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Executive
Preview

Chemistry

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for the IB Diploma

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MULTI-COMPONENT SAMPLE
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Third edition Digital Access

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Dear Teacher,

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Welcome to the new edition of our Chemistry for the IB Diploma series, providing full
support for the new course for examination from 2025. This new series has been designed
to flexibly meet all of your teaching needs, including extra support for the new assessment.
This preview will help you understand how the coursebook, the workbook and the teacher’s
resource work together to best meet the needs of your classroom, timetable and students.
This Executive Preview contains sample content from the series, including:
•
•

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A guide explaining how to use the series
A guide explaining how to use each resource
In developing this new edition, we carried out extensive global research with IB Chemistry
teachers – through lesson observations, interviews and work on the Cambridge Panel, our
online teacher research community. Teachers just like you have helped our experienced
authors shape these new resources, ensuring that they meet the real teaching needs of the
IB Chemistry classroom.
The coursebook has been specifically written to support English as a second language
learners with key subject words, glossary definitions in context and accessible language
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throughout. We have also provided new features that help with active learning, assessment for
learning and student reflection. Numerous exam-style questions with answers in the digital
coursebook, which accompanies the print coursebook, ensure your students feel confident
approaching the assessment and have all the tools they need to succeed in their examination.
Core to the series is the brand-new digital teacher’s resource. It will help you support
your learners and confidently teach to the new IB Chemistry guide, whether you are new
to teaching the subject or more experienced. For each topic there are lesson ideas and
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activities, common misconceptions to look out for, worksheets, PowerPoint presentations,

answers to the coursebook, extra wrap-up activities and more. Also included is a practical
guide to help your students develop their academic writing.
Please take five minutes to find out how our resources will support you and your learners.
To view the full series, you can visit our website or speak to your local sales representative.
You can find their contact details here:
cambridge.org/gb/education/find-your-sales-consultant
Best wishes,

Micaela Inderst
Senior Commissioning Editor for the IB Diploma
Cambridge University Press

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CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

How to use this series

This suite of resources supports students and teachers of the Chemistry course for the IB
Diploma programme. All of the books in the series work together to help students develop the
necessary knowledge and scientific skills required for this subject.

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The coursebook with digital access provides full coverage of
the latest IB Chemistry guide.
It clearly explains facts, concepts and practical techniques, and
uses real world examples of scientific principles. A wealth of
formative questions within each chapter help students develop

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their understanding, and own their learning. A dedicated
chapter in the digital coursebook helps teachers and students
unpack the new assessment, while exam-style questions provide
essential practice and self-assessment. Answers are provided on
Cambridge GO so support self-study and home-schooling.

The workbook builds upon the coursebook with

digital access with further exercises and exam-style Chemistry
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questions, carefully constructed to help learners for the IB Diploma

develop the skills that they need as they progress

WORKBOOK

Jacqueline Paris

through their IB Chemistry Diploma course. The

exercises also help students develop understanding
of the meaning of various command words used
in questions, and provide practice in responding
appropriately to these.
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Second edition Digital Access

Together with IB teachers

The teacher’s resource supports and enhances the coursebook

with digital access and the workbook. This resource includes
teaching plans, overview of required background knowledge,
learning objectives and success criteria, common misconceptions,
and a wealth of ideas to support lesson planning, assessment
and differentiation. detailed lesson ideas. It also includes editable
worksheets for vocabulary support and exam practice (with
answers) and exemplar PowerPoint presentations, to help you
plan and deliver your best teaching.

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

Chemistry

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for the IB Diploma

PL COURSEBOOK

Steve Owen
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Third edition Digital Access

4 Together with IB teachers

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CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

Contents
How to use this series vii
How to use this book viii
Unit 1 The nature of matter 1 Unit 2 Bonding and structure  103

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1 The particulate nature of matter 2 6 The ionic model 104
1.1 Elements, compounds and mixtures  3 6.1 Ionic and covalent bonding 105
1.2 Kinetic molecular theory  11 6.2 Formation of ions 105
1.3 Temperature and kinetic energy  13 6.3 The formation of ionic compounds 108
1.4

2.1
2.2

3.1
3.2
Changes of state

2 The nuclear atom

Isotopes21

3 Electron configurations
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The structure of atoms 18

The electromagnetic spectrum

The hydrogen atom spectrum
31
32
17

30
14 6.4

6.5

6.6

7.1
7.2
Ionic bonding and the structure of
ionic compounds110
Physical properties of ionic
compounds111
Lattice enthalpy and strength
of ionic bonding

7 The covalent model

Covalent bonds
113

118
Shapes of molecules: VSEPR theory 128
117
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3.3 Electron configurations 37
3.4 Putting electrons into orbitals: 7.3 Lone pairs and bond angles 131
Aufbau principle44 7.4 Multiple bonds and bond angles 132
3.5 Ionisation energy 47 7.5 Polarity135
7.6 Pauling electronegativities 136
4 Counting particles by mass: 7.7 Intermolecular forces 138
The mole  57
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7.8 Melting points and boiling points 145

4.1 Relative masses 58 7.9 Solubility147
4.2 Moles60 7.10 Covalent network structures 152
4.3 The mass of a molecule 64 7.11 The expanded octet 156
4.4 Empirical and molecular formulas 66 7.12 Formal charge 157
4.5 Solutions75 7.13 Shapes of molecules and ions with
4.6 Avogadro’s law 86 an expanded octet 161
7.14 Hybridisation165
5 Ideal gases 89 7.15 Sigma and pi bonds 169
5.1 Real gases and ideal gases 90 7.16 Resonance and delocalisation 173
5.2 Macroscopic properties of ideal gases 92
5.3 Calculations involving ideal gases 94

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Contents

8 The metallic model 182 13 Energy cycles in reactions 353

8.1 Classifying elements as metals 183 13.1 Bond enthalpies 354
8.2 Metallic bonding 184 13.2 Hess’s law 360
8.3 Properties of metals and their uses 186 13.3 Using standard enthalpy change
8.4 Transition metals 188 of combustion data  368
13.4 Standard enthalpy changes
9 From models to materials 192 of formation372
9.1 Alloys193 13.5 Energy cycles for ionic compounds 375
9.2 Polymers195
14 Energy from fuels 485
9.3 Bonding and electronegativity 209

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14.1 Combustion reactions 386
14.2 Fuels390
Unit 3 Classification of matter 215
14.3 Renewable and non-renewable
energy sources397
10 The periodic table 216
14.4 Fuel cells 399
10.1
10.2
10.3
10.4
10.5
10.6
The periodic table
Periodicity222

Oxidation state
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The transition metals (d block)

11 Functional groups: Classification

of organic compounds 
217

The chemistry of Group 1 and Group 17 231

Oxides235
240
244

256
15 Entropy and spontaneity
15.1
15.2
15.3
Entropy404
Spontaneous reactions
Gibbs energy and equilibrium

Unit 5 How much, how fast,

how far?
16 How much? The amount of
408
414
403

419
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11.1 The structures of organic molecules 258
11.2 Homologous series and functional chemical change 420
groups263 16.1 The meaning of chemical equations 421
11.3 Naming organic molecules 270 16.2 Yield and atom economy of
11.4 Isomers383 chemical reactions 433
11.5 Spectroscopic identification of 16.3 Titrations436
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organic compounds397 16.4 Linked reactions 440

Unit 4 What drives chemical 17 How fast? The rate of

reactions?331 chemical change 447
17.1 What is ‘rate’ of reaction? 448
12 Measuring enthalpy changes 332
17.2 Experiments to measure the rate
12.1 Heat and temperature 333 of reaction449
12.2 Exothermic and endothermic reactions 333 17.3 Collision theory 452
12.3 Enthalpy changes and standard 17.4 Factors affecting reaction rate 453
conditions337
17.5 The rate equation 460
12.4 Measuring enthalpy changes 338
17.6 Mechanisms of reactions 471
17.7 Variation of the rate constant with
temperature482

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CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

18 How far? The extent 20.4 Voltaic cells 600

of chemical change490 20.5 Rechargeable batteries 605
18.1 Reversible reactions and equilibrium 491 20.6 Electrolysis 610
18.2 The position of equilibrium 494 20.7 Redox reactions in organic chemistry 613
18.3 Equilibrium constants 498 20.8 Reduction reactions 619
18.4 Calculations involving equilibrium 20.9 Standard electrode potentials 623
constants 504 20.10 Electrolysis of aqueous solutions 637
18.5 Relationship between equilibrium
constants and Gibbs energy 515 21 Electron sharing reactions 643
21.1 Radicals644

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Unit 6 Mechanisms of chemical 21.2 The radical substitution mechanism 647
change521
22 Electron-pair sharing reactions 651
19 Proton transfer reactions 522 22.1 Nucleophilic substitution reactions 652
19.1 Acids, bases and salts 523 22.2 Addition reactions 656
19.2
19.3
19.4
19.5
19.6
19.7
19.8
19.9
19.10
Reactions of acids

pH532

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Brønsted–Lowry acids and bases

Strong and weak acids and bases

The dissociation of water
Calculating pH values
Acid–base titrations
pOH547
Ionisation constants for acids
526
529

534
539
541
543
22.3
22.4
22.5

22.6

Glossary681
Index693
Lewis acids and bases
Nucleophilic substitution mechanisms
Electrophilic addition reactions
of alkenes
Electrophilic substitution reactions
661
663

670
676
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Acknowledgements694
and bases 549
19.11 The base ionisation constant, Kb553
19.12 The strength of an acid and its
conjugate base 558
19.13 The pH of salt solutions 559
19.14 More pH curves 563
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19.15 Buffer solutions 572

20 Electron transfer reactions 585

20.1 Redox reactions 586
20.2 Redox equations 591
20.3 The activity series 598

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How to use this book

How to use this book

Throughout this book, you will find lots of different features that will help your learning. These are explained below.

UNIT INTRODUCTION
A unit is made up of a number of chapters. The key concepts for all the chapters covered in a unit are
summarised in the Unit opening chapter as the introduction.

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LEARNING OBJECTIVES
Each chapter in the book begins with a list of learning objectives. These set the scene for each chapter,
help with navigation through the coursebook and indicate the important concepts in each topic. A bulleted

GUIDING QUESTIONS

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list at the beginning of each section clearly shows the learning objectives for the section.

These are questions on subject knowledge you

will need before starting each chapter.

Link
EXAM TIPS
These short hints provide useful information that
will help tackle the tasks in the exam.

SCIENCE IN CONTEXT
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These are a mix of questions and explanation that refer This feature presents real-world examples
to other Chapters or sections of the book. and applications of the content in a chapter,
encouraging you to look further into topics.
The content in this book is divided into Standard and You will note that some of these features end with
Higher Level material. Either a chevron or a vertical line questions intended to stimulate further thinking,
running down the margin of all Higher Level material, prompting you to look at some of the benefits
allows you to easily identify Higher Level from Standard and problems of these applications.
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material.

NATURE OF SCIENCE
Key terms are highlighted in orange bold font at their
first appearance in the book so you can immediately Nature of Science is an overarching theme of the
recognise them. At the end of the book, there is a IB Chemistry Diploma course The theme examines
glossary that defines all the key terms. the processes and concepts that are central to
scientific endeavour, and how science serves and
connects with the wider community. Throughout
KEY POINTS the book, there are ‘Nature of Science’ paragraphs
that discuss particular concepts or discoveries from
This feature contains important key learning
the point of view of one or more aspects of Nature
points (facts) and/or equations to reinforce your
of Science.
understanding and engagement.

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CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

THEORY OF KNOWLEDGE TEST YOUR UNDERSTANDING

This section stimulates thought about critical thinking These questions appear within each chapter,
and how we can say we know what we claim to know. to help you develop your understanding. The
You will note that some of these features end with questions can be used as the basis for class
questions intended to get you thinking and discussing discussions or homework assignments. If you
these important Theory of Knowledge issues. can answer these questions, it means you have
understood the important points of a section.

INTERNATIONAL MINDEDNESS
Throughout this Chemistry for the IB Diploma WORKED EXAMPLE

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course, the international mindedness feature Many worked examples appear throughout the text to
highlights international concerns. Chemistry is help you understand how to tackle different types of
a truly international endeavour, being practised questions.
across all continents, frequently in international
or even global partnerships. Many problems that
chemistry aims to solve are international and will

EXTENSION
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require globally implemented solutions.

This feature highlights information in the book that

is extension content and is not part of the syllabus.

EXAM-STYLE QUESTIONS
REFLECTION
The questions appear at the end of each chapter.
The purpose is for you as a learner to reflect
on the development of your skills proficiency
and your progress against the objectives. The
reflection questions are intended to encourage
your critical thinking and inquiry-based learning.
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Exam-style questions at the end of each topic provide essential practice and self-assessment. These are signposted
in the print coursebook and can be found in the digital version of the coursebook.

