The Free High School Science Texts: Textbooks for High School Students Chemistry

The Free High School Science Texts: Textbooks for High School Students Chemistry
FHSST Authors
The Free High School Science Texts:
Textbooks for High School Students
Studying the Sciences
Chemistry
Grades 10 - 12
Version 0
November 9, 2008
ii
Copyright 2007 “Free High School Science Texts”
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FHSST Core Team
Mark Horner ; Samuel Halliday ; Sarah Blyth ; Rory Adams ; Spencer Wheaton
FHSST Editors
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Whitfield
FHSST Contributors
Rory Adams ; Prashant Arora ; Richard Baxter ; Dr. Sarah Blyth ; Sebastian Bodenstein ;
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Daniels ; Sean Dobbs ; Fernando Durrell ; Dr. Dan Dwyer ; Frans van Eeden ; Giovanni
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Andrew Kubik ; Dr. Marco van Leeuwen ; Dr. Anton Machacek ; Dr. Komal Maheshwari ;
Kosma von Maltitz ; Nicole Masureik ; John Mathew ; JoEllen McBride ; Nikolai Meures ;
Riana Meyer ; Jenny Miller ; Abdul Mirza ; Asogan Moodaly ; Jothi Moodley ; Nolene Naidu ;
Tyrone Negus ; Thomas O’Donnell ; Dr. Markus Oldenburg ; Dr. Jaynie Padayachee ;
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iii
iv
Contents
I
II
Introduction
1
Matter and Materials
3
1 Classification of Matter - Grade 10
1.1
1.2
5
Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5
1.1.1
Heterogeneous mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . .
6
1.1.2
Homogeneous mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . .
6
1.1.3
Separating mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7
Pure Substances: Elements and Compounds . . . . . . . . . . . . . . . . . . . .
9
1.2.1
Elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
9
1.2.2
Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
9
1.3
Giving names and formulae to substances . . . . . . . . . . . . . . . . . . . . . 10
1.4
Metals, Semi-metals and Non-metals . . . . . . . . . . . . . . . . . . . . . . . . 13
1.4.1
Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13
1.4.2
Non-metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14
1.4.3
Semi-metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14
1.5
Electrical conductors, semi-conductors and insulators . . . . . . . . . . . . . . . 14
1.6
Thermal Conductors and Insulators . . . . . . . . . . . . . . . . . . . . . . . . . 15
1.7
Magnetic and Non-magnetic Materials . . . . . . . . . . . . . . . . . . . . . . . 17
1.8
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18
2 What are the objects around us made of? - Grade 10
21
2.1
Introduction: The atom as the building block of matter . . . . . . . . . . . . . . 21
2.2
Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21
2.2.1
Representing molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . 21
2.3
Intramolecular and intermolecular forces . . . . . . . . . . . . . . . . . . . . . . 25
2.4
The Kinetic Theory of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26
2.5
The Properties of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28
2.6
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31
3 The Atom - Grade 10
3.1
35
Models of the Atom . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35
3.1.1
The Plum Pudding Model . . . . . . . . . . . . . . . . . . . . . . . . . . 35
3.1.2
Rutherford’s model of the atom
v
. . . . . . . . . . . . . . . . . . . . . . 36
CONTENTS
3.1.3
3.2
3.3
CONTENTS
The Bohr Model . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37
How big is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.2.1
How heavy is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.2.2
How big is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
Atomic structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.3.1
The Electron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39
3.3.2
The Nucleus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39
3.4
Atomic number and atomic mass number . . . . . . . . . . . . . . . . . . . . . 40
3.5
Isotopes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42
3.6
3.7
3.8
3.9
3.5.1
What is an isotope? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42
3.5.2
Relative atomic mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . 45
Energy quantisation and electron configuration . . . . . . . . . . . . . . . . . . 46
3.6.1
The energy of electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . 46
3.6.2
Energy quantisation and line emission spectra . . . . . . . . . . . . . . . 47
3.6.3
Electron configuration . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47
3.6.4
Core and valence electrons . . . . . . . . . . . . . . . . . . . . . . . . . 51
3.6.5
The importance of understanding electron configuration . . . . . . . . . 51
Ionisation Energy and the Periodic Table . . . . . . . . . . . . . . . . . . . . . . 53
3.7.1
Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 53
3.7.2
Ionisation Energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 55
The Arrangement of Atoms in the Periodic Table . . . . . . . . . . . . . . . . . 56
3.8.1
Groups in the periodic table
. . . . . . . . . . . . . . . . . . . . . . . . 56
3.8.2
Periods in the periodic table . . . . . . . . . . . . . . . . . . . . . . . . 58
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59
4 Atomic Combinations - Grade 11
63
4.1
Why do atoms bond? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63
4.2
Energy and bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63
4.3
What happens when atoms bond? . . . . . . . . . . . . . . . . . . . . . . . . . 65
4.4
Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65
4.4.1
The nature of the covalent bond . . . . . . . . . . . . . . . . . . . . . . 65
4.5
Lewis notation and molecular structure . . . . . . . . . . . . . . . . . . . . . . . 69
4.6
Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72
4.7
4.8
4.6.1
Non-polar and polar covalent bonds . . . . . . . . . . . . . . . . . . . . 73
4.6.2
Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73
Ionic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74
4.7.1
The nature of the ionic bond . . . . . . . . . . . . . . . . . . . . . . . . 74
4.7.2
The crystal lattice structure of ionic compounds . . . . . . . . . . . . . . 76
4.7.3
Properties of Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . 76
Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 76
4.8.1
The nature of the metallic bond . . . . . . . . . . . . . . . . . . . . . . 76
4.8.2
The properties of metals . . . . . . . . . . . . . . . . . . . . . . . . . . 77
vi
CONTENTS
4.9
CONTENTS
Writing chemical formulae
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 78
4.9.1
The formulae of covalent compounds . . . . . . . . . . . . . . . . . . . . 78
4.9.2
The formulae of ionic compounds . . . . . . . . . . . . . . . . . . . . . 80
4.10 The Shape of Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 82
4.10.1 Valence Shell Electron Pair Repulsion (VSEPR) theory . . . . . . . . . . 82
4.10.2 Determining the shape of a molecule . . . . . . . . . . . . . . . . . . . . 82
4.11 Oxidation numbers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85
4.12 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 88
5 Intermolecular Forces - Grade 11
91
5.1
Types of Intermolecular Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . 91
5.2
Understanding intermolecular forces . . . . . . . . . . . . . . . . . . . . . . . . 94
5.3
Intermolecular forces in liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . 96
5.4
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97
6 Solutions and solubility - Grade 11
101
6.1
Types of solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 101
6.2
Forces and solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102
6.3
Solubility . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103
6.4
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106
7 Atomic Nuclei - Grade 11
107
7.1
Nuclear structure and stability . . . . . . . . . . . . . . . . . . . . . . . . . . . 107
7.2
The Discovery of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 107
7.3
Radioactivity and Types of Radiation . . . . . . . . . . . . . . . . . . . . . . . . 108
7.4
7.3.1
Alpha (α) particles and alpha decay . . . . . . . . . . . . . . . . . . . . 109
7.3.2
Beta (β) particles and beta decay . . . . . . . . . . . . . . . . . . . . . 109
7.3.3
Gamma (γ) rays and gamma decay . . . . . . . . . . . . . . . . . . . . . 110
Sources of radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112
7.4.1
Natural background radiation . . . . . . . . . . . . . . . . . . . . . . . . 112
7.4.2
Man-made sources of radiation . . . . . . . . . . . . . . . . . . . . . . . 113
7.5
The ’half-life’ of an element . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113
7.6
The Dangers of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116
7.7
The Uses of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117
7.8
Nuclear Fission
7.9
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118
7.8.1
The Atomic bomb - an abuse of nuclear fission . . . . . . . . . . . . . . 119
7.8.2
Nuclear power - harnessing energy . . . . . . . . . . . . . . . . . . . . . 120
Nuclear Fusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120
7.10 Nucleosynthesis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121
7.10.1 Age of Nucleosynthesis (225 s - 103 s) . . . . . . . . . . . . . . . . . . . 121
7.10.2 Age of Ions (103 s - 1013 s) . . . . . . . . . . . . . . . . . . . . . . . . . 122
7.10.3 Age of Atoms (1013 s - 1015 s) . . . . . . . . . . . . . . . . . . . . . . . 122
7.10.4 Age of Stars and Galaxies (the universe today) . . . . . . . . . . . . . . 122
7.11 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122
vii
CONTENTS
CONTENTS
8 Thermal Properties and Ideal Gases - Grade 11
125
8.1
A review of the kinetic theory of matter . . . . . . . . . . . . . . . . . . . . . . 125
8.2
Boyle’s Law: Pressure and volume of an enclosed gas . . . . . . . . . . . . . . . 126
8.3
Charles’s Law: Volume and Temperature of an enclosed gas . . . . . . . . . . . 132
8.4
The relationship between temperature and pressure . . . . . . . . . . . . . . . . 136
8.5
The general gas equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 137
8.6
The ideal gas equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 140
8.7
Molar volume of gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 145
8.8
Ideal gases and non-ideal gas behaviour . . . . . . . . . . . . . . . . . . . . . . 146
8.9
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 147
9 Organic Molecules - Grade 12
151
9.1
What is organic chemistry? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151
9.2
Sources of carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151
9.3
Unique properties of carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4
Representing organic compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4.1
Molecular formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4.2
Structural formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 153
9.4.3
Condensed structural formula . . . . . . . . . . . . . . . . . . . . . . . . 153
9.5
Isomerism in organic compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 154
9.6
Functional groups . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 155
9.7
The Hydrocarbons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 155
9.7.1
The Alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 158
9.7.2
Naming the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 159
9.7.3
