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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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
xiv
Chapter 1
Classification of Matter - Grade 10
All the objects that we see in the world around us, are made of matter. Matter makes up the
air we breathe, the ground we walk on, the food we eat and the animals and plants that live
around us. Even our own human bodies are made of matter!
Different objects can be made of different types of matter, or materials. For example, a cupboard (an object) is made of wood, nails and hinges (the materials). The properties of the
materials will affect the properties of the object. In the example of the cupboard, the strength
of the wood and metals make the cupboard strong and durable. In the same way, the raincoats
that you wear during bad weather, are made of a material that is waterproof. The electrical wires
in your home are made of metal because metals are a type of material that is able to conduct
electricity. It is very important to understand the properties of materials, so that we can use
them in our homes, in industry and in other applications. In this chapter, we will be looking at
different types of materials and their properties.
The diagram below shows one way in which matter can be classified (grouped) according to
its different properties. As you read further in this chapter, you will see that there are also
other ways of classifying materials, for example according to whether they are good electrical
conductors.
MATTER
MIXTURES
Homogeneous
PURE SUBSTANCES
Heterogeneous
Elements
Metals
Magnetic
Compounds
Non-metals
Non-magnetic
Figure 1.1: The classification of matter
1.1
Mixtures
We see mixtures all the time in our everyday lives. A stew, for example, is a mixture of different
foods such as meat and vegetables; sea water is a mixture of water, salt and other substances,
and air is a mixture of gases such as carbon dioxide, oxygen and nitrogen.
5
1.1
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
Definition: Mixture
A mixture is a combination of more than one substance, where these substances are not
bonded to each other.
In a mixture, the substances that make up the mixture:
• are not in a fixed ratio
Imagine, for example, that you have a 250 ml beaker of water. It doesn’t matter whether
you add 20 g, 40 g, 100 g or any other mass of sand to the water; it will still be called a
mixture of sand and water.
• keep their physical properties
In the example we used of the sand and water, neither of these substances has changed in
any way when they are mixed together. Even though the sand is in water, it still has the
same properties as when it was out of the water.
• can be separated by mechanical means
To separate something by ’mechanical means’, means that there is no chemical process
involved. In our sand and water example, it is possible to separate the mixture by simply
pouring the water through a filter. Something physical is done to the mixture, rather than
something chemical.
Some other examples of mixtures include blood (a mixture of blood cells, platelets and plasma),
steel (a mixture of iron and other materials) and the gold that is used to make jewellery. The
gold in jewellery is not pure gold but is a mixture of metals. The carat of the gold gives an idea
of how much gold is in the item.
We can group mixtures further by dividing them into those that are heterogeneous and those
that are homogeneous.
1.1.1
Heterogeneous mixtures
A heterogeneous mixture does not have a definite composition. Think of a pizza, that is a
mixture of cheese, tomato, mushrooms and peppers. Each slice will probably be slightly different
from the next because the toppings like the mushrooms and peppers are not evenly distributed.
Another example would be granite, a type of rock. Granite is made up of lots of different mineral
substances including quartz and feldspar. But these minerals are not spread evenly through the
rock and so some parts of the rock may have more quartz than others. Another example is
a mixture of oil and water. Although you may add one substance to the other, they will stay
separate in the mixture. We say that these heterogeneous mixtures are non-uniform, in other
words they are not exactly the same throughout.
Definition: Heterogeneous mixture
A heterogeneous mixture is one that is non-uniform, and where the different components
of the mixture can be seen.
1.1.2
Homogeneous mixtures
A homogeneous mixture has a definite composition, and specific properties. In a homogeneous
mixture, the different parts cannot be seen. A solution of salt dissolved in water is an example
of a homogeneous mixture. When the salt dissolves, it will spread evenly through the water so
that all parts of the solution are the same, and you can no longer see the salt as being separate
from the water. Think also of a powdered drink that you mix with water. Provided you give the
container a good shake after you have added the powder to the water, the drink will have the
same sweet taste for anyone who drinks it, it won’t matter whether they take a sip from the top
6
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.1
or from the bottom. The air we breathe is another example of a homogeneous mixture since it is
made up of different gases which are in a constant ratio, and which can’t be distinguished from
each other.
Definition: Homogeneous mixture
A homogeneous mixture is one that is uniform, and where the different components of the
mixture cannot be seen.
An alloy is a homogeneous mixture of two or more elements, at least one of which is a metal,
where the resulting material has metallic properties. Alloys are usually made to improve on the
properties of the elements that make them up. Steel for example, is much stronger than iron,
which is its main component.
1.1.3
Separating mixtures
Sometimes it is important to be able to separate a mixture. There are lots of different ways to
do this. These are some examples:
• Filtration
A piece of filter paper in a funnel can be used to separate a mixture of sand and water.
• Heating / evaporation
Sometimes, heating a solution causes the water to evaporate, leaving the other part of the
mixture behind. You can try this using a salt solution.
• Centrifugation
This is a laboratory process which uses the centrifugal force of spinning objects to separate
out the heavier substances from a mixture. This process is used to separate the cells and
plasma in blood. When the test tubes that hold the blood are spun round in the machine,
the heavier cells sink to the bottom of the test tube. Can you think of a reason why it
might be important to have a way of separating blood in this way?
• Dialysis
This is an interesting way of separating a mixture because it can be used in some important
applications. Dialysis works using a process called diffusion. Diffusion takes place when
one substance in a mixture moves from an area where it has a high concentration to an
area where its concentration is lower. This movement takes place across a semi-permeable
membrane. A semi-permeable membrane is a barrier that lets some things move across it,
but not others. This process is very important for people whose kidneys are not functioning
properly, an illness called renal failure.
teresting Normally, healthy kidneys remove waste products from the blood. When a person
Interesting
Fact
Fact
has renal failure, their kidneys cannot do this any more, and this can be lifethreatening. Using dialysis, the blood of the patient flows on one side of a
semi-permeable membrane. On the other side there will be a fluid that has no
waste products but lots of other important substances such as potassium ions
(K + ) that the person will need. Waste products from the blood diffuse from
where their concentration is high (i.e. in the person’s blood) into the ’clean’
fluid on the other side of the membrane. The potassium ions will move in the
opposite direction from the fluid into the blood. Through this process, waste
products are taken out of the blood so that the person stays healthy.
7
1.1
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
Activity :: Investigation : The separation of a salt solution
Aim:
To demonstrate that a homogeneous salt solution can be separated using physical
methods.
Apparatus:
glass beaker, salt, water, retort stand, bunsen burner.
Method:
1. Pour a small amount of water (about 20 ml) into a beaker.
2. Measure a teaspoon of salt and pour this into the water.
3. Stir until the salt dissolves completely. This is now called a salt solution. This
salt solution is a homogeneous mixture.
4. Place the beaker on a retort stand over a bunsen burner and heat gently. You
should increase the heat until the water almost boils.
5. Watch the beaker until all the water has evaporated. What do you see in the
beaker?
H2 O
salt
solution
water evaporates
when the solution
is heated
salt crystals
remain at the
bottom of the beaker
stand
bunsen
burner
Results:
The water evaporates from the beaker and tiny grains of salt remain at the
bottom.
Conclusion:
The sodium chloride solution, which was a homogeneous mixture of salt and
water, has been separated using heating and evaporation.
Activity :: Discussion : Separating mixtures
Work in groups of 3-4
Imagine that you have been given a container which holds a mixture of sand,
iron filings (small pieces of iron metal), salt and small stones of different sizes. Is
this a homogeneous or a heterogeneous mixture? In your group, discuss how you
would go about separating this mixture into the four materials that it contains.
8
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.2
Exercise: Mixtures
1. Which of the following subtances are mixtures?
(a)
(b)
(c)
(d)
(e)
(f)
tap water
brass (an alloy of copper and zinc)
concrete
aluminium
Coca cola
distilled water
2. In each of the examples above, say whether the mixture is homogeneous or
heterogeneous
1.2
Pure Substances: Elements and Compounds
Any material that is not a mixture, is called a pure substance. Pure substances include elements
and compounds. It is much more difficult to break down pure substances into their parts, and
complex chemical methods are needed to do this.
1.2.1
Elements
An element is a chemical substance that can’t be divided or changed into other chemical
substances by any ordinary chemical means. The smallest unit of an element is the atom.
Definition: Element
An element is a substance that cannot be broken down into other substances through
chemical means.
There are 109 known elements. Most of these are natural, but some are man-made. The
elements we know are represented in the Periodic Table of the Elements, where each element
is abbreviated to a chemical symbol. Examples of elements are magnesium (Mg), hydrogen (H),
oxygen (O) and carbon (C). On the Periodic Table you will notice that some of the abbreviations
do not seem to match the elements they represent. The element iron, for example, has the
chemical formula Fe. This is because the elements were originally given Latin names. Iron has
the abbreviation Fe because its Latin name is ’ferrum’. In the same way, sodium’s Latin name
is ’natrium’ (Na) and gold’s is ’aurum’ (Au).
1.2.2
Compounds
A compound is a chemical substance that forms when two or more elements combine in a fixed
ratio. Water (H2 O), for example, is a compound that is made up of two hydrogen atoms for
every one oxygen atom. Sodium chloride (NaCl) is a compound made up of one sodium atom
for every chlorine atom. An important characteristic of a compound is that it has a chemical
formula, which describes the ratio in which the atoms of each element in the compound occur.
Definition: Compound
A substance made up of two or more elements that are joined together in a fixed ratio.
Diagram 1.2 might help you to understand the difference between the terms element, mixture
and compound. Iron (Fe) and sulfur (S) are two elements. When they are added together, they
9
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
Fe
A mixture of iron and sulfur
S
Fe S
S
An atom
of the element iron
(Fe)
S
S
S
Fe
Fe
Fe
Fe
Fe
Fe
S
S
Fe
S
S
Fe
Fe
S
An atom
of the element sulfur (S)
Fe S
1.3
The compound iron sulfide
(FeS)
Figure 1.2: Understanding the difference between a mixture and a compound
form a mixture or iron and sulfur. The iron and sulfur are not joined together. However, if
the mixture is heated, a new compound is formed, which is called iron sulfide (FeS). In this
compound, the iron and sulfur are joined to each other in a ratio of 1:1. In other words, one
atom of iron is joined to one atom of sulfur in the compound iron sulfide.
Exercise: Elements, mixtures and compounds
1. In the following table, tick whether each of the substances listed is a mixture
or a pure substance. If it is a mixture, also say whether it is a homogeneous or
heterogeneous mixture.
Substance
Mixture or pure
Homogeneous
heterogeneous
mixture
or
fizzy colddrink
steel
oxygen
iron filings
smoke
limestone (CaCO3 )
2. In each of the following cases, say whether the substance is an element, a
mixture or a compound.
(a)
(b)
(c)
(d)
(e)
1.3
Cu
iron and sulfur
Al
H2 SO4
SO3
Giving names and formulae to substances
It is easy to describe elements and mixtures. But how are compounds named? In the example
of iron sulfide that was used earlier, which element is named first, and which ’ending’ is given
to the compound name (in this case, the ending is -ide)?
The following are some guidelines for naming compounds:
10
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.3
1. The compound name will always include the names of the elements that are part of it.
• A compound of iron (Fe) and sulfur (S) is iron sulf ide (FeS)
• A compound of potassium (K) and bromine (S) is potassium bromide (KBr)
• A compound of sodium (Na) and chlorine (Cl) is sodium chlor ide (NaCl)
2. In a compound, the element that is to the left and lower down on the Periodic Table,
is used first when naming the compound. In the example of NaCl, sodium is a group 1
element on the left hand side of the table, while chlorine is in group 7 on the right of the
table. Sodium therefore comes first in the compound name. The same is true for FeS and
KBr.
3. The symbols of the elements can be used to represent compounds e.g. FeS, NaCl and
KBr. These are called chemical formulae. In these three examples, the ratio of the
elements in each compound is 1:1. So, for FeS, there is one atom of iron for every atom
of sulfur in the compound.
4. A compound may contain compound ions. Some of the more common compound ions
and their names are shown below.
Name of compound ion
Carbonate
sulphate
Hydroxide
Ammonium
Nitrate
Hydrogen carbonate
Phosphate
Chlorate
Cyanide
Chromate
Permanganate
formula
CO3 2−
SO4 2−
OH−
NH4 +
NO3 −
HCO3 −
PO4 3−
ClO3 −
CN−
CrO4 2−
MnO4 −
5. When there are only two elements in the compound, the compound is often given a suffix
(ending) of -ide. You would have seen this in some of the examples we have used so far.
When a non-metal is combined with oxygen to form a negative ion (anion) which then
combines with a positive ion (cation) from hydrogen or a metal, then the suffix of the
name will be ...ate or ...ite. NO−
3 for example, is a negative ion, which may combine with
a cation such as hydrogen (HNO3 ) or a metal like potassium (KNO3 ). The NO−
3 anion
has the name nitrate. SO3 in a formula is sulphite, e.g. sodium sulphite (Na2 SO3 ). SO4
is sulphate and PO4 is phosphate.
6. Prefixes can be used to describe the ratio of the elements that are in the compound. You
should know the following prefixes: ’mono’ (one), ’di’ (two) and ’tri’ (three).
• CO (carbon monoxide) - There is one atom of oxygen for every one atom of carbon
• N O2 (nitrogen dioxide) - There are two atoms of oxygen for every one atom of
nitrogen
• SO3 (sulfur trioxide) - There are three atoms of oxygen for every one atom of sulfur
Important:
When numbers are written as ’subscripts’ in compounds (i.e. they are written below the
element symbol), this tells us how many atoms of that element there are in relation to other
elements in the compound. For example in nitrogen dioxide (NO2 ) there are two oxygen
atoms for every one atom of nitrogen. In sulfur trioxide (SO3 ), there are three oxygen atoms
for every one atom of sulfur in the compound. Later, when we start looking at chemical
equations, you will notice that sometimes there are numbers before the compound name.
For example, 2H2 O means that there are two molecules of water, and that in each molecule
there are two hydrogen atoms for every one oxygen atom.
11
1.3
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
Exercise: Naming compounds
1. The formula for calcium carbonate is CaCO3 .
(a) Is calcium carbonate a mixture or a compound? Give a reason for your
answer.
(b) What is the ratio of Ca:C:O atoms in the formula?
2. Give the name of each of the following substances.
(a)
(b)
(c)
(d)
(e)
(f)
KBr
HCl
KMnO4
NO2
NH4 OH
Na2 SO4
3. Give the chemical formula for each of the following compounds.
(a)
(b)
(c)
(d)
(e)
potassium nitrate
sodium iodide
barium sulphate
nitrogen dioxide
sodium monosulphate
4. Refer to the diagram below, showing sodium chloride and water, and then
answer the questions that follow.
(a)
(b)
(c)
(d)
What is the chemical formula for water?
What is the chemical formula for sodium chloride?
Label the water and sodium chloride in the diagram.
Which of the following statements most accurately describes the picture?
i. The picture shows a mixture of an element and a compound
ii. The picture shows a mixture of two compounds
iii. The picture shows two compounds that have been chemically bonded
to each other
5. What is the formula of this molecule?
H
H
H
A
B
C
D
C
C
H
H
C6 H2 O
C2 H6 O
2C6HO
2 CH6 O
12
O
H
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.4
1.4
Metals, Semi-metals and Non-metals
The elements in the Periodic Table can also be divided according to whether they are metals,
semi-metals or non-metals. On the right hand side of the Periodic Table is a dark ’zigzag’ line.
This line separates all the elements that are metals from those that are non-metals. Metals are
found on the left of the line, and non-metals are those on the right. Metals, semi-metals and
non-metals all have their own specific properties.
1.4.1
Metals
Examples of metals include copper (Cu), zinc (Zn), gold (Au) and silver (Ag). On the Periodic
Table, the metals are on the left of the zig-zag line. There are a large number of elements that
are metals. The following are some of the properties of metals:
• Thermal conductors
Metals are good conductors of heat and are therefore used in cooking utensils such as pots
and pans.
• Electrical conductors
Metals are good conductors of electricity, and are therefore used in electrical conducting
wires.
• Shiny metallic lustre
Metals have a characteristic shiny appearance and are often used to make jewellery.
• Malleable
This means that they can be bent into shape without breaking.
• Ductile
Metals can stretched into thin wires such as copper, which can then be used to conduct
electricity.
• Melting point
Metals usually have a high melting point and can therefore be used to make cooking pots
and other equipment that needs to become very hot, without being damaged.
You can see how the properties of metals make them very useful in certain applications.
Activity :: Group Work : Looking at metals
1. Collect a number of metal items from your home or school. Some examples
are listed below:
•
•
•
•
•
•
hammer
electrical wiring
cooking pots
jewellery
burglar bars
coins
2. In groups of 3-4, combine your collection of metal objects.
3. What is the function of each of these objects?
4. Discuss why you think metal was used to make each object. You should consider
the properties of metals when you answer this question.
13
1.5
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.4.2
Non-metals
In contrast to metals, non-metals are poor thermal conductors, good electrical insulators (meaning that they do not conduct electrical charge) and are neither malleable nor ductile. The
non-metals are found on the right hand side of the Periodic Table, and include elements such as
sulfur (S), phosphorus (P), nitrogen (N) and oxygen (O).
1.4.3
Semi-metals
Semi-metals have mostly non-metallic properties. One of their distinguishing characteristics is
that their conductivity increases as their temperature increases. This is the opposite of what
happens in metals. The semi-metals include elements such as silicon (Si) and germanium (Ge).
Notice where these elements are positioned in the Periodic Table.
1.5
Electrical conductors, semi-conductors and insulators
An electrical conductor is a substance that allows an electrical current to pass through it.
Many electrical conductors are metals, but non-metals can also be good conductors. Copper is
one of the best electrical conductors, and this is why it is used to make conducting wire. In
reality, silver actually has an even higher electrical conductivity than copper, but because silver
is so expensive, it is not practical to use it for electrical wiring because such large amounts are
needed. In the overhead power lines that we see above us, aluminium is used. The aluminium
usually surrounds a steel core which adds tensile strength to the metal so that it doesn’t break
when it is stretched across distances. Occasionally gold is used to make wire, not because it is
a particularly good conductor, but because it is very resistant to surface corrosion. Corrosion is
when a material starts to deteriorate at the surface because of its reactions with the surroundings, for example oxygen and water in the air.
An insulator is a non-conducting material that does not carry any charge. Examples of insulators
would be plastic and wood. Do you understand now why electrical wires are normally covered
with plastic insulation? Semi-conductors behave like insulators when they are cold, and like
conductors when they are hot. The elements silicon and germanium are examples of semiconductors.
Definition: Conductors and insulators
A conductor allows the easy movement or flow of something such as heat or electrical charge
through it. Insulators are the opposite to conductors because they inhibit or reduce the flow
of heat, electrical charge, sound etc through them.
Activity :: Experiment : Electrical conductivity
Aim:
To investigate the electrical conductivity of a number of substances
Apparatus:
• two or three cells
• light bulb
• crocodile clips
• wire leads
• a selection of test substances (e.g. a piece of plastic, aluminium can, metal
pencil sharpener, metal magnet, wood, chalk).
14
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.6
light bulb
battery
test substance
X
crocodile clip
Method:
1. Set up the circuit as shown above, so that the test substance is held between
the two crocodile clips. The wire leads should be connected to the cells and
the light bulb should also be connected into the circuit.
2. Place the test substances one by one between the crocodile clips and see what
happens to the light bulb.
Results:
Record your results in the table below:
Test substance
Metal/non-metal
Does
glow?
bulb
Conductor or
insulator
Conclusions:
In the substances that were tested, the metals were able to conduct electricity
and the non-metals were not. Metals are good electrical conductors and non-metals
are not.
1.6
Thermal Conductors and Insulators
A thermal conductor is a material that allows energy in the form of heat, to be transferred
within the material, without any movement of the material itself. An easy way to understand
this concept is through a simple demonstration.
Activity :: Demonstration : Thermal conductivity
Aim:
To demonstrate the ability of different substances to conduct heat.
Apparatus:
15
1.6
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
You will need two cups (made from the same material e.g. plastic); a metal
spoon and a plastic spoon.
