HBSC 2103

Table of Contents
Course Guide xiăxv
Topic 1 The Particulate Nature of Matter 1
1.1 What is Matter? 2
1.2 States of Matter 3
1.3 Properties of Matter in Solid, Liquid and Gaseous States 4
1.3.1 Solids 5
1.3.2 Liquids 6
1.3.3 Gases 8
1.4 States of Matter and Kinetic Particle Theory 10
1.5 Changes of State of Matter 13
1.5.1 Melting 13
1.5.2 Freezing 13
1.5.3 Evaporation 14
1.5.4 Condensation 14
1.5.5 Sublimation 14
1.6 Cloud and Rain Formation 15
1.7 Practical Investigation in the Primary Science Curriculum 18
Summary 19
Key Terms 20
References 21
Topic 2 Atoms, Elements and Compounds 22
2.1 Nature of Atoms, Elements, Compounds and Mixtures 23
2.1.1 Pure Substances 24
2.1.2 Mixtures 28
2.1.3 Separating Components of a Mixture 29
2.2 Differences between Metals and Non-metals 36
2.3 Chemical Symbols of Elements 38
2.4 Formulae of Molecules for Elements and Compounds 40
2.5 Similarities and Differences of Elements, Compounds and
Mixtures 42
2.6 Alloys 44

iv  TABLE OF CONTENTS
2.7 Solutions 47
2.7.1 Solution, Solute and Solvent 48
2.7.2 Saturated Solution 51
2.7.3 Some Factors Affecting Solubility 53
Summary 55
Key Terms 56
References 56
Topic 3 Atomic Structure 58
3.1 History of Atomic Model 59
3.2 Subatomic Particles 60
3.3 Atomic Number, Nucleon Number and Mass Number 61
3.4 Isotopes 64
3.5 Electronic Configuration of Atoms 67
3.6 Valence Electrons 70
Summary 71
Key Terms 72
References 73
Topic 4 Periodic Table 74
4.1 History of the Periodic Table 75
4.1.1 Antoine Lavoisier (1743ă1794) 75
4.1.2 Johann Dobereiner (1780ă1849) 76
4.1.3 John Newlands (1837ă1898) 77
4.1.4 Lothar Meyer (1830ă1895) 78
4.1.5 Dmitri Mendeleev (1834ă1907) 79
4.1.6 Henry JG Moseley (1887ă1915) 80
4.1.7 Modern Periodic Table 81
4.2 Analysis of the Periodic Table 83
4.2.1 Groups 83
4.2.2 Periods 84
4.3 Electronic Structures and the Periodic Table 86
4.4 Transition Elements 89
4.4.1 Properties 89
4.4.2 Industrial Uses 92
4.5 Group 1 93
4.5.1 Electronic Structure 93
4.5.2 Group Trends 94
4.5.3 Physical Properties 96
4.5.4 Chemical Properties 96

TABLE OF CONTENTS  v
4.6 Group 17 97
4.6.1 Electronic Structure 97
4.6.2 Group Trends 98
4.6.3 Physical Properties 100
4.6.4 Chemical Properties 100
4.7 Noble Gases 101
4.8 Period 3 103
4.8.1 Chloride and Hydride for Elements in Period 3 104
Summary 106
Key Terms 109
References 109
Topic 5 Chemical Bonding 110
5.1 The Stability of Noble Gas Structure 111
5.2 The Octet Rule 112
5.3 The Formation of Ions 113
5.4 Ionic Bonds 116
5.4.1 Formation of Ionic Bonds in Magnesium Chloride 116
5.4.2 Dot-and-cross Diagram (Lewis Diagram) 118
5.4.3 Chemical Formulae of Ionic Compounds 120
5.4.4 Structure and Properties of Ionic Compounds 121
5.5 Covalent Bonds 122
5.5.1 Molecules of Elements 123
5.5.2 Molecules of Compounds 125
5.5.3 Structure and Properties of Covalent Compounds 128
5.5.4 Giant Molecular Compounds 129
5.6 Metallic Bonds 133
5.6.1 Strength of Metallic Bonds 133
5.6.2 Conductors and Insulators 134
5.7 Intermolecular Forces 134
5.7.1 Van der Waal Forces 135
Summary 138
Key Terms 139
References 140
Topic 6 Chemical Calculations 141
6.1 Writing and Balancing Chemical Equations 142
6.2 Relative Atomic Mass and Relative Molecular Mass 146
6.3 The Mole Concept 149
6.3.1 The Mole and Avogadro Constant 149
6.3.2 Moles of Gases 154
6.3.3 Moles and Solutions 154

vi  TABLE OF CONTENTS
6.4 Empirical Formula and Molecular Formula 155
6.4.1 Empirical Formula of Compound 158
6.4.2 Molecular Formula of Compound 158
6.4.3 Calculating the Formula of a Compound 159
Summary 163
Key Terms 164
References 164
Topic 7 Acids and Bases 165
7.1 Acid, Base and Alkali 166
7.1.1 Arrhenius Theory of Acids and Bases 166
7.1.2 Bronsted-Lowry Theory of Acids and Bases 167
7.1.3 Alkali 169
7.1.4 Hydrated Protons and Hydronium Ions 169
7.1.5 Acid Strength and Base Strength 170
7.1.6 Autoionisation of Water 171
7.2 pH Scale 172
7.3 Characteristic Properties of Acids, Bases and Alkalis 174
7.3.1 Properties of Acids 174
7.3.2 Properties of Bases and Alkalis 175
7.4 Tests for Acids and Alkalis 177
7.5 Concentration of Acid and Base 181
7.5.1 Relationship between the Number of Moles and
Molarity 181
7.5.2 Preparation of a Standard Solution 183
7.5.3 Solution Liquefaction 185
Summary 187
Key Terms 188
References 189
Topic 8 Salts 190
8.1 Preparation of salts 191
8.1.1 Preparation of Soluble Salts 192
8.1.2 Preparation of Insoluble Salts 198
8.2 Identification of Cations and Anions 200
8.2.1 Identification of Anions 202
8.2.2 Identification of Cations 202
8.3 Identification of Gases 204
8.4 Crystallisation 206

TABLE OF CONTENTS  vii
8.5 Qualitative Analysis of Salts 208
8.5.1 Preliminary Examination of the Salt 209
8.5.2 Identification of Anions and Cations in the Salt 214
8.5.3 Confirmatory Tests for Specific Anions and Cations 214
Summary 216
Key Terms 217
References 218

Topic
1
The Particulate
Nature of
Matter
LEARNING OUTCOMES
By the end of this topic, you should be able to:
1. Describe the concept of matter;
2. Identify the properties of the three states of matter;
3. Describe the states of matter and the kinetic particle theory;
4. Discuss the changes of states using appropriate examples and
activities;
5. Explain the formation of cloud and rain; and
6. Conduct appropriate investigations to explain the changes of states
of matter.
INTRODUCTION
Have you ever wondered what matter is? Why is ice a solid, water a liquid and
steam a gas? Why does ice melt and water evaporate? How do clouds and rain
form? We will find out the answers to these questions in the following
discussions. We will examine matter, states of matter and their properties, the
Kinetic Particle Theory, changes of states, formation of clouds and rain through
practical investigations and activities in the Primary Science Curriculum.

2 TOPIC 1 THE PARTICULATE NATURE OF MATTER
WHAT IS MATTER?
1.1
Let us start this topic by learning the meaning of matter first. Do you know that all
things around us are matter? In fact, we can say there is matter everywhere. For
instance, there is matter in your hair, the air you breathe and the water you drink.
There is also matter in the clothes you are wearing. Matter can exist as solid, liquid
or gas. So, how do you define matter?
Matter can be defined as anything
that has mass and occupies space.
Thus, even you yourself are matter because you have mass and occupy space. But
bear in mind that you must not confuse matter with weight. This is because mass is
a measure of amount in a given sample. In other words, it can be said that mass is
a measure of quantity of material in a given object. As for weight, it is a measure of
the gravitational pull of an object on earth.
Note that air is another example of matter. We can show that air has mass and
occupies space through the following activities:
(a) Blow up two balloons, A and B, to about the same size. Put a
piece of sticky tape on balloon B. Then, balance the two
balloons. Gently push a pin through the sticky tape and then
pull it out.
(b) Fill a basin with water. Hold a cup upside down. Push it into the
water.
What did you observe? What can you say about air in these two activities? Now
you should be convinced with the fact that air has mass and occupies space, or in
other words, air is also matter.

TOPIC 1 THE PARTICULATE NATURE OF MATTER 3
STATES OF MATTER
As mentioned before, matter normally exists in one of these three physical states
solid, liquid or gas. Can you describe these three states of matter?
Solids are firm and have a definite form. Wood, glass, iron nails, cotton and
paper are all examples of solids.
Liquids, on the other hand, are not rigid. If liquid is poured onto a table, it will
flow all over the surface. Examples of liquids are water, milk and oil.
Lastly, let us talk about gas. Gases can be found everywhere around us.
However, they are not visible or cannot be seen by the naked eye. Can you think
of the examples of gases? Some examples of gases are oxygen, hydrogen,
nitrogen and carbon dioxide. Thus, gases have no fixed shape and spread out to
fill any container.
Based on the above discussion, we can say that everything we know is made of
matter whether it is in solid, liquid or gaseous form. In fact, a single object can be
in three different physical states. One good example is water.
Water can exist as a solid, liquid or gas at different temperatures. At 20C, water
exists in the form of a solid, which is ice. At 30C, water is a liquid. At 120C,
water exists as a gas in the form of steam. Figure 1.1 shows you the three states of
water.
Ice (solid) Water (liquid) Steam (gas)
Figure 1.1: Water in three different physical states
Source: http://images.google.com
1.2

ACTIVITY 1.1
Look around you. What are five examples of solid, liquid and gas?
Name them in the table below.
Examples of Solids, Liquids and Gases
No. Solids No. Liquids No. Gases
1. 1. 1.
2. 2. 2.
3. 3. 3.
4. 4. 4.
5. 5. 5.
PROPERTIES OF MATTER IN SOLID,
LIQUID AND GASEOUS STATES
1.3
Can you recall the three different states of matter discussed earlier? The three
states are solid, liquid and gas, and they each have different physical properties.
What can we say about physical properties?
Physical properties are characteristics that do not
change the identity and composition of the
substance. The physical properties include colour,
odour, density, melting point, boiling point, and
hardness.
Let us study these three states further in the next sections starting with solids.