Properties of the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . 163
9.7.4
Reactions of the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . 163
9.7.5
The alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166
9.7.6
Naming the alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166
9.7.7
The properties of the alkenes . . . . . . . . . . . . . . . . . . . . . . . . 169
9.7.8
Reactions of the alkenes
9.7.9
The Alkynes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171
. . . . . . . . . . . . . . . . . . . . . . . . . . 169
9.7.10 Naming the alkynes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171
9.8
9.9
The Alcohols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172
9.8.1
Naming the alcohols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 173
9.8.2
Physical and chemical properties of the alcohols . . . . . . . . . . . . . . 175
Carboxylic Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 176
9.9.1
Physical Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177
9.9.2
Derivatives of carboxylic acids: The esters . . . . . . . . . . . . . . . . . 178
9.10 The Amino Group . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178
9.11 The Carbonyl Group . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178
9.12 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 179
viii
CONTENTS
CONTENTS
10 Organic Macromolecules - Grade 12
185
10.1 Polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 185
10.2 How do polymers form? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 186
10.2.1 Addition polymerisation . . . . . . . . . . . . . . . . . . . . . . . . . . . 186
10.2.2 Condensation polymerisation . . . . . . . . . . . . . . . . . . . . . . . . 188
10.3 The chemical properties of polymers . . . . . . . . . . . . . . . . . . . . . . . . 190
10.4 Types of polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191
10.5 Plastics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191
10.5.1 The uses of plastics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 192
10.5.2 Thermoplastics and thermosetting plastics . . . . . . . . . . . . . . . . . 194
10.5.3 Plastics and the environment . . . . . . . . . . . . . . . . . . . . . . . . 195
10.6 Biological Macromolecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 196
10.6.1 Carbohydrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 197
10.6.2 Proteins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 199
10.6.3 Nucleic Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 202
10.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 204
III
Chemical Change
209
11 Physical and Chemical Change - Grade 10
211
11.1 Physical changes in matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 211
11.2 Chemical Changes in Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212
11.2.1 Decomposition reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 213
11.2.2 Synthesis reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 214
11.3 Energy changes in chemical reactions . . . . . . . . . . . . . . . . . . . . . . . . 217
11.4 Conservation of atoms and mass in reactions . . . . . . . . . . . . . . . . . . . . 217
11.5 Law of constant composition . . . . . . . . . . . . . . . . . . . . . . . . . . . . 219
11.6 Volume relationships in gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . 219
11.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 220
12 Representing Chemical Change - Grade 10
223
12.1 Chemical symbols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 223
12.2 Writing chemical formulae
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224
12.3 Balancing chemical equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224
12.3.1 The law of conservation of mass . . . . . . . . . . . . . . . . . . . . . . 224
12.3.2 Steps to balance a chemical equation
. . . . . . . . . . . . . . . . . . . 226
12.4 State symbols and other information . . . . . . . . . . . . . . . . . . . . . . . . 230
12.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 232
13 Quantitative Aspects of Chemical Change - Grade 11
233
13.1 The Mole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 233
13.2 Molar Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 235
13.3 An equation to calculate moles and mass in chemical reactions . . . . . . . . . . 237
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CONTENTS
13.4 Molecules and compounds
CONTENTS
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 239
13.5 The Composition of Substances . . . . . . . . . . . . . . . . . . . . . . . . . . . 242
13.6 Molar Volumes of Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 246
13.7 Molar concentrations in liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . 247
13.8 Stoichiometric calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 249
13.9 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252
14 Energy Changes In Chemical Reactions - Grade 11
255
14.1 What causes the energy changes in chemical reactions? . . . . . . . . . . . . . . 255
14.2 Exothermic and endothermic reactions . . . . . . . . . . . . . . . . . . . . . . . 255
14.3 The heat of reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 257
14.4 Examples of endothermic and exothermic reactions . . . . . . . . . . . . . . . . 259
14.5 Spontaneous and non-spontaneous reactions . . . . . . . . . . . . . . . . . . . . 260
14.6 Activation energy and the activated complex . . . . . . . . . . . . . . . . . . . . 261
14.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 264
15 Types of Reactions - Grade 11
267
15.1 Acid-base reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.1 What are acids and bases? . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.2 Defining acids and bases . . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.3 Conjugate acid-base pairs . . . . . . . . . . . . . . . . . . . . . . . . . . 269
15.1.4 Acid-base reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 270
15.1.5 Acid-carbonate reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 274
15.2 Redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 276
15.2.1 Oxidation and reduction
. . . . . . . . . . . . . . . . . . . . . . . . . . 277
15.2.2 Redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 278
15.3 Addition, substitution and elimination reactions . . . . . . . . . . . . . . . . . . 280
15.3.1 Addition reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 280
15.3.2 Elimination reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 281
15.3.3 Substitution reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 282
15.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 283
16 Reaction Rates - Grade 12
287
16.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 287
16.2 Factors affecting reaction rates . . . . . . . . . . . . . . . . . . . . . . . . . . . 289
16.3 Reaction rates and collision theory . . . . . . . . . . . . . . . . . . . . . . . . . 293
16.4 Measuring Rates of Reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 295
16.5 Mechanism of reaction and catalysis . . . . . . . . . . . . . . . . . . . . . . . . 297
16.6 Chemical equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 300
16.6.1 Open and closed systems . . . . . . . . . . . . . . . . . . . . . . . . . . 302
16.6.2 Reversible reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 302
16.6.3 Chemical equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . 303
16.7 The equilibrium constant . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 304
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CONTENTS
16.7.1 Calculating the equilibrium constant . . . . . . . . . . . . . . . . . . . . 305
16.7.2 The meaning of kc values . . . . . . . . . . . . . . . . . . . . . . . . . . 306
16.8 Le Chatelier’s principle . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 310
16.8.1 The effect of concentration on equilibrium . . . . . . . . . . . . . . . . . 310
16.8.2 The effect of temperature on equilibrium . . . . . . . . . . . . . . . . . . 310
16.8.3 The effect of pressure on equilibrium . . . . . . . . . . . . . . . . . . . . 312
16.9 Industrial applications . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 315
16.10Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 316
17 Electrochemical Reactions - Grade 12
319
17.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 319
17.2 The Galvanic Cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 320
17.2.1 Half-cell reactions in the Zn-Cu cell . . . . . . . . . . . . . . . . . . . . 321
17.2.2 Components of the Zn-Cu cell . . . . . . . . . . . . . . . . . . . . . . . 322
17.2.3 The Galvanic cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 323
17.2.4 Uses and applications of the galvanic cell . . . . . . . . . . . . . . . . . 324
17.3 The Electrolytic cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 325
17.3.1 The electrolysis of copper sulphate . . . . . . . . . . . . . . . . . . . . . 326
17.3.2 The electrolysis of water . . . . . . . . . . . . . . . . . . . . . . . . . . 327
17.3.3 A comparison of galvanic and electrolytic cells . . . . . . . . . . . . . . . 328
17.4 Standard Electrode Potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . 328
17.4.1 The different reactivities of metals . . . . . . . . . . . . . . . . . . . . . 329
17.4.2 Equilibrium reactions in half cells . . . . . . . . . . . . . . . . . . . . . . 329
17.4.3 Measuring electrode potential . . . . . . . . . . . . . . . . . . . . . . . . 330
17.4.4 The standard hydrogen electrode . . . . . . . . . . . . . . . . . . . . . . 330
17.4.5 Standard electrode potentials . . . . . . . . . . . . . . . . . . . . . . . . 333
17.4.6 Combining half cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 337
17.4.7 Uses of standard electrode potential . . . . . . . . . . . . . . . . . . . . 338
17.5 Balancing redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342
17.6 Applications of electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . 347
17.6.1 Electroplating . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 347
17.6.2 The production of chlorine . . . . . . . . . . . . . . . . . . . . . . . . . 348
17.6.3 Extraction of aluminium
. . . . . . . . . . . . . . . . . . . . . . . . . . 349
17.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 349
IV
Chemical Systems
353
18 The Water Cycle - Grade 10
355
18.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355
18.2 The importance of water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355
18.3 The movement of water through the water cycle . . . . . . . . . . . . . . . . . . 356
18.4 The microscopic structure of water . . . . . . . . . . . . . . . . . . . . . . . . . 359
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18.4.1 The polar nature of water . . . . . . . . . . . . . . . . . . . . . . . . . . 359
18.4.2 Hydrogen bonding in water molecules . . . . . . . . . . . . . . . . . . . 359
18.5 The unique properties of water . . . . . . . . . . . . . . . . . . . . . . . . . . . 360
18.6 Water conservation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 363
18.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 366
19 Global Cycles: The Nitrogen Cycle - Grade 10
369
19.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 369
19.2 Nitrogen fixation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 369
19.3 Nitrification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 371
19.4 Denitrification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372
19.5 Human Influences on the Nitrogen Cycle . . . . . . . . . . . . . . . . . . . . . . 372
19.6 The industrial fixation of nitrogen . . . . . . . . . . . . . . . . . . . . . . . . . 373