Method:
• Pour boiling water into the two cups so that they are about half full.
• At the same time, place a metal spoon into one cup and a plastic spoon in the
other.
• Note which spoon heats up more quickly
Results:
The metal spoon heats up more quickly than the plastic spoon. In other words,
the metal conducts heat well, but the plastic does not.
Conclusion:
Metal is a good thermal conductor, while plastic is a poor thermal conductor.
This explains why cooking pots are metal, but their handles are often plastic or
wooden. The pot itself must be metal so that heat from the cooking surface can
heat up the pot to cook the food inside it, but the handle is made from a poor
thermal conductor so that the heat does not burn the hand of the person who is
cooking.
An insulator is a material that does not allow a transfer of electricity or energy. Materials that
are poor thermal conductors can also be described as being good insulators.
teresting Water is a better thermal conductor than air and conducts heat away from the
Interesting
Fact
Fact
body about 20 times more efficiently than air. A person who is not wearing
a wetsuit, will lose heat very quickly to the water around them and can be
vulnerable to hypothermia. Wetsuits help to preserve body heat by trapping a
layer of water against the skin. This water is then warmed by body heat and acts
as an insulator. Wetsuits are made out of closed-cell, foam neoprene. Neoprene
is a synthetic rubber that contains small bubbles of nitrogen gas when made for
use as wetsuit material. Nitrogen gas has very low thermal conductivity, so it
does not allow heat from the body (or the water trapped between the body and
the wetsuit) to be lost to the water outside of the wetsuit. In this way a person
in a wetsuit is able to keep their body temperature much higher than they would
otherwise.
Activity :: Investigation : A closer look at thermal conductivity
Look at the table below, which shows the thermal conductivity of a number
of different materials, and then answer the questions that follow. The higher the
number in the second column, the better the material is at conducting heat (i.e. it is
a good thermal conductor). Remember that a material that conducts heat efficiently,
will also lose heat more quickly than an insulating material.
16
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
Material
Silver
Stainless steel
Standard glass
Concrete
Red brick
Water
Snow
Wood
Polystyrene
Air
1.7
Thermal Conductivity (W/m/K)
429
16
1.05
0.9 - 2
0.69
0.58
0.5 - 0.25
0.04 - 0.12
0.03
0.024
Use this information to answer the following questions:
1. Name two materials that are good thermal conductors.
2. Name two materials that are good insulators.
3. Explain why:
(a) cooler boxes are often made of polystyrene
(b) homes that are made from wood need less internal heating during the
winter months.
(c) igloos (homes made from snow) are so good at maintaining warm temperatures, even in freezing conditions.
teresting It is a known fact that well-insulated buildings need less energy for heating than
Interesting
Fact
Fact
do buildings that have no insulation. Two building materials that are being used
more and more worldwide, are mineral wool and polystyrene. Mineral wool
is a good insulator because it holds air still in the matrix of the wool so that
heat is not lost. Since air is a poor conductor and a good insulator, this helps
to keep energy within the building. Polystyrene is also a good insulator and is
able to keep cool things cool and hot things hot! It has the added advantage of
being resistant to moisture, mould and mildew.
Remember that concepts such as conductivity and insulation are not only relevant in the building,
industrial and home environments. Think for example of the layer of blubber or fat that we find
in animals. In very cold environments, fat and blubber not only provide protection, but also act
as an insulator to help the animal to keep its body temperature at the right level. This is known
as thermoregulation.
1.7
Magnetic and Non-magnetic Materials
We have now looked at a number of ways in which matter can be grouped, such as into metals,
semi-metals and non-metals; electrical conductors and insulators, and thermal conductors and
insulators. One way in which we can further group metals, is to divide them into those that are
magnetic and those that are non-magnetic.
Definition: Magnetism
Magnetism is one of the phenomena by which materials exert attractive or repulsive forces
on other materials.
17
1.8
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
A metal is said to be ferromagnetic if it can be magnetised (i.e. made into a magnet). If you
hold a magnet very close to a metal object, it may happen that its own electrical field will be
induced and the object becomes magnetic. Some metals keep their magnetism for longer than
others. Look at iron and steel for example. Iron loses its magnetism quite quickly if it is taken
away from the magnet. Steel on the other hand will stay magnetic for a longer time. Steel is
often used to make permanent magnets that can be used for a variety of purposes.
Magnets are used to sort the metals in a scrap yard, in compasses to find direction, in the magnetic strips of video tapes and ATM cards where information must be stored, in computers and
TV’s, as well as in generators and electric motors.
Activity :: Investigation : Magnetism
You can test whether an object is magnetic or not by holding another magnet
close to it. If the object is attracted to the magnet, then it too is magnetic.
Find some objects in your classroom or your home and test whether they are
magnetic or not. Then complete the table below:
Object
Magnetic
magnetic
or
non-
Activity :: Group Discussion : Properties of materials
In groups of 4-5, discuss how our knowledge of the properties of materials has
allowed society to:
• develop advanced computer technology
• provide homes with electricity
• find ways to conserve energy
1.8
Summary
• All the objects and substances that we see in the world are made of matter.
• This matter can be classified according to whether it is a mixture or a pure substance.
• A mixture is a combination of one or more substances that are not chemically bonded
to each other. Examples of mixtures are air (a mixture of different gases) and blood (a
mixture of cells, platelets and plasma).
• The main characteristics of mixtures are that the substances that make them up are not
in a fixed ratio, they keep their individual properties and they can be separated from each
other using mechanical means.
18
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
1.8
• A heterogeneous mixture is non-uniform and the different parts of the mixture can be
seen. An example would be a mixture of sand and salt.
• A homogeneous mixture is uniform, and the different components of the mixture can’t
be seen. An example would be a salt solution. A salt solution is a mixture of salt and
water. The salt dissolves in the water, meaning that you can’t see the individual salt
particles. They are interspersed between the water molecules. Another example is a metal
alloy such as steel.
• Mixtures can be separated using a number of methods such as filtration, heating, evaporation, centrifugation and dialysis.
• Pure substances can be further divided into elements and compounds.
• An element is a substance that can’t be broken down into simpler substances through
chemical means.
• All the elements are recorded in the Periodic Table of the Elements. Each element has
its own chemical symbol. Examples are iron (Fe), sulfur (S), calcium (Ca), magnesium
(Mg) and fluorine (F).
• A compound is a substance that is made up of two or more elements that are chemically
bonded to each other in a fixed ratio. Examples of compounds are sodium chloride (NaCl),
iron sulfide (FeS), calcium carbonate (CaCO3 ) and water (H2 O).
• When naming compounds and writing their chemical formula, it is important to know
the elements that are in the compound, how many atoms of each of these elements will
combine in the compound and where the elements are in the Periodic Table. A number of
rules can then be followed to name the compound.
• Another way of classifying matter is into metals (e.g. iron, gold, copper), semi-metals
(e.g. silicon and germanium) and non-metals (e.g. sulfur, phosphorus and nitrogen).
• Metals are good electrical and thermal conductors, they have a shiny lustre, they are
malleable and ductile, and they have a high melting point. These properties make metals
very useful in electrical wires, cooking utensils, jewellery and many other applications.
• A further way of classifying matter is into electrical conductors, semi-conductors and
insulators.
• An electrical conductor allows an electrical current to pass through it. Most metals are
good electrical conductors.
• An electrical insulator is not able to carry an electrical current. Examples are plastic,
wood, cotton material and ceramic.
• Materials may also be classified as thermal conductors or thermal insulators depending
on whether or not they are able to conduct heat.
• Materials may also be either magnetic or non-magnetic.
Exercise: Summary
1. For each of the following multiple choice questions, choose one correct answer
from the list provided.
A Which of the following can be classified as a mixture:
i. sugar
ii. table salt
iii. air
iv. Iron
B An element can be defined as:
19
1.8
CHAPTER 1. CLASSIFICATION OF MATTER - GRADE 10
i. A substance that cannot be separated into two or more substances by
ordinary chemical (or physical) means
ii. A substance with constant composition
iii. A substance that contains two or more substances, in definite proportion by weight
iv. A uniform substance
2. Classify each of the following substances as an element, a compound, a solution(homogeneous mixture), or a heterogeneous mixture: salt, pure water, soil,
salt water, pure air, carbon dioxide, gold and bronze
3. Look at the table below. In the first column (A) is a list of substances. In the
second column (B) is a description of the group that each of these substances
belongs in. Match up the substance in Column A with the description in
Column B.
Column A
iron
H2 S
sugar solution
sand and stones
steel
Column B
a compound containing 2 elements
a heterogeneous mixture
a metal alloy
an element
a homogeneous mixture
4. You are given a test tube that contains a mixture of iron filings and sulfur. You
are asked to weigh the amount of iron in the sample.
a Suggest one method that you could use to separate the iron filings from
the sulfur.
b What property of metals allows you to do this?
5. Given the following descriptions, write the chemical formula for each of the
following substances:
a silver metal
b a compound that contains only potassium and bromine
c a gas that contains the elements carbon and oxygen in a ratio of 1:2
6. Give the names of each of the following compounds:
a NaBr
b BaSO4
c SO2
7. For each of the following materials, say what properties of the material make
it important in carrying out its particular function.
a
b
c
d
e
f
tar on roads
iron burglar bars
plastic furniture
metal jewellery
clay for building
cotton clothing
20
Chapter 2
What are the objects around us
made of? - Grade 10
2.1
Introduction: The atom as the building block of matter
We have now seen that different materials have different properties. Some materials are metals
and some are non-metals; some are electrical or thermal conductors, while others are not. Depending on the properties of these materials, they can be used in lots of useful applications. But
what is it exactly that makes up these materials? In other words, if we were to break down a
material into the parts that make it up, what would we find? And how is it that a material’s
microscopic structure is able to give it all these different properties?
The answer lies in the smallest building block of matter: the atom. It is the type of atoms, and
the way in which they are arranged in a material, that affects the properties of that substance.
It is not often that substances are found in atomic form. Normally, atoms are bonded to other
atoms to form compounds or molecules. It is only in the noble gases (e.g. helium, neon and
argon) that atoms are found individually and are not bonded to other atoms. We will look at
the reasons for this in a later chapter.
2.2
Molecules
Definition: Molecule
A molecule is a group of two or more atoms that are attracted to each other by relatively
strong forces or bonds
Almost everything around us is made up of molecules. Water is made up of molecules, each of
which has two hydrogen atoms joined to one oxygen atom. Oxygen is a molecule that is made
up of two oxygen atoms that are joined to one another. Even the food that we eat is made
up of molecules that contain atoms of elements such as carbon, hydrogen and oxygen that are
joined to one another in different ways. All of these are known as small molecules because
there are only a few atoms in each molecule. Giant molecules are those where there may be
millions of atoms per molecule. Examples of giant molecules are diamonds, which are made up
of millions of carbon atoms bonded to each other, and metals, which are made up of millions of
metal atoms bonded to each other.
2.2.1
Representing molecules
The structure of a molecule can be shown in many different ways. Sometimes it is easiest to
show what a molecule looks like by using different types of diagrams, but at other times, we
may decide to simply represent a molecule using its chemical formula or its written name.
21
2.2
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
1. Using formulae to show the structure of a molecule
A chemical formula is an abbreviated (shortened) way of describing a molecule, or some
other chemical substance. In chapter 1, we saw how chemical compounds can be represented using element symbols from the Periodic Table. A chemical formula can also tell
us the number of atoms of each element that are in a molecule, and their ratio in that
molecule.
For example, the chemical formula for a molecule of carbon dioxide is:
CO2
The formula above is called the molecular formula of that compound. The formula tells
us that in one molecule of carbon dioxide, there is one atom of carbon and two atoms of
oxygen. The ratio of carbon atoms to oxygen atoms is 1:2.
Definition: Molecular formula
A concise way of expressing information about the atoms that make up a particular chemical
compound. The molecular formula gives the exact number of each type of atom in the
molecule.
A molecule of glucose has the molecular formula:
C6 H12 O6
In each glucose molecule, there are six carbon atoms, twelve hydrogen atoms and six oxygen atoms. The ratio of carbon:hydrogen:oxygen is 6:12:6. We can simplify this ratio to
write 1:2:1, or if we were to use the element symbols, the formula would be written as
CH2 O. This is called the empirical formula of the molecule.
Definition: Empirical formula
This is a way of expressing the relative number of each type of atom in a chemical compound.
In most cases, the empirical formula does not show the exact number of atoms, but rather
the simplest ratio of the atoms in the compound.
The empirical formula is useful when we want to write the formula for a giant molecule.
Since giant molecules may consist of millions of atoms, it is impossible to say exactly how
many atoms are in each molecule. It makes sense then to represent these molecules using
their empirical formula. So, in the case of a metal such as copper, we would simply write
Cu, or if we were to represent a molecule of sodium chloride, we would simply write NaCl.
Chemical formulae therefore tell us something about the types of atoms that are in a
molecule and the ratio in which these atoms occur in the molecule, but they don’t give us
any idea of what the molecule actually looks like, in other words its shape. Another useful
way of representing molecules is to use diagrams.
Another type of formula that can be used to describe a molecule is its structural formula.
A structural formula uses a graphical representation to show a molecule’s structure (figure
2.1).
2. Using diagrams to show the structure of a molecule
Diagrams of molecules are very useful because they give us an idea of the space that is
occupied by the molecule, and they also help us to picture how the atoms are arranged in
the molecule. There are two types of diagrams that are commonly used:
22
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
CH3
(a) C4 H10
(b) C2 H5
2.2
CH
CH3
CH3
Figure 2.1: Diagram showing (a) the molecular, (b) the empirical and (c) the structural formula
of isobutane
• Ball and stick models
This is a 3-dimensional molecular model that uses ’balls’ to represent atoms and
’sticks’ to represent the bonds between them. The centres of the atoms (the balls)
are connected by straight lines which represent the bonds between them. A simplified
example is shown in figure 2.2.
oxygen atom
hydrogen atom
Figure 2.2: A ball and stick model of a water molecule
• Space-filling model
This is also a 3-dimensional molecular model. The atoms are represented by multicoloured spheres. Space-filling models of water and ammonia are shown in figures
2.3 and 2.4.
Figures 2.3 and 2.4 are some examples of simple molecules that are represented in different ways.
oxygen atom
O
hydrogen atoms
H
H
Figure 2.3: A space-filling model and structural formula of a water molecule. Each molecule
is made up of two hydrogen atoms that are attached to one oxygen atom. This is a simple
molecule.
Figure 2.5 shows the bonds between the carbon atoms in diamond, which is a giant
molecule. Each carbon atom is joined to four others, and this pattern repeats itself until
a complex lattice structure is formed. Each black ball in the diagram represents a carbon
atom, and each line represents the bond between two carbon atoms.
teresting Diamonds are most often thought of in terms of their use in the jewellery industry.
Interesting
Fact
Fact
However, about 80% of mined diamonds are unsuitable for use as gemstones and
are therefore used in industry because of their strength and hardness. These
23
2.2
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
nitrogen atom
hydrogen atom
N
H
H
H
Figure 2.4: A space-filling model and structural formula of a molecule of ammonia. Each
molecule is made up of one nitrogen atom and three hydrogen atoms. This is a simple molecule.
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
b
Figure 2.5: Diagrams showing the microscopic structure of diamond. The diagram on the left
shows part of a diamond lattice, made up of numerous carbon atoms. The diagram on the right
shows how each carbon atom in the lattice is joined to four others. This forms the basis of the
lattice structure. Diamond is a giant molecule.
properties of diamonds are due to the strong covalent bonds betwene the carbon
atoms in diamond. The most common uses for diamonds in industry are in
cutting, drilling, grinding, and polishing.
Exercise: Atoms and molecules
1. In each of the following, say whether the chemical substance is made up of
single atoms, simple molecules or giant molecules.
(a) ammonia gas (NH3 )
(b) zinc metal (Zn)
(c) graphite (C)
(d) nitric acid (HNO3 )
(e) neon gas (Ne2 )
2. Refer to the diagram below and then answer the questions that follow:
24
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
O
C
2.3
O
(a) Identify the molecule.
(b) Write the molecular formula for the molecule.
(c) Is the molecule a simple or giant molecule?
3. Represent each of the following molecules using its chemical formula, structural
formula and ball and stick model.
(a) H2
(b) NH3
(c) sulfur dioxide
2.3
Intramolecular and intermolecular forces
When atoms join to form molecules, they are held together by chemical bonds. The type of
bond, and the strength of the bond, depends on the atoms that are involved. These bonds are
called intramolecular forces because they are bonding forces inside a molecule (’intra’ means
’within’ or ’inside’). Sometimes we simply call these intramolecular forces chemical bonds.
Definition: Intramolecular force
The force between the atoms of a molecule, which holds them together.
Examples of the types of chemical bonds that can exist between atoms inside a molecule are
shown below. These will be looked at in more detail in chapter 4.
• Covalent bond
Covalent bonds exist between non-metal atoms e.g. There are covalent bonds between
the carbon and oxygen atoms in a molecule of carbon dioxide.
• Ionic bond
Ionic bonds occur between non-metal and metal atoms e.g. There are ionic bonds between
the sodium and chlorine atoms in a molecule of sodium chloride.
• Metallic bond
Metallic bonds join metal atoms e.g. There are metallic bonds between copper atoms in
a piece of copper metal.
Intermolecular forces are those bonds that hold molecules together. A glass of water for
example, contains many molecules of water. These molecules are held together by intermolecular
forces. The strength of the intermolecular forces is important because they affect properties such
as melting point and boiling point. For example, the stronger the intermolecular forces, the higher
the melting point and boiling point for that substance. The strength of the intermolecular forces
increases as the size of the molecule increases.
25
2.4
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
Definition: Intermolecular force
A force between molecules, which holds them together.
Diagram 2.6 may help you to understand the difference between intramolecular forces and intermolecular forces.
intermolecular forces
intramolecular forces
H
O
O
H
O
O
H
O
O
Figure 2.6: Two representations showing the intermolecular and intramolecular forces in water:
space-filling model and structural formula.
It should be clearer now that there are two types of forces that hold matter together. In the case
of water, there are intramolecular forces that hold the two hydrogen atoms to the oxygen atom
in each molecule of water. There are also intramolecular forces between each of these water
molecules. As mentioned earlier, these forces are very important because they affect many of
the properties of matter such as boiling point, melting point and a number of other properties.
Before we go on to look at some of these examples, it is important that we first take a look at
the Kinetic Theory of Matter.
Exercise: Intramolecular and intermolecular forces
1. Using ammonia gas as an example...
(a) Explain what is meant by an intramolecular force or chemical bond.
(b) Explain what is meant by an intermolecular force.
2. Draw a diagram showing three molecules of carbon dioxide. On the diagram,
show where the intramolecular and intermolecular forces are.
3. Why is it important to understand the types of forces that exist between atoms
and between molecules? Try to use some practical examples in your answer.
2.4
The Kinetic Theory of Matter
The kinetic theory of matter is used to explain why matter exists in different phases (i.e. solid,
liquid and gas), and how matter can change from one phase to the next. The kinetic theory of
matter also helps us to understand other properties of matter. It is important to realise that
what we will go on to describe is only a theory. It cannot be proved beyond doubt, but the fact
that it helps us to explain our observations of changes in phase, and other properties of matter,
suggests that it probably is more than just a theory.
Broadly, the Kinetic Theory of Matter says that:
26
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
2.4
• Matter is made up of particles that are constantly moving.
• All particles have energy, but the energy varies depending on whether the substance is a
solid, liquid or gas. Solid particles have the least energy and gas particles have the most
amount of energy.
• The temperature of a substance is a measure of the average kinetic energy of the particles.
• A change in phase may occur when the energy of the particles is changed.
• There are spaces between the particles of matter.
• There are attractive forces between particles and these become stronger as the particles
move closer together. These attractive forces will either be intramolecular forces (if the
particles are atoms) or intermolecular forces (if the particles are molecules). When the
particles are extremely close, repulsive forces start to act.
Table 2.1 summarises the characteristics of the particles that are in each phase of matter.
Table 2.1: Table summarising the general features of solids, liquids and gases.
