Year 10 Science
Minerals to Metals
- Chemistry and the Mining Industry
A king he was on carven throne
In many-pillared halls of stone
With golden roof and silver floor,
And runes of power upon the door.
The light of sun and star and moon
In shining lamps of crystal hewn
Undimmed by cloud or shade of night
There shone forever fair and bright.
There hammer on the anvil smote,
There chisel clove, and graver wrote;
There forged was blade, and bound was hilt;
The delver mined, the mason built.
There beryl, pearl, and opal pale,
And metal wrought like fishes' mail,
Buckler and corselet, axe and sword,
And shining spears were laid in hoard.
- JRR Tolkien, The Fellowship of the Ring
Common Assessment Tasks
Due Date
Formal Practical Reports (4)
Homework Tasks
1
20%
Continuous 20%
Oresome Froth Flotation
10%
Minerals to Metals Tests (2)
50%
Name:
________________________________
Form:
________________________________
Teacher:
________________________________
Second edition
2
Table Of Contents
Chapter 1
1-1
1-2
1-3
1-5
Minerals to Metals Standards ………………………..……………..………………………..…………………
Notes to the Student , Task descriptions …….………………………….....…..…………………………..
Course timeline and task description ..……….………..……………………………………………………..
Reference Websites, Setting out your workbook ………..……………………………………………….
Page 3
Page 4
Page 5
Page 6
Chapter 2
2-1
2-2
2-3
2-4
2-5
2-6
2-7
2-8
2-9
2-10
2-11
2-12
The Importance of Metals………………….…………………..……………..………………………….……………
The Properties of Metals ........………………..…….....…………………………………………………………..
Metallic Bonding …….…………………………………………….………………………………………….………
The Periodic Table
……………………………………………………………………………………………..…
The Modern Periodic Table
………………………………………………………………………………….…
Review Activity: Electronic Configuration and the Periodic Table ……………………………….
Particle Puzzle ……….………….………………………………….……………………………………………….…
Ionic Bonding …………………….……………………………………………………………………..…………….….
Class Exercise: Ionic Bonding …………………….…………………………………………..…….………….….
Writing Ionic Formulas …………………….…………………………………………………………..…………….
Chemistry in the Balance …………………….…………………………………………………………..……………
Balancing: Bringing it all Together …………………….…………………………………………………………
Page 7
Page 8
Page 9
Page 11
Page 14
Page 15
Page 14
Page 15
Page 18
Page 21
Page 20
Page 21
Chapter 3
3-1
3-2
3-3
3-4
3-5
3-6
3-8
3-9
3-10
3-11
3-12
3-13
3-14
3-15
3-16
3-17
3-18
3-19
3-20
3-21
3-22
Alkali Metals …………………….…………………………………………………………………………………………
Patterns within Group 1 …………………….……………………..…………………………………………………
The Transition Metals …….………………….……………………..…………………………………………………
Research: Metals and their Properties …….………………….………………………………………………
Alloys – Modifying the Properties of Metals …….………….……………………………………………….
Practical: Reaction of Metals with Water …….………….……………………………………………………
Practical: Displacement of Metals …….………….…………………………………………….……….……….
Homework Exercise .…….………….…………………………………………………….………….………………...
Extracting Metals from their Ores – Redox Chemistry …….………….…………………………………
Video: Redox Rocks! …….………….……………………………………………………………………………….…
Demonstration Reaction: Mighty Thermite …….………….…………………………….………………….
Extraction Methods .…….………….…………………………………………………….…………………………...
Chemical Reactions involving Metals .…….………….………………………..……………………………....
Mineral Calculations – Composition of Common Ores .…….………….………………….……..……...
Oresome Froth - Student Worksheet .…………………………………………………………………..……...
Extraction of Copper from Ore .………………………………………………….………..….….………..……...
Teacher Demonstration: Electroplating of Solutions .……..……………..…………………..…..….....
Practical 7: Extraction of Copper from its Ore .……..……………..…………………..….…………..…...
Production of Iron – The Blast Furnace ….……..……………..…………………..……….……………..…..
Demonstration: Corrosion of Iron ….……..……………..……….………………...………………………......
Summary Minerals to Metals Questions ….……..……………..……….………………...…..………….......
Page 22
Page 23
Page 24
Page 25
Page 31
Page 33
Page 35
Page 39
Page 40
Page 42
Page 45
Page 47
Page 49
Page 51
Page 52
Page 54
Page 55
Page 56
Page 60
Page 63
Page 70
Chapter 4
4-1
Chemical Data .…….………….…………………………………………………….……………………….…………..
Page 71
3
1-1 Science Standards: Minerals To Metals
Properties of Metals
Different types of chemical reactions are used to produce a range of products



4




Describe the general properties of metals
Relate the chemical and physical properties of metals to their usefulness to society
Recall that metals can be arranged in a reactivity series based on the reactions of the
metals and their compounds: potassium, sodium, lithium, calcium, magnesium,
aluminium, zinc, iron, copper, silver and gold
Describe how reactions with water and dilute acids can be used to deduce the following
order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron, and
copper
Use the position of a metal within the reactivity series to predict displacement reactions
between metals and their oxides, and between metals and their salts in aqueous
solutions
Establish the chemical formula of an ionic compound
Balance chemical equations
Extraction Methods






Describe methods (pyrometallurgy / electrolysis) used to extract different kinds of
metals
Relate the difficulty in refining a metal to its position on chemical reactivity series
Identify the role of ionic bonding involving metals and their compounds in terms of
electron transfer
Identify oxidation and reduction in chemical reactions
Identify the components of an electrolytic cell and the reactions occurring at the
electrodes
Describe environmental/economic issues associated with mining and uses of a resource
Corrosion of Iron




Describe the chemical processes involved in the Blast Furnace
Recall the conditions under which iron rusts
Describe how the rusting of iron may be prevented by grease, oil, paint and galvanising
Understand the sacrificial protection of iron in terms of the reactivity series
1-2 To The Student
Electronics
Minerals to Metals
Traffic
Water
Genetics/Evolution
Immunology
This booklet is one of a series for Year 10 Science. If you miss a class for any reason, you are encouraged to ask
a classmate or your teacher of the set homework. Common assessment tasks are the major tasks that all
students in Year 10 must complete to satisfactorily complete the course. There are scanned pages from
Chemistry for Higher Tier on the intranet for you to read and complete your homework questions – you can
access these from home. The intranet will also have PowerPoint and other essential files for you to use to help
you prepare for the unit test.
1-3 Task Descriptions
You will work through the activities in this booklet by completing different types of tasks. They are described
in the table below and are easily identifiable in the workbook by their symbols. Along the way there will be
definitions of key terms – you are to learn them. Creating a glossary in your exercise book for these terms and
adding new terms as they arise is a recommended method. These, and only these, terms will be used in tests
and the examination.

Workbook exercises: Exercises to be completed in your exercise books

Common Assessment Tasks: These are the major tasks that all students must complete and include unit
tests and assignments. Task descriptions and marking schemes can be found on the curriculum portal.