19.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 374
20 The Hydrosphere - Grade 10
377
20.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 377
20.2 Interactions of the hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . 377
20.3 Exploring the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 378
20.4 The Importance of the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . 379
20.5 Ions in aqueous solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 379
20.5.1 Dissociation in water . . . . . . . . . . . . . . . . . . . . . . . . . . . . 380
20.5.2 Ions and water hardness . . . . . . . . . . . . . . . . . . . . . . . . . . . 382
20.5.3 The pH scale . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 382
20.5.4 Acid rain . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 384
20.6 Electrolytes, ionisation and conductivity . . . . . . . . . . . . . . . . . . . . . . 386
20.6.1 Electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 386
20.6.2 Non-electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 387
20.6.3 Factors that affect the conductivity of water . . . . . . . . . . . . . . . . 387
20.7 Precipitation reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 389
20.8 Testing for common anions in solution . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.1 Test for a chloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.2 Test for a sulphate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.3 Test for a carbonate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 392
20.8.4 Test for bromides and iodides . . . . . . . . . . . . . . . . . . . . . . . . 392
20.9 Threats to the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 393
20.10Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 394
21 The Lithosphere - Grade 11
397
21.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 397
21.2 The chemistry of the earth’s crust . . . . . . . . . . . . . . . . . . . . . . . . . 398
21.3 A brief history of mineral use . . . . . . . . . . . . . . . . . . . . . . . . . . . . 399
21.4 Energy resources and their uses . . . . . . . . . . . . . . . . . . . . . . . . . . . 400
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21.5 Mining and Mineral Processing: Gold . . . . . . . . . . . . . . . . . . . . . . . . 401
21.5.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 401
21.5.2 Mining the Gold . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 401
21.5.3 Processing the gold ore . . . . . . . . . . . . . . . . . . . . . . . . . . . 401
21.5.4 Characteristics and uses of gold . . . . . . . . . . . . . . . . . . . . . . . 402
21.5.5 Environmental impacts of gold mining . . . . . . . . . . . . . . . . . . . 404
21.6 Mining and mineral processing: Iron . . . . . . . . . . . . . . . . . . . . . . . . 406
21.6.1 Iron mining and iron ore processing . . . . . . . . . . . . . . . . . . . . . 406
21.6.2 Types of iron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 407
21.6.3 Iron in South Africa . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 408
21.7 Mining and mineral processing: Phosphates . . . . . . . . . . . . . . . . . . . . 409
21.7.1 Mining phosphates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 409
21.7.2 Uses of phosphates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 409
21.8 Energy resources and their uses: Coal . . . . . . . . . . . . . . . . . . . . . . . 411
21.8.1 The formation of coal . . . . . . . . . . . . . . . . . . . . . . . . . . . . 411
21.8.2 How coal is removed from the ground . . . . . . . . . . . . . . . . . . . 411
21.8.3 The uses of coal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 412
21.8.4 Coal and the South African economy . . . . . . . . . . . . . . . . . . . . 412
21.8.5 The environmental impacts of coal mining . . . . . . . . . . . . . . . . . 413
21.9 Energy resources and their uses: Oil . . . . . . . . . . . . . . . . . . . . . . . . 414
21.9.1 How oil is formed . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 414
21.9.2 Extracting oil . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 414
21.9.3 Other oil products . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 415
21.9.4 The environmental impacts of oil extraction and use . . . . . . . . . . . 415
21.10Alternative energy resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 415
21.11Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 417
22 The Atmosphere - Grade 11
421
22.1 The composition of the atmosphere . . . . . . . . . . . . . . . . . . . . . . . . 421
22.2 The structure of the atmosphere . . . . . . . . . . . . . . . . . . . . . . . . . . 422
22.2.1 The troposphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 422
22.2.2 The stratosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 422
22.2.3 The mesosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424
22.2.4 The thermosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424
22.3 Greenhouse gases and global warming . . . . . . . . . . . . . . . . . . . . . . . 426
22.3.1 The heating of the atmosphere . . . . . . . . . . . . . . . . . . . . . . . 426
22.3.2 The greenhouse gases and global warming . . . . . . . . . . . . . . . . . 426
22.3.3 The consequences of global warming . . . . . . . . . . . . . . . . . . . . 429
22.3.4 Taking action to combat global warming . . . . . . . . . . . . . . . . . . 430
22.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 431
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23 The Chemical Industry - Grade 12
435
23.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 435
23.2 Sasol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 435
23.2.1 Sasol today: Technology and production . . . . . . . . . . . . . . . . . . 436
23.2.2 Sasol and the environment . . . . . . . . . . . . . . . . . . . . . . . . . 440
23.3 The Chloralkali Industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 442
23.3.1 The Industrial Production of Chlorine and Sodium Hydroxide . . . . . . . 442
23.3.2 Soaps and Detergents . . . . . . . . . . . . . . . . . . . . . . . . . . . . 446
23.4 The Fertiliser Industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 450
23.4.1 The value of nutrients . . . . . . . . . . . . . . . . . . . . . . . . . . . . 450
23.4.2 The Role of fertilisers . . . . . . . . . . . . . . . . . . . . . . . . . . . . 450
23.4.3 The Industrial Production of Fertilisers . . . . . . . . . . . . . . . . . . . 451
23.4.4 Fertilisers and the Environment: Eutrophication . . . . . . . . . . . . . . 454
23.5 Electrochemistry and batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . 456
23.5.1 How batteries work . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 456
23.5.2 Battery capacity and energy . . . . . . . . . . . . . . . . . . . . . . . . 457
23.5.3 Lead-acid batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 457
23.5.4 The zinc-carbon dry cell . . . . . . . . . . . . . . . . . . . . . . . . . . . 459
23.5.5 Environmental considerations . . . . . . . . . . . . . . . . . . . . . . . . 460
23.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 461
A GNU Free Documentation License
467
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Chapter 3
The Atom - Grade 10
We have now looked at many examples of the types of matter and materials that exist around
us, and we have investigated some of the ways that materials are classified. But what is it that
makes up these materials? And what makes one material different from another? In order to
understand this, we need to take a closer look at the building block of matter, the atom. Atoms
are the basis of all the structures and organisms in the universe. The planets, the sun, grass and
trees, the air we breathe, and people are all made up of different combinations of atoms.
3.1
Models of the Atom
It is important to realise that a lot of what we know about the structure of atoms has been
developed over a long period of time. This is often how scientific knowledge develops, with
one person building on the ideas of someone else. We are going to look at how our modern
understanding of the atom has evolved over time.
The idea of atoms was invented by two Greek philosophers, Democritus and Leucippus in the
fifth century BC. The Greek word ατ oµoν (atom) means indivisible because they believed that
atoms could not be broken into smaller pieces.
Nowadays, we know that atoms are made up of a positively charged nucleus in the centre
surrounded by negatively charged electrons. However, in the past, before the structure of the
atom was properly understood, scientists came up with lots of different models or pictures to
describe what atoms look like.
Definition: Model
A model is a representation of a system in the real world. Models help us to understand
systems and their properties. For example, an atomic model represents what the structure
of an atom could look like, based on what we know about how atoms behave. It is not
necessarily a true picture of the exact structure of an atom.
3.1.1
The Plum Pudding Model
After the electron was discovered by J.J. Thomson in 1897, people realised that atoms were made
up of even smaller particles than they had previously thought. However, the atomic nucleus had
not been discovered yet, and so the ’plum pudding model’ was put forward in 1904. In this
model, the atom is made up of negative electrons that float in a soup of positive charge, much
like plums in a pudding or raisins in a fruit cake (figure 3.1). In 1906, Thomson was awarded
the Nobel Prize for his work in this field. However, even with the Plum Pudding Model, there
was still no understanding of how these electrons in the atom were arranged.
35
3.1
CHAPTER 3. THE ATOM - GRADE 10
-
-
-
electrons
-
-
-
’soup’ of positive charge
-
Figure 3.1: A schematic diagram to show what the atom looked like according to the Plum
Pudding model
The discovery of radiation was the next step along the path to building an accurate picture of
atomic structure. In the early twentieth century, Marie Curie and her husband discovered that
some elements (the radioactive elements) emit particles, which are able to pass through matter
in a similar way to X-rays (read more about this in chapter 7). It was Ernest Rutherford who, in
1911, used this discovery to revise the model of the atom.
3.1.2
Rutherford’s model of the atom
Radioactive elements emit different types of particles. Some of these are positively charged alpha
(α) particles. Rutherford carried out a series of experiments where he bombarded sheets of gold
foil with these particles, to try to get a better understanding of where the positive charge in the
atom was. A simplified diagram of his experiment is shown in figure 3.2.
C
B
b
gold sheet
b
radioactive
substance
A
α particles
A
α particles
b
b
b
b
b
b
b
C
(a)
B
C
screen
b
b
b
b
b
b
b
b
b
b
b
b
A
b
b
b
b
B
b
b
b
b
nucleus of
gold atom
(b)
Figure 3.2: Rutherford’s gold foil experiment. Figure (a) shows the path of the α particles after
they hit the gold sheet. Figure (b) shows the arrangement of atoms in the gold sheets, and the
path of the α particles in relation to this.
Rutherford set up his experiment so that a beam of alpha particles was directed at the gold
sheets. Behind the gold sheets, was a screen made of zinc sulfide. This screen allowed Rutherford to see where the alpha particles were landing. Rutherford knew that the electrons in the gold
atoms would not really affect the path of the alpha particles, because the mass of an electron is
so much smaller than that of a proton. He reasoned that the positively charged protons would
be the ones to repel the positively charged alpha particles and alter their path.
