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Cambridge IGCSE™
Chemistry
Fourth Edition
Bryan Earl
Doug Wilford
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ISBN: 978 1 3983 1050 6
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Contents
Acknowledgements vi
How to use this book ix
Scientific enquiry x
1 States of matter 1
1.1 Solids, liquids and gases 1
1.2 The kinetic particle theory of matter 2
1.3 Changes of state 4
1.4 The effects of temperature and pressure on the volume of a gas 6
1.5 Diffusion 7
2 Atoms, elements and compounds 10
2.1 Elements 10
2.2 Compounds 14
2.3 Mixtures 18
2.4 Inside atoms 20
3 Bonding and structure 30
3.1 Ionic bonding 30
3.2 Covalent bonding 38
3.3 Metallic bonding 50
4 Stoichiometry – chemical calculations 54
4.1 Relative atomic mass 54
4.2 Calculating moles 55
4.3 Moles and compounds 57
4.4 Calculating formulae 60
4.5 Moles and chemical equations 62
5 Electrochemistr y 69
5.1 Electricity and chemistry 69
5.2 Electrolysis of lead(II) bromide (Teacher demonstration) 70
5.3 Electrolysis of aluminium oxide 72
5.4 Electrolysis of aqueous solutions 75
5.5 Electrolysis of copper(II) sulfate aqueous solution 78
5.6 Fuel cells 81
5.7 Electroplating 82
6 Chemical energetics 87
6.1 Substances from petroleum 87
6.2 Fossil fuels 90
6.3 What is a fuel? 90
6.4 Alternative sources to fossil fuels 91
6.5 Exothermic and endothermic reactions 91
iii
7 Chemical reactions 100
7.1 Reactions 100
7.2 Factors that affect the rate of a reaction 101
7.3 Enzymes 107
7.4 Reversible reactions and equilibrium 108
7.5 Ammonia – an important nitrogen-containing chemical 109
7.6 Industrial manufacture of sulfuric acid – the Contact process 111
8 Acids, bases and salts 116
8.1 Acids and alkalis 116
8.2 Formation of salts 120
8.3 Methods of preparing soluble salts 121
8.4 Preparing insoluble salts 124
8.5 Testing for different salts 124
8.6 Water of crystallisation 126
9 The Periodic Table 134
9.1 Development of the Periodic Table 134
9.2 Electronic configuration and the Periodic Table 137
9.3 Group I – the alkali metals 137
9.4 Group VII – the halogens 141
9.5 Group VIII – the noble gases 145
9.6 Transition elements 145
9.7 The position of hydrogen 147
10 Metals 150
10.1 Properties of metals 150
10.2 Metal reactions 151
10.3 Reactivity of metals and their uses 153
10.4 Identifying metal ions 156
10.5 Extraction of metals 157
10.6 Metal corrosion 161
10.7 Alloys 165
11 Chemistry of the environment 170
11.1 Water 170
11.2 Artificial fertilisers 175
11.3 The air 177
11.4 Atmospheric pollution 179
12 Organic chemistry 1 188
12.1 Alkanes 188
12.2 The chemical behaviour of alkanes 191
12.3 Alkenes 194
12.4 The chemical behaviour of alkenes 196
12.5 Polymers 198
iv
13 Organic chemistry 2 206
13.1 Functional groups 206
13.2 Alcohols (R–OH) 206
13.3 Carboxylic acids 211
13.4 Esters 212
13.5 Condensation polymers 214
13.6 Natural polyamides 216
14 Experimental techniques and chemical analysis 221
14.1 Apparatus used for measurement in chemistry 221
14.2 Separating mixtures 224
14.3 Qualitative analysis 233
Theory past paper questions 240
Alternative to Practical past paper questions 255
Glossary 262
Index 269
Periodic Table of Elements 275
v
Acknowledgements
The authors would like to thank Irene, Katharine, Lambert Photography/Science Photo Library, b ©
Michael and Barbara for their patience, support and midosemsem – Fotolia; p.33 © Science Source/
encouragement throughout the production of this Science Photo Library; p.36 l © Martyn F. Chillmaid,
textbook. We would also like to thank the editorial r © Martyn F. Chillmaid; p.37 (left column) l ©
and publishing teams at Hodder Education who have Martyn F. Chillmaid, r © Martyn F. Chillmaid, (right
supported us on the journey over the past year. column) Courtesy of the University of Illinois at
Source acknowledgements Urbana-Champaign Archives; p.38 © Andrew
pp. 13, 38, 40, 41, 42, 189, 195, 196, 198, 199, 200 Lambert Photography/Science Photo Library; p.40 ©
and 213 Andrew Lambert Photography/Science Photo Library;
p.41 l © Andrew Lambert Photography/Science
The molecular models shown were made using Photo Library, r © Andrew Lambert Photography/
the Molymod® system available from Molymod® Science Photo Library; p.42 © Andrew Lambert
Molecular Models, Spiring Enterprises Limited, Photography/Science Photo Library; p.44 ©
Billingshurst, West Sussex RH14 9NF England. E.R.Degginger/Science Photo Library; p.46 l ©
Photo credits Philippe Plailly/Science Photo Library, r © Robert
r right, l left, t top, m middle, b bottom Harding Photo Library; p.48 tl © Richard Megna/
p.1 © bl Dolphinartin/stock.adobe.com, tr © Fundamental/Science Photo Library, bl © Sheila
12ee12/stock.adobe.com, br © Olaf Schubert/Image Terry/Science Photo Library, tr © Michael Pettigrew
Broker/Alamy Stock Photo; p.2 l © Donald L Fackler – Fotolia, br © Kage Mikrofotografie GBR/Science
Jr/Alamy Stock Photo, r © Andrew Lambert Photo Library; p.49 © overthehill – Fotolia; p.56 l
Photography/Science Photo Library; p.3 tl © © Andrew Lambert Photography/Science Photo
Geoscience Features Picture Library/Dr.Booth, bl © Library, r © Andrew Lambert Photography/Science
Javier Larrea/Agefotostock/Alamy Stock Photo, br © Photo Library; p.57 © Andrew Lambert Photography/
Power And Syred/Science Photo Library; p.7 l © Science Photo Library; p.63 © Dirk Wiersma/Science
Andrew Lambert Photography/Science Photo Library, Photo Library; p.69 tl © Alexander Maksimenko
r © Photiconix/Alamy Stock Vector; p.8 l © Andrew – Fotolia, tr © Mkos83/stock.adobe.com, b © Pavel
Lambert Photography/Science Photo Library, r © Losevsky/stock.adobe.com; p.72 © Howard Davies/
Andrew Lambert Photography/Science Photo Library; Alamy Stock Photo; p.73 © Kekyalyaynen/
p.10 © NASA/ESA/STSCI/Hubble Heritage Team/ Shutterstock.com; p.74 l © dpa picture alliance
Science Photo Library; p.11 t © Robert Harding/ archive / Alamy Stock Photo, r © Rex Features; p.75
Getty Images, m © FlemishDreams – Fotolia, b © l © Kathy Gould/iStock/Thinkstock, r © Trevor
Echo 23 Media/Alamy Stock Photo; p.12 tl © Bay Clifford Photography/Science Photo Library; p.78 tr
Ismoyo/AFP/Getty Images, bl © Christie’s Images / © Andrew Lambert Photography/Science Photo
The Bridgeman Art Library, tr © BL Images Ltd / Library, br © Clynt Garnham Renewable Energy /
Brian Lawrence / Alamy Stock Photo; p.13 © Andrew Alamy Stock Photo; p.81 © Frankie Martin/NASA/
Lambert Photography/Science Photo Library; p.15 © Science Photo Library; p.82 © Adam Hart-Davis/
Andrew Lambert Photography / Science Photo Science Photo Library; p.83 bl © Andrew Lambert
Library; p.18 tr © Galyna Andrushko – Fotolia, br © Photography/Science Photo Library, br © Andrew
Andrew Lambert Photography / Science Photo Lambert Photography/Science Photo Library; p.87 tl
Library; p.19 © Andrew Lambert Photography/ © Dinodia Photos / Alamy Stock Photo, bl © Oleg
Science Photo Library; p.20 (left column) tl © Zhukov – Fotolia, mt © Matt K – Fotolia, mb ©
Vvvita/stock.adobe.com, tr © Weyo/stock.adobe. Dmitry Sitin – Fotolia, tr © iofoto – Fotolia, br ©
com, bl © Shutterstock/Stay_Positive, bm © yang yu – Fotolia; p.88 l © Andrew Lambert
Magraphics/stock.adobe.com, br sommai/stock. Photography/Science Photo Library, r © Sindo
adobe.com, (right column) tr © Science Photo Farina/Science Photo Library; p.89 © Stevie Grand/
Library/Eye of Science, br © moodboard/Thinkstock; Science Photo Library; p.90 © GeoScience Features
p.21 tr © Digital Instruments/Veeco/Science Photo Picture Library; p.101 tl © Illustrez-vous - Fotolia,
Library; p.24 br © Henry Westheim Photography / bl © Arnaud Santinil – Fotolia, tr © Yakovlevadaria/
Alamy Stock Photo; p.30 tl © Andrew Lambert stock.adobe.com, br © National Motor Museum,
Photography/ Science Photo Library, tr © Andrew Motoring Picture Library / Alamy Stock Photo;
vi
p.102 l © Andrew Lambert Photography/Science Elenathewise/stock.adobe.com; p.147 tl © Andrew
Photo Library, tr © Andrew Lambert Photography/ Lambert Photography/Science Photo Library, tr ©
Science Photo Library, br © Andrew Lambert Vlorzor/stock.adobe.com; p.150 tr © Andrew
Photography/Science Photo Library; p.104 tr © Lambert Photography/Science Photo Library, br ©
Andrew Lambert Photography/Science Photo Library, Digital Vision/Getty Images; p.151 tl © Neil
bl © Andrew Lambert Photography/Science Photo Emmerson/Robert Harding Picture Library Ltd /
Library, br © Andrew Lambert Photography/Science Alamy Stock Photo, tr © Andrew Lambert
Photo Library; p.107 © Alfred Pasieka/Science Photography/Science Photo Library; p.153 tl ©
Photo Library; p.108 © Daniel Kalker/DPA/PA govicinity – Fotalia.com, ml © M.bonotto/stock.
Images; p.109 © Andrew Lambert Photography/ adobe.com, bl © Antony Nettle/Alamy Stock Photo;
Science Photo Library; p.111 © Boyer/Roger Viollet/ p.154 © Andrew Lambert Photography/Science
Getty Images; p.116 tl © Yvdavid/stock.adobe.com, Photo Library; p.155 © Leah-Anne Thompson/stock.