SELF-ASSESSMENT CHECKLIST
These appear at the end of each Chapter as a series of statements. You might find it helpful to rate how confident
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you are for each of these statements when you are revising. You should re-visit any topics that you rated ‘Needs
more work’ or ‘Almost there’.

Needs Almost Confident

I can Section
more work there to move on

Free online material

Additional material to support the Chemistry for the IB Diploma course is available online.
This includes Assessment guidance – a dedicated chapter in the digital coursebook helps teachers and students unpack
the new assessment and model exam specimen papers. Additionally, answers to the Test your understanding and
Exam-style questions are also available.
Visit Cambridge GO and register to access these resources at www.cambridge.org/GO.

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Unit 1

The nature

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of matter PL
INTRODUCTION
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We have all heard of atoms: the particles from which everything is made. Democritus and his teacher,
Leucippus, fifth-century BCE Greek philosophers, are usually credited with first suggesting the idea of the
atom as the smallest indivisible particle from which all matter is made, but the modern understanding of
science in terms of atoms only really began in the 19th century with the work of John Dalton (1766–1844).
An understanding of atoms and atomic structure is now regarded as fundamental to chemistry, but we
usually talk about atomic theory, so does that mean that atoms may not really exist? We will not look
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specifically at the evidence for the existence of atoms, but does the fact that everything in this book and
other scientific literature is explained by assuming the existence of atoms provide that evidence?
So, assuming that atomic theory is the best way of explaining the world around us, what do we know about
atoms? Atoms are most definitely small – there are many more atoms in a drop of water than there are stars
in the Milky Way, and there are probably more atoms in a glass of water than there are stars in the universe
(although no one is sure how many stars there are in the universe). We know that there are different types of
atoms, but how many are there? A simple answer would be as many as there are elements, but there are also
isotopes, and which isotope we are talking about can make a big difference to the properties of the element
and to the world – a country with a storage vault full of uranium-235 (which can be used for making nuclear
weapons) will be viewed very differently by other governments from one with uranium-238!

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Chapter 1

The particulate

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nature of matter PL
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LEARNING OBJECTIVES
In this chapter you will:
• understand the terms element, compound and mixture
• understand the differences between heterogeneous and homogeneous mixtures
• understand how to separate the components of a mixture
• use kinetic molecular theory to understand the properties of solids, liquids and gases
• understand that temperature in K is proportional to the average kinetic energy of particles
• understand how to convert temperatures between K and °C
• use state symbols in chemical equations
• use kinetic molecular theory to explain changes of state.

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

An element can be defined in different ways and, for the

GUIDING QUESTIONS moment, we will define it in terms of its properties:
• What are the differences between elements,
compounds and mixtures? KEY POINT
• How can the components of a mixture An element is a chemical substance that cannot
be separated? be broken down into a simpler substance by
chemical means.
• How can kinetic molecular theory be used
to explain the properties of solids, liquids
and gases?
Gold only contains gold atoms and sulfur only contains

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sulfur atoms, and because of this, these cannot be
broken down into anything simpler than gold atoms or
Introduction sulfur atoms using chemical reactions.

The song ‘Woodstock’, released in 1970, includes the In Chapter 2, we will look at the structure of atoms, and
words ‘we are stardust’ and, strangely enough, this is pretty this will allow us to define an element in terms of the
particles that make up the atom:

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much true. The lightest elements (mostly hydrogen and
helium with some lithium) were formed in the immediate
aftermath of the Big Bang, but the other elements that
we, and everything around us, are made of were formed
in stars. In this chapter, we will look at the distinction
between elements, compounds and mixtures, explain their
properties in terms of kinetic molecular theory and look
at how to separate the components of mixtures. The
distinction between elements, compounds and mixtures
is fundamental to an understanding of chemistry and,
although in subsequent chapters we will mention very little
KEY POINT
An element is a pure substance in which each
atom has the same number of protons in the
nucleus (see Chapter 2).

So, for example, gold is an element and all samples of

pure gold contain only atoms that have 79 protons in
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about mixtures, it is important to remember that most the nucleus.
substances in everyday life are actually mixtures. The symbols for elements are shown in the
periodic table (Chapter 10, section 10.1). In a sample
of an element, the atoms may be present as individual
1.1 Elements, atoms (e.g. helium, He), be chemically bonded as
individual molecules (e.g. oxygen, O2, or ozone, O3) or be

compounds and chemically bonded as part of a giant structure (e.g. gold,

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Au, or carbon, C). Some representations of elements are

shown in Figure 1.1. The key thing to notice is that, in
mixtures each part of Figure 1.1, all the atoms are the same.

Elements a b c d

Elements are the primary constituents of matter. There

are 118 elements that have been discovered so far, and
these are shown in the periodic table. Of these, about
90 occur naturally in reasonable amounts, and the rest Figure 1.1: Some elements. a This could be a noble gas,
are present in only trace amounts or are artificially such as helium, which consists of just individual atoms.
made. By far the most abundant element in the universe b This could be gaseous oxygen, consisting of O2
is hydrogen, followed by helium, but in the Earth’s crust molecules, in which the oxygen atoms are chemically
oxygen is the most abundant and astatine is the least bonded to each other. c This could be a metal, such as
abundant. Astatine has no stable isotopes and scientists gold. d This could be carbon – the lines represent bonds
estimate that, at any one time, there is probably less than between atoms.
30 g present in the whole of the Earth’s crust.

13
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

THEORY OF KNOWLEDGE KEY POINT

What is an element? The elements in a compound are chemically
The concept of an element is fundamental to combined, and therefore, compounds can
the study of chemistry, but, strangely enough, only be converted into their elements again by
chemists do not necessarily agree on the chemical reactions.
definition of an element. If we say that oxygen
is an element, that is fine, but do we mean an
For example, hydrogen can be obtained from water by
O atom, a sample of oxygen gas, which contains
reacting it with sodium, or hydrogen and oxygen could
O2 molecules, or even ozone, which contains
both be produced by electrolysis (passing electricity
O3 molecules?
through water).

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The physical properties and chemical properties of a
Chemistry is partly a study of how chemical elements compound are different from those of the elements from
combine to make the world and the universe around which it is formed.
us. When different elements combine chemically, they
form compounds.
KEY POINT

Compounds

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In water (H2O), a compound, there are always exactly
twice as many hydrogen atoms as oxygen atoms – this
ratio never varies for a particular compound. If the ratio
is different, it is a different compound, for example,
if the ratio is 1:1, the compound is hydrogen peroxide
(H2O2) and not water.
Chemical properties how a substance behaves
in a chemical reaction.
Physical properties all the other properties
of a substance, such as melting point, density,
hardness and electrical conductivity.

For example, hydrogen an explosive gas, combines with

oxygen, a highly reactive gas, to form water, which
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KEY POINT is a liquid at room temperature. Water reacts in very
different ways to hydrogen and oxygen (it has different
A compound is a pure substance formed when chemical properties) – you would not try to put a fire
two or more elements combine chemically in a out with hydrogen or oxygen!
fixed ratio.
Similarly, when sodium (a highly reactive metal) is heated
with chlorine (a toxic gas), a white, crystalline substance,
The atoms (or ions) in a compound are chemically sodium chloride (common salt), is formed, which reacts
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bonded to each other – this may be covalent bonding in very different ways to sodium and chlorine.
(see Chapter 7, section 7.1) or ionic bonding (see
Chapter 6, section 6.1). Some representations of the
structures of compounds are shown in Figure 1.2.
Mixtures
Sometimes the chemical bonds (lines) will be shown (two Elements and compounds are pure substances, but most
of the structures of water) and sometimes not. things around us are not pure, they are mixtures. We
breathe in air, which is a mixture; all the foods we eat are
a b
mixtures; oxygen is carried around our body by blood,
O another mixture.
H H
The components of a mixture can be elements or
compounds – or even mixtures! Air is a mixture of mostly
Figure 1.2: Some representations of compounds. The key elements (nitrogen, oxygen, argon) with smaller amounts
thing to notice here is that more than one type of atom is of compounds (carbon dioxide, water vapour etc.).
present in each structure. a Three different ways of showing Representations of mixtures are shown in Figure 1.3.
the structure of water. b An ionic compound, such as
sodium chloride.

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1 The particulate nature of matter

a b chloride in 100 cm3 water, 10 g of sodium chloride in

100 cm3 of water etc., up to the limit of solubility (how
much dissolves at a certain temperature).

Link
Figure 1.3: Some mixtures. a A mixture of gases. b An alloy.
Alloys are mixtures of metals with other metals (or
non-metals). In alloys, there is metallic bonding
KEY POINT throughout the structure (see Chapter 8, section 8.1), so
the components of the mixture are actually chemically
The components of a mixture are not chemically bonded to each other. An alloy is, however, still regarded
bonded together, so they retain their as a mixture because it will not have a fixed composition –

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individual properties. the metals can be mixed together in various proportions.

Mixtures can be homogeneous

In a mixture of iron and sulfur, the two elements retain or heterogeneous
their individual chemical and physical properties, so
iron is magnetic and will react with dilute acids to form

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hydrogen gas, and sulfur is yellow and burns in air to
form sulfur dioxide. When the mixture is heated and
forms the compound iron sulfide, this has a different
appearance, is not magnetic (Figure 1.4) and, for
example, reacts with acids to form the extremely smelly
and toxic gas hydrogen sulfide – the compound has
different properties to its elements.
KEY POINT
A homogeneous mixture has the same (uniform)
composition throughout the mixture – it consists
of only one phase.
Solutions and mixtures of gases are
homogeneous mixtures.
A heterogeneous mixture does not have uniform
composition – it consists of separate phases.
M
The term phase can be used in different ways in
chemistry; here, it refers to a region that is the same
throughout, in terms of chemical composition and
physical properties. In a heterogeneous mixture, there
will be distinct boundaries between different phases.
Figure 1.4: The iron in a mixture of iron and sulfur (left)
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retains its magnetic properties, but iron sulfide (right) is An example of a homogeneous mixture is a solution.
not magnetic. The concentration is the same throughout: if several
1cm3 samples of a solution of sodium chloride are taken
from a beaker and evaporated separately to dryness, the
KEY POINT same mass of solid sodium chloride will be recovered
from each sample.
The components of a mixture can be mixed
together in any proportion. One example of a heterogeneous mixture is sand in a
sample of water. Sand and water can be distinguished
from each other – they are separate phases. Other
When atoms combine to form compounds, they do so in examples include milk and orange juice. Orange juice
fixed ratios (according to the formula of the compound), is a complex mixture, containing an aqueous phase
but there are no such limitations on making a mixture, with various substances dissolved or suspended in it.
and iron and sulfur can be mixed together in absolutely Suspended material in orange juice includes cellulose,
any proportions. Solutions are mixtures, and a solution proteins, lipids and pectins. If you leave some freshly
of sodium chloride could be made by dissolving 1 g squeezed orange juice to stand, some parts will settle out,
of sodium chloride in 100 cm3 of water, 2 g of sodium but will it become completely clear? Milk, as a colloid, is
discussed in detail in the Science in Context section below.

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

Mixtures of solids are always heterogeneous mixtures. When looking at mixtures that are liquids or gases, if
For example, a mixture of iron and sulfur is a the mixture is clear, so that you can see through it, it is
heterogeneous mixture. Even though the mixture may a homogeneous mixture; if it is cloudy/opaque, so that
have been made very carefully, so that there are the same some/all of the light is scattered as it passes through it,
masses of iron and sulfur in each cubic centimetre, the then the mixture is heterogeneous.
composition is not uniform because there are distinct
particles of iron and sulfur (you may need to use a
magnifying glass to see them), and each particle of iron
and sulfur represents a different phase.

SCIENCE IN CONTEXT

E
Solutions and mixtures
Tea or coffee without milk are solutions. This is usually
easier to see with tea, but, if you dilute your black
coffee in a glass cup with some water, you will be able
to see that, although it is coloured, it is clear, so that

PL
light passes through it without being scattered, and
therefore, it is a solution and a homogeneous mixture
(although, if you used a cafetiere or a not-very-good
filter, you may still have a few coffee grounds in it,
which would make it a heterogeneous mixture!). If you
add sugar and stir it well, the sugar dissolves, and so,
you still have a homogeneous mixture; however,
if you add milk (Figure 1.5), your coffee goes
cloudy – this is now a heterogeneous mixture. Milk is
a type of mixture called a colloid (or colloidal system)
and contains very small droplets of fat and solid
Figure 1.5: White coffee and doughnuts are
heterogeneous mixtures.