36
CHAPTER 3. THE ATOM - GRADE 10
3.1
What he discovered was that most of the alpha particles passed through the foil undisturbed,
and could be detected on the screen directly behind the foil (A). Some of the particles ended up
being slightly deflected onto other parts of the screen (B). But what was even more interesting
was that some of the particles were deflected straight back in the direction from where they
had come (C)! These were the particles that had been repelled by the positive protons in the
gold atoms. If the Plum Pudding model of the atom were true, then Rutherford would have
expected much more repulsion since the positive charge, according to that model, is distributed
throughout the atom. But this was not the case. The fact that most particles passed straight
through suggested that the positive charge was concentrated in one part of the atom only.
Rutherford’s work led to a change in ideas around the atom. His new model described the
atom as a tiny, dense, positively charged core called a nucleus, surrounded by lighter, negatively
charged electrons. Another way of thinking about this model was that the atom was seen to be
like a mini solar system where the electrons orbit the nucleus like planets orbiting around the
sun. A simplified picture of this is shown in figure 3.3.
b
b
b
electron orbiting the nucleus
b
b
b
nucleus (containing protons and neutrons)
Figure 3.3: Rutherford’s model of the atom
3.1.3
The Bohr Model
There were, however, some problems with this model: for example it could not explain the very
interesting observation that atoms only emit light at certain wavelengths or frequencies. Niels
Bohr solved this problem by proposing that the electrons could only orbit the nucleus in certain
special orbits at different energy levels around the nucleus. The exact energies of the orbitals in
each energy level depends on the type of atom. Helium for example, has different energy levels
to Carbon. If an electron jumps down from a higher energy level to a lower energy level, then
light is emitted from the atom. The energy of the light emitted is the same as the gap in the
energy between the two energy levels. You can read more about this in section 3.6. The distance
between the nucleus and the electron in the lowest energy level of a hydrogen atom is known as
the Bohr radius.
teresting Light has the properties of both a particle and a wave! Einstein discovered that
Interesting
Fact
Fact
light comes in energy packets which are called photons. When an electron in
an atom changes energy levels, a photon of light is emitted. This photon has the
same energy as the difference between the two electron energy levels.
37
3.2
3.2
CHAPTER 3. THE ATOM - GRADE 10
How big is an atom?
It is difficult sometimes to imagine the size of an atom, or its mass, because we cannot see them,
and also because we are not used to working with such small measurements.
3.2.1
How heavy is an atom?
It is possible to determine the mass of a single atom in kilograms. But to do this, you would
need very modern mass spectrometers, and the values you would get would be very clumsy and
difficult to use. The mass of a carbon atom, for example, is about 1.99 x 10−26 kg, while the
mass of an atom of hydrogen is about 1.67 x 10−27 kg. Looking at these very small numbers
makes it difficult to compare how much bigger the mass of one atom is when compared to another.
To make the situation simpler, scientists use a different unit of mass when they are describing
the mass of an atom. This unit is called the atomic mass unit (amu). We can abbreviate
(shorten) this unit to just ’u’. If we give carbon an atomic mass of 12 u, then the mass of an
atom of hydrogen will be 1 u. You can check this by dividing the mass of a carbon atom in
kilograms (see above) by the mass of a hydrogen atom in kilograms (you will need to use a
calculator for this!). If you do this calculation, you will see that the mass of a carbon atom is
twelve times greater than the mass of a hydrogen atom. When we use atomic mass units instead
of kilograms, it becomes easier to see this. Atomic mass units are therefore not giving us the
actual mass of an atom, but rather its mass relative to the mass of other atoms in the Periodic
Table. The atomic masses of some elements are shown in table 3.1 below.
Table 3.1: The atomic mass of a number of elements
Element
Atomic mass (u)
Nitrogen (N)
14
Bromine (Br)
80
Magnesium (Mg)
24
Potassium (K)
39
Calcium (Ca)
40
Oxygen (O)
16
The actual value of 1 atomic mass unit is 1.67 x 10−24 g or 1.67 x 10−27 kg. This is a very tiny
mass!
3.2.2
pm stands for
picometres. 1
pm = 10−12
m
How big is an atom?
Atomic diameter also varies depending on the element. On average, the diameter of an atom
ranges from 100 pm (Helium) to 670 pm (Caesium). Using different units, 100 pm = 1 Angstrom,
and 1 Angstrom = 10−10 m. That is the same as saying that 1 Angstrom = 0,0000000010 m
or that 100 pm = 0,0000000010 m! In other words, the diameter of an atom ranges from
0.0000000010 m to 0.0000000067 m. This is very small indeed.
3.3
Atomic structure
As a result of the models that we discussed in section 3.1, scientists now have a good idea of
what an atom looks like. This knowledge is important because it helps us to understand things
like why materials have different properties and why some materials bond with others. Let us
now take a closer look at the microscopic structure of the atom.
So far, we have discussed that atoms are made up of a positively charged nucleus surrounded
by one or more negatively charged electrons. These electrons orbit the nucleus.
38
CHAPTER 3. THE ATOM - GRADE 10
3.3.1
3.3
The Electron
The electron is a very light particle. It has a mass of 9.11 x 10−31 kg. Scientists believe that the
electron can be treated as a point particle or elementary particle meaning that it can’t be broken
down into anything smaller. The electron also carries one unit of negative electric charge which
is the same as 1.6 x 10−19 C (Coulombs).
3.3.2
The Nucleus
Unlike the electron, the nucleus can be broken up into smaller building blocks called protons
and neutrons. Together, the protons and neutrons are called nucleons.
The Proton
Each proton carries one unit of positive electric charge. Since we know that atoms are electrically
neutral, i.e. do not carry any extra charge, then the number of protons in an atom has to be the
same as the number of electrons to balance out the positive and negative charge to zero. The
total positive charge of a nucleus is equal to the number of protons in the nucleus. The proton
is much heavier than the electron (10 000 times heavier!) and has a mass of 1.6726 x 10−27 kg.
When we talk about the atomic mass of an atom, we are mostly referring to the combined mass
of the protons and neutrons, i.e. the nucleons.
The Neutron
The neutron is electrically neutral i.e. it carries no charge at all. Like the proton, it is much
heavier than the electron and its mass is 1.6749 x 10−27 kg (slightly heavier than the proton).
teresting Rutherford predicted (in 1920) that another kind of particle must be present in
Interesting
Fact
Fact
the nucleus along with the proton. He predicted this because if there were only
positively charged protons in the nucleus, then it should break into bits because
of the repulsive forces between the like-charged protons! Also, if protons were
the only particles in the nucleus, then a helium nucleus (atomic number 2) would
have two protons and therefore only twice the mass of hydrogen. However, it is
actually four times heavier than hydrogen. This suggested that there must be
something else inside the nucleus as well as the protons. To make sure that the
atom stays electrically neutral, this particle would have to be neutral itself. In
1932 James Chadwick discovered the neutron and measured its mass.
Mass (kg)
Units of charge
Charge (C)
proton
1.6726 x 10−27
+1
1.6 x 10−19
neutron
1.6749 x 10−27
0
0
electron
9.11 x 10−31
-1
-1.6 x 10−19
Table 3.2: Summary of the particles inside the atom
teresting Unlike the electron which is thought to be a point particle and unable to be
Interesting
Fact
Fact
broken up into smaller pieces, the proton and neutron can be divided. Protons
and neutrons are built up of smaller particles called quarks. The proton and
neutron are made up of 3 quarks each.
39
3.4
3.4
CHAPTER 3. THE ATOM - GRADE 10
Atomic number and atomic mass number
The chemical properties of an element are determined by the charge of its nucleus, i.e. by the
number of protons. This number is called the atomic number and is denoted by the letter Z.
Definition: Atomic number (Z)
The number of protons in an atom
The mass of an atom depends on how many nucleons its nucleus contains. The number of
nucleons, i.e. the total number of protons plus neutrons, is called the atomic mass number
and is denoted by the letter A.
Definition: Atomic mass number (A)
The number of protons and neutrons in the nucleus of an atom
Standard notation shows the chemical symbol, the atomic mass number and the atomic number
of an element as follows:
number of nucleons
A
ZX
chemical symbol
number of protons
For example, the iron nucleus which has 26 protons and 30 neutrons, is denoted as
56
26 Fe
,
where the total nuclear charge is Z = 26 and the mass number A = 56. The number of neutrons
is simply the difference N = A − Z.
40
CHAPTER 3. THE ATOM - GRADE 10
3.4
Important:
Don’t confuse the notation we have used above, with the way this information appears
on the Periodic Table. On the Periodic Table, the atomic number usually appears in the
top lefthand corner of the block or immediately above the element’s symbol. The number
below the element’s symbol is its relative atomic mass. This is not exactly the same as
the atomic mass number. This will be explained in section 3.5. The example of iron is used
again below.
26
Fe
55.85
You will notice in the example of iron that the atomic mass number is more or less the same as
its atomic mass. Generally, an atom that contains n protons and neutrons (i.e. Z = n), will have
a mass approximately equal to n u. The reason is that a C-12 atom has 6 protons, 6 neutrons
and 6 electrons, with the protons and neutrons having about the same mass and the electron
mass being negligible in comparison.
Exercise: The structure of the atom
1. Explain the meaning of each of the following terms:
(a) nucleus
(b) electron
(c) atomic mass
2. Complete the following table: (Note: You will see that the atomic masses on
the Periodic Table are not whole numbers. This will be explained later. For
now, you can round off to the nearest whole number.)
Element Atomic
Atomic
Number
Number
Number
mass
number
of pro- of elec- of neutons
trons
trons
Mg
24
12
O
8
17
Ni
28
40
20
Zn
0
C
12
6
3. Use standard notation to represent the following elements:
(a) potassium
(b) copper
(c) chlorine
4. For the element
35
17 Cl,
give the number of ...