bl © Picture Store/stock.adobe.com, tr © adobe.com; p.156 l © Andrew Lambert Photography/
Valery121283/stock.adobe.com, br © Iurii Science Photo Library, r © Andrew Lambert
Kachkovskyi/stock.adobe.com; p.117 tl © Alex Photography/Science Photo Library; p.158 tl ©
Yeung/stock.adobe.com, ml © Leah-Anne Thompson/ GeoScience Features Picture Library, bl ©
stock.adobe.com, bl © Illustrez-vous – Fotolia, r © GeoScience Features Picture Library, r © GeoScience
Andrew Lambert Photography/Science Photo Library; Features Picture Library; p.159 Tata Steel’s Blast
p.118 tl © Andrew Lambert Photography/Science Furnace 4 at its Port Talbot works in the UK
Photo Library, bl © Andrew Lambert Photography/ (courtesy of Tata Steel UK Limited); p.160 t ©
Science Photo Library; p.119 © Alex White – Fotolia; Stockr/stock.adobe.com, b © Martin Bond/Science
p.121 © Andrew Lambert Photography/Science Photo Library; p.161 l © Hank Morgan/Science Photo
Photo Library; p.122 © Andrew Lambert Library, r © Digital Vision/Getty Images; p.162 l ©
Photography/Science Photo Library; p.123 l © Decha Helene Rogers/ Art Directors & TRIP / Alamy Stock
Thapanya/Shutterstock.com, tr © Andrew Lambert Photo, tr © Visual China Group/Getty Images, br ©
Photography/Science Photo Library, br © Andrew Aleksandr/stock.adobe.com; p.163 tl © g215 –
Lambert Photography/Science Photo Library; p.124 Fotolia, ml © Gravicapa – Fotolia, bl © Blueblood/
© Andrew Lambert Photography/Science Photo stock.adobe.com, r © steven gillis hd9 imaging/
Library; p.125 l © Andrew Lambert Photography/ Alamy Stock Photo; p.164 tl © Leslie Garland Picture
Science Photo Library, r © Andrew Lambert Library/Alamy Stock Photo, br © byggarn.se –
Photography/Science Photo Library; p.126 © Andrew Fotolia; p.165 tr © Photo20ast/stock.adobe.com, br
Lambert Photography/Science Photo Library; p.127 © Kage Mikrofotografie GBR/Science Photo Library;
© Andrew Lambert Photography/Science Photo p.170 bl © Patricia Elfreth – Fotolia, br © BSIP SA/
Library; p.135 © Science Photo Library; p.136 tl © Alamy Stock Photo; p172. tl © Can Balcioglu –
Windsor – Fotolia, bl © jon11 – Fotolia, tr © Debby Fotolia, ml © Galina Barskaya – Fotolia, bl ©
Moxon, br © Mihaela Maiorescu - Fotolia; p.137 SciencePhotos/Alamy Stock Photo, br © Andrew
© Forance.stock.adobe.com; p.138 tl © Andrew Lambert Photography/Science Photo Library; p.173
Lambert Photography/Science Photo Library, bl © © niyazz – Fotolia; p.174 l © Sze Fei Wong/iStock/
Andrew Lambert Photography/Science Photo Library, Thinkstock, r © AoshiVN/iStock/Thinkstock; p.175 ©
r © Andrew Lambert Photography/Science Photo Andrew Lambert Photography/Science Photo Library;
Library; p.139 © Charles D. Winters/Science Photo p.176 bl © Dudarev Mikhail – Fotolia; p.177 l ©
Library; p.140 © Andrew Lambert Photography/ Universal Images Group/Getty Images, r © Argus
Science Photo Library; p.141 tr © Andrew Lambert – Fotolia; p.180 tl © Astrid & Hanns-Frieder Michler/
Photography/Science Photo Library, br © Andrew Science Photo Library, bl © VisualHongKong / Alamy
Lambert Photography/Science Photo Library; p.142 Stock Photo; p.181 tr © Photodisc/Getty Images, ml
© Martyn F. Chillmaid; p.143 tl © Mark Richardson/ © Peter Treanor/Alamy Stock Photo, br © LaSa/
Alamy Stock Photo, bl © Richard B. Levine/Alamy stock.adobe.com; p.182 © Falko Matte – Fotolia;
Stock Photo, tr © Uniquely India/PhotosIndia.com p.184 © Mark Bourdillon/Alamy Stock Photo; p.188
LLC / Alamy Stock Photo, br © ZUMA Press, Inc. / © Centaur – Fotolia; p.189 tl–bl © Andrew Lambert
Alamy Stock Photo; p.144 © Science Photo Library; Photography/ Science Photo Library; p.191 tr ©
p.146 tl © Chris Leachman/stock.adobe.com, bl © nikkytok – Fotolia, br © Mr.Hackenbush/Alamy Stock
Koichi Kamoshida/Getty Images, tr © Squirrel7707/ Photo; p.192 © Andrey Popov/stock.adobe.com;
stock.adobe.com; mr © Yakov/stock.adobe.com, br © p.193 © NASA; p.195 tl–bl © Andrew Lambert
vii
Photography/Science Photo Library; p.196 © Andrew Shcheglov – Fotolia; p.221 © burnel11 – Fotolia;
Lambert Photography/Science Photo Library; p.197 p.221 tl © Martyn F Chillmaid, tr © Martyn F
© Andrew Lambert Photography/Science Photo Chillmaid, br © Martyn F Chillmaid; p.223 © Maurice
Library; p.198 © Andrew Lambert Photography/ Savage/Alamy Stock Photo; p.225 © Andrew
Science Photo Library; p.199 tl © Josie Elias/Alamy Lambert Photography/Science Photo Library; p.226
Stock Photo, tr © Andrew Lambert Photography/ bl © Vectorfusionart/stock.adobe.com, tr © Klaus
Science Photo Library, br © Martyn F. Chillmaid. Guldbrandsen/Science Photo Library, br © Klaus
Thanks to Molymod.com for providing the model; Guldbrandsen/Science Photo Library; p.227 l ©
p.200 tl © Steve Cukrov – Fotolia, bl © Andrew Andrew Lambert Photography/Science Photo Library,
Lambert Photography/Science Photo Library, br © tr © Ricardo Funari / Brazilphotos/Alamy Stock
Andrew Lambert Photography/Science Photo Library; Photo, br © Andrew Lambert Photography/Science
p.201 l © smikeymikey1 – Fotolia, r © Henning Photo Library; p.228 tl © Sipa Press / Rex Features,
Kaiser/ DPA/PA Images; p.202 tl © Sally Morgan/ bl © KPA/Zuma/Rex Features, r © Sciencephotos/
Ecoscene, bl © Andrew Lambert Photography/ Alamy Stock Photo; p.229 tr © BOC Gases, bl © Paul
Science Photo Library, tr © Richard Carey/stock. Rapson/Science Photo Library; p.230 l © Andrew
adobe.com; p.206 © M.studio – Fotolia; p.208 © Lambert Photography/Science Photo Library, r ©
Ian Dagnall/Alamy Stock Photo; p.209 tr © Biophoto Jenny Matthews/Alamy Stock Photo; p.231 ©
Associates/Science Photo Library, br © Science Andrew Lambert Library; p.233 © Doug Taylor /
Photo Library; p.210 l © Andrew Lambert Alamy Stock Photo; p.234 © Ggw/stock.adobe.com;
Photography/Science Photo Library, r © Andrew p.235 l © Turtle Rock Scientific/Science Source/
Lambert Photography/Science Photo Library; p.213 Science Photo Library m © Science Photo Library, r
tl–bl (table) © Martyn F Chillmaid. Thanks to © Giphotostock/Science Photo Library; p.236 l ©
Molymod.com for providing the models, b © Paul Andrew Lambert Photography/Science Photo Library,
Cooper / Rex Features; p.215 l © Andrew Lambert r © Martyn F. Chillmaid/Science
Photography/Science Photo Library, r © Leonid Photo Library.
viii
How to use this book As you read through the book, you will notice that
some text is shaded yellow. This indicates that the
To make your study of Chemistry for Cambridge highlighted material is Supplement content only.
IGCSE™ as rewarding and successful as possible, Text that is not shaded covers the Core syllabus. If
this textbook, endorsed by Cambridge Assessment you are studying the Extended syllabus, you should
International Education, offers the following look at both the Core and Supplement sections.
important features:
Practical skills
FOCUS POINTS
These boxes identify the key practical skills
Each topic starts with a bullet point summary of you need to understand as part of completing
what you will encounter within the topic. the course.
This is followed by a short outline of the topic Key definitions
so that you know what to expect over the next These provide explanations of the meanings of key words
few pages. as required by the syllabus.
Test yourself Worked examples
These questions appear regularly throughout the These boxes give step-by-step guidance on how to approach
chapter so you can check your understanding as you different sorts of calculations, with follow-up questions so
progress. you can practise these skills.
Revision checklist Going further
At the end of each chapter, a revision checklist will These boxes take your learning further than is
allow you to recap what you have learned in each required by the Cambridge syllabus so that you have
topic and double-check that you understand the key the opportunity to stretch yourself.
concepts before moving on.
Answers are provided online with the accompanying
Exam-style questions Cambridge IGCSE Chemistry Teacher’s Guide. A Practical
Skills Workbook is also available to further support
Each chapter is followed by exam-style questions you in developing your practical skills as part of
to help familiarise learners with the style of carrying out experiments.
questions they may see in their examinations.
These will also prove useful in consolidating your
learning. Past paper questions are also provided
at the end of the book.
ix
Scientific enquiry
Throughout your IGCSE Chemistry course, you will obtained. Most importantly, it is essential that
need to carry out experiments and investigations you carry out a risk assessment before you do any
aimed at developing some of the skills and abilities practical work.
that scientists use to solve real-life problems. 3 Make and record observations – the data you
need to answer the questions you have set
Simple experiments may be designed to measure, yourself can only be found if you have planned
for example, the temperature of a solution or the your investigation sensibly and carefully. For
rate of a chemical reaction. Longer experiments, example, you might start to use a measuring
or investigations, may be designed to allow you to cylinder to collect a gas, but as you develop
actually see the relationship between two or more your ideas you may realise a burette might be
physical quantities such as how rate of reaction more appropriate and more accurate. Be careful
varies with temperature and concentration. not to dismantle the equipment/apparatus until
you have completed your analysis of the data,
Investigations are likely to come about from the and you are sure you do not need to repeat any
topic you are currently studying in class, and your of the measurements! If you have to reset your
teacher may provide you with suggestions. For all equipment/apparatus, it may add further errors
investigations, both your teacher and you must to the results you have obtained.
consider the safety aspects of the chemicals and
apparatus involved. You should never simply carry Ensure that all your data, numeric or
out a chemistry investigation without consideration observational, is displayed in a clear format. This
of the hazards of the chemicals or getting the will often be in the form of headed tables with
approval of your teacher. the correct units being shown to the appropriate
degree of precision.
To carry out an investigation, you will need to: 4 Interpret and evaluate observations – the
1 Select and safely use suitable techniques, results you obtain from any investigation must
be displayed carefully and to the accuracy of
apparatus and materials – your aim must be to the equipment you have chosen to use. Your
safely collect sufficient evidence using the most choice of presentation will help you interpret
appropriate apparatus for the technique you have your evidence and make conclusions. Often
chosen. Being able to draw and label diagrams your presentation will be in the form of a graph
correctly to show how the equipment will be used or a table. For some graphs, you may need to
is also important. Your techniques will need to be calculate gradients or use it to find values at a
explained clearly to do a proper risk assessment. specific point during the investigation by drawing
For example, how to carry out a titration, how intercepts. Good chemists keep looking at the
you are going to follow the rate of the reaction data and alter the way in which it is obtained to
you are using, or how to test for ions and gases. get more accurate results. You should be able to
2 Plan your experiment – this is an important part evaluate whether your data is good or bad. If it is
of doing science and involves working out what good, were there any anomalous results? Why did
you are going to do to try to find answers to you get them?
the questions you have set yourself. Predictions 5 Evaluate methods and suggest possible
based on work you have been studying or improvements – at the end of your investigation,
are doing in class may help you develop the you must be able to evaluate the equipment,
investigation in terms of the number and type methods and techniques that you have used.
of observations or data needed. You will also Think about any sources of errors that could have
need to be able to identify the independent affected your results by the use of the wrong
and dependent variables. For example, if you equipment. Consider, if you were able to carry out
are trying to find out how temperature affects the investigation again, what you would change.
the rate of a reaction, the temperature will be The more data you obtain, the easier it is to spot
the independent variable, but the dependent anomalous results.
variable might be the volume of gas collected.
Other variables such as concentration need to be
controlled so that they will not affect the data
x
A written report of any chemical investigation label each axis of a graph with the name and unit
would normally be made of these fixed components: of the quantity being plotted. Make sure that the
» First, state the aim of the work at the very scale you use allows the points to fill up as much
of the graph paper as possible.
beginning to inform your teacher what you were
doing the investigation to find out. Clearly explain the calculations involved in
» A list of all items of equipment/apparatus used the interpretation of your data and give the
and a record of the smallest division of the scale significant figures appropriate to the equipment
of each measuring device you have used (see used.
Chapter 14). For example, burettes can be read » Conclusions can be obtained from the graphs
to two decimal places, to the nearest 0.05 cm3, and calculations you carry out. Your conclusions
where the second decimal place is either a 0 from the data obtained might be different from
if the bottom of the meniscus is on the scale those that you expected. Even so it is very
division, or a 5 if it is between the divisions. important for any scientist to come to terms with
If the meniscus was between 24.10 cm3 and the findings of their experimental result, good
24.20 cm3, the reading would be 24.15 cm3. or bad!
» You must show that you have considered the » In the evaluation you should make a comparison
safety of yourself and others before you carry between the conclusions of your investigation
out any practical work. Provide a list of all the and your expectations: how close or how different
chemicals you will use, as well as the ones you were they? You should comment on the reliability
will produce, and do a risk assessment to check and accuracy of the observations and the data
on all the hazards of the chemicals. The results obtained. Could you have improved the method
of your risk assessment might indicate that you to give better or more accurate results? Would a
need to work in a well-ventilated room or in a pH probe have been better than using universal
fume cupboard. In some cases you may need the indicator to find the point of neutralisation in a
assistance of your teacher. If in doubt, always ask titration? Were there any anomalous points on
your teacher for advice. your graphs, or any unusual data or observations?
» Clearly state the details of the methods used, Highlight these and try to give an explanation.
starting with the wearing of eye protection. The
methods should be shown as numbered steps and Suggestions for
should be made as clear as possible. Ideas of the investigations
number of measurements that will be made and
their frequency should be stated. Observations Some suggested investigations are outlined in this
should be clear and you should use changes book as follows:
in colour and physical state as part of your 1 Find which vinegars contain the most acid.
observations.