Calculate

Write the answer in this booklet

Explain/Discuss/Debate:
5
1-4 Minerals to Metals Timeline
Week 4
Week 3
Week 2
Week 1
Concepts and Skills
6
 Establish the chemical formula
of an ionic compound
 Balance chemical equations
 Describe the general properties
of metals
• Relate the chemical and
physical properties of metals to
their usefulness to society
 Use the structure and metallic
bonding to explain the
conductivity, strength and
malleability of metals
• Describe reactions of metals
with water and dilute acids
 Describe how reactions with
water and dilute acids can be
used to deduce the order of
metal reactivity
 Recall that metals can be
arranged in a reactivity series
based on the reactions of the
metals and their compounds
 Use the position of a metal
within the reactivity series to
predict displacement reactions
between metals and their
oxides, and between metals and
their salts in aqueous solutions
 Describe methods used to
extract different kinds of
geological resources
(pyrometallurgy / electrolysis)
 Relate the difficulty in refining a
metal to its position on chemical
reactivity series
 Identify the role of ionic
bonding involving metals and
their compounds in terms of
electron transfer
 Identify oxidation and reduction
in chemical reactions
 Identify the components of an
electrolytic cell and reactions
occurring at the electrodes
 Describe environmental/
economic issues associated with
mining and uses of a resource
Class Activities
Homework
Intro PowerPoint: Importance
and Properties of metals – Notes
2-1
Review of Chemistry concepts
2-7 Ionic Bonding
2-8 Writing Ionic Formulas
2-9 Chemistry in the Balance
2-10 Balancing: Bringing it all
Together
Demo Prac 1: Alloys
Prac 2: Reactions of Metals with
Water
Prac 3: Reactions of Metals with
Acids
H/W Activity Ex. 2.2 – 2.6
Results, equations, prac
questions, conclusions
Prac 1,2 Formal report –
hardcopy
Prac 3 Formal report
Chem for Higher Tier (CHT)
PDF on Portal
3-1 Alkali Metals
3-2 Patterns within Group 1
3-3 The Transition Metals
3-4 Metals and Properties
Test: Formulas and Balancing
Equations (15%)
Prac 4: Displacement of Metals
3-9 Homework Exercise to
Summarise Practicals
Mineral Calculations (RAM)
3-10 Extracting Metals from
their Ores
PowerPoint: Extraction Methods
IT Simulation: Oresome Froth
Flotation
Video: Redox Rocks!
3-12 Teacher Demo: Mighty
Thermite
Teacher Demo: Pyrometallurgy
3-13 Extraction Methods
3-14 Chemical Reactions
involving Metals
3-15 Mineral Calculations
3-18 Demo: Electroplating of
Solutions
3-20 Electroplating of Solutions
Prac 5: Extraction of Copper from
its Ore
Complete Oresome Froth
worksheet and include a
screenshot of most profitable
setup – hardcopy submission.
3-16 Chemical Reactions
involving Metals
Extracting Metals from Ores:
p. 76 – 77, Q 1 – 5
Principles of Electrolysis: p.
82-83 Q1-4
Week 5
 Describe the chemical processes
involved in the Blast Furnace
 Recall the conditions under
which iron rusts
 Describe how the rusting of iron
may be prevented by grease, oil,
paint and galvanising
 Understand the sacrificial
protection of iron in terms of
the reactivity series
3-20 Production of Iron – The
Blast Furnace
Extracting Iron, p. 81 Q1- 6
Corrosion: p. 88-89 Q 1-6
Demo Prac 6: Corrosion of Iron
PowerPoint: Corrosion of Iron
Revision for Unit Test
3-24 Summary Minerals to
Metals Questions
Test 2
1-5 References
Chemistry for Higher Tier: R. Gallagher, Paul Ingram, Oxford University Press
IGCSE Chemistry: Bryan Earl and Doug Wilford, Hodder Education
Interactive Periodic Table: http://www.ptable.com/
1-6 Setting Out Your Exercise Book
Your workbook is one of the main ways for you to communicate with your teachers. Follow these steps each
and every time you use your exercise book.
1. Rule margins in red and ensure they are wide enough for teacher correction
2. Write the date and title of the lesson – this often helps you focus. Use red colour.
3. Diagrams are drawn in pencil, and are large enough to follow easily.
4. Rule a line after each class and homework AND use your workbook sequentially – write on both sides of
each page
1-7 The Science Toolbox
A scientist uses a number of tools to answer key
scientific questions. Many help organise and illustrate
data, including:
Diagrams. Use pencil, simple 2D format, labelled
Experimentation: Fair tests
Flow charts – with title, labelled axes and scales
Graphs
Maps
Photographs
Reports
Surveys
Tables
You should have the correct equipment
with you to complete these tasks.
This workbook
Your notebook
Compass
Protractor
Calculator
Ruler
Pens, pencil and eraser
7
 2-1 The Importance of Metals
To accompany Powerpoint: Minerals to Metals
Australia is one of the world’s biggest producers of minerals. In the 1850’s the
country experienced a rapid rise in immigration as a result of a mining boom
– the Gold Rush. Gold brought an influx of wealth and by that time, Australia
was producing 40% of the world’s gold. Today, minerals continue to
provide much of the wealth of this nation, and governments have
continued to support the expansion of the mining industry.
Minerals and the Australian Mining Industry
Minerals are naturally occurring inorganic solids with a crystalline
structure and a fixed chemical composition and known formula. Ores are
concentrations of minerals in rock that are high enough to be economically
extracted for use. All ores are minerals, but all minerals are not necessarily ores. Figure 1: Holey dollar coin
Principal Mineral Formulas – Complete this table showing the chemical
from early Sydney settlement
composition of common minerals.
Copper
Chalcopyrite
CuFeS2
Calcium Limestone
Mercury
Titanium
Iron
Haematite
Fe2O3
Aluminium
Iron
Magnetite
Fe3O4
Zinc
PbS
Tin
Lead
Cinnabar
HgS
TiO2
Bauxite
Al2O3
ZnS
Cassiterite
Australia is a major primary producer of many metals, such as those listed below. Write two uses of each metal,
and the world ranking of Australian production.
Metal
Aluminium
Silver
Gold
Iron
Copper
Lead
Titanium
Zinc
Tin
8
Use
World
Ranking
Which metal is appropriately used for each application: Titanium / lead / copper / gold / iron / zinc /
aluminium / silver ?
Metal
Use
Metal
Teapot
Artificial hip joint
Bridge
Roofing material coating
Food wrap foil
Car battery electrode
Electrical wiring
Yellowish jewellery
Use
Other uses of metals
There are other uses of metals that are less obvious. Metals are used as catalysts to speed up reactions. Nickel
is used as a
in the manufacture of margarine. Platinum is used in car exhausts to reduce pollution.
Metals are used to colour materials such as in s
g
and c
. Copper is
also used as an a
in swimming pools.
2-2 The Properties of Metals
Metals: How do metals differ from non-metals?
There are key properties that distinguish metals from other substances.
What are the chemical and physical properties of most metals?
1._____________________________________________________________________________
2._____________________________________________________________________________
3._____________________________________________________________________________
4._____________________________________________________________________________
5._____________________________________________________________________________
6._____________________________________________________________________________
7._____________________________________________________________________________
8._____________________________________________________________________________
9._____________________________________________________________________________
What is the structure of metals?
Earlier, we looked at the properties of metals. Why do metals have these properties?
The difference and similarities in the properties of metals are due to their atomic composition.
Many metals are quite dense, which means that their atoms are packed closely together. When the atoms of
metals get very close to each other, the outer electron shells merge together. Outer electrons are then able to
move not only around one atom but also between all the atoms. The outer electrons of the metal atoms
separate from the atoms and create a ‘sea of electrons’. These electrons are d
and so are free to
move through the whole structure.
The metal atoms become positively charged ions and are attracted to the sea of electrons. This attraction is
called m
.
9
2-3 Metallic Bonding
How do metals conduct heat and electricity?
Delocalised electrons in metallic bonding allow metals to conduct heat and electricity. For example, when a
metal is heated, the delocalised electrons gain k
. These electrons then move faster and
transfer the gained energy throughout the metal. This makes h
in metals very efficient.
Delocalised electrons also conduct e
through metals in a similar way.
Why do metals have high melting points?
The properties of metals are related to their structure. Metals often have high m
and b
. Gold, for example, has a melting point of 1064°C and a boiling point of 2807°C. This
property is due to the strong attraction between the p
charged metal ions and the sea of
electrons. In metal extraction and other industrial processes, furnaces often run continuously to maintain the
high temperatures needed to work with molten metals.
Why are metals strong?
Metals are usually s
, not b
. When a metal is hit, the layers of metal ions are able to slide
over each other, and so the structure does not shatter. The metallic bonds do not break because the delocalised
electrons are free to move throughout the structure. This also explains why metals are malleable (easy to
shape) and ductile (can be drawn into wires).
The positive ions result from metal atoms ‘losing’ their outer shell electrons and these electrons form a sea of
delocalised negative charge. It is the electrostatic force of attraction between the positive metal ions and the
delocalised electrons that holds the atoms together. This attraction is called metallic bonding.
It is important to note that the positive ions (cations) stay in fixed positions in
the metal lattice (because they are in a solid). Only the outer electrons are free
to move. It is also worth noting that these cations are composed of a nucleus,
containing protons and neutrons, and inner shell electrons.
Draw a diagram of a cross-section of a sodium cation, Na+, as it would be in a
metallic lattice. (Sodium atoms have 11 protons, 12 neutrons and 11 electrons.)
Statement
Metals are excellent conductors of heat
Metals are malleable but not ductile
Metallic bonding is weak
Bonding is due to attraction between positive ions
Metals have free electrons that conduct electricity
Metals are excellent insulators of heat
Metals are described as ions in an electron sea
When hit, metal ions can slip past one another
10
True or False?
Minerals to metals – Forging Links in Chemistry and Mining
Metals have an important role in any modern society. They are used commercially in the manufacture of ships,
planes, cars and industrial machines. To strengthen concrete, iron is a major material used in the construction
of buildings. Metal wires of pure copper and aluminium are used for the supply of electricity. Copper is used for
water pipes. Aluminium is used in alloys for manufacturing of different parts of ships and planes and different
heavy machinery.
Extraction of metals and minerals is the fifth-largest industry in the world. In Australia, the mining industry
employs 130 000 people. Mining companies employ chemists, geologists and engineers to explore and extract
metals.
The invention of extraction technologies made metals available to humans. The invention of these technologies
empowered civilisations and changed the history of the world.
Metal
Gold
Copper
Iron
Aluminium
Ore
Chalcopyrite
Haematite
Bauxite
Extraction Method
None
Heat in air
Smelt in furnace
Technology Invention
< 5000 BCE
BCE
1500 BCE
1860
Elements: a reminder
Everything in the world is made from elements. They combine with each other to give millions of compounds.
Your body is built up from over 40 different elements - including enough iron to make a big nail!
Most of the elements are metals. There are over 80 of these. Chances are you've seen and touched a good many
of them already. Think of the iron in gates and railings, aluminium cooking foil, the tin coating on'tin' cans,
silver and gold in jewellery, and magnesium in the lab.
A quick reminder of the Periodic Table:
• The elements are arranged in order of increasing atomic number.
• The zig-zag line separates the metals from the non-metals in the table, with the metals on the left.
• The numbered columns are called groups. The number tells you how many outer electrons the atoms of
that group have.
• The elements in a group behave in a similar way.
• The numbered rows are periods. This time the number tells you how many electron shells there are.
Look at the block of transition metals. These metals have a lot in common with each other, but are quite
different from the Group 1 and 2 metals. For a start, they are much less reactive.
Some of the metals in the table are very common, and we use them a lot. For example aluminium makes up 8%
of the Earth's crust, and iron 6%, and we use them everywhere. But some are very rare, and you will probably
never see them, or even say their names!
11
 2-4 The Periodic Table – Homework Review
Look closely the periodic table on the back cover. In Year 9, you encountered the idea that the position of an
element on the periodic table is related to its properties. You also identified the link between the properties
and the valence (outer shell) electrons.
1. a. Write the electron configurations for the following Group II elements
Be (4)
Mg (12)
Ca (20)
b. Write a general statement about the relationship between the group number and the number of outer shell
electrons of the atoms of the elements in that group.
Because the chemistry of an element depends on the outer shell electrons that its atoms have, atoms in the
same group often have similar chemical properties, although this is only noticeable for elements that are near
each other in the group. The further apart they are, the less likely they are to have similar chemical properties
(look at the marked differences between the elements at the top and the bottom of group IV).
2. A row (from left to right) is called a period. The period from potassium to krypton is called Period 4.
a. Write the electron configurations for the following elements from Period 2 (atomic numbers in brackets):
Li (3) ______________________
Be (4) ______________________
B (5)
______________________
b. Write a general statement about the relationship between the period number and the number of shells with
electrons in the atoms of the elements in that period.
Once again because the chemistry of an element depends on the outer shell electrons that its atoms have, as the
number of outer shell electrons increases from left to right across a period, so there often are gradual trends in
the chemical properties of the elements as you move across a period.
3. Look at the relative masses of the elements from potassium across to krypton in Period 4.
a. How does the atomic number change across Period 4?
b. How does the order of the elements in the periodic table relate to atomic number?
4.
12
Where are the metals found in the periodic table? Where are the non-metals?
 2-5 The Modern Periodic Table – Homework Review
A. The metals aluminium, calcium, magnesium, potassium and sodium are obtained by electrolysis, which
involves passing electricity through their molten oxides or chlorides. Use a coloured pencil to identify these on
the table.
Copper, iron, lead, mercury, silver, tin and zinc are obtained by heating the metal ores in a furnace either alone
or with coke. Use a different colour to identify these.
B. On the outline of the periodic table below, use different colours or some other code to show elements on
your periodic table with similar properties. Specifically, label the following blocks: Reactive metals, Hard (less
reactive) metals, Reactive non-metals, Non-reactive (inert) non-metals.
13
2-6 Electronic Configuration and the Periodic Table – Homework Review
Group 1
1
H
1
Group 2
Group 3
Group 4
Group 5
Group 6
Group 7
Electronic configuration means the arrangement of electrons in an atom. The electronic configurations
Group 8
4
He
2
for the first 20 elements are shown here, this time in rows and columns, to match the periodic table:
1. Draw the electronic structure for each element (this is shown for neon)
2. Write out the electronic structure (this is shown for neon – 2,8)
7
Li
3
9
Be
4
11
B
5
12
C
6
14
N
7
16
O
8
19
F
9
20
Ne
10
2,8
23
Na
11
24
Mg
12
39
K
19
40
Ca
20
14
27
Al
13
28
Si
14
31
P
15
32
S
16
35
Cl
17
40
Ar
18
 2-7 Particle Puzzle – Homework Review
All these words have something to do with atoms and molecules. Unscramble them and fit them into the
puzzle.
15
1.
Particle with a negative charge
RENTLOCE
2.
Made up of atoms which all have the same atomic number
MELTEEN
3.
Needed for motion
GREENY
4.
Charged atom or group of atoms
NOI
5.
Same number of protons - different number of neutrons
POSITOE
6.
Anything which takes up space
TETRAM
7.
Smallest bit of a substance which can normally exist
UMOCELLE
8.
Number of protons in a hydrogen atom
NEO
9.
Charge on an electron
GENAVITE
10. Nuclear particle with no electric charge
TUNNORE
11. Central part of the atom
SCLUUNE
12. Very small bit of matter
RAPTLICE
13. Charge on a proton
PISOVITE
14. A particle in the nucleus
TORNOP
 2-8 Ionic Bonding
Reference: Chemistry for Higher Tier, pp. 26 -29, 32-33.
In Year 9 Chemistry, you were introduced to the Bohr atomic model, which described the arrangement of
electrons in shells surrounding a positively charged nucleus. It is this arrangement of electrons that determines
the properties of all elements and explains the way chemicals react. Therefore, an atom’s electronic structure is
of high interest to chemists. You will need your periodic table to complete this activity.
1. Label each of the following atomic models, including atomic number (Z), the number of protons. In a neutral
atom, the number of protons equals the number of electrons. Sodium is given as an example.
Atom
Atomic number
Electronic arrangement
Atom
Sodium
Atomic number
Electronic arrangement
11
2,8,1
Noble Gases
The Noble gases, in increasing atomic number, are:
These are distinguishable from all other elements, in that they are chemically unreactive and generally do not
form compounds. They are stable because
Only the
16
have complete outer shells. For all other atoms, their outer shells are
. That is why they react. Atoms can obtain full outer (valence) shells by
.
All these other elements will react in such a way as to obtain a similar electronic arrangement as the Noble
Gases. Metals react differently to non-metals, and we will consider examples of sodium and chlorine.
Sodium
The sodium atom has
arrangement by
. A sodium atom has
protons, 11 neutrons, and
become a
ion, Na+.
in its valence shell. It will react to have a complete outer shell
. It will then become a
protons, 11 neutrons, and
electrons. The sodium ion has
electrons. Therefore, it has lost
electron
and
Cl
Chlorine
A chlorine atom has
outer shell electrons. To produce a full
valence shell, it reacts by gaining
from
another
atom, and becomes
.
A chloride atom has
protons, 18 neutrons, and
electrons. The chloride ion has
protons, 18 neutrons,
and
electrons. Therefore, it has gained
electron
and become a
ion, with symbol
.
Ions
As a general rule, metals lose electrons, while non-metals
electrons, and ions are formed. Ions are charged when they have a
different number of protons and electrons.
The Ionic Bond
If metals can lose electrons to fill their outer shell and non-metals are available to accept them, a chemical
reaction can occur. (Note: the proton number does not change in these chemical reactions, but electrons are
transferred)
Li+
F-

+
Lithium – Li
(element)
Ionic Bond
Fluorine – F
(element)
Proton
Neutron
Electron
Lithium Fluoride - LiF
(compound)
Above is an illustration of lithium reacting with fluorine. Lithium metal
one
electron to the non-metal, fluorine. Lithium ions and fluoride ions are formed, with symbols:
and
. These positive and negative ions are electrostatically attracted to form an ionic bond.
2. Draw a similar diagram for the reaction with sodium (11) and chlorine (17) to form sodium chloride.
17
Ionic Compounds
When sodium reacts with chlorine, billions of ions form, and are
They cluster together, so that each ion is surrounded by
ions of
electrostatically held together by strong
bonds.
to each other.
charge. They are
A giant structure of ions is formed, which contains
of ions. This structure is a
compound of
, or common salt and its formula is
. While the ions
have positive and negative charges, the total charge on the giant molecule has
.
Other Ionic Compounds
Magnesium atoms have
valence electrons, and oxygen has
. When magnesium burns in
the presence of air (oxygen), its outer shell electrons are transferred to oxygen. Magnesium then becomes a
ion, with charge
, and oxygen becomes a negative oxide ion, with charge
.
3. Draw the electron shell structure below:
4. Calcium and fluorine can also be involved in metal / non-metal reactions. Calcium has
valence
electrons, and fluorine has
. (Hint: the reaction is similar to that of magnesium and chlorine)
These two new charges that have been produced are
electrostatically attracted and form an ionic compound, calcium
fluoride, with formula
.
Calcium Fluoride Crystal
18
 2-9 Class Exercise: Ionic Bonding
(Teacher note: The Atoms Kit is available to help demonstrate this concept)
1. View the following shell structures, and identify each element.
+

