + oxygen  metal oxide
For example:
The most common example would be the corrosion of iron (rust).
iron + oxygen  iron(III) oxide
4Fe(s) + 3O2(g)  2Fe2O3(s)
magnesium + oxygen  magnesium oxide
2Mg(s) + O2(g)  2MgO(s)
Observable changes:

Release of heat

Fire

Colour change
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Homework: Copy Table 8.1,page 122, Chemistry Contexts 1
Table 7.1 p127 new edition

Describe and justify the criteria used to place metals into
an order of activity based on their ease of reaction with
oxygen , water and dilute acids.

Identify the reaction of metals with acids as requiring the
transfer of electrons.
Note: The above points are covered in the notes below.
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Oxidation – Reduction reactions
Oxidation: involves the loss of electrons.
Reduction: involves the gain of electrons.
Hence oxidation-reduction reactions involve one reactant losing
electrons and the other reactant gaining electrons.
For example in the reaction:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
In this reaction the magnesium atom loses 2 electrons to become a
magnesium ion, while 2 hydrogen ions gain an electron each to
become hydrogen gas.
This electron transfer can be represented as 2 half equations:
Mg(s)  Mg2+(aq) + 2e-
Oxidation
And
2H+(aq) + 2e-  H2(g)
Reduction
Adding the half equations together gives the overall equation.
Metal ions can also displace electrons from other metals. For
example:
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
This is known as a metal displacement reaction.
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Metal displacement reactions along with the reactions with water,
oxygen and dilute acids allow metals to be place into an activity
series:
K Ba Ca Na Mg Al Zn Fe Ni Sn Pb Cu Ag Hg Pt Au
Strongest
Reducing agent
Strongest
Oxidising
agent
Note:
An oxidising agent (or oxidant) causes another substance to be
oxidised. Therefore, the oxidant is itself reduced.
Conversely:
A reducing agent (or reductant) causes another substance to be
reduced. Therefore, the reductant is itself oxidised.
Homework:
Questions 2, 3, page132, Chemistry Contexts 1
Questions 9 & 10 p140 new edition

Perform
a
first-hand
investigation
incorporating
information from secondary sources to determine the
metal activity series.
Obtain the practical sheets – Determining a Metal Activity Series
and using these sheets write up your experiments including the
following points:

Aim

Risk Assessment

Method

Results

Discussion
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
Have your write up approved before you do the
practical.


You will then perform the experiment in the lab.
Identify the importance of first ionisation energy in
determining the reactivity of metals.

Outline the relationship between the relative activities of
metals and their position on the Periodic Table.
Ionisation Energy
Remember that protons are positively charged and found in the
nucleus of an atom with the negative electrons in shell around the
nucleus. Therefore, there is an electrostatic attraction between the
protons and electrons. To remove an electron this electrostatic
attraction must be overcome.
The ionisation energy of an element is the energy required to
remove an electron from a gaseous atom of the element.
The first ionisation energy, is the energy required to remove the
first electron. The second ionisation energy, is the energy required
to remove the second electron and so on.
Ionisation energies decrease down a group and increase across a
period. How does this affect the reactivity of metals?
A comparison of the activity series for metals reveals that the most
reactive metals are generally found on the left side of the periodic
table with the least reactive metals tending to be in the middle of
the periodic table.
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
Outline examples of the selection of metals for different
purposes based on their reactivity, with a particular
emphasis on current developments in the use of metals.
Gold

Gold is the least reactive of all metals.

It readily keeps its lustre which makes it excellent for
jewellery.

Because it is not reactive, and has excellent electrical
conductivity,
gold
is
used
in
electrical
connections
in
computers and electronic circuits.

Gold has also been used in the space area as it has excellent
reflective properties.
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Magnesium

Magnesium is a highly reactive metal which makes it suitable
for some very specific applications.

Magnesium is attached to the steel hulls of ships. Since
magnesium is more reactive than the iron in the steel, hence
it will corrode before the steel, protecting the ship from
corrosion. In this instance the magnesium is called a sacrificial
anode.
Zinc

Zinc is used in the production of galvanised iron, where steel
is dipped into molten zinc.
The zinc serves two purposes:
-
It reacts with air to form an impervious layer
-
It also reacts as a sacrificial anode
Zinc is also used in dry cell batteries. Here, zinc is oxidised
and electrons are released.
Tin

Tin is used to coat base metal objects. Such as the tin cans
food is packaged in.
Chromium

Chromium is used to coat metal objects for protection and
appearance. For example, chrome plating on motor vehicles
and other objects.
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Aluminium

Aluminium is suitable for use in areas like the construction
industry. Not only because of its strength and lightness of
weight, but also aluminium forms and impervious oxide layer
which it prevents further corrosion.
Copper

Copper is unreactive and hence corrosion resistant.

It is used extensively in the plumbing and electrical industries.
Homework:
Construct balanced chemical equations and half-equations showing
electron transfer for the following reactions:
1. sodium + hydrochloric acid
2. potassium + sulfuric acid
3. calcium + hydrochloric acid
4. magnesium + sulfuric acid
5. aluminium + sulfuric acid
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