» Presentation of results and calculations. If you (Chapter 8)
made several measurements of a quantity, draw 2 Find the molar volume of hydrogen by reacting
up a table in which to record your results. Use
the column headings, or start of rows, to name magnesium with hydrochloric acid. (Chapter 4)
exactly what the measurement is and state the 3 Determination of the enthalpy of combustion
units used; for example, in a rates of reaction
experiment, ‘Mass of calcium carbonate/g’. of ethanol. (Chapter 6)
4 The effect of changing the surface area of a
Give numeric values to the number of
significant figures appropriate to the equipment reactant on the rate of reaction. (Chapter 7)
being used, for example, a mass could be 5 Use the anion and cation methods of
recorded to 0.5 g or 0.05 g depending on the
resolution of the top-pan balance you use. Take identification to find the ions present in tap
averages and remember that anomalous or non- water. (Chapter 14)
concordant results should not be used in their 6 Show that ammonia is a weak base by measuring
calculation. If you decide to make a graph of your its pH and conductivity, and comparing your
results, you will need at least six data points results with those from a solution of sodium
taken over as large a range as possible; be sure to hydroxide with the same concentration. (Chapter 8)
7 Determine the melting point of stearic acid.
(Chapter 1)
8 How can sodium chloride be obtained from rock
salt? (Chapter 14)
xi
9 What are the effects of acid rain on a variety of A further example involves the well-known
building materials? (Chapter 11) Russian chemist Dimitri Mendeleev. He realised that
the physical and chemical properties of the known
10 Which is the best temperature, between 34 and elements were related to their atomic mass in a
40°C, for the fermentation of sugar to take place? ‘periodic’ way, and arranged them so that groups of
(Chapter 12) elements with similar properties fell into vertical
columns in his table. However, in devising his
11 What are the chemical properties of the weak table, Mendeleev did not conform completely to the
organic acid, ethanoic acid? (Chapter 13) order of atomic mass, with some elements swapped
around. It took time for his ideas to gain acceptance
12 Do foodstuffs contain carbon? (Chapter 12) because the increase in atomic mass was not regular
when moving from one element to another. We
Ideas and evidence in science now know, with the development of atomic theory
and a better understanding of chemical processes,
In some of the investigations you perform in the that the elements in the Periodic Table are not all
school laboratory, you may find that you do not in atomic mass order. It took until 1934, with an
interpret your data in the same way as your friends understanding of atomic number and post-Russian
do; perhaps you will argue with them as to the best revolution, for the Periodic Table to be finally
way to explain your results and try to convince accepted in the form you see today.
them that your interpretation is right. Scientific
controversy frequently arises through people There are many different types of scientists
interpreting evidence differently. with specialties in their own areas of work such
as chemistry and physics, but they all work in the
For example, our ideas about atoms have changed same way. They come up with new theories and
over time. Scientists have developed new models ideas, they carry out work to find the evidence to
of atoms over the centuries as they collected new establish whether their ideas are correct and, if not,
experimental evidence. If we go back to the Greeks why. Scientists rely on other scientists checking
in the 5th century bc, they thought matter was their work, often improving the ideas of everyone
composed of indivisible building blocks which they and moving science forward. The use of new ideas
called atomos. However, the idea was essentially is often beneficial to everyone in the world, for
forgotten for more than 2000 years. Then John example, the discovery of vaccines for Covid-19, or
Dalton published his ideas about atoms in 1800. He the push to improve battery manufacture for use
suggested that all matter was made of tiny particles in electric cars which would in turn help solve one
called atoms, which he imagined as tiny spheres of the biggest problems we have to face: global
that could not be divided. It then took another warming. Scientists are working hard to stop global
100 years before Joseph Thomson, Ernest Rutherford warming but their ideas are not always embraced
and James Chadwick carried out experiments and because of economic and political factors.
discovered that there was a structure within the
atom. This saw the continuous development of what
we know today as atomic theory.
xii
1 States of matter
FOCUS POINTS
★ What is the structure of matter?
★ What are the three states of matter?
★ How does kinetic particle theory help us understand how matter behaves?
In this first chapter, you will look at the three states of matter: solids, liquids and gases. The structure
of these states of matter and how the structures can be changed from one to another is key to
understanding the states of matter.
You will use the kinetic particle theory to help explain how matter behaves, so you can understand
the difference in the properties of the three states of matter and how the properties are linked to
the strength of bonds between the particles they contain. Why, for example, can you compress gases
but cannot compress a solid? By the end of this chapter, you should be able to answer this question
and use the ideas involved to help you to understand many everyday observations, such as why car
windows mist up on a cold morning or why dew forms on grass at night.
1.1 Solids, liquids and gases
Chemistry is about what matter is like and how it
behaves, and our explanations and predictions of
its behaviour. What is matter? This word is used to
cover all the substances and materials from which
the physical universe is composed. There are many
millions of different substances known, and all of
them can be categorised as solids, liquids or gases
(Figure 1.1). These are what we call the three states
of matter.
b Liquid
a Solid c Gas
▲ Figure 1.1 Water in three different states
1
1 States of matter
A solid, at a given temperature, has a definite The main points of the theory are:
volume and shape which may be affected by changes » All matter is made up of tiny, moving particles,
in temperature. Solids usually increase slightly in size
when heated, called expansion (Figure 1.2), invisible to your eye. Different substances have
and usually decrease in size if cooled, called different types of particles (atoms, molecules or
contraction. ions) of varying sizes.
» The particles move all the time. The higher the
A liquid, at a given temperature, has a temperature, the faster they move on average.
fixed volume and will take up the shape of any » Heavier particles move more slowly than lighter
container into which it is poured. Like a solid, a ones at a given temperature.
liquid’s volume is slightly affected by changes in The kinetic particle theory can be used as a
temperature. scientific model to explain how the arrangement
of particles relates to the properties of the three
A gas, at a given temperature, has neither a states of matter.
definite shape nor a definite volume. It will take up
the shape of any container into which it is placed Explaining the states of matter
and will spread out evenly within it. Unlike solids
and liquids, the volumes of gases are affected In a solid, the particles attract one another. There
greatly by changes in temperature. are attractive forces between the particles which
hold them close together. The particles have little
Liquids and gases, unlike solids, are compressible. freedom of movement and can only vibrate about
This means that their volume can be reduced by a fixed position. They are arranged in a regular
the application of pressure. Gases are much more manner, which explains why many solids form
compressible than liquids. crystals.
It is possible to model such crystals by using
spheres to represent the particles. For example,
Figure 1.3a shows spheres built up in a regular way
to represent the structure of a chrome alum crystal.
The shape is very similar to that of a part of an
actual chrome alum crystal (Figure 1.3b).
▲ Figure 1.2 Without expansion gaps between the rails, the a A model of a chrome alum crystal
track would buckle in hot weather
1.2 The kinetic particle
theory of matter
The kinetic particle theory helps to explain the
way that matter behaves. It is based on the idea
that all matter is made up of tiny particles. This
theory explains the physical properties of matter in
terms of the movement of the particles from which
it is made.
2
1.2 The kinetic particle theory of matter
b An actual chrome alum crystal ▲ Figure 1.5 Sodium chloride crystals
▲ Figure 1.3
In a liquid, the particles are still close together but
Studies using X-ray crystallography (Figure 1.4) have they move around in a random way and often collide
confirmed how particles are arranged in crystal with one another. The forces of attraction between
structures. When crystals of a pure substance the particles in a liquid are weaker than those in
form under a given set of conditions, the particles a solid. Particles in the liquid form of a substance
are always arranged (or packed) in the same way. have more energy on average than the particles in
However, the particles may be packed in different the solid form of the same substance.
ways in crystals of different substances. For
example, common salt (sodium chloride) has its In a gas, the particles are relatively far apart.
particles arranged to give cubic crystals as shown in They are free to move anywhere within the container
Figure 1.5. in which they are held. They move randomly at very
high velocities, much more rapidly than those in a
liquid. They collide with each other, but less often
than in a liquid, and they also collide with the walls
of the container. They exert virtually no forces of
attraction on each other because they are relatively
far apart. Such forces, however, are very significant.
If they did not exist, we could not have solids or
liquids (see Changes of state, p. 4).
The arrangement of particles in solids, liquids
and gases is shown in Figure 1.6.
Solid
Particles only vibrate about fixed positions.
Regular structure.
Liquid
Particles have some freedom and can move
around each other. Collide often.
▲ Figure 1.4 A modern X-ray crystallography instrument Gas
used for studying crystal structure Particles move freely and at random in all
the space available. Collide less often than
in liquid.
▲ Figure 1.6 The arrangement of particles in solids, liquids
and gases
3
1 States of matter
Test yourself of attraction between themselves and the other
particles in the liquid and they escape to form a gas.
1 When a metal (such as copper) is heated, it The liquid begins to evaporate as a gas is formed.
expands. Explain what happens to the metal
particles as the solid metal expands. Eventually, a temperature is reached at which
the particles are trying to escape from the liquid
1.3 Changes of state so quickly that bubbles of gas actually start to form
inside the liquid. This temperature is called the
The kinetic particle theory model can be used to boiling point of the substance. At the boiling point,
explain how a substance changes from one state to the pressure of the gas created above the liquid
another. If a solid is heated, the particles vibrate equals that of the air, which is atmospheric pressure.
faster as they gain energy. This makes them ‘push’
their neighbouring particles further away. This Liquids with high boiling points have stronger
causes an increase in the volume of the solid, forces between their particles than liquids with low
such that the solid expands, and we can say that boiling points.
expansion has taken place.
When a gas is cooled, the average energy of the
Eventually, the heat energy causes the forces of particles decreases and the particles move closer
attraction to weaken. The regular pattern of the together. The forces of attraction between the particles
structure breaks down, and the particles can now now become significant and cause the gas to condense
move around each other. The solid has melted. The into a liquid. When a liquid is cooled, it freezes to form
temperature at which this takes place is called the a solid. Energy is released in each of these changes.
melting point of the substance. The temperature
of a melting pure solid will not rise until it has all Changes of state are examples of physical
melted. When the substance has become a liquid, changes. Whenever a physical change of state
there are still very significant forces of attraction occurs, the temperature remains constant during the
between the particles, which is why the substance change. During a physical change, no new substance
is a liquid and not a gas. is formed.
Solids which have high melting points have Heating and cooling curves
stronger forces of attraction between their particles
than those which have low melting points. A list of The graph shown in Figure 1.7 was drawn by
some substances with their corresponding melting plotting the temperature of water as it was
and boiling points is shown in Table 1.1. heated steadily from −15°C to 110°C. You can see
from the curve that changes of state have taken
▼ Table 1.1 Melting points and boiling points of substances place. When the temperature was first measured,
only ice was present. After a short time, the curve
Substance Melting point/°C Boiling point/°C flattens showing that even though heat energy is
Aluminium 661 2467 being put in, the temperature remains constant.
Ethanol −117
Magnesium oxide 827 79 Temperature/ºC110
Mercury −30 3627 100
Methane
Oxygen −182 357 Liquid and All
Sodium chloride −218 −164 gas (liquid gas
Sulfur 801 −183
Water 1413 water and
113 445
0 100 water
vapour)
0 All Solid and liquid All
15 solid (ice and liquid liquid
water) (liquid
water)
If a liquid is heated, the average energy of the (ice)
particles increases and the particles will move around
even faster. Some particles at the surface of the Time/minutes
liquid have enough energy to overcome the forces
▲ Figure 1.7 Graph of temperature against time for the
change from ice at −15°C to water to steam
4
1.3 Changes of state
Practical skills
Changes of state Clamp
For safe experiments/demonstrations which Thermometer
are related to this chapter, please refer to the
Cambridge IGCSE Chemistry Practical Skills Boiling tube
Workbook, which is also part of this series.
Stearic acid
Safety
Beaker
l Eye protection must be worn.
l Take care when handling and using hot water. Warm water
The apparatus on the right was set up to obtain 1 Why was it important to remove the boiling
a cooling curve for stearic acid. The stearic acid tube with the stearic acid from the water?
was placed into a boiling tube which was then
placed in a beaker of water that was heated 2 Why was the stearic acid stirred with the
to 80°C, which is above the melting point of thermometer?
stearic acid.
3 Why were temperature readings taken every
The boiling tube was then removed from the minute for 12 minutes?
beaker and the temperature of the stearic acid
was recorded every minute for 12 minutes using The following data was obtained from the
the thermometer to stir the stearic acid while it experiment:
was a liquid.
6 7 8 9 10 11 12
Time/mins 0 1 2 3 4 5 69 69 69 67 64 62 60
Temperature/°C 79 76 73 70 69 69 6 a At what temperature did the stearic acid
begin to change state?
4 Draw and label axes for plotting this data.
5 Plot the points and draw a line of best fit. b How could you tell this from your graph?
c Explain what is happening at this
temperature.
In ice, the particles of water are close together melting point. A sharp melting point therefore
and are attracted to one another. For ice to melt, indicates that it is a pure sample. The addition or
the particles must obtain sufficient energy to presence of impurities lowers the melting point.
overcome the forces of attraction between the
water particles, so that relative movement can You can find the melting point of a substance
take place. The heat energy is being used to using the apparatus shown in Figure 1.8. The
overcome these forces. addition or presence of impurities lowers the
melting point. A mixture of substances also has a
The temperature will begin to rise again lower melting point than a pure substance, and the
only after all the ice has melted. Generally, the melting point will be over a range of temperatures
heating curve for a pure solid always stops rising and not sharp.
at its melting point and gives rise to a sharp
5
1 States of matter
Rubber band Thermometer What do you think has caused the difference
Oil between the balloons in Figure 1.9? The pressure
Melting point inside a balloon is caused by the gas particles
tube striking the inside surface of the balloon. At a
Solid higher temperature, there is an increased pressure
inside the balloon (Figure 1.10). This is due to the
gas particles having more energy and therefore
moving around faster, which results in the particles
striking the inside surface of the balloon more
frequently, which leads to an increase in pressure.