Property of matter Gas
Liquid
Gas
Particles
Atoms or molecules
Atoms or molecules
Atoms or molecules
Energy and move- Particles have high Particles have less Low energy - partiment of particles
energy and are con- energy than in the cles vibrate around a
stantly moving
gas phase
fixed point
Spaces between par- Large spaces be- Smaller spaces than Very little space
ticles
cause of high energy in gases
between particles.
Particles are tightly
packed together
Attractive forces be- Weak forces because Stronger forces than Very strong forces.
tween particles
of the large distance in gas. Liquids can Solids have a fixed
between particles
be poured.
volume.
Changes in phase
In general a gas A liquid becomes a Solids become liqbecomes a liquid gas if its tempera- uids or gases if their
or solid when it is ture is increased. It temperature is incooled.
Particles becomes a solid if creased.
have less energy its temperature deand therefore move creases.
closer together so
that the attractive forces become
stronger, and the
gas becomes a liquid
or a solid
Let’s look at an example that involves the three phases of water: ice (solid), water (liquid) and
water vapour (gas).
solid
liquid
gas
Figure 2.7: The three phases of matter
In a solid (e.g. ice), the water molecules have very little energy and can’t move away from each
other. The molecules are held close together in a regular pattern called a lattice. If the ice is
27
2.5
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
heated, the energy of the molecules increases. This means that some of the water molecules are
able to overcome the intermolecular forces that are holding them together, and the molecules
move further apart to form liquid water. This is why liquid water is able to flow, because the
molecules are more free to move than they were in the solid lattice. If the molecules are heated
further, the liquid water will become water vapour, which is a gas. Gas particles have lots of
energy and are far away from each other. That is why it is difficult to keep a gas in a specific
area! The attractive forces between the particles are very weak and they are only loosely held
together. Figure 2.8 shows the changes in phase that may occur in matter, and the names that
describe these processes.
Gas
co n
eva dens
p o a ti o
rat n
io n
on
a ti
lim o n
s u b a ti
re- blim
su
Liquid
Solid
freezing
melting
Figure 2.8: Changes in phase
2.5
The Properties of Matter
Let us now look at what we have learned about chemical bonds, intermolecular forces and the
kinetic theory of matter, and see whether this can help us to understand some of the macroscopic
properties of materials.
1. Melting point
Definition: Melting point
The temperature at which a solid changes its phase or state to become a liquid. The reverse
process (change in phase from liquid to solid) is called freezing.
In order for a solid to melt, the energy of the particles must increase enough to overcome
the bonds that are holding the particles together. It makes sense then that a solid which is
held together by strong bonds will have a higher melting point than one where the bonds
are weak, because more energy (heat) is needed to break the bonds. In the examples we
have looked at, metals, ionic solids and some atomic lattices (e.g. diamond) have high
melting points, whereas the melting points for molecular solids and other atomic lattices
(e.g. graphite) are much lower. Generally, the intermolecular forces between molecular
solids are weaker than those between ionic and metallic solids.
2. Boiling point
Definition: Boiling point
The temperature at which a liquid changes its phase to become a gas.
28
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
2.5
When the temperature of a liquid increases, the average kinetic energy of the particles also
increases, and they are able to overcome the bonding forces that are holding them in the
liquid. When boiling point is reached, evaporation takes place and some particles in the
liquid become a gas. In other words, the energy of the particles is too great for them to
be held in a liquid anymore. The stronger the bonds within a liquid, the higher the boiling
point needs to be in order to break these bonds. Metallic and ionic compounds have high
boiling points while the boiling point for molecular liquids is lower.
The data in table 2.2 below may help you to understand some of the concepts we have
explained. Not all of the substances in the table are solids at room temperature, so for
now, let’s just focus on the boiling points for each of these substances. Of the substances
listed, ethanol has the weakest intermolecular forces, and sodium chloride and mercury
have the strongest. What do you notice?
Substance
Ethanol (C2 H6 O)
Water
Mercury
Sodium chloride
Melting point (0 C)
-114,3
0
-38,83
801
Boiling point (0 C)
78,4
100
356,73
1465
Table 2.2: The melting and boiling points for a number of substances
You will have seen that substances such as ethanol, with relatively weak intermolecular
forces, have the lowest boiling point, while substances with stronger intermolecular forces
such as sodium chloride and mercury, must be heated much more if the particles are to
have enough energy to overcome the forces that are holding them together in the liquid
or solid phase.
Exercise: Forces and boiling point
The table below gives the molecular formula and the boiling point for a
number of organic compounds called alkanes. Refer to the table and then
answer the questions that follow.
Organic compound Molecular formula Boiling point (0 C)
Methane
CH2
-161.6
Ethane
C2 H6
-88.6
Propane
C3 H8
-45
Butane
C4 H10
-0.5
Pentane
C5 H12
36.1
Hexane
C6 H14
69
Heptane
C7 H16
98.42
Octane
C8 H18
125.52
Data from: http://www.wikipedia.com
(a) Draw a graph to show the relationship between the number of carbon atoms
in each alkane, and its boiling point (Number of carbon atoms will go on
the x-axis and boiling point on the y-axis).
(b) Describe what you see.
(c) Suggest a reason for what you have observed.
(d) Why was it enough for us to use ’number of carbon atoms’ as a measure
of the molecular weight of the molecules?
3. Density and viscosity
Density is a measure of the mass of a substance per unit volume. The density of a solid
is generally higher than that of a liquid because the particles are hold much more closely
29
2.5
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
together and therefore there are more particles packed together in a particular volume. In
other words, there is a greater mass of the substance in a particular volume. In general,
density increases as the strength of the intermolecular forces increases. Viscosity is a
measure of how resistant a liquid is to changing its form. Viscosity is also sometimes
described as the ’thickness’ of a fluid. Think for example of syrup and how slowly it pours
from one container into another. Now compare this to how easy it is to pour water. The
viscosity of syrup is greater than the viscosity of water. Once again, the stronger the
intermolecular forces in the liquid, the greater its viscosity.
It should be clear now that we can explain a lot of the macroscopic properties of matter (i.e.
the characteristics we can see or observe) by understanding their microscopic structure and
the way in which the atoms and molecules that make up matter are held together.
Activity :: Investigation : Determining the density of liquids:
Density is a very important property because it helps us to identify different
materials. Every material, depending on the elements that make it up, and the
arrangement of its atoms, will have a different density.
The equation for density is:
Density = Mass/Volume
Discussion questions:
To calculate the density of liquids and solids, we need to be able to first determine
their mass and volume. As a group, think about the following questions:
• How would you determine the mass of a liquid?
• How would you determine the volume of an irregular solid?
Apparatus:
Laboratory mass balance, 10 ml and 100 ml graduated cylinders, thread, distilled
water, two different liquids.
Method:
Determine the density of the distilled water and two liquids as follows:
1. Measure and record the mass of a 10 ml graduated cyclinder.
2. Pour an amount of distilled water into the cylinder.
3. Measure and record the combined mass of the water and cylinder.
4. Record the volume of distilled water in the cylinder
5. Empty, clean and dry the graduated cylinder.
6. Repeat the above steps for the other two liquids you have.
7. Complete the table below.
Liquid
Distilled water
Liquid 1
Liquid 2
Mass (g)
Volume (ml)
Density (g/ml)
Activity :: Investigation : Determining the density of irregular solids:
Apparatus:
Use the same materials and equpiment as before (for the liquids). Also find a
number of solids that have an irregular shape.
Method:
Determine the density of irregular solids as follows:
30
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
2.6
1. Measure and record the mass of one of the irregular solids.
2. Tie a piece of thread around the solid.
3. Pour some water into a 100 ml graduated cylinder and record the volume.
4. Gently lower the solid into the water, keeping hold of the thread. Record the
combined volume of the solid and the water.
5. Dtermine the volume of the solid by subtracting the combined volume from the
original volume of the water only.
6. Repeat these steps for the second object.
7. Complete the table below.
Solid
Solid 1
Solid 2
Solid 3
2.6
Mass (g)
Volume (ml)
Density (g/ml)
Summary
• The smallest unit of matter is the atom. Atoms can combine to form molecules.
• A molecule is a group of two or more atoms that are attracted to each other by chemical
bonds.
• A small molecule consists of a few atoms per molecule. A giant molecule consists of
millions of atoms per molecule, for example metals and diamonds.
• The structure of a molecule can be represented in a number of ways.
• The chemical formula of a molecule is an abbreviated way of showing a molecule, using
the symbols for the elements in the molecule. There are two types of chemical formulae:
molecular and empirical formula.
• The molecular formula of a molecule gives the exact number of atoms of each element
that are in the molecule.
• The empirical formula of a molecule gives the relative number of atoms of each element
in the molecule.
• Molecules can also be represented using diagrams.
• A ball and stick diagram is a 3-dimensional molecular model that uses ’balls’ to represent
atoms and ’sticks’ to represent the bonds between them.
• A space-filling model is also a 3-dimensional molecular model. The atoms are represented
by multi-coloured spheres.
• In a molecule, atoms are held together by chemical bonds or intramolecular forces.
Covalent bonds, ionic bonds and metallic bonds are examples of chemical bonds.
• A covalent bond exists between non-metal atoms. An ionic bond exists between nonmetal and metal atoms, and a metallic bond exists between metal atoms.
• Intermolecular forces are the bonds that hold molecules together.
• The kinetic theory of matter attempts to explain the behaviour of matter in different
phases.
• The theory says that all matter is composed of particles which have a certain amount
of energy which allows them to move at different speeds depending on the temperature
(energy). There are spaces between the particles, and also attractive forces between
particles when they come close together.
31
2.6
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
• Understanding chemical bonds, intermolecular forces and the kinetic theory of matter, can
help to explain many of the macroscopic properties of matter.
• Melting point is the temperature at which a solid changes its phase to become a liquid.
The reverse process (change in phase from liquid to solid) is called freezing. The stronger
the chemical bonds and intermolecular forces in a substance, the higher the melting point
will be.
• Boiling point is the temperature at which a liquid changes phase to become a gas. The
stronger the chemical bonds and intermolecular forces in a substance, the higher the boiling
point will be.
• Density is a measure of the mass of a substance per unit volume.
• Viscosity is a measure of how resistant a liquid is to changing its form.
Exercise: Summary exercise
1. Give one word or term for each of the following descriptions.
(a) The property that determines how easily a liquid flows.
(b) The change in phase from liquid to gas.
(c) A composition of two or more atoms that act as a unit.
(d) Chemical formula that gives the relative number of atoms of each element
that are in a molecule.
2. For each of the following questions, choose the one correct answer from the
list provided.
A Ammonia, an ingredient in household cleaners, can be broken down to
form one part nitrogen (N) and three parts hydrogen (H). This means that
ammonia...
i. is a colourless gas
ii. is not a compound
iii. cannot be an element
iv. has the formula N3 H
B If one substance A has a melting point that is lower than the melting point
of substance B, this suggests that...
i. A will be a liquid at room temperature.
ii. The chemical bonds in substance A are weaker than those in substance
B.
iii. The chemical bonds in substance A are stronger than those in substance B.
iv. B will be a gas at room temperature.
3. Boiling point is an important concept to understand.
a Define ’boiling point’.
b What change in phase takes place when a liquid reaches its boiling point?
c What is the boiling point of water?
d Use the kinetic theory of matter and your knowledge of intermolecular
forces, to explain why water changes phase at this temperature.
4. Refer to the table below which gives the melting and boiling points of a
number of elements, and then answer the questions that follow. (Data from
http://www.chemicalelements.com)
Element
copper
magnesium
oxygen
carbon
helium
sulfur
Melting point
1083
650
-218.4
3500
-272
112.8
32
Boiling point (0 C)
2567
1107
-183
4827
-268.6
444.6
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
a What state of matter (i.e. solid, liquid or gas) will each of these elements
be in at room temperature?
b Which of these elements has the strongest forces between its atoms? Give
a reason for your answer.
c Which of these elements has the weakest forces between its atoms? Give
a reason for your answer.
33
2.6
2.6
CHAPTER 2. WHAT ARE THE OBJECTS AROUND US MADE OF? - GRADE 10
34
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
Chapter 11
Physical and Chemical Change Grade 10
Matter is all around us. The desks we sit at, the air we breathe and the water we drink, are all
examples of matter. But matter doesn’t always stay the same. It can change in many different
ways. In this chapter, we are going to take a closer look at physical and chemical changes that
occur in matter.
11.1
Physical changes in matter
A physical change is one where the particles of the substances that are involved in the change
are not broken up in any way. When water is heated for example, the temperature and energy
of the water molecules increases and the liquid water evaporates to form water vapour. When
this happens, some kind of change has taken place, but the molecular structure of the water has
not changed. This is an example of a physical change.
H2 O(l) → H2 O(g)
Conduction (the transfer of energy through a material) is another example of a physical change.
As energy is transferred from one material to another, the energy of each material is changed,
but not its chemical makeup. Dissolving one substance in another is also a physical change.
Definition: Physical change
A change that can be seen or felt, but that doesn’t involve the break up of the particles in
the reaction. During a physical change, the form of matter may change, but not its identity.
A change in temperature is an example of a physical change.
There are some important things to remember about physical changes in matter:
• Arrangement of particles
When a physical change occurs, the particles (e.g. atoms, molecules) may re-arrange
themselves without actually breaking up in any way. In the example of evaporation that
we used earlier, the water molecules move further apart as their temperature (and therefore
energy) increases. The same would be true if ice were to melt. In the solid phase, water
molecules are packed close together in a very ordered way, but when the ice is heated, the
molecules overcome the forces holding them together and they move apart. Once again,
the particles have re-arranged themselves, but have not broken up.
H2 O(s) → H2 O(l)
211
11.2
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
solid
gas
liquid
Figure 11.1: The arrangement of water molecules in the three phases of matter
Figure 11.1 shows this more clearly. In each phase of water, the water molecule itself stays
the same, but the way the molecules are arranged has changed.
In a physical change, the total mass, the number of atoms and the number of molecules
will always stay the same.
• Energy changes
Energy changes may take place when there is a physical change in matter, but these energy
changes are normally smaller than the energy changes that take place during a chemical
change.
• Reversibility
Physical changes in matter are usually easier to reverse than chemical changes. Water
vapour for example, can be changed back to liquid water if the temperature is lowered.
Liquid water can be changed into ice by simply increasing the temperature, and so on.
11.2
Chemical Changes in Matter
When a chemical change takes place, new substances are formed in a chemical reaction. These
new products may have very different properties from the substances that were there at the start
of the reaction.
The breakdown of copper(II) chloride to form copper and chlorine is an example of chemical
change. A simplified diagram of this reaction is shown in figure 11.2. In this reaction, the initial
substance is copper(II) chloride but, once the reaction is complete, the products are copper and
chlorine.
Cl
Cu
Cl
Cu
+
Cl
Cl
CuCl2 → Cu + Cl2
Figure 11.2: The decomposition of copper(II) chloride to form copper and chlorine
Definition: Chemical change
The formation of new substances in a chemical reaction. One type of matter is changed
into something different.
There are some important things to remember about chemical changes:
• Arrangement of particles
212
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
11.2
During a chemical change, the particles themselves are changed in some way. In the
example of copper (II) chloride that was used earlier, the CuCl2 molecules were split
up into their component atoms. The number of particles will change because each one
CuCl2 molecule breaks down into one copper atom (Cu) and one chlorine molecule (Cl2 ).
However, what you should have noticed, is that the number of atoms of each element
stays the same, as does the total mass of the atoms. This will be discussed in more detail
in a later section.
• Energy changes
The energy changes that take place during a chemical reaction are much greater than those
that take place during a physical change in matter. During a chemical reaction, energy
is used up in order to break bonds, and then energy is released when the new product is
formed. This will be discussed in more detail in section ??.
• Reversibility
Chemical changes are far more difficult to reverse than physical changes.
Two types of chemical reactions are decomposition reactions and synthesis reactions.
11.2.1
Decomposition reactions
A decomposition reaction occurs when a chemical compound is broken down into elements or
smaller compounds. The generalised equation for a decomposition reaction is:
AB → A + B
One example of such a reaction is the decomposition of hydrogen peroxide (figure 11.3) to form
hydrogen and oxygen according to the following equation:
2H2 O2 → 2H2 O + O2
H
H
O
O
H
H
H
O
H
O
H
H
+
O
O
Figure 11.3: The decomposition of H2 O2 to form H2 O and O2
The decomposition of mercury (II) oxide is another example.
Activity :: Experiment : The decomposition of mercury (II) oxide
Aim:
To observe the decomposition of mercury (II) oxide when it is heated.
Note: Because this experiment involves mercury, which is a poisonous substance,
it should be done in a fume cupboard, and all the products of the reaction must be
very carefully disposed of.
Apparatus:
Mercury (II) oxide (an orange-red product); two test tubes; a large beaker; stopper and delivery tube; Bunsen burner; wooden splinter.
213
O
O
11.2
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
delivery
tube
c cb
b
c
b
rubber
stopper
b bc
c
c
b
bc bc
c b
c
cb
b
bc bc
c bc
b
cbc
b
bubbles of
oxygen gas
collecting in
second test
tube
mercury
(II) oxide
water
bunsen
burner
Method:
1. Put a small amount of mercury (II) oxide in a test tube and heat it gently over
a Bunsen burner. Then allow it to cool. What do you notice about the colour
of the mercury (II) oxide?
2. Heat the test tube again, and note what happens. Do you notice anything on
the walls of the test tube? Record these observations.
3. Test for the presence of oxygen using a glowing splinter.
Results:
• During the first heating of mercury (II) oxide, the only change that took place
was a change in colour from orange-red to black and then back to its original
colour.
• When the test tube was heated again, deposits of mercury formed on the inner
surface of the test tube. What colour is this mercury?
• The glowing splinter burst into flame when it was placed in the test tube,
meaning that oxygen is present.
Conclusions:
When mercury is heated, it decomposes to form mercury and oxygen. The
chemical decomposition reaction that takes place can be written as follows:
2HgO → 2Hg + O2
11.2.2
Synthesis reactions
During a synthesis reaction, a new product is formed from smaller elements or compounds.
The generalised equation for a synthesis reaction is as follows:
A + B → AB
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CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
11.2
One example of a synthesis reaction is the burning of magnesium in oxygen to form magnesium
oxide. The equation for the reaction is:
2M g + O2 → 2M gO
Figure 11.4 shows the chemical changes that take place at a microscopic level during this chemical
reaction.
Mg
Mg
+
O
O
Mg
O
Mg
Figure 11.4: The synthesis of magnesium oxide (MgO) from magnesium and oxygen
Activity :: Experiment : Chemical reactions involving iron and sulfur
Aim:
To demonstrate the synthesis of iron sulfide from iron and sulfur.
Apparatus:
5.6 g iron filings and 3.2 g powdered sulfur; porcelain dish; test tube; bunsen
burner
Method:
1. Before you carry out the experiment, write a balanced equation for the reaction
you expect will take place.
2. Measure the quantity of iron and sulfur that you need and mix them in a
porcelain dish.
3. Take some of this mixture and place it in the test tube. The test tube should
be about 1/3 full.
4. This reaction should ideally take place in a fume cupboard. Heat the test tube
containing the mixture over the Bunsen burner. Increase the heat if no reaction
takes place. Once the reaction begins, you will need to remove the test tube
from the flame. Record your observations.
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11.2
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
5. Wait for the product to cool before breaking the test tube with a hammer.
Make sure that the test tube is rolled in paper before you do this, otherwise
the glass will shatter everywhere and you may be hurt.
6. What does the product look like? Does it look anything like the original reactants? Does it have any of the properties of the reactants (e.g. the magnetism
of iron)?
Results:
• After you removed the test tube from the flame, the mixture glowed a bright
red colour. The reaction is exothermic and produces energy.
• The product, iron sulfide, is a dark colour and does not share any of the
properties of the original reactants. It is an entirely new product.
Conclusions:
A synthesis reaction has taken place. The equation for the reaction is:
F e + S → F eS
Activity :: Investigation : Physical or chemical change?
Apparatus:
Bunsen burner, 4 test tubes, a test tube rack and a test tube holder, small
spatula, pipette, magnet, a birthday candle, NaCl (table salt), 0.1M AgNO3 , 6M
HCl, magnesium ribbon, iron filings, sulfur.