To consider:
1 Other heterogeneous mixtures you will come
M
protein particles dispersed throughout an aqueous across in a coffee shop include whipped cream,
phase. These particles scatter light (the Tyndall effect) hot chocolate and muffins… can you think
and, therefore, white coffee is not clear but opaque. of any more?
2 How do the methods for separating
homogeneous mixtures differ from methods for
separating heterogeneous mixtures?
SA

3 Is it possible to separate all the components

from white coffee or doughnuts?

TEST YOUR UNDERSTANDING

1 Classify each of the following as an element, d vanadium
a compound or a mixture:
e ammonia
a water
f air
b oxygen
g hydrogen chloride
c potassium iodide
h magnesium oxide.

16
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

CONTINUED
2 Classify each of the diagrams shown as an a a mixture of carbon dioxide gas and
element, compound or mixture: helium gas
a b c d e b a mixture of solid copper(II) oxide and solid
calcium carbonate
c potassium hydroxide solution
d mayonnaise.
3 Classify each of the following as a
heterogeneous or a homogeneous mixture:

E
Separating the components of a mixture The equation for the reaction is
CuO((ss))++ H
CuO SO44((aq
H22SO aq))→ CuSO44((aq
→ CuSO aq))++ H O((ll))
H22O
The components of a mixture can be separated from
each other by physical processes – physical processes
Link
involve chemical reactions.

Filtration

PL
are things like filtration and distillation, which do not

In a chemistry laboratory, filtration is usually used to

separate an insoluble solid from a liquid. It can also be
used to separate a solid from a gas. The apparatus most
often used for filtration is shown in Figure 1.6. The solid
left in the filter paper is called the residue, and the liquid
that passes through the filter paper is called the filtrate. The
filter paper acts as a physical barrier to the pieces of solid
Copper(II) oxide is a base and reacts with sulfuric acid
in a neutralisation reaction (Chapter 19).

INTERNATIONAL MINDEDNESS
Diesel engines
Diesel engines are used extensively in heavy-duty
commercial vehicles, such as lorries and buses,
as well as cars. One of the major environmental
M
but allows the liquid to pass through gaps between fibres. concerns with the use of vehicles with diesel
filter paper engines is that they emit significantly more
particulate matter (soot) than gasoline (petrol)
engines, and this can be damaging to health.
residue filter funnel Diesel engines are, therefore, fitted with particulate
filters, to filter out as much of the particulate matter
as possible. Different countries have different
SA

regulations on emissions from diesel vehicles.

filtrate

Figure 1.6: Filtration can be used to separate a solid from Evaporation

a liquid. Evaporation can be used to remove a solvent from a
solution to leave the solute. If a solution of sodium
Filtration is used as part of the process in the chloride is heated, water will evaporate/boil off to leave
preparation of copper(II) sulfate. Copper(II) sulfate solid sodium chloride (Figure 1.7).
solution can be made by the reaction between copper(II)
oxide (a black solid that is insoluble in water) and dilute
sulfuric acid. Excess (more than enough to react with
all the sulfuric acid) copper(II) oxide is added to hot
sulfuric acid. The excess copper(II) oxide is then filtered
off. In this case, copper(II) oxide is the residue and
copper(II) sulfate solution is the filtrate.

17
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

This technique is called solvent extraction, and we talk

about a solute partitioning itself between two solvents.

evaporating solution
dish separatory funnel

aqueous layer (less caffeine)

dichloromethane layer (more caffeine)

Figure 1.7: Evaporation of the solvent can be used Figure 1.8: A separatory funnel is used in the extraction of

E
to obtain a solute from a solution. If larger crystals are caffeine from tea.
required, only some of the water should be boiled off and
then the solution should be left to crystallise.
Link
Solvation Caffeine is more soluble in dichloromethane than in

PL
Solvation can be used to separate a mixture of two or
more substances, due to differences in solubility.
For example, a mixture of solid copper(II) oxide
(insoluble in water) and sodium chloride (soluble in
water) can be separated by putting the mixture into a
beaker of warm distilled/deionised water and stirring
to make sure that all the sodium chloride has dissolved.
The mixture is filtered: copper(II) oxide is the residue
and sodium chloride solution is the filtrate. Copper(II)
oxide is washed with distilled water to remove any
water. A general rule for solubility is ‘like dissolves
like’. The intermolecular forces are more similar
between caffeine (dipole–dipole interactions) and
dichloromethane (dipole–dipole interactions) than
between caffeine (dipole–dipole interactions) and
water (hydrogen bonds). Intermolecular forces will be
discussed in Chapter 7.

EXAM TIP
M
traces of sodium chloride solution and then dried in a The word solvation is used here because it is
warm oven (distilled water will evaporate). Solid sodium used on the IB syllabus, but it is not actually the
chloride can be obtained from the solution by heating it correct word. Solvation will be discussed further
in an evaporating dish until all the water evaporates. in Chapter 7. The process here is probably best
described as dissolving.
Note that distilled/deionised water must be used because
tap water contains dissolved solids and, when heated,
will leave a residue of these solids, so that the copper(II)
SA

oxide and sodium chloride obtained will not be pure.

Distillation
Distillation could be used, for example, to separate water
Application of solvation to the extraction of caffeine from a sodium chloride solution. The sodium chloride
A common laboratory experiment is the extraction of solution is heated, water evaporates and condenses
caffeine from tea. The basic principles of the technique again in the condenser, so that it can be collected in the
are that tea leaves are boiled with water to make an collection vessel (it is called the distillate).
aqueous solution, which is shaken with dichloromethane Suitable apparatus for distillation is shown in Figure 1.9.
(an organic liquid with the formula CH2Cl2) in a
If heating is continued for a long enough time, only
separatory funnel (Figure 1.8). Dichloromethane is
solid sodium chloride will be left in the round-bottomed
not soluble in water and remains as a separate phase
flask, and all the water will be in the collection vessel.
in the separatory funnel. Caffeine is more soluble in
dichloromethane than in water and distributes itself The difference between distillation and evaporation
between the water layer and the dichloromethane layer, is that, in distillation, the solvent is boiled off, but
with much more in the dichloromethane layer. The then condensed again, so that it can be collected. So,
dichloromethane layer can then be run off and the evaporation would generally be used when it is the
solvent evaporated to leave caffeine (a white solid). solute that is the desired product and distillation when it
is the solvent that is required.

18
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

water out

condenser

E
water in

solution (e.g.
sodium chloride

PL heat

Figure 1.9: The experimental set-up for distillation.

KEY POINT
solution)
distillate

Distillation can also be used to separate a mixture of

two liquids, as long as there is a large enough difference
between their boiling points (about 70 °C).
M
Distillation (simple distillation) can be used to
separate the solute and solvent from a solution The liquid with the lower boiling point (more volatile)
(where the solute was a solid) or to separate a will go into the vapour phase more easily and will be
mixture of two liquids with sufficiently different collected in the collection vessel (Figure 1.9), whereas
boiling points. the liquid with the higher boiling point will be left in the
round-bottomed flask. If the boiling points of the two
liquids are too close, then complete separation will not
SA

INTERNATIONAL MINDEDNESS be obtained, and a mixture will distil over.

This technique is used extensively in organic chemistry
Fresh water
for extracting the more volatile liquid product of a
Seawater is a mixture, and distillation can be reaction from the reaction mixture and for purifying a
used to obtain water without salt from seawater. liquid product of a reaction. When used for purification,
The process of removing salt from seawater is the pure liquid is collected in the collection vessel,
called desalination. Desalination is very important and any non-volatile impurities will be left in the
in some parts of the world, where sufficient round-bottomed flask. The purity of the liquid could
freshwater is not available, ‘for example, in be tested by using chromatography (see the Paper
parts of Southwest Asia and North Africa. Water chromatography section below).
obtained by desalination can be used for human
consumption, agriculture or in industry.

19
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

Paper chromatography At the simplest level, the number of spots present on

Paper chromatography may be used, for example, to a chromatogram indicates the number of components
separate the various dyes in coloured inks, to separate of the mixture (although other tests might need to be
a mixture of sugars or amino acids or to test the purity done, to check whether a particular spot is indeed a
of a substance. The experimental set-up for paper pure substance).
chromatography is shown in Figure 1.10.
closed container closed container KEY POINT
Chromatography can be used to test the purity
solvent front
of the product of a reaction. The presence of
chromatography
paper
more than one spot indicates that the substance
is impure.

E
components
of mixture
It can take quite a bit of research and trial and error to
find a suitable solvent for chromatography that provides
good separation of the components of the mixture. The
polarity (see Chapter 7) of the substances influences
spot of
mixture

•
pencil
line

solvent

PL
Figure 1.10: A paper chromatography experiment.
pencil line

solvent

The process of the solvent travelling up the paper to

produce a chromatogram is called development.

To carry out a paper chromatography experiment:

A line is drawn with a pencil (not a pen, as the
the choice of solvent, but there are also other factors
involved. The solvent does not have to be a pure liquid,
and very often mixtures are used.
All chromatography techniques involve a stationary
phase and a mobile phase. The components in a mixture
are separated because of their differences in affinity for
the stationary and mobile phases. This is explained in
Chapter 7.

Location of spots
M
If the substances to be separated are colourless (e.g.
inks may move with the solvent) across a piece amino acids or sugars), then some method must be
of chromatography paper about 1 cm from used to locate the spots on the paper or TLC plate. The
the bottom. spots may be located using a locating agent. Amino
• A sample of the mixture is placed on the pencil line acids, which are colourless, may be located by spraying
and allowed to dry. with ninhydrin, which makes them show up as pink
or purple spots. Other methods that are useful for
• The paper is suspended in a container with a small
SA

organic solutes are exposing the paper or plate to iodine

amount of solvent at the bottom, so that the end vapour (the spots become brown) or spraying the plate
of the paper dips into the solvent (the original with concentrated sulfuric acid then heating it (the
sample spot must be above the top of the solvent; spots appear as brown–black). Spots may also often
otherwise it will just dissolve into the solvent). be located by the use of an ultraviolet lamp, as some
• The container is closed, so that the atmosphere substances fluoresce under ultraviolet light.
becomes saturated with the solvent – this prevents
evaporation of the solvent from the surface of
the paper.
• The solvent is drawn up the paper by capillary action.
• The process is stopped when the solvent front is
about 1 cm from the top of the paper. A pencil line
is drawn to record the position of the solvent front
and the paper is dried.

20
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

TEST YOUR UNDERSTANDING

4 Select a technique that could be used to separate the components of the following mixtures:
a sand from water
b potassium chloride from a potassium chloride solution
c different indicators in universal indicator solution
d a mixture of ethoxyethane (CH3CH2OCH2CH3, boiling point 34 °C) and 1-(hexyloxy)hexane
(CH3(CH2)5O(CH2)5CH3, boiling point 220 °C).

E
5 Explain how you would separate a mixture of potassium bromide (soluble in water) and calcium
carbonate (insoluble in water).
6 Explain how you would separate iodine from an aqueous iodine solution, given that iodine is much more
soluble in hexane than in water, and hexane is immiscible with water. Immiscible means that hexane and
water do not mix – they form separate layers.

1.2 Kinetic molecular

theory PL
Kinetic molecular theory (often just called kinetic theory)
is a model that was developed originally to explain the
properties of gases, but it is usually also extended to
describe liquids and solids. Within this model, we describe
solid liquid

Figure 1.11: The three states of matter.

EXAM TIP
gas
M
all matter as being made up of individual particles that
are in constant motion (hence, the word ‘kinetic’). In diagrams showing the states of matter,
remember the following:
NATURE OF SCIENCE • Solid: the particles should be arranged
regularly and touching.
Models are used throughout science. A model
is a way of making sense of the world around • Liquid: the particles are arranged randomly
SA

us. Models may either be qualitative, as here, or but still mostly touching.
quantitative (involving numbers and equations). • Gas: the particles are arranged randomly
The validity of a particular model can be tested and are shown far apart.
by looking at how closely predictions made using
the model agree with experimental observations.
Solid: the particles are generally regularly arranged and,
due to relatively strong forces of attraction between
The three states of matter most commonly encountered them, are only able to vibrate about mean positions.
are solid, liquid and gas, and these differ in terms of the As the forces between the particles are relatively strong,
arrangement and movement of particles. The particles solids have fixed shapes.
making up a substance may be individual atoms or Liquid: the particles have weaker forces between them,
molecules or ions. Simple diagrams of the three states of and so are able to move around each other. As the forces
matter are shown in Figure 1.11, in which the individual between the particles are weaker than those in solids,
particles are represented by spheres. liquids take the shape of the container they are in. There
are, however, still forces between the particles, so they
stay together and do not fill the container.