Heat ▲ Figure 1.10 The gas particles striking the surface
▲ Figure 1.8 If a substance, such as the solid in the melting create the pressure
point tube, is heated slowly, this apparatus can be used to Since the balloon is made from an elastic material,
find the melting point of the substance the increased pressure causes the balloon to
stretch and the volume increases. An increase
In the same way, if you want to boil a liquid, such as in volume of a gas with increased temperature
water, you have to give it some extra energy. This can is a property of all gases. French scientist J.A.C.
be seen on the graph in Figure 1.7, where the curve Charles made an observation like this in 1781 and
levels out at 100°C – the boiling point of water. concluded that when the temperature of a gas
increased, the volume also increased at a fixed
Solids and liquids can be identified from their pressure. We can extend this idea to suggest
characteristic melting and boiling points. that changing the pressure of a fixed volume of
a gas must have an effect on the temperature of
The reverse processes of condensing and freezing the gas. If you have ever used a bicycle pump to
occur when a substance is cooled. Energy is given blow up a bicycle tyre then you may have felt the
out when the gas condenses to the liquid and the pump get hotter the more you used it. As you
liquid freezes to give the solid. use the pump you increase pressure on the air in
the pump. Such an increase in pressure causes
1.4 The effects of the gas molecules to move closer together so the
temperature and pressure molecules to collide more frequently and more
on the volume of a gas frictional forces come into play, which causes the
temperature to rise. In addition, as the molecules
▲ Figure 1.9 Temperature changes the volume of the air in a are forced closer to one another, intermolecular
balloon. Higher temperatures increase the volume of the bonds form, again increasing the temperature of
balloon and cold temperatures reduce its volume. the gas. As the temperature of the gas increases,
this also causes the molecules to move faster,
6 causing even more collisions.
Test yourself 1.5 Diffusion
2 Why do gases expand more than solids for the Diffusion can be explained by the kinetic
same increase in temperature? particle theory. This theory states that all matter
is made up of many small particles which are
3 Ice on a car windscreen will disappear as you constantly moving. In a solid, as we have seen,
drive along, even without the heater on. Explain the particles simply vibrate about a fixed point.
why this happens. However, in a gas, the particles move randomly
past one another, colliding with each other.
4 When salt is placed on ice, the ice melts.
Explain why this happens.
5 Draw and label a graph of water at 100°C being
allowed to cool to −5°C.
1.5 Diffusion ▲ Figure 1.12 Hydrochloric acid (left) and ammonia
(right) diffuse at different rates
When you go through the door of a restaurant, you
can often smell the food being cooked. For this Gases diffuse at different rates. If one piece of
to happen, gas particles must be leaving the pans cotton wool is soaked in concentrated ammonia
the food is being cooked in and be spreading out solution and another is soaked in concentrated
through the air in the restaurant. This spreading out hydrochloric acid and these are put at opposite
of a gas is called diffusion and it takes place in a ends of a dry glass tube, then after a few minutes
haphazard and random way. a white cloud of ammonium chloride appears.
Figure 1.12 shows the position at which the two
All gases diffuse to fill the space available. gases meet and react. The white cloud forms in
Figure 1.11 shows two gas jars on top of each other. the position shown because the ammonia particles
Liquid bromine has been placed in the bottom gas are lighter; they have a smaller relative molecular
jar (left photo) and then left for a day (right photo). mass (Chapter 4, p. 54) than the hydrogen chloride
The brown-red fumes are gaseous bromine that has particles (released from the hydrochloric acid) and
spread evenly throughout both the gas jars from the so move faster, such that the gas diffuses more
liquid present in the lower gas jar. quickly. (See Chemistry Practical Skills Workbook for
more detail of this experiment.) This experiment
▲ Figure 1.11 After 24 hours the bromine fumes have is a teacher demonstration only, which must be
diffused throughout both gas jars carried out in a fume cupboard. If considering
carrying out this practical, teachers should refer
to the Practical Skills Workbook for full guidance
and safety notes.
7
1 States of matter
Diffusion also takes place in liquids (Figure 1.13) Revision checklist
but it is a much slower process than in gases. This
is because the particles of a liquid move much more After studying Chapter 1 you should be able to:
slowly. ✔ State the three states of matter and describe the
▲ Figure 1.13 Diffusion of green food colouring can take structure arrangement of the particles in each.
days to reach the stage shown on the right ✔ Explain the properties of the three states of
When diffusion takes place between a liquid and matter using ideas about the separation and
a gas, it is known as intimate mixing. Kinetic movement of particles.
particle theory can be used to explain this process. ✔ Name the changes of state and describe what
It states that collisions are taking place randomly happens to the particles in a substance when they
between particles in a liquid or a gas and that there occur.
is sufficient space between the particles of one
substance for the particles of the other substance to ✔ Explain what is happening when a substance
move into. changes state.
Test yourself ✔ Describe what happens to a given amount
of gas when temperature and/or pressure
6 When a jar of coffee is opened, people can often decreases and increases.
smell it from anywhere in the room. Use the
kinetic particle theory to explain how this happens. ✔ Use the kinetic particle theory to explain the
effects of pressure and temperature on the
7 Describe, with the aid of diagrams, the diffusion of volume of a gas.
a drop of green food colouring added to the bottom
of a beaker. ✔ Describe the process of diffusion and explain
why gases diffuse.
8 Explain why diffusion is faster in gases than in
liquids. ✔ Describe and explain the rate of diffusion of a
gas in terms of its relative molecular mass.
9 Explain why a gas with a low relative molecular
mass can diffuse faster than a gas with a
high relative molecular mass at the same
temperature.
8
Exam-style questions
Exam-style questions
1 a Sketch diagrams to show the arrangement of a Describe what you would observe after:
particles in: i a few hours [1]
i solid oxygen [1] ii several days. [1]
ii liquid oxygen [1]
iii oxygen gas. [1] b Explain your answer to Question 5a using
b Describe how the particles move in these your ideas of the kinetic particle theory. [2]
three states of matter. [3] c State the physical process that takes place in
this experiment. [1]
c Explain, using the kinetic particle theory,
what happens to the particles in oxygen
as it is cooled down. [3] 6 The apparatus shown below was set up.
2 Explain the meaning of each of the following Stopper White cloud Glass tube
terms. In your answer include an example to help
with your explanation.
a expansion [2]
b contraction [2]
c physical change [2]
d diffusion [2] Cotton wool soaked Cotton wool soaked
e random motion [2] in concentrated in concentrated
hydrochloric acid ammonia solution
3 a Explain why solids do not diffuse. [2] When this apparatus is used, the following
b Give two examples of diffusion of gases
and liquids found in your house. [2] things are observed. Explain why each of these
is observed.
4 Explain the following, using the ideas you have a A white cloud is formed. [1]
learned about the kinetic particle theory:
a When you take a block of butter out of the b It took a few minutes before the white
cloud formed. [1]
fridge, it is quite hard. However, after 15 c The white cloud formed further from the
minutes it is soft enough to spread. [2]
b When you come home from school and open cotton wool soaked in ammonia than that
soaked in hydrochloric acid. [2]
the door, you can smell food being cooked. [2] d Cooling the concentrated ammonia and
c A football is blown up until it is hard on
a hot summer’s day. In the evening the hydrochloric acid before carrying out the
experiment increased the time taken for
football feels softer. [2] the white cloud to form. [1]
d When a person wearing perfume enters a
room, it takes several minutes for the smell
to reach the back of the room. [2]
5 Some green food colouring was carefully added to
the bottom of a beaker of water using a syringe.
The beaker was then covered and left for several
days.
Beaker
Water
Green
food
colouring
9
2 Atoms, elements and compounds
FOCUS POINTS
★ How are elements, molecules, ions, compounds and mixtures different from each other?
★ How do the properties of the particles in an atom lead to an atom’s structure?
★ What do oxidation and reduction mean?
★ What is an isotope?
In Chapter 1, you saw that all matter is made up of particles. In this chapter you will look closely at
these particles and see that they are made up of atoms. Atoms are the smallest part of elements.
An element is made up of one type of atom and can be either a metal or a non-metal. Metals and
non-metals have different properties.
You will look at how atoms of different elements can combine to form substances called compounds,
and how this combining occurs in a chemical reaction. By the end of the chapter you should be able to
write a simple word or symbol equation to represent these reactions.
You will see that although atoms are the smallest part of an element that shares the chemical
properties of that element, they are made from even smaller particles. By learning about the
properties and behaviour of these smaller particles (electrons, protons and neutrons), you will be
able to see how they affect the chemical properties of elements and compounds.
The universe is made up of a very large number of In 1803, John Dalton suggested that each element
substances (Figure 2.1), and our own part of the was composed of its own kind of particles, which he
universe is no exception. When we examine this vast called atoms. Atoms are much too small to be seen.
array of substances more closely, it is found that they We now know that about 20 × 106 of them would
are made up of some basic substances which were stretch over a length of only 1 cm.
given the name elements in 1661 by Robert Boyle.
2.1 Elements
▲ Figure 2.1 Structures in the universe, such as stars,
planets and meteorites, are made of millions of As well as not being able to be broken down into
substances. These are made up mainly from just 91 a simpler substance, each element is made up of
elements, all of which occur naturally on the Earth only one kind of atom. The word atom comes from
the Greek word atomos meaning ‘unsplittable’. For
example, aluminium is an element which is made up
of only aluminium atoms. It is not possible to obtain
a simpler substance chemically from the aluminium
atoms. You can only combine it with other elements
to make more complex substances, such as aluminium
oxide, aluminium nitrate or aluminium sulfate.
One hundred and eighteen elements have now
been identified. Twenty of these do not occur in
nature and have been made artificially by scientists.
They include elements such as curium and flerovium.
Ninety-eight of the elements occur naturally and
range from some very reactive gases, such as
fluorine and chlorine, to gold and platinum, which
10
2.1 Elements
are unreactive elements. A physical property is any ▼ Table 2.1 Physical data for some metallic and
characteristic of a substance that we can measure. non-metallic elements at room temperature and
The elements have different properties that we can pressure
measure, and we can then classify them according to
those properties. Element Metal or Density/ Melting Boiling
Aluminium non-metal g cm−3 point/°C point/°C
All elements can be classified according to their
various properties. A simple way to do this is to Metal 2.70 660 2580
classify them as metals or non-metals (Figures 2.2
and 2.3). Table 2.1 shows the physical property Copper Metal 8.92 1083 2567
data for some common metallic and non-metallic
elements. You will notice from Table 2.1 that many Gold Metal 19.29 1065 2807
metals have high densities, high melting points and
high boiling points, and that most non-metals have Iron Metal 7.87 1535 2750
low densities, low melting points and low boiling
points. Table 2.2 summarises the different properties Lead Metal 11.34 328 1740
of metals and non-metals.
Magnesium Metal 1.74 649 1107
Nickel Metal 8.90 1453 2732
Silver Metal 10.50 962 2212
Zinc Metal 7.14 420 907
Carbon Non-metal 2.25 Sublimes at 3642
Hydrogen Non-metal 0.07a −259 −253
Nitrogen Non-metal 0.88b −210 −196
Oxygen Non-metal 1.15c −218 −183
Sulfur Non-metal 2.07 113 445
Source: Earl B., Wilford L.D.R. Chemistry data book. Nelson
Blackie, 1991 a: at −254°C; b: at −197°C; c: at −184°C.
a Gold is very decorative The elements also have chemical properties, which
are characteristics or behaviours that may be
b Aluminium has many uses in the aerospace industry observed when the substance undergoes a chemical
change or reaction. A discussion of the chemical
c These coins contain nickel properties of some metals and non-metals is given
▲ Figure 2.2 Some metals in Chapters 9 and 10.
▼ Table 2.2 How the properties of metals and non-metals
compare
Property Metal Non-metal
Solid, liquid or gas
Physical state at Usually solid
room temperature (occasionally Poor – usually
soft or brittle
liquid) Dull
Malleability Good Usually low
Usually low
Ductility Good Usually low
Very poor
Appearance Shiny (lustrous)
(solids)
Melting point Usually high
Boiling point Usually high
Density Usually high
Conductivity Good
(thermal and
electrical)
11
2 Atoms, elements and compounds
Test yourself c Neon is used in advertising signs
▲ Figure 2.3 Some non-metals
1 Using Tables 2.1 and 2.2, pick the ‘odd one out’ in
the following group and explain why it is different Atoms – the smallest particles
from the others.