(a) protons
(b) neutrons
(c) electrons
... in the atom.
41
3.5
CHAPTER 3. THE ATOM - GRADE 10
5. Which of the following atoms has 7 electrons?
(a) 52 He
(b) 13
6 C
(c) 73 Li
(d) 15
7 N
6. In each of the following cases, give the number or the element symbol represented by ’X’.
(a) 40
18 X
(b) x20 Ca
(c) 31
x P
7. Complete the following table:
A
Z
N
235
92 U
238
92 U
In these two different forms of Uranium...
(a) What is the same?
(b) What is different?
Uranium can occur in different forms, called isotopes. You will learn more
about isotopes in section 3.5.
3.5
3.5.1
Isotopes
What is an isotope?
If a few neutrons are added to or removed from a nucleus, the chemical properties of the atom
will stay the same because its charge is still the same. Therefore, the chemical properties of an
element depend on the number of protons inside the atom. This means that such an atom should
remain in the same place in the Periodic table. For example, no matter how many neutrons we
add or subtract from a nucleus with 6 protons, that element will always be called carbon and
have the element symbol C (see the Table of Elements). Atoms which have the same number
of protons, but a different number of neutrons, are called isotopes.
Definition: Isotope
The isotope of a particular element, is made up of atoms which have the same number of
protons as the atoms in the orginal element, but a different number of neutrons.
The different isotopes of an element have the same atomic number Z but different mass numbers
A because they have a different number of neutrons N . The chemical properties of the different
isotopes of an element are the same, but they might vary in how stable their nucleus is. Note
that if an element is written for example as C-12, the ’12’ is the atomic mass of that atom. So,
Cl-35 has an atomic mass of 35 u, while Cl-37 has an atomic mass of 37 u.
teresting In Greek, “same place” reads as ὶσoς τ òπoς (isos topos). This is why atoms
Interesting
Fact
Fact
which have the same number of protons, but different numbers of neutrons, are
called isotopes. They are in the same place on the Periodic Table!
The following worked examples will help you to understand the concept of an isotope better.
42
CHAPTER 3. THE ATOM - GRADE 10
3.5
Worked Example 1: Isotopes
Question: For the element
234
92 U
(uranium), use standard notation to describe:
1. the isotope with 2 fewer neutrons
2. the isotope with 4 more neutrons
Answer
Step 1 : Go over the definition of isotope
We know that isotopes of any element have the same number of protons (same
atomic number) in each atom which means that they have the same chemical symbol. However, they have a different number of neutrons, and therefore a different
mass number.
Step 2 : Rewrite the notation for the isotopes
Therefore, any isotope of uranium will have the symbol:
U
Also, since the number of protons in uranium isotopes is always the same, we can
write down the atomic number:
92 U
Now, if the isotope we want has 2 fewer neutrons than
original mass number and subtract 2, which gives:
234
92 U,
then we take the
232
92 U
Following the steps above, we can write the isotope with 4 more neutrons as:
238
92 U
Worked Example 2: Isotopes
Question: Which of the following are isotopes of
•
40
19 K
•
42
20 Ca
•
40
18 Ar
40
20 Ca?
Answer
Step 1 : Go over the definition of isotope:
We know that isotopes have the same atomic number but different mass numbers.
Step 2 : Determine which of the elements listed fits the definition of an
isotope.
You need to look for the element that has the same atomic number but a different
atomic mass number. The only element is 12
20 Ca. What is different is that there are
2 more neutrons than in the original element.
43
3.5
CHAPTER 3. THE ATOM - GRADE 10
Worked Example 3: Isotopes
Question: For the sulfur isotope
1.
2.
3.
4.
33
16 S,
give the number of...
protons
nucleons
electrons
neutrons
Answer
Step 1 : Determine the number of protons by looking at the atomic number,
Z.
Z = 16, therefore the number of protons is 16 (answer to (a)).
Step 2 : Determine the number of nucleons by looking at the atomic mass
number, A.
A = 33, therefore the number of nucleons is 33 (answer to (b)).
Step 3 : Determine the number of electrons
The atom is neutral, and therefore the number of electrons is the same as the number of protons. The number of electrons is 16 (answer to (c)).
Step 4 : Calculate the number of neutrons
N = A − Z = 33 − 16 = 17
The number of neutrons is 17 (answer to (d)).
Exercise: Isotopes
1. Atom A has 5 protons and 5 neutrons, and atom B has 6 protons and 5 neutrons.
These atoms are...
(a) allotropes
(b) isotopes
(c) isomers
(d) atoms of different elements
34
2. For the sulfur isotopes, 32
16 S and 16 S, give the number of...
(a) protons
(b) nucleons
(c) electrons
(d) neutrons
3. Which of the following are isotopes of Cl35 ?
(a) 17
35 Cl
(b) 35
17 Cl
(c) 37
17 Cl
4. Which of the following are isotopes of U-235? (X represents an element symbol)
(a) 238
92 X
(b) 238
90 X
(c) 235
92 X
44
CHAPTER 3. THE ATOM - GRADE 10
3.5.2
3.5
Relative atomic mass
It is important to realise that the atomic mass of isotopes of the same element will be different
because they have a different number of nucleons. Chlorine, for example, has two common
isotopes which are chlorine-35 and chlorine-37. Chlorine-35 has an atomic mass of 35 u, while
chlorine-37 has an atomic mass of 37 u. In the world around us, both of these isotopes occur
naturally. It doesn’t make sense to say that the element chlorine has an atomic mass of 35
u, or that it has an atomic mass of 37 u. Neither of these are absolutely true since the mass
varies depending on the form in which the element occurs. We need to look at how much more
common one is than the other in order to calculate the relative atomic mass for the element
chlorine. This is then the number that will appear on the Periodic Table.
Definition: Relative atomic mass
Relative atomic mass is the average mass of one atom of all the naturally occurring isotopes
of a particular chemical element, expressed in atomic mass units.
Worked Example 4: The relative atomic mass of an isotopic element
Question: The element chlorine has two isotopes, chlorine-35 and chlorine-37. The
abundance of these isotopes when they occur naturally is 75% chlorine-35 and 25%
chlorine-37. Calculate the average relative atomic mass for chlorine.
Answer
Step 1 : Calculate the mass contribution of chlorine-35 to the average relative atomic mass
Contribution of Cl-35 = (75/100 x 35) = 26.25 u
Step 2 : Calculate the contribution of chlorine-37 to the average relative
atomic mass
Contribution of Cl-37 = (25/100 x 37) = 9.25 u
Step 3 : Add the two values to arrive at the average relative atomic mass of
chlorine
Relative atomic mass of chlorine = 26.25 u + 9.25 u = 35.5 u.
If you look on the periodic table, the average relative atomic mass for chlorine is 35,5
u. You will notice that for many elements, the relative atomic mass that is shown
is not a whole number. You should now understand that this number is the average
relative atomic mass for those elements that have naturally occurring isotopes.
Exercise: Isotopes
You are given a sample that contains carbon-12 and carbon-14.
1. Complete the table below:
45
3.6
CHAPTER 3. THE ATOM - GRADE 10
Isotope
Carbon-12
Carbon-14
Chlorine-35
Chlorine-37
Z
A
Protons
Neutrons
Electrons
2. If the sample you have contains 90% carbon-12 and 10% carbon-14, calculate
the relative atomic mass of an atom in that sample.
3. In another sample, you have 22.5% Cl-37 and 77.5% Cl-35. Calculate the
relative atomic mass of an atom in that sample.
Activity :: Group Discussion : The changing nature of scientific knowledge
Scientific knowledge is not static: it changes and evolves over time as scientists
build on the ideas of others to come up with revised (and often improved) theories and
ideas. In this chapter for example, we saw how peoples’ understanding of atomic
structure changed as more information was gathered about the atom. There are
many more examples like this one in the field of science. Think for example, about
our knowledge of the solar system and the origin of the universe, or about the particle
and wave nature of light.
Often, these changes in scientific thinking can be very controversial because they
disturb what people have come to know and accept. It is important that we realise
that what we know now about science may also change. An important part of
being a scientist is to be a critical thinker. This means that you need to question
information that you are given and decide whether it is accurate and whether it can
be accepted as true. At the same time, you need to learn to be open to new ideas
and not to become stuck in what you believe is right... there might just be something
new waiting around the corner that you have not thought about!
In groups of 4-5, discuss the following questions:
• Think about some other examples where scientific knowledge has changed because of new ideas and discoveries:
–
–
–
–
What were these new ideas?
Were they controversial? If so, why?
What role (if any) did technology play in developing these new ideas?
How have these ideas affected the way we understand the world?
• Many people come up with their own ideas about how the world works. The
same is true in science. So how do we, and other scientists, know what to
believe and what not to? How do we know when new ideas are ’good’ science
or ’bad’ science? In your groups, discuss some of the things that would need
to be done to check whether a new idea or theory was worth listening to, or
whether it was not.
• Present your ideas to the rest of the class.
3.6
3.6.1
Energy quantisation and electron configuration
The energy of electrons
You will remember from our earlier discussions, that an atom is made up of a central nucleus,
which contains protons and neutrons, and that this nucleus is surrounded by electrons. Although
46
CHAPTER 3. THE ATOM - GRADE 10
3.6
these electrons all have the same charge and the same mass, each electron in an atom has a
different amount of energy. Electrons that have the lowest energy are found closest to the nucleus
where the attractive force of the positively charged nucleus is the greatest. Those electrons that
have higher energy, and which are able to overcome the attractive force of the nucleus, are found
further away.