1.3.1 Solids
As we can see in Figure 1.1 (the water example), solid has a fixed shape. It is hard
and the shape cannot be changed easily. Hence, solids cannot be compressed. It
exhibits a regular arrangement of particles and it is rigid. Many of them have a
definite three-dimensional shape with surfaces at specific angles to each other.
For example, table salt (sodium chloride) at room temperature as well as sugar
have cubic shapes with faces at 90. Figures 1.2 (a) and (b) show you the shape of
solid salt and sugar.
Figure 1.2(a): Shape of table salt
Source: Burns (1992), p. 22
Figure 1.2(b): Shape of sugar crystals
Source: www.encarta.msn.com/media
Do you know that solids have definite size, mass and weight at a given
temperature? For instance, a piece of iron nail can be of different sizes; two
centimetres long, or five centimetres long. When weighed on a beam balance, the
mass and weight of these two sizes of iron nails are also different. However,
these iron nails can be resized or reshaped under certain conditions and
temperature. Furthermore, solids also have a fixed volume. How do we measure
it? The volume of a solid can be measured using a measuring cylinder as shown
in Figure 1.3.

Figure 1.3: Measuring volume of solid using a measuring cylinder
Based on Figure 1.3, we can see that when a piece of marble is put into a
measuring cylinder containing 22.4 cm3 of water, the level of the water will rise
to 26.6 cm3. The difference between the first water level, before the marble is put
into the cylinder, and the water level after the marble is put into the cylinder, is
the total volume of the marble, that is, 4.2 cm3.
Lastly, solids do not flow easily. For instance, when a solid is placed into a
container, it cannot completely fill up the container. Instead, there will be spaces
in between the solid and the container.
1.3.2 Liquids
Now let us proceed to learn more about liquids. Unlike solids, liquids do not
have specific shapes of their own. Liquids take the shape of the container they
are in. For example, if we pour water into a glass or a container as shown in
Figure 1.4, it will take the shape of that container.

Figure 1.4: Water in the glass
Likewise, if we pour apple juice or any other liquid into a bottle or a paper cup, it
will take the shape of that particular container. In conclusion, we can say that no
matter how you change the shape of the container, the liquid will take the shape
of that particular container as shown in Figure 1.5.

Figure 1.5: Picture of different shapes of containers filled with liquids
Furthermore, just like solids, liquids too have a definite mass and volume. A litre
of liquid will not expand to fill a two-litre container. However, liquids are not as
hard as solids. They cannot be compressed to fill any sizes of containers. A liquid
can flow when it is poured.

For instance, if we pour water into a glass, we can see the water flowing. When it
rains, we can see droplets of water as shown in Figure 1.6a. When a droplet of
water drops into a water source, it causes a ripple as shown in Figure 1.6b.

Figure 1.6a: Droplets of rain Figure 1.6b: A droplet of water causing
a ripple in the water
ACTIVITY 1.2
Mercury is a liquid and is used to measure temperature in
thermometers. Discuss the properties of mercury that make it suitable to
be used in themometers.
1.3.3 Gases
Lastly, let us learn the properties of gases. Gases have definite mass but no
definite shape of their own and volume. They completely fill the containers they
are in. However, gases can flow easily and compressed into different types or
sizes of containers as shown in Figure 1.7.

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Figure 1.7: Hot air balloons, gas canister, air compressor and gas cylinders

As a conclusion, we can summarise the different physical properties of solids,
liquids and gases as shown in Table 1.1.
Table 1.1: Comparison of the Physical Properties of Solids, Liquids and Gases
Physical Properties Solids Liquids Gases
Mass Definite Definite Definite
Volume Definite Definite Fill the container
they are in
Shape Definite Take the shape of
the container they
are in
Take the shape of
the container they
are in
Ability to flow Unable to flow Flows Flows easily

STATES OF MATTER AND KINETIC
PARTICLE THEORY
1.4
Let us now learn the history of matter. More than 2,000 years ago, a Greek
philosopher called Democritus suggested this hypothesis:

All matter, living and non-living, is made of tiny particles too small to be seen.
His idea was that if you keep cutting something into smaller and smaller pieces,
you will eventually come to the smallest particles, which are the building blocks
of matter. He used the word amos (which in Greek means „cannot be divided‰)
to describe the smallest particles. This is where the word „atom‰ comes from. In
addition, these particles are always in constant motion (you will learn more
about atoms in Topic 2).
Since then, scientists have done many tests with matter, and the results have
always agreed with DemocritusÊ hypothesis. Such a hypothesis that is supported
by many experimental results is called a theory. So the hypothesis that matter is
made up of tiny particles too small to be seen which are in constant motion is
now called the kinetic particle theory of matter.
What does the kinetic particle theory of matter state?
When do we use this theory? The kinetic particle theory can be used to explain
and differentiate the properties of the three states of mattersolids, liquids and
gases as shown in Table 1.2.

Table 1.2: Three States of Matter Solids, Liquids and Gases
States of Matter/
Aspect Solid Liquid Gas
Arrangement of
Particles
Particles are usually
arranged in a regular
pattern; they are
closely packed and
are located at fixed
positions. The closer
the molecules are, the
harder they will be.
Particles are
arranged close
together but not
tightly and orderly
in a fixed pattern.
Particles are arranged
randomly. There is no
orderly pattern.
Forces of
Attraction
between
Particles
There are strong
forces called chemical
bonds holding the
particles in fixed
positions.
There are strong
forces holding the
particles together
but not enough to
hold them in fixed
positions.
There are attractive
forces between the
particles but they are
very weak. These
forces are insufficient
to hold gas particles
together to form any
definite shape or
volume. Thus, the
particles can move
freely to fill the
container. A litre of
gas can expand to fill
a two or more litre
container.
Motion of the
Particles
Since the particles are
in a solid form, they
are arranged and
packed tightly; there
is little motion of the
particles. The only
movements are tiny
vibrations to and
from a fixed position.
The particles vibrate
faster when they are
heated.
Since the particles
are not tightly
packed, they are
able to move about
randomly
throughout the
liquid. Particles
move faster when
they are heated.
Since the particles are
very far apart, they
move quickly, freely
and randomly in all
directions. Particles
collide with each
other and also with
the walls of the
container, and bounce
off in all directions.
Particles move faster
when they are heated.
Kinetic Energy Low Moderate High

Do you know that the kinetic particle theory can also be used to explain the
process of diffusion? What is diffusion? Let us examine this situation: if someone
opens a bottle of perfume in front of the classroom, you will soon smell the
perfume in other parts of the room. The fragrance spreads through the air in all
directions. The gradual mixing of substances is called diffusion.
Based on the previous situation, let us explain how perfume diffuses. When the
lid is on, the gas particles remain inside the bottle. When the lid is taken off, the
liquid perfume evaporates easily. Since there are only weak forces between the
particles, they can spread out, moving away from the crowded bottle to places
where there are fewer particles of perfume. Eventually, the particles spread
evenly throughout the air in the room. The process of diffusion is shown in
Figure 1.8.
Figure 1.8: The process of diffusion
ACTIVITY 1.3
Put a Petri dish on a sheet of white paper and half fill it with water.
Let it stand for a while to let the water become perfectly still. Use a
pair of tweezers or a spatula to place a single crystal of potassium
permanganate in the centre of the dish. Then, leave the dish
undisturbed for five minutes. What can you see? Explain what you
have observed in terms of particles.

CHANGES OF STATE OF MATTER
1.5
Now, let us move on to learn about changes of states of matter. We start off by
learning the basic concept, a substance can be changed from one state into
another when it is heated or cooled. The changes in the state of substance can be
explained using the kinetic theory model discussed earlier.
Each change of state involves a physical process and change in energy of the
substance. These physical processes are explained in the following.
1.5.1 Melting
Let us conduct a demonstration where an ice cube is dropped into a cup of hot
water to show how it melts. What will happen? Yes, the ice will melt, which
means that when it happens, the ice has changed its form from solid to liquid.
This physical process is known as melting. Based on the demonstration, can you
define melting? Melting is a process where solid changes to its liquid state when
heated at a certain temperature and pressure.
During the melting process, heat energy is absorbed. Thus, forces and bonds are
broken during this process too.
1.5.2 Freezing
What is freezing? Freezing is a reverse process of melting. You can reverse
melting simply by putting or placing water in a freezer. Water, which is in liquid
form, will change to ice, a solid form. This physical process is known as freezing.
You can use the kinetic theory of matter to explain the changes from liquid to
solid due to cooling. Therefore, freezing is the process where liquid changes to its
solid state when it is cooled at a certain temperature and pressure. Conversely,
during freezing, heat energy is released and not absorbed.

1.5.3 Evaporation
The next changing state of matter is evaporation. How do you explain
evaporation? What happens when a small bowl of water is left out overnight?
Yes, the water in the bowl would have dried up the next day. Water, which is in
liquid form, will change to vapour, a gas form. This physical process is known as
evaporation.
Once again, you can use the kinetic theory of matter to explain the changes from
liquid to gas due to heating at room temperature. Evaporation is the process
where atoms or molecules, which are in liquid state, gain sufficient energy to
enter the gaseous state. During evaporation, heat energy is absorbed.
1.5.4 Condensation
Now, we move on to condensation. This is how you can demonstrate it: cover a
cup of hot water with a saucer for a few minutes, then observe what is on the
surface of the saucer when the saucer is removed from the cup. You will notice
that there are drops of water on the surface of the saucer. Why this process
happened? Hot water in vapour form condenses into droplets of water in liquid
form when it is cooled. This defines condensationit is the process where gas or
vapour changes to its liquid state when cooled at certain temperature and
pressure. During condensation, heat energy is released.
1.5.5 Sublimation
For sublimation, we can use a sample of dry ice (solid carbon dioxide) and then
touch it (you have to be very careful with it as dry ice may cause frost bite,
damaging the skin very much in the manner of a burn). We will notice traces of
vapour released from the surface of the dry ice but strangely it does not melt.
After a while, the size of dry ice decreases gradually. Why? Dry ice does not melt
but changes from solid state directly to gaseous state. This physical process is
called sublimation.

Sublimation is the process where solid changes to vapour or gas, without going
through the melting process. During sublimation, heat energy is absorbed. Can
you think of other examples? Other examples of substances that undergo
sublimation are iodine and ammonium chloride.
As a conclusion, the five physical processes involved in the change of state are
summarised in Figure 1.9.
Figure 1.9: Changes in states of matter
CLOUD AND RAIN FORMATION
1.6
When we look up at the sky, we may see different types of clouds. When the
clouds look dark, it is a sign that it is going to rain. What are these clouds made
up of? How do they form? The link to these questions is water.
How do we link water to clouds? Well, water can evaporate from plants, animals,
puddles, soil and other ground surfaces, and from oceans, lakes, rivers and
streams to form clouds as shown in Figure 1.10.