Method:
1. Place a small amount of wax from a birthday candle into a test tube and heat
it over the bunsen burner until it melts. Leave it to cool.
2. Add a small spatula of NaCl to 5 ml water in a test tube and shake. Then use
the pipette to add 10 drops of AgNO3 to the sodium chloride solution.
3. Take a 5 cm piece of magnesium ribbon and tear it into 1 cm pieces. Place
two of these pieces into a test tube and add a few drops of 6M HCl. NOTE:
Be very careful when you handle this acid because it can cause major burns.
4. Take about 0.5 g iron filings and 0.5 g sulfur. Test each substance with a
magnet. Mix the two samples in a test tube, and run a magnet alongside the
outside of the test tube.
5. Now heat the test tube that contains the iron and sulfur. What changes do
you see? What happens now, if you run a magnet along the outside of the test
tube?
6. In each of the above cases, record your observations.
Questions:
Decide whether each of the following changes are physical or chemical and give
a reason for your answer in each case. Record your answers in the table below:
Description
Physical
chemical
change
melting candle wax
dissolving NaCl
mixing NaCl with AgNO3
tearing magnesium ribbon
adding HCl to magnesium ribbon
mixing iron and sulfur
heating iron and sulfur
216
or
Reason
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
11.3
11.3
Energy changes in chemical reactions
All reactions involve some change in energy. During a physical change in matter, such as the
evaporation of liquid water to water vapour, the energy of the water molecules increases. However, the change in energy is much smaller than in chemical reactions.
When a chemical reaction occurs, some bonds will break, while new bonds may form. Energy
changes in chemical reactions result from the breaking and forming of bonds. For bonds to
break, energy must be absorbed. When new bonds form, energy will be released because the
new product has a lower energy than the ’inbetween’ stage of the reaction when the bonds in
the reactants have just been broken.
In some reactions, the energy that must be absorbed to break the bonds in the reactants, is less
than the total energy that is released when new bonds are formed. This means that in the overall
reaction, energy is released. This type of reaction is known as an exothermic reaction. In other
reactions, the energy that must be absorbed to break the bonds in the reactants, is more than
the total energy that is released when new bonds are formed. This means that in the overall
reaction, energy must be absorbed from the surroundings. This type of reaction is known as an
endothermic reaction. In the earlier part of this chapter, most decomposition reactions were
endothermic, and heating was needed for the reaction to occur. Most of the synthesis reactions
were exothermic, meaning that energy was given off in the form of heat or light.
More simply, we can describe the energy changes that take place during a chemical reaction as:
Total energy absorbed to break bonds - Total energy released when new bonds form
So, for example, in the reaction...
2M g + O2 → 2M gO
Energy is needed to break the O-O bonds in the oxygen molecule so that new Mg-O bonds can
be formed, and energy is released when the product (MgO) forms.
Despite all the energy changes that seem to take place during reactions, it is important to
remember that energy cannot be created or destroyed. Energy that enters a system will have
come from the surrounding environment, and energy that leaves a system will again become part
of that environment. This principle is known as the principle of conservation of energy.
Definition: Conservation of energy principle
Energy cannot be created or destroyed. It can only be changed from one form to another.
Chemical reactions may produce some very visible, and often violent, changes. An explosion,
for example, is a sudden increase in volume and release of energy when high temperatures are
generated and gases are released. For example, NH4 NO3 can be heated to generate nitrous oxide.
Under these conditions, it is highly sensitive and can detonate easily in an explosive exothermic
reaction.
11.4
Conservation of atoms and mass in reactions
The total mass of all the substances taking part in a chemical reaction is conserved during a
chemical reaction. This is known as the law of conservation of mass. The total number of
atoms of each element also remains the same during a reaction, although these may be arranged
differently in the products.
We will use two of our earlier examples of chemical reactions to demonstrate this:
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11.4
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
• The decomposition of hydrogen peroxide into water and oxygen
2H2 O2 → 2H2 O + O2
H
H
O
H
O
H
H
O
H
O
H
H
+
O
O
O
Left hand side of the equation
Total atomic mass = (4 × 1) + (4 × 16) = 68 u
Number of atoms of each element = (4 × H) + (4 × O)
Right hand side of the equation
Total atomic mass = (4 × 1) + (2 × 16) + (2 × 16) = 68 u
Number of atoms of each element = (4 × H) + (4 × O)
Both the atomic mass and the number of atoms of each element are conserved in the
reaction.
• The synthesis of magnesium and oxygen to form magnesium oxide
2M g + O2 → 2M gO
Mg
Mg
+
O
O
Mg
O
Mg
O
Left hand side of the equation
Total atomic mass = (2 × 24.3) + (2 × 16) = 80.6 u
Number of atoms of each element = (2 × Mg) + (2 × O)
Right hand side of the equation
Total atomic mass = (2 × 24.3) + (2 × 16) = 80.6 u
Number of atoms of each element = (2 × Mg) + (2 × O)
Both the atomic mass and the number of atoms of each element are conserved in the
reaction.
Activity :: Demonstration : The conservation of atoms in chemical reactions
Materials:
• Coloured marbles or small balls to represent atoms. Each colour will represent
a different element.
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CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
11.5
• Prestik
Method:
1. Choose a reaction from any that have been used in this chapter or any other
balanced chemical reaction that you can think of. To help to explain this
activity, we will use the decomposition reaction of calcium carbonate to produce
carbon dioxide and calcium oxide.
CaCO3 → CO2 + CaO
2. Stick marbles together to represent the reactants and put these on one side of
your table. In this example you may for example join one red marble (calcium),
one green marble (carbon) and three yellow marbles (oxygen) together to form
the molecule calcium carbonate (CaCO3 ).
3. Leaving your reactants on the table, use marbles to make the product molecules
and place these on the other side of the table.
4. Now count the number of atoms on each side of the table. What do you notice?
5. Observe whether there is any difference between the molecules in the reactants
and the molecules in the products.
Discussion
You should have noticed that the number of atoms in the reactants is the same
as the number of atoms in the product. The number of atoms is conserved during
the reaction. However, you will also see that the molecules in the reactants and
products is not the same. The arrangement of atoms is not conserved during the
reaction.
11.5
Law of constant composition
In any given chemical compound, the elements always combine in the same proportion with each
other. This is the law of constant proportions.
The law of constant composition says that, in any particular chemical compound, all samples
of that compound will be made up of the same elements in the same proportion or ratio. For
example, any water molecule is always made up of two hydrogen atoms and one oxygen atom in
a 2:1 ratio. If we look at the relative masses of oxygen and hydrogen in a water molecule, we
see that 94% of the mass of a water molecule is accounted for by oxygen, and the remaining 6%
is the mass of hydrogen. This mass proportion will be the same for any water molecule.
This does not mean that hydrogen and oxygen always combine in a 2:1 ratio to form H2 O.
Multiple proportions are possible. For example, hydrogen and oxygen may combine in different proportions to form H2 O2 rather than H2 O. In H2 O2 , the H:O ratio is 1:1 and the mass
ratio of hydrogen to oxygen is 1:16. This will be the same for any molecule of hydrogen peroxide.
11.6
Volume relationships in gases
In a chemical reaction between gases, the relative volumes of the gases in the reaction are present
in a ratio of small whole numbers if all the gases are at the same temperature and pressure. This
relationship is also known as Gay-Lussac’s Law.
For example, in the reaction between hydrogen and oxygen to produce water, two volumes of
H2 react with 1 volume of O2 to produce 2 volumes of H2 O.
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11.7
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
2H2 + O2 → 2H2 O
In the reaction to produce ammonia, one volume of nitrogen gas reacts with three volumes of
hydrogen gas to produce two volumes of ammonia gas.
N2 + 3H2 → 2N H3
This relationship will also be true for all other chemical reactions.
11.7
Summary
• Matter does not stay the same. It may undergo physical or chemical changes
• A physical change means that the form of matter may change, but not its identity. For
example, when water evaporates, the energy and the arrangement of water molecules will
change, but not the structure of the water molecule itself.
• During a physical change, the arrangement of particles may change but the mass, number
of atoms and number of molecules will stay the same.
• Physical changes involve small changes in energy, and are easily reversible.
• A chemical change occurs when one form of matter changes into something else. A
chemical reaction involves the formation of new substances with different properties.
For example, carbon dioxide reacts with water to form carbonic acid.
CO2 + H2 O → H2 CO3
• A chemical change may involve a decomposition or synthesis reaction. During chemical
change, the mass and number of atoms is conserved, but the number of molecules is not
always the same.
• Chemical reactions involve larger changes in energy. During a reaction, energy is needed
to break bonds in the reactants, and energy is released when new products form. If the
energy released is greater than the energy absorbed, then the reaction is exothermic. If the
energy released is less than the energy absorbed, then the reaction is endothermic. These
chemical reactions are not easily reversible.
• Decomposition reactions are usually endothermic and synthesis reactions are usually
exothermic.
• The law of conservation of mass states that the total mass of all the substances taking
part in a chemical reaction is conserved and the number of atoms of each element in the
reaction does not change when a new product is formed.
• The conservation of energy principle states that energy cannot be created or destroyed,
it can only change from one form to another.
• The law of constant composition states that in any particular compound, all samples of
that compound will be made up of the same elements in the same proportion or ratio.
• Gay-Lussac’s Law states that in a chemical reaction between gases, the relative volumes
of the gases in the reaction are present in a ratio of small whole numbers if all the gases
are at the same temperature and pressure.
Exercise: Summary exercise
1. Complete the following table by saying whether each of the descriptions is an
example of a physical or chemical change:
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CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
Description
Physical
chemical
11.7
or
hot and cold water mix together
milk turns sour
a car starts to rust
food digests in the stomach
alcohol disappears when it is placed on your skin
warming food in a microwave
separating sand and gravel
fireworks exploding
2. For each of the following reactions, say whether it is an example of a synthesis
or decomposition reaction:
(a)
(b)
(c)
(d)
(N H4 )2 CO3 → 2N H3 + CO2 + H2 O
4F e + 3O2 → 2F e2 O3
N2 (g) + 3H2 (g) → 2N H3
CaCO3 (s) → CaO + CO2
3. For the following equation:
CaCO3 → CO2 + CaO
Show that the ’law of conservation of mass’ applies.
221
11.7
CHAPTER 11. PHYSICAL AND CHEMICAL CHANGE - GRADE 10
222
Chapter 12
Representing Chemical Change Grade 10
As we have already mentioned, a number of changes can occur when elements react with one
another. These changes may either be physical or chemical. One way of representing these
changes is through balanced chemical equations. A chemical equation describes a chemical
reaction by using symbols for the elements involved. For example, if we look at the reaction
between iron (Fe) and sulfur (S) to form iron sulfide (FeS), we could represent these changes
either in words or using chemical symbols:
iron + sulfur → iron sulfide
or
F e + S → F eS
Another example would be:
ammonia + oxygen → nitric oxide + water
or
4N H3 + 5O2 → 4N O + 6H2 O
Compounds on the left of the arrow are called the reactants and these are needed for the reaction to take place. In this equation, the reactants are ammonia and oxygen. The compounds on
the right are called the products and these are what is formed from the reaction.
In order to be able to write a balanced chemical equation, there are a number of important
things that need to be done:
1. Know the chemical symbols for the elements involved in the reaction
2. Be able to write the chemical formulae for different reactants and products
3. Balance chemical equations by understanding the laws that govern chemical change
4. Know the state symbols for the equation
We will look at each of these steps separately in the next sections.
12.1
Chemical symbols
It is very important to know the chemical symbols for common elements in the Periodic Table
so that you are able to write chemical equations and to recognise different compounds.
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12.2
CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Exercise: Revising common chemical symbols
• Write down the chemical symbols and names of all the elements that you know.
• Compare your list with another learner and add any symbols and names that
you don’t have.
• Spend some time, either in class or at home, learning the symbols for at least the
first twenty elements in the periodic table. You should also learn the symbols
for other common elements that are not in the first twenty.
• Write a short test for someone else in the class and then exchange tests with
them so that you each have the chance to answer one.
12.2
Writing chemical formulae
A chemical formula is a concise way of giving information about the atoms that make up a
particular chemical compound. A chemical formula shows each element by its symbol, and also
shows how many atoms of each element are found in that compound. The number of atoms (if
greater than one) is shown as a subscript.
Examples:
CH4 (methane)
Number of atoms: (1 x carbon) + (4 x hydrogen) = 5 atoms in one methane molecule
H2 SO4 (sulfuric acid)
Number of atoms: (2 x hydrogen) + (1 x sulfur) + (4 x oxygen) = 7 atoms in one molecule of
sulfuric acid
A chemical formula may also give information about how the atoms are arranged in a molecule
if it is written in a particular way. A molecule of ethane, for example, has the chemical formula
C2 H6 . This formula tells us how many atoms of each element are in the molecule, but doesn’t
tell us anything about how these atoms are arranged. In fact, each carbon atom in the ethane
molecule is bonded to three hydrogen atoms. Another way of writing the formula for ethane is
CH3 CH3 . The number of atoms of each element has not changed, but this formula gives us
more information about how the atoms are arranged in relation to each other.
The slightly tricky part of writing chemical formulae comes when you have to work out the ratio
in which the elements combine. For example, you may know that sodium (Na) and chlorine (Cl)
react to form sodium chloride, but how do you know that in each molecule of sodium chloride
there is only one atom of sodium for every one atom of chlorine? It all comes down to the
valency of an atom or group of atoms. Valency is the number of bonds that an element can
form with another element. Working out the chemical formulae of chemical compounds using
their valency, will be covered in chapter 4. For now, we will use formulae that you already know.
12.3
Balancing chemical equations
12.3.1
The law of conservation of mass
In order to balance a chemical equation, it is important to understand the law of conservation
of mass.
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CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
12.3
Definition: The law of conservation of mass
The mass of a closed system of substances will remain constant, regardless of the processes
acting inside the system. Matter can change form, but cannot be created or destroyed. For
any chemical process in a closed system, the mass of the reactants must equal the mass of
the products.
In a chemical equation then, the mass of the reactants must be equal to the mass of the products. In order to make sure that this is the case, the number of atoms of each element in the
reactants must be equal to the number of atoms of those same elements in the products. Some
examples are shown below:
Example 1:
F e + S → F eS
Fe
+
S
Fe
S
Reactants
Atomic mass of reactants = 55.8 u + 32.1 u = 87.9 u
Number of atoms of each element in the reactants: (1 × Fe) and (1 × S)
Products
Atomic mass of product = 55.8 u + 32.1 u = 87.9 u
Number of atoms of each element in the products: (1 × Fe) and (1 × S)
Since the number of atoms of each element is the same in the reactants and in the products, we
say that the equation is balanced.
Example 2:
H2 + O2 → H2 O
H
H
+
O
O
H
O
H
Reactants
Atomic mass of reactants = (1 + 1) + (16 + 16) = 34 u
Number of atoms of each element in the reactants: (2 × H) and (2 × O)
Product
Atomic mass of product = (1 + 1 + 16) = 18 u
Number of atoms of each element in the products: (2 × H) and (1 × O)
Since the total atomic mass of the reactants and the products is not the same, and since there are
more oxygen atoms in the reactants than there are in the product, the equation is not balanced.
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12.3
CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Example 3:
N aOH + HCl → N aCl + H2 O
Na
O
H
+
H
Cl
Na
Cl
+
H
O
H
Reactants
Atomic mass of reactants = (23 + 16 + 1) + (1 + 35.4) = 76.4 u
Number of atoms of each element in the reactants: (1 × Na) + (1 × O) + (2 × H) + (1 × Cl)
Products
Atomic mass of products = (23 + 35.4) + (1 + 1 + 16) = 76.4 u
Number of atoms of each element in the products: (1 × Na) + (1 × O) + (2 × H) + (1 × Cl)
Since the number of atoms of each element is the same in the reactants and in the products, we
say that the equation is balanced.
We now need to find a way to balance those equations that are not balanced so that the number
of atoms of each element in the reactants is the same as that for the products. This can be
done by changing the coefficients of the molecules until the atoms on each side of the arrow
are balanced. You will see later in chapter 13 that these coefficients tell us something about the
mole ratio in which substances react. They also tell us about the volume relationship between
gases in the reactants and products.
Important: Coefficients
Remember that if you put a number in front of a molecule, that number applies to the whole
molecule. For example, if you write 2H2 O, this means that there are 2 molecules of water. In
other words, there are 4 hydrogen atoms and 2 oxygen atoms. If we write 3HCl, this means that
there are 3 molecules of HCl. In other words there are 3 hydrogen atoms and 3 chlorine atoms
in total. In the first example, 2 is the coefficient and in the second example, 3 is the coefficient.
12.3.2
Steps to balance a chemical equation
When balancing a chemical equation, there are a number of steps that need to be followed.
• STEP 1: Identify the reactants and the products in the reaction, and write their chemical
formulae.
• STEP 2: Write the equation by putting the reactants on the left of the arrow, and the
products on the right.
• STEP 3: Count the number of atoms of each element in the reactants and the number of
atoms of each element in the products.
• STEP 4: If the equation is not balanced, change the coefficients of the molecules until the
number of atoms of each element on either side of the equation balance.
• STEP 5: Check that the atoms are in fact balanced.
• STEP 6 (we will look at this a little later): Add any extra details to the equation e.g.
phase.
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CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Worked Example 49: Balancing chemical equations 1
Question: Balance the following equation:
M g + HCl → M gCl2 + H2
Answer
Step 1 : Because the equation has been written for you, you can move
straight on to counting the number of atoms of each element in the reactants
and products
Reactants: Mg = 1 atom; H = 1 atom and Cl = 1 atom
Products: Mg = 1 atom; H = 2 atoms and Cl = 2 atoms
Step 2 : Balance the equation
The equation is not balanced since there are 2 chlorine atoms in the product and only
1 in the reactants. If we add a coefficient of 2 to the HCl to increase the number of
H and Cl atoms in the reactants, the equation will look like this:
M g + 2HCl → M gCl2 + H2
Step 3 : Check that the atoms are balanced
If we count the atoms on each side of the equation, we find the following:
Reactants: Mg = 1; H = 2; Cl = 2
Products: Mg = 1; H = 2; Cl = 2
The equation is balanced. The final equation is:
M g + 2HCl → M gCl2 + H2
Worked Example 50: Balancing chemical equations 2
Question: Balance the following equation:
CH4 + O2 → CO2 + H2 O
Answer
Step 1 : Count the number of atoms of each element in the reactants and
products
Reactants: C = 1; H = 4; O = 2
Products: C = 1; H = 2; O = 3
Step 2 : Balance the equation
If we add a coefficient of 2 to H2 O, then the number of hydrogen atoms in the
reactants will be 4, which is the same as for the reactants. The equation will be:
CH4 + O2 → CO2 + 2H2 O
Step 3 : Check that the atoms balance
Reactants: C = 1; H = 4; O = 2
Products: C = 1; H = 4; O = 4
You will see that, although the number of hydrogen atoms now balances, there are
more oxygen atoms in the products. You now need to repeat the previous step. If
we put a coefficient of 2 in front of O2 , then we will increase the number of oxygen
atoms in the reactants by 2. The new equation is:
CH4 + 2O2 → CO2 + 2H2 O
When we check the number of atoms again, we find that the number of atoms
of each element in the reactants is the same as the number in the products. The
equation is now balanced.
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12.3
12.3
CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Worked Example 51: Balancing chemical equations 3
Question: Nitrogen gas reacts with hydrogen gas to form ammonia. Write a balanced chemical equation for this reaction.
Answer
Step 1 : Identify the reactants and the products, and write their chemical
formulae
The reactants are nitrogen (N2 ) and hydrogen (H2 ), and the product is ammonia
(NH3 ).