3.6.2
Energy quantisation and line emission spectra
If the energy of an atom is increased (for example when a substance is heated), the energy of the
electrons inside the atom can be increased (when an electron has a higher energy than normal
it is said to be ”excited”). For the excited electron to go back to its original energy (called the
ground state), it needs to release energy. It releases energy by emitting light. If one heats up
different elements, one will see that for each element, light is emitted only at certain frequencies
(or wavelengths). Instead of a smooth continuum of frequencies, we see lines (called emission
lines) at particular frequencies. These frequencies correspond to the energy of the emitted light.
If electrons could be excited to any energy and lose any amount of energy, there would be a
continuous spread of light frequencies emitted. However, the sharp lines we see mean that there
are only certain particular energies that an electron can be excited to, or can lose, for each
element.
You can think of this like going up a flight of steps: you can’t lift your foot by any amount to
go from the ground to the first step. If you lift your foot too low you’ll bump into the step and
be stuck on the ground level. You have to lift your foot just the right amount (the height of the
step) to go to the next step, and so on. The same goes for electrons and the amount of energy
they can have. This is called quantisation of energy because there are only certain quantities
of energy that an electron can have in an atom. Like steps, we can think of these quantities as
energy levels in the atom. The energy of the light released when an electron drops down from
a higher energy level to a lower energy level is the same as the difference in energy between the
two levels.
3.6.3
Electron configuration
Electrons are arranged in energy levels around the nucleus of an atom. Electrons that are in the
energy level that is closest to the nucleus, will have the lowest energy and those further away
will have a higher energy. Each energy level can only hold a certain number of electrons, and
an electron will only be found in the second energy level once the first energy level is full. The
same rule applies for the higher energy levels. You will need to learn the following rules:
• The 1st energy level can hold a maximum of 2 electrons
• The 2nd energy level can hold a maximum of 8 electrons
• The 3rd energy level can hold a maximum of 8 electrons
• If the number of electrons in the atom is greater than 18, they will need to move to the
4th energy level.
In the following examples, the energy levels are shown as concentric circles around the central
nucleus.
1. Lithium
Lithium (Li) has an atomic number of 3, meaning that in a neutral atom, the number of
electrons will also be 3. The first two electrons are found in the first energy level, while
the third electron is found in the second energy level (figure 3.11).
2. Fluorine
Fluorine (F) has an atomic number of 9, meaning that a neutral atom also has 9 electrons.
The first 2 electrons are found in the first energy level, while the other 7 are found in the
second energy level (figure 3.12).
47
3.6
CHAPTER 3. THE ATOM - GRADE 10
second energy level
electrons
first energy level
nucleus, containing 3 protons and 4 neutrons
Figure 3.4: The arrangement of electrons in a lithium atom.
Figure 3.5: The arrangement of electrons in a fluorine atom.
3. Argon
Argon has an atomic number of 18, meaning that a neutral atom also has 18 electrons.
The first 2 electrons are found in the first energy level, the next 8 are found in the second
energy level, and the last 8 are found in the third energy level (figure 3.6).
Figure 3.6: The arrangement of electrons in an argon atom.
But the situation is slightly more complicated than this. Within each energy level, the electrons
move in orbitals. An orbital defines the spaces or regions where electrons move.
Definition: Atomic orbital
An atomic orbital is the region in which an electron may be found around a single atom.
There are different orbital shapes, but we will be dealing with only two. These are the ’s’ and ’p’
orbitals (there are also ’d’ and ’f’ orbitals). The first energy level contains only one ’s’ orbital,
the second energy level contains one ’s’ orbital and three ’p’ orbitals and the third energy level
also contains one ’s’ orbital and three ’p’ orbitals. Within each energy level, the ’s’ orbital is at
a lower energy than the ’p’ orbitals. This arrangement is shown in figure 3.7.
When we want to show how electrons are arranged in an atom, we need to remember the
following principles:
48
CHAPTER 3. THE ATOM - GRADE 10
3.6
4s
3p
Third main
energy level
2p
Second main
energy level
3s
E
N
E
R
G
Y
2s
First main
energy level
1s
Figure 3.7: The positions of the first ten orbits of an atom on an energy diagram. Note that
each block is able to hold two electrons.
49
3.6
CHAPTER 3. THE ATOM - GRADE 10
• Each orbital can only hold two electrons. Electrons that occur together in an orbital are
called an electron pair. These electrons spin in opposite directions around their own axes.
• An electron will always try to enter an orbital with the lowest possible energy.
• An electron will occupy an orbital on its own, rather than share an orbital with another
electron. An electron would also rather occupy a lower energy orbital with another electron,
before occupying a higher energy orbital. In other words, within one energy level, electrons
will fill an ’s’ orbital before starting to fill ’p’ orbitals.
The way that electrons are arranged in an atom is called its electron configuration.
Definition: Electron configuration
Electron configuration is the arrangement of electrons in an atom, molecule, or other physical
structure.
An element’s electron configuration can be represented using Aufbau diagrams or energy level
diagrams. An Aufbau diagram uses arrows to represent electrons. You can use the following
steps to help you to draw an Aufbau diagram:
1. Determine the number of electrons that the atom has.
2. Fill the ’s’ orbital in the first energy level (the 1s orbital) with the first two electrons.
3. Fill the ’s’ orbital in the second energy level (the 2s orbital) with the second two electrons.
4. Put one electron in each of the three ’p’ orbitals in the second energy level (the 2p orbitals),
and then if there are still electrons remaining, go back and place a second electron in each
of the 2p orbitals to complete the electron pairs.
5. Carry on in this way through each of the successive energy levels until all the electrons
have been drawn.
Important:
When there are two electrons in an orbital, the electrons are called an electron pair. If the
orbital only has one electron, this electron is said to be an unpaired electron. Electron
pairs are shown with arrows in opposite directions. This is because when two electrons
occupy the same orbital, they spin in opposite directions on their axes.
An Aufbau diagram for the element Lithium is shown in figure 3.8.
2s
1s
Figure 3.8: The electron configuration of Lithium, shown on an Aufbau diagram
A special type of notation is used to show an atom’s electron configuration. The notation describes the energy levels, orbitals and the number of electrons in each. For example, the electron
configuration of lithium is 1s2 2s1 . The number and letter describe the energy level and orbital,
and the number above the orbital shows how many electrons are in that orbital.
Aufbau diagrams for the elements fluorine and argon are shown in figures 3.9 and 3.10 respectively. Using standard notation, the electron configuration of fluorine is 1s2 2s2 2p5 and the
electron configuration of argon is 1s2 2s2 2p6 3s2 3p6 .
50
CHAPTER 3. THE ATOM - GRADE 10
3.6
4s
3p
3s
2p
E
N
E
R
G
Y
2s
1s
Fluorine
Figure 3.9: An Aufbau diagram showing the electron configuration of fluorine
3.6.4
Core and valence electrons
Electrons in the outermost energy level of an atom are called valence electrons. The electrons
that are in the energy shells closer to the nucleus are called core electrons. Core electrons are
all the electrons in an atom, excluding the valence electrons. An element that has its valence
energy level full is more stable and less likely to react than other elements with a valence energy
level that is not full.
Definition: Valence electrons
The electrons in the outer energy level of an atom
Definition: Core electrons
All the electrons in an atom, excluding the valence electrons
3.6.5
The importance of understanding electron configuration
By this stage, you may well be wondering why it is important for you to understand how electrons
are arranged around the nucleus of an atom. Remember that during chemical reactions, when
atoms come into contact with one another, it is the electrons of these atoms that will interact
first. More specifically, it is the valence electrons of the atoms that will determine how they
51
3.6
CHAPTER 3. THE ATOM - GRADE 10
4s
3p
3s
2p
2s
1s
Argon
Figure 3.10: An Aufbau diagram showing the electron configuration of argon
react with one another.
To take this a step further, an atom is at its most stable (and therefore unreactive) when all
its orbitals are full. On the other hand, an atom is least stable (and therefore most reactive)
when its valence electron orbitals are not full. This will make more sense when we go on to
look at chemical bonding in a later chapter. To put it simply, the valence electrons are largely
responsible for an element’s chemical behaviour, and elements that have the same number of
valence electrons often have similar chemical properties.
Exercise: Energy diagrams and electrons
1. Draw Aufbau diagrams to show the electron configuration of each of the following elements:
(a) magnesium
(b) potassium
(c) sulfur
(d) neon
(e) nitrogen
2. Use the Aufbau diagrams you drew to help you complete the following table:
52
CHAPTER 3. THE ATOM - GRADE 10
Element
No.
of
energy
levels
3.7
No.
of
core electrons
No.
of
valence
electrons
Electron
configuration
(standard
notation)
Mg
K
S
Ne
N
3. Rank the elements used above in order of increasing reactivity. Give reasons
for the order you give.
Activity :: Group work : Building a model of an atom
Earlier in this chapter, we talked about different ’models’ of the atom. In science,
one of the uses of models is that they can help us to understand the structure of
something that we can’t see. In the case of the atom, models help us to build a
picture in our heads of what the atom looks like.
Models are often simplified. The small toy cars that you may have played with as
a child are models. They give you a good idea of what a real car looks like, but they
are much smaller and much simpler. A model cannot always be absolutely accurate
and it is important that we realise this so that we don’t build up a false idea about
something.
In groups of 4-5, you are going to build a model of an atom. Before you start,
think about these questions:
• What information do I know about the structure of the atom? (e.g. what parts
make it up? how big is it?)
• What materials can I use to represent these parts of the atom as accurately as
I can?