Everything is made up of billions of extremely small
zinc copper oxygen lead atoms. The smallest atom is hydrogen, and we
2 Using Tables 2.1 and 2.2, pick the ‘odd one out’ in represent each hydrogen atom as a sphere having
a diameter of 0.000 000 07 mm (or 7 × 10−8 mm)
the following group and explain why it is different (Table 2.3). Atoms of different elements have
from the others. different diameters as well as different masses.
carbon nitrogen iron sulfur
3 Using Tables 2.1 and 2.2, pick the ‘odd one out’ in ▼ Table 2.3 Sizes of atoms
the following group of properties of metals and
explain why it is different from the others. Atom Diameter of atom/mm Masses/g
– high melting point Hydrogen 7 × 10−8 1.67 × 10−24
– high density Oxygen 12 × 10−8 2.66 × 10−23
– soft or brittle Sulfur 5.32 × 10−23
– good electrical conductivity 20.8 × 10−8
a A premature baby needs oxygen Chemists use shorthand symbols to label the
elements and their atoms. The symbol consists of
b Artists often use charcoal (carbon) to produce an initial one, two or three letters, the first of which is always
sketch a capital. The initial letter of the element’s name
is often used and, where several elements have the
12 same initial letter, another letter from the name
is added. For example, C is used for carbon, Ca for
calcium and Cl for chlorine. Some symbols seem to
have no relationship to the name of the element,
for example, Na for sodium and Pb for lead. These
symbols come from their Latin names: natrium
for sodium and plumbum for lead. A list of some
common elements and their symbols is given in
Table 2.4.
2.1 Elements
▼ Table 2.4 Some common elements and their symbols. Molecules
The Latin names of some of the elements are given in
brackets The atoms of some elements are joined together
in small groups. These small groups of atoms are
Physical state at room called molecules. The atoms of some elements
temperature and are always joined in pairs, for example, hydrogen,
Element Symbol pressure oxygen, nitrogen, fluorine, chlorine, bromine and
Aluminium Al Solid iodine. They are known as diatomic molecules. In
Argon Ar Gas chemical shorthand the molecule of chlorine shown
Barium Ba Solid ointhFeigr uerleem2.e4ntiss,wsruicthteansapshCols2.phTohreuastaonmds of some
Boron B Solid sulfur, join
Bromine Br Liquid in larger numbers, four and eight respectively, which
Calcium Ca Solid we write as P4 and S8.
Carbon C Solid
Chlorine Cl Gas The complete list of the elements with their
Chromium Cr Solid corresponding symbols is shown in the Periodic
Copper (Cuprum) Cu Solid Table on p. 135.
Fluorine F Gas The gaseous elements helium, neon, argon,
Germanium Ge Solid krypton, xenon and radon (which are all gases at 0°C
Gold (Aurum) Au Solid at sea level and atmospheric pressure) are composed
Helium He Gas of separate, individual atoms. When an element
Hydrogen H Gas exists as separate atoms, then the molecules are
Iodine I Solid said to be monatomic. In chemical shorthand these
Iron (Ferrum) Fe Solid monatomic molecules are written as He, Ne, Ar, Kr,
Lead (Plumbum) Pb Solid Xe and Rn respectively.
Magnesium Mg Solid
Mercury (Hydragyrum) Hg Liquid Cl Cl
Neon Ne Gas
Nitrogen N Gas a Represented by a letter-and-stick model
Oxygen O Gas
Phosphorus P Solid b Represented by a space-filling model
Potassium (Kalium) K Solid ▲ Figure 2.4 A chlorine molecule
Silicon Si Solid
Silver (Argentum) Ag Solid
Sodium (Natrium) Na Solid
Sulfur S Solid
Tin (Stannum) Sn Solid
Zinc Zn Solid
13
2 Atoms, elements and compounds
Molecules are not always formed by atoms of the
same type joining together as elemental molecules.
Most molecules consist of atoms of different
elements, for example, water exists as molecules
containing oxygen and hydrogen atoms. We will
learn more about these in the next section.
Test yourself b A model of water showing 2 H atoms and one O atom.
Models such as this can be built to show what a compound
4 How many atoms of hydrogen would have to be looks like
placed side by side along the edge of your ruler to
fill just one of the 1 mm divisions? ▲ Figure 2.5
5 How would you use chemical shorthand to write Water molecules contain two atoms of hydrogen and
a representation of the molecules of iodine and one atom of oxygen, and water has the chemical
fluorine? finortmheulmaoHle2Oc.ulIef,tnhoerneuims boenrlyisonreeqautiroemd of an element
in the
6 Using the Periodic Table on p. 135, write down the formula, only its symbol, as in the case of oxygen in
symbols for each of these elements and give their the water molecule H2O.
physical state at room temperature.
a chromium Test yourself
b krypton
c osmium 7 What is the formula for the molecule shown in the
diagram which contain carbon (black sphere) and
2.2 Compounds hydrogen (white spheres)?
Compounds are pure substances which are formed
when two or more elements chemically combine
together. A pure substance is a material that has
a constant composition (is homogeneous) and has
consistent properties throughout. Water is a simple
compound formed from the elements hydrogen and
oxygen (Figure 2.5). This combining of the elements
can be represented by a word equation:
hydrogen + oxygen → water
Hydrogen Oxygen Hydrogen and Water Elements other than hydrogen will also react
a pure a pure oxygen mixed a pure with oxygen to form compounds called oxides.
element element together compound For example, magnesium reacts violently with
formed from oxygen gas to form the white powder magnesium
hydrogen oxide (Figure 2.6). This reaction is accompanied
burning in by a release of energy as new chemical bonds are
oxygen formed.
O H O
H H
HH O HH O
O O
H
HH H O HH HH
H H O
H OO H
O H H O
H H H
a The element hydrogen reacts with the element oxygen to
produce the compound water
14
2.2 Compounds
Any chemical process that involves reduction and
oxidation is known as a redox reaction. For example,
to extract iron from iron(III) oxide, the oxygen has
to be removed. The reduction of iron(III) oxide can
be done in a blast furnace using carbon monoxide.
The iron(III) oxide loses oxygen to the carbon
monoxide and is reduced to iron.
Carbon monoxide is the reducing agent. A
reducing agent is a substance that reduces
another substance during a redox reaction. In the
reaction, carbon monoxide is oxidised to carbon
dioxide by the iron(III) oxide. In this process,
the iron(III) oxide is the oxidising agent. An
oxidising agent is a substance which oxidises
another substance during a redox reaction.
We can write the redox reaction as:
iron(III) oxide + carbon monoxide → iron + carbon
dioxide
▲ Figure 2.6 Magnesium burns brightly in oxygen to Test yourself
produce magnesium oxide
8 Zinc is extracted from its ore zinc blende in a
When a new substance is formed during a chemical furnace by a redox reaction. What does the term
reaction, a chemical change has taken place. ‘redox reaction’ mean?
magnesium + oxygen → magnesium oxide 9 Identify the oxidising and reducing agents in
When substances such as hydrogen and magnesium each of the following reactions:
combine with oxygen in this way, they are said to have a copper(II) oxide + hydrogen → copper + water
been oxidised, and this process is known as oxidation. b tin(II) oxide + carbon → tin + carbon dioxide
c PbO(s) + H2(g) → Pb(s) + H2O(l).
Reduction is the opposite of oxidation. In this
process oxygen is removed rather than added. For a further discussion of oxidation and reduction
see Chapter 3 (p. 31) and Chapter 5 (p. 71).
Key definitions
Oxidation is gain of oxygen. Key definitions
Reduction is loss of oxygen. Redox reactions involve simultaneous oxidation and
reduction.
An oxidising agent is a substance that oxidises another
substance and is itself reduced.
A reducing agent is a substance that reduces another
substance and is itself oxidised.
15
2 Atoms, elements and compounds
Practical skills
Heating copper Mass of crucible = 12.90 g
Mass of crucible + copper = 14.18 g
For safe experiments/demonstrations which
are related to this chapter, please refer to the Mass of copper = g
Cambridge IGCSE Chemistry Practical Skills
Workbook, which is also part of this series. Mass of crucible + contents after heating = 14.30 g
Safety Colour of contents after heating = black
● Eye protection must be worn. 1 Draw a labelled diagram of the experimental
● Take care when handling hot apparatus. set-up used in this experiment.
● Handle the copper with tongs or tweezers, not
2 Calculate the change in mass that has taken
your fingers. place during the heating.
A student wants to find out what happens when 3 Explain what has caused the change in mass.
copper is heated in air. In order to do this, 4 What is the black substance left on the copper
they carried out the following experiment and
recorded their results. after heating?
5 Write a word and balanced chemical equation
● First, they found the mass of an empty crucible
(a suitably prepared beer-bottle top (metal) is to show the process that has taken place.
an alternative to a porcelain crucible). 6 a How could you modify the experiment
● They added a piece of copper to the crucible to ensure there was no loss of substance
and found the mass again. taking place during the heating process?
b What are the other possible sources of
● They then heated the crucible strongly for error?
approximately two minutes. 7 Predict what would happen, in terms of mass
change and colour change, if calcium were
● After they allowed it to cool, they then found heated in air in the same way as the copper.
the mass after heating.
Formulae Test yourself
The formula of a compound is made up from the 10 Write down the ratio of the atoms present in the
symbols of the elements that make up the compound formula for each of the compounds shown in
and numbers that show the ratio of the different Table 2.5.
atoms the compound is made from. Carbon dioxide
has tchaerbfoonrmautolamCfOo2r, which tells you that it contains ▼ Table 2.5 Names and formulae of some common
one every two oxygen atoms. The 2 compounds
in the formula indicates that there are two oxygen Compound Formula
atoms present in each molecule of carbon dioxide. Ammonia
Calcium hydroxide NH3
Worked example Carbon dioxide Ca(OH)2
Copper sulfate CO2
Write the ratio of atoms in sodium sulfate – Na2SO4. Ethanol (alcohol) CuSO4
Glucose C2H5OH
Substance Formula Ratio of atoms Hydrochloric acid C6H12O6
Sodium sulfate Na2SO4 Na : S : O Nitric acid HCl
2 : 1 : 4 Sodium carbonate
Sodium hydroxide HNO3
Sulfuric acid Na2CO3
NaOH
H2SO4
16
2.2 Compounds
The ratio of atoms within a chemical compound Balanced equations often include symbols that
is usually constant. Compounds are made up of show the physical state of each of the reactants and
fixed proportions of elements: they have a fixed products:
composition. Chemists call this the Law of constant (s) = solid
composition. (l) = liquid
(g) = gas
For further discussion of formulae, see p. 35. (aq) = aqueous (water) solution
We can use the reaction between iron and sulfur as an
Balancing chemical equations example. The word equation to represent this reaction is:
Word equations are a useful way of representing iron + sulfur heat→ iron(II) sulfide
chemical reactions, but a better and more useful When we replace the words with symbols for the
way of seeing what happens during a chemical reactants and the products, and include their
reaction is to produce a balanced chemical physical state symbols, we get:
equation. This type of equation gives the formulae
of the substances that are reacting, the reactants, FeF(es()s +F) e+( SsS)(s(+)s)S (s)heat→heatF→eS(Fse)S(s)
and the new substances formed during the Since there is the same number of each type of atom
chemical reaction, the products, as well as showing on both sides of the equation, this is a balanced
the relative numbers of each of the particles chemical equation.
involved.
Worked example
Write, for the reaction between magnesium and oxygen change the formula of magnesium oxide, so to produce
producing magnesium oxide: the necessary two oxygen atoms on the right-hand side,
a the word equation we will need 2MgO – this means 2 × MgO formula units.
b the balanced chemical equation. The equation now becomes:
a The word equation is: Mg(s) + O2 (g) heat→ 2MgO(s)
magnesium + oxygen heat→ magnesium oxide There are now two atoms of magnesium on the right-
hand side and only one on the left. To balance the
b When we replace the words with symbols for the reactants equation, we place a 2 in front of the magnesium, and
and the products and include their physical state symbols, obtain the following balanced chemical equation:
it is important to remember that oxygen is a diatomic
molecule: 2Mg(s) + O2 (g) heat→ 2MgO(s)
Mg(s) + O2 (g) heat→ MgO(s) This balanced chemical equation now shows us that
two atoms of magnesium react with one molecule
In the equation there are two oxygen atoms on the left- of oxygen gas when heated to produce two units of
hand side (O2) but only one on the right (MgO). We cannot magnesium oxide.
Test yourself
11 Write the word and balanced chemical equations
for the reactions which take place between:
a calcium and oxygen
b copper and oxygen.
17
2 Atoms, elements and compounds
Going further ▲ Figure 2.7 Sea water is a common mixture. It is a water
solution of substances such as sodium chloride as well as
Instrumental techniques gases such as oxygen and carbon dioxide
Elements and compounds can be detected and The difference between mixtures
identified by a variety of instrumental methods. and compounds
Scientists have developed instrumental techniques
that allow us to probe and discover which elements There are differences between compounds and
are present in a substance as well as how the atoms mixtures, which can be seen by looking at the
are arranged within the substance. reaction between iron filings and sulfur. A mixture
Many of the instrumental methods that have been of iron filings (powdered iron) and sulfur (Figure 2.8
developed are quite sophisticated. Some methods are bottom right), looks different from either of the
suited to identifying elements. For example, atomic individual elements (Figure 2.8 top). This mixture
absorption spectroscopy allows the element to be has the properties of both iron and sulfur; for
identified and also allows the quantity of the element example, a magnet can be used to separate the iron
that is present to be found. filings from the sulfur (Figure 2.9).
Some methods are particularly suited to the
identification of compounds. For example, infrared
spectroscopy is used to identify compounds by showing
the presence of particular groupings of atoms.