21
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

Gas: the particles are assumed to have no forces between of molecules in air (mostly nitrogen and oxygen) at 25 °C
them (see Chapter 5) and move around randomly in all and atmospheric pressure is almost 500 ms−1; the particles
directions. There are no forces between the particles, so collide, on average, every 150 ps (1.5 × 10−10 s) and only
they are free to move around anywhere in a container, and travel about 7 × 10−8 m between collisions.
thus, ‘fill’ the container. To give you some idea of how The properties of the three states of matter are
quickly the particles in a gas are moving: the average speed summarised in Table 1.1.

Solids Liquids Gases

Distance between particles close together close but further apart than in solids far apart
Arrangement regular random random

E
Shape fixed shape no fixed shape; no fixed shape;
take the shape of the container fill the container
Volume fixed fixed not fixed
Movement vibrate move around each other move around in all directions
Speed of movement slowest faster fastest
Energy
Forces of attraction

PL
lowest
strongest
Table 1.1: Properties of the three states of matter.

Temperature
There are two temperature scales that are used
commonly in everyday life: the Fahrenheit scale
(melting point of ice = 32 °F and boiling point of
higher
weaker
highest
weakest

The absolute, or Kelvin, scale of temperature starts at

absolute zero, which is the lowest temperature possible.
All molecular motion does not actually stop at absolute
zero (this would contravene the Heisenberg uncertainty
principle), but it is the temperature at which everything
M
water = 212 °F), which is used predominantly in the would be in its lowest energy state. It is not possible to
USA and a few other countries, and the Celsius or actually reach absolute zero, but scientists have managed
centigrade scale (melting point of ice = 0 °C and boiling to get very close – below one nanokelvin!
point of water = 100 °C), which is used in the rest of the
world. In science, however, we much more commonly KEY POINT
use the absolute, or Kelvin, scale of temperature. For
calculations involving temperatures in science, it is Absolute zero the lowest temperature possible,
SA

usually essential to use temperatures in kelvin. corresponds to 0 K or −273.15 °C (usually taken

as −273 °C).
KEY POINT A change of 1 °C is the same as a change of 1 K.
The kelvin is the SI unit of temperature.

EXAM TIP
SI stands for Système International and is the
internationally accepted system of units used in science. A temperature change in °C is the same as one
Within the SI, seven base units are defined by reference in K.
to seven fundamental constants (such as the speed of
light in a vacuum and the Planck constant), which have
agreed specific values. Other SI base units include the
second, the metre and the mole.

22
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

nitrogen ammonia water The particles in gases and liquids are constantly colliding
boiling point melting point boiling point and, therefore, the particles will not all be moving at the
same speed, and there will be a spread of kinetic energies
0 77 195 273 373 temperature in K
for the particles, which is why we use the term average
kinetic energy. The distribution of kinetic energies in
–273 –196 –76 0 100 temperature in °C
a sample of gas at two different temperatures is shown
absolute water (ice) in Figure 1.13. At higher temperature, there are fewer
zero melting point particles with lower kinetic energy and more particles with
higher kinetic energy, and so, the average kinetic energy of
Figure 1.12 Some temperatures in K and °C. the particles is greater. This will be explored in Chapter 17.

lower
Figure 1.12 compares some temperatures in K and °C.

E
temperature higher
The fact that a change of 1 °C is the same as a change temperature

Fraction of particles
of 1 K makes it quite straightforward to convert
temperatures between the two scales.

KEY POINTS
To convert °C into K, add 273.
To convert K into °C, subtract 273.

WORKED EXAMPLE 1.1

Convert a temperature of 25 °C into kelvin.
Answer
PL 0
0 Kinetic energy

Figure 1.13: The distribution of kinetic energies in a sample

of gas is called the Maxwell–Boltzmann distribution.

If two gases are at the same temperature, their particles

will have the same average kinetic energy. This does not
M
To do this, we add 273 to the temperature in °C:
25 + 273 = 298 K mean that the average speed of the particles is the same.
Kinetic energy is calculated using the following equation:
Ek = 1 2 mv 2
WORKED EXAMPLE 1.2
where m is the mass of the particle and v is the speed.
Convert a temperature of 350 K into °C. This means that, the lighter the particles, the higher the
SA

Answer average speed at a particular temperature. The average

To do this, subtract 273 from the temperature in K: speed of carbon dioxide molecules (relative mass 44.01)
350 − 273 = 77 °C at 25 °C is about 380 m s−1, whereas the average speed of
hydrogen molecules (relative mass 2.02) is about 1770 m s−1.

1.3 Temperature and TEST YOUR UNDERSTANDING

kinetic energy 7 Convert each of the following temperatures

in °C to temperatures in K.
a 25 °C
KEY POINT
b 500 °C
Temperature is a measure of the average (mean)
kinetic energy (Ek) of the particles in a substance. c −100 °C
• The higher the temperature, the higher the d −145 °C
average kinetic energy of the particles.

23
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

CONTINUED KEY POINT

8 Convert each of the following temperatures A substance will be:
in K to temperatures in °C.
• a solid if the temperature is below its melting
a 323 K point
b 100 K • a liquid if the temperature is between its
melting point and its boiling point
c 50 K
• a gas if the temperature is above its boiling
d 500 K point

E
9 What is wrong with the temperatures −50 K
and −300 °C?
So, for example, bromine melts at −7.2 °C and boils
at 58.8 °C; therefore, below −7.2 °C bromine will be a
solid, between −7.2 °C and 58.8 °C it will be a liquid,

1.4 Changes of state and above 58.8 °C it will be a gas. There is no universally

PL
When one state of matter becomes another state of
matter, we describe this as a change of state. Changes
of state are summarised in Figure 1.14. Converting one
state of matter into another usually involves heating (the
change of state is an endothermic process) or cooling
the substance (the change of state is an exothermic
process) but can also be achieved by changing pressure.
Endothermic and exothermic processes will be
considered in Chapter 12.
accepted definition of ‘room temperature’, but it is
often taken as 25 °C, and so bromine is one of only two
elements that is a liquid at room temperature.
Boiling and evaporation both involve a change in state
from liquid to gas, but they are not the same thing –
boiling only occurs at a certain temperature (the boiling
point), but evaporation of the liquid can occur at any
temperature between the melting and boiling points.

Using state symbols in equations

M
heating – energy is supplied
particles gain energy State symbols are used in chemical equations to indicate
the physical state that the substances are in.
sublimation deposition
boiling
solid melting liquid evaporating gas
KEY POINT
freezing condensing The state symbols are:
SA

cooling – energy taken out (s) = solid

particles lose energy
(l) = liquid
Figure 1.14: Changes of state. Note that evaporation can
occur at any temperature, but boiling occurs at a fixed (g) = gas
temperature.
(aq) = aqueous (dissolved in water)

Sublimation is the change of state when a substance

goes directly from the solid state to the gaseous state, We saw the following chemical equation
without going through the liquid state. Both iodine and in the previous section on Filtration
solid carbon dioxide (dry ice) sublime at atmospheric CuO ( s ) + H2SO4 ( aq ) → CuSO4 ( aq ) + H2O ( l )
pressure. The reverse process is called deposition.
This indicates that solid copper(II) oxide (CuO)
The temperatures at which a substance changes state are reacts with sulfuric acid, which is an aqueous solution
called its melting point (change from solid to liquid) and (dissolved in water) to form an aqueous solution of
boiling point (change from liquid to gas). copper(II) sulfate and liquid water.

24
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

Changes of state may be described using equations The average kinetic energy of the particles increases,
including state symbols, for example: until the boiling point of the liquid is reached. At this
point (80 °C), the continued supply of heat energy
Melting of ice to form water: H2O ( s ) → H2O ( l )
is used to overcome the forces of attraction between
Boiling/evaporation of liquid bromine: Br2 ( l ) → Br2 ( g ) the particles completely and the temperature of the
Sublimation of iodine: I2 ( s ) → I2 ( g )
substance remains constant, until all the liquid has been
converted into gas. The continued supply of heat energy
How to write balanced chemical equations will be then increases the average kinetic energy of the particles
discussed in Chapter 16. and, therefore, the temperature of the gas. The particles
move around faster and faster, as the temperature of the
gas increases.
Temperature during changes

E
of state TEST YOUR UNDERSTANDING
If a pure substance is heated slowly, from below its 10 State the names of the following changes
melting point to above its boiling point, a graph of of state:
temperature against time can be obtained (Figure 1.15).
a from solid to liquid
100
90
80
70
boiling
point = 80 °C

melting

PL
boiling gas

liquid and gas

b
c
from solid to gas
from gas to liquid
Temperature / °C

point = 50 °C liquid present

60 d from gas to solid.
melting
50
40
11 Use data in the table to determine whether
solid liquid and solid each of the elements will be a solid, liquid or
30
present gas at the specified temperature:
20
10
Substance Melting point Boiling point
M
0 / °C / °C
0 5 10 15 20 25 30 35
Time/minutes Magnesium 650 1090
Figure 1.15: A heating curve showing changes of state. Fluorine −220 −188
Polonium 254 962
As a solid is heated, its particles vibrate more violently. Mercury −39 357
The particles gain kinetic energy and the temperature of
SA

the solid rises. At 50 °C, the solid in Figure 1.15 begins a magnesium at 100 °C
to melt – at this stage, there is solid and liquid present
together, and the temperature remains constant until all b fluorine at −200 °C
the solid has melted. All the heat energy being supplied c polonium at 1000 °C
is used to partially overcome the forces of attraction
between particles, so that they can move around d mercury at 25 °C.
each other. Another way of saying this is that, at the
melting point, all the heat energy being supplied goes
into increasing the potential energy of the substance
(overcoming forces between particles) and not to
increasing the kinetic energy, so the temperature does
not change.
When all the solid has melted, the continued supply of
heat energy causes the kinetic energy of the particles to
increase, so that the particles in the liquid move around
each other more quickly and the temperature increases.

25
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

SELF-ASSESSMENT CHECKLIST
Think about the topics covered in this chapter. Which parts are you most confident with?
Which topics require some extra practice?

Needs Nearly Confident

I can... Section
more work there to move on
explain the terms element, compound and mixture, and
1.1
distinguish between them
explain the difference between heterogeneous and
1.1

E
homogeneous mixtures and give examples of each
explain the different methods for separating the
components of a mixture and suggest a suitable method 1.1
for separating a particular mixture
explain the properties of solids, liquids and gases in

PL
terms of kinetic molecular theory
state the relationship between temperature in K and the
average kinetic energy of particles
convert temperatures between K and °C
use state symbols in chemical equations
explain changes of state in terms of kinetic
molecular theory.
1.2

1.3

1.3
1.4

1.4
M
REFLECTION
To what extent do you feel that you have met many of the ideas in this chapter before? Can you highlight
specific areas that are new to you? Are you confident with these areas? Can you use your knowledge to
identify heterogeneous and homogeneous mixtures around your home or school? Do you think that you
could explain the difference between elements, compounds and mixtures to another student?
SA

EXAM-STYLE QUESTIONS
You can find questions in the style of IB exams in the digital coursebook.

26
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The
CHEMISTRY FOR THE IB DIPLOMA: EXAM-STYLE particulate nature of matter
QUESTIONS

Exam-style questions
A periodic table is required to answer some of these questions.
The multiple-choice questions can all be answered without the aid of a calculator.

Chapter 1

E
1 Which of the following contains an element, a compound and a mixture? [1]
A H2O(l), H2(g), FeS(s)
B Cl2(aq), Br2(g), NaBr(l)
C CH4(g), I2(l), CO2(l)
D NaCl(aq), CO(g), H2S(g)
2

3
Which of the following is a homogeneous mixture?
A a mixture of sand and sodium chloride
B a sodium chloride solution
C a mixture of hexane and water
D a mixture of sulfur and iron
Consider the following process: I2 (g) → I2 (s)
The name of this process is
A condensation
B sublimation
PL [1]

[1]
M
C deposition
D vaporisation
4 A substance, X, which is a solid at room temperature, is heated and the temperature monitored.
The graph of temperature against time is shown.
100
90 D
SA

C
80
70
Temperature / °C

B
60
A
50
40
30
20
10
0
0 5 10 15 20 25 30 35
Time / minutes

At which point are a solid and a liquid present? [1]

A
B
C
D

Chemistry for the IB Diploma – Owen © Cambridge University Press 2023 1 27

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

EXAM-STYLE QUESTIONS

5 Rubidium has a melting point of 39 °C and a boiling point of 688 °C.

What are the melting and boiling points of rubidium in kelvin? [1]

Melting point / K Boiling point / K

A −234 415
B 234 415
C 312 961
D 39 688

6 In which of the following is the temperature in K higher than the temperature in °C? [1]

E
A 100 °C 250 K
B 150 °C 500 K
C −100 °C 100 K
D 0 °C 250 K
7
lead(II) nitrate.