The Structure of Ionic Compounds
In the compound, sodium chloride, the ions are arranged in a regular arrangement, or
. They are
held together by strong
. Since the ions are arranged in a regular pattern
all ionic substances are
.
The Properties of Ionic Compounds
1. Ionic Compounds have
Magnesium oxide has a much higher melting point than sodium chloride because
2. Ionic Compounds are usually
3. Ionic Compounds can conduct
19
 2-10 Revision Exercise: Writing Ionic Formulas
(1 period)
1. Use the table of common ions to figure out the formula for the compounds listed. Sodium chloride has been
done for you as an example.
Name of substance
Formula
Ions present
NaCl
Na+, Cl -
Sodium chloride
Ratio of ions
Working out space
1:1
Lithium chloride
Lithium fluoride
Iron sulfide
Magnesium chloride
Magnesium oxide
Calcium oxide
Calcium fluoride
Potassium chloride
Sodium oxide
Aluminium oxide
2. Write the formula for the compounds formed when the cations and anions bond. Sodium chloride and
sodium sulfate have been done for you as an example.
Negative Ions (Anions)
Br
Na+
Positive Ions (Cations)
NH4+
Ca2+
Pb2+
Zn2+
Mg2+
Fe3+
Al3+
20
NaCl
NO3
Cl
O
SO42
CO32
PO43
 2-11 Chemistry in the Balance
(1 period)
What do students find the hardest thing to believe about chemistry?
Balance the following equations by writing in all the coefficients, including ‘1’ where appropriate. For each
equation, find its sequence of coefficients in the key list at the bottom of the page and write the corresponding
letter next to the number of the equation. Read down the letters to find the answer. (A subscript after a bracket
multiplies all the atoms inside the bracket.) Hint: compound ions such as NO3- , PO43- and SO42- are easier to
identify as a group – treat them as one object on both sides of the equation. Further explanation of
how to balance equations can be found at Khan Academy.
1.
Ba(NO3)2
+
K2SO4

KNO3
+
BaSO4
2.
CH4
+
O2

H2O
+
CO2
3.
CuO
+
HNO3

Cu(NO3)2
+
H2O
4.
Fe
+
Al2O3

Fe2O3
+
Al
5.
HCl
+
CaO

CaCl2
+
H2O
6.
Al2O3
+
C

CO
+
Al
7.
FeCl3
+
AgNO3

Fe(NO3)3
+
AgCl
8.
NaOH
+
H3PO4

Na3PO4
+
H2O
9.
Fe2O3
+
C

Fe
+
CO
10.
Pb(NO3)2
+
Na2SO4

NaNO3
+
PbSO4
11.
NaOH
+
HCl

NaCl
+
H2O
12.
Na
+
H2O

H2
+
NaOH
13.
H2SO4
+
Cu(OH)2

CuSO4
+
H2O
14.
BaO
+
HNO3

Ba(NO3)2
+
H2O
15.
Pb(NO3)2
+
KCl

PbCl2
+
KNO3
16.
Al(OH)3
+
HNO3

H2O
+
Al(NO3)3
17.
H2O
+
K

H2
+
KOH
Key list
A.
1, 1, 1, 1
F.
1, 2, 1, 2
U.
1, 3, 3, 1
W. 2, 1, 2, 1
B.
1, 1, 1, 2
H.
1, 2, 2, 1
S.
1, 3, 3, 2
N.
2, 2, 1, 2
C.
1, 1, 2, 1
T.
1, 3, 1, 3
I.
2, 1, 1, 1
G.
3, 1, 1, 2
E.
1, 2, 1, 1
Y.
1, 3, 2, 3
M. 2, 1, 1, 2
R.
3, 1, 1, 3
21
 2-12 Balancing: Bringing it all Together
Balancing chemical equations is a form of chemistry accounting, where all atoms are identified before and after
a reaction takes place. The principle is that, to follow the law of conservation of mass, balance the atoms before
and after the reaction occurs.
Hints:
1. Write the chemical formula first, using the methods already seen earlier. If these are correct, do not
change them.
2. Balance the numbers of atoms before and after the reaction takes place, by placing numbers in front of
each compound. As noted in step 1, it is essential do not change the formulas.
3. Compound ions such as NO3-, PO43- and SO42- are easier to identify as a group – treat them as one object
on both sides of the equation.
4. If you have even number of atoms on one side of the equation and odd on the other, double everything
and continue balancing.
Write the following in symbol form and balance the following equations.
1.
Copper (II) oxide + nitric acid ® copper (II) nitrate + water
2.
Calcium hydroxide + carbon dioxide
3.
Calcium carbonate + water + carbon dioxide
4.
Silver carbonate + hydrochloric acid
5.
Potassium + water
6.
Calcium + water
7.
Silver + copper nitrate
8.
Aluminium carbonate + hydrochloric acid  aluminium chloride + carbon dioxide + water.
22




calcium carbonate + water

calcium hydrogen carbonate
silver chloride + water + carbon dioxide
potassium hydroxide + hydrogen gas
calcium hydroxide + hydrogen gas

copper + silver nitrate
 3-1 Alkali Metals
CHT p. 60 – 1
The alkali metals are the six metals of
in the periodic table: lithium, sodium, potassium,
rubidium, caesium, and francium. Only the first three are safe enough to keep in the school lab, but all have to
be treated with respect and care. Rubidium and caesium are
. If they are added to
water in a glass trough the violent reaction will shatter the glass. Francium is even more dangerous – and
. Hardly any exists.
Physical properties
The alkali metals are not typical metals.




Like all metals, they are good conductors of heat and electricity.
They are softer than most other metals. Sodium and potassium can be
.
They are ’lighter’ than most other metals - they have
. Lithium, sodium and potassium
float on water (and immediately react with it).
They have low
points compared with most metals.
Chemical properties
The alkali metals are the most reactive of all metals. Here are some of their main reactions.
Reaction with water
Alkali metals react violently with water, producing
and a metal hydroxide.
The reaction for sodium is:
sodium + water  sodium hydroxide + hydrogen
The hydrogen bubbles off. The hydroxide is alkaline (
) so the indicator changes colour.
All alkali metals react vigorously with water, releasing hydrogen gas and forming hydroxides. The hydroxides
give alkaline solutions.
Reaction with non-metals
With chlorine – Put heated alkali metals into chlorine gas - and they will burst into flame. They burn
brightly, giving white solids called chlorides. The reaction with sodium is:
sodium + chlorine sodium chloride
Sodium chloride is also known as common
colourless solution.
. It dissolves in water to give a
With oxygen – Alkali metals burst into flame when you heat them. If you place them in oxygen they
burn even more fiercely, forming white solids called
. This time the flames have different
colours. Lithium burns with a red flame, sodium burns yellow, and potassium’s flame is lilac.
Ionic compounds
The compounds formed in these reactions with metals and non-metals are ionic. Sodium chloride is made of
sodium ions (Na+) and chloride ions (Cl-). Sodium oxide is made of sodium ions and oxide ions (O2-). So, during
the reaction, sodium atoms have become sodium ions with a charge of 1+.
All alkali metals form ionic compounds in which the metal ion has a charge of 1+. The compounds are white
solids that dissolve in water to form
.
Questions – Answer in complete sentences in notebook
1. Name each of the Group 1 elements and give their symbols.
2. What is the family name for the Group 1 elements? Why are they called that?
3. Which best describes the Group 1 metals:
a. soft or hard? b. high or low density? c. high or low melting point? d. reactive/unreactive with water?
23
 3-2 Patterns within Group 1
CHT p. 62 – 3
Trends in physical properties
The Group 1 metals are a family. Like all families, each member is a little different. Look at this table:
Metal
Lithium
Sodium
Potassium
Rubidium
Caesium
Symbol
This metal is silvery and ...
Density (g/mL)
Melts at ... (oC)
Boils at... (oC)
Li
Na
K
Rb
Cs
Soft
A little softer
Softer still
Even softer
Softest (of 5)
0.53
0.97
0.86
1.53
1.88
181
98
63
39
29
1342
883
760
686
669
As you can see, each property shows an overall increase or decrease as you go down the table. This kind of
pattern is called a trend. You can summarise the trends in the table like this:
lithium
sodium
potassium
rubidium
caesium
Softness increases
Density increases
Melting point increases
Boiling point increases
Why they have similar chemical properties
You know that the members of the Group 1 family react in a similar way with
. Why is this? The answer is simple: they all have the same number of outer shell electrons. Atoms
with the same number of outer shell electrons react in a
.
The trend in reactivity
Compare the reactions of Group 1 family with water:
The increase in violence shows the metals are getting
more
. Reactivity increases as you
go down Group 1.
Metal
Lithium
Sodium
Potassium
Rubidium
Caesium
What you see
A lot of fizz around the floating metal
It shoots around on the surface of water
It melts and the hydrogen bursts into flames
Sparks fly everywhere
A violent explosion
Explaining the trend in reactivity
As you go down Group 1 the atoms get larger, because they add electron shells. They react in order to obtain a
full outer shell. They do this by
an electron. They become an ion with a charge of 1+. The
larger the atom, the further the outer shell is from the positively charged nucleus and it is less attracted. So the
easier it is to lose an electron. So the more reactive the metal becomes.
Comparing Groups 1 and 2
The elements of Group 2 are quite similar to the alkali metals. For example, magnesium and calcium react with
water too – but much more
. Hardness
Melting point
,
and reactivity
.
When Group 2 metals react, they have to lose
to obtain a full outer shell. This is more
difficult than losing just one. And that is why they are less reactive than the Group 1 metals.
As you go down Group 2 the atoms add electron shells, as usual. So the outer shell gets further from the
nucleus. That makes it easier to lose electrons from it. So reactivity increases down the group.
Questions – Answer in complete sentences in notebook
1. The Group I metals show a trend in melting points.
a. What does that mean?
b. Describe two other physical trends for the group.
2. Find one measurement that does not fit the trend exactly, in the table on the opposite page.
3. Group 1 metals all have similar chemical properties. Why is this?
4. When a Group 1 metal reacts, what happens to the outer shell electron of its atoms?
5. a. Which is more strongly held, the outer electron in lithium or in sodium? Explain why you think so.
b. Sodium is more reactive than lithium. Why?
24
 3-3 The Transition Metals
CHT p. 64 – 5
The transition metals are the middle block of 30 elements in the periodic table. They include most of the metals
you find in everyday use in the kitchen, in school, and around town.
Physical properties
Here are three of the transition metals:
Iron – widely used. Grey with a
metallic lustre
Copper – reddish with a metallic
lustre
Nickel – silvery grey with a
metallic lustre
The transition metals have these physical properties:
•
,
and
. You cannot cut them with a knife, as you can the Group 1 metals.
• high melting points. (Mercury is an exception. It is a liquid at room temperature. M.Pt -39°C.)
• malleable and ductile.
• good conductors of heat and electricity. Silver is the best electrical conductor of them all, and copper is
next.
•
density. Unlike sodium, they sink in water, since their
density is greater than the density of water (1 g/mL).
Their chemical properties
• They are much
reactive than the Group 1 metals. For
example: Copper does not react with water, nor does it burn in air.
• Their low reactivity means they do not
very readily in
air or water. (Iron is an exception. It rusts very easily, and we spend a
fortune every year on rust prevention.
• Most of them form
compounds. (Most compounds
of Group 1 and 2 metals are white.) Colours used in stained glass
windows usually come from transition metal compounds.
The uses of Transition metals
The properties of the transition metals suit them to a wide range of uses.
• Their hardness and strength make them suitable for building structures.
Iron is by far the most used for
. (It is abundant and is not expensive.) It is
usually used in the form of steels.
• They are used for things that need
or
current to pass through easily.
Steel is used for hot water radiators. Copper is used for electric cables.
• Many transition metals and their compounds acts as
– they speed reactions up without
being chemically changed themselves. Iron is a catalyst in the Haber process to make ammonia from
nitrogen and hydrogen. Nickel is a catalyst for the hydrogenation of vegetable oils to
.
Questions – Answer in complete sentences in notebook
1. Name five transition metals (not those listed above).
2. Which best describes the transition metals:
a. soft or hard? b. high or low density? c. high or low melting point? d. reactive/unreactive with water?
3. What is unusual about mercury?
4. Most paints contain compounds of transition elements. Suggest a reason why.
5. Suggest reasons why copper is used in hot water pipes while iron is not.
25
Questions
Answer the following questions in complete sentences.
1. You are given a sample of an element. Write down three clues that would tell you it was a metal.
2. Metals tend to be solid at room temperature. Name one that is not.
3. Think of two ways in which gold behaves:
a. like iron
b. differently from iron
4. The table below shows four groups of ‘substances’. This table was published by the French chemist, Antoine
Lavoisier in 1789. Lavoisier thought these ‘substances’ were elements. The modern names of some of them
are given in brackets.
Acid-making elements
Sulphur
Phosphorus
Charcoal (carbon)
Gas-like elements
Light
Caloric (heat)
Oxygen
Azote (nitrogen)
Hydrogen
Metallic elements
Cobalt
Mercury
Copper
Nickel
Gold
Zinc
Iron
Silver
Lead
Tin
Manganese
Tungsten
Platina (platinum)
Earthy elements
Lime (calcium oxide)
Magnesia (magnesium oxide)
Barytes (barium sulphate)
Argilla (aluminium oxide)
Sitex (silicon dioxide)
a. Write a definition of:
i. an element
ii. a compound
b. Name a substance in the table that is a compound – not an element.
c. Why do you think Lavoisier thought this substance was an element?
d. Name a ‘substance’ in the table that is neither an element nor a compound.
5. The back cover of this booklet has a periodic table. Look at the row of elements from lithium to neon.
a. What is this row of the periodic table called?
b. Which element in this row is the least reactive? Explain why.
Look at the column of elements from beryllium (Be) to calcium (Ca).
c. What is this column of the periodic table called?
d. Which is the most reactive element shown in this column? Explain your answer in terms of electronic
structure.
e. Describe how the atomic structures of the first 20 elements relate to their positions in the periodic table.
Draw diagrams to illustrate some examples.
6. Read the following passage about the physical properties of metals.
Elements are divided into metals and non-metals. All metals are electrical conductors. Many of them have a high
density and they are usually ductile and malleable. All these properties influence the way the metals are
used. Some metals are sonorous and this leads to special uses for them.
a. Explain the meaning of the highlighted words in italics.
b. Copper is ductile. How is this property useful in everyday life?
c. Aluminium is hammered and bent to make large structures for use in ships and aeroplanes. What property is
important in the shaping of this metal?
d. Name one metal that has a low density.
e. Some metals are cast into bells. What property must the chosen metals have?
f. Copy and add the correct word: Metals are good conductors of
and electricity. Name one
other physical property of metals and give two examples of how this property is useful.
26
 3-4 Research: Metals and their Properties
Use resources on the web to find information about the properties of metals in the table. Metal prices can be
found at the London Metals Exchange: http://www.lme.co.uk/
Approximate
cost ($A/kg)
World
production
(thousands of
tonnes per year)
Density
(grams
per mL)
Melting
point (oC)
Heat conductivity
Resistivity
(write in the actual
values and then rank
from 1 to 9)
(write in the actual
values and then rank
from 1 to 9)
Al
Cu
Au
Fe
Pb
Ni
Ag
Sn
Zn
Use the information on the data table to construct rankings of the properties listed. (The ranking for cost has
been done as an example.)
Highest
cost
Au
Highest annual
production
Highest
density
Highest melting
point
Best conductor of
heat
Best conductor of
electricity
Ag
Sn
Ni
Cu
Zn
Al
Pb
Fe
27
Refer to your rankings table to answer the following questions.
1. Which metal is the best conductor of heat?
2. Give two reasons why we don’t use this metal to make saucepans.
3. Which metals are the two best conductors of electricity?
4. Give a reason why one of these good conductors of electricity is used more than the other.
5. a. What method of transport would make use of a low density metal?
b. Which metal would be best to use for this method of transport?
6. Food cans are made of iron coated with tin. Compare the melting point of tin with that of iron. How could the
tin be put onto the iron?
7. a. What are plumb lines used for?
b. Which metal is used to make these? Find out why this metal was assigned its symbol.
c. Why was this particular metal used for plumb lines?
8. a. Name three uses for which plastics have replaced metals. You may need to discuss this with someone older
who can remember from what items were ‘traditionally’ made.
b. Why do you think this has happened?
28