Infrared spectroscopy is used by the pharmaceutical
industry to identify and discriminate between drugs
that are similar in structure, for example, penicillin-
type drugs. Used both with organic and inorganic
molecules, this method assumes that each compound
has a unique infrared spectrum. Samples can be
solid, liquid or gas and are usually tiny. However, Ne,
He, O2, N2 or H2 cannot be used.
This method is also used to monitor environmental
pollution and has biological uses in monitoring tissue
physiology including oxygenation, respiratory status
and blood-flow damage.
Forensic scientists make use of both these techniques
because they are very accurate, but they only require
tiny amounts of sample – often only small amounts of
sample are found at crime scenes. Other techniques
utilised are nuclear magnetic resonance spectroscopy
and ultraviolet/visible spectroscopy.
2.3 Mixtures ▲ Figure 2.8 The elements sulfur and iron at the top of the
photograph, and (below) black iron(II) sulfide on the left
Many everyday things are not pure substances: and a mixture of the two elements on the right
they are mixtures. A mixture contains more than
one substance, which could be elements and/or Substances in a mixture have not undergone a
compounds. Examples of common mixtures are: chemical reaction and it is possible to separate
» sea water (Figure 2.7) them, provided that there is a suitable difference
» air, which is a mixture of elements such as in their physical properties. If the mixture of iron
oxygen, nitrogen and neon, and compounds such
as carbon dioxide (see Chapter 11, p. 178)
» alloys such as brass, which is a mixture of copper
and zinc (for a further discussion of alloys, see
Chapter 10, p. 165).
18
2.3 Mixtures
and sulfur is heated, a chemical reaction occurs reaction, which is called an endothermic reaction
and a new substance is formed. The product of the (Chapter 6, p. 92). You will learn more about the
reaction is iron(II) sulfide (Figure 2.8 bottom left), different types of reactions in Chapter 6.
and the word equation for this reaction is:
▼ Table 2.6 Different properties of iron, sulfur, an iron/
iron + sulfur heat→ iron(II) sulfide sulfur mixture and iron(II) sulfide
Substance Appearance Effect of a Effect of dilute
Iron magnet
Dark grey hydrochloric
powder Attracted
to it acid
Very little action
when cold. When
warm, a gas is
produced with a
lot of bubbling
(effervescence)
Sulfur Yellow None No effect when
powder hot or cold
Iron/sulfur Dirty yellow Iron powder Iron powder
mixture powder attracted reacts as above
to it
▲ Figure 2.9 A magnet will separate the iron from the Iron(II) Black solid No effect A foul-smelling
mixture sulfide gas is produced
with some
During the reaction, heat energy is given out as effervescence
new chemical bonds are formed. This is called an
exothermic reaction and accompanies a chemical In iron(II) sulfide, FeS, one atom of iron has
change (Chapter 6, p. 92). The iron(II) sulfide combined with one atom of sulfur. In a mixture of
formed has very different properties to the mixture iron and sulfur, no such ratio exists as the atoms
of iron and sulfur (Table 2.6); for example, iron(II) have not chemically combined. Table 2.7 compares
sulfide would not be attracted towards a magnet. mixtures and compounds. Some common mixtures
Some chemical reactions take in heat during the are discussed in Chapter 10 (p. 165) and Chapter 11
(p. 177).
▼ Table 2.7 The major differences between mixtures and compounds
Mixture Compound
It contains two or more substances.
The composition can vary. It is a single substance.
No chemical change takes place when a mixture is formed.
The composition is always the same.
The properties are those of the individual elements/
compounds. When the new substance is formed it involves chemical
The components may be separated quite easily by physical change.
means.
The properties are very different to those of the component
elements.
The components can only be separated by one or more
chemical reactions.
Test yourself 13 Which of the following are not mixtures: milk, tin,
sulfur, cough linctus, brass, gold?
12 Make a list of some other common mixtures and
then use your research to find out and state what
they are mixtures of.
19
2 Atoms, elements and compounds
Going further
Other types of mixtures The glass fibres give the plastic extra strength so that it
does not break when it is bent or moulded into shape. The
There are mixtures which are formed by mixing two finished material has the strength and flexibility of the
substances (or phases) which cannot mix. Gels, sols, glass fibres as well as the lightness of plastic (Figure 2.11).
foams and emulsions are all examples of just such
mixtures. Look closely at the substances in Figure 2.10,
which shows examples of these different types of mixture.
ab
c
a Glass-reinforced plastic (GRP) consists of glass fibres
(rod shapes) embedded in plastic, in this case polyester
d
a This jelly is an example of a ‘gel’ b The glass-reinforced plastic used to make boats like
b Emulsion paint is an example of a ‘sol’ this is a composite material
c These foams have been formed by trapping bubbles of
gas in liquids ▲ Figure 2.11
d Emulsions are formed by mixing immiscible liquids
▲ Figure 2.10
Composite materials are those that combine the
properties of two constituents in order to get the
exact properties needed for a particular job. Glass-
reinforced fibre is an example of a composite material
combining the properties of two different materials. It
is made by embedding short fibres of glass in a matrix
of plastic.
2.4 Inside atoms any one element were identical and that each
atom was ‘indivisible’. Scientists in those days
Everything you see around you is made out of also believed that atoms were solid particles, like
tiny particles, which we call atoms (Figure 2.12). marbles.
When John Dalton developed his atomic theory,
over 200 years ago, he stated that the atoms of
20
2.4 Inside atoms
Electrons are tiny and relatively light; approximately
1837 electrons are equal in mass to the mass of
one proton or one neutron. A summary of each type
of particle, its mass and relative charge is shown
in Table 2.8. You will notice that the masses of all
these particles are measured in atomic mass units
(amu). This is because they are so light that their
masses cannot be measured usefully in grams.
▼ Table 2.8 Characteristics of a proton, a neutron and an
electron
▲ Figure 2.12 An electron micrograph of atoms taken using Particle Symbol Relative mass/ Relative charge
a very powerful microscope called an electron microscope Proton p amu +1
Neutron n 1 0
However, in the last hundred years or so it has Electron e 1 −1
been proved by great scientists, such as Niels Bohr, 1/1837
Albert Einstein, Henry Moseley, Joseph Thomson,
Ernest Rutherford and James Chadwick, that atoms Although atoms contain electrically charged
are in fact made up of even smaller ‘sub-atomic’ particles, the atoms themselves are electrically
particles. Seventy sub-atomic particles have now neutral (they have no overall electric charge). This
been discovered, and the most important of these is because atoms contain equal numbers of electrons
are electrons, protons and neutrons. and protons. For example, Figure 2.14 represents
the atom of the non-metallic element helium. The
These three sub-atomic particles are found in atom of helium possesses two protons, two neutrons
distinct and separate regions of the atom. The and two electrons. The electrical charge of the
protons and neutrons are found in the centre of the protons in the nucleus is, therefore, balanced by the
atom, which is called the nucleus. Neutrons have opposite charge of the two electrons.
no charge and protons are positively charged. The
nucleus occupies only a very small volume of the Key
atom and is very dense. Electrons
Protons
The rest of the atom surrounding the nucleus is Neutrons
where electrons are found. Electrons are negatively
charged and move around the nucleus very quickly ▲ Figure 2.14 An atom of helium has two protons, two
at specific distances from the nucleus in electron electrons and two neutrons
shells or energy levels. The electrons are held in
the shells within the atom by an electrostatic force Proton number and mass number
of attraction between themselves and the positive
charge of protons in the nucleus (Figure 2.13). Each The number of protons in the nucleus of an atom is
shell can contain only a fixed number of electrons: called the proton number (or atomic number) and is
the first shell can hold up to two electrons, the given the symbol Z. The helium atom in Figure 2.14,
second shell can hold up to eight electrons, the for example, has a proton number of 2, since it has
third shell can hold up to 18, and so on. two protons in its nucleus. Each element has its own
proton number and no two elements have the same
Nucleus, containing proton number. For example, the element lithium
neutrons and protons has a proton number of 3 since it has three protons
Electron shells in its nucleus.
▲ Figure 2.13 Diagram of an atom
21
2 Atoms, elements and compounds
Neutrons and protons have a similar mass, The number of neutrons present can be calculated
whereas electrons possess very little mass. So the by rearranging the relationship between the proton
mass of any atom depends on the number of protons number, mass number and number of neutrons to
and neutrons in its nucleus. The total number give:
of protons and neutrons found in the nucleus of
an atom is called the mass number (or nucleon number of neutrons = mass number − proton number
number) and is given the symbol A. (A) (Z)
Key definitions Worked example
Proton number or atomic number is the number of
protons in the nucleus of an atom. 1 What is the number of neutrons in one atom of 2142Mg?
number of neutrons = mass number − proton number
Mass number or nucleon number is the total number of 12 = 24(A) – 12(Z)
protons and neutrons in the nucleus of an atom.
2 What is the number of neutrons in one atom of 20872Pb?
mass number = proton number + number of neutrons number of neutrons = mass number − proton number
(A) (Z) 125 = 207(A) – 82(Z)
The helium atom in Figure 2.14 has a mass number Table 2.9 shows the number of protons, neutrons and
of 4, since it has two protons and two neutrons in electrons in the atoms of some common elements.
its nucleus. If we consider the metallic element
lithium, it has three protons and four neutrons in its
nucleus. It therefore has a mass number of 7.
The proton number and mass number of an
element are usually written in the following
shorthand way:
Mass number (A ) ←←42He ← Symbol of the element
Proton number (Z )
▼ Table 2.9 Number of protons, neutrons and electrons in some elements
Element Symbol Proton Number of Number of Number of Mass number
Hydrogen H number electrons protons neutrons 1
Helium He 1 1 1 0 4
Carbon C 2 2 2 2 12
Nitrogen N 6 6 6 6 14
Oxygen O 7 7 7 7 16
Fluorine F 8 8 8 8 19
Neon Ne 9 9 9 10 20
Sodium Na 10 10 10 10 23
Magnesium Mg 11 11 11 12 24
Sulfur S 12 12 12 12 32
Potassium K 16 16 16 16 39
Calcium Ca 19 19 19 20 40
Iron Fe 20 20 20 20 56
Zinc Zn 26 26 26 30 65
30 30 30 35
22
2.4 Inside atoms
Ions ▼ Table 2.10 Some common ions
An ion is an electrically charged particle. When an Name Formula
atom loses one or more electrons, it is no longer Lithium ion Li+
electrically neutral and becomes a positively charged Sodium ion Na+
ion. This is called a cation. Potassium ion K+
Magnesium ion Mg2+
For example, when potassium is involved in a Calcium ion Ca2+
chemical reaction, each atom loses an electron to Aluminium ion Al3+
form a positive ion, K+. Zinc ion Zn2+
Ammonium ion NH+4
19K+ 19 protons = 19+ Fluoride ion F−
18 electrons = 18− Chloride ion Cl−
Overall charge = 1+ Bromide ion Br −
Hydroxide ion OH−
When an atom gains one or more electrons, it Oxide ion O2−
becomes a negatively charged ion. This is called Sulfide ion S2−
an anion. For example, in some of the chemical Carbonate ion CO23−
reactions involving oxygen, each oxygen atom gains Nitrate ion NO−3
two electrons to form a negative ion, O2−. Sulfate ion SO24−
8O2− 8 protons = 8+
10 electrons = 10−
Overall charge = 2−
The process of gaining or losing electrons is known Test yourself
as ionisation.
14 a State three differences between an electron
Table 2.10 shows some common ions. You will and a proton.
notice that:
b State two similarities between a neutron and
» some ions contain more than one type of atom, a proton.
for ieoxnammpaylepNoOss3−ess
» an more than one unit of charge c Why are atoms electrically neutral?
(either negative or positive), for example, Al3+, 15 Copy and complete the following table by writing
O2− and SO42−.
the symbol for the element shown from the given
Worked example number of particles present. The first one is done
for you.
1 Show how the sodium ions shown in Table 2.10 are Element Symbol Particles present
formed. Nitrogen 147N 7p, 7n, 7e
11 protons = 11+ Aluminium 13p, 14n, 13e
Potassium 19p, 20n, 19e
1 1Na+ Ov1e0r electrons = 10− Argon 18p, 22n, 18e
all charge = 1+
2 Show how the sulfide ions shown in Table 2.10 are 16 Copy and complete the following table giving the
formed. charge on the ions shown.