Figure 1.10: Formation of clouds

The formation of clouds involved condensation and evaporation. Condensation
occurs when water vapour (gas) in the air changes into liquid due to cooling.
These water droplets are formed when water vapour condenses around a
condensation centre a tiny particle of smoke, dust, ash or salt. Visible clouds are
tiny water droplets suspended in the air. Clouds form under certain conditions,
such as when more water vapour evaporates from the earth into the atmosphere
than condenses on the earth, and when there are dusts, smoke or other particles
suspended in the air, water vapour condenses onto these particles in the air.
Clouds float in the air and are moved by the wind. Note that there are different
types of clouds and not all clouds produce rain.
How can you tell that it is going to rain? There are a few signs that tell you that
rain is imminent (Figure 1.11). However, this cannot always be true as clouds can
always be moved by the wind.

Sometimes, there is lightning
before rain
Heavy dark clouds Yes ...it is raining!
Figure 1.11: Signs just before raining
What can you say about rain? Rain is liquid water that falls from clouds. Rain
occurs when the water droplets in a cloud get too heavy to stay suspended in the
sky and so fall due to gravity. In a super-cooled atmosphere, water droplets and
ice crystals in a cloud interact to produce more ice crystals. However, these
crystals from the cloud will melt as they fall. Otherwise, hail can happen. But this
rarely occurs. Have you ever encountered one?
ACTIVITY 1.4
You can try out this activity to make rain. Pour some hot water into a
clear plastic jar. Cover the top of the jar with a plastic sheet. Press the
centre of the plastic sheet down so that it forms a funnel shape. Put
some ice-cubes onto the plastic sheet. Observe what happens inside the
tube. Explain how it happened.

PRACTICAL INVESTIGATION IN THE
PRIMARY SCIENCE CURRICULUM
Lastly, let us look at two practical experiments that you can conduct to explain
the changes of states of matter (water). These two experiments are to investigate
the boiling point of water and its evaporation. Just follow these steps and
instructions.
Experiment 1: The Boiling Point of Water
Materials Required:
1. Round-bottomed flask
2. Thermometer
3. Retort stand
4. Wire gauge
5. Bunsen burner
6. Lighter
Procedures:
1. Prepare the apparatus as shown in Figure 1.12.
Figure 1.12: Boiling point of water
1.7

2. Take about 100 ml of distilled water. Fill it into the round-bottomed flask.
3. Set up the thermometer as shown in Figure 1.12.
4. Heat the flask and record the temperature for every five minutes until the
water is boiling.
5. Record the temperature when it boils. This is the boiling point of the water
(the temperature would remain a constant at this time).
6. Repeat the same experiment using different liquids (that are safe for this
experiment) with different boiling points.
Experiment 2: Evaporation
Prepare two Petri dishes. Pour about 10 ml (two teaspoons) of water in each of
the dish. Place one dish in the sunlight. If the sun is not shining strongly enough,
place the dish closely to a source of light. Place the other dish in the shade.
Observe each dish every four hours, then overnight. Record what happens to the
water.
Answer these questions.
(a) Where did the water go?
(b) From which dish did the water disappear faster?
(c) What caused the water to disappear?
The process of water „going‰ into the air is evaporation. List some other
examples of evaporation. What happens to the water after it evaporates?
Matter is anything that has mass and takes up space.
Matter can be classified as solids, liquids or gases.
Solids have fixed shape, fixed volume, are hard and cannot be compressed.
Liquids have no fixed shape but take the shape of the container they are in,
have fixed volume, are not hard and can flow.

Gases have no fixed shape and volume, and take the shape of the container,
flow easily and can easily be compressed.
The Kinetic Particle Theory of Matter states that matter is made up of tiny
particles too small to be seen which are in constant motion.
The Kinetic Particle Theory can be used to explain the properties of solids,
liquids and gases.
Changes of states can be demonstrated using appropriate examples and
activities such as using ice cube for melting, putting water into freezer for
freezing, and so on.
Visible clouds are tiny water droplets suspended in the air and are formed
when water vapour condenses around a condensation centre a tiny particle
of smoke, dust, ash or salt.
Rain occurs when the water droplets in a cloud get too heavy to stay
suspended in the sky and so fall due to gravity.
Appropriate experiments can be conducted to explain the changes of states of
matter such as boiling and evaporating the water.
Cloud
Condensation
Evaporation
Freezing
Gas
Kinetic particle theory
Liquid
Matter
Melting
Rain
Solid
Sublimation

Brady, J. E., Senese, F. (2004). Chemistry: Matter and its changes (4th ed.).
New York: John Wiley Sons, Inc.
Kots, J. C., Treichel, P. M., Weaver, G. C. (2006). Chemistry: The chemical
reactivity (2nd ed.). Victoria, Australia: Thomson Learning.
Timberlake, K. C. (2006). An introduction to general, organic, and biological
chemistry (9th ed.). San Francisco, CA: Pearson-Benjamin Cummings.
Sumdahl, S. S., Sumdahl, S. A. (2003). Introductory chemistry: A foundation
(6th ed.). Boston, NY: Houghton Mifflins Co.

Topic
2
Atoms,
Elements and
Compounds
LEARNING OUTCOMES
1. Identify how matter is classified based on its atoms;
2. Differentiate between metals and non-metals;
3. Write the chemical symbols for elements;
4 Write the formula of molecules for elements and compounds;
5. Identify similarities and differences between elements, compounds
and mixtures;
6. Describe alloys; and
7. Differentiate between solution, solute and solvent.
INTRODUCTION
In this new topic, you will be introduced to atoms, elements and compounds.
Let us recap what we have learnt so far; we learnt that the things that we see
around us are all matter. They occupy space and have mass. They also have
different forms and appearances. Some substances are gases, some are liquids
and some are solids; some are hard and shiny but others are soft and dull.
Different substances behave differently. For example, iron rusts but gold does
not, and copper conducts electricity while sulphur does not. How can these

TOPIC 2 ATOMS, ELEMENTS AND COMPOUNDS 23
observations be explained? What are these substances made up of that give
them different forms and appearances? Well, you will find the answers later as
we begin this topic by identifying how matter is classified based on its atoms,
and learning how to differentiate between metals and non-metals. Then, we
will learn how to write the chemical symbols for elements and the formula of
molecules for elements and compounds. Later, we will learn how to identify
similarities and differences between elements, compounds and mixtures
followed by an explanation of alloys. Lastly, we will look at solution, solute and
solvent, and learn how to prepare them. Are you ready? Let us start the lesson!
NATURE OF ATOMS, ELEMENTS,
COMPOUNDS AND MIXTURES
2.1
In Topic 1, we learned that matter is anything that has mass and occupies space.
Matter, whether it is living or non-living, is made up of atoms the almost tiny
and small building blocks of matter. The properties of matter relate not only to
the kinds of atoms it contains (composition) but also to the arrangement of these
atoms (structure). This helps us to classify and describe the many different kinds
of matter that can be found around us. All the many kinds of matter can be
classified and described in two ways: (i) according to its physical state as a gas,
liquid or solid (which has been discussed in Topic 1); and (ii) according to its
composition either pure substances or mixtures, as shown in Figure 2.1.
Figure 2.1: Classification of matter
Let us look at these classifications further, starting with pure substances.

24 TOPIC 2 ATOMS, ELEMENTS AND COMPOUNDS
2.1.1 Pure Substances
What does a pure substance stand for?
A pure substance (usually referred to simply as a substance) is
matter that has a fixed composition and distinct properties.
For example, water and ordinary table salt (sodium chloride), which make up the
primary components of seawater, are pure substances. Pure substances, in turn,
can either be elements or compounds. These two types of pure substances will be
discussed further in the following sections.
(a) Element
What is an element?
An element is the simplest substance with the following three features:
(i) It consists of only one type of atom;
(ii) It cannot be broken down into simpler substances either by physical
or chemical means; and
(iii) It can exist either as individual atoms or molecules as shown in
Figure 2.2.
An element is a substance which cannot be broken down
into simpler substances by chemical or physical methods.

Figure 2.2: Atoms and molecules of an element
Can you give some examples of individual atoms? Examples of elements
that consist of individual atoms include aluminium, zinc, iron, calcium and
gold. What is your definition of a molecule?
A molecule consists of two or more atoms of the same element,
or different elements, which are chemically bound together.
Can you give some examples of molecules? Examples of elements that
consist of molecules include oxygen, hydrogen and nitrogen. For example,
an oxygen molecule consists of two oxygen atoms whereas a hydrogen
molecule consists of two hydrogen atoms which are held together in
specific shapes.

Do you know that only about 90 of the 115 presently known elements occur
naturally? The remaining ones have been produced artificially by nuclear
chemists using high energy particle accelerators. These elements can be
further grouped into metals, non-metals and semi-metals as explained in
Table 2.1.
Table 2.1: Three Group of Elements
Group Description Example
Metal There are 90 types of metals. Potassium, mercury, lead,
magnesium, silver and sodium.
Non-metal There are 18 types of non-metals. Hydrogen, chlorine, bromine,
phosphorus, carbon and oxygen.
Semi-metals
There are seven semi-metals or
metalloids whose properties are
intermediate between metals and
non-metals.
Boron, silicon, germanium,
arsenic, antimony, tellurium and
astatine.
We will learn more about metals and non-metals in subtopic 2.3.
(b) Compound
Let us move on to learn about compound. Firstly, let us define it. Can you
give its definition?
A compound is a substance which consists of two or
more elements chemically combined together.

We can also say that a compound is a pure substance that is formed when
two or more different elements combine chemically; thus, it contains two or
more kinds of elements bonded together as shown in Figure 2.3.
Figure 2.3: Molecules of a compound
We can determine a compound by inspecting these three features:
(a) It can be broken down into a simpler type of matter (elements) by chemical
means (but not by physical means);
(b) It has properties that are different from its component elements; and
(c) It has a constant composition throughout and always contains the same
ratio of its component atoms.
Can you provide the examples of compounds? Some examples of compounds
include carbon dioxide and sodium chloride. When one atom of carbon combines
with two atoms of oxygen, carbon dioxide is formed. Such transformation is said
to be a chemical reaction. Similarly, when an atom of sodium combines with
an atom of chlorine, sodium chloride is formed. Sodium chloride has different
properties from sodium and chlorine.

2.1.2 Mixtures
Before we discuss further, let us first look at the definition of mixture. Can you
define it?
Mixtures are two or more substances that are mixed
together but not chemically joined.
Or to give a more detailed definition, we can say that mixtures are combinations
of two or more substances that are mixed physically in which each substance
retains its own chemical identity and hence its own properties, just like the one
shown in Figure 2.4.
Figure 2.4: Mixture of elements and compounds
While pure substances have fixed composition, the composition of mixtures
can vary. For example, a cup of sweetened tea can contain either a little sugar or
a lot. The substances making up a mixture such as sugar and water are called
components of the mixture.
Mixture can be further classified as either homogenous or heterogeneous
(refer Figure 2.1). What are the differences between them? Well, let us look at
homogenous mixture first.