3-5 Alloys – Modifying the Properties of Metals
(1 period)
Teacher Demonstration
Aims: 1. Make solder from elements and cast the solder in a mould
2. Compare some of the properties of solder with those of tin and lead.
Safety
!
This demonstration involves heating hazards. Wear goggles
and aprons. Care should be taken when handling hot items.
Bench protectors are to be used in this experiment.
Equipment



Equal quantities of lead foil and tin foil
Bunsen burner and heatproof mat
Evaporating dish with sand
Silica crucible
Tripod
Tongs
Wood splint
Pipe clay triangle
Sheet of steel
Introduction
Most metallic materials used today are not pure metals but mixtures of a
metal with one or more other elements. These mixtures are called
alloys. The elements in the mixture and the amount of each element
present affect the properties of the alloy, so it is possible to make alloys
that have the specific properties needed for a particular job. Alloys are
usually made by mixing a known ratio of each of the elements in their
molten (melted) states.
As in pure metals, alloy the
bonding is metallic. However,
Figure 3: Metallic structure
the presence of ions of different
sizes to other ions (e.g. zinc ions
amongst copper ions in brass) causes distortion in the regular
arrangement of the ions. As a result it is difficult for ions to slide past
each other, making an alloy harder and less malleable than any of the
pure metals it contains.
Alloys usually do not conduct electricity as well, because it is more
difficult for the electrons to flow through the irregular lattice. If the
various ions present are significantly different in size (or if they are
non-metals), the electrons may not be able to flow through at all because
the outer shells no longer overlap.
Figure 2: Alloy structure
Alloys melt at lower temperatures than any of the elements that
compose them, because the lattice is already partially disrupted, and
does not take as much energy to fully disrupt it. Solder melts more
readily than either of its components, tin and lead.
If the added atoms are much smaller than the atoms in the network,
like the carbon atoms added to iron to make steel, they can fit into
the holes between the layers of atoms in the network. When this
happens we call it an interstitial alloy. Steel has carbon atoms in the
interstitial spaces between iron atoms. Although steel is mostly iron,
it is not an element. Steel has 0.2 - 2.1% carbon, by mass. Compared
Figure 4: Steel is an interstitial alloy. Carbon to iron, steel is harder and stronger.
atoms sit between larger iron atoms.
29
Preparation Questions
Minerals to Metals PowerPoint
An alloy is a mixture of a metal with at least one other element. Steel is a common example of an alloy. It
contains iron mixed with carbon and other elements. Adding other elements to a metal changes its
and so changes its properties. The final alloy may have very different properties to the original metal. By
changing the amount of each element in an alloy, material scientists can custom-make alloy properties to fit a
given job.
Alloys have been used for thousands of years. Bronze, an alloy of
years ago. Other well-known alloys include:
and
, was used 5000
Brass: an alloy of copper and zinc. It does not tarnish and is used for door knobs, buttons and musical
instruments.
Solder is an alloy of
and
. It is used in electronics to attach components to circuit boards.
Alloy wheels are made from aluminium (90%) with other elements such as silicon, magnesium, manganese,
and iron to build strength and corrosion resistance.
Why is steel stronger than iron?
Steel is an alloy of iron and other elements, including carbon, nickel and
than pure iron and can be used for everything from bridges…
.
. Steel is stronger
The atoms in pure iron are arranged in densely-packed layers. These layers can slide over each other, making
pure iron a very soft material. The atoms of other elements are different sizes. When other elements are added
to iron, their atoms distort the regular structure of the iron atoms. It is more difficult for the layers of iron
atoms in steel to
each other and so this alloy is stronger than pure iron.
Draw the structure of pure iron and steel alloy below:
Pure iron
Steel alloy
What’s so clever about smart alloys?
Shape memory alloy is a type of smart material made from metals that returns to its original shape after being
deformed.
is a type of shape memory alloy made from nickel and titanium. This material can be
used to make a pair of glasses that ‘remembers’ its shape and does not break when crushed. Nitinol has also
been used to hold badly broken bones in place while they heal.
30
Practical Part A: Producing and Casting Solder
1. Place lead foil in a crucible, place on a pipe clay triangle and tripod, and heat
with a moderate flame using a Bunsen burner.
2. Slowly add an equal mass of tin foil, stirring the metal mixture with a wooden
splint.
3. Prepare a mould by making a shape in the moulding sand.
4. To make a casting of your alloy, carefully pour the molten metal mixture into
your prepared mould and allow it to cool.
Part B: Properties of Tin, Lead and Solder
1. Appearance: Record the colour and lustre differences of each metal.
2. Hardness: A simple test for comparing the hardness of materials relies on the fact that a harder material
will scratch a softer material. Compare the hardness of solder, tin and lead, using scratch test. Cross out the
incorrect phrase in the statements below:
Tin (scratches/is scratched by) lead.
Tin (scratches/is scratched by) solder.
Lead (scratches/is scratched by) solder.
3. Melting Point (Must be conducted with a Bench Protector)
a. Place a sheet of steel on the top of a tripod. Note: This is not to be placed hot on a bench.
b. Place small beads of solder, tin and lead on the steel sheet close together near the centre.
c. Place the Bunsen burner directly below the centre of the steel sheet and heat the steel sheet gently. Care
must be taken to leave the sheet of steel on the tripod. Do not remove it or place it on the bench.
d. Note the time taken for each metal bead to melt. What
does the time indicate regarding the melting points of each
material?
e. When all beads have melted, turn the burner off and note
the time taken for each molten bead to solidify.
Tin
Lead
Solder
Colour
Lustre
Time to melt
31
4. Melting Point
The melting point of tin / lead mixtures (by mass) is shown on the graph below:
Temperature
(oC)
400o
C
200o
C
100% Tin
as in pure
100% Lead
% Metal by mass
a. What would be the melting point of a tin-lead alloy of 80% tin, 20% lead mixture?
b. Which tin-lead mixture will produce the lowest melting point?
5. Using Alloys
There are thousands of alloys in use in our society. Here are a few.
a. Fill in the common names of the alloys mentioned below. (You may need to check the internet or your
dictionary). The choices are stainless steel, cupro-nickel, duralumin, brass, bronze, dental amalgam, solder,
tungsten/steel, magnesium alloy, alnico.
Alloy
Alloy Metals
70% Copper, 30% zinc
Aluminium, copper, magnesium,
manganese
Magnesium alloy
Magnesium, aluminium
Does not corrode. Attractive gold
colour. Easily worked.
Strong and light (low density).
Resists corrosion.
Light and easily cast. Hard, high
melting point.
Iron, carbon, tungsten, chromium
Very hard and resistant to wear.
70% Iron, 1% carbon, 20%
chromium and 9% nickel
95% Copper, 5% tin
Does not corrode. Shiny, strong and
hard.
Aluminium, nickel
70% Tin, 30% lead
75% Copper, 25% nickel
Mercury, silver
32
Special Properties
Shiny, attractive, does not corrode.
Excellent magnetic properties
Solid at room temperature, but melts
easily.
Cheap, light, hard. Does not corrode.
Soft, then sets hard and chemically
inert.
b. Choose an alloy suitable for the following jobs (give reasons for your choice)
Alloy
Reason
High Speed Drill
Joining Electrical Components
Jet Engine Parts
Kitchen Sink
Statues
Magnet
Coins
6. Match the part of the statements below with the correct answer in column B to make eight facts about alloys.
Write the complete sentences in the space below.
Statement
Answer
An alloy is a metal made by
A. coins
Alloys often combine
B. copper and zinc
Electrical solder melts at a temperature lower than
C. aluminium and copper
Brass is an alloy made from
D. lead and tin
Duralumin is an alloy made from
E. mixing two or more metals
Cupro-nickel is the alloy used to make
F. the properties of the metals
Bronze is an alloy used to make
G. alloys
Most metal objects used today are made of
H. statues
33
 3-6 Practical: Reaction of Metals with Water
(2 periods)
Aim: To investigate the reactions involving a variety of metals with water.
!
Safety
Wear goggles and aprons during this practical. Do not taste
any chemicals. Reactive metals are corrosive and should not
be handled with your skin. Test only one metal at a time.
Metals vary their reactivity in nature. Some may react explosively when exposed to water, while others very
stable and will not react, except with the strongest of acids. This property of relative reactivity affects the
location of mineral deposits on earth. In this practical, and the one following, trends in a series of reactivity can
be determined by comparing the rates of reaction.
Materials
Supplied to the teacher only
 Metals: sodium, calcium, potassium (to be distributed in small pieces only)
Supplied to each bench
 Dropper bottle of phenolphthalein (Basic solution indicator)
 Metals: magnesium, iron, copper, lead, zinc
Supplied to each group


Forceps (to pick up Na, K, and Ca pieces)
5 test-tubes


Hot water bath
Paper towel
400 mL beaker (for sodium, potassium and calcium experiments)
Watch glass (to go over 400mL beaker)
Sand paper (to clean metals)
Procedure
1. Your formal report is to be written into your practical book.
2. The Requirements and Procedure sections should refer to ‘As recorded in Minerals to Metals booklet’.
3. In the Results section, rule up a table to include your results of all metal reaction with water.
Test
Water + sodium
metal
Observation
Inference
In this column, write in detail what you
see happening: colour changes, time
taken to complete, changes of state, etc.
In this column, write what you think is
being produced, based on your observations
4. Half fill the 200 mL beaker with hot water from the urn. This will be the water bath.
5. Clean a piece of zinc with the sand paper. This is only to be done over a bench mat to protect the benches.
6. Put about 2 cm deep of water into a test-tube and add a few drops of phenolphthalein (the latter goes pink if
a basic solution involving the hydroxide ion (OH-) is formed).
7. Add the zinc to the water and check for a gas given off and/or a pink colour.
8. If no reaction has been observed, place the test-tube in the water-bath for a few minutes to check if it will
occur when heated.
34
9. Record your observations (or lack of reaction).
10. Record your inferences in your table. Has a reaction occurred? If so, what has been produced?
11. Repeat the steps 5 - 10 above for the remaining metals: magnesium, iron, copper, and lead.
12. Put about 2 cm of cold water in a 400 mL beaker. Add a few drops of phenolphthalein.
13. Collect a small sample of sodium, potassium and calcium from your teacher. One piece of metal at the time.
14. Use the forceps to transfer the calcium to the water and cover with a watch glass.
15. Record your results. List the metals from most reactive to least, based on your observations.
16. Repeat the calcium procedure for sodium and then potassium (only with your teacher present).
17. Write a Conclusion and Discussion for this experiment,
Questions
1. Why is it necessary to keep sodium and calcium under paraffin oil?
2. Why was it important to ensure that the metal surfaces were cleaned with sandpaper?
3. What is the function of the water bath?
4. You have observed some metals and their relative reactivites. Look for any patterns of these elements in a
periodic table and predict the chemical reactivity of the metals: Cs, Rb, Ag, Au, Pt.
5. Write a balanced equation for any occurred in the practical.
6. Complete the General Reaction and equations involving reactive metals and water below:
Reactive Metal + Water

Na(s)
+
H2O(l)

Ca(s)
+
H2O(l)

Zn(s)
+
H2O(l)

K(s)
+
H2O(l)

Li(s)
+
H2O(l)

Cu(s)
+
H2O(l)

35
 3-7 Practical: Reaction of Metals with Acids
(2 periods)
Aims: To investigate the reactions of metals with acids
To develop a series of metals in order of their reactivity
A formal report for this experiment is to be written into your practical book.
!
Safety
Wear goggles and aprons throughout this practical. All acids
are corrosive and should be handled with care. Do not taste
any chemicals, and clean up any spills with paper towelling.
Materials
For the teacher’s bench only
For each group
 10 M hydrochloric acid
 Hot water bath
For each bench
Set of 5 test-tubes