PL
Lead(II) iodide (PbI2) can be made by adding a solution of potassium iodide to a solution of

The equation for the reaction is:

2KI(aq) + Pb(NO3 )2 (aq) → PbI2 (s) + 2KNO3 (aq)
Which method could be used to most easily separate lead iodide from the reaction mixture?
A distillation
B filtration
C evaporation
D solvation
[1]
M
8 Spots of four substances were put on the baseline of a piece of chromatography paper
in the positions marked with an × in the diagram. The resulting chromatogram is shown.

solvent front
SA

A B C D

Which of the substances is pure? [1]

A C
B D

28
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate
CHEMISTRY FOR THE IB DIPLOMA: EXAM-STYLE nature of matter
QUESTIONS

9 Ethanol (boiling point 78 °C) is a very good solvent for many organic compounds.
Which liquid is most likely to be completely separated from an ethanol solution using distillation? [1]
A hexane (boiling point 69 °C)
B 3-ethylpentane (boiling point 93 °C)
C propan-2-ol (boiling point 82 °C)
D propane-1,2,3-triol (boiling point 289 °C)
10 Which of the following statements about kinetic molecular theory is correct? [1]
A In gases, the particles vibrate about mean positions.
B In liquids, there are no forces between particles.

E
C The particles in a gas all have the same kinetic energy.
D The average kinetic energy of particles increases as temperature increases.
11 The melting and boiling points of some substance are shown in the table. [1]

Melting point / K Boiling point / K

a
b
ethyl benzoate (C9H10O2)
anthracene (C14H10 )
propanone (C3H6O)
ethene (C2H4)
PL
238
489
178
104

Explain which substances are liquids at 25 °C.

486
614
329
169

Ethyl benzoate is soluble in propanone. Molly is given a solution containing 10 g of

ethyl benzoate and 100 g of propanone. Explain how she could separate the components
of this mixture.
[2]

[3]
M
c Anthracene is only slightly soluble in propanone with its solubility being about 1 g per
100 g of propanone. 10 g of anthracene and 100 g of propanone are shaken together for
a few minutes.
Explain how all the anthracene could be extracted from this mixture. [2]
d Rosalie makes the statement: ‘there is no temperature at which all four substances will
be liquids’. Evaluate this statement. [1]
SA

12 Zinc sulfate can be made by reacting excess zinc with dilute sulfuric acid according to the
following equation:
Zn(s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g)
Excess means that there will still be some zinc left in the reaction mixture when the reaction
has finished, but all of the sulfuric acid should have reacted.
a Classify each of the reactants as an element, compound or mixture. [1]
b Explain how you could obtain a solid sample of zinc sulfate from the reaction mixture. [2]
c Kinetic molecular theory is a model that can be used to explain the properties of solids,
liquids and gases. Describe the difference in the motion of the particles in zinc and hydrogen
at room temperature. [2]
d The melting point of zinc is 420 °C. Sketch a graph showing how the temperature of a sample
of zinc varies with time, as it is heated slowly from 400 °C to 440 °C. Identify the physical state
of zinc in each region of your graph. [2]

29
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: COURSEBOOK

EXAM-STYLE QUESTIONS

13 A student has been given 50 cm3 of a solution that contains 1 g of caffeine and 1 g of
sodium chloride.
a Explain why pure caffeine cannot be extracted from this mixture by heating to evaporate
off the water. [1]
b Some data about three solvents is given in the table.

Solvent Caffeine solubility Solvent miscibility Sodium chloride

/ g per 100 g of solvent with water solubility in solvent
water 2.3 miscible soluble
propanone 1.5 miscible insoluble

E
trichloromethane 11.6 immiscible insoluble

i Give two reasons why trichloromethane can be used to extract caffeine from the
mixture but propanone cannot. [2]
ii Describe how caffeine can be extracted from the mixture. [3]

PL
M
SA

30
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Chemistry

E
for the IB Diploma

PL WORKBOOK

Jacqueline Paris
M
SA

Second edition Digital Access

Together with IB teachers

Original material © Cambridge University Press & Assessment 2023. This material is not final and is subject to further changes prior to publication.
 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

Contents
How to use this series v
How to use this book vi
Unit 1 The nature of matter 1 5 Ideal gases 43

E
5.1 Real gases and ideal gases 44
1 The particulate nature of matter  2 5.2 Macroscopic properties of ideal
1.1 Elements, compounds and mixtures 4 gases and Exercise 5.3 Calculations
1.2 Kinetic molecular theory 6 involving ideal gases  46
1.3 Temperature and kinetic energy 7 5.3 Calculations involving ideal gases 46
1.4

2.1
2.2
Changes of state

2 The nuclear atom

3 Electron configurations
3.1
3.2
3.3
PL
The structure of atoms
Isotopes15
14

The electromagnetic spectrum 

The hydrogen atom spectrum
Electron configurations 
21
21
22
13

19
8
Unit 2

6.1
6.2
6.3
6.4

6.5
Bonding and structure
6 The ionic model
Ionic and covalent bonding
Formation of ions
The formation of ionic compounds
Ionic bonding and the structure
of ionic compounds
53
54
54

56
Physical properties of ionic compounds 57
51
52
M
6.6 Exercise 6.6 Lattice enthalpy and
3.4 Putting electrons into orbitals:
the strength of ionic bonding 58
Aufbau principle24
3.5 Ionisation energy 25 7 The covalent model 62
4 Counting particles by mass: 7.1 Covalent bonds 65
7.2 Shapes of molecules: VSEPR theory 67
The mole  30
SA

7.3 Lone pairs and bond angles 67

4.1 Relative masses 32
7.4 Multiple bonds and bond angles 68
4.2 Moles32
7.5 and Polarity and
4.3 The mass of a molecule 33
7.6  Pauling electronegativities 69
4.4 Empirical and molecular formulas 34
7.7 Intermolecular forces 69
4.5 Solutions36
7.8 Melting points and boiling points 70
4.6 Avogadro’s law 39
7.9 Solubility71
7.10 Covalent network structures 73
7.11 and The expanded octet and
7.12  Formal charge 74

46
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Contents

7.13 Shapes of molecules with Unit 4 What drives chemical

an expanded octet75
reactions?133
7.14 Hybridisation76
7.15 Sigma and pi bonds 77 12 Measuring enthalpy change 134
7.16 Resonance and delocalisation 78 12.1 Heat and temperature 135
12.2 Exothermic and endothermic
8 The metallic model 84 reactions136
8.1, 8.2 Classifying elements as metals, 12.3 Enthalpy changes and standard
and 8.3 Metallic bonding, Properties of conditions137
metals and their uses 85 12.4 Measuring enthalpy changes 138

E
8.4 Transition metals 86
13 Energy cycles in reactions 143
9 From models to materials 89 13.1 Bond enthalpies 144
9.1 Alloys90 13.2 Hess’s law 145
9.2 Polymers91 13.3 Using standard enthalpy
9.3

Unit 3

10.2
10.3

10.4
Classification of matter
10 The periodic table
10.1 The periodic table
Periodicity103
PL
Bonding and electronegativity

The chemistry of Group 1 and

Group 17
Oxides105
93

102

104
99
100
13.4

13.5
change of combustion data
Using standard enthalpy
changes of formation
Energy cycles for ionic compounds

14 Energy from fuels

14.1
14.2
Combustion reactions
Fuels159
14.3 and Renewable and non-renewable
energy sources and
158
157
147

148
150
M
10.5 Oxidation state 106 14.4  Fuel cells 160
10.6 The transition metals (d block) 107
15 Entropy and spontaneity 164
11 Functional groups: Classification 15.1 Entropy165
of organic compounds 112 15.2 Spontaneous reactions 166
11.1 The structures of organic molecules 114 15.3
SA

Gibbs energy and equilibrium 168

11.2 Homologous series and
functional groups 116 Unit 5 How much, how fast,
11.3 Naming organic molecules 119 how far? 173
11.4 Isomers120
11.5 Spectroscopic identification 16 How much? The amount
of organic compounds 123 of chemical change174
16.1 The meaning of chemical equations 175
16.2 Yield and atom economy of
chemical reactions178
16.3 Titrations179
16.4 Linked reactions 181

47
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CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

17 How fast? The rate 19.12 The strength of an acid and its
conjugate base  229
of chemical change185
19.13 The pH of salt solutions 230
17.1 and What is ‘rate’ of reaction and
19.14 More pH curves 231
17.2  Experiments to measure the rate
of reaction187 19.15 Buffer solutions 232
17.3 Collision theory  189 20 Electron transfer reactions 238
17.4 Factors affecting reaction rate 189
20.1 Redox reactions 240
17.5 The rate equation 191
20.2 Redox equations 241
17.6 Mechanisms of reactions 193
20.3 Redox titrations 243
17.7 Variation of the rate constant

E
20.4 The activity series 244
with temperature196
20.5 Voltaic cells 246
18 How far? The extent 20.6 Rechargeable batteries 248
of chemical change205 20.7 Electrolysis249
18.1 Reversible reactions and equilibrium 206 20.8 Redox reactions in organic chemistry 250
18.2
18.3
18.4

18.5 PL
The position of equilibrium
Equilibrium constants
Calculations involving
equilibrium constants210
Relationship between equilibrium
constants and Gibbs energy

Unit 6 Mechanisms of chemical

change217
207
208

211
20.9
20.10
20.11
Reduction reactions
Standard electrode potentials
Electrolysis of aqueous solutions

21 Electron sharing reactions

21.1
21.2
Radicals261
252
253
255

The radical substitution mechanism 262

22 Electron-pair sharing reactions

261

266
M
22.1 Nucleophilic substitution reactions 267
19 Proton transfer reactions 218 22.2 Addition reactions 268
19.1 and Acids, bases and salts, and 22.3 Lewis acids and bases 269
19.2  Reactions of acids 221 22.4 Nucleophilic substitution mechanisms 270
19.3 Brønsted–Lowry acids and bases 222 22.5 Electrophilic addition reactions
19.4 pH223 of alkenes 271
SA

19.5 Strong and weak acids and bases 224 22.6 Electrophilic substitution reactions 272
19.6 The dissociation of water  225
19.7 Calculating pH values 225 Glossary277
19.8 Acid–base titrations 226
19.9 pOH227
19.10 Ionisation constants for acids
and bases  227
19.11 The base ionisation constant, Kb228

48
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

How to use this series

How to use this series

This suite of resources supports students and teachers of the Chemistry course for the
IB Diploma programme. All of the books in the series work together to help students
develop the necessary knowledge and scientific skills required for this subject.

The coursebook with digital access provides full coverage of

E
the latest IB Chemistry guide.
It clearly explains facts, concepts and practical techniques, and
uses real world examples of scientific principles. A wealth of
formative questions within each chapter help students develop
their understanding, and own their learning. A dedicated

PL
chapter in the digital coursebook helps teachers and students
unpack the new assessment, while exam-style questions provide
essential practice and self-assessment. Answers are provided on
Cambridge GO, to support self-study and home-schooling.

The workbook with digital access builds upon the

coursebook with digital access with further exercises
Chemistry
M
and exam-style questions, carefully constructed to for the IB Diploma
help students develop the skills that they need as WORKBOOK

they progress through their IB Chemistry Diploma Jacqueline Paris

course. The exercises also help students develop

understanding of the meaning of various command
words used in questions, and provide practice in
responding appropriately to these.
SA

Second edition Digital Access

Together with IB teachers

The Teacher’s resource supports and enhances the coursebook

with digital access and the workbook with digital access.
This resource includes teaching plans, overviews of required
background knowledge, learning objectives and success criteria,
common misconceptions, and a wealth of ideas to support lesson
planning and delivery, assessment and differentiation. It also
includes editable worksheets for vocabulary support and exam
practice (with answers) and exemplar PowerPoint presentations,
to help plan and deliver the best teaching.

49
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

How to use this book

A chapter outline appears at the start of every chapter to introduce the learning aims
and help you navigate the content.

CHAPTER OUTLINE KEY TERMS

In this chapter you will: Definitions of key
vocabulary are given

E
• describe the structure of the atom and the relative charges and masses at the beginning of
of protons, neutrons and electrons each chapter.
• describe how protons, neutrons and electrons behave in electric fields You will also find
• deduce the number of protons, neutrons and electrons in atoms and ions. definitions of these
words in the glossary.

Exercises

EXAM-STYLE QUESTIONS
PL
Exercises help you to practice skills that are important for studying Standard Level and
Higher Level Chemistry.

Questions at the end of each chapter are more demanding exam-style questions,
TIP
Tip boxes will help
you complete the
exercises, and give
you support in areas
that you might find
difficult.
M
some of which may require use of knowledge from previous chapters. Answers to
these questions can be found in digital form on Cambridge GO.
Visit Cambridge GO and register to access these resources at
www.cambridge.org/GO

KEY EQUATIONS
SA

In these boxes you find chemical equations in the form of symbols

and formula.