16 protons = 16+
16S2− Ov1e8ra elllecchtarorgnes = 18− Proton Number of Charge on
= 2− number electrons ion
11 10
Element
Sodium 9 10
Fluorine
Magnesium 12 10
23
2 Atoms, elements and compounds
Isotopes large amounts of energy, but also certain types
of dangerous radiations that can in some cases be
In some elements, not all of the atoms in a sample useful to society. For example, uranium-235 is used
of the element are identical. Some atoms of the as a source of power in nuclear reactors in nuclear
same element can contain different numbers of power stations and cobalt-60 is used in radiotherapy
neutrons and so have different mass numbers. Atoms treatment in hospitals (Figure 2.16); these are both
of the same element which have different neutron radioisotopes.
numbers are called isotopes. The two isotopes of
chlorine are shown in Figure 2.15. ▼ Table 2.11 Some atoms and their isotopes
Element Symbol Particles present
Hydrogen 1e, 1p, 0n
(Deuterium) 11H 1e, 1p, 1n
(Tritium) 1e, 1p, 2n
Carbon 21H 6e, 6p, 6n
6e, 6p, 7n
Oxygen 31H 6e, 6p, 8n
8e, 8p, 8n
Strontium 162C 8e, 8p, 9n
8e, 8p, 10n
Uranium 163C 38e, 38p, 48n
38e, 38p, 50n
164C 38e, 38p, 52n
92e, 92p, 143n
35 37 186O 92e, 92p, 146n
17 17
Cl Cl 17 O
8
17 Protons 17 Protons 188O
86 Sr
38
18 Neutrons 20 Neutrons 88 Sr
38
▲ Figure 2.15 The nuclei of two isotopes of chlorine 90 Sr
38
U235
92
Key definition U238
92
Isotopes are atoms of the same element that have
the same number of protons but different numbers of
neutrons.
Isotopes of the same element have the same
chemical properties because they have the same
number of electrons and therefore the same
electronic configuration (p. 26).
The only effect of the extra neutron is to alter the ▲ Figure 2.16 Cobalt-60 is used in radiotherapy treatment.
mass of the atom and any properties that depend A beam of gamma rays produced by the radioactive isotope
on it, such as density. Other examples of atoms and is directed into the patient’s body to kill tumour tissue.
their isotopes are shown in Table 2.11.
There are two types of isotopes: those which
are stable and those which are unstable. The
isotopes which are unstable, as a result of the extra
neutrons in their nuclei, are radioactive and are
called radioisotopes. The nuclei of these atoms
break up spontaneously with the release of not only
24
2.4 Inside atoms
Going further
The mass spectrometer The spectrometer is set up to ensure that when the
ions leave the acceleration area they all have the same
How do we know isotopes exist? They were first kinetic energy, regardless of the mass of the ions. This
discovered by scientists using apparatus called a mass means that the lighter ions travel faster than the heavier
spectrometer (Figure 2.17). The first mass spectrometer ones, and effectively separate the ions according to their
was built by the British scientist Francis Aston in 1919 mass. Having left the acceleration area, the time for
and enabled scientists to compare the relative masses the ions to reach the detector is recorded. The detector
of atoms accurately for the first time. counts the number of each of the ions which fall upon
it and so a measure of the percentage abundance of
Vacuum chamber each isotope is obtained. A typical mass spectrum for
chlorine is shown in Figure 2.18.
Sample Ion
inlet detector 100
Ionisation Acceleration Drift region Abundance/% 75 35Cl (75%)
area area 50
Flight path
time measurement 25 37Cl (25%)
▲ Figure 2.17 A diagram of a mass spectrometer 0 10 20 30 40 50
A vacuum exists inside the mass spectrometer. A Mass
sample of the vapour of the element is injected into the
ionisation chamber where it is bombarded by electrons. ▲ Figure 2.18 The mass spectrum for chlorine
The collisions which take place between these electrons
and the injected atoms cause an electron to be lost from
the atom, which becomes a positive ion with a +1 charge.
These positive ions are then accelerated towards a
negatively charged plate, in the acceleration area.
Relative atomic mass Worked example
The average mass of a large number of atoms of an
element is called its relative atomic pmeracsesn(tAagr)e.
This quantity takes into account the What is the relative atomic mass of chlorine? Chlorine has
two isotopes:
abundance of all the isotopes of an element which 1375CI 37 CI
exist. 75 17
% abundance 25
Key definition Hence the ‘average mass’ or Ar of a chlorine atom is:
Relative atomic mass, Acor,misptahreedavteor1a/g1e2tmh aosf sthoef the (75 × 35) + (25 × 37) = 35.5
isotopes of an element mass 100
of an atom of 12C. Ar = 35.5
1
In 1961 the International Union of Pure and = 35.5 amu
Applied Chemistry (IUPAC) recommended that the
asttaonmdaorfdcuasrbeodnf-o1r2twheasAtr askceanletwoahsacvaerbaomna-1s2s.oAfn
1o2f tahmeun.aTthueraAllryoofcacnurerlienmg eanttomisstohfe average mass Test yourself
an element on
17 Calculate the number of neutrons in the following
atoms:
a scale where 12C has a mass of exactly 12 units. a 1237AI b c 4901Zr d 19706Os.
P31
average mass of isotopes of the element
15
Ar = 1 × mass of one atom of carbon-12 18 Given that the percentage abundance of 1200Ne is
12 90% and that of 1202Ne is 10%, calculate the Ar of neon.
Note: 1 of the mass of one carbon-12 atom = 1 amu.
12
25
2 Atoms, elements and compounds
The arrangement of electrons in 8p
atoms 8n
The nucleus of an atom contains the heavier sub- ▲ Figure 2.19 Arrangement of electrons in an oxygen atom
atomic particles – the protons and the neutrons. The
electrons, the lightest of the sub-atomic particles, H
move around the nucleus at great distances from the
nucleus relative to their size. They move very fast in 1p
electron shells, very much like the planets orbit the
Sun. Hydrogen
Li
It is not possible to give the exact position
of an electron in an electron shell. However, we 3p
can state that electrons can only occupy certain, 4n
definite electron shells and that they cannot exist
between them. Also, as mentioned earlier, each of Lithium
the electron shells can hold only a certain number
of electrons. Na
11p
» First shell holds up to two electrons. 12n
» Second shell holds up to eight electrons. Sodium
» Third shell holds up to 18 electrons.
Ar
There are further electron shells which contain 18p
increasing numbers of electrons. 22n
The third electron shell can be occupied by a Argon
maximum of 18 electrons. However, when eight
electrons have occupied this shell, a certain K
stability is given to the atom and the next two 19p
electrons go into the fourth electron shell, and then 20n
the remaining ten electrons complete the third shell.
Potassium
The electrons fill the electron shell starting from
the shell nearest to the nucleus, which has the ▲ Figure 2.20 Electronic configurations of hydrogen,
lowest energy. When this is full (with two electrons), lithium, sodium, argon and potassium
the next electron goes into the second electron
shell. When this shell is full with eight electrons,
then the electrons begin to fill the third and fourth
shells as stated above.
For example, a 168O atom has a proton number of 8
and therefore has eight electrons. Two of the eight
electrons enter the first electron shell, leaving six
to occupy the second electron shell, as shown in
Figure 2.19. The electronic configuration for oxygen
can be written in a shorthand way as 2,6.
There are 118 elements, and Table 2.12 shows
the way in which the electrons are arranged in the
first 20 of these elements. The way in which the
electrons are distributed is called the electronic
configuration (electronic structure). Figure 2.20
shows the electronic configuration of a selection
of atoms.
26
2.4 Inside atoms
▼ Table 2.12 Electronic configuration of the first 20 elements
Element Symbol Proton number Number of electrons Electronic configuration
Hydrogen H 1 1 1
Helium He 2 2 2
Lithium Li 3 3 2,1
Beryllium Be 4 4 2,2
Boron B 5 5 2,3
Carbon C 6 6 2,4
Nitrogen N 7 7 2,5
Oxygen O 8 8 2,6
Fluorine F 9 9 2,7
Neon Ne 10 10 2,8
Sodium Na 11 11 2,8,1
Magnesium Mg 12 12 2,8,2
Aluminium Al 13 13 2,8,3
Silicon Si 14 14 2,8,4
Phosphorus P 15 15 2,8,5
Sulfur S 16 16 2,8,6
Chlorine Cl 17 17 2,8,7
Argon Ar 18 18 2,8,8
Potassium K 19 19 2,8,8,1
Calcium Ca 20 20 2,8,8,2
From Table 2.12 you can see helium has a full outer chemical bonding, which we will discuss further in
shell of two electrons, and neon has a full outer the next chapter.
shell of 8 electrons. These two elements are very
unreactive for this reason. In Chapter 9 we will see Test yourself
that the number of outer shell electrons is related to
the position in the Periodic Table. Helium and neon 19 How many electrons may be accommodated in
are part of a group of elements known as the noble each of the first three electron shells?
or inert gases and they are generally very stable and
unreactive (p. 145). This is linked to their full outer 20 What is the same about the electronic
shells. When elements react to form compounds, configurations of:
they do so to achieve full electron shells, and this a lithium, sodium and potassium?
idea forms the basis of the electronic theory of b beryllium, magnesium and calcium?
21 An element X has a proton number of 13. What is
the electronic configuration of X?
27
2 Atoms, elements and compounds
Revision checklist ✔ Define an oxidising agent and a reducing agent.
After studying Chapter 2 you should be able to: ✔ Describe the structure of the atom as a central
✔ Describe the differences between elements, nucleus containing neutrons and protons
surrounded by electrons in shells.
compounds and mixtures.
✔ Interpret and use symbols for given atoms. ✔ State the relative charges and relative masses of a
✔ State the formulae of the elements and compounds proton, a neutron.
you have dealt with. ✔ Define proton number and atomic number as well
✔ Define the molecular formula of a compound as as mass number.
the number and type of different atoms in one ✔ Determine the electronic configuration of
molecule. elements with the proton number 1 to 20.
✔ Deduce the formula of a simple compound from
the relative numbers of atoms present in a ✔ Describe the formation of positive ions, known as
model or a diagrammatic representation of the cations, and negative ions, known as anions.
compound.
✔ Construct word equation and symbol equations to ✔ State what isotopes are.
show how reactants form products, including state
symbols. ✔ State that isotopes of the same element have
✔ Define redox reactions as involving both oxidation the same electronic configuration and so have
and reduction. the same chemical properties.
✔ Define oxidation as oxygen gain and reduction as
oxygen loss. ✔ Calculate the relative atomic mass (Ar) of an
✔ Identify redox reactions as reactions involving gain element from given data of the relative masses
and loss of oxygen. and abundance of their different isotopes.
28
Exam-style questions
Exam-style questions
1 a Define the terms: 5 State, at room temperature and pressure (r.t.p.),
i proton [3] which of the substances listed below is/are:
ii neutron [2] a a solid element
iii electron. [3] b a liquid element
b An atom X has a proton number of 19 and c a gaseous mixture
relative atomic mass of 39. d a solid mixture
i How many electrons, protons and neutrons e a liquid compound
are there in atom X? [3] f a solid compound.
ii How many electrons will there be in the bromine, carbon dioxide, helium,
outer electron shell of atom X? [1] steel, air, oil, marble, copper, water,
iii What is the electronic configuration of sand, tin, bronze, mercury, salt [11]
atom X? [1] 6 a How many atoms of the different elements are
there in the formulae of these compounds?
2 a 6391Ga and 7311Ga are isotopes of gallium. b iBiivviv iiv iii ai l nmcspeaauomthinethgpecrmtaapenihcnroeoa,tonrlanhC,iicnece1uCi2,imdHta6fCHr,oc2aH2ls5iHOltdOu4oeN1 ,Hlw1,fO Ca Ci3Htn ue3g(C,NOe(ONqO3uHH)2a4 )t2iSoOn4s . [[[[[[[4233333]]]]]]]
Use this example to explain what you
understand by the term isotope. [3]
b A sample 73o11fGagaaltlioumms.isCa6l0c%ula63t91Geathaetormelsative
and 40%
atomic mass of this sample of gallium. [2]
3 Define the following terms using specific
examples to help with your explanation:
a element [2] iiviiii v i FZMHLein2g((((s(sgs)s))))+++++OCOOC2l2(2O2((g(g2gg)()))g→)→→→→LZHFi2neM2OOOC(g((lsO3sg)() )(s s )) + C(s) [[[[[22223]]]]]
b metal [3]
c non-metal [2]
d compound [2]
e molecule [2]
f mixture. [2]
4 State which of the substances listed below are:
a metallic elements
b non-metallic elements
c compounds
d mixtures.
silicon, sea water, calcium, argon,
water, air, carbon monoxide, iron,
sodium chloride, diamond, brass,
copper, dilute sulfuric acid, sulfur,
oil, nitrogen, ammonia [17]
29
3 Bonding and structure
FOCUS POINTS
★ In what different ways do elements combine to form compounds?
★ How do the bonds within a compound affect its properties?
★ What is a redox reaction?
In the previous chapter, you established that
both metal and non-metal elements are
made up of atoms, which can combine to form
compounds. In this chapter you will look more
closely at how different elements combine, or
bond, to form compounds. You will see that,
generally, when elements combine they can
form either ionic bonds with one another and
are known as ionic compounds, or they can form
covalent bonds with one another and are known
as covalent compounds. You will also look at
a third type of bonding, found in metals only,
called metallic bonding.
You will find out how the type of bonding affects
the properties of these compounds. You will
also learn more about formulae and how they
can be obtained from given data and you will
obtain an understanding of redox reactions and
the processes associated with them.
3.1 Ionic bonding ▲ Figure 3.1 The properties of salt are very different from
those of the sodium and chlorine it was made from. To
Ionic bonds are usually found in compounds get salt into your diet you would not eat sodium or inhale
that contain metals combined with non-metals. chlorine!