Forceps
1 M nitric acid
1 M hydrochloric acid
1 M sulfuric acid
Ice-cream container for the disposal of wastes
Metals: magnesium, iron, copper, lead, zinc
Test-tube rack
Sand paper
Stoppers for test-tubes
String for testing hydrogen
Procedure
1. Half fill the 200mL beaker with hot water from the urn for a water bath
2. Set up your written report in your practical book, as in Practical 1. Your results section should include a table
as shown below:
Test
Observation
Inference
Hydrochloric acid
+ zinc
In this column, write in detail
what you see happening: colour
changes, time taken to complete,
changes of state, etc.
In this column, write what you think is being
produced, based on your observations
3. Place a piece of zinc in a test-tube and add about 2 cm of 1M acid of your choice.
4. If a gas is given off, place a stopper lightly on the top of the test-tube and collect the gas for a short time. (Do
not put the stopper on tightly as the test-tube may crack)
5. Test for hydrogen gas by removing the stopper and applying a lighted taper. If no reaction occurs, warm the
test-tube in a water bath, and then test for hydrogen if a gas is evolved.
6. If a reaction does not occur, put the metal aside for later testing.
7. Record your observations and inferences in your Results table.
8. Repeat steps 3 - 7 above for the remaining metals.
36
Teacher Demonstration
9. Two metals will not show evidence of a chemical reaction so far. Your teacher will demonstrate the addition
of these to concentrated 10 M acids. Record as before.
10. Write a Conclusion for this experiment.
11. Discussion – consider other factors that could affect the outcome of each experiment.
Conclusion
1. Why are some of the reactions heated in a water bath?
2. Why is concentrated hydrochloric acid used for some metals?
3. Practicals 1 and 2 both involve reacting metals with water and acid. In completing both practicals, a
reactivity series of metals can be produced. List the metals in order of their reactivity from most to lest
reactive.
4. Complete the General Reaction involving reactive metals and acids below.
Metal + Acid

5. Write balanced equations for the reactions that occurred during the practical.
Zn(s)
+
H2SO4(aq)

Cu(s)
+
HCl(aq)

Mg(s)
+
HNO3(aq)

Pb(s)
+
HCl(aq)

Fe(s)
+
HNO3(aq)

37

3-8 Practical: Displacement of Metals
(1 period)
Aim: To investigate and predict the reaction of a metal placed in a solution of a metal salt.
A formal report for this experiment is to be written into your practical book.
Introduction: If you consider the following equations, you will notice that one is the reverse of the other.
Cu(s)
+
Zn(NO3)2(aq)

Zn(s)
+
Cu(NO3)2(aq)
Zn(s)
+
Cu(NO3)2(aq)

Cu(s)
+
Zn(NO3)2(aq),
One of the reactions will spontaneously proceed, while the other will not occur. Which? Today’s experiment
should help you identify a method to predict if displacement reactions will occur.
!
Safety
Wear goggles & aprons during this practical, which involve
heating hazards. Care should be taken when handling hot
items. Bench protectors must be used in this experiment.
Materials



Solutions of lead nitrate, copper sulfate, zinc sulfate and magnesium sulfate
Metals: lead, magnesium, zinc and copper
4 test-tubes
Test-tube rack
Marker pen
Method
1.
2.
3.
4.
Place one of each of the metals in a separate test-tube.
Add 2 cm of the solution you have been allotted by your teacher.
Label your test-tubes and set aside until the end of the lesson.
Record your observation and compare with those of the rest of the class in the table below (in some cases
this will be ‘No reaction’, and indicated by a cross X).
5. Record the results of the rest of the class in order to complete your table.
Metal
Copper
Lead
Magnesium
Zinc
38
Copper Solution
Lead Solution
Magnesium Solution
Zinc Solution
6. Write a conclusion below, and establish a general rule (we will call this: The Law of Displacement) that
should be applied to predict the likelihood of a displacement reaction occurring.
7. Predict if the following reactions will occur (Yes / No)
a. Cu(s)
+
2AgNO3(aq)

2Ag(s)
+
Cu(NO3)2(aq)
b. Mg(s)
+
Zn(NO3)2(aq)

Zn(s)
+
Mg(NO3)2(aq)
c. 3Pb(s) +
2Al2O3(s)

4Al(s)
+
3PbO2(s)
d. 2K(s)
+
FeS(s)

Fe(s)
+
K2S(aq)
e. Fe(s)
+
MgSO4(aq)

Mg(s)
+
FeSO4(aq)
8. Complete the following reactions by writing the correct formula and balancing the equation. Which two
reactions will not occur? (Write ‘No Reaction’)
a.
zinc
+ lead nitrate solution

b.
iron
+ zinc sulphate solution

c.
lead
+ copper nitrate solution

d.
lead
+ potassium nitrate solution 
e.
magnesium + zinc chloride solution

f.
copper
+ sodium chloride solution

g.
zinc
+ iron sulphate solution

h.
gold
+ silver nitrate solution

i.
silver
+ zinc chloride solution

j.
magnesium + calcium nitrate solution

oxidation loss of a electron, it makes them stable e.g. metal wanting to lose an electron taken by a non metal
Copper who underwent oxidation met up nitrate got the electron and they were happy in the bar, magnesium
is not happy because he is alone and only has 2 on the outer shell, magnesium sees copper and wants the
nitrate and takes the nitrate and gives one of his electron to nitrate and gives another electron to another
nitrate sitting in the bar
9. Three metals X, Y and Z have the following reactions. Y will displace X from a solution of its salt. Z will
displace both X and Y from solutions of their salts. Place the three metals in order of reactivity, starting with
the least reactive.
39
10. Here is a list of metals in order of decreasing reactivity. Q and R are mystery metals.
K > Q > Ca > Mg > Al > Zn > R > Fe > Cu
a. Will Q react with cold water?
b. Will R react with cold water?
c. Will R react with dilute hydrochloric acid?
d. Will R displace copper from copper sulphate solution?
e. Will Q displace potassium from a solution of potassium nitrate?
40
 3-9 Homework Exercise to Summarise Practicals 1 – 3
You have now completed a series of practicals that demonstrate chemistry of metals and their reactions. You
can now establish the metal reactivity series from observing their behaviour in acids and water. In addition,
you have also, for the first time, been able to predict if a reaction proceeds, as shown in displacement reactions.
This activity summarises the practicals that have helped you sort metallic elements according to reactivity.
1. Set up a table in your book with the following headings (leave room for 13 metals).
Metal
Reaction with
water
Reaction with
acids
Reaction with
Zn(NO3)2
Reaction with
Pb(NO3)2
Reaction with
CuSO4
2. Using the list below, the results of your practicals and any other appropriate source, fill in the metals given in
order of reactivity (the most reactive at the top).
Zinc, aluminium, tin, copper, potassium, platinum, lead, sodium, iron, magnesium, calcium, silver, gold
3. Fill in the table using a ‘’or an ‘X’ to indicate whether you consider that a reaction would occur.
Answer the following questions in complete sentences.
4. Which metal is the most reactive?
5. Which metal is the least reactive?
6. What would happen when zinc is added to hydrochloric acid?
7. What would happen when silver is added to hydrochloric acid?
8. Predict what would happen when copper metal is added to silver nitrate solution?
9. Which metals will react with cold water?
10. Which metal was more reactive, copper or silver?
11. Which is more reactive: lead or zinc?
12. Which metals are less reactive than copper?
13. Which metals are less reactive than zinc?
41

3-10 Extracting Metals from their Ores – Redox Chemistry
CHT p. 76 - 7
Getting the metal out
An ore is rock from which a metal is obtained. After mining the ore the next step is
the metal from it. This can be quite easy or very tough, depending on how reactive the metal is.
Native metals A few metals (the most unreactive ones like silver, gold, platinum, and sometimes copper) are
found as
. The metal is obtained by separating it from its impurities, just like removing
stones from soil. This is a physical process and does not involve any chemical reactions.
Other metals The other metals exist as
in their ores. This means the metal has to be
extracted using a chemical reaction.
Extraction by removing oxygen
Often the metal compound is an
or a compound that can easily be converted to an oxide. The
metal is obtained by removing the oxygen from the oxide. This process is called
.
Reduction is the removal of oxygen. A substance that can remove oxygen is called a reducing agent
ore
reduction

metal oxide
metal
(the oxygen is removed using a reducing agent)
Carbon and carbon monoxide are often used as reducing agents. As shown on the reactivity series, carbon is
more reactive than zinc. So carbon and carbon monoxide can remove oxygen from zinc, iron, and lead oxides.
For example carbon monoxide is used to extract iron from iron ore:
iron(III) oxide + carbon monoxide
 iron + carbon dioxide
Fe2O3(s) + 3CO(g)  2Fe(l) + 3CO2 (g)
Extraction by electrolysis
Electrolysis means
. It is the most powerful way to extract a metal from its
compounds. First the compound is melted or dissolved so that the ions are free to move. Then a current is
applied to separate them. Since electricity is expensive, this method is used only for more reactive metals that
are hard to extract in other ways, such as sodium from molten sodium chloride:
sodium chloride  sodium + chlorine
2NaCl(l)  2Na(l) + Cl2 (g)
Again oxygen is removed, which means reduction takes place, But how? First the aluminium oxide is separated
into aluminium and oxide ions. Then the aluminium ions gain electrons to become aluminium atoms:
Na+ + e-
 Na
So this brings us to a wider definition of reduction: Reduction is a gain of electrons.
A mnemonic (memory aid) to help learn what happens to these electrons is:
OIL RIG (Oxidation is
- Reduction is
) of electrons.
Reduction is the opposite of oxidation
Look again at the reduction of iron(III) oxide:
iron(III) oxide + carbon monoxide
 iron + carbon dioxide
The iron(III) oxide lost oxygen. It was reduced. But the carbon monoxide gained oxygen. It was
. Reduction is the opposite of oxidation. They always take place together.
42
Reduction and oxidation in electrolysis
For a substance to be electrolysed, it needs to be either molten or in a solution. The sodium ions then gain
electrons. They are reduced.
But oxidation and reduction always take place together. So, what happens to the chloride ions? They lose
electrons and produce chlorine gas:
2Cl-

Cl2 + 2e-
And this brings us to a wider definition for oxidation: Oxidation is the loss of electrons.
Redox reactions
Since oxidation and reduction always take place together, the overall reaction is called a redox reaction.
(Reduction plus oxidation.) All chemical extraction methods involve redox reactions.
Displacement reactions are also of redox type. One metal element is oxidised while the other metal ion is
reduced.
Zn(s) +
Pb(NO3)2(aq)

Pb(s)
+
Zn(NO3)2(aq)
The zinc loses electrons and is oxidised. The lead accepts those electrons and is reduced.
Zn  Zn2+ + 2e-
/
Pb  Pb2+ + 2e-
Displacement reactions are Redox
In some reactions, metals compete with each other for other anions. This type of reaction is known as a
displacement reaction. With this type of competitive reaction, the reactivity series can be used to predict
which of the metals will ‘win’.
In a displacement reaction, a more reactive metal will displace a less reactive metal from a solution of its salt.
Zinc is above copper in the reactivity series. If a piece of zinc metal is left to stand in a solution of copper
nitrate, the copper (II) nitrate slowly loses its blue colour. The zinc continues to displace the copper from the
solution and eventually becomes colourless zinc nitrate.
zinc + copper(II) nitrate  zinc + copper nitrate
Zn(s) + Cu(NO3)2(aq)  Zn(NO3)2(aq) + Cu(s)
This is a redox reaction involving the transfer of two electrons from the zinc metal to the copper ions. The zinc
is oxidised to zinc ions in aqueous solution, and each atom loses two electrons.
Zn(s)  Zn2+(aq) + 2eThe copper ions are reduced to copper metal, gaining two electrons that were donated by the zinc. This
transfer of electrons takes place at the surface of the zinc.
Cu2+(aq) + 2e-  Cu(s)
 Questions
1. Why is no chemical reaction needed to extract gold?
2. Lead is extracted by heating its oxide with carbon:
lead oxide + carbon  lead + carbon monoxide
a. Why can carbon be used for this reaction?
b. One substance is reduced. Which one?
c. Which substance is the reducing agent?
d. Which substance is oxidised during the reaction?
3. Give two definitions for each of these:
a. reduction b. oxidation
5 Sodium is extracted from rock salt (sodium chloride).
a. Electrolysis is needed for this. Explain why.
b. Write a word equation for the reaction.
c. Say which ions are oxidised, and which reduced.
43
 3-11 Video: Redox Rocks!
(1 period)
So far in this topic, you have studied the reactivity series of metals. This video shows how displacement of
metals will happen. Metals that are lower down on the list will always displace anything above it. Four metals
(Au, Al, Zn and Cu ) of different reactivity are examined and the different processes needed to extract them
from ores and to prepare pure metals are examined.
When copper reacts with oxygen, it forms copper oxide. The copper has two available electrons in
; the oxygen has room for
. The oxygen takes the two electrons from the
copper and both atoms have become
. We say that the copper has been
and the oxygen has been
.
The term REDOX comes from combining the words Oxidation and Reduction. As these two processes always
occur together REDOX is the term used to describe these reactions. This video describes the processes of
oxidation and reduction in an industrial context.
Oxidation can be described as:

gain of

loss of

loss of
Reduction can be described as:

loss of

gain of

gain of
OIL RIG is the mnemonic used to help remember this:
O
R
I
I
L
G
In order to extract the metal from its oxide, we reverse this reaction and put the electrons back on the ions.
For example, to extract copper, we can place an
the electrons more strongly than iron.
nail into a copper solution. The copper attracts
Write the half-equation for the reduction of copper ions:
Write the half-equation for the formation of iron ions:
In the beaker, we see copper metal forming a coating on the iron nail surface. This is an example of a
reaction.
44
Extracting Copper
How would you describe copper’s position on the reactivity series?
Copper can be found in nature as pure metal due to its position in the reactivity series. How would you explain
this fact?
Underline two common copper minerals: Galena chalcopyrite magnetite rutile sphalerite cuprite
Copper forms compounds with sulfur. Roasting at high temperature in air reduces them. Balance the
equations:
Cu2S(l) + O2(g) 
Cu2O(l) + SO2(g)
Cu2O(l) + Cu2S(l) 
Cu(l) + SO2(g)
What is the half-equation for the reduction of copper?
What is the half-equation for the oxidation of sulphur?
Combine these two half-equations to give the REDOX reaction. (Electrons cancel on both sides of the equation,
since the number of electrons
equals the number of electrons
.)
Extracting Zinc
Zinc is more reactive than copper because it holds its electrons more weakly. It is never found as a pure metal.
How would you describe zinc’s position on the reactivity series relative to copper?
Is zinc ever found in nature as a pure metal? Why or why not?
Underline the common zinc mineral: Galena chalcopyrite magnetite rutile sphalerite cuprite
Zinc ore is crushed and roasted in a furnace, but insoluble zinc oxide is formed. Adding strong acid produces
soluble ions, available for electrolysis. Electric current is passed through the solution.
Zinc ions are produced at the negative terminal:
. Zinc ions are reduced to form pure
. The positive terminal is the
at the cathode.
Write the half-equation for the reduction of zinc:
45
Write the half-equation for the oxidation of hydrogen:
Combine these two half-equations to give the REDOX reaction. (Electrons cancel on both sides of the equation)
The most common use of zinc is the coating of iron/steel products, know as g
.
Aluminium
Copper and zinc can be via the reduction of its compounds using electrolysis of solutions. However, aluminium
is produced from its molten ore.
How would you describe aluminium’s position on the reactivity series relative to zinc and copper?
Aluminium is abundant in the earth’s rocks, such as clays and gemstones. Aluminium is commonly found in
ore and its mineral is
: hydrated aluminium oxide – Al2O3.3H2O.
(We will not study the Bayer Process in this topic)
Aluminium is reduced from alumina at
temperature and uses large quantities of electricity. The
smelter runs with a current of
amps. Alumina melts at about
Celsius. Additives
are added to reduce this melting temperature and save on energy costs.
What is the half-equation for the reduction of
Aluminium at the cathode? The positive aluminium
ions attract to the negative electrode and form
aluminium metal.
What is the half-equation for the oxidation of carbon
at the anode?
Combine these two half-equations to give the REDOX
reaction. (electrons cancel on both sides of the equation)
Because aluminium is high on the reactivity table, it is
into atoms.
Aluminium ingots are melted into shapes or sheets to produce foil, cans, w
.
46
to convert aluminium ions back
and
 3-12 Demonstration Reaction: Mighty Thermite
Thermite is a pyrotechnic displacement reaction between a reactive metal
powder and a metal oxide. It produces a highly exothermic oxidationreduction reaction known as a thermite reaction.
Red iron(III) oxide is the most common iron oxide used in thermite, although
magnetite also works. Aluminium is the more reactive metal that displaces the
iron from its oxide. Molten iron metal is produced.
Although the reactants are stable at room temperature, they react with an
extremely intense exothermic reaction once they are ignited. Temperatures of
up to 2,500° can be obtained in the thermite reaction. Thermite contains its
own supply of oxygen and does not require any external source of air for the
process. Consequently, it cannot be smothered and may ignite in any
environment, given sufficient initial heat. It will burn well while wet and cannot be easily extinguished with
water.
!
Safety
This demonstration produces intense heat and light. It should
be done in a fume cupboard, with heat-proof mats covering all
base. Do not stare at the light.
Your teacher will demonstrate the thermite reaction. A strip of magnesium metal is used as a fuse, which
provides heat to initiate the reaction. This reaction gives off intense heat and light; so avoid staring at the
flame.
Record five observations of the reaction.
Overall Equation
Thermite is an example of exothermic redox reaction. Write down the oxidation and reduction ionic equations.
Oxidation:
Reduction:
Thermite is not explosive, but exposes a very small area of metal to extremely high temperatures. The intense
heat generated can be used to cut through metal or weld metal components together. For example, broken iron
railway tracks, where the repair can take place on site without removing the track.
a. What is the function of the aluminium metal in this reaction?
b. Could copper metal be used instead of the aluminium? Explain.
c. Could copper oxide be used instead of iron oxide? Explain.
47
 Questions
1. Predict whether or not the following reactions will take place:
a.
b.
c.
d.
magnesium + copper oxide
iron + aluminium oxide
calcium+ magnesium oxide
Complete the word equations, and write balanced chemical and ionic equations for those reactions
that do take place.
2. Use the following list of metals to answer the following questions: iron, calcium, potassium, gold, aluminium,
magnesium, sodium, zinc, platinum, titanium.
a.
b.
c.
d.
e.
f.
Which of the metals is found native?
Which of the metals is found in nature as the ore: i. rock salt? ii. rutile?
Which metal has a carbonate found in nature called chalk?
Which of the metals will not react with oxygen to form an oxide?
Which of the metals will react violently with cold water?
Choose one of the metals in your answer to e and write a balanced chemical equation for the
reaction that takes place.
g. Which of the metals has a protective oxide coating on its surface?
h. Which of the metals reacts very slowly with cold water but extremely vigorously with steam?
i. Which of the metals is used to galvanise iron?
3. Explain the following:
a.
b.
c.
d.
Metals such as gold and silver occur native in the Earth's crust.
The parts of shipwrecks made of iron rust more slowly in very deep water.
Zinc bars are attached to the structure of oilrigs to prevent them from rusting.
Copper roofs quickly become covered with a green coating when exposed to the atmosphere.
Pyrite mineral
48
3-13 Extraction Methods
Most metals are found in the Earth’s crust combined with other elements in rocks known as ores. Minerals are
the metal-bearing compounds and have a fixed formula. e.g. Iron is commonly found combined with oxygen in
ores called haematite (Fe2O3) and magnetite (Fe3O4).
Metals need to be extracted from minerals before they can be turned into elements. The minerals have to be
separated from the ore body, then the metal is chemically reduced using one of two processes: Pyrometallurgy
or Electrolysis.
Metals can be found in the Earth’s crust combined with other elements or uncombined as pure substances.
Some unreactive metals, like gold, silver and copper, can be found uncombined as elements and are said to
occur ‘native’. Most metals are found combined with other elements, as compounds in ores. These metals must
be extracted from their ores before they can be made useful.
The location of ore deposits is related to the reactivity series. Highly reactive elements formed soluble salts
that dissolved in the ocean, while unreactive metals are found native.
1. Complete the table below from the PowerPoint slide.
Li
Exist as salts in
oceans
increasing reactivity
Carbonates in
limestone
Forms silicates in
rocks
Found as oxides in
rocks
Found as sulfides
in volcanic regions
K
Na
Ca
Mg
Metals above carbon in the reactivity
series must be extracted using
electrolysis. (Electrolysis can also be
used to purify copper.)
Al
(carbon)
Zn
Fe
Sn
Metals less reactive than carbon can be
extracted from their ores by reduction
using carbon, coke or charcoal.
Pb
(hydrogen)
Cu
Found native as
unreacted
elements
Ag
Au
Platinum, gold, silver can occur native
and do not need to be extracted.
Pt
Extraction methods vary for different metals. Metals that are found native do not require further extraction.
Most metals need to be separated from other elements from compounds using chemical reactions. There are
two main ways of extracting metals from ores:
Pyrometallurgy – burning ores with carbon (reduction)
Electrolysis – using DC electric current in solutions
How does reactivity affect the method of extraction?
49
 3-14 Chemical Reactions involving Metals
The reactivity series aids us in predicting the viability of a chemical reaction. You need to know these reaction
types and their products. Write the chemical formulas of all products below and balance the equations.
Reactions of metals with oxygen
Metals react in the presence of oxygen to form metal oxides. Reactions with oxygen are called oxidation
reactions. Silver, gold and platinum do not react with oxygen in the air.
Metal +
Oxygen
 Metal oxide
Al(s)
+
O2(g)

Zn(s)
+
O2(g)

Fe(s)
+
O2(g)

Sn(s)
+
O2(g)

Pb(s)
+
O2(g)

Cu(s)
+
O2(g)

Metal Reactions with water
Very reactive metals react vigorously with water to produce a metal hydroxide and hydrogen gas.
Metal
+
Water
 Metal Hydroxide + Hydrogen
Li(s)
+
H2O(l)

Na(s)
+
H2O(l)

K(s)
+
H2O(l)

Metals near the middle of the series react with steam. (Note the gaseous state of water)
Metal
Water (Steam)  Metal oxide + Hydrogen
+
Mg(s) +
Al(s)
+
H2O(g)

H2O(g)

Zn(s)
+
H2O(g)

Fe(s)
+
H2O(g)

Metal Reactions with dilute acid
Many metals will react with acids, and effervescing hydrogen gas is evolved.
Metal + Acid  Hydrogen gas + Salt
Na(s)
+
HCl(aq)

Al(s)
+
HCl(aq)

Zn(s)
+
H2SO4(aq)

Fe(s)
+
HNO3(aq)

Pb(s)
+
H2SO4(aq)

50
Reaction of metal with:
Element
Water
Dilute HCl
Symbol
Air (when heated)
Potassium
Sodium
K
Reacts with cold water
Violent
reaction
Calcium
Magnesium
Reacts with hot water
Aluminium
Protected by oxide layer
Zinc
Iron
Tin
Lead
Reacts with steam
Na
Reacts, but gets less
vigorous
Mg
Al
Reacts, but
gets less
vigorous
Reacts slowly with steam
Zn
Fe
Reacts slowly
Copper
Silver
Ca
Sn
Pb
Cu
No reaction
No reaction
Gold
No reaction
Ag
Au
1. Complete the following general equations, and provide a balanced example of each.
Acid + reactive metal
Water + reactive metal
Steam + metal
2. Refer to the above list of metals and their reactions.
a. Which metal is stored in oil?
b. i. Which metals react with cold water?
ii. Which gas is given off in this reaction?
c. i. Choose one metal that will not react with cold water but will react with steam.
ii. Name two products of this reaction.
d. i. Name one metal that reacts slowly with dilute hydrochloric acid.
Chalcopyrite crystal
ii. Name the two products of this reaction.
e. Which of the metals will not react with oxygen when heated
f. Which metal forms an oxide when heated in air but does not
react with dilute acid?
51
 3-15 Mineral Calculations – Composition of Common Ores
Relative Atomic Masses (RAMs) are found in the Periodic Table on the back page.
The theoretical percentage by mass of a metal in a compound can be calculated by using the relative atomic
masses (RAMs). e.g. Calculate the percentage of calcium in calcium carbonate:
% Ca in CaCO3
=
RAM (Ca)
RAM (Ca) + RAM (C) + 3×RAM (O)
=
40
40 + 12 + (3 ×16)
=
40%
×100
1
×100
1
1. Calculate the percentage by mass of the metals in each of the minerals below:
Metal
Mineral
Formula
Percent
Metal
Metal
Mineral
Sphalerite mineral
Formula
Copper
Chalcopyrite
CuFeS2
Mercury
Cinnabar
HgS
Calcium
Limestone
CaCO3
Titanium
Rutile
TiO2
Iron
Haematite
Fe2O3
Aluminium
Bauxite
Al2O3
Iron
Magnetite
Fe3O4
Zinc
Sphalerite
ZnS
Lead
Galena
PbS
Tin
Cassiterite
SnO2
Percent
Metal
The percentage composition can be used to calculate the mass of metal produced from minerals.
2. How much iron can be obtained from an iron ore stockpile containing 80 tonnes of haematite (Fe2O3)?
3. Galena is mined at Broken Hill, NSW. If 50 tonnes of galena is mined per day, how much lead can be
extracted?
4. A metallurgist wants to produce 10 tonnes of aluminium. What mass of bauxite must be needed to obtain
this amount?
5. How much limestone is needed to produce 500 g of calcium metal?
6. What is the minimum amount of sphalerite required to obtain 50 kg of zinc?
7. How much mercury is present in 200 g of cinnabar?
52
  3-16 Oresome Froth - Student Worksheet
(2 periods)
Open the Oresome Froth Flotation program and complete the activity. The file is located on the school intranet.
This worksheet is to record your Oresome Froth investigation into the effect of variables on the process of froth
flotation. Complete this worksheet and include a screenshot of your most profitable setup – print it and name it
for submission.
Level I - Trainee Metallurgist
1. Investigate the effect of particle size
Air rate =
Collector =
Milling cost
Particle Size
pH =
Product grade
Product recovery Total operating cost
Conclusion:
2. Investigate the effect of amount of collector
Air rate =
Milling cost
Particle size =
Amount of collector
pH =
Product grade
Product recovery Total operating cost
Conclusion:
3. Investigate the effect of air rate
Amount of Collector =
Milling cost
Particle size =
Air rate
pH =
Product grade
Product recovery Total operating cost
Conclusion:
53
4. Investigate the effect of pH
Amount of collector =
Particle size =
Milling cost
pH
Air rate =
Product grade
Product recovery Total operating cost
Conclusion:
5. Optimum conditions
Particle size
pH
Amount of
collector
Air rate
Product grade
Product
recovery
Total operating
cost
Level 2 - Metallurgist
Particle size
Air rate (l/s)
pH
Amount of collector (kg/tonne)
Product grade
Product recovery
Total operating cost
Cu Cell 1
Cu Cell 2
Level 3 – Chief Metallurgist
Particle size
Air rate (l/s)
pH
Amount of collector (kg/tonne)
Product grade
Product recovery
Total operating cost
Cu Cell 1
Cu Cell 2
Zn Cell 1
Zn Cell 2
Summary. Describe the balance of conditions that are required to produce copper from its ore.
Source: Minerals Council of Australia
54
 3-17 Teacher Demonstration: Froth Flotation of Copper Ore
(1 period)
Copper is a very valuable metal for use in our modern society. More than half of the copper consumed is used
in electrical generators and motors, electrical wiring, television sets, computers and almost everything
electrical.
Copper is mined in at the Olympic Dam copper-uranium-gold deposit in
South Australia and the Mount Isa copper-lead-zinc deposit in
Queensland. Many copper ores contain less than 1% copper, and
would be uneconomical to extract without effective concentration
technologies.
Mining companies use the froth flotation process to concentrate the
metal-bearing mineral in an ore. It is a physical process that uses less
energy to separate the mineral from the waste material (gangue). The crude
ore is ground to a fine powder and mixed with water and frothing reagents.
When air is blown through the mixture, mineral particles cling to the bubbles
and rise to form froth on the surface. The waste material (gangue) settles to
the bottom and can be removed. The froth is skimmed off, and the water and
chemicals are distilled or otherwise removed for recycling, leaving a clean concentrate. This flotation process is
used to concentrate a number of minerals, such as silver, lead and copper.
The copper minerals and waste rock are separated at the mill using froth flotation. The copper ore slurry from
the grinding mills is mixed with ground-up limestone to give a basic pH, pine oil used as frother to make
bubbles, an alcohol to strengthen the bubbles, and a collector chemical called potassium amyl xanthate.
The ore is ground to a certain particle size – fine enough to release the mineral grain from
the rock, but producing too small a particle size is expensive and difficult to
recovery.
The xanthates are added to the slurry in relatively small quantities.
Xanthate is a long hydrocarbon (5 carbons) chain molecule. One end of the
chain is polar and sticks to sulfide minerals while the other end is nonpolar,
containing the hydrocarbon chain is hydrophobic – it hates being in the water and is attracted to the nonpolar
hydrocarbon pine oil molecules.
The froth must be strong enough to support the weight of the mineral floated and yet not be tenacious and
non-flowing.
The effectiveness is dramatically affected by pH and flotation circuits are often operated at a pH range of 7.5 to
11.5. Lime is often used to raise the pH. Raising the pH causes the polar end to ionize more and to
preferentially stick to chalcopyrite (CuFeS2) and leave the pyrite (FeS2) alone. Air is blown into the tanks and
agitated like a giant blender, producing a foamy froth. The chalcopyrite grains become coated with xanthate
molecules with their hydrophobic ends waving
around trying desperately to get out of the water.
They attach themselves to the oily air bubbles
which become coated with chalcopyrite grains as
they rise to the surface and flow over the edge of
the tank. In this manner through a series of steps
the copper ore is concentrated to an eventual value
of nearly 30% copper. Waste rock particles do not
adhere to the bubbles and drop to the bottom of the
tank. The waste ground-up rock that comes out of
the bottom of the tanks at the tail end of this
process is called "tailings”.
55
 3-18 Teacher Demonstration: Extraction of Copper from Ore
(1 period)
Aim: To model the extraction of copper metal from copper oxide ore
Introduction
Some metals can be extracted by strongly roasting minerals in the presence of coke, a semi-refined form of
coal. The modern blast furnace uses this same technology to extract iron.
!
Safety
Wear goggles and aprons during this practical, which involves
heating hazards. Care should be taken when handling hot items
with tongs. Bench protectors are to be used in this experiment.
Teacher Note: Ensure that the windows are open to allow adequate ventilation.
Materials