Step 2 : Write the equation so that the reactants are on the left and products
on the right of the arrow
The equation is as follows:
N 2 + H2 → N H 3
Step 3 : Count the atoms of each element in the reactants and products
Reactants: N = 2; H = 2
Products: N = 1; H = 3
Step 4 : Balance the equation
In order to balance the number of nitrogen atoms, we could rewrite the equation as:
N2 + H2 → 2N H3
Step 5 : Check that the atoms are balanced
In the above equation, the nitrogen atoms now balance, but the hydrogen atoms
don’t (there are 2 hydrogen atoms in the reactants and 6 in the product). If we put
a coefficient of 3 in front of the hydrogen (H2 ), then the hydrogen atoms and the
nitrogen atoms balance. The final equation is:
N2 + 3H2 → 2N H3
Worked Example 52: Balancing chemical equations 4
Question: In our bodies, sugar (C6 H12 O6 ) reacts with the oxygen we breathe in
to produce carbon dioxide, water and energy. Write the balanced equation for this
reaction.
Answer
Step 1 : Identify the reactants and products in the reaction, and write their
chemical formulae.
Reactants: sugar (C6 H12 O6 ) and oxygen (O2 )
Products: carbon dioxide (CO2 ) and water (H2 O)
Step 2 : Write the equation by putting the reactants on the left of the arrow,
and the products on the right
C6 H12 O6 + O2 → CO2 + H2 O
Step 3 : Count the number of atoms of each element in the reactants and
the number of atoms of each element in the products
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CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Reactants: C=6; H=12; O=8;
Products: C=1; H=2; O=3;
Step 4 : Change the coefficents of the molecules until the number of atoms
of each element on either side of the equation balance.
It is easier to start with carbon as it only appears once on each side. If we add a 6
in front of CO2 , the equation looks like this:
C6 H12 O6 + O2 → 6CO2 + H2 O
Reactants: C=6; H=12; O=8;
Products: C=6; H=2; O=13;
Step 5 : Change the coefficients again to try to balance the equation.
Let’s try to get the number of hydrogens the same this time.
C6 H12 O6 + O2 → 6CO2 + 6H2 O
Reactants: C=6; H=12; O=8;
Products: C=6; H=12; O=18;
Step 6 : Now we just need to balance the oxygen atoms.
C6 H12 O6 + 6O2 → 6CO2 + 6H2 O
Reactants: C=6; H=12; O=18;
Products: C=6; H=12; O=18;
Exercise: Balancing simple chemical equations
Balance the following equations:
1. Hydrogen fuel cells are extremely important in the development of alternative
energy sources. Many of these cells work by reacting hydrogen and oxygen gases
together to form water, a reaction which also produces electricity. Balance the
following equation:
H2 (g) + O2 (g) → H2 O(l)
2. The synthesis of ammonia (NH3 ), made famous by the German chemist Fritz
Haber in the early 20th century, is one of the most important reactions in the
chemical industry. Balance the following equation used to produce ammonia:
N2 (g) + H2 (g) → N H3 (g)
3.
4.
5.
6.
7.
8.
M g + P4 → M g3 P2
Ca + H2 O → Ca(OH)2 + H2
CuCO3 + H2 SO4 → CuSO4 + H2 O + CO2
CaCl2 + N a2 CO3 → CaCO3 + N aCl
C12 H22 O11 + O2 → CO2 + H2 O
Barium chloride reacts with sulphuric acid to produce barium sulphate and
hydrochloric acid.
9. Ethane (C2 H6 ) reacts with oxygen to form carbon dioxide and steam.
10. Ammonium carbonate is often used as a smelling salt. Balance the following
reaction for the decomposition of ammonium carbonate:
(N H4 )2 CO3 (s) → N H3 (aq) + CO2 (g) + H2 O(l)
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12.4
CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
12.4
State symbols and other information
The state (phase) of the compounds can be expressed in the chemical equation. This is done by
placing the correct label on the right hand side of the formula. There are only four labels that
can be used:
1. (g) for gaseous compounds
2. (l) for liquids
3. (s) for solid compounds
4. (aq) for an aqueous (water) solution
Occasionally, a catalyst is added to the reaction. A catalyst is a substance that speeds up the
reaction without undergoing any change to itself. In a chemical equation, this is shown by using
the symbol of the catalyst above the arrow in the equation.
To show that heat was needed for the reaction, a Greek delta (∆) is placed above the arrow in
the same way as the catalyst.
Important: You may remember from chapter 11 that energy cannot be created or destroyed
during a chemical reaction but it may change form. In an exothermic reaction, ∆H is less
than zero, and in an endothermic reaction, ∆H is greater than zero. This value is often
written at the end of a chemical equation.
Worked Example 53: Balancing chemical equations 4
Question: Solid zinc metal reacts with aqueous hydrochloric acid to form an aqueous solution of zinc chloride (ZnCl2 )and hydrogen gas. Write a balanced equation
for this reaction.
Answer
Step 1 : Identify the reactants and products and their chemical formulae
The reactants are zinc (Zn) and hydrochloric acid (HCl). The products are zinc
chloride (ZnCl2 ) and hydrogen (H2 ).
Step 2 : Place the reactants on the left of the equation and the products on
the right hand side of the arrow.
Zn + HCl → ZnCl2 + H2
Step 3 : Balance the equation
You will notice that the zinc atoms balance but the chlorine and hydrogen atoms
don’t. Since there are two chlorine atoms on the right and only one on the left, we
will give HCl a coefficient of 2 so that there will be two chlorine atoms on each side
of the equation.
Zn + 2HCl → ZnCl2 + H2
Step 4 : Check that all the atoms balance
When you look at the equation again, you will see that all the atoms are now balanced.
Step 5 : Ensure all details (e.g. state symbols) are added
In the initial description, you were told that zinc was a metal, hydrochloric acid and
zinc chloride were in aqueous solutions and hydrogen was a gas.
Zn(s) + 2HCl(aq) → ZnCl2 (aq) + H2 (g)
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CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Worked Example 54: Balancing chemical equations 5 (advanced)
Question: Balance the following equation:
(N H4 )2 SO4 + N aOH → N H3 + H2 O + N a2 SO4
In this example, the first two steps are not necessary because the reactants and
products have already been given.
Answer
Step 1 : Balance the equation
With a complex equation, it is always best to start with atoms that appear only once
on each side i.e. Na, N and S atoms. Since the S atoms already balance, we will
start with Na and N atoms. There are two Na atoms on the right and one on the
left. We will add a second Na atom by giving NaOH a coefficient of two. There are
two N atoms on the left and one on the right. To balance the N atoms, NH3 will
be given a coefficient of two. The equation now looks as follows:
(N H4 )2 SO4 + 2N aOH → 2N H3 + H2 O + N a2 SO4
Step 2 : Check that all atoms balance
N, Na and S atoms balance, but O and H atoms do not. There are six O atoms and
ten H atoms on the left, and five O atoms and eight H atoms on the right. We need
to add one O atom and two H atoms on the right to balance the equation. This
is done by adding another H2 O molecule on the right hand side. We now need to
check the equation again:
(N H4 )2 SO4 + 2N aOH → 2N H3 + 2H2 O + N a2 SO4
The equation is now balanced.
Exercise: Balancing more advanced chemical equations
Write balanced equations for each of the following reactions:
1. Al2 O3 (s) + H2 SO4 (aq) → Al2 (SO4 )3 (aq) + 3H2 O(l)
2. M g(OH)2 (aq) + HN O3 (aq) → M g(N O3 )2 (aq) + 2H2 O(l)
3. Lead(ll)nitrate solution reacts with potassium iodide solution.
4. When heated, aluminium reacts with solid copper oxide to produce copper
metal and aluminium oxide (Al2 O3 ).
5. When calcium chloride solution is mixed with silver nitrate solution, a white
precipitate (solid) of silver chloride appears. Calcium nitrate (Ca(NO3 )2 ) is
also produced in the solution.
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12.4
12.5
12.5
CHAPTER 12. REPRESENTING CHEMICAL CHANGE - GRADE 10
Summary
• A chemical equation uses symbols to describe a chemical reaction.
• In a chemical equation, reactants are written on the left hand side of the equation, and
the products on the right. The arrow is used to show the direction of the reaction.
• When representing chemical change, it is important to be able to write the chemical
formula of a compound.
• In any chemical reaction, the law of conservation of mass applies. This means that
the total atomic mass of the reactants must be the same as the total atomic mass of the
products. This also means that the number of atoms of each element in the reactants
must be the same as the number of atoms of each element in the product.
• If the number of atoms of each element in the reactants is the same as the number of
atoms of each element in the product, then the equation is balanced.
• If the number of atoms of each element in the reactants is not the same as the number of
atoms of each element in the product, then the equation is not balanced.
• In order to balance an equation, coefficients can be placed in front of the reactants and
products until the number of atoms of each element is the same on both sides of the
equation.
Exercise: Summary exercise
Balance each of the following chemical equations:
1. N H4 + H2 O → N H4 OH
2. Sodium chloride and water react to form sodium hydroxide, chlorine and hydrogen.
3. Propane is a fuel that is commonly used as a heat source for engines and homes.
Balance the following equation for the combustion of propane:
C3 H8 (l) + O2 (g) → CO2 (g) + H2 O(l)
4. Aspartame, an artificial sweetener, has the formula C14 H18 N2 O5 . Write the
balanced equation for its combustion (reaction with O2 ) to form CO2 gas,
liquid H2 O, and N2 gas.
5. F e2 (SO4 )3 + K(SCN ) → K3 F e(SCN )6 + K2 SO4
6. Chemical weapons were banned by the Geneva Protocol in 1925. According
to this protocol, all chemicals that release suffocating and poisonous gases
are not to be used as weapons. White phosphorus, a very reactive allotrope
of phosphorus, was recently used during a military attack. Phosphorus burns
vigorously in oxygen. Many people got severe burns and some died as a result.
The equation for this spontaneous reaction is:
P4 (s) + O2 (g) → P2 O5 (s)
(a) Balance the chemical equation.
(b) Prove that the law of conservation of mass is obeyed during this chemical
reaction.
(c) Name the product formed during this reaction.
(d) Classify the reaction as endothermic or exothermic. Give a reason for your
answer.
(e) Classify the reaction as a sythesis or decomposition reaction. Give a reason
for your answer.
(DoE Exemplar Paper 2 2007)
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Chapter 18
The Water Cycle - Grade 10
18.1
Introduction
You may have heard the word ’cycle’ many times before. Think for example of the word ’bicycle’
or the regular ’cycle tests’ that you may have at school. A cycle is a series of events that repeats
itself. In the case of a bicycle, the wheel turns through a full circle before beginning the motion
again, while cycle tests happen regularly, normally every week or every two weeks. Because a
cycle repeats itself, it doesn’t have a beginning or an end.
Our Earth is a closed system. This means that it can exchange energy with its surroundings
(i.e. the rest of the solar system), but no new matter is brought into the system. For this reason,
it is important that all the elements and molecules on Earth are recycled so that they are never
completely used up. In the next two sections, we are going to take a closer look at two cycles
that are very important for life on Earth. They are the water cycle and the nitrogen cycle.
18.2
The importance of water
For many people, it is so easy to take water for granted, and yet life on Earth would not exist
were it not for this extraordinary compound. Not only is it believed that the first forms of life
actually started in water, but most of the cells in living organisms contain between 70% and
95% water. Here in the cells, water acts as a solvent and helps to transport vital materials such
as food and oxygen to where they are needed, and also removes waste products such as carbon
dioxide and ammonia from the body. For many animals and plants, water is their home. Think
for example of fish and amphibians that live either all or part of the time in rivers, dams and the
oceans. In other words, if water did not exist, no life would be possible.
Apart from allowing life to exist, water also has a number of other functions. Water shapes the
landscape around us by wearing away at rocks and also transports and deposits sediments on
floodplains and along coastal regions. Water also plays a very important role in helping to regulate Earth’s climate. We will discuss this again later in the chapter. As humans we use water in
our homes, in industry, in mining, irrigation and even as a source of electricitiy in hydro-electric
schemes. In fact, if we were able to view Earth from space, we would see that almost three
quarters of our planet’s surface is covered in water. It is because of this that Earth is sometimes
called the ’Blue Planet’. Most of this water is stored in the oceans, with the rest found in ice
(e.g. glaciers), groundwater (e.g. boreholes), surface water (e.g. rivers, lakes, estuaries, dams)
and in the atmosphere as clouds and water vapour.
teresting In the search for life on other planets, one of the first things that scientists look
Interesting
Fact
Fact
for is water. However, most planets are either too close to the sun (and therefore
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18.3
CHAPTER 18. THE WATER CYCLE - GRADE 10
too hot) for water to exist in liquid form, or they are too far away and therefore
too cold. So, even if water were to be found, the conditions are unlikely to allow
it to exist in a form that can support the diversity of life that we see on Earth.
18.3
The movement of water through the water cycle
The water cycle is the continuous movement of water over, above, and beneath the Earth’s
surface. As water moves, it changes phase between liquid (water), solid (ice) and gas (water
vapour). It is powered by solar energy and, because it is a cycle, it has no beginning or end.
Definition: The Water Cycle
The water cycle is the continuous circulation of water across the Earth. The water cycle is
driven by solar radiation and it includes the atmosphere, land, surface water and groundwater. As water moves through the cycle, it changes state between liquid, solid, and gas
phases. The actual movement of water from one part of the cycle to another (e.g. from
river to ocean) is the result of processes such as evaporation, precipitation, infiltration and
runoff.
The movement of water through the water cycle is shown in figure 18.1. In the figure, each
process within this cycle is numbered. Each process will be described below.
1. The source of energy
The water cycle is driven by the sun, which provides the heat energy that is needed for
many of the other processes to take place.
2. Evaporation
When water on the earth’s surface is heated by the sun, the average energy of the water
molecules increases and some of the molecules are able to leave the liquid phase and
become water vapour. This is called evaporation. Evaporation is the change of water from
a liquid to a gas as it moves from the ground, or from bodies of water like the ocean,
rivers and dams, into the atmosphere.
3. Transpiration
Transpiration is the evaporation of water from the aerial parts of plants, especially the
leaves but also from the stems, flowers and fruits. This is another way that liquid water
can enter the atmosphere as a gas.
4. Condensation
When evaporation takes place, water vapour rises in the atmosphere and cools as the
altitude (height above the ground) increases. As the temperature drops, the energy of the
water vapour molecules also decreases, until the molecules don’t have enough energy to
stay in the gas phase. At this point, condensation occurs. Condensation is the change of
water from water vapour (gas) into liquid water droplets in the air. Clouds, fog and mist
are all examples of condensation. A cloud is actually a collection of lots and lots of tiny
water droplets. This mostly takes place in the upper atmosphere but can also take place
close to the ground if there is a significant temperature change.
teresting Have you ever tried breathing out on a very cold day? It looks as though
Interesting
Fact
Fact
you are breathing out smoke! The moist air that you breathe out is much
warmer than the air outside your body. As this warm, moist air comes into
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.3
1
SUN
Condensation forms
clouds
4
Rain falls onto the
soil, flows into the
rivers or seeps into
the soil
5
Rain falls
directly into
rivers, dams
or the
oceans
7
2
Surface water
Evaporation
3
Some rain seeps into
the soil and becomes
part of the ground water
supply
6
Ground water may feed into
rivers or will eventually lead
into the sea
Figure 18.1: The water cycle
contact with the colder air outside, its temperature drops very quickly and
the water vapour in the air you breathe out condenses. The ’smoke’ that
you see is actually formed in much the same way as clouds form in the upper
atmosphere.
5. Precipitation
Precipitation occurs when water falls back to the earth’s surface in the form of rain or
snow. Rain will fall as soon as a cloud becomes too saturated with water droplets. Snow is
similar to rain, except that it is frozen. Snow only falls if temperatures in the atmosphere
are around freezing. The freeing point of water is 00 C).
6. Infiltration
If precipitation occurs, some of this water will filter into the soil and collect underground.
This is called infiltration. This water may evaporate again from the soil at a later stage,
or the underground water may seep into another water body.
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CHAPTER 18. THE WATER CYCLE - GRADE 10
7. Surface runoff
This refers to the many ways that water moves across the land. This includes surface runoff
such as when water flows along a road and into a drain, or when water flows straight across
the sand. It also includes channel runoff when water flows in rivers and streams. As it
flows, the water may infiltrate into the ground, evaporate into the air, become stored in
lakes or reservoirs, or be extracted for agricultural or other human uses.
Important: It is important to realise that the water cycle is all about energy exchanges.
The sun is the original energy source. Energy from the sun heats the water and causes
evaporation. This energy is stored in water vapour as latent heat. When the water vapour
condenses again, the latent heat is released, and helps to drive circulation in the atmosphere.
The liquid water falls to earth, and will evaporate again at a later stage. The atmospheric
circulation patterns that occur because of these exchanges of heat are very important in
influencing climate patterns.
Activity :: Experiment : The Water Cycle
Materials:
Tile or piece of plastic (e.g. lid of ice-cream container) to make a hill slope; glass
fish tank with a lid; beaker with ice cubes; lamp; water
Set up a model of the water cycle as follows:
lamp
ice cubes
glass tank
slope
water
1. Lean the plastic against one side so that it creates a ’hill slope’ as shown in the
diagram.
2. Pour water into the bottom of the tank until about a quarter of the hill slope
is covered.
3. Close the fish tank lid.
4. Place the beaker with ice on the lid directly above the hill slope.
5. Turn the lamp on and position it so that it shines over the water.
6. Leave the model like this for 20-30 minutes and then observe what happens.
Make sure that you don’t touch the lamp as it will be very hot!
Observation questions:
1. Which parts of the water cycle can you see taking place in the model?
2. Which parts of the water cycle are not represented in the model?
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.4
3. Can you think of how those parts that are not shown could be represented?
4. What is the energy source in the model? What would the energy source be in
reality?
5. What do you think the function of the ice is in the beaker?
18.4
The microscopic structure of water
In many ways, water behaves very differently from other liquids. These properties are directly
related to the microscopic structure of water, and more specifically to the shape of the molecule
and its polar nature, and to the bonds that hold water molecules together.
18.4.1
The polar nature of water
Every water molecule is made up of one oxygen atom that is bonded to two hydrogen atoms.
When atoms bond, the nucleus of each atom has an attractive force on the electrons of the other
atoms. This ’pull’ is stronger in some atoms than in others and is called the electronegativity of
the atom. In a water molecule, the oxygen atom has a higher electronegativty than the hydrogen
atoms and therefore attracts the electrons more strongly. The result is that the oxygen atom
has a slightly negative charge and the two hydrogen atoms each have a slightly positive charge.
The water molecule is said to be polar because the electrical charge is not evenly distributed
in the molecule. One part of the molecule has a different charge to other parts. You will learn
more about this in chapter 4.
Oxygen
(slightly negative charge)
Hydrogen
O
(slightly positive charge)
H
H
Hydrogen
(slightly positive charge)
Figure 18.2: Diagrams showing the structure of a water molecule. Each molecule is made up of
two hydrogen atoms that are attached to one oxygen atom.
18.4.2
Hydrogen bonding in water molecules
In every water molecule, the forces that hold the individual atoms together are called intramolecular forces. But there are also forces between different water molecules. These are
called intermolecular forces (figure 18.3). You will learn more about these at a later stage, but
for now it is enough to know that in water, molecules are held together by hydrogen bonds.
Hydrogen bonds are a much stronger type of intermolecular force than those found in many
other substances, and this affects the properties of water.
Important: Intramolecular and intermolecular forces
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18.5
CHAPTER 18. THE WATER CYCLE - GRADE 10
If you find these terms confusing, remember that ’intra’ means within (i.e. the forces within
a molecule). An introvert is someone who doesn’t express emotions and feelings outwardly.
They tend to be quieter and keep to themselves. ’Inter’ means between (i.e. the forces between
molecules). An international cricket match is a match between two different countries.
intermolecular forces
H
intramolecular forces
O
O
H
O
O
H
O
O
Figure 18.3: Intermolecular and intramolecular forces in water. Note that the diagram on the
left only shows intermolecular forces. The intramolecular forces are between the atoms of each
water molecule.
18.5
The unique properties of water
Because of its polar nature and the strong hydrogen bonds between its molecules, water has
some special properties that are quite different to those of other substances.
1. Absorption of infra-red radiation
The polar nature of the water molecule means that it is able to absorb infra-red radiation
(heat) from the sun. As a result of this, the oceans and other water bodies act as heat
reservoirs, and are able to help moderate the Earth’s climate.