(a) Homogeneous Mixture

A homogenous mixture is a mixture in which the mixing is uniform
and therefore has a constant composition throughout the mixture. The
components are indistinguishable.
For instance, the air is a homogenous mixture of the gaseous substances
such as nitrogen, oxygen and smaller amounts of other substances. The
nitrogen in the air has all the properties of pure nitrogen because both the
pure substance and the mixture contain the same nitrogen molecules. Salt,
water and many other substances dissolve in water to form homogenous
mixtures. Do you know that homogenous mixtures are also called
solutions? This will be further discussed in subtopic 2.7.
(b) Heterogenous Mixture
How about a heterogeneous mixture? What can you say about it?

A heterogeneous mixture is a mixture in which the mixing is not
uniform and therefore has regions of different compositions. The
components are distinguishable.
Can you think of some examples of heterogeneous mixtures? Some
examples are when you mix sand with sugar, water with petrol, dust with
air, and sulphur with iron filings.
2.1.3 Separating Components of a Mixture
Are you aware that a mixture can be separated into its components by physical
means? This is because each component of a mixture retains its own properties.
So, how do we do that? We can do that by using some of these methods:
Filtration;
Evaporation;
Distillation;
Fractional distillation;
Crystallisation; and
Chromatography.

(a) Filtration
Firstly, this method is suitable for an insoluble solid and liquid mixture as
Figure 2.5: Filtration process
Source: http://www.saskschools.ca
As you can see in Figure 2.5, the mixture is passed through a filter. The
residue is the substance that remains on the filter paper. The filtrate is the
substance that flows through the filter paper. A mixture of sand and water
is a good example of this method. Filtration yields sand as the residue and
water as the filtrate.
(b) Evaporation
The second method of separation is evaporation. Let us define this method
first.
Evaporation is a method of separating a solid that has
been dissolved in a solvent.

When do we use this method? This method is suitable for a soluble solid
and liquid mixture. Figure 2.6 shows you how to create evaporation. If a
mixture is heated or left over a few days, the solvent or liquid evaporates,
leaving the solid as residue.
Figure 2.6: Evaporation process
Source: http://www.allrefer.com
For example, a mixture of salt and water can be separated by evaporating
the water and leaving the solid salt behind.
(c) Distillation
The next method of separation is distillation. Let us define its meaning first.
Distillation is a process to separate a substance (in the
form of a solution) from its solvent.

When do we use this method? This method is suitable for a homogenous
mixture or solution. The liquid to be separated is evaporated by boiling,
and its vapour is then collected through condensation as illustrated in
Figure 2.7. The condensed vapour, which is in the form of purified liquid, is
called the distillate.
Figure 2.7: Distillation process
Source: http://cbskilkenny.ie/quickrevise/Experiments/Chemistryvids/
Distillation/
For example, seawater can be separated by distillation. Water has a much
lower boiling point than table salt as water is more volatile. If we boil
a solution of salt and water, water will evaporate and leave behind salt.
The water vapour is converted back to liquid form on the walls of the
condenser.

(d) Fractional Distillation
Now, let us move on to the fourth method which is fractional distillation.
What do you know about it? Can you define this method?
Fractional distillation is a method to separate a mixture of
compounds by their boiling points. This is done by
heating them to certain temperatures.
When do we use this method? This method is suitable for liquid-liquid
mixtures with different boiling points. When heated, the component of the
mixture with the lower boiling point will evaporate first and be distilled,
followed by the component with the next higher boiling point and so on, as
Figure 2.8: Fractional distillation
Source: http://www.chemistrydaily.com

For example, the liquid-liquid mixtures in crude oil can be separated by
fractional distillation into its components: petrol at 70C, followed by
naphtha at 140C, kerosene at 180C, diesel at 260C, and so forth.
(e) Crystallisation
Now, let us learn about crystallisation. What does it mean?
Crystallisation is a process of forming crystals from a
uniform solution.
For example, a copper (II) sulphate solution can be separated into its
components, that is, copper (II) sulphate and water, by heating the solution
until it is concentrated. Then, we do filtering, and cool down the hot filtrate
to obtain solid copper (II) sulphate in the form of crystals. To obtain the best
crystals, the crystallisation should be conducted slowly. A cold filtration
separates the crystals from the solvent, which is water.
(f) Chromatography
Lastly, let us look at the final method, which is chromatography. The
differing abilities of substances to adhere to the surfaces of various solids
such as paper and starch make it possible to separate mixtures. This is
the basis of chromatography. But what is the formal definition of
chromatography?
Chromatography refers to a set of methods used to separate
different compounds which normally involve separating
chemicals and identifying them by colour.

Ink is a good example for this method. The components in ink, which is a
dye mixture, can be separated by paper chromatography as shown in
Figure 2.9.
Figure 2.9: Chromatography
Identifying Elements, Compounds and Mixtures
Given the following substances, (i) circle the substances which are
elements, (ii) box up the substances which are compounds, and (iii) tick
off the substances which are mixtures. Good luck!
Bar of
soap
Sulphur
powder
Iron
filings and
gold
filings
Magnesium
ribbon
Sand Sugar Petrol
Water Plastic Pewter Bronze
Copper (II)
oxide and
carbon
Sodium
Calcium
chloride
Ethanol Air Iodine
Sulphuric
acid
Aluminium
foil
Charcoal
Marble
chips
Vinegar Wood Gold ring Silver ring
Rice and
salt
Dilute
hydrochloric
acid
Crude oil
ACTIVITY 2.1

DIFFERENCES BETWEEN METALS AND
NON-METALS
2.2
In the previous subtopic, you have learnt that elements can be grouped into three
groups: metal, non-metal and semi-metal. However, for this subtopic, we wil
only cover the two main groups, which are metals and non-metals. What can
we say about metals?
Metals are the largest category of elements and are easy to characterise by their
appearance. All except mercury are solid at room temperature and most have the
silvery shine. In addition, metals are generally malleable, rather than brittle (can
be pounded into thin sheets) and ductile, which means it can be twisted and
drawn into wires without breaking. Metals are also good conductors of heat and
electricity. Metals react with non-metals to form ionic compounds.
For example, the reaction of aluminium with bromine produces aluminium
bromide, an ionic compound.
2Al (s) + 3Br2 (l) 2AlBr3 (s)
Most metal oxides are basic oxides, which when dissolved in water react to form
metal hydroxides, as in the following example:
CaO (s) + H2O (l) Ca(OH)2 (aq)
Metal oxides also demonstrate their basic ability by reacting with acid to
form salt and water as illustrated in the reaction of magnesium oxide with
hydrochloric acid. This forms magnesium chloride and water.
MgO (s) + 2HCl (aq) MgCl2 (aq) + H2O (l)
Can you figure out some examples of metal? You can refer to Figure 2.10, which
shows some examples of metal. Are you familiar with these metals?

Figure 2.10: Metals
Like metals, non-metals are easy to characterise by their appearance. Non-metals
are gases, liquids or solids at room temperature. They are not silvery in
appearance and several are brightly coloured. The solid non-metals are brittle
rather than malleable, and they are poor conductors of heat and electricity. The
melting points of non-metals are generally lower than those of metals. Like
metals, non-metals react with metals to form ionic compounds.
Compounds composed entirely of non-metals are molecular compounds that
tend to be gases, liquids, or low melting point solids. Among the examples are
hydrogen chloride and carbon dioxide. Most non-metal oxides are acidic oxides,
which when dissolved in water react to form acids, as in the following example
(carbon dioxide):
CO2 (g) + H2O (l) H2CO3 (aq)
Carbon dioxide dissolves in water to form carbonic acid. Non-metal oxides also
dissolve in basic solutions to form salt as shown in the following example:
CO2(g) + 2NaOH(aq) Na2CO3(aq) + H2O(l)

Carbon dioxide reacts with a base, sodium hydroxide, to form a salt, sodium
carbonate, and water. The different physical properties of metals and non-metals
are summarised in Table 2.2.
Table 2.2: Properties of Metals and Non-Metals
Physical Properties Metals Non-Metals
Appearance of the surface Shiny Dull
Conductivity of electricity Good Poor
Conductivity of heat Good Poor
Melting point High Low
Density High Low
Malleability Malleable Non-malleable
Ductility Ductile Not ductile
CHEMICAL SYMBOLS OF ELEMENTS
2.3
How do we represent elements? A set of symbols written in the form of one or
two letters are used to represent the atoms of a particular element, just as shown
in the examples in the earlier subtopic. The first letter of an elementÊs symbol is
always capitalised, and the second letter, if any, is lowercase. For example, the
chemical symbol for the element calcium is Ca. Many of the symbols comprise
only one or two letters of the elementÊs English name such as H for hydrogen,
C for carbon, S for sulphur and so forth. Other symbols are derived from Latin
or other languages such as Na for sodium (Latin, natrium), Pb for lead (Latin,
plumbum), W for tungsten (German,wolfram).

Table 2.3 shows you examples of elements that are represented by the first letter
of its name.
Table 2.3: One-Letter Symbols of Elements
Element Symbol
Hydrogen H
Nitrogen N
Phosphorus P
Fluorine F
Iodine I
Sulphur S
Oxygen O
Table 2.4, on the other hand, shows you examples of elements that are
represented by two-letter symbols.
Table 2.4: Two-Letter Symbols of Elements
Element Symbol
Bromine Br
Magnesium Mg
Manganese Mn
Calcium Ca
Chlorine Cl
Neon Ne
Nickel Ni

Lastly, Table 2.5 shows examples of elements whose symbols are derived from
Latin names.
Table 2.5: Latin-Based Symbols of Elements
Element Latin Name Symbol
Silver Argentum Ag
Copper Cuprum Cu
Mercury Hydragyrum Hg
Potassium Kalium K
Tin Stannum Sn
Iron Ferrum Fe
Lead Plumbum Pb
SELF-CHECK 2.1
Identify the symbols for the following elements:
Xenon Osmium Nickel Fluorine Astatine
Barium Radium Platinum Plutonium Silicon
FORMULAE OF MOLECULES FOR
ELEMENTS AND COMPOUNDS
2.4
As discussed earlier, the atom is the smallest representative sample of an element.
Atoms can be combined to form molecules. Many elements found in nature are in
molecular form, that, is two or more of the same type of atoms bounded together.
For example, the oxygen normally found in air consists of molecules that contain
two oxygen atoms. Any molecule that is made up of two atoms is called a diatomic
molecule. This molecular form of oxygen can be represented by a chemical
formula. The chemical formula for a substance shows the chemical composition of
the elements present and the ratio in which the atoms of the elements occur. For a
substance composed of molecules, the chemical formula that indicates the actual

number and types of atoms in the molecule is called the molecular formula. The
molecular formula for oxygen is O2. The subscript in the formula tells us that two
oxygen atoms are present in each molecule. Other examples of elements that
normally occur as diatomic molecules are shown in Table 2.6.
Table 2.6: Molecular Formula of Elements
Element Molecular formula
Bromine Br2
Nitrogen N2
Fluorine F2
Iodine I2
Chlorine Cl2
Hydrogen H2
Compounds that are composed of molecules are called molecular compounds and
they contain more than one type of atom. For example, a molecule of water
consists of two hydrogen atoms and one oxygen atom. Its molecular formula is
H2O. An absence of subscript on the O indicates one atom of oxygen per water
molecule. Other examples of compounds that exist as molecules are shown in
Table 2.7.
Table 2.7: Molecular Formula of Compounds
Compound Molecular Formula
Hydrogen peroxide H2O2
Hydrogen chloride HCl
Carbon dioxide CO2
Carbon monoxide CO
Methane CH4
Ethene C2H4
Notice how the composition of each compound is given by its chemical formula.
Also, notice that these substances are composed only of non-metallic elements.