• How will I put all these different parts together in my model?
As a group, share your ideas and then plan how you will build your model. Once
you have built your model, discuss the following questions:
• Does our model give a good idea of what the atom actually looks like?
• In what ways is our model inaccurate? For example, we know that electrons
move around the atom’s nucleus, but in your model, it might not have been
possible for you to show this.
• Are there any ways in which our model could be improved?
Now look at what other groups have done. Discuss the same questions for each
of the models you see and record your answers.
3.7
3.7.1
Ionisation Energy and the Periodic Table
Ions
In the previous section, we focused our attention on the electron configuration of neutral atoms.
In a neutral atom, the number of protons is the same as the number of electrons. But what
53
3.7
CHAPTER 3. THE ATOM - GRADE 10
happens if an atom gains or loses electrons? Does it mean that the atom will still be part of the
same element?
A change in the number of electrons of an atom does not change the type of atom that it is.
However, the charge of the atom will change. If electrons are added, then the atom will become
more negative. If electrons are taken away, then the atom will become more positive. The atom
that is formed in either of these cases is called an ion. Put simply, an ion is a charged atom.
Definition: Ion
An ion is a charged atom. A positively charged ion is called a cation e.g. N a+ , and a
negatively charged ion is called an anion e.g. F − . The charge on an ion depends on the
number of electrons that have been lost or gained.
Look at the following examples. Notice the number of valence electrons in the neutral atom,
the number of electrons that are lost or gained, and the final charge of the ion that is formed.
Lithium
A lithium atoms loses one electrons to form a positive ion (figure 3.11).
1 electron lost
Li
Li+
Li atom with 3 electrons
Li+ ion with only 2 electrons
Figure 3.11: The arrangement of electrons in a lithium ion.
In this example, the lithium atom loses an electron to form the cation Li+ .
Fluorine
A fluorine atom gains one electron to form a negative ion (figure 3.12).
1 electron gained
F
F−
Figure 3.12: The arrangement of electrons in a fluorine ion.
Activity :: Investigation : The formation of ions
54
CHAPTER 3. THE ATOM - GRADE 10
3.7
1. Use the diagram for lithium as a guide and draw similar diagrams to show how
each of the following ions is formed:
(a) Mg2+
(b) Na+
(c) Cl−
(d) O2−
2. Do you notice anything interesting about the charge on each of these ions?
Hint: Look at the number of valence electrons in the neutral atom and the
charge on the final ion.
Observations:
Once you have completed the activity, you should notice that:
• In each case the number of electrons that is either gained or lost, is the same as the number
of electrons that are needed for the atoms to achieve a full or an empty valence energy
level.
• If you look at an energy level diagram for sodium (Na), you will see that in a neutral
atom, there is only one valence electron. In order to achieve an empty valence level, and
therefore a more stable state for the atom, this electron will be lost.
• In the case of oxygen (O), there are six valence electrons. To fill the valence energy level,
it makes more sense for this atom to gain two electrons. A negative ion is formed.
3.7.2
Ionisation Energy
Ionisation energy is the energy that is needed to remove one electron from an atom. The ionisation energy will be different for different atoms.
The second ionisation energy is the energy that is needed to remove a second electron from an
atom, and so on. As an energy level becomes more full, it becomes more and more difficult to
remove an electron and the ionisation energy increases. On the Periodic Table of the Elements,
a group is a vertical column of the elements, and a period is a horizontal row. In the periodic
table, ionisation energy increases across a period, but decreases as you move down a group.
The lower the ionisation energy, the more reactive the element will be because there is a greater
chance of electrons being involved in chemical reactions. We will look at this in more detail in
the next section.
Exercise: The formation of ions
Match the information in column A with the information in column B by writing
only the letter (A to I) next to the question number (1 to 7)
55
3.8
CHAPTER 3. THE ATOM - GRADE 10
1. A positive ion that has 3 less electrons
than its neutral atom
2. An ion that has 1 more electron than
its neutral atom
3.
The anion that is formed when
bromine gains an electron
4. The cation that is formed from a magnesium atom
5. An example of a compound ion
6. A positive ion with the electron configuration of argon
7. A negative ion with the electron configuration of neon
A. Mg2+
B. Cl−
C. CO2−
3
D. Al3+
E. Br2−
F. K+
G. Mg+
H. O2−
I. Br−
3.8
The Arrangement of Atoms in the Periodic Table
The periodic table of the elements is a tabular method of showing the chemical elements.
Most of the work that was done to arrive at the periodic table that we know, can be attributed to
a man called Dmitri Mendeleev in 1869. Mendeleev was a Russian chemist who designed the
table in such a way that recurring (”periodic”) trends in the properties of the elements could be
shown. Using the trends he observed, he even left gaps for those elements that he thought were
’missing’. He even predicted the properties that he thought the missing elements would have
when they were discovered. Many of these elements were indeed discovered and Mendeleev’s
predictions were proved to be correct.
To show the recurring properties that he had observed, Mendeleev began new rows in his table
so that elements with similar properties were in the same vertical columns, called groups. Each
row was referred to as a period. One important feature to note in the periodic table is that all
the non-metals are to the right of the zig-zag line drawn under the element boron. The rest of
the elements are metals, with the exception of hydrogen which occurs in the first block of the
table despite being a non-metal.
Group
group number
1
8
H
2
3
4
5
6
7
He
Li
Be
B
C
N
O
F
Ne
Na Mg
Al
Si
P
S
Cl
Ar
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Br
Kr
K
Ca Sc
Ti
V
Period
Figure 3.13: A simplified diagram showing part of the Periodic Table
3.8.1
Groups in the periodic table
A group is a vertical column in the periodic table, and is considered to be the most important
way of classifying the elements. If you look at a periodic table, you will see the groups numbered
56
CHAPTER 3. THE ATOM - GRADE 10
3.8
at the top of each column. The groups are numbered from left to right as follows: 1, 2, then an
open space which contains the transition elements, followed by groups 3 to 8. These numbers
are normally represented using roman numerals. In some periodic tables, all the groups are
numbered from left to right from number 1 to number 18. In some groups, the elements display
very similar chemical properties, and the groups are even given separate names to identify them.
The characteristics of each group are mostly determined by the electron configuration of the
atoms of the element.
• Group 1: These elements are known as the alkali metals and they are very reactive.
Hydrogen
Lithium
Sodium
Potassium
Figure 3.14: Electron diagrams for some of the Group 1 elements
Activity :: Investigation : The properties of elements
Refer to figure 3.14.
1. Use a Periodic Table to help you to complete the last two diagrams for
sodium (Na) and potassium (K).
2. What do you notice about the number of electrons in the valence energy
level in each case?
3. Explain why elements from group 1 are more reactive than elements from
group 2 on the periodic table (Hint: Think back to ’ionisation energy’).
• Group 2: These elements are known as the alkali earth metals. Each element only has
two valence electrons and so in chemical reactions, the group 2 elements tend to lose these
electrons so that the energy shells are complete. These elements are less reactive than
those in group 1 because it is more difficult to lose two electrons than it is to lose one.
Group 3 elements have three valence electrons.
Important: The number of valence electrons of an element corresponds to its group number
on the periodic table.
• Group 7: These elements are known as the halogens. Each element is missing just one
electron from its outer energy shell. These elements tend to gain electrons to fill this shell,
rather than losing them.
• Group 8: These elements are the noble gases. All of the energy shells of the halogens are
full, and so these elements are very unreactive.
• Transition metals: The differences between groups in the transition metals are not usually
dramatic.
57
3.8
CHAPTER 3. THE ATOM - GRADE 10
Helium
Lithium
Figure 3.15: Electron diagrams for two of the noble gases, helium (He) and neon (Ne).
It is worth noting that in each of the groups described above, the atomic diameter of the
elements increases as you move down the group. This is because, while the number of valence
electrons is the same in each element, the number of core electrons increases as one moves down
the group.
3.8.2
Periods in the periodic table
A period is a horizontal row in the periodic table of the elements. Some of the trends that can
be observed within a period are highlighted below:
• As you move from one group to the next within a period, the number of valence electrons
increases by one each time.
• Within a single period, all the valence electrons occur in the same energy shell. If the
period increases, so does the energy shell in which the valence electrons occur.
• In general, the diameter of atoms decreases as one moves from left to right across a period.
Consider the attractive force between the positively charged nucleus and the negatively
charged electrons in an atom. As you move across a period, the number of protons in
each atom increases. The number of electrons also increases, but these electrons will still
be in the same energy shell. As the number of protons increases, the force of attraction
between the nucleus and the electrons will increase and the atomic diameter will decrease.
• Ionisation energy increases as one moves from left to right across a period. As the valence
electron shell moves closer to being full, it becomes more difficult to remove electrons. The
opposite is true when you move down a group in the table because more energy shells are
being added. The electrons that are closer to the nucleus ’shield’ the outer electrons from
the attractive force of the positive nucleus. Because these electrons are not being held
to the nucleus as strongly, it is easier for them to be removed and the ionisation energy
decreases.
• In general, the reactivity of the elements decreases from left to right across a period.
Exercise: Trends in ionisation energy
Refer to the data table below which gives the ionisation energy (in kJ/mol) and
atomic number (Z) for a number of elements in the periodic table:
58
CHAPTER 3. THE ATOM - GRADE 10
Z
1
2
3
4
5
6
7
8
9
Ionisation energy
1310
2360
517
895
797
1087
1397
1307
1673
3.9
Z
10
11
12
13
14
15
16
17
18
Ionisation energy
2072
494
734
575
783
1051
994
1250
1540
1. Draw a line graph to show the relationship between atomic number (on the
x-axis) and ionisation energy (y-axis).