50
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Unit 1

The nature of matter

E
PL
M
SA

Original material © Cambridge University Press & Assessment 2023. This material is not final and is subject to further changes prior to publication.
 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Chapter 1

The particulate nature

of matter
CHAPTER OUTLINE
In this chapter you will:

E
• understand the terms element, compound and mixture
• understand the differences between heterogeneous and
homogeneous mixtures
• understand how to separate the components of a mixture
•

•
•
•
liquids and gases

energy of particles
PL
use kinetic molecular theory to understand the properties of solids,

understand that temperature in K is proportional to the average kinetic

understand how to convert temperatures between K and °C

use state symbols in chemical equations
use kinetic molecular theory to explain changes of state.
M
KEY TERMS
Make sure you understand the following key terms before you do
the exercises.
atom: the smallest part of an element that can still be recognised as that
SA

element; in the simplest picture of the atom, the electrons orbit around the
central nucleus; the nucleus is made up of protons and neutrons (except for
a hydrogen atom, which has no neutrons)
element: a chemical substance that cannot be broken down into a simpler
substance by chemical means. Each atom has the same number of protons in
the nucleus
compound: a pure substance formed when two or more elements combine
chemically in a fixed ratio
mixture: two or more substances mixed together. The components of a
mixture can be mixed together in any proportion (although there are limits for
solutions). The components of a mixture are not chemically bonded together,
and so, retain their individual properties. The components of a mixture can be
separated from each other by physical processes

52
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

CONTINUED
molecule: an electrically neutral particle consisting of two or more atoms
chemically bonded together
heterogeneous mixture: a mixture of two or more substances, that does not
have uniform composition and consists of separate phases. A heterogeneous
mixture can be separated by mechanical means. An example is a mixture of
two solids
chemical properties: how a substance behaves in chemical reactions
chromatography: a technique used to separate the components of a mixture

E
due to their different affinities for another substance and/or solubility in a solvent
deposition: the change of state from a gas to a solid
filtration: a separation technique used to separate insoluble solids from a
liquid or solution

PL
physical properties: properties such as melting point, solubility and electrical
conductivity, relating to the physical state of a substance and the physical
changes it can undergo
solvation: a process used to separate a mixture of two or more substances,
due to differences in solubility
states of matter: solid, liquid and gas
state symbols: used to indicate the physical state of an element or
compound; these may be either written as subscripts after the chemical
M
formula or in normal type: (aq) = aqueous (dissolved in water); (g) = gas;
(l) = liquid; (s) = solid
boiling: change of state from a liquid to a gas at the boiling point of
the substance
boiling point: the temperature at which a liquid boils under a specific set of
conditions - usually we will be considering the boiling point at atmospheric
SA

pressure
distillation: a separation technique used to separate the solvent from a
solution or separate liquid components of a mixture that have different
boiling points
sublimation: the change of state from a solid to a gas
melting: the change of state from a solid to a liquid
freezing: the change of state from a liquid to a solid
melting point: the temperature at which melting occurs
homogeneous mixture: a mixture of two or more substances with the same
(uniform) composition throughout the mixture – it consists of only one phase.
Examples are solutions or a mixture of gases
solution: that which is formed when a solute dissolves in a solvent

53
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

CONTINUED
evaporation: the change of state from a liquid to a gas that can occur
at any temperature above the melting point
solute: a substance that is dissolved in another substance (the solvent) to form
a solution
solvent: a substance that dissolves another substance (the solute);
the solvent should be present in excess of the solute
temperature: a measure of the average kinetic energy of particles

E
Exercise 1.1 Elements, compounds
and mixtures

2 PL
This exercise will check you understand the key terms element, compound, mixture,
atom and molecule, which are important fundamental ideas in chemistry.

Approximately how many different elements are there?

Some elements exist as individual atoms, some as a small group of atoms bonded
together into a molecule and others are bonded together into a giant structure.

b
Name two elements that exist as giant structures at 25 °C.

Name an element that exists as a single atom.

M
c Name an element that exists as a molecule made of two atoms joined together
(a diatomic molecule).

3 Identify which of the following formulas represent atoms and which

represent molecules: TIP

a He An atom is a single
SA

particle.
b O2
A molecule is made
c H2O up of more than one
atom.
d C
The atoms in a
4 Identify which of the following formulas represent elements and which molecule can be of
represent compounds: the same element.
a He

b O2

c H2O

d C

54
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

5 This statement is incorrect, explain why:

TIP
Elements are made of atoms and compounds are made of molecules.
Question 5 is linked
6 An alloy is a mixture of a metal and other elements. Give one way in which the to ideas in Chapter 6.
composition of an alloy differs from that of a compound.

7 Compounds have both different chemical properties and physical properties from
the elements from which they are formed.

a What is meant by the term physical properties?

b What is meant by the term chemical properties?

E
8 Most everyday substances are mixtures although they are often labelled as pure.
Pure orange juice is a common example. The manufacturers simply mean that
nothing has been added to the orange juice. In chemistry, the term pure is
not used in the same way.

In chemistry, what is meant by the term pure?

PL
Why do the components of a mixture retain their individual properties?

10 Group the following substances into elements, mixtures and compounds:

air, water, sodium chloride solution, sodium chloride crystals, iron, chlorine gas,
carbon dioxide gas.

11 What name is given to a mixture that has a uniform composition and only consists
of one phase?

12 What name is given to a mixture that does not have a uniform composition and
TIP
Solid, liquid, gas
and solution are all
examples of phases.
M
consists of separate phases?

13 Why is a mixture of the solids sodium chloride and sand not a

homogeneous mixture?

14 When a small amount of salt and water is mixed together, it forms a homogeneous
mixture, but this is not true when flour is mixed with water, why?
SA

15 Are chemical or physical processes typically used to separate the components

of a mixture?

55
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CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

16 Match the name of the separation technique with the type of mixture it can be
used to separate.

Technique Types of mixture

A filtration 1 substances with very different solubilities
in a solvent
B distillation 2 an insoluble solid from a liquid
C evaporation 3 a solute with very different solubilities in
two different solvents

E
D solvation 4 the solute from a solution
E solvent extraction 5 liquids with a large difference in their
solubilities in different solvents
F paper chromatography 6 a mixture of substances with small
differences in their solubilities in a solvent

Exercise 1.2 Kinetic molecular theory

and gases.

1
PL
Kinetic molecular theory is used to explain the observed properties of solids, liquids

Complete Table 1.1, which describes the arrangement and movement of particles
in solids, liquids and gases.
M
Solids Liquids Gases
diagram showing
the arrangement of
the particles
relative distance of
the particles from
SA

one another
relative energy of
the particles
movement
of particles
relative force of
attraction between
the particles

Table 1.1: Arrangement and movement of particles.

2 Which of the descriptions of particles in Table 1.1 can explain the fixed shape
of solids and the lack of a fixed shape in liquids and gases?

56
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

3 Which of the descriptions in Table 1.1 explain why, at a given temperature,

the volume of a gas is not fixed but the volume of solids and liquids are?

4 Younger students are often confused by the observed properties of a powder.

A powder can flow like a liquid and take up the shape of its container but
does not completely spread out into a puddle like a liquid.

a How would you explain that a powder is a solid?

b How would you explain the ability of a powder to flow like a liquid?

5 Which scale is the SI scale for temperature?

E
6 On the kelvin scale, what does zero K (or absolute zero) represent?

7 Complete Table 1.2 to show equivalent temperatures on the kelvin and

Celsius scales.

Celsius scale Kelvin scale

0

40

946
PL 373

74

500
3
TIP
Absolute zero equals
−273.15°C, but you
can use −273 °C
for your chemical
M
calculations.
Table 1.2: Equivalent temperatures on the kelvin and Celsius scales.

8 Temperature is used in some chemical calculations. When it is, the kelvin scale is
always used, unless the calculation involves a temperature change.

Explain why either Celsius or kelvin can be used to measure temperature change.
SA

Exercise 1.3 Temperature and

kinetic energy
Not all of the particles in a sample have the same amount of energy, and so, they do
not all move with the same speed. In this exercise, you will explore the distribution of
kinetic energies at different temperatures.

1 Consider a sample of oxygen at a constant temperature.

a Do all the oxygen particles have the same kinetic energy? Explain your answer.

b Do all the particles of the gas move at the same speed? Explain your answer.

57
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CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

2 Consider a mixture of the gases nitrogen and helium at a constant temperature.

a The average kinetic energy of the particles will be higher for which gas?

b The average speed of the particles will be higher for which gas?

3 Describe how the following change when the temperature of a gas is increased:
TIP
a the average kinetic energy of the particles
The most probable
b the average speed of the particles kinetic energy is the
energy at the peak of
c the most probable kinetic energy of the particles a Maxwell–Boltzmann

E
d the fraction of particles with the most probable kinetic energy. distribution curve.

Exercise 1.4 Changes of state

PL
Heating or cooling a substance can cause it to change state, as these processes involve
the breaking or formation of forces of attraction between the particles. In this exercise,
you will check that you understand these processes and can work out the state of a
substance at a given temperature from its melting point and boiling point.
Figure 1.1 summarises the changes of state.

sublimation
heating – energy is supplied
particles gain energy

deposition

boiling
M
solid melting liquid evaporating gas
freezing condensing

cooling – energy taken out

particles lose energy
SA

Figure 1.1: The changes of state.

1 Which change of state does not take place only at a fixed temperature for a
given pressure? TIP

2 Identify which changes of state are exothermic and which are endothermic. The same name for
the temperature at
3 What name is given to the temperature at which a substance changes from a liquid which the change in
to a solid? question 3 happens
is used, no matter in
4 What name is given to the temperature at which a substance changes between which direction the
gas and liquid? change happens.
5 Carbon dioxide and iodine are two examples of substances that undergo sublimation.

a What is meant by the term sublimation?

b What term is used to describe the reverse of sublimation?

58
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

6 Complete the table to show whether a substance is a solid, liquid or gas at the
temperature stated in the column header.

Melting Boiling State at State at State at

Substance
point / °C point / °C −50 °C 115 °C 200 K
A 15 125
B 253 578
C −83 78
D −169 −87

E
7 Figure 1.2 shows the cooling curve for a substance.

280
260 A
240
220
200
B
C

PL
Temperature/°C

180
160 E
140 D
120
100
80 F
60
40
20
M
0
0 5 10 15 20 25 30 35 40 45 50
Time/mins

Figure 1.2: The cooling curve for a substance.

a Label the diagram to show the following:

SA

i the region where the substance is a solid

ii the region where the substance is a liquid

iii the region where the substance is a gas

iv the region where the substance is freezing

v the region where the substance is condensing

vi the melting point of the substance

vii the boiling point of the substance.

b Explain, in terms of the movement and arrangement of the particles, why the
temperature of the substance remains the same during a change of state.

59
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

EXAM-STYLE QUESTIONS
1 Which of the following lists substances that are all made up of molecules?
A C, O2, CO2
B Na, Cl2, NaCl
C H2, He, Li
D P4, S8, O3 [1]
2 Which of the following statements is true of heterogeneous mixtures?
A Their components cannot be separated by physical means.
B They have the same composition throughout the mixture.

E
C The components are in a fixed ratio.
D The components are in separate phases. [1]
3 Which of the following is not a heterogeneous mixture?
A cola
B tea with milk
C tea with sugar
D milk

A boiling, condensing
B condensing, boiling
C evaporation, cooling
D boiling, cooling
PL
4 Which of the following shows the correct sequence of the changes of state involved in distillation?

5 What is the name given to the separation technique that is used to separate the components of
[1]

[1]
M
a mixture that have different solubilities in a solvent at different temperatures?
A distillation
B recrystallisation
C evaporation
D paper chromatography [1]
6 Mercury is a liquid at 25 °C, which of the following could be its melting and boiling points?
SA

Melting point Boiling point

A −38.9 °C 83.7 K
B −38.9 K 629.7 °C
C −38.9 K 356.7 K
D −38.9 °C 356.7 °C
 [1]

60
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

1 The particulate nature of matter

CONTINUED
7 Which is the correct equation for sublimation?
A CO2(s) CO2(g)
B CO2(g) CO2(s)
C H2O(s) H2O(l)
D CO2(g) CO2(aq) [1]
8 Which statement is correct about melting?
A The average kinetic energy of the particles increases, but the temperature stays the same.
B The average kinetic energy of the particles increases, and the temperature increases.

E
C The average kinetic energy of the particles stays the same, but the temperature increases.
D The average kinetic energy of the particles stays the same, and the temperature stays the same. [1]
9 Ammonia liquid boils at −33 °C and freezes at −78 °C at atmospheric pressure.
a Predict the state of ammonia at
i −50 °C
ii −80 °C
iii 200 K.