2. Specific heat
Definition: Specific heat
Specific heat is the amount of heat energy that is needed to increase the temperature of a
substance by one degree.
Water has a high specific heat, meaning that a lot of energy must be absorbed by water
before its temperature changes.
Activity :: Demonstration : The high specific heat of water
(a) Pour about 100 ml of water into a glass beaker.
(b) Place the beaker on a stand and heat it over a bunsen burner for about 2
minutes.
(c) After this time, carefully touch the side of the beaker (Make sure you touch
the glass very lightly because it will be very hot and may burn you!). Then
use the end of a finger to test the temperature of the water.
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.5
What do you notice? Which of the two (glass or water) is the hottest?
You have probably observed this phenomenon if you have boiled water in a pot on the
stove. The metal of the pot heats up very quickly, and can burn your fingers if you touch
it, while the water may take several minutes before its temperature increases even slightly.
How can we explain this in terms of hydrogen bonding? Remember that increasing the
temperature of a substance means that its particles will move more quickly. However,
before they can move faster, the bonds between them must be broken. In the case of
water, these bonds are strong hydrogen bonds, and so a lot of energy is needed just to
break these, before the particles can start moving faster.
It is the high specific heat of water and its ability to absorb infra-red radiation that allows
it to regulate climate. Have you noticed how places that are closer to the sea have less
extreme daily temperatures than those that are inland? During the day, the oceans heat
up slowly, and so the air moving from the oceans across land is cool. Land temperatures
are cooler than they would be if they were further from the sea. At night, the oceans lose
the heat that they have absorbed very slowly, and so sea breezes blowing across the land
are relatively warm. This means that at night, coastal regions are generally slightly warmer
than areas that are further from the sea.
By contrast, places further from the sea experience higher maximum temperatures, and
lower minimum temperatures. In other words, their temperature range is higher than that
for coastal regions. The same principle also applies on a global scale. The large amount of
water across Earth’s surface helps to regulate temperatures by storing infra-red radiation
(heat) from the sun, and then releasing it very slowly so that it never becomes too hot or
too cold, and life is able to exist comfortably. In a similar way, water also helps to keep
the temperature of the internal environment of living organisms relatively constant. This
is very important. In humans, for example, a change in body temperature of only a few
degrees can be deadly.
3. Melting point and boiling point
The melting point of water is 00 C and its boiling point is 1000C. This large difference
between the melting and boiling point is very important because it means that water can
exist as a liquid over a large range of temperatures. The three phases of water are shown
in figure 18.4.
4. High heat of vaporisation
Definition: Heat of vaporisation
Heat of vaporisation is the energy that is needed to change a given quantity of a substance
into a gas.
The strength of the hydrogen bonds between water molecules also means that it has a
high heat of vaporisation. ’Heat of vaporisation’ is the heat energy that is needed to
change water from the liquid to the gas phase. Because the bonds between molecules are
strong, water has to be heated to 1000 C before it changes phase. At this temperature,
the molecules have enough energy to break the bonds that hold the molecules together.
The heat of vaporisation for water is 40.65 kJ/mol. It is very lucky for life on earth that
water does have a high heat of vaporisation. Can you imagine what a problem it would
be if water’s heat of vaporisation was much lower? All the water that makes up the cells
in our bodies would evaporate and most of the water on earth would no longer be able to
exist as a liquid!
5. Less dense solid phase
Another unusual property of water is that its solid phase (ice) is less dense than its liquid
phase. You can observe this if you put ice into a glass of water. The ice doesn’t sink to
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18.5
CHAPTER 18. THE WATER CYCLE - GRADE 10
Gas (water vapour)
co n
eva dens
p o a ti o
rat n
io n
on
a ti
lim o n
s u b a ti
re- blim
su
Liquid
Solid (ice)
freezing
melting
Figure 18.4: Changes in phase of water
the bottom of the glass, but floats on top of the liquid. This phenomenon is also related to
the hydrogen bonds between water molecules. While other materials contract when they
solidify, water expands. The ability of ice to float as it solidifies is a very important factor
in the environment. If ice sank, then eventually all ponds, lakes, and even the oceans would
freeze solid as soon as temperatures dropped below freezing, making life as we know it
impossible on Earth. During summer, only the upper few inches of the ocean would thaw.
Instead, when a deep body of water cools, the floating ice insulates the liquid water below,
preventing it from freezing and allowing life to exist under the frozen surface.
Figure 18.5: Ice cubes floating in water
teresting Antarctica, the ’frozen continent’, has one of the world’s largest and deepest
Interesting
Fact
Fact
freshwater lakes. And this lake is hidden beneath 4 kilometres of ice! Lake
Vostok is 200 km long and 50 km wide. The thick, glacial blanket of ice acts
as an insulator, preventing the water from freezing.
6. Water as a solvent
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.6
Water is also a very good solvent, meaning that it is easy for other substances to dissolve
in it. It is very seldom, in fact, that we find pure water. Most of the time, the water that
we drink and use has all kinds of substances dissolved in it. It is these that make water
taste different in different areas. So why, then, is it important that water is such a good
solvent? We will look at just a few examples.
• Firstly, think about the animals and plants that live in aquatic environments such
as rivers, dams or in the sea. All of these living organisms either need oxygen for
respiration or carbon dioxide for photosynthesis, or both. How do they get these
gases from the water in which they live? Oxygen and carbon dioxide are just two
of the substances that dissolve easily in water, and this is how plants and animals
obtain the gases that they need to survive. Instead of being available as gases in the
atmosphere, they are present in solution in the surrounding water.
• Secondly, consider the fact that all plants need nitrogen to grow, and that they absorb
this nitrogen from compounds such as nitrates and nitrates that are present in the
soil. The question remains, however, as to how these nitrates and nitrites are able to
be present in the soil at all, when most of the Earth’s nitrogen is in a gaseous form
in the atmosphere. Part of the answer lies in the fact that nitrogen oxides, which
are formed during flashes of lightning, can be dissolved in rainwater and transported
into the soil in this way, to be absorbed by plants. The other part of the answer lies
in the activities of nitrogen-fixing bacteria in the soil, but this is a topic that we will
return to in a later section.
It should be clear now, that water is an amazing compound, and that without its unique properties, life on Earth would definitely not be possible.
Exercise: The properties of water
1. A learner returns home from school on a hot afternoon. In order to get cold
water to drink, she adds ice cubes to a glass of water. She makes the following
observations:
• The ice cubes float in the water.
• After a while the water becomes cold and the ice cubes melt.
(a) What property of ice cubes allows them to float in the water?
(b) Briefly explain why the water gets cold when the ice cubes melt.
(c) Briefly describe how the property you mentioned earlier affects the survival
of aquatic life during winter.
2. Which properties of water allow it to remain in its liquid phase over a large
temperature range? Explain why this is important for life on earth.
18.6
Water conservation
Water is a very precious substance and yet far too often, earth’s water resources are abused and
taken for granted. How many times have you walked past polluted rivers and streams, or seen
the flow of water in a river reduced to almost nothing because of its extraction for industrial and
other uses? And if you were able to test the quality of the water you see, you would probably
be shocked. Often our water resources are contaminated with chemicals such as pesticides and
fertilisers. If water is to continue playing all the important functions that were discussed earlier,
it is vital that we reduce the impact of humans on these resources.
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CHAPTER 18. THE WATER CYCLE - GRADE 10
Activity :: Group work : Human impacts on the water cycle
Read the following extract from an article, entitled ’The Effects of Urbanisation
on the Water Cycle’ by Susan Donaldson, and then answer the questions that follow.
As our communities grow, we notice many visible changes including housing developments, road networks, expansion of services and more. These
changes have an impact on our precious water resources, with pollution
of water being one of many such impacts. To understand these impacts
you will need to have a good knowledge of the water cycle!
It is interesting to note that the oceans contain most of earth’s water
(about 97%). Of the freshwater supplies on earth, 78% is tied up in polar
ice caps and snow, leaving only a very small fraction available for use by
humans. Of the available fresh water, 98% is present as groundwater,
while the remaining 2% is in the form of surface water. Because our
usable water supply is so limited, it is vitally important to protect water
quality. Within the water cycle, there is no ’new’ water ever produced on
the earth. The water we use today has been in existence for billions of
years. The water cycle continually renews and refreshes this finite water
supply.
So how exactly does urbanisation affect the water cycle? The increase
in hard surfaces (e.g. roads, roofs, parking lots) decreases the amount
of water that can soak into the ground. This increases the amount of
surface runoff. The runoff water will collect many of the pollutants that
have accumulated on these surfaces (e.g. oil from cars) and carry them
into other water bodies such as rivers or the ocean. Because there is less
infiltration, peak flows of stormwater runoff are larger and arrive earlier,
increasing the size of urban floods. If groundwater supplies are reduced
enough, this may affect stream flows during dry weather periods because
it is the groundwater that seeps to the surface at these times.
Atmospheric pollution can also have an impact because condensing water
vapour will pick up these pollutants (e.g. SO2 , CO2 and NO2 ) and return
them to earth into other water bodies. However, while the effects of
urbanisation on water quality can be major, these impacts can be reduced
if wise decisions are made during the process of development.
Questions
1. In groups, try to explain...
(a) what is meant by ’urbanisation’
(b) how urbanisation can affect water quality
2. Explain why it is so important to preserve the quality of our water supplies.
3. The article gives some examples of human impacts on water quality. In what
other ways do human activities affect water quality?
4. What do you think some of the consequences of these impacts might be for
humans and other forms of life?
5. Imagine that you are the city manager in your own city or the city closest to
you. What changes would you introduce to try to protect the quality of water
resources in your urban area?
6. What measures could be introduced in rural areas to protect water quality?
Apart from the pollution of water resources, the overuse of water is also a problem. In looking
at the water cycle, it is easy sometimes to think that water is a never-ending resource. In a sense
this is true because water cannot be destroyed. However, the availability of water may vary from
place to place. In South Africa for example, many regions are extremely dry and receive very
little rainfall. The same is true for many other parts of the world, where the scarcity of water
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.6
is a life and death issue. The present threat of global warming is also likely to affect water
resources. Some climate models suggest that rising temperatures could increase the variability
of climate and decrease rainfall in South Africa. With this in mind, and remembering that South
Africa is already a dry country, it is vitally important that we manage our water use carefully. In
addition to this, the less water there is available, the more likely it is that water quality will also
decrease. A decrease in water quality limits how water can be used and developed.
At present, the demands being placed on South Africa’s water resources are large. Table 18.1
shows the water requirements that were predicted for the year 2000. The figures in the table
were taken from South Africa’s National Water Resource Strategy, produced by the Department
of Water Affairs and Forestry in 2004. In the table, ’rural’ means water for domestic use and
stock watering in rural areas, while ’urban’ means water for domestic, industrial and commercial
use in the urban area. ’Afforestation’ is included because many plantations reduce stream flow
because of the large amounts of water they need to survive.
Table 18.1: The predicted water requirements for various water management areas in South
Africa for 2000 (million m3 /annum)
Water management
area
Irrigation
Urban
Rural
Limpopo
Thukela
Upper Vaal
Upper Orange
Breede
Country total
238
204
114
780
577
7920
34
52
635
126
39
2897
28
31
43
60
11
574
Mining
and bulk
industrial
14
46
173
2
0
755
Power
generation
7
1
80
0
0
297
Afforestation Total
1
0
0
0
6
428
Activity :: Case Study : South Africa’s water requirements
Refer to table 18.1 and then answer the following questions:
1. Which water management area in South Africa has the highest need for water...
(a)
(b)
(c)
(d)
in the mining and industry sector?
for power generation?
in the irrigation sector?
Suggest reasons for each of your answers above.
2. For South Africa as a whole...
(a) Which activity uses the most water?
(b) Which activity uses the least water?
3. Complete the following table, by calculating the percentage (%) that each
activity contributes to the total water requirements in South Africa for the year
2000.
Water use activity
Irrigation
Urban
Rural
Mining and bulk industry
Power generation
Afforestation
% of SA’s total water requirements
365
322
334
1045
968
633
12871
18.7
CHAPTER 18. THE WATER CYCLE - GRADE 10
Table 18.2: The available water yield in South Africa in 2000 for various water management
areas (million m3 /annum)
Water management Surface
Ground Irrigation Urban
Mining Total loarea
water
and
cal yield
bulk
industrial
Limpopo
160
98
8
15
0
281
Thukela
666
15
23
24
9
737
Upper Vaal
598
32
11
343
146
1130
Upper Orange
4311
65
34
37
0
4447
Breede
687
109
54
16
0
866
Country total
10240
1088
675
970
254
13227
Now look at table 18.2, which shows the amount of water available in South Africa during 2000.
In the table, ’usable return flow’ means the amount of water that can be reused after it has been
used for irrigation, urban or mining.
Activity :: Case Study : Water conservation
Refer to table 18.2 and then answer the following questions:
1. Explain what is meant by...
(a) surface water
(b) ground water
2. Which water management area has the...
(a)
(b)
(c)
(d)
lowest surface water yield?
highest surface water yield?
lowest total yield?
highest total yield?
3. Look at the country’s total water requirements for 2000 and the total available
yield.
(a) Calculate what percentage of the country’s water yield is already being
used up.
(b) Do you think that the country’s total water requirements will increase or
decrease in the coming years? Give a reason for your answer.
4. South Africa is already placing a huge strain on existing water resources. In
groups of 3-4, discuss ways that the country’s demand for water could be
reduced. Present your ideas to the rest of the class for discussion.
18.7
Summary
• Water is critical for the survival of life on Earth. It is an important part of the cells of
living organisms and is used by humans in homes, industry, mining and agriculture.
• Water moves between the land and sky in the water cycle. The water cycle describes
the changes in phase that take place in water as it circulates across the Earth. The water
cycle is driven by solar radiation.
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CHAPTER 18. THE WATER CYCLE - GRADE 10
18.7
• Some of the important processes that form part of the water cycle are evaporation, transpiration, condensation, precipitation, infiltration and surface runoff. Together these processes
ensure that water is cycled between the land and sky.
• It is the microscopic structure of water that determines its unique properties.
• Water molecules are polar and are held together by hydrogen bonds. These characteristics
affect the properties of water.
• Some of the unique properties of water include its ability to absorb infra-red radiation, its
high specific heat, high heat of vaporisation and the fact that the solid phase of water is
less dense that its liquid phase.
• These properties of water help it to sustain life on Earth by moderating climate, regulating
the internal environment of living organisms and allowing liquid water to exist below ice,
even if temperatures are below zero.
• Water is also a good solvent. This property means that it is a good transport medium
in the cells of living organisms, and that it can dissolve gases and other compounds that
may be needed by aquatic plants and animals.
• Human activities threaten the quality of water resources through pollution and altered
runoff patterns.
• As human populations grow, there is a greater demand for water. In many areas, this
demand exceeds the amount of water available for use. Managing water wisely is important
in ensuring that there will always be water available both for human use, and to maintain
natural ecosystems.
Exercise: Summary Exercise
1. Give a word or term for each of the following phrases:
(a)
(b)
(c)
(d)
The
The
The
The
continuous circulation of water across the earth.
change in phase of water from gas to liquid.
movement of water across a land surface.
temperature at which water changes from liquid to gas.
2. In each of the following multiple choice questions, choose the one correct answer
from the list provided.
(a) Many of the unique properties of water (e.g. its high specific heat and
high boiling point) are due to:
i. strong covalent bonds between the hydrogen and oxygen atoms in each
water molecule
ii. the equal distribution of charge in a water molecule
iii. strong hydrogen bonds between water molecules
iv. the linear arrangement of atoms in a water molecule
(b) Which of the following statements is false?
i. Most of the water on earth is in the oceans.
ii. The hardening of surfaces in urban areas results in increased surface
runoff.
iii. Water conservation is important because water cannot be recycled.
iv. Irrigation is one of the largest water users in South Africa.
3. The sketch below shows a process that leads to rainfall in town X. The town
has been relying only on rainfall for its water supply because it has no access
to rivers or tap water. A group of people told the community that they will
never run out of rainwater because it will never stop raining.
367
18.7
CHAPTER 18. THE WATER CYCLE - GRADE 10
Cloud
P2
P1
Town X
Sea
(a) List the processes labelled P1 and P2 that lead to rainfall in town X.
(b) Is this group of people correct in saying that town X will never run out of
rainwater? Justify your answer using the sketch.
Recently, the amount of rainwater has decreased significantly. Various
reasons have been given to explain the drought. Some of the community
members are blaming this group who told them that it will never stop
raining.
(c) What scientific arguments can you use to convince the community members that this group of people should not be blamed for the drought?
(d) What possible strategies can the community leaders adopt to ensure that
they have a regular supply of water.
368
Chapter 19
Global Cycles: The Nitrogen Cycle
- Grade 10
19.1
Introduction
The earth’s atmosphere is made up of about 78% nitrogen, making it the largest pool of this
gas. Nitrogen is essential for many biological processes. It is in all amino acids, proteins and
nucleic acids. As you will see in a later chapter, these compounds are needed to build tissues,
transport substances around the body, and control what happens in living organisms. In plants,
much of the nitrogen is used in chlorophyll molecules which are needed for photosynthesis and
growth.
So, if nitrogen is so essential for life, how does it go from being a gas in the atmosphere to being
part of living organisms such as plants and animals? The problem with nitrogen is that it is an
’inert’ gas, which means that it is unavailable to living organisms in its gaseous form. This is
because of the strong triple bond between its atoms that makes it difficult to break. Something
needs to happen to the nitrogen gas to change it into a form that it can be used. And at some
later stage, these new compounds must be converted back into nitrogen gas so that the amount
of nitrogen in the atmosphere stays the same. This process of changing nitrogen into different
forms is called the nitrogen cycle (figure 19.1).
Definition: The nitrogen cycle
The nitrogen cycle is a biogeochemical cycle that describes how nitrogen and nitrogencontaining compounds are changed in nature.
Very broadly, the nitrogen cycle is made up of the following processes:
• Nitrogen fixation - The process of converting inert nitrogen gas into more useable nitrogen
compounds such as ammonia.
• Nitrification - The conversion of ammonia into nitrites and then into nitrates, which can
be absorbed and used by plants.
• Denitrification - The conversion of nitrates back into nitrogen gas in the atmosphere.
We are going to look at each of these processes in more detail.
19.2
Nitrogen fixation
Nitrogen fixation is needed to change gaseous nitrogen into forms such as ammonia that are more
useful to living organisms. Some fixation occurs in lightning strikes and in industrial processes,
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19.2
CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
Nitrogen in the Atmosphere
Industrial
fixation
Nitrogen fixation
by bacteria
Animals obtain
nitrates from plants
Decomposers
e.g. bacteria
Ammonia
(NH3-)
Nitrification by
nitrifying bacteria
Plant consumption
Atmosphere
Soil
Denitrification returns nitrogen to the atmosphere
Lightning
fixation
Nitrites
(NO-2)
Nitrification
Nitrates
(NO-3)
Figure 19.1: A simplified diagram of the nitrogen cycle
but most fixation is done by different types of bacteria living either in the soil or in parts of the
plants.
1. Biological fixation
Some bacteria are able to fix nitrogen. They use an enzyme called nitrogenase to combine
gaseous nitrogen with hydrogen to form ammonia. The bacteria then use some of this
ammonia to produce their own organic compounds, while what is left of the ammonia
becomes available in the soil.
Some of these bacteria are free-living, in other words they live in the soil. Others live in
the root nodules of legumes (e.g. soy, peas and beans). Here they form a mutualistic
relationship with the plant. The bacteria get carbohydrates (food) from the plant and,
in exchange, produce ammonia which can be converted into nitrogen compounds that are
essential for the survival of the plant. In nutrient-poor soils, planting lots of legumes can
help to enrich the soil with nitrogen compounds.
A simplified equation for biological nitrogen fixation is:
N2 + 8H + + 8e− → 2N H3 + H2
Energy is used in the process, but this is not shown in the above equation.
Another important source of ammonia in the soil is decomposition. When animals and
plants die, the nitrogen compounds that were present in them are broken down and converted into ammonia. This process is carried out by decomposition bacteria and fungi in
the soil.