SIMILARITIES AND DIFFERENCES OF
ELEMENTS, COMPOUNDS AND MIXTURES
You have learnt the differences of metals and non-metals. How about elements,
compounds and mixtures? Let us study Table 2.8, which shows the similarities
and differences between elements, compounds and mixtures.
Table 2.8: Differences between Elements, Compounds and Mixtures
Aspect Element Compound Mixture
Formation Can be obtained
by breaking
down a
compound.
Formed when
elements are
combined in a
chemical reaction.
Formed by mixing
different components
together physically.
Energy involved
in the formation
Heat or light energy is
usually released or
absorbed
No energy in the form
of light or heat is
absorbed or released.
Constituents A pure substance
that is made up
of one type of
atom.
A substance made up
of two or more
different elements
which are chemically
combined.
Consists of elements,
compounds or both
which are combined
physically.
Ratio of
constituents
The proportion of
elements is fixed.
The components or
elements of a
compound combine in
a fixed proportion by
mass.
The components or
substances in a mixture
can be mixed in any
proportion by mass.
Ability to be
broken down
into simpler
substances
Cannot be broken
down into
simpler
substances.
Can be broken down
into simpler
substances by
chemical means using
heat or electricity.
The products are
either simpler
compounds or the
elements which make
up the compound.
Can be separated into
its components by
physical means.
The components are
only physically
separated. They are not
converted or broken
down into other
substances.
Properties of the
constituents
Each element has
its own distinct
properties.
The properties of a
compound are
different from their
constituent elements.
Has the properties of its
constituent elements.
2.5

ACTIVITY 2.2
You are given the following: aluminium foil, iron filings, sulphur,
magnet, evaporating dish, bunsen burner, and test tube holder.
(i) Put a small amount of iron filings into a beaker. Note the colour of
the iron filings. Wrap one end of a magnet in a paper towel and
dip it into the iron filings. Record your observations. (Silvery-white
metal attracted to the magnet.)
(ii) Put a small amount of sulphur in a beaker. Note the colour of the
sulphur. Wrap one end of a magnet in a paper towel and dip it
into the sulphur. Record your observations. (Yellow crystals not
attracted to the magnet.)
(iii) Mix the contents of the two beakers. Note the colour of the
mixture. Wrap one end of a magnet in a paper towel and dip it
into the mixture. Record your observations. (A combination of
silvery-white metal and yellow crystals. Iron filings still attracted
to the magnet but not the sulphur.)
(iv) Line the inside of an evaporating dish with aluminium foil and set
it aside for later use.
(v) Put a small amount of the mixture into a test tube. Heat the
mixture. Record your observation. (Blackish grey compound
observed).
(vi) Pour the contents of the test tube into the evaporating dish lined
with aluminium foil. Observe the product and fill in the table
below.
Substance Observation Response to Magnet
Iron “
Sulphur X
Mixture “
Compound
Compare the properties of the individual elements, mixture and
compound. Give examples to support your answer.

ALLOYS
Now, let us move on to learn about alloys. What is the definition of alloys?
These mixtures are homogenous and are prepared by heating and bonding the
metals together. The resultant alloy has completely different properties from the
starting metals.
Figure 2.11 shows you the structure of atoms in a metal which are packed
together very closely.
Figure 2.11: A metal structure
As a result, most metals have a high density. Besides that, the layers of atoms in a
metal can slide over each other easily and cause the properties of metals such as
being malleable and ductile as shown in Figure 2.12.
Figure 2.12: The metal structure before and after a force is applied to it
2.6
An alloy is a mixture of two or more elements; both
can be metals, or one can be a metal and another a
non-metal.

In an alloy, the atoms of different metals have different sizes. This makes the
layers of atoms slide over each other even harder due to the disruption of the
initial orderly layers of metal atoms. Usually a small quantity of other metals
need to be added to a pure metal to make it even harder, stronger and tougher.
Figure 2.13 shows the structure of an alloy in which the big foreign metal atoms
disrupt the orderly atom distribution of the initial metal and stop them from
sliding.
Figure 2.13: Structure of an alloy
Why do we need an alloy? The main reason of having an alloy is to enhance
the physical properties of metals. For example, although iron is a strong and
malleable metal, it suffers from a disadvantage of being prone to rust. Rust is an
oxide of iron. Rust destroys the upper layers of iron and makes the metalÊs
surface crumbly and weak. To avoid rust, iron can be mixed with nickel and
chromium to make steel. Steel or stainless steel is a highly malleable and strong
substance, which is rust proof. Thus, mixing two or more metals can bring
advantages and allow for more applications. Some other common alloys are steel,
brass, and duraluminium.
There are three types of common alloys: aluminium, iron and copper as
explained in Table 2.9.

Table 2.9: Three Types of Alloys
Type of
Alloy Description Example
Aluminium Aluminium is a bright
metal and conducts heat
and electricity well. But it
is not a strong metal. An
alloy of aluminium that
retains its good
properties and makes it
strong would be ideal for
many applications.
Alloys of aluminium that
are light but strong are
duraluminium and
magnalium.
Duraluminium:
Also called duralium or duralumin.
The proportions of the metals in
duralium are Aluminium 95%,
Copper 4%, Magnesium 0.5%, and
Manganese 0.5%. Duralium
comprises mostly aluminium, but it is
found to be very strong and corrosion
resistant. Duralium is used in aircraft
engines, car engines, pressure cooker,
and industrial cauldrons, under sea
vessels and ships.
Magnalium:
Contains Aluminium 95%, and
Magnesium 5%. Magnalium is a very
hard alloy and also light in weight.
Magnalium can be machined easily. It
is used in many instruments and
structures.
Iron To make use of ironÊs
good properties and
eliminate the possibility
of rusting, iron has to be
alloyed with other
metals. The most
important alloy of iron is
steel.
Steel:
This alloy contains 99.5% iron, and
0.5% carbon. The different carbon
contents give steel a grade; sometimes
it can be as high as 1.5%. Steel is a
much harder substance than iron. Steel
is used for making nails, screws,
railway lines, bridges, buildings, and
so on. The applications of steel are
limitless.
Stainless steel:
Nickel and chromium are added to
steel to make steel shine and attractive
in appearance. Varying proportions of
nickel and chromium can give
different grades of steel. Besides being
shiny, stainless steel is strong and
corrosion resistant. It is widely used in
making utensils, equipment and tools,
and extensively used in many
industries as containers.

Copper Copper is a relatively soft
metal and is prone to be
oxidised in the air. An
oxidised copper surface is
dull and unattractive. To
overcome these
drawbacks of copper, it
can be alloyed with
stannum, zinc and nickel
to give brass, bronze and
German silver.
Brass:
Brass is an alloy of copper and zinc
which generally has 80% copper and
20% zinc. By varying these
proportions, we can obtain various
grades of brass. Brass is stronger and
more malleable than copper. It is
golden in colour. Brass is used to make
nuts and bolts, tubes, decorative items
such as vases, jewellery, lamps, and so
on.
Bronze:
Bronze contains 90% copper and 10%
stannum (tin). Bronze is strong and is
used to make coins, medals, statues,
decorative items, and so on.
German silver:
German silver has 60% copper, 20%
zinc and 20% nickel. It has a silvery
shiny look, hence the name German
silver. It is used for electroplating so
that items look decorative. This also
prevents atmospheric corrosion.
SOLUTIONS
2.7
Now, we come to the final subtopic of this topic, which is about solutions. What
can you expect to learn in the following? Well, there are three segments in this
subtopic, and they focus on solution, solute and solvent, saturated solution, and
factors that affect solubility. Let us begin the lesson with solution, solute and
solvent.

2.7.1 Solution, Solute and Solvent
Do you know what solution stands for? Can you define it?
A solution is a homogeneous mixture of two or more substances.
How does it form? A solution is formed when tiny individual particles (1 mm
in diameter) of one substance are uniformly dispersed among the individual
particles of the other substance. An example of a solution is sugar water.
Individual molecules of sugar are uniformly distributed among the molecules of
water. Sugar dissolves or breaks down in water. The dissolved substance, which
is sugar, is called the solute while the liquid which dissolves the sugar, which is
water, is called the solvent.
Water is actually a universal solvent which can dissolve many substances. Based
on the given example, we can say that a solution is a mixture obtained by
dissolving substances, called the solute, in another substance called the solvent.
Solutions may be mixtures of two or more solids, liquids or gases, or any one of
these in another. Figure 2.14 describes the formation of a solution.