When this type of bond is formed, electrons are
transferred from the metal atoms to the non-metal Sodium has just one electron in its outer shell
atoms during the chemical reaction. In doing this, (so1hu1eNtelalr(2e1,7l8eC,cl1t2)r.,o8Cn,h7ol)o.frWeinahecehnhastoshdseiesuvemetnwateoolemecletirmsotenrnsatnisnsrfieetarsrceotd,utttehore
the atoms become more stable due to their full outer the outer shell of a chlorine atom (Figure 3.2).
shells. For example, consider what happens when
sodium and chlorine react together and combine to
make sodium chloride (Figure 3.1).
sodium + chlorine → sodium chloride
30
3.1 Ionic bonding
In this way both the atoms obtain full outer the electrons from one atom are shown as crosses
shells and become ‘like’ the noble gas nearest to and those of the other atom as dots. In the reaction
them in the Periodic Table (see Figure 9.3, p. 135). the sodium atom has become a sodium ion, in a
One way of showing this transfer of electrons is by process known as ionisation. This sodium ion has an
using a dot-and-cross diagram. In these diagrams, electronic configuration like the noble gas neon.
Sodium atom Chlorine atom Sodium ion Chloride ion
(Na) (Cl) (Na ) (Cl )
11 protons 11 17 protons 17
10 electrons 10 18 electrons 18
Overall charge 1 Overall charge 1
▲ Figure 3.2 Ionic bonding in sodium chloride Only the outer electrons are important in bonding,
so we can simplify the diagrams by missing out the
To lose electrons in this way is called oxidation. inner shells (Figure 3.3).
The chlorine atom has become a chloride ion with
an electronic configuration like argon. To gain The charges on the sodium and chloride ions are
electrons in this way is called reduction. equal but opposite. They balance each other and
the resulting formula for sodium chloride is NaCl.
In the chemical process producing sodium These oppositely charged ions attract each other
chloride, both oxidation and reduction have taken and are pulled, or bonded, to one another by strong
place and so this is known as a redox reaction. electrostatic forces. This type of bonding is called
ionic bonding.
Key definitions
Oxidation involves loss of electrons. Key definition
Reduction involves gain of electrons. An ionic bond is a strong electrostatic attraction
Redox reactions involve simultaneous oxidation between oppositely charged ions.
and reduction.
All students should be able to describe the
A further discussion of oxidation and reduction formation of ionic bonds between elements of Group
in terms of electron transfer takes place in I and Group VII and draw dot-and-cross diagrams.
Chapter 5 (p. 71). Extended learners should be able to do this for
any metallic and non-metallic element, including
drawing dot-and-cross diagrams.
Na Cl [Na] Cl
▲ Figure 3.3 Simplified diagram of ionic bonding in sodium chloride Sodium chloride (NaCl)
31
3 Bonding and structure [Mg] 2 2
O
Mg O
Magnesium oxide (MgO)
▲ Figure 3.4 Simplified diagram of ionic bonding in magnesium oxide
Figure 3.4 shows the electron transfers that Magnesium obtains a full outer shell by losing
take place between a magnesium atom and an two electrons. These are transferred to the oxygen
oxygen atom during the formation of magnesium atom. In magnesium oxide, the Mg2+ and O2− are
oxide. oppositely charged and are attracted to one
another. The formula for magnesium oxide is MgO.
Cl Cl
Ca
[Ca]2
Cl Cl
Calcium chloride (CaCl2)
▲ Figure 3.5 The transfer of electrons that takes place during the formation of calcium chloride
When calcium and chlorine react, the calcium formed containing two chloride ions (Cl−) for each
atom gives each of the two chlorine atoms one calcium ion (Ca2+). The chemical formula is CaCl2.
electron (Figure 3.5). In this case, a compound is
Test yourself b lithium chloride (LiCl)
1 Define the term ‘ionic bond’. c magnesium fluoride (MgF2)
2 Why do inert (noble) gases not bond with themselves d calcium oxide (CaO).
and other elements?
3 Draw diagrams to represent the bonding in each
of the following ionic compounds:
a potassium fluoride (KF)
32
3.1 Ionic bonding
Going further
Scientists, using X-ray diffraction (Figure 3.6a), have is possible to use them to investigate extremely small
obtained photographs that indicate the way in which the structures.
ions are arranged (Figure 3.6b). The electron density When X-rays are passed through a crystal of sodium
map of sodium chloride is shown in Figure 3.6c. chloride, for example, you get a pattern of spots called
Figure 3.6d shows the structure of sodium chloride as a diffraction pattern (Figure 3.6b). This pattern can
determined by the X-ray diffraction technique. The study be recorded on photographic film and used to work
of crystals using X-ray diffraction was pioneered by Sir out how the ions or atoms are arranged in the crystal.
William Bragg and his son Sir Lawrence Bragg in 1912. Crystals give particular diffraction patterns depending
X-rays are a form of electromagnetic radiation. They on their structure, and this makes X-ray diffraction a
have a much shorter wavelength than light therefore it particularly powerful technique in the investigation of
crystal structures.
Crystal
X-rays
a X-ray diffraction technique Diffraction pattern
Na on a photographic film
Cl
b X-ray diffraction photograph of sodium chloride
100 200 500 Cl
1000 Na
2000 5000 100 Na surrounded Cl surrounded
300 by 6 Cl ions by 6 N a ions
500
1000 d The structure of sodium chloride
2000
5000
c Electron density map of sodium chloride
▲ Figure 3.6
33
3 Bonding and structure
Ionic structures Properties of ionic compounds
Ionic structures are solids at room temperature and Ionic compounds have the following properties.
have high melting and boiling points. » They are usually solids at room temperature,
The ions are packed together in a regular with high melting and boiling points.
arrangement called a lattice. Within the lattice, This is due to the strong electrostatic forces
oppositely charged ions attract one another strongly.
holding the crystal lattice together. A lot of
Figure 3.6d shows only a tiny part of a small energy is therefore needed to separate the
crystal of sodium chloride. Many millions of sodium ions and melt the substance.
ions and chloride ions would be arranged in this » They are usually hard substances.
way in a crystal of sodium chloride to make up the » They mainly dissolve in water.
giant ionic lattice structure. Each sodium ion in This is because water molecules are able to
the lattice is surrounded by six chloride ions, and bond with both the positive and the negative
each chloride ion is surrounded by six sodium ions. ions, which breaks up the lattice and keeps the
ions apart. Figure 3.8 shows the interaction
Going further between water molecules (the solvent) and
sodium and chloride ions from sodium chloride
Not all ionic substances form the same structures. (the solute). For a further discussion of the
Caesium chloride (CsCl), for example, forms a solubility of substances see, Chapter 8, p. 120.
different structure due to the larger size of the » They usually conduct electricity when in the
caesium ion compared with that of the sodium ion. This molten state or in aqueous solution.
gives rise to the structure shown in Figure 3.7, which is The forces of attraction between the ions are
called a body-centred cubic structure. Each caesium weakened and the ions are free to move to the
ion is surrounded by eight chloride ions and, in turn, appropriate electrode. This allows an electric
each chloride ion is surrounded by eight caesium ions. current to be passed through the molten
compound (see Chapter 5, p. 71).
Cs CI » They usually cannot conduct electricity when solid,
a The arrangement of ions because the ions are not free to move.
Na+
Cl− O H
Na+ H
Cl−
Cs surrounded by 8 Cl− ions ▲ Figure 3.8 Salt (an ionic compound) dissolving in water
In the same way, each Cl−
is surrounded by 8 Cs ions
b The unit cell
▲ Figure 3.7 The structure of caesium chloride
34
3.1 Ionic bonding
Test yourself
4 Link the terms in the boxes on the left with the definitions on the right.
Oxidation An atom or group of atoms which has lost one or more electrons
Reduction Involves loss of electrons
Redox A strong electrostatic force of attraction between oppositely charged ions
Negative ion Involves gain of electrons
Positive ion Involves both oxidation and reduction
Ionic bond An atom or group of atoms which has gained one or more electrons
Formulae of ionic substances can form the Cu+ ion and the Cu2+ ion, with oxidation
numbers of 1 and 2 respectively. Therefore it can
We saw previously (see p. 31) that ionic compounds form two different compounds with chlorine, CuCl
contain positive and negative ions, whose charges
balance. For example, sodium chloride contains one abnydusCiunCgl2R. oWmeacnannuamlseoradlisstinintghueisirhntahmeedsi:ffCeurCelnicse
Na+ ion for every Cl− ion, giving rise to the formula
NaCl. This method can be used to write down cSoimpiplaerrl(yI,)icrohnlofroidrme santhdeCFueC2l+2 is copper(II) chloride.
formulae which show the ratio of the number of ions and Fe3+ ions and so
can also form two different compounds with, for
present in any ionic compound. example, cchhlloorriindee:).FeCl2 (iron(II) chloride) and FeCl3
The formula of amtabgyneeasicuhmMcgh2l+oiroidneciosmMbginCiln2.gTwhiisth (iron(III)
formula is arrived Table 3.1 shows the oxidation numbers of a
two Cl− ions, and once again the charges balance. The
use of oxidation numbers is useful when working out series of ions you will normally meet in your study
the formula of a compound. An oxidation number shows of chemistry.
You will notice that Table 3.1 includes groups
how oxidised or reduced an ion is compared to its atom. of atoms which have net charges. For example,
Na+ has an oxidation number of +1 because it is formed the nitrate ion is a single unit composed of one
by the loss of one electron (Chapter 5, p. 71) from a nitrogen atom and three oxygen atoms, and has
sodium atom but Mg2+ has an oxidation number of +2
because it is formed when a magnesium atom loses one single negative charge. The formula, therefore,
taonwfo2otmoicnauegittnstrihaedatseeituttimhohenensbNitprOrara3cetksheeeantws.tobTfuheoleidrsnebivnpeedlraMiyccgaem(tdNeaOisgn3nt)hbe2.rasaitYucotmkuheetiwroseinwll.airteh
two electrons (Table 3.1). Na+ can bond (combine) with The ratio of the atoms present is therefore:
only one Cl− ion, whereas Mg2+ can bond with two
Cl− ions. A chloride ion has an oxidation state of –1
because it is formed when a chlorine atom gains one
extra electron.
Some elements, such as copper and iron, possess Mg (N O3)2
two ions with different oxidation numbers. Copper 1Mg : 2N : 6O
▼ Table 3.1 Oxidation numbers (valencies) of some elements (ions) and groups of atoms
Oxidation numbers (valencies)
1 2 3
Metals Lithium (Li+) Magnesium (Mg2+) Aluminium (Al3+)
Sodium (Na+) Calcium (Ca2+) Iron(III) (Fe3+)
Potassium (K+) Copper(II) (Cu2+)
Silver (Ag+) Zinc (Zn2+)
Copper(I) (Cu+) Iron(II) (Fe2+)
Lead (Pb2+)
Barium (Ba2+)
35
3 Bonding and structure
Oxidation numbers (valencies)
1 2 3
(O2−)
Non-metals Fluoride (F−) Oxide (S2−)
Groups of atoms Sulfide
Chloride (Cl−)
Bromide (Br−)
Hydrogen (H+)
Hydroxide (OH−) Carbonate (CO32−) Phosphate (PO43−)
Sulfate (SO42−)
Nitrate (NO3−) Dichromate(VI) (Cr2O72−)
(NH4+)
Ammonium (HCO3−)
(MnO4−)
Hydrogencarbonate
Manganate (VII)
Oxidation number ▲ Figure 3.9 Iron(II) sulfate is pale green, while iron(III)
sulfate is yellow
Each atom in an element or compound is assigned
an oxidation number to show how much it is » An increase in the oxidation number, for
reduced or oxidised. The following points should be example, from +2 to +3 as in the case of Fe2+ to
remembered when using oxidation numbers. Fe3+, is oxidation.
» Roman numerals (I, II, III, IV, V, VI, VII, VIII) » However, a reduction in the oxidation number,
are used in writing the oxidation number of an for example, from +6 to +3 as in the case of Cr6+
element. (in CrO42−) to Cr3+, is reduction.
» This number is placed after the element that Key definitions
iirtorne(fIeIrIs)tcoh.lFoorirdeexaanmdplneo,tthireonn(a3m) echfloorriFdeeC.l3 is Oxidation involves an increase in oxidation number.
» The oxidation number of the free element is Reduction involves a decrease in oxidation number.
always 0, for example, in metals such as zinc
and copper. During a redox reaction, the substance that brings
» In simple monatomic ions, the oxidation number about oxidation is called an oxidising agent and is
is the same as the charge on the ion. So iodine itself reduced during the process. A substance that
has an oxidation number of 0 in I2 but an brings about reduction is a reducing agent and is
oxidation number of −1 in I−. itself oxidised during the process.
» Compounds have no charge overall. Hence
the oxidation numbers of all the individual
elements in a compound must add up to 0. The
oxidation numbers of elements in compounds
can vary, as seen in Table 3.1. It is possible
to recognise which of the different oxidation
numbers a metal element is in by the colour of
its compounds (Figure 3.9).
36
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