2. Industrial nitrogen fixation
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CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
19.3
In the Haber-Bosch process, nitrogen (N2 ) is converted together with hydrogen gas (H2 )
into ammonia (NH3 ) fertiliser. This is an artificial process.
3. Lightning
In the atmosphere, lightning and photons are important in the reaction between nitrogen
(N2 ) and oxygen (O2 ) to form nitric oxide (NO) and then nitrates.
teresting It is interesting to note that by cultivating legumes, using the Haber-Bosch
Interesting
Fact
Fact
process to manufacture chemical fertilisers and increasing pollution from vehicles
and industry, humans have more than doubled the amount of nitrogen that would
normally be changed from nitrogen gas into a biologically useful form. This has
serious environmental consequences.
19.3
Nitrification
Nitrification involves two biological oxidation reactions: firstly, the oxidation of ammonia with
oxygen to form nitrite (NO−
2 ) and secondly the oxidation of these nitrites into nitrates.
1. N H3 + O2 → N O2− + 3H + + 2e− (production of nitrites)
2. N O2− + H2 O → N O3− + 2H + + 2e− (production of nitrates)
Nitrification is an important step in the nitrogen cycle in soil because it converts the ammonia
(from the nitrogen fixing part of the cycle) into nitrates, which are easily absorbed by the roots
of plants. This absorption of nitrates by plants is called assimilation. Once the nitrates have
been assimilated by the plants, they become part of the plants’ proteins. These plant proteins
are then available to be eaten by animals. In other words, animals (including humans) obtain
their own nitrogen by feeding on plants. Nitrification is performed by bacteria in the soil, called
nitrifying bacteria.
Activity :: Case Study : Nitrates in drinking water
Read the information below and then carry out your own research to help you
answer the questions that follow.
The negatively charged nitrate ion is not held onto soil particles and so can
be easily washed out of the soil. This is called leaching. In this way, valuable
nitrogen can be lost from the soil, reducing the soil’s fertility. The nitrates can
then accumulate in groundwater, and eventually in drinking water. There are strict
regulations that control how much nitrate can be present in drinking water, because
nitrates can be reduced to highly reactive nitrites by microorganisms in the gut.
Nitrites are absorbed from the gut and bind to haemoglobin (the pigment in blood
that helps to transport oxygen around the body). This reduces the ability of the
haemoglobin to carry oxygen. In young babies this can lead to respiratory distress,
a condition known as ”blue baby syndrome”.
1. How is nitrate concentration in water measured?
2. What concentration of nitrates in drinking water is considered acceptable? You
can use drinking water standards for any part of the world, if you can’t find any
for South Africa.
3. What is ’blue baby syndrome’ and what are the symptoms of the disease?
371
19.4
19.4
CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
Denitrification
Denitrification is the process of reducing nitrate and nitrite into gaseous nitrogen. The process
is carried out by denitrification bacteria. The nitrogen that is produced is returned to the atmosphere to complete the nitrogen cycle.
The equation for the reaction is:
2N O3− + 10e− + 12H + → N2 + 6H2 O
19.5
Human Influences on the Nitrogen Cycle
Humans have contributed significantly to the nitrogen cycle in a number of ways.
• Both artificial fertilisation and the planting of nitrogen fixing crops, increase the amount
of nitrogen in the soil. In some ways this has positive effects because it increases the fertility
of the soil, and means that agricultural productivity is high. On the other hand, however, if
there is too much nitrogen in the soil, it can run off into nearby water courses such as rivers,
or can become part of the groundwater supply as we mentioned earlier. Increased nitrogen
in rivers and dams can lead to a problem called eutrophication. Eutrophication is a process
where water bodies such as rivers, estuaries, dams and slow-moving streams receive excess
nutrients (e.g. nitrogen and phosphorus compounds) that stimulate excessive plant growth.
Sometimes this can cause certain plant species to be favoured over the others and one
species may ’take over’ the ecosystem, resulting in a decrease in plant diversity. This
is called a ’bloom’. Eutrophication also affects water quality. When the plants die and
decompose, large amounts of oxygen are used up and this can cause other animals in the
water to die.
Activity :: Case Study : Fertiliser use in South Africa
Refer to the data table below, which shows the average fertiliser use (in
kilograms per hectare or kg/ha) over a number of years for South Africa and
the world. Then answer the questions that follow:
SA
World
1965
27.9
34.0
1970
42.2
48.9
1975
57.7
63.9
1980
80.3
80.6
1985
66.6
86.7
1990
54.9
90.9
1995
48.5
84.9
2000
47.1
88.2
2002
61.4
91.9
1. On the same set of axes, draw two line graphs to show how fertiliser use
has changed in SA and the world between 1965 and 2002.
2. Describe the trend you see for...
(a) the world
(b) South Africa
3. Suggest a reason why the world’s fertiliser use has changed in this way over
time.
4. Do you see the same pattern for South Africa?
5. Try to suggest a reason for the differences you see in the fertiliser use data
for South Africa.
6. One of the problems with increased fertiliser use is that there is a greater
chance of nutrient runoff into rivers and dams, and therefore a greater
danger of eutrophication. In groups of 5-6, discuss the following questions:
(a) What could farmers do to try to reduce the risk of nutrient runoff from
fields into water systems? Try to think of at least 3 different strategies
that they could use.
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CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
19.6
(b) Imagine you are going to give a presentation on eutrophication to a
group of farmers who know nothing about it. How will you educate
them about the dangers? How will you convince them that it is in their
interests to change their farming practices? Present your ideas to the
class.
• Atmospheric pollution is another problem. The main culprits are nitrous oxide (N2 O),
nitric oxide (NO) and nitrogen dioxide (NO2 ). Most of these gases result either from emissions from agricultural soils (and particularly artificial fertilisers), or from the combustion
of fossil fuels in industry or motor vehicles. The combustion (burning) of nitrogen-bearing
fuels such as coal and oil releases this nitrogen as N2 or NO gases. Both NO2 and NO can
combine with water droplets in the atmosphere to form acid rain. Furthermore, both NO
and NO2 contribute to the depletion of the ozone layer and some are greenhouse gases.
In high concentrations these gases can contribute towards global warming.
19.6
The industrial fixation of nitrogen
A number of industrial processes are able to fix nitrogen into different compounds and then
convert these compounds into fertilisers. In the descriptions below, you will see how atmospheric
nitrogen is fixed to produce ammonia, how ammonia is then reacted with oxygen to form nitric
acid and how nitric acid and ammonia are then used to produce the fertiliser, ammonium nitrate.
• Preparation of ammonia (NH3 )
The industrial preparation of ammonia is known as the Haber-Bosch process. At a high
pressure and a temperature of approximately 5000 C, and in the presence of a suitable
catalyst (usually iron), nitrogen and hydrogen react according to the following equation:
N2 + 3H2 → 2N H3
Ammonia is used in the preparation of artficial fertilisers such as (NH4 )2 SO4 and is also
used in cleaning agents and cooling installations.
teresting Fritz Haber and Carl Bosch were the two men responsible for developing
Interesting
Fact
Fact
the Haber-Bosch process. In 1918, Haber was awarded the Nobel Prize in
Chemistry for his work. The Haber-Bosch process was a milestone in industrial chemistry because it meant that nitrogenous fertilisers were cheaper
and much more easily available. At the time, this was very important in
providing food for the growing human population.
Haber also played a major role in the development of chemical warfare in
World War I. Part of this work included the development of gas masks with
absorbent filters. He also led the teams that developed chlorine gas and
other deadly gases for use in trench warfare. His wife, Clara Immerwahr,
also a chemist, opposed his work on poison gas and committed suicide with
his service weapon in their garden. During the 1920s, scientists working at
his institute also developed the cyanide gas formulation Zyklon B, which
was used as an insecticide and also later, after he left the programme, in the
Nazi extermination camps.
Haber was Jewish by birth, but converted from Judaism in order to be more
accepted in Germany. Despite this, he was forced to leave the country in
1933 because he was Jewish ’by definition’ (his mother was Jewish). He died
in 1934 at the age of 65. Many members of his extended family died in the
Nazi concentration camps, possibly gassed by Zyklon B.
373
19.7
CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
• Preparation of nitric acid (HNO3 )
Nitric acid is used to prepare fertilisers and explosives. The industrial preparation of nitric
acid is known as the Ostwald process. The Ostwald process involves the conversion of
ammonia into nitric acid in various stages:
Firstly, ammonia is heated with oxygen in the presence of a platinum catalyst to form nitric
oxide and water.
4N H3 (g) + 5O2 (g) → 4N O(g) + 6H2 O(g)
Secondly, nitric oxide reacts with oxygen to form nitrogen dioxide. This gas is then readily
absorbed by the water to produce nitric acid. A portion of nitrogen dioxide is reduced back
to nitric oxide.
2N O(g) + O2 (g) → 2N O2 (g)
3N O2 (g) + H2 O(l) → 2HN O3 (aq) + N O(g)
The NO is recycled, and the acid is concentrated to the required strength by a process
called distillation.
• Preparation of ammonium nitrate
Ammonium nitrate is used as a fertiliser, as an explosive and also in the preparation of
’laughing gas’ which is used as an anaesthetic. Ammonium nitrate is prepared by reacting
ammonia with nitric acid:
N H3 + HN O3 → N H4 N O3
Activity :: Debate : Fertiliser use
Divide the class into two groups to debate the following topic:
Increasing the use of artificial fertilisers is the best solution to meet the growing
food needs of the world’s human population.
One group should take the position of agreeing with the statement, and the other
should disagree. In your groups, discuss reasons why you have the opinion that you
do, and record some notes of your discussion. Your teacher will then explain to you
how to proceed with the debate.
19.7
Summary
• Nitrogen is essential for life on earth, since it forms part of amino acids, proteins and
nucleic acids.
• The atmosphere is composed mostly of nitrogen gas, but the gas is inert, meaning that
it is not available to living organisms in its gaseous form.
• The nitrogen cycle describes how nitrogen and nitrogen-containing compounds are changed
into different forms in nature.
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CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
19.7
• The nitrogen cycle consists of three major processes: nitrogen fixation, nitrification and
denitrification.
• Nitrogen fixation is the conversion of atmospheric nitrogen into compounds such as
ammonia, that are more easily used.
• Nitrogen can be fixed biologically through the actions of bacteria, industrially through
the Haber-Bosch process or by lightning.
• Nitrification converts ammonia into nitrites and nitrates, which can be easily assimilated
by plants.
• Denitrification converts nitrites and nitrates back into gaseous nitrogen to complete the
nitrogen cycle.
• Humans have had a number of impacts on the nitrogen cycle. The production of artificial
fertilisers for example, means that there is a greater chance of runoff into water systems.
In some cases, eutrophication may occur.
• Eutrophication is the enrichment of water systems with excess nutrients, which may
stimulate excessive plant growth at the expense of other parts of the ecosystem.
• Many nitrogen gases such as NO, N2 O and NO2 are released by agricultural soils and
artificial fertilisers. These gases may combine with water vapour in the atmosphere and
result in acid rain. Some of these gases are also greenhouse gases and may contribute
towards global warming.
• A number of industrial processes are used to produce articifical fertilisers.
• The Haber-Bosch process converts atmsopheric nitrogen into ammonia.
• The Ostwald process reacts ammonia with oxygen to produce nitric acid, which is used
in the preparation of fertilisers and explosives.
• If ammonia and nitric acid react, the product is ammonium nitrate, which is used as a
fertiliser and as an explosive.
Exercise: Summary Exercise
1. Look at the diagram and the descriptions of the nitrogen cycle earlier in the
chapter:
(a) Would you describe the changes that take place in the nitrogen cycle as
chemical or physical changes? Explain your answer.
(b) Are the changes that take place in the water cycle physical or chemical
changes? Explain your answer.
2. Explain what is meant by each of the following terms:
(a) nitrogen fixing
(b) fertiliser
(c) eutrophication
3. Explain why the fixing of atmospheric nitrogen is so important for the survival
of life on earth.
4. Refer to the diagram below and then answer the questions that follow:
375
19.7
CHAPTER 19. GLOBAL CYCLES: THE NITROGEN CYCLE - GRADE 10
N2
(1)
(2)
(3)
(4)
(5)
(a) Explain the role of decomposers in the nitrogen cycle.
(b) If the process taking place at (3) is nitrification, then label the processes
at (1) and (5).
(c) Identify the nitrogen products at (2) and (4).
(d) On the diagram, indicate the type of bacteria that are involved in each
stage of the nitrogen cycle.
(e) In industry, what process is used to produce the compound at 2?
(f) Does the diagram above show a ’cycle’ ? Explain your answer.
5. NO and NO2 are both nitrogen compounds:
(a) Explain how each of these compounds is formed?
(b) What effect does each of these compounds have in the environment?
6. There are a number of arguments both ’for’ and ’against’ the use of artificial
fertilisers. Draw a table to summarise the advantages and disadvantages of
their use.
376
Chapter 20
The Hydrosphere - Grade 10
20.1
Introduction
As far as we know, the Earth we live on is the only planet that is able to support life. Among
other things, Earth is just the right distance from the sun to have temperatures that are suitable
for life to exist. Also, the Earth’s atmosphere has exactly the right type of gases in the right
amounts for life to survive. Our planet also has water on its surface, which is something very
unique. In fact, Earth is often called the ’Blue Planet’ because most of it is covered in water.
This water is made up of freshwater in rivers and lakes, the saltwater of the oceans and estuaries,
groundwater and water vapour. Together, all these water bodies are called the hydrosphere.
20.2
Interactions of the hydrosphere
It is important to realise that the hydrosphere interacts with other global systems, including the
atmosphere, lithosphere and biosphere.
• Atmosphere
When water is heated (e.g. by energy from the sun), it evaporates and forms water vapour.
When water vapour cools again, it condenses to form liquid water which eventually returns
to the surface by precipitation e.g. rain or snow. This cycle of water moving through the
atmosphere, and the energy changes that accompany it, is what drives weather patterns
on earth.
• Lithosphere
In the lithosphere (the ocean and continental crust at the Earth’s surface), water is an
important weathering agent, which means that it helps to break rock down into rock
fragments and then soil. These fragments may then be transported by water to another
place, where they are deposited. This is called erosion. These two process i.e. weathering
and erosion, help to shape the earth’s surface. You can see this for example in rivers.
In the upper streams, rocks are eroded and sediments are transported down the river and
deposited on the wide flood plains lower down. On a bigger scale, river valleys in mountains
have been carved out by the action of water, and cliffs and caves on rocky beach coastlines,
are also the result of weathering and erosion by water.
• Biosphere
In the biosphere, land plants absorb water through their roots and then transport this
through their vascular (transport) system to stems and leaves. This water is needed in
photosynthesis, the food production process in plants. Transpiration (evaporation of water
from the leaf surface) then returns water back to the atmosphere.
377
20.3
20.3
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Exploring the Hydrosphere
The large amount of water on our planet is something quite unique. In fact, about 71% of
the earth is covered by water. Of this, almost 97% is found in the oceans as saltwater, about
2.2% occurs as a solid in ice sheets, while the remaining amount (less than 1%) is available as
freshwater. So from a human perspective, despite the vast amount of water on the planet, only a
very small amount is actually available for human consumption (e.g. drinking water). Before we
go on to look more closely at the chemistry of the hydrosphere, we are going to spend some time
exploring a part of the hydrosphere, in order to start appreciating what a complex and beautiful
part of the world it is.
Activity :: Investigation : Investigating the hydrosphere
1. Choosing a study site:
For this exercise, you can choose any part of the hydrosphere that you would
like to explore. This may be a rock pool, a lake, river, wetland or even just a
small pond. The guidelines below will apply best to a river investigation, but
you can ask similar questions and gather similar data in other areas. When
choosing your study site, consider how accessible it is (how easy is it to get
to?) and the problems you may experience (e.g. tides, rain).
2. Collecting data:
Your teacher will provide you with the equipment you need to collect the following data. You should have at least one study site where you will collect
data, but you might decide to have more if you want to compare your results
in different areas. This works best in a river, where you can choose sites down
its length.
(a) Chemical data
Measure and record data such as temperature, pH, conductivity and dissolved oxygen at each of your sites. You may not know exactly what
these measurements mean right now, but it will become clearer later in the
chapter.
(b) Hydrological data
Measure the water velocity of the river and observe how the volume of water
in the river changes as you move down its length. You can also collect a
water sample in a clear bottle, hold it to the light and see whether the
water is clear or whether it has particles in it.
(c) Biological data
What types of animals and plants are found in or near this part of the
hydrosphere? Are they specially adapted to their environment?
Record your data in a table like the one shown below:
Site 1
Site 2
Site 3
Temperature
pH
Conductivity
Dissolved oxygen
Animals and plants
3. Interpreting the data:
Once you have collected and recorded your data, think about the following
questions:
• How does the data you have collected vary at different sites?
• Can you explain these differences?
• What effect do you think temperature, dissolved oxygen and pH have on
animals and plants that are living in the hydrosphere?
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20.4
• Water is seldom ’pure’. It usually has lots of things dissolved (e.g. Mg,
Ca and NO−
3 ions) or suspended (e.g. soil particles, debris) in it. Where
do these substances come from?
• Are there any human activities near this part of the hydrosphere? What
effect could these activities have on the hydrosphere?
20.4
The Importance of the Hydrosphere
It is so easy sometimes to take our hydrosphere for granted, and we seldom take the time to
really think about the role that this part of the planet plays in keeping us alive. Below are just
some of the very important functions of water in the hydrosphere:
• Water is a part of living cells
Each cell in a living organism is made up of almost 75% water, and this allows the cell
to function normally. In fact, most of the chemical reactions that occur in life, involve
substances that are dissolved in water. Without water, cells would not be able to carry
out their normal functions, and life could not exist.
• Water provides a habitat
The hydrosphere provides an important place for many animals and plants to live. Many
−
+
gases (e.g. CO2 , O2 ), nutrients e.g. nitrate (NO−
3 ), nitrite (NO2 ) and ammonium (NH4 )
2+
2+
ions, as well as other ions (e.g. Ca and Mg ) are dissolved in water. The presence of
these substances is critical for life to exist in water.
• Regulating climate
You may remember from chapter ?? that one of water’s unique characteristics is its high
specific heat. This means that water takes a long time to heat up, and also a long time
to cool down. This is important in helping to regulate temperatures on earth so that they
stay within a range that is acceptable for life to exist. Ocean currents also help to disperse
heat.
• Human needs
Humans use water in a number of ways. Drinking water is obviously very important, but
water is also used domestically (e.g. washing and cleaning) and in industry. Water can
also be used to generate electricity through hydropower.
These are just a few of the very important functions that water plays on our planet. Many of the
functions of water relate to its chemistry and to the way in which it is able to dissolve substances
in it.
20.5
Ions in aqueous solution
As we mentioned earlier, water is seldom pure. Because of the structure of the water molecule,
it is able to dissolve substances in it. This is very important because if water wasn’t able to do
this, life would not be able to survive. In rivers and the oceans for example, dissolved oxygen
means that organisms are still able to respire (breathe). For plants, dissolved nutrients are also
available. In the human body, water is able to carry dissolved substances from one part of the
body to another.
Many of the substances that dissolve are ionic, and when they dissolve they form ions in solution.
We are going to look at how water is able to dissolve ionic compounds, and how these ions
maintain a balance in the human body, how they affect water hardness, and how specific ions
determine the pH of solutions.
379
20.5
20.5.1
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Dissociation in water
You may remember from chapter 5 that water is a polar molecule (figure 20.1). This means
that one part of the molecule has a slightly positive charge and the other part has a slightly
negative charge.
δ+
H2 O
δ−
Figure 20.1: Water is a polar molecule
It is the polar nature of water that allows ionic compounds to dissolve in it. In the case of
sodium chloride (NaCl) for example, the positive sodium ions (Na+ ) will be attracted to the
negative pole of the water molecule, while the negative chloride ions (Cl− ) will be attracted to
the positive pole of the water molecule. In the process, the ionic bonds between the sodium
and chloride ions are weakened and the water molecules are able to work their way between the
individual ions, surrounding them and slowly dissolving the compound. This process is called
dissociation. A simplified representation of this is shown in figure 20.2.