Figure 2.14: Solute, solvent and solution
Air is a good example of gas-gas mixture consisting of other gases such as
oxygen, nitrogen and carbon dioxide. Can you guess what type of solution is in a
bottle of soft drink? If you open a bottle of soft drink, you will notice bubbles
coming out of the liquid. All carbonated beverages have particles of carbon
dioxide gas dissolved in them. A soft drink is an example of a gas-liquid solution.
We have learnt about metal alloys earlier in this topic. Can you still remember?
They are good examples of solid-solid solutions where the solute is a solid and
the solvent is also a solid. Do you know that there are two types of solution
alloys? They are substitution and interstitial as shown in Figure 2.15.
Figure 2.15: Types of alloys

How are substitution alloys being formed? Substitution alloys are formed when
atoms of the solute take the position which is normally occupied by a solvent
atom. They are formed when two metallic components have similar atomic radii
and chemical-bonding characteristics.
For example, let us consider the alloy of silver and gold. Interstitial alloys are
formed when atoms of the solute occupy the interstitial positions. For an
interstitial alloy to form, the component present in the interstitial positions
between the solvent atoms must have a much smaller covalent radius than the
solvent atoms. For example, steel; it is the combination of an iron alloy and
carbon.
For your quick reference, these three types of solutions are summarised in
Table 2.10.
Table 2.10: Types of Solutions
Type of Solution Solute Solvent Examples
Gas Gas Gas Air (oxygen, nitrogen, argon and other
gases)
Liquid Gas
Liquid
Solid
Liquid
Liquid
Liquid
Carbonated water (carbon dioxide in
water)
Petroleum (mixture of hydrocarbons)
Seawater (Sodium chloride in water)
Solid Solid
Liquid
Solid
Solid
Metal alloys such as brass
Dental amalgam (mercury in silver)
Source: Adapted from McMurray and Fay (2001), p. 432
ACTIVITY 2.3
Petroleum is a solution of a liquid in liquid. Discuss what liquids are in
petroleum and how you can separate these liquids in the laboratory.

2.7.2 Saturated Solution
Do you know that a solution can be categorised into three categories? They
can be a dilute solution, concentrated solution or saturated solution. Can you
differentiate them? Let us start with dilute solution.
A dilute solution has little solute particles dissolved in the solvent. The solvent
can dissolve the solute particles more easily. When we add more solute into
the solvent, it can still dissolve. At this level, we call this solution a concentrated
solution. However, if we keep on adding the solute into the solvent until it
reaches a level where the solute cannot dissolve any more at that particular
temperature, then we call this solution a saturated solution as shown in
Figure 2.16.
Figure 2.16: Dilute, concentrated and saturated solutions

If you take solid sodium chloride, NaCl or table salt, and add it to water,
dissolution occurs rapidly at first but then slows down as more and more NaCl is
added. Eventually, the dissolution stops because a dynamic equilibrium is
reached where the number of Na+ and Cl ions leaving the crystal or solute form
to go into solution is equal to the number of ions returning from the solution
form to the crystal. At this point, the solution is said to be saturated in that solute.
Dissolve
Solute + Solvent Solution
Crystallise
A saturated solution is obtained when the solution is in equilibrium with
undissolved solid. Additional solute will not dissolve if added to such a solution.
The amount of solute needed to form a saturated solution in a given quantity
of solvent is known as the solubility of the solute. This is usually expressed in
grams of solute in 100 g of solvent.
For example, the solubility of NaCl in water is 36 g per 100 mL of water at 20C.
This is the maximum amount of NaCl that can be dissolved in water to give a
stable, equilibrium solution at that temperature. If we dissolve 40 g of NaCl per
100 mL of water at 20C, there is undissolved solute in the solution. We say
the solvent has reached its saturation point. The solution is called a saturated
solution as no more solute can dissolve in the solvent.
It is also possible to dissolve less solute than that needed to form a saturated
solution. If we dissolve 30.0 g of NaCl per 100 mL of water at 20C, the solution is
said to be unsaturated because it has the capacity to dissolve more solute as

Figure 2.17: Saturated and unsaturated solutions
2.7.3 Some Factors Affecting Solubility
As discussed earlier, when a solute dissolves in the solvent, a solution is formed.
But when only a small amount (or none at all) of a solute can be dissolved in the
water, the solute is insoluble. Then, when the solute is slightly dissolved in the
water, we get a suspension.
For example, chalkboard dust is insoluble in water and suspends on the water
surface forming a suspension. So, the question is how much solute can dissolve
in a solvent? Well, it all depends on the following three factors:
(a) Size of the solute particles;
(b) Type of the solvent; and
(c) Temperature of the solvent.

The smaller the size of the solute particles, the faster the solute dissolves in the
solvent. For example, when you make coffee or tea for a drink, you can see that
small crystals of sugar dissolve faster than a cube of sugar in your hot drink. The
type of solvent will also affect solubility. In some cases, the solute may not
dissolve in a particular solvent but may dissolve in other solvents. For example,
sugar will dissolve in water but may not dissolve in other types of solvent like
paraffin. Most solutes dissolve in water, hence water is a universal solvent. The
solubility of most molecular and ionic solids increases with increasing
temperature of the solvent. Again, you can observe this when making hot drinks.
ACTIVITY 2.4
Investigation: How much salt crystals of different sizes can be
dissolved in a liquid?
You are given two 150 ml beakers, glass rod, rock salt, and fine salt.
Using an electronic balance, weigh about 20 g each of rock salt and fine
salt. Put the rock salt in beaker A and fine salt in beaker B with 100 mL
of distilled water. Record the time and stir the contents of each beaker
gently using a glass rod. Each time the rock salt and fine salt have
dissolved, add another 20 g each of the salts into the respective beakers.
Record how many grams are used for both salts until the salts do not
dissolve. Record which salt dissolves more. Record the time taken.
Compare the different rates and the amount of salts dissolved.
SELF-CHECK 2.2
1. Compare elements and compounds. Give ONE example for each
of them.
2. Define mixture.
3. Name and describe THREE characteristics of metals and non-metals.
4. Describe how you would prepare a saturated solution of copper
(II) sulphate.

Atoms are the small building blocks of matter. The composition and structure
of atoms determine whether they are pure substances, or mixtures.
Pure substances are matter that has a fixed composition and distinct
properties. They can be classified into elements and compounds.
Mixtures are two or more substances that are mixed together but not
chemically joined. They can be classified as heterogeneous or homogenous.
An element is a substance which cannot be broken down into simpler
substances by chemical or physical methods. It can be classified into two
main groups: metals and non-metals. Each group has its own properties.
A compound is a substance which consists of two or more elements
chemically combined together.
A heterogeneous mixture is a mixture that does not have a uniform
composition.
A homogeneous mixture is a mixture that has a uniform composition.
Elements, compounds and mixtures can be different in terms of formation,
constituents, energy and many more.
Metal and non-metal elements have different properties such as appearance
of the surface, conductivity of electricity, conductivity of heat, melting point
and so on.
Elements can be represented by one-letter symbols, two-letter symbols, and
symbols derived from their Latin names.
The chemical molecular formula for a substance shows the chemical
composition of elements present and the ratio in which the atoms of the
elements occur.
The differences between elements, compounds and mixtures can be
determined in terms of formation, energy involved in the formation,
constituents and so on.

Alloys are mixtures of two or more metals.
A solution is a homogeneous mixture in which one substance (the solute) is
dissolved in another substance (the solvent).
A saturated solution is obtained when the solution is in equilibrium with
undissolved solid.
They are three factors that affect solubility: size of the solute particles, type of
the solvent, and temperature of the solvent.
Alloys
Atoms
Compounds
Elements
Metals
Mixtures
Non-metals
Saturated solution
Size
Solute
Solution
Solvent
Suspension
Temperature
Type
Brady, J. E., Senese, F. (2004). Chemistry: Matter and its changes (4th ed.).
New York: John Wiley Sons, Inc.
Briggs, J. G. R. (1992). Science in focus chemistry for GCE ÂOÊ Level. Singapore:
Pearson Education.
Brown, T. L., Lemay, H. E., Bursten, B. E. (2000). Chemistry: The central science
(8th ed.). New Jersey: Prentice Hall.
Kots, J. C., Treichel, P. M., Weaver, G. C. (2006). Chemistry: The chemical
reactivity (2nd ed.). Victoria, Australia: Thomson Learning.

McMurray, J., Fay, R. C. (2001). Chemistry (3rd ed.). New Jersey: Prentice Hall.
Timberlake, K. C. (2006). An introduction to general, organic, and biological
chemistry (9th ed.). San Francisco, CA: Pearson-Benjamin Cummings.
Whitten, K. W., Davis, R. E., Peck, M. L., Stanley, G. G. (2010). Chemistry
(9th ed.). Belmont: Brooks/Cole.

Topic
3
Atomic
Structure
LEARNING OUTCOMES
1. Explain the history of atomic model;
2. Describe subatomic particles;
3. Differentiate between atomic number, nucleon number and mass
number;
4. Summarise the concepts of isotopes;
5. Show how to configure the electronic configuration of atom; and
6. Summarise the concept of valence electrons.
INTRODUCTION
Hi there and welcome to the third topic of this module. Before we learn more
about atomic structure, let us recall the definition of matter. Matter is defined as
anything that has mass and takes up space or volume. Matter is made up of
discrete atoms. We will continue to learn more about atoms in this topic.
In this topic, we will learn about the history of atomic model, subatomic
particles, atomic number, nucleon number and mass number. Then, we will
study isotopes, and the electronic configuration of atoms. Lastly, we will learn
about valence electrons. Let us start the lesson!

TOPIC 3 ATOMIC STRUCTURE 59
HISTORY OF ATOMIC MODEL
We begin this topic by tracking the atomic model history. According to the
atomic theory, matter is made up of much smaller particles known as atoms.
Do you know that the history of atoms started from John DaltonÊs model? Later,
James Chadwick provided a complete portrayal of the components in an atom.
Let us look at Table 3.1, which shows the different atomic models and their
explanations.
Table 3.1: Models of Atom
Contributor Model Explanation
John Dalton Billiard Ball Model (1805)
Source:
https://reichchemistry.wikispaces.com
DaltonÊs atomic model
was portrayed as a small
indivisible ball similar to a
very tiny ball.
Joseph John
Thomson
Plum Pudding Model (1897)
Source: http://www.scienceclarified.com
Thomson discovered the
electron, a negatively-charged
particle. The atom
was described as a sphere
of positive charge with
electrons embedded in it.
Ernest
Rutherford
Solar System Model/Rutherford Model
(1911)
Source: http://www.faqs.org
Rutherford discovered the
proton, a positively-charged
particle in an
atom. The proton and
most of the mass of the
atom were concentrated in
the central region called
the nucleus. The electrons
moved in the spherical
space outside the nucleus.
3.1

60 TOPIC 3 ATOMIC STRUCTURE
Neils Bohr Bohr Model (1913)
Source:
http://www.hsctut.materials.unsw.edu.au
According to Bohr, the
electrons in an atom were
not randomly distributed
around the atomic
nucleus, but moved
around the nucleus in
fixed orbits (shell). Each
orbit formed a circle and
had a fixed distance from
the nucleus.
Source: http://www.csmate.colostate.edu
SUBATOMIC PARTICLES
In the study of atomic structure, we will look first at the subatomic particles, also
known as fundamental particles. These are the basic building blocks of all atoms.
Atoms consist principally of three subatomic particles: electrons, protons and
neutrons.
Both the protons and neutrons reside in the nucleus and they are called nucleon.
As seen in Figure 3.1, the electrons reside in orbits around the nucleus.
Figure 3.1: Electrons orbiting around the nucleus
Source: http://www.cfo.doe.gov
3.2

Let us examine these particles in detail. The relative mass and charge of the three
subatomic particles are shown in Table 3.2.
Table 3.2: Symbols, Relative Electric Charge and Relative Masses of Subatomic Particles
Element Symbol Relative Electric Charge Relative Mass
Proton p +1 1
Neutron n 0 Approximately 0.0005
Electron e 1 1
The mass of an electron is very small compared with the mass of either a proton
or a neutron. The charge on a proton is equal in magnitude, but opposite in sign
to the charge on an electron. Since the masses of protons and neutrons are greater
than those of electrons, the mass of an atom is mostly concentrated in the
nucleus. An atom consists of an equal number of electrons and protons. Hence,
an atom is electrically neutral.
ACTIVITY 3.1
ATOMIC NUMBER, NUCLEON NUMBER
AND MASS NUMBER
Now, let us move on to atomic number, nucleon number and mass number.
Firstly, what does atomic number mean?