Carbon block
Tongs
Sodium carbonate flux
Blow-pipe
Tin ore
Paraffin oil
Method
1. Set up the apparatus as shown in the picture.
2. Scrape a small amount of carbon powder into the
hollow in the carbon block.
3. In the hollow, make a paste of:




1 spatula of tin (or copper) oxide
½ spatula of sodium carbonate
A few drops of paraffin oil
Carbon powder - scratched from the block, or ½ spatula of graphite (carbon) powder
Note: Sodium carbonate is used as a flux. This is a substance that, mixed with the oxide, lowers its melting point
and enables the reaction to occur at the Bunsen flame temperature. The paraffin oil is used to make a paste so
that fine hot oxide does not blow away. Paraffin burns to produce carbon dioxide and water and hence does
not affect the reaction.
4. Light the Bunsen burner and adjust the flame to a reducing flame (with the hole half closed and a large cone
of blue unburnt gas in the centre).
5. Place a rubber tip over the mouth-piece of the pipe (for hygienic reasons). Place the tip of the pipe into the
blue unburnt cone of gas in the centre of the flame. Use the blow-pipe to blow a steady stream of air and
unburnt gas over the ore.
Note: Care must be taken to ensure that your face and body is kept well clear of the flame.
6. When you have achieved globules of metal, allow the carbon and metal to cool to finger temperature (it
should not feel at all warm). View the metal globules block under a stereo-microscope. The charcoal blocks
will probably keep burning during and after the exercise. They can be cooled by dripping water on them. Do
not put them on hot benches until they have been cooled with water.
7. Record your observations – in detail.
56
 3-19 Teacher Demonstration: Electroplating of Solutions
(1 period)
CHT p. 82-3, p. 86-7
Electrolysis is a relatively modern method of producing elements by the use of electricity. The elements
potassium, sodium, barium, calcium and magnesium were all discovered using electrolysis by English chemist
Humphry Davy – all had to await the invention of electricity to be discovered.
Passing an electric current through compounds decomposes them. This can only happen if the compound is:
• made of ions
• molten, or dissolved in water so that the ions are free to move.
Electroplating is used to plate kitchen or bathroom taps with chromium to make them look good and to protect
them from rust. Steel is plated with tin to make ‘cans’ for food.
Materials
250 mL beaker
2 leads with alligator clips
nickel strip
steel wool
100 mL nickel sulfate solution
copper strip
power pack
1. Use the steel wool to clean a copper strip and a nickel strip. Place the strips in a
small beaker containing 40 mL of copper sulfate solution as shown.
2. Connect the copper strip to the positive side of a 6 Volt D.C. power supply and the
nickel strip to the negative side. Switch on and observe the changes that occur at the
negative electrode for a period of 5 minutes.
3. Fully label the diagram: positive/negative electrode, anode/cathode, ions present
in solution.
4. What happens at the negative electrode when the switch is closed and electricity
is passed through nickel sulfate solution? This is a reduction reaction.
5. Write an equation to show what happens to the Ni2+ ions in the solution.
Questions
1. What type of compound undergoes electrolysis?
2. For the electrolysis of molten lead bromide, draw diagrams to show:
a. how the ions move when the switch is closed
b. what happens at the positive electrode
c. what happens at the negative electrode
3. Name the products at the electrodes when these molten compounds are electrolysed:
a. magnesium chloride
b. calcium bromide
4. a. Draw a labelled diagram to show how you would plate an iron nail with nickel.
b. Write equations to show the changes that would take place at the positive and negative electrodes.
57
 3-20 Practical 7: Extraction of Copper from its Ore
(3 periods)
CHT p. 86-7
Aim: To extract copper and copper sulfate from a copper ore (Malachite)
Introduction
Once mined, a mineral undergoes a series of physical and chemical processes to produce the valuable metal we
require in a modern society. This practical looks at the steps involved in producing copper from its mineral,
malachite.
!
Safety
Wear goggles and aprons during this practical. Sulfuric
acid is corrosive and should be handled carefully. Wash
your hands at conclusion of the practical.
Preparation Question: Read the practical and complete the flow chart below:
Materials
For the teacher’s bench only