2. Describe any trends that you observe.
3. Explain why...
(a) the ionisation energy for Z=2 is higher than for Z=1
(b) the ionisation energy for Z=3 is lower than for Z=2
(c) the ionisation energy increases between Z=5 and Z=7
Exercise: Elements in the Periodic Table
Refer to the elements listed below:
Lithium (Li); Chlorine (Cl); Magnesium (Mg); Neon (Ne); Oxygen (O); Calcium
(Ca); Carbon (C)
Which of the elements listed above:
1. belongs to Group 1
2. is a halogen
3. is a noble gas
4. is an alkali metal
5. has an atomic number of 12
6. has 4 neutrons in the nucleus of its atoms
7. contains electrons in the 4th energy level
8. has only one valence electron
9. has all its energy orbitals full
10. will have chemical properties that are most similar
11. will form positive ions
3.9
Summary
• Much of what we know today about the atom, has been the result of the work of a
number of scientists who have added to each other’s work to give us a good understanding
of atomic structure.
59
3.9
CHAPTER 3. THE ATOM - GRADE 10
• Some of the important scientific contributors include J.J.Thomson (discovery of the electron, which led to the Plum Pudding Model of the atom), Ernest Rutherford (discovery
that positive charge is concentrated in the centre of the atom) and Niels Bohr (the
arrangement of electrons around the nucleus in energy levels).
• Because of the very small mass of atoms, their mass in measured in atomic mass units
(u). 1 u = 1.67 × 10−24 g.
• An atom is made up of a central nucleus (containing protons and neutrons), surrounded
by electrons.
• The atomic number (Z) is the number of protons in an atom.
• The atomic mass number (A) is the number of protons and neutrons in the nucleus of
an atom.
• The standard notation that is used to write an element, is A
Z X, where X is the element
symbol, A is the atomic mass number and Z is the atomic number.
• The isotope of a particular element is made up of atoms which have the same number
of protons as the atoms in the original element, but a different number of neutrons. This
means that not all atoms of an element will have the same atomic mass.
• The relative atomic mass of an element is the average mass of one atom of all the
naturally occurring isotopes of a particular chemical element, expressed in atomic mass
units. The relative atomic mass is written under the elements’ symbol on the Periodic
Table.
• The energy of electrons in an atom is quantised. Electrons occur in specific energy levels
around an atom’s nucleus.
• Within each energy level, an electron may move within a particular shape of orbital. An
orbital defines the space in which an electron is most likely to be found. There are different
orbital shapes, including s, p, d and f orbitals.
• Energy diagrams such as Aufbau diagrams are used to show the electron configuration
of atoms.
• The electrons in the outermost energy level are called valence electrons.
• The electrons that are not valence electrons are called core electrons.
• Atoms whose outermost energy level is full, are less chemically reactive and therefore more
stable, than those atoms whose outer energy level is not full.
• An ion is a charged atom. A cation is a positively charged ion and an anion is a negatively
charged ion.
• When forming an ion, an atom will lose or gain the number of electrons that will make its
valence energy level full.
• An element’s ionisation energy is the energy that is needed to remove one electron from
an atom.
• Ionisation energy increases across a period in the Periodic Table.
• Ionisation energy decreases down a group in the Periodic Table.
Exercise: Summary
1. Write down only the word/term for each of the following descriptions.
(a) The sum of the number of protons and neutrons in an atom
60
CHAPTER 3. THE ATOM - GRADE 10
3.9
(b) The defined space around an atom’s nucleus, where an electron is most
likely to be found
2. For each of the following, say whether the statement is True or False. If it is
False, re-write the statement correctly.
22
(a) 20
10 Ne and 10 Ne each have 10 protons, 12 electrons and 12 neutrons.
(b) The atomic mass of any atom of a particular element is always the same.
(c) It is safer to use helium gas rather than hydrogen gas in balloons.
(d) Group 1 elements readily form negative ions.
3. Multiple choice questions: In each of the following, choose the one correct
answer.
(a) The three basic components of an atom are:
i. protons, neutrons, and ions
ii. protons, neutrons, and electrons
iii. protons, neutrinos, and ions
iv. protium, deuterium, and tritium
(b) The charge of an atom is...
i. positive
ii. neutral
iii. negative
(c) If Rutherford had used neutrons instead of alpha particles in his scattering
experiment, the neutrons would...
i. not deflect because they have no charge
ii. have deflected more often
iii. have been attracted to the nucleus easily
iv. have given the same results
(d) Consider the isotope 234
92 U. Which of the following statements is true?
i. The element is an isotope of 234
94 Pu
ii. The element contains 234 neutrons
iii. The element has the same electron configuration as 238
92 U
iv. The element has an atomic mass number of 92
(e) The electron configuration of an atom of chlorine can be represented using
the following notation:
i. 1s2 2s8 3s7
ii. 1s2 2s2 2p6 3s2 3p5
iii. 1s2 2s2 2p6 3s2 3p6
iv. 1s2 2s2 2p5
4. The following table shows the first ionisation energies for the elements of period
1 and 2.
Period
Element
First ionisation energy (kJ.mol−1 )
1
H
He
Li
Be
B
C
N
O
F
Ne
1312
2372
520
899
801
1086
1402
1314
1681
2081
2
(a) What is the meaning of the term first ionisation energy ?
(b) Identify the pattern of first ionisation energies in a period.
(c) Which TWO elements exert the strongest attractive forces on their electrons? Use the data in the table to give a reason for your answer.
61
3.9
CHAPTER 3. THE ATOM - GRADE 10
(d) Draw Aufbau diagrams for the TWO elements you listed in the previous
question and explain why these elements are so stable.
(e) It is safer to use helium gas than hydrogen gas in balloons. Which property
of helium makes it a safer option?
(f) ’Group 1 elements readily form positive ions’.
Is this statement correct? Explain your answer by referring to the table.
62
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You may extract a single document from such a collection, and distribute it individually under
this License, provided you insert a copy of this License into the extracted document, and follow
this License in all other respects regarding verbatim copying of that document.
AGGREGATION WITH INDEPENDENT WORKS
A compilation of the Document or its derivatives with other separate and independent documents
or works, in or on a volume of a storage or distribution medium, is called an “aggregate” if the
copyright resulting from the compilation is not used to limit the legal rights of the compilation’s
users beyond what the individual works permit. When the Document is included an aggregate,
this License does not apply to the other works in the aggregate which are not themselves derivative
works of the Document.
If the Cover Text requirement of section A is applicable to these copies of the Document, then if
the Document is less than one half of the entire aggregate, the Document’s Cover Texts may be
placed on covers that bracket the Document within the aggregate, or the electronic equivalent
of covers if the Document is in electronic form. Otherwise they must appear on printed covers
that bracket the whole aggregate.
TRANSLATION
Translation is considered a kind of modification, so you may distribute translations of the Document under the terms of section A. Replacing Invariant Sections with translations requires
special permission from their copyright holders, but you may include translations of some or
all Invariant Sections in addition to the original versions of these Invariant Sections. You may
include a translation of this License, and all the license notices in the Document, and any Warranty Disclaimers, provided that you also include the original English version of this License and
the original versions of those notices and disclaimers. In case of a disagreement between the
translation and the original version of this License or a notice or disclaimer, the original version
will prevail.
If a section in the Document is Entitled “Acknowledgements”, “Dedications”, or “History”, the
requirement (section A) to Preserve its Title (section A) will typically require changing the actual
title.
TERMINATION
You may not copy, modify, sub-license, or distribute the Document except as expressly provided
for under this License. Any other attempt to copy, modify, sub-license or distribute the Document
is void, and will automatically terminate your rights under this License. However, parties who
have received copies, or rights, from you under this License will not have their licenses terminated
so long as such parties remain in full compliance.
FUTURE REVISIONS OF THIS LICENSE
The Free Software Foundation may publish new, revised versions of the GNU Free Documentation
License from time to time. Such new versions will be similar in spirit to the present version, but
may differ in detail to address new problems or concerns. See http://www.gnu.org/copyleft/.
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APPENDIX A. GNU FREE DOCUMENTATION LICENSE
Each version of the License is given a distinguishing version number. If the Document specifies
that a particular numbered version of this License “or any later version” applies to it, you have the
option of following the terms and conditions either of that specified version or of any later version
that has been published (not as a draft) by the Free Software Foundation. If the Document does
not specify a version number of this License, you may choose any version ever published (not as
a draft) by the Free Software Foundation.
ADDENDUM: How to use this License for your documents
To use this License in a document you have written, include a copy of the License in the document
and put the following copyright and license notices just after the title page:
c YEAR YOUR NAME. Permission is granted to copy, distribute and/or
Copyright modify this document under the terms of the GNU Free Documentation License,
Version 1.2 or any later version published by the Free Software Foundation; with no
Invariant Sections, no Front-Cover Texts, and no Back-Cover Texts. A copy of the
license is included in the section entitled “GNU Free Documentation License”.
If you have Invariant Sections, Front-Cover Texts and Back-Cover Texts, replace the “with...Texts.”
line with this:
with the Invariant Sections being LIST THEIR TITLES, with the Front-Cover Texts being LIST,
and with the Back-Cover Texts being LIST.
If you have Invariant Sections without Cover Texts, or some other combination of the three,
merge those two alternatives to suit the situation.
If your document contains nontrivial examples of program code, we recommend releasing these
examples in parallel under your choice of free software license, such as the GNU General Public
License, to permit their use in free software.
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