3.3
The atomic number refers to the number of protons in an
atom which is represented by the symbol Z.

Can you find out more on the Internet about the similarities and
differences between the charge of a proton and an electron?
Compare and then make a summary on the differences between the
charge of a proton and an electron.

Do you know that the atomic number of an element is the identity of the
element? This is because the number of proton in the nucleus of every atom in an
element is always the same.
For example, each hydrogen atom contains only one proton and its atomic
number is 1. On the other hand, the carbon atom has six protons and its atomic
number is 6, whereas the atomic number of oxygen is 8 because it contains eight
protons in its nucleus.
What about a neutral atom? For a neutral atom, the number of its electrons is
equal to the number of its proton or the atomic number. In other words, the
proton number of an atom can also represent the number of electrons.
How about the nucleon number? What does it mean?

The nucleon number of an element is the total number of
protons and neutrons in the nucleus of an atom.
Do you know that nucleon number is sometimes referred to as the mass number?
This is because since the mass of an atom is very small, the nucleon number of an
atom is almost the same as the mass of the atom. The mass number is represented
by the symbol A as shown below.

Nucleon number (A) = Number of protons (Z) + Number of neutrons
How do we use the atomic number and nucleon number? Figure 3.2 shows the
standard representation for an atom of any element by using atomic number
(proton number) and nucleon number.

Figure 3.2: Nucleon and proton numbers of an element
Let us see an example of a standard representation for an atom as shown in
Figure 3.3. It shows you the nucleon and proton numbers contained in the
helium atom.
Figure 3.3: Nucleon and proton numbers of a Helium atom
ACTIVITY 3.2
Calculate the number of protons, electrons and neutrons and fill in
the table below.
Atom Nucleon
Number
Proton
Number
No. of
Proton
No. of
Electron
No. of
Neutron
Helium 4 2 2 2 2
Oxygen 16 8
Sodium 23 11
Chlorine 35 17

ISOTOPES
3.4
Let us learn about isotopes now. Firstly, do you know that there are atoms which
have the same number of protons but different number of neutrons? We call
these atoms as isotopes. Based on the previous statement, can you define
isotopes?

Wecandefineisotopesasatoms ofthesameelementwith
thesamenumberofprotonsbutwithdifferentnumberof
neutrons.
For example, there are three distinct kinds of hydrogen atoms, commonly called
hydrogen, deuterium and tritium, as shown in Figure 3.4. Each contains one
proton in the atomic nucleus.
Figure 3.4: The three isotopes of hydrogen
Source: http://www.pppl.gov
Other examples of isotopes are the Carbon-12 and Carbon-14 isotopes. Both have
the same number of protons, which is 6, but different number of neutrons.
Carbon-12 has six neutrons, whereas Carbon-14 has eight neutrons.
Do you know that there are similarities and differences between isotopes of the
same element? These similarities and differences between isotopes of the same
element are summarised in Table 3.3.

Table 3.3: Similarities and Differences between Isotopes of the Same Element
Isotopes of the Same Element
Similarities Differences
(i) Same proton number
(ii) Same number of electrons in an
atom
(iii) Same electron arrangement
(i) Different nucleon numbers
(ii) Different neutron numbers in an
atom

How about isotopes for other elements? You can refer to Table 3.4, which shows
isotopes for hydrogen, oxygen, chlorine, carbon and sodium.
Table 3.4: Isotopes of Some Elements
Element Isotopes of
Element Symbol Nucleon
Number
Proton
Number
Neutron
Number
Hydrogen
Hydrogen 3J
3 1 1 0
Deuterium 4J
3 12 1 1
Tritium 5J
3 13 1 2
Oxygen
Oxygen16 38Q
: 16 8 8
Oxygen17 39Q
: 17 8 9
Oxygen18 3:Q
: 18 8 10
Chlorine
Chlorine35 57En
39 35 17 18
Chlorine37 59En
39 37 17 20
Carbon
Carbon12 34E
8 12 6 6
Carbon13 35E
8 13 6 7
Carbon14 36E
8 14 6 8
Sodium
Sodium23 45Pc
33 23 11 12
Sodium24 46Pc
33 24 11 13

Do you know that isotopes can be used in various fields? Let us look at Table 3.5,
which describes the specific isotopes and their usages.
Table 3.5: Specific Isotopes and their Usages
Isotope Field Usage
Iodine-31 Medical The treatment of thyroid disease.
Krypton-85 Industry To control the thickness of plastic sheets in the
plastic industry.
Uranium-235 Power Resources Nuclear power stations.
Carbon-14 Agriculture To carry out experiments or studies regarding
photosynthesis and protein synthesis.
Phosporous-32 General Research Used in fertilisers to study the metabolism of
phosphorus in plants.
ACTIVITY 3.3
Find at least THREE different isotopes for each of these areas:
medicine, power resources, agriculture, and general research. You
can research using the Internet, books or encyclopaedias to get the
answers. Good luck!

ELECTRONIC CONFIGURATION OF ATOMS
As mentioned earlier, atoms are made up of protons, neutrons and electrons.
Where are they located? The protons and neutrons are located in the nucleus of
an atom. How about electrons? The electrons are not randomly located but are
actually arranged in shells or energy levels around the nucleus of an atom.
The shells of an atom are numbered 1, 2, 3 and so on, starting from the one
closest to the nucleus. Each shell can occupy a certain number of electrons. For
atoms with the proton numbers of 1 to 20, the first shell can hold a maximum of
two electrons. As for the second shell, it can hold a maximum of eight electrons.
This is followed by the third shell, where it can also hold a maximum of eight
electrons.
Now, let us look at Figure 3.5, which shows the potassium atom.
Figure 3.5: Electron configuration of the potassium atom
3.5

How do we find out the proton number of the atom? For a neutral atom, the
number of electrons is the same as the number of protons. Referring to the
periodic table, the potassium atom has 19 electrons. Let us arrange the electrons
in shells by following this rule: electrons occupy the shells closest to the nucleus
first, and they occupy a new shell when a previous one has been occupied.
Therefore, the first shell of the potassium atom has a maximum of two electrons;
the second and the third shells each have a maximum of eight electrons. The
outer shell has one electron. These are summarised as follows:
Number of electrons in the first shell: 2
Number of electrons in the second shell: 8
Number of electrons in the third shell: 8
Number of electrons in the last shell: 1
The electron configuration of potassium = 2.8.8.1
You can refer to Figure 3.6 for a traditional representation of an atomÊs electronic
configuration. It is a dot and cross diagram. Figure 3.6 shows a nitrogen atom
which has seven electrons two electrons in the first shell and five electrons in
the second shell. Therefore, the electron configuration of the nitrogen atom is 2.5.
Figure 3.6: Dot and cross diagram of nitrogen atom

SELF-CHECK 3.1
1. Given that 62Ec
F-CHECK 2.1
42 57En
39 8Nk
5 ,
(a) Write the electron configuration for the following
elements.
(b) Draw a dot and cross diagram for each of the elements.
2. Complete the table below.
Element
Proton
Number
Electron
Number
Number of Electron in Shell: Electron
1st 2nd 3rd 4th Arrangement
Hydrogen 1 1
Helium 2 2
Lithium 3 3
Beryllium 4 4
Boron 5 5
Carbon 6 6
Nitrogen 7 7
Oxygen 8 8
Fluorine 9 9
Neon 10 10
Sodium 11 11
Magnesium 12 12
Aluminium 13 13
Silicon 14 14
Phosphorus 15 15
Sulphur 16 16
Chlorine 17 17
Argon 18 18
Potassium 19 19
Calcium 20 20

VALENCE ELECTRONS
As we have learnt earlier, if the number of electrons is less than 20, the first shell
can hold a maximum of two electrons, the second shell eight electrons and the
third shell eight electrons.
What if the number of electrons is more than 20? Then, the third shell can hold a
maximum of up to 18 electrons. However, for the purpose of this module, the
focus will be on elements with less than 20 electrons only.
So, let us continue our lesson on valence electrons. Firstly, what are valence
electrons?

Do you know that from the electron arrangement, we can determine the number
of valence electrons in an atom? Let us look at an example.
A chlorine atom has an electron arrangement of 2.8.7. There are seven electrons
in the outermost occupied shell of the chlorine atom. Thus, the number of
valence electrons in a chlorine atom is 7.
Before we end this subtopic as well as this topic, let us summarise the
relationship between the number of valence electrons and group/period number.
The summary is shown in Table 3.6.

Table 3.6: Number of Valence Electrons and Group Number
Number of Valence
Electrons 1 2 3 4 5 6 7
8
(Except
Helium)
Group 1 2 13 14 15 16 17 18
3.6
Valence electrons are the electrons found in the outermost
shell of an atom. It is the furthest shell from the nucleus.

Based on Table 3.6, we can deduce that for elements with one or two valence
electrons, the group number of these elements is equal to the number of valence
electrons contained inside the elements.
As for elements with three to eight valence electrons, the group number of these
elements is equal to the number of valence electrons plus the number 10. An
exception to the rule is Helium as it is placed in Group 18, despite having an
electron arrangement of two.
How about the number of shells and period number? You can refer to Table 3.7
which shows the number of shells and period number based on each group.
Table 3.7: Number of Shells and Period Number
Number of
Shells Occupied
with Electrons
1 2 3 4 5 6 7
Group 1 2 3 4 5 6 7
Based on Table 3.7, we can see that the period number of an element is equal to
the number of shells occupied with electrons in an atom of the particular
element. This is quite similar to Table 3.6, right?
As a conclusion, we can say that for elements with one or two valence electrons,
the group number of these elements is equal to the number of valence electrons
contained inside the elements, and as for the period number of an element, it is
equal to the number of shells occupied with electrons in an atom of the particular
element.
The atomic theory states that all matter is made up of atoms.
The history of atoms started when an atom was portrayed as a tiny ball.
Later, the electron, a negatively-charged particle, was discovered. This was
followed by the discovery of protons and nucleus.
Atoms consist of three subatomic particles: electrons, protons and neutrons.
The atomic number of an element is the number of its protons in the nucleus.
It is represented by the symbol Z.