6M Nitric acid
For each group








Malachite Ore
Malachite ore
2 beakers: 250mL, 50mL
2M sulfuric acid
Deionised water
Power supply
Measuring cylinder 50mL
Heating equipment
Measuring cylinder 10mL
Mg(NO3)2, 8 M HNO3
Iron nail, Ni electrodes
Filter funnel
Filter paper
Hand lens, white tile, and stopper
Sand paper, evaporating dish
1.
Sulfuric acid
2.
Solid waste
to bin
3.
Filtrate
Method
1. Crush
4.
5.
6.
Product:
Product:
Product:
Consider a simple and effective method to break the ore into smaller pieces, and then carry it out.
2. Dissolve
a. Place the crushed ore in a beaker and slowly and carefully add 4 mL of sulfuric acid.
b. Wait until the frothing stops and add another 2mL.
c. Continue to add small amounts of acid, until the frothing stops and the mineral is reacted.
58
Describe the reaction, including the reaction rate and the colour of the reactants and products.
3. Filter
a. Cool the mixture and allow a few minutes for the
mixture to settle.
b. Set up a filtration apparatus, and slowly decant (pour
without disturbing the solid) the mixture through the
filter paper and collect in a 50 mL measuring cylinder.
c. Examine the solid waste in the filter paper, and then
discard in bin.
Identify the colour of the filtrate (solution), and describe
the filter paper contents (residue).
4. Displacement
a. Place the large rubber stopper into a 50 mL beaker, pour the filtered solution, and add deionised water until
the volume in the beaker is 40 mL.
Figure: Galena (PbS)
Note: The solution will still have excess acid and should not be handled.
b. Place a clean iron nail and a piece of nickel in the solution and leave
for 5 minutes.
c. Use tweezers to remove the metals and observe any changes.
Describe the reactions on the surface of both metals.
What is deposited on the nail?
Explain any differences that occurred between the iron and nickel.
59
5. Electrolysis
a. Connect the nickel electrolysis electrodes to the D.C.
power supply and set it to 6 Volts.
b. Switch on the power supply and observe the reactions
occurring at the electrodes.
c. Why is it essential to run this experiment on D.C.?
d. What is the function of the rubber stopper?
e. At which electrode is copper deposited?
f. Describe what occurs at the other electrode.
g. After 5 minutes, switch off the power, disconnect the electrodes and allow them to dry.
h. Take your metallic electrode and check it with your teacher, who will place it on a white tile and add a few
drops of nitric acid.
i. Describe this reaction and comment on its inference.
j. Repeat steps b. and c. using clean nickel electrodes submerged in a magnesium nitrate solution.
k. After 5 minutes, switch off the power and disconnect the electrodes.
Describe this reaction and compare it to that which occurred in the copper solution. Explain any similarities /
differences.
60
7. Summary
Write a Conclusion for the practical, noting the changes of copper during each step.
Write an equation for the reaction occurring in Step 5.
Suggest how you could calculate the percentage mass of copper present in the original ore.
Suggest an appropriate reduction method for each of the following metal compounds.
(Electrolysis, Roasting, Roasting with Carbon)
Sodium chloride
Silver sulfide
Zinc carbonate
Iron oxide
Magnesium chloride
Aluminium oxide
Structure of Galena
61
 3-21 Production of Iron – The Blast Furnace
PowerPoint: Minerals to Metals. Complete the missing terms and equations and label the diagram.
Metals towards the middle of the reactivity series may be extracted by reducing the metal oxide with carbon.
Metals extracted using this method include iron,
, zinc, chromium and
.
Iron is extracted from its oxides, haematite (Fe2O3) or magnetite (Fe3O4), in a blast furnace. The blast furnace is
a 50 metre high steel tower lined with heat-resistant bricks. It is loaded with the
of
(usually haematite),
(made by heating coal) and
(calcium carbonate).
Large amounts hot air are blown in near the bottom of the furnace through holes (
) which makes the
charge glow red hot, as the coke burns in the preheated air to produce carbon dioxide gas. This is a
. This reaction is very
and provides most of the heat for the furnace.
carbon + oxygen  carbon dioxide
Equation:
A number of chemical reactions then follow. The limestone begins to
at high temperature:
calcium carbonate  calcium oxide + carbon dioxide
Equation:
The carbon dioxide gas produced reacts with more hot coke higher up in the furnace, producing carbon
monoxide in an endothermic reaction.
carbon dioxide + coke  carbon monoxide
Equation:
Carbon monoxide is a
. It rises up the furnace and surrounds the hot iron oxide ore. It
reduces this ore to produce the iron metal. The carbon monoxide is itself
(oxygen is added to it).
The overall reaction is a REDOX reaction. The dense
trickles to the bottom of the furnace.
This takes place at a temperature of around 700°C.
iron(III) oxide + carbon monoxide  iron + carbon dioxide
Equation:
Removing impurities
The most common impurity in iron ore is sand (silica), silicon dioxide, SiO2, but others such as sulfur,
magnesium,
and manganese are also found in smaller amounts. These ores contain at least
60% iron, mixed with sand (SiO2). The presence of these in iron make it brittle and easily corroded. Removing
these impurities improves the quality of the iron and strengthens it for use as a construction material.
Limestone (CaCO3) is commonly used to remove the major impurity of
.
Calcium oxide is a base and this reacts with acidic impurities such as silicon dioxide in the iron to form a slag,
which is mainly calcium silicate. Silica is acidic, and the limestone reacts with it in a
reaction.
calcium oxide + silicon dioxide  calcium silicate
Equation:
The slag trickles down the furnace, but because it is less dense than the molten iron, it floats on top of it.
The molten iron, as well as the molten slag, may be drained off into torpedo cradles at regular intervals when
they fill at the bottom.
Waste material
The hot waste gases, mainly nitrogen and oxides of carbon (CO and CO2) escape from the top of the furnace.
is in the air blast and does not take part in any reaction. They are used in a heat exchange
process to
and so help to reduce the energy costs of the process. Slag is the other
62
major waste material, and many tonnes are produced daily. Once cooled and solidified, builders use slag for
or as road fill material.
Clearly label the diagram of the Blast Furnace.
The extraction of iron is a continuous process, running almost every day of the year, and is much cheaper to
run than the costly electrolytic method.
What types of steel are there?
Molten iron is impure, with about 4% carbon, and is known as pig iron. When cooled and solidified, pig iron is
and rusts easily. Reducing the
content and adding other metals into steel – an
alloy, further refines most iron.
Varying the amount of carbon gives steel different properties. For example, higher carbon content makes
harder steel. High carbon steel (> 0.5% carbon) is hard and tough, but
. It is used for cutting tools.
Mild steel (< 0.25% carbon) is
. It has good tensile strength but poor
resistance to corrosion. It is commonly used as a general-purpose material.
63
Stainless steel is an alloy of iron that contains at least 15%
, with smaller amounts of nickel
and magnesium. It is hard and tough, resists wear and corrosion and used for
and kitchen tap
ware.
Questions
1. Name the raw materials for extracting iron. Give their chemical formulas.
2. Write an equation for the reaction that reducing iron from its ore.
3. The calcium carbonate in the blast furnace helps to purify the iron. Explain how, with equations.
4. What is the ‘blast’ of the blast furnace?
5. Name each of the waste gases from the blast furnace.
6. The slag and waste gases are both useful. How?
7. What is the difference between cast iron and steel?
8. What are some properties of pig iron?
9. Steels with higher carbon content tend to be (more/less) hard than others.
10. What is stainless steel made of? Name some common uses of stainless steel.
11. Iron is extracted from its ores haematite and magnetite. The ore is mixed with limestone and coke and
reduced to the metal in a blast furnace. The following is a brief outline of the reactions involved.
coke + oxygen  gas X
gas X + coke  gas Y
iron(III) oxide + gas Y  iron + gas X
a. Name the gases X and Y.
b. Give a chemical test to identify gas X.
c. Write balanced chemical equations for the reactions shown above.
The added limestone is involved in the following reactions:
limestone  calcium oxide + gas X
calcium oxide + silicon dioxide  slag
d. Give the chemical names for limestone and slag.
e. Write balanced chemical equations for the reactions shown above.
f. Describe what happens to the liquid iron and slag when they reach the bottom of the furnace.
g. Why do you think that the furnace used in the extraction of iron is called a blast furnace?
64
 3-22 Demonstration: Corrosion of Iron
Aim: To investigate a variety of factors involved in the corrosion of iron.
Introduction:
Corrosion is a complex reaction in which metals are converted to their oxides or other compounds. Corrosion
reactions involve the exchange of electrons as the metal sheds some valence electrons and becomes an ion.
The corrosion of iron is of most interest in modern developed societies. Rusting of iron structures costs billions
of dollars annually and attempts are sought to find methods to economically reduce it. Other metals exist which
corrode much slower than iron, but iron’s properties of strength, abundance, low mining and production costs
makes it difficult to find a suitable substitute. The chemical reactions involved in the corrosion of iron are not
completely understood. It is known, however, that iron metal is oxidised to iron (Il) ions, Fe2+ and hydroxide
ions, OH- , produced by the action of water and oxygen gas on the iron surface. One possible equation for the
overall reaction is shown below.
2Fe + O2 + 2H2O
2Fe2+ + 4OH-
Equipment
Test tubes/rack
Petri Dishes
Iron nails
0.1M iron (II) Sulphate (FeSO4)
Zinc strip (Zn)
Tin strip (Sn)
Steel Wool
Litmus paper
Phenolphthalein solution
0.1M potassium ferricyanide (K3Fe(CN)6)
Copper Wire (Cu)
Aluminium strip (Al)
Preparation: Write down the formulas of the following compounds.
Iron hydroxide
Hydrochloric acid
Copper (II) carbonate
Sodium carbonate
Copper (II) oxide
Sulfuric acid
Zinc
Iron sulfate
Sodium hydroxide
Sodium chloride
!
Safety
Goggles and aprons are to be used during this
experiment. If chemicals are splashed on your skin,
wash under a tap. Clean up any spills.
Procedure Day 1: Experiment Setup
Part A: Reaction of iron with aqueous solutions
1. Use steel wool to scratch clean four iron nails. Place each nail in a separate, small test tube, sliding the nail in
carefully to avoid breaking the tube.
2. Obtain one of the three sets of chemicals as assigned by your teacher. Label each of four test tubes with the
chemical name and fill each with a solution so that each nail is just covered.
3. Use litmus to determine whether each solution is acidic, basic, or neutral. Record the results of these tests in
the table.
4. Allow the test tubes to stand overnight. Continue with observations at the next class.
65
Part B: Effects of Stress / Protection by other metals
5. Use steel wool to scratch clean iron nails - five straight and one bent. Place two nails in the same petri dish,
beside, but not touching, one another.
6. To the other four nails, as listed below, twist a metal strip around the nail. Make sure there is tight contact
with the nail and metal. Place these nails also in pairs in a petri dish, beside, but not touching, one another.
7. Pour the assigned warm agar-suspension into the petri dishes to a depth of at least 2 mm to cover the nails,
being careful not to disturb them. The nails must not touch each other.
Set 1: Agar solution – no additives
Straight nail
Bent nail
Al-wrapped nail
Cu-wrapped nail
Zn-wrapped nail
Sn-wrapped nail
Set 2: Agar solution – phenolphthalein additive
Straight nail
Bent nail
Al-wrapped nail
Cu-wrapped nail
Zn-wrapped nail
Sn-wrapped nail
Set 3: Agar solution – Potassium ferricyanide additive
Straight nail
Bent nail
Al-wrapped nail
Cu-wrapped nail
Zn-wrapped nail
Sn-wrapped nail
8. The tests will be left overnight and observed during the next class for corrosion effects.
Procedure Day 2: Observations of corrosion
Part A: Reaction of iron with aqueous solutions
1. Test for the presence of iron (II) ions, Fe2+. Pour 1 mL (1 cm depth) of 0.1M iron (lI) sulfate in a test tube.
Add one drop of potassium ferricyanide. Record your observations.
2. Test for the presence of hydroxide ions, OH-. To another test tube, add 1 mL of 0.1M sodium hydroxide
solution. Add one drop of phenolphthalein. Record your observations in the table.
3. Test each solution pH, using universal indicator paper. Record results in the table.
4. Observe the nails placed in test tubes with the different tests, being careful not to disturb them. View all the
class agar and test tubes and record your observations of any changes in the table below. Be detailed, noting
the colours, their intensity, and position around each nail. Estimate the amount of corrosion on a scale of 0 –
10. (Where 10 is extreme and zero is no corrosion.)
5. Your teacher will test each of the test tubes containing the nails for the presence of ferrous ions by adding 1
or 2 drops of potassium ferricyanide to each tube. Record your observations. The presence of ferrous ions in
the test tubes is evidence that corrosion has occurred.
66
Test
pH
Initial Observations
Observations After 24 Hours
Set 1
NaCl
Distilled
water and air
NaOH
Sn wrap/H2O
Set 2
KOH
Boiled water,
oil layer
HCl
Cu wrap/H2O
Set 3
FeSO4
H2SO4
Al wrap/H2O
Zn wrap/H2O
67
Part B: Effects of Stress: Protection by other metals
6. Draw coloured diagrams of each petri dish in the diagrams below to identify areas where corrosion occurred.
Agar gel
68
Potassium ferricyanide
Phenolphthalein
7. List the chemicals used in Part A for which there was no evidence of corrosion.
8. List the chemicals used in Part A for which there was evidence of corrosion.
9. What must be present for corrosion of iron to take place?
10. Which metals appeared to protect the iron nail against corrosion? Discuss in detail.
11. Which metals promote the corrosion rate?
12. Explain how the colours developed in the petri dishes and identify the products of corrosion for each
reaction.
13. Are there any solutions that seem to inhibit corrosion? Explain these effects.
14. Are there any solutions that seem to promote corrosion? Explain these effects.
15. What effect does bending seem to have on the tendency of iron to corrode?
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 3-23 Summary Minerals to Metals Questions
1. Only a few elements are found uncombined in the Earth's crust. Gold is one example. The rest occur as
compounds, and have to be extracted from their ores. This is usually carried out by heating with carbon, or by
electrolysis. Some information about the extraction of three different metals is shown below.
Metal
Iron
Aluminium
Sodium
Formula of main ore
Fe2O3
Al2O3
NaCl
Method of extraction
Heating with carbon
Electrolysis
Electrolysis
a. Give the chemical name of each ore.
b. Arrange the three metals in order of reactivity.
c. How are the more reactive metals extracted from their ores?
d. i. How is the least reactive metal extracted from its ore?
ii. Why is this a reduction reaction?
iii. Why can't this method be used for the more reactive metals?
e. Which of the methods would you use to extract: i. potassium? ii. lead? iii. magnesium?
f. Gold is a metal found native in the Earth's crust. Explain what native means.
2. Explain why the following metals are suitable for the given uses. (There should be more than one reason in
each case.)
a. aluminium for window frames b. iron for bridges
c. copper for electrical wiring
d. lead for roofing
e. zinc for coating steel
3. Many metals are more useful when mixed with other elements than when they are pure.
a What name is given to the mixtures?
b. i. What metals are found in these mixtures? brass
solder
stainless steel
ii Describe the useful properties of the mixtures.
iii. Give one use for each of these mixtures.
c. Name another mixture of metals and the useful properties it has.
4. a. Draw a diagram of the blast furnace. Show clearly on your diagram:
i. where air is ‘blasted’ into the furnace.
ii. where the molten iron is removed.
iii. where the second liquid is removed.
b. i. Name the three raw materials added at the top of the furnace.
ii. What is the purpose of each material?
c. i. What is the name of the second liquid that is removed from the bottom of the furnace?
ii. When it solidifies, does it have any uses? If so, name one.
d. i. Name a waste gas that comes out at the top of the furnace.
ii. Does this gas have a use? If so, what?
e. Write an equation for the chemical reaction that produces the iron.
5. Gold mines in South Africa can be up to 4km deep. These mines are very expensive and dangerous. No
other metal is mined under these conditions. Can you suggest a reason for this?
70
1 tonne (1000 kg) of
rock + copper ore
10 kg of concentrated
copper ore
Step 1

Step 2

2.5 kg of 99%
pure copper
Almost 2.5 kg of
99.9% pure copper
Step 3

5.
The diagram above shows stages in obtaining copper from a low-grade ore. This ore contains copper(II)
sulphide, CuS. It may also contain small amounts of silver, gold, platinum, iron, zinc, cadmium, and arsenic.
a What is an ore?
b. What is a low-grade ore?
c. What name is given to the waste rock in an ore?
d. i. What process is used in Step 1 to concentrate the sulphide ore? Explain how it works.
ii The waste material from this process is in the form of mud. Explain why.
e. i. What process is carried out in step 2, to extract the copper from the ore?
ii. Write an equation for the reaction.
iii. What is this type of reaction called?
iv. The copper is 99% pure. Suggest some impurities it may contain.
f. i. What process is carried out at step 3 to purify the metal?
ii. What will the main cost in this process be?
iii. As well as pure copper, this process may produce other valuable substances. Explain.
g. List all the environmental problems that may arise in going from A to D.
6. Metal extraction should be carried out in the most economical way, yet calcium and magnesium are both
extracted from their ores by the expensive process of electrolysis. Suggest why this method is used.
7. In the Middle Ages, alchemists found that if they heated cinnabar (which is an ore of mercury) they obtained
silvery globules of mercury. What does this suggest about the reactivity of mercury metal?
8. Chromium metal can be extracted from its ore by a displacement reaction using aluminium:
chromium oxide + aluminium  chromium + aluminium oxide
What does this reaction tell you about the relative reactivities of chromium and aluminium?
Suggest why carbon is not used to extract chromium.
9. Metal
Date of Discovery
Iron
Ancient
Copper
Ancient
Zinc
Ancient
Aluminium
1826
Does there seem to be any link between date of
discovery and method of extraction?
10. A solution of zinc bromide is electrolysed as shown
in the diagram.
a. Label the positive and negative electrodes.
b. Which electrodes will the zinc and bromide ions
migrate to?
c. Identify the reactions occurring at each electrode.
71
4-1 Chemical Data
Positive and Negative Ions – Valency Table
+1
+2
Lithium
Li+
+3
Magnesium
Mg2+
Sodium
Na+
Potassium
K+
Silver
Ag+
Copper (I)
Ammonium
Cu+
NH4+
Hydrogen
H+
Al3+
Tin (IV)
Sn4+
Calcium
Ca2+
Chromium (III)
Cr3+
Lead (IV)
Pb4+
Barium
Ba2+
Iron (III)
Fe3+
Iron (II)
Fe2+
Copper (II)
Zinc
Cu2+
Zn2+
Tin (II)
Sn2+
Lead (II)
Pb2+
Mercury
Hg2+
Manganese (II)
Mn2+
Nickel
Ni2+
-1
-2
-3
Fluoride
F-
Oxide
O2-
Chloride
Cl-
Sulfide
S2-
Bromide
Br-
Sulfate
SO42-
Iodide
I-
Carbonate
CO32-
Hydroxide
OH-
Nitrate
NO3-
Permanganate
MnO4-
Hydrogen sulfate
HSO4-
Hydrogen carbonate
HCO3-
-
Important chemical formulas you should know
Formula
72
Common Name
CO2
Carbon dioxide
HCl
Hydrochloric acid
H2SO4
Sulfuric acid
HNO3
Nitric acid
NaOH
Sodium hydroxide
NH3
Ammonia
O2
Oxygen
N2
Nitrogen
H2
Hydrogen
Cl2
Chlorine
F2
Fluorine
I2
Iodine
+4
Aluminium
Phosphate
PO43-
Andromeda Galaxy
Supernova remnant W49B
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Firefox nebula