PL
b Sketch a graph of temperature against time as a sample of ammonia is cooled from 0 °C to −50 °C.
10 Some seaweeds accumulate iodide ions in their leaves, and so, are a good source of iodine.
The seaweed must first be dried and then heated to burn off the organic matter. The remaining
ash is then boiled in water and allowed to cool. The iodide ions dissolve in the water.
a Suggest a suitable technique that could be used to separate the iodide solution from
any insoluble impurities.
b State the type of mixture that remains after the insoluble impurities have been removed.
[3]
[4]

[1]
[1]
M
c When dilute sulfuric acid and hydrogen peroxide are added to the mixture, an aqueous solution
of iodine is produced:
2H+ + H2O2 + 2I− I2 + 2H2O
Give the state symbols for I2 and H2O in the equation above. [2]
d Iodine is not particularly soluble in water. It is much more soluble in organic solvents such
as cyclohexane. Outline a method that could be used to separate the iodine from the solution. [3]
SA

11 The statements below describe the analysis of a mixture of amino acids by paper chromatography.
a Place the statements in the correct order:
i spray the plate with a locating agent
ii mark the position of the solvent front
iii place a small sample of the unknown sample on the bottom of a piece of chromatography paper
iv place the paper into a tank containing a suitable solvent
v allow the solvent to rise up the paper [3]

61
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CHEMISTRY FOR THE IB DIPLOMA: WORKBOOK

CONTINUED
b Figure 1.3 shows the results of the experiment.

M = Mixture of amino acids

A = Glycine
B = Lysine
C = Alanine

E
M A B C

Figure 1.3: Chromatogram of an amino acid mixture.

PL
Which amino acids did the mixture contain?
Why do substances A, B and C each only produce one spot on the chromatogram?
[1]
[1]
M
SA

62
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

Chemistry

E
for the IB Diploma
PL
M
SA

Digital Teacher’s Resource

Together with IB teachers

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Table of contents
About the authors
How to use this series
How to use this Teacher’s resource
About the syllabus

E
About the assessment
How to Integrate TOK in your Science lesson
Academic writing and the international baccalaureate diploma

Teaching notes
Unit 1: The nature of matter

1 The particulate nature of matter

1.1
1.2
1.3
1.4
PL
Elements, compounds and mixtures
Kinetic molecular theory
Temperature and kinetic energy
Changes of state
1
2
5
5
5
M
2 The nuclear atom 9
2.1 The structure of atoms 10
2.2 Isotopes 13

3 Electron configuration 17
3.1 The electromagnetic spectrum 19
SA

3.2 The hydrogen atom spectrum 19

3.3 Electron configurations 21
3.4 Putting electrons into orbitals: Aufbau principle 21
3.5 Ionisation energy 24

4 Counting particles by mass: The mole 28

4.1 Relative masses 30
4.2 Moles 30
4.3 The mass of a molecule 30
4.4 Empirical and molecular formulas 33
4.5 Solutions 35
4.6 Avogadro’s law 38

32
1
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CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

5 Ideal gases 41
5.1 Real gases and ideal gases 42
5.2 Macroscopic properties of ideal gases 42
5.3 Calculations involving ideal gases 45

Unit 2: Bonding and structure

6 The ionic model 49

6.1 Ionic and covalent bonding 50
6.2 Formation of ions 50
6.3 The formation of ionic compounds 51

E
6.4 Ionic bonding and the structure of ionic compounds 53
6.5 Physical properties of ionic compounds 53
6.6 Lattice enthalpy and the strength of ionic bonding 53

7 The covalent model 55

7.1
7.2
7.3
7.4
7.5
7.7
7.8
7.9
7.10
7.11
Covalent bonds

Intermolecular forces PL
Shapes of molecules: VSEPR theory
Lone pairs and bond angles
Multiple bonds and bond angles
Polarity and Pauling electronegativities

Melting points and boiling points

Solubility
Covalent network structures
The expanded octet and
57
57
57
57
59
59
59
59
59
61
M
7.12 Formal charge 61
7.13 Shapes of molecules with an expanded octet 61
7.14 Hybridisation 61
7.15 Sigma and pi bonds 61
7.16 Resonance and delocalisation 61

8 The metallic model 84 64

SA

8.1 Classifying elements as metals, 65

8.2 Metallic bonding 66
8.3 Properties of metals and their uses 66
8.4 Transition metals 68

9 From models to materials 70

9.1 Alloys 71
9.2 Polymers 74
9.3 Bonding and electronegativity 74

2 33
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CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Unit 3 Classification of matter

10 The periodic table 77

10.1 The periodic table 79
10.2 Periodicity 81
10.3 The chemistry of Group 1 and Group 17 81
10.4 Oxides 83
10.5 Oxidation state 85
10.6 The transition metals (d block) 87

11 Functional groups: Classification of organic compounds 91

E
11.1 The structures of organic molecules 93
11.2 Homologous series and functional groups 93
11.3 Naming organic molecules 96
11.4 Isomers 99
11.5 Spectroscopic identification of organic compounds 101

12 Measuring enthalpy change

12.1
12.2
12.3
12.4
Heat and temperature
PL
Unit 4 What drives chemical reactions?

Exothermic and endothermic reactions

Enthalpy changes and standard conditions
Measuring enthalpy changes

13 Energy cycles in reactions

13.1 Bond enthalpies
107
108
108
110
110

112
113
M
13.2 Hess’s law 115
13.3 Using standard enthalpy change of combustion data 117
13.4 Using standard enthalpy changes of formation 117
13.5 Energy cycles for ionic compounds 117

14 Energy from fuels 119

SA

14.1 Combustion reactions 120

14.2 Fuels 121
14.3 and Renewable and non-renewable energy sources and
14.4 Fuel cells 123

15 Entropy and spontaneity 125

15.1 Entropy 126
15.2 Spontaneous reactions 126
15.3 Gibbs energy and equilibrium 128

Unit 5 How much, how fast, how far?

16 How much? The amount of chemical change 131

16.1 The meaning of chemical equations 132
16.2 Yield and atom economy of chemical reactions 134
16.3 Titrations 136
16.4 Linked reactions 136

34
3
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CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

17 How fast? The rate of chemical change 138

17.1 and What is ‘rate’ of reaction and 140
17.2 Experiments to measure the rate of reaction 140
17.3 Collision theory 141
17.4 Factors affecting reaction rate 141
17.5 The rate equation 143
17.6 Mechanisms of reactions 143
17.7 Variation of the rate constant with temperature 143

18 How far? The extent of chemical change 146

18.1 Reversible reactions and equilibrium 148

E
18.2 The position of equilibrium 149
18.3 Equilibrium constants 151
18.4 Calculations involving equilibrium constants 151
18.5 Relationship between equilibrium constants and Gibbs energy 151

Unit 6 Mechanisms of chemical change

19 Proton transfer reactions

19.1
19.2
19.3
19.4
19.5
19.6
19.7
Reactions of acids

pH
PL
and Acids, bases and salts

Brønsted–Lowry acids and bases

Strong and weak acids and bases

The dissociation of water
Calculating pH values
154
157
157
157
159
159
159
159
M
19.8 Acid–base titrations 162
19.9 pOH 163
19.10 Ionisation constants for acids and bases 163
19.11 The base ionisation constant, Kb 165
19.12 The strength of an acid and its conjugate base 167
19.13 The pH of salt solutions
SA

19.14 More pH curves

19.15 Buffer solutions

20 Electron transfer reactions 171

20.1 Redox reactions 174
20.2 Redox equations 174
20.3 Redox titrations 174
20.4 The activity series 176
20.5 Voltaic cells 176
20.6 Rechargeable batteries 176
20.7 Electrolysis 179
20.8 Redox reactions in organic chemistry 180
20.9 Reduction reactions 180
20.10 Standard electrode potentials 183
20.11 Electrolysis of aqueous solutions 183

4 35
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CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

21 Electron sharing reactions 187

21.1 Radicals 188
21.2 The radical substitution mechanism 188

22 Electron-pair sharing reactions 191

22.1 Nucleophilic substitution reactions 193
22.2 Addition reactions 193
22.3 Lewis acids and bases 195
22.4 Nucleophilic substitution mechanisms 197
22.5 Electrophilic addition reactions of alkenes 199
22.6 Electrophilic substitution reactions 201

E
Digital resources
The following items are available on Cambridge GO. For more information on how to access and
use your digital resource, please see inside front cover.

Worksheets
PowerPoints
End of Chapter tests
Specimen papers
Coursebook answers
Workbook answers
Worksheet answers
PL
M
End of chapter tests answers
Specimen paper answers
Glossary
Acknowledgements
SA

36
5
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

1 The particulate
nature of matter
Teaching plan

E
Sub-chapter Approximate Learning content Resources
number of
learning
hours
1.1 Elements, 2–3 Recall the definitions of elements, Coursebook
compounds and compounds and mixtures.
Section 1.1
mixtures

PL
Distinguish between the properties of
an element, compound or mixture.
Understand the difference between
homogeneous and heterogeneous
mixtures.
Describe experimental techniques to
separate mixtures.
Test your understanding
Questions 2 and 3
Workbook
Exercise 1.1
Teacher’s resource
PowerPoint 1 Slides 2–5
Worksheet 1.1 Questions
M
1, 3 and 5,
End of chapter test
Questions 1–6, 9, 10

1.2 Kinetic 2 Determine the state symbols in Coursebook

molecular theory chemical equations.
2 Sections 1.2–1.4
1.3 Temperature Recall the names of the changes
SA

and kinetic 2 of state. Test your understanding

energy Question 12
1.4 Changes Explain the physical properties of
matter and changes of states using Workbook
of state
kinetic molecular theory. Exercises 1.2–1.4
Understand that temperature in kelvin Teacher’s resource
is a measure of average kinetic energy
of particles. PowerPoint 1 Slide 6

Know how to convert between Celsius Worksheet 1.1 Questions

and kelvin scales. 2 and 4
End of chapter test
Questions 7 and 8

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

BACKGROUND KNOWLEDGE

• Understand how to classify substances as • Know the names of the interconversions of

elements, compounds or mixtures. the three states of matter.
• Describe simple techniques for separating • Know there are two different scales for
mixtures (filtration, distillation, evaporation measuring temperatures: degree Celsius
and paper chromatography). and kelvin.
• Draw particle diagrams and use them to • Know how to use a data logger and a
explain the properties of solids, liquids temperature probe, plot line graphs and
and gases. draw lines of best fit.

E
Syllabus overview
• The first part of the syllabus covers the concepts of elements, compounds, mixtures and the application
of kinetic molecular theory to explain the particle models of states of matter. Chemists should know the

•
PL
differences between compounds and mixtures and how to construct names and formulas of compounds.
This will facilitate the study of chemical reactions using balanced symbol equations and how to solve
problems using molar ratios of reactants and products (Chapters 4 and 16).
There are many opportunities for students to practice fundamental laboratory techniques, covering the
various methods for separating mixtures. Students should be encouraged to think about how to test for
the purity of products after separation and research into how to purify products further. When measuring
melting/cooling curves of substances, students also practice mathematical skills of presenting their data
graphically and analysing the results to extract information on melting/boiling points.
Simulations can be used to illustrate molecular movement of particles. This gives an introduction on how
the kinetic energy of particles is distributed in a sample of gas at a fixed temperature and the concept of
activation energy in a chemical reaction (Chapter 17).
M
1.1 Elements, compounds and mixtures
LEARNING PLAN

Learning objectives Success criteria

SA

Understand the terms element, compound Students should be able to explain the terms
and mixture element, compound and mixture and distinguish
between them.
Understand the differences between
heterogeneous and homogeneous mixtures Students should be able to explain the difference
between heterogeneous and homogeneous
Understand how to separate the components mixtures and give examples of each.
of a mixture
Students should be able to explain the different
methods for separating the components of
a mixture and suggest a suitable method for
separating a particular mixture.