The mass number of an element is the total number of neutrons and protons
in the nucleus of the atom. It is also called the nucleon number and is
represented by the symbol A.
Isotopes are atoms of the same element with the same number of protons but
with a different number of neutrons.
Isotopes are mostly used in various fields such as medicine, industry,
agriculture, power resources and general research.
The electrons are not randomly located but are arranged in shells or energy
levels around the nucleus of an atom.
The first shell can hold a maximum of two electrons. The first shell will be
filled first.
The second shell can hold a maximum of eight electrons. The third shell can
hold a maximum of eight electrons.
A valence electron is the electron of the outermost shell. The number of
valence electrons in an atom can be determined from its electron
arrangement.
Atomic number
Electron shells
Electron configuration
Isotopes
Mass number
Atomic number
Electron shells
Electron configuration
Isotopes
Mass number

Briggs, J. G. R. (2003). Science in focus chemistry for GCE ÂOÊ Level. Singapore:
Pearson Education Asia Pte Ltd.
Conoley, C., Hills, P. (2002). Chemistry (2nd ed.). London: Harper-Collins.
Hewitt, P. G. (1998). Conceptual physics (8th ed.). Reading, Massachusetts:
Addison-Wesley.
Kementerian Pendidikan Malaysia Bahagian Pendidikan Guru. (1995) Buku
sumber pengajaran pembelajaran sains sekolah rendah: Strategi pengajaran
dan pembelajaran sains. Kuala Lumpur: Kementerian Pendidikan Malaysia.
Ralph, A. B. (2003). Fundamentals of chemistry. New Jersey: Prentice Hall.
Whitten, K. W., Davis, R. E., Peck, M. L., Stanley, G. G. (2010). Chemistry
(9th ed.). Belmont: Brooks/Cole.

Topic
4
Periodic Table
LEARNING OUTCOMES
1. Analyse the periodic table;
2. Summarise the electronic structures and periodic table;
3. Identify properties and usages of transition elements;
4. Identify the electronic structure, group trends, physical properties
and chemical properties of Group 1 and Group 17;
5. Summarise noble gases; and
6. Identify the properties and classification of Period 3 elements.
INTRODUCTION
Hello and letÊs start Topic 4! In this topic, you will learn about the periodic table.
Before we go further, do you know that there are 118 discovered elements in
nature? Most of these elements are naturally occurring elements. However, a few
of these elements are made up artificially in nuclear reactors. Elements with the
same chemical properties were grouped together by chemists, resulting in
the development of the periodic table. This systematic method of classifying
elements has enabled us to study and generalise the chemical and physical
properties of elements in the same group.

We will learn more about the periodic table as we track back its history and
study how the groups and periods of the periodic table can be analysed. This is
followed by the electronic structures and the periodic table, and properties and
usages of transition elements.

TOPIC 4 PERIODIC TABLE
75
Then, we will examine the electronic structure, group trends, physical properties
and chemical properties of Group 1 and Group 17. Last but not least, we will look
at noble gases and Period 3 elements. Are you ready now? Let us start the
journey!
HISTORY OF THE PERIODIC TABLE
4.1
Let us now review the history of the periodic table as well as the events which
led to the development of the modern periodic table. Do you know that the
majority of the elements that we know today were actually discovered during the
18th and 19th century? You will notice that elements with similar properties were
grouped together systematically in a table. This marked the beginning of the
development of the periodic table.
Chemists such as Lavoiser, Dobereiner, Newlands, Meyer, Mendelev and Mosely
contributed to the development of the periodic table in use today. We will now
read about their respective contributions.
4.1.1 Antoine Lavoisier (1743–1794)
Do you know that Antoine Lavoisier (Figure 4.1) was the first scientist to classify
elements into four groups? He classified substances, including light and heat,
into metals and non-metals.
Figure 4.1: Antoine Lavoisier (17431794)
Source: http://www.sciencephoto.com
However, his classification was not successful due to wrong information. For
example, non-elements such as heat and light, and compounds such as silica,
magnesia, chalk, barita and alumina were included in his classification table.

7 6 TOPIC 4 PERIODIC TABLE
Table 4.1: Antoine LavoisierÊs 1789 Classification of Substances
Acid-making Gas-like Elements Metallic Elements Earthy Elements
Sulphur Light Cobalt, Mercury,
Tin
Lime (Calcium
Oxide)
Phosphorus Caloric (Heat) Copper, Nickel, Iron Magnesia
(Magnesium Oxide)
Charcoal (Carbon) Oxygen Gold, Lead, Silver,
Zinc
Barytes (Barium
Sulphate)
Azote (Nitrogen) Manganese,
Tungsten
Argilla (Aluminium
Oxide)
Hydrogen Platina (Platinum) Silvex (Silicon
Dioxide)
Source: http://www.docbrown.info/page12/gifs/Lavoisier1789.gif
4.1.2 Johann Dobereiner (1780–1849)
Johann Dobereiner (Figure 4.2) divided the elements into groups. Each group
consists of three elements with similar chemical properties and is called a triad.
In each triad, the atomic weight of the middle element is the average of the other
two elements. According to the Law of Triad, the atomic mass of sodium is the
mean of the total atomic mass of lithium and potassium. Thus, the atomic mass of
sodium is 23 (refer to Table 4.2).
Figure 4.2: Johann Dobereiner (17801849)
Source: http://elements-table.com/history/

77
Table 4.2: Law of Triad
Element Symbol A (Atomic Mass)
Lithium Li 7
Sodium Na 23
Potassium K 39
Mean of Li + K = (7 + 39)/2 = 46/2 = 23 (The value of Na)
However, this classification was unsuccessful because the classification was
limited to a few elements only. Then, other scientists realised that there was a
relationship between the properties and atomic masses of the elements, as shown
in Table 4.3.
Table 4.3: Relationship between the Properties and Atomic Masses of the Elements
Triads III IV
Elements Copper
Cu
Silver
Ag
Gold
Au
Zinc
Zn
Cadmium
Cd
Mercury
Hg
Atomic weights 635 108 197 65 112.5 200
Mean Weights 130.25 132.5
Source: http://www.tutornext.com/ws/402-g-limit
4.1.3 John Newlands (1837–1898)
Another chemist that contributed to the existence of the periodic table was John
Newlands (Figure 4.3).
Figure 4.3: John Newlands (18371898)
Source: http://elements-table.com

Newlands arranged all the known elements horizontally in the ascending
order of their atomic masses. Each row consisted of seven elements. He found
that elements with similar properties recurred at every eighth element. This
arrangement was known as the Law of Octaves.

However, this law was only obeyed by the first 17 elements. Thus, it was not
successful. There were no positions allocated for elements yet to be discovered.
However, Newlands contibution to the development of the periodic table was
very important as he was the first chemist who discovered the existence of
periodicity in the elements.
4.1.4 Lothar Meyer (1830–1895)
Lothar Meyer (Figure 4.4) plotted a graph of atomic volume against atomic mass
for all known elements. He found that elements with the same chemical
properties occupied the same relative positions on the curve. He showed that the
properties of the elements were in a periodic pattern with their atomic masses.
Hence, Meyer also proved that the properties of the elements recur periodically.
Figure 4.4: John Newlands (18371898)
Source: http://www.wou.edu

79
4.1.5 Dmitri Mendeleev (1834–1907)
Dmitri Mendeleev (Figure 4.5) showed that the properties of elements changed
periodically with their atomic mass. He arranged the elements in the order of
increasing atomic mass and grouped them according to similar chemical
properties. He was able to predict the properties of undiscovered elements and
left gap for these elements.
Figure 4.5: Dmitri Mendeleev (18341907)
Source: http://chemistry.about.com
Mendeleev had also correctly predicted the properties of the elements gallium,
scandium and germanium which were only discovered much later. MendeleevÊs
table was used as a blueprint for the modern periodic table. Figure 4.6 shows
MendeleevÊs periodic table.
Figure 4.6: MendeleevÊs periodic table
Source: http://www.msnucleus.org

4.1.6 Henry J. G. Moseley (1887–1915)
Henry J. G. Moseley (Figure 4.7) studied the x-ray spectrum of elements. He
concluded that the proton numbers should be used as a basis for the periodic
change of chemical properties instead of the atomic mass. He rearranged the
elements in the ascending order of their proton numbers.
Figure 4.7: Henry J. G. Moseley (18871915)
Source: http://en.wikipedia.org
Similar to Mendeleev, Mosely left gaps for elements yet to be discovered. He
produced a periodic table which was almost the same as MendeleevÊs periodic
table. Thus, he confirmed the work of Mendeleev.
Due to MoseleyÊs work, the periodic table was successfully developed and being
used today. The modern periodic table is based on the arrangement of elements
in the ascending order of their proton numbers. Finally, the periodic table is as
what we see today.

81
4.1.7 Modern Periodic Table
Based on our earlier discussions about the early history of the periodic table,
what can you conclude about it? How would you define the periodic table?

The periodic table is a classification of elements whereby
elements with the same chemical properties are placed in
the same group. This makes the study of the chemistry of
these elements easier and more systematic.
Later, Glenn Seaborg (Figure 4.8) discovered that the transuranium elements
have atomic numbers from 94 to 102, resulting in the redesign of the periodic
table.
Figure 4.8: Gleen Seaborg
Source: http://www.wired.com
Technically, both the lanthanide and actinide series of elements are to be placed
between the alkaline earth metal and the transition metal.
However, by doing this, the periodic table would be too wide. Thus, the
lanthanide and actinide series of elements were placed under the rest of the
periodic table. This is the periodic table that we use today. Dr Seaborg and his
colleagues were also responsible for identifying more than 100 isotopes of
elements.

Figure 4.9 shows the modern periodic table. From here on, we will do an in-depth
study of the periodic table. Based on calculation, there are 118 elements in
the current periodic table but for the purpose of study for this module, only 111
elements will be considered.
Figure 4.9: The modern periodic table
Source: http://www.webelements.com/
SELF-CHECK 4.1
1. List the name of the chemists who played a significant role in the
early development of the periodic table.
2. What was the conclusion of the study by Henry J. G. Moseley?
3. Define the periodic table in your own words.
4. Differentiate between the old version and the modern version of
the periodic table.

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