38
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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Common misconceptions
Misconception How to identify How to overcome
Students confuse the meaning Ask students to assign various Draw particle diagrams to
of compounds and molecules names and formulas of elements show which names/formulas
and compounds to a Venn are molecules or compounds.
diagram of two circles labelled Molecules can be elements
compounds and molecules. (O2) or compounds (H2O) and
only covalent compounds
are molecules.
Students confuse physical and Show students pictures A physical change is one in which
chemical changes of different processes (for no new chemicals are formed, for

E
example, physical processes, example, dissolving and changes
including melting, freezing of states (in separating mixtures).
and sublimation, and chemical A chemical reaction involves
processes, including rusting, making new substances. Teachers
fireworks and cooking an egg) can demonstrate some examples
and ask them to distinguish when elements are combined

Starter ideas
PL
the physical from the
chemical changes.
in chemical reactions to form
compounds. For example,
burning Na in Cl2 or Mg in O2.
Use particle diagrams to show
that the microscopic make-up
of the reactants and products
is different, and the atoms are
bonded together differently.
M
1 Recap prior knowledge from pre-IB (10 minutes)
Resources: Test your understanding questions 2 and 3 in the Coursebook.
Description and purpose: Students define element, compound and mixture. They should then sort out the
listed substances and diagrammatic representations into the three categories. This activity assesses students’
prior knowledge.
What to do next: If most of the students can define element, compound and mixture and identify them
SA

correctly, teachers can ask them to give more examples of each. Make sure to emphasise the keywords in
the definitions. If students find it difficult to distinguish amongst the three categories, help by pointing out
that elements can be found in the periodic table, elements combine chemically to form compounds and give
examples of names and formulas of various compounds. Most of things we meet daily are mixtures and can
be separated by physical methods.
Language focus: Learners are encouraged to pay attention to definitions of the key terms.

Main teaching ideas

2 Teacher demonstrations on the formation of compounds from constituent elements
(20 minutes)
Resources: Search for the websites mentioned in the Description and purpose line for apparatus/chemicals
required for each demonstration.
Description and purpose: Iron and sulphur (search the ‘Royal Society of Chemistry’ website with the
keywords ‘iron and sulfur reaction’)
Sodium and chlorine (search the ‘University of Washington’ website with the keywords ‘sodium and
chlorine reaction’)

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

These experiments can be performed to show how elements retain their properties in a mixture but change
their properties when forming compounds. Ask students to record the observations (changes in physical
states, colours, endothermic vs exothermic etc.) during the reaction and write word and symbol equations to
represent the chemical processes.
Differentiation ideas
Support: Provide students with a table to record the appearance of the reactants and products of the
reactions, and their observations during the reactions. Stress the importance of forming new substances
in chemical reactions to form compounds.
Stretch and challenge: Students can be asked to construct balanced chemical equations for these
reactions with state symbols.
Language focus: Recording observations. This is one way of checking the correct use of terminology.

E
3 Student practical (2 × 45 minutes)
Resources: A mixture of sand and water, sodium chloride solution, a mixture of ink and water, a mixture
of food dyes. Apparatus required for filtration (for example, funnel, filter paper, clamp, boss head, stand,
beaker), simple distillation (for example, round-bottomed flask, thermometer, bung, Liebig condenser,
beaker, Bunsen burner, heat-proof mat), evaporation (evaporating basin, gauze, Bunsen burner, heat-proof

PL
mat) and paper chromatography (chromatography paper, beaker, pencil, ruler, small capillary tube).
Description and purpose: Ask students to separate various mixtures, including sand and water (filtration),
table salt dissolved in water (simple distillation to keep the water or evaporation to obtain only the salt
crystal), ink and water (simple distillation), a mixture of food dyes (paper chromatography). The practical
could be run at different stations set up around a laboratory.
Differentiation ideas
Support: Providing exact step-by-step methods with diagrams to guide students through the practical.
Stretch and challenge: Students design their own methods and carry out the experiments once their
methods are approved by a teacher.
M
Plenary ideas
1 How to separate mixtures (10 minutes)
Resources: Fill in the information in the following table on how to separate mixtures. The first two rows have
been completed as an example.

Separation of … Homogeneous or Technique Example

SA

heterogeneous mixtures
two liquids homogeneous simple distillation: the ink and water: water
liquid with a lower boiling will boil first
point will boil first
two liquids heterogeneous two liquids that are water and
immiscible and have dichloromethane:
different densities can be dichloromethane is
separated into layers in a denser, so it will come
separatory funnel out of the funnel first
a solid and a liquid homogeneous
a solid and a liquid heterogeneous
two solids heterogeneous

Description and purpose: This exercise gives students an opportunity to summarise, recall and apply their
knowledge.
Language focus: Take note of the language used when making a summary of the experimental methods.

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 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

1.2 Kinetic molecular theory; 1.3 Temperature and

kinetic energy and 1.4 Changes of state
LEARNING PLAN

Learning objectives Success criteria

Use kinetic molecular theory to understand Students should be able to explain the properties
theproperties of solids, liquids and gases of solids, liquids and gases in terms of kinetic
molecular theory.

E
Understand that temperature in K is
proportional to the average kinetic energy Students can recall that temperature in K is
of particles proportional to the average kinetic energy of
particles.
Understand how to convert temperatures
between K and °C Students should be able to convert temperatures
between K and °C.

Common misconceptions
Misconception
PL
Use state symbols in chemical equations
Use kinetic molecular theory to explain
changes of state

How to identify
Students can apply state symbols in chemical
equations.
Students should be able to explain changes of
state using kinetic molecular theory.

How to overcome
M
Students confuse boiling and Ask students to explain the In both processes, liquids change
evaporation differences between boiling to gases. Boiling occurs at a
and evaporation. particular temperature (boiling
point) and throughout the whole of
the liquid. Evaporation can occur
at all temperatures but only on the
surface of the liquid.
SA

Changes in temperature on Ask students what is ΔT = 30 °C The intervals on both the

kelvin and Celsius scales are when converted to kelvin scale. temperature scales are the same,
muddled up so the changes in temperature can
have either K or °C as units but the
numerical values remain the same.
A change of 20 °C to 50 °C (30 °C)
has the same value as a change
of 293 K to 323 K (30 K) when
converted to the kelvin scale.
Students struggle to Show students a model of Students often have the
understand what is in the the giant ionic lattice of NaCl misconception that the space is
space between particles in the or a model of the molecular filled with air. Air is a mixture of
particle model structure of ice and ask them many gas molecules/atoms, and
what is in between the particles these entities are themselves too
in the model. big to fit into the space between
particles in the NaCl ions or H2O
molecules in the solid state.

5 Chemistry for the IB Diploma – Bonsall © Cambridge University Press 2023 41

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CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Starter ideas
1 Solid, liquid and gas particle diagrams (10 minutes)
Resources: A piece of A4 paper showing three equal-sized square boxes.
Description and purpose: Students complete diagrams showing the arrangement of particles in a solid,
liquid and gas and then name the processes for the changes of state (including sublimation and deposition).
This activity recaps students’ knowledge from pre-IB.
What to do next: Students should be clearly aware of how the particle arrangements are represented in
these diagrams.

2 Recognise the states of matter based on melting and boiling points (10 minutes)
Resources: Test your understanding Question 12 in the Coursebook.

E
Description and purpose: Ask students to identify the states of matter at given temperatures. Students should
be able to apply their knowledge of melting and boiling points to recognise the states of matter.
What to do next: Show the melting/boiling points on a number line to order them, if students find this
activity difficult.

Main teaching ideas

PL
1 Practical on the freezing of stearic acid (45 minutes)
Resources: A detailed list of apparatus and chemicals can be found by searching the ‘Royal Society of
Chemistry’ website with the keywords ‘freezing of stearic acid’. Graph paper is required for analysing
the results.
Description and purpose: Students need to heat up 3 spatulas of stearic acid until it melts. Then allow the
acid to cool and take a temperature reading every 10 seconds with a temperature probe and a data logger.
Plot a graph of temperature of stearic acid (after it completely melts) against time. Ask students to explain
the shape of the cooling curve and identify the freezing point of the acid.
Support: Provide a step-by-step method with a titled table to write down results. The temperature against
time graph could be plotted using Google sheets.
M
Challenge: Students can design their own method to carry out the experiment and plot data on a piece of
graph paper.

2 Explaining the changes of states of matter in terms of the changes in the

arrangement, movement and energy of the particles and the bonds in between
the particles (20 minutes)
Resources: Demonstrations showing changes of states (for example, ice melting, water boiling, steam
SA

condensing, dry ice subliming)

Description and purpose: Ask students to apply the concept of intermolecular forces/bonds to explain the
changes of states of water. Students will self-assess their explanations with key words.
Differentiation ideas
Support: Teachers can help students to review their answers and provide feedback on the use of
keywords.
Stretch and challenge: Students can identify which processes are endothermic and which are exothermic.
Students can look into the different types of intermolecular forces and other types of bonding between
particles.
Students can find out why sublimation occurs for some substances using the phase diagram (search on
Chemguide.co.uk with the keywords ‘phase diagram’).
Language focus: Using scientific terminology and constructing logical long answers.

42
6 Chemistry for the IB Diploma – Bonsall © Cambridge University Press 2023
Original material © Cambridge University Press & Assessment 2023. This material is not final and is subject to further changes prior to publication.
 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Plenary ideas
1 True or false (5 minutes)
Resources: A table listing various properties of solids liquids and gases, for example (Worksheet question 4):

True False
solids can be compressed
liquids have no fixed volume
gases have no fixed volume
gases have a high density
solids have fixed shape

E
liquids can diffuse

Description and purpose: Students mark true/false in the table. This activity allows students to apply their
knowledge of the particle models to draw conclusion on the macroscopic properties of different states
of matter.

PL
2 Labelling a cooling curve and explaining the stages of temperature change
using the kinetic molecular model (10 minutes)
Resources: A cooling curve, for example, Figure 1.1.

Tc
Cooling curve
Temperature

condensation
point

Tf freezing
M
point

Time

Description and purpose: Students label the cooling curve with states of matter, changes of states, and
identify the condensation and freezing points of the substance. Students also need to provide explanations
on why temperature of the liquid goes down as cooling occurs but stays constant during freezing.
SA

Assessment ideas
• Ask students to give examples of elements, compounds, mixtures, solids, liquids and gases around the
classroom/lab.
• Suggest a suitable method for separating different types of mixtures.
• Label diagrams of lab apparatus and set-ups for filtration, evaporation, distillation and reflux.
• Students can design and carry out an experiment to obtain pure salt from rock salt.
• Calculations involving conversions between Celsius and kelvin temperature scales.
• Label the different stages of a melting and boiling curve.
• Test your understanding questions from the Coursebook.
• Define key words from the chapter.
• Explain the changes of states that occur during separation of mixtures. Ask students to use Post-it notes to
assess their peers’ answers.
• Give students explanations (containing common mistakes, missing out keywords) on the changes in the
states of matter using the kinetic molecular theory and ask them to mark against a mark scheme.

7 Chemistry for the IB Diploma – Bonsall © Cambridge University Press 2023 43

Original material © Cambridge University Press & Assessment 2023. This material is not final and is subject to further changes prior to publication.
 Any references or material related to answers, grades, papers or examinations are based on the opinion of the author(s).

CHEMISTRY FOR THE IB DIPLOMA: TEACHER’S RESOURCE

Homework ideas
• Exam-style questions from the Coursebook, for example, questions 11–13.
• Exercises 1.1–1.4 from the Workbook.
• Carry out a paper chromatography experiment at home to separate the dyes in sweets. An example can be
found by searching the ‘Royal Society of Chemistry’ website with the keywords ‘chromatography of sweets’.
• Use Word Art to create an image for all the keywords in this chapter.
• Create flashcards on definitions of elements, compounds and mixtures and the different techniques used for
separating components of mixtures.
• Sorting cards into solids, liquids and gases, or elements, compounds and mixtures. An example can be found
by searching the ‘Royal Society of Chemistry’ website with the keywords ‘lesson plans’ and ‘particle models’.

E
Links to digital resources
• Demonstrations on chemical changes (forming compounds from elements): Iron and sulphur – search the
‘Royal Society of Chemistry’ website with the keywords ‘iron and sulfur reaction’
• Sodium and chlorine – search the Royal Society of Chemistry website with the keywords ‘sodium
and chlorine’
•

•
PL
Experiment on freezing stearic acid: search the ‘Royal Society of Chemistry’ website with the keywords
‘freezing of stearic acid’
Simulations on particle movements during changes of states: search on phet.colorado.edu for ‘states
of matter simulation’
Home experiment to separate dyes in sweets using paper chromatography: search the ‘Royal Society of
Chemistry’ website with the keywords ‘chromatography of sweets’
Revision notes on kinetic molecular theory and the states of matter (2016 syllabus): search on ibchem.com
with the keywords ‘kinetic molecular theory’ and ‘states of matter’
Introduction to the phase diagrams: search on chemguide.co.uk with the keywords ‘phase diagram’

CROSS-CURRICULAR LINKS
M
• Maths: Basic arithmetic calculations, plotting and interpreting graphs.
• Physics: Use and convert between kelvin and Celsius temperature scales. Molecular theory of
solids, liquids and gases. Describe phase changes using particle behaviour.
• TOK: How does scientific knowledge progress?
SA

44
8 Chemistry for the IB Diploma – Bonsall © Cambridge University Press 2023
Original material © Cambridge University Press & Assessment 2023. This material is not final and is subject to further changes prior to publication.