Definition: Dissociation
Dissociation in chemistry and biochemistry is a general process in which ionic compounds
separate or split into smaller molecules or ions, usually in a reversible manner.
Cl−
δ+
H2 O
δ−
Cl− δ + H2 O δ − N a+ δ − H2 O δ + Cl−
δ−
H2 O
δ+
Cl−
Figure 20.2: Sodium chloride dissolves in water
The dissolution of sodium chloride can be represented by the following equation:
N aCl(s) → N a+ (aq) + Cl− (aq)
The symbols s (solid), l (liquid), g (gas) and aq (material is dissolved in water) are written after
the chemical formula to show the state or phase of the material. The dissolution of potassium
sulphate into potassium and sulphate ions is shown below as another example:
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CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.5
K2 SO4 (s) → 2K + (aq) + SO42− (aq)
Remember that molecular substances (e.g. covalent compounds) may also dissolve, but most
will not form ions. One example is sugar.
C6 H12 O6 (s) ⇔ C6 H12 O6 (aq)
There are exceptions to this and some molecular substances will form ions when they dissolve.
Hydrogen chloride for example can ionise to form hydrogen and chloride ions.
HCl(g) → H + (aq) + Cl− (aq)
teresting The ability of ionic compounds to dissolve in water is extremely important in
Interesting
Fact
Fact
the human body! The body is made up of cells, each of which is surrounded
by a membrane. Dissolved ions are found inside and outside of body cells, in
different concentrations. Some of these ions are positive (e.g. Mg2+ ) and some
are negative (e.g. Cl− ). If there is a difference in the charge that is inside and
outside the cell, then there is a potential difference across the cell membrane.
This is called the membrane potential of the cell. The membrane potential
acts like a battery and affects the movement of all charged substances across the
membrane. Membrane potentials play a role in muscle functioning, digestion,
excretion and in maintaining blood pH, to name just a few. The movement of
ions across the membrane can also be converted into an electric signal that can
be transferred along neurons (nerve cells), which control body processes. If ionic
substances were not able to dissociate in water, then none of these processes
would be possible! It is also important to realise that our bodies can lose ions
such as Na+ , K+ , Ca2+ , Mg2+ , and Cl− , for example when we sweat during
exercise. Sports drinks such as Lucozade and Powerade are designed to replace
these lost ions so that the body’s normal functioning is not affected.
Exercise: Ions in solution
1. For each of the following, say whether the substance is ionic or molecular.
(a)
(b)
(c)
(d)
potassium nitrate (KNO3 )
ethanol (C2 H5 OH)
sucrose sugar (C12 H22 O11
sodium bromide (NaBr)
2. Write a balanced equation to show how each of the following ionic compounds
dissociate in water.
(a)
(b)
(c)
(d)
sodium sulphate (Na2 SO4 )
potassium bromide (KBr)
potassium permanganate (KMNO4 )
sodium phosphate (Na3 PO4 )
381
20.5
20.5.2
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Ions and water hardness
Definition: Water hardness
Water hardness is a measure of the mineral content of water. Minerals are substances such
as calcite, quartz and mica that occur naturally as a result of geological processes.
Hard water is water that has a high mineral content. Water that has a low mineral content
is known as soft water. If water has a high mineral content, it usually contains high levels of
metal ions, mainly calcium (Ca) and magnesium (Mg). The calcium enters the water from either
CaCO3 (limestone or chalk) or from mineral deposits of CaSO4 . The main source of magnesium
is a sedimentary rock called dolomite, CaMg(CO3 )2 . Hard water may also contain other metals
as well as bicarbonates and sulphates.
teresting The simplest way to check whether water is hard or soft is to use the lather/froth
Interesting
Fact
Fact
test. If the water is very soft, soap will lather more easily when it is rubbed
against the skin. With hard water this won’t happen. Toothpaste will also not
froth well in hard water.
A water softener works on the principle of ion exchange. Hard water passes through a media
bed, usually made of resin beads that are supersaturated with sodium. As the water passes
through the beads, the hardness minerals (e.g. calcium and magnesium) attach themselves to
the beads. The sodium that was originally on the beads is released into the water. When the
resin becomes saturated with calcium and magnesium, it must be recharged. A salt solution is
passed through the resin. The sodium replaces the calcium and magnesium, and these ions are
released into the waste water and discharged.
20.5.3
The pH scale
The concentration of specific ions in solution, affects whether the solution is acidic or basic. You
will learn about acids and bases in chapter 15. Acids and bases can be described as substances
that either increase or decrease the concentration of hydrogen (H+ or H3 O+) ions in a solution.
An acid increases the hydrogen ion concentration in a solution, while a base decreases the
hydrogen ion concentration. pH is used to measure whether a substance is acidic or basic
(alkaline).
Definition: pH
pH is a measure of the acidity or alkalinity of a solution. The pH scale ranges from 0 to 14.
Solutions with a pH less than seven are acidic, while those with a pH greater than seven
are basic (alkaline). pH 7 is considered neutral.
pH can be calculated using the following equation:
pH = −log[H + ]
or
pH = −log[H3 O+ ]
The brackets in the above equation are used to show concentration in mol.dm−3 .
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CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.5
Worked Example 93: pH calculations
Question: Calculate the pH of a solution where the concentration of hydrogen ions
is 1 × 10−7 mol.dm−3 .
Answer
Step 1 : Determine the concentration of hydrogen ions in mol.dm−3
In this example, the concentration has been given and is 1 × 10−7 mol.dm−3
Step 2 : Substitute this value into the pH equation and calculate the pH
value
pH = -log[H+]
= -log(1 × 10−7 )
=7
Worked Example 94: pH calculations
Question: In a solution of ethanoic acid, the following equilibrium is established:
CH3 COOH(aq) + H2 O ⇔ CH3 COO− (aq) + H3 O+
The concentration of CH3 COO− ions is found to be 0.003 mol.dm−3 . Calculate the
pH of the solution.
Answer
Step 1 : Determine the concentration of hydrogen ions in the solution
According to the balanced equation for this reaction, the mole ratio of CH3 COO−
ions to H3 O+ ions is the same, therefore the concentration of these two ions in the
solution will also be the same. So, [H3 O+ ] = 0.003 dm−3 .
Step 2 : Substitute this value into the pH equation and calculate the pH
value
pH = -log[H3O+ ]
= -log(0.003)
= 2.52
Understanding pH is very important. In living organisms, it is necessary to maintain a constant
pH so that chemical reactions can occur under optimal conditions.
Important: It may also be useful for calculations involving the pH scale, to know that the
following equation can also be used:
[H3 O+ ][OH− ] = 1 × 10−14
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20.5
CHAPTER 20. THE HYDROSPHERE - GRADE 10
teresting A build up of acid in the human body can be very dangerous. Lactic acidosis
Interesting
Fact
Fact
is a condition caused by the buildup of lactic acid in the body. It leads to
acidification of the blood (acidosis) and can make a person very ill. Some of
the symptoms of lactic acidosis are deep and rapid breathing, vomiting, and
abdominal pain. In the fight against HIV, lactic acidosis is a problem. One of
the antiretrovirals (ARV’s) that is used in anti-HIV treatment is Stavudine (also
known as Zerit or d4T). One of the side effects of Stavudine is lactic acidosis,
particularly in overweight women. If it is not treated quickly, it can result in
death.
In agriculture, farmers need to know the pH of their soils so that they are able to plant the
right kinds of crops. The pH of soils can vary depending on a number of factors such as
rainwater, the kinds of rocks and materials from which the soil was formed and also human
influences such as pollution and fertilisers. The pH of rain water can also vary and this too has
an effect on agriculture, buildings, water courses, animals and plants. Rainwater is naturally
acidic because carbon dioxide in the atmosphere combines with water to form carbonic acid.
Unpolluted rainwater has a pH of approximately 5.6. However, human activities can alter the
acidity of rain and this can cause serious problems such as acid rain.
Exercise: Calculating pH
1. Calculate the pH of each of the following solutions:
(a) A 0.2 mol.dm−3 KOH solution
(b) A 0.5 mol.dm−3 HCl solution
2. What is the concentration (in mol.dm−3 ) of H3 O+ ions in a NaOH solution
which has a pH of 12?
3. The concentrations of hydronium and hydroxyl ions in a typical sample of
seawater are 10−8 mol.dm−3 and 10−6 mol.dm−3 respectively.
(a) Is the seawater acidic or basic?
(b) What is the pH of the seawater?
(c) Give a possible explanation for the pH of the seawater.
(IEB Paper 2, 2002)
20.5.4
Acid rain
The acidity of rainwater comes from the natural presence of three substances (CO2 , NO, and
SO2 ) in the lowest layer of the atmosphere. These gases are able to dissolve in water and
therefore make rain more acidic than it would otherwise be. Of these gases, carbon dioxide
(CO2 ) has the highest concentration and therefore contributes the most to the natural acidity
of rainwater. We will look at each of these gases in turn.
Definition: Acid rain
Acid rain refers to the deposition of acidic components in rain, snow and dew. Acid rain
occurs when sulfur dioxide and nitrogen oxides are emitted into the atmosphere, undergo
chemical transformations, and are absorbed by water droplets in clouds. The droplets then
fall to earth as rain, snow, mist, dry dust, hail, or sleet. This increases the acidity of the
soil, and affects the chemical balance of lakes and streams.
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CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.5
1. Carbon dioxide
Carbon dioxide reacts with water in the atmosphere to form carbonic acid (H2 CO3 ).
CO2 + H2 O → H2 CO3
The carbonic acid dissociates to form hydrogen and hydrogen carbonate ions. It is the
presence of hydrogen ions that lowers the pH of the solution, making the rain acidic.
H2 CO3 → H + + HCO3−
2. Nitric oxide
Nitric oxide (NO) also contributes to the natural acidity of rainwater and is formed during
lightning storms when nitrogen and oxygen react. In air, NO is oxidised to form nitrogen
dioxide (NO2 ). It is the nitrogen dioxide which then reacts with water in the atmosphere
to form nitric acid (HNO3 ).
3N O2 (g) + H2 O → 2HN O3 (aq) + N O(g)
The nitric acid dissociates in water to produce hydrogen ions and nitrate ions. This again
lowers the pH of the solution, making it acidic.
HN O3 → H + + N O3−
3. Sulfur dioxide
Sulfur dioxide in the atmosphere first reacts with oxygen to form sulfur trioxide, before
reacting with water to form sulfuric acid.
2SO2 + O2 → 2SO3
SO3 + H2 O → H2 SO4
Sulfuric acid dissociates in a similar way to the previous reactions.
H2 SO4 → HSO4− + H +
Although these reactions do take place naturally, human activities can greatly increase the concentration of these gases in the atmosphere, so that rain becomes far more acidic than it would
otherwise be. The burning of fossil fuels in industries, vehicles etc is one of the biggest culprits.
If the acidity of the rain drops below 5, it is referred to as acid rain.
Acid rain can have a very damaging effect on the environment. In rivers, dams and lakes, increased acidity can mean that some species of animals and plants will not survive. Acid rain can
also degrade soil minerals, producing metal ions that are washed into water systems. Some of
these ions may be toxic e.g. Al3+ . From an economic perspective, altered soil pH can drastically
affect agricultural productivity.
Acid rain can also affect buildings and monuments, many of which are made from marble and
limestone. A chemical reaction takes place between CaCO3 (limestone) and sulfuric acid to
produce aqueous ions which can be easily washed away. The same reaction can occur in the
lithosphere where limestone rocks are present e.g. limestone caves can be eroded by acidic
rainwater.
H2 SO4 + CaCO3 → CaSO4 .H2 O + CO2
385
20.6
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Activity :: Investigation : Acid rain
You are going to test the effect of ’acid rain’ on a number of substances.
Materials needed:
samples of chalk, marble, zinc, iron, lead, dilute sulfuric acid, test tubes, beaker,
glass dropper
Method:
1. Place a small sample of each of the following substances in a separate test
tube: chalk, marble, zinc, iron and lead
2. To each test tube, add a few drops of dilute sulfuric acid.
3. Observe what happens and record your results.
Discussion questions:
• In which of the test tubes did reactions take place? What happened to the
sample substances?
• What do your results tell you about the effect that acid rain could have on each
of the following: buildings, soils, rocks and geology, water ecosystems?
• What precautions could be taken to reduce the potential impact of acid rain?
20.6
Electrolytes, ionisation and conductivity
Conductivity in aqueous solutions, is a measure of the ability of water to conduct an electric
current. The more ions there are in the solution, the higher its conductivity.
Definition: Conductivity
Conductivity is a measure of a solution’s ability to conduct an electric current.
20.6.1
Electrolytes
An electrolyte is a material that increases the conductivity of water when dissolved in it.
Electrolytes can be further divided into strong electrolytes and weak electrolytes.
Definition: Electrolyte
An electrolyte is a substance that contains free ions and behaves as an electrically conductive
medium. Because they generally consist of ions in solution, electrolytes are also known as
ionic solutions.
1. Strong electrolytes
A strong electrolyte is a material that ionises completely when it is dissolved in water:
AB(s,l,g) → A+ (aq) + B − (aq)
This is a chemical change because the original compound has been split into its component ions and bonds have been broken. In a strong electrolyte, we say that the extent
of ionisation is high. In other words, the original material dissociates completely so that
there is a high concentration of ions in the solution. An example is a solution of potassium
nitrate:
KN O3 (s) → K + (aq) + N O3− (aq)
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CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.6
2. Weak electrolytes
A weak electrolyte is a material that goes into solution and will be surrounded by water
molecules when it is added to water. However, not all of the molecules will dissociate into
ions. The extent of ionisation of a weak electrolyte is low and therefore the concentration
of ions in the solution is also low.
AB(s,l,g) → AB(aq) ⇔ A+ (aq) + B − (aq)
The following example shows that, in the final solution of a weak electrolyte, some of the
original compound plus some dissolved ions are present.
C2 H3 O2 H(l) → C2 H3 O2 H ⇔ C2 H3 O2− (aq) + H + (aq)
20.6.2
Non-electrolytes
A non-electrolyte is a material that does not increase the conductivity of water when dissolved
in it. The substance goes into solution and becomes surrounded by water molecules, so that the
molecules of the chemical become separated from each other. However, although the substance
does dissolve, it is not changed in any way and no chemical bonds are broken. The change is a
physical change. In the oxygen example below, the reaction is shown to be reversible because
oxygen is only partially soluble in water and comes out of solution very easily.
C2 H5 OH(l) → C2 H5 OH(aq)
O2 (g) ⇔ O2 (aq)
20.6.3
Factors that affect the conductivity of water
The conductivity of water is therefore affected by the following factors:
• The type of substance that dissolves in water
Whether a material is a strong electrolyte (e.g. potassium nitrate, KNO3 ), a weak electrolyte (e.g. acetate, C2 H3 O2 H) or a non-electrolyte (e.g. sugar, alcohol, oil) will affect
the conductivity of water because the concentration of ions in solution will be different in
each case.
• The concentration of ions in solution
The higher the concentration of ions in solution, the higher its conductivity will be.
• Temperature
The warmer the solution the higher the solubility of the material being dissolved, and
therefore the higher the conductivity as well.
Activity :: Experiment : Electrical conductivity
Aim:
To investigate the electrical conductivities of different substances and solutions.
Apparatus:
solid salt (NaCl) crystals; different liquids such as distilled water, tap water,
seawater, benzene and alcohol; solutions of salts e.g. NaCl, KBr; a solution of
an acid (e.g. HCl) and a solution of a base (e.g. NaOH); torch cells; ammeter;
conducting wire, crocodile clips and 2 carbon rods.
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20.6
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Method:
Set up the experiment by connecting the circuit as shown in the diagram below.
In the diagram, ’X’ represents the substance or solution that you will be testing.
When you are using the solid crystals, the crocodile clips can be attached directly to
each end of the crystal. When you are using solutions, two carbon rods are placed
into the liquid, and the clips are attached to each of the rods. In each case, complete
the circuit and allow the current to flow for about 30 seconds. Observe whether the
ammeter shows a reading.
battery
Ammeter
A
test substance
X
crocodile clip
Results:
Record your observations in a table similar to the one below:
Test substance
Ammeter reading
What do you notice? Can you explain these observations?
Remember that for electricity to flow, there needs to be a movement of charged
particles e.g. ions. With the solid NaCl crystals, there was no flow of electricity
recorded on the ammeter. Although the solid is made up of ions, they are held
together very tightly within the crystal lattice, and therefore no current will flow.
Distilled water, benzene and alcohol also don’t conduct a current because they are
covalent compounds and therefore do not contain ions.
The ammeter should have recorded a current when the salt solutions and the acid
and base solutions were connected in the circuit. In solution, salts dissociate into
their ions, so that these are free to move in the solution. Acids and bases behave
in a similar way, and dissociate to form hydronium and oxonium ions. Look at the
following examples:
KBr → K+ + Br−
NaCl → Na+ + Cl−
HCl + H2 O → H3 O+ + Cl−
NaOH → Na+ + OH−
Conclusions:
Solutions that contain free-moving ions are able to conduct electricity because of
the movement of charged particles. Solutions that do not contain free-moving ions
do not conduct electricity.
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CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.7
teresting Conductivity in streams and rivers is affected by the geology of the area where
Interesting
Fact
Fact
the water is flowing through. Streams that run through areas with granite
bedrock tend to have lower conductivity because granite is made of materials
that do not ionise when washed into the water. On the other hand, streams
that run through areas with clay soils tend to have higher conductivity because
the materials ionise when they are washed into the water. Pollution can also affect conductivity. A failing sewage system or an inflow of fertiliser runoff would
raise the conductivity because of the presence of chloride, phosphate, and nitrate
(ions) while an oil spill (non-ionic) would lower the conductivity. It is very important that conductivity is kept within a certain acceptable range so that the
organisms living in these water systems are able to survive.
20.7
Precipitation reactions
Sometimes, ions in solution may react with each other to form a new substance that is insoluble.
This is called a precipitate.
Definition: Precipitate
A precipitate is the solid that forms in a solution during a chemical reaction.
Activity :: Demonstration : The reaction of ions in solution
Apparatus and materials:
4 test tubes; copper(II) chloride solution; sodium carbonate solution; sodium
sulphate solution
CuCl2
CuCl2
Na2 CO3
Na2 SO4
Method:
1. Prepare 2 test tubes with approximately 5 ml of dilute Cu(II)chloride solution
in each
2. Prepare 1 test tube with 5 ml sodium carbonate solution
3. Prepare 1 test tube with 5 ml sodium sulphate solution
4. Carefully pour the sodium carbonate solution into one of the test tubes containing copper(II) chloride and observe what happens
5. Carefully pour the sodium sulphate solution into the second test tube containing
copper(II) chloride and observe what happens
Results:
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20.7
CHAPTER 20. THE HYDROSPHERE - GRADE 10
1. A light blue precipitate forms when sodium carbonate reacts with copper(II)
chloride
2. No precipitate forms when sodium sulphate reacts with copper(II) chloride
It is important to understand what happened in the previous demonstration. We will look at
what happens in each reaction, step by step.
1. Reaction 1: Sodium carbonate reacts with copper(II) chloride
When these compounds react, a number of ions are present in solution: Cu2+ , Cl− , N a+
and CO32− .
Because there are lots of ions in solution, they will collide with each other and may recombine in different ways. The product that forms may be insoluble, in which case a precipitate
will form, or the product will be soluble, in which case the ions will go back into solution.
Let’s see how the ions in this example could have combined with each other:
Cu2+ + CO2−
3 → CuCO3
Cu2+ + 2Cl− → CuCl2
Na+ + Cl− → NaCl
Na+ + CO2−
3 → Na2 CO3
You can automatically exclude the reactions where sodium carbonate and copper(II) chloride are the products because these were the initial reactants. You also know that sodium
chloride (NaCl) is soluble in water, so the remaining product (copper carbonate) must be
the one that is insoluble. It is also possible to look up which salts are soluble and which
are insoluble. If you do this, you will find that most carbonates are insoluble, therefore the
precipitate that forms in this reaction must be CuCO3 . The reaction that has taken place
between the ions in solution is as follows:
2N a+ + CO32− + Cu2+ + 2Cl− → CuCO3 + 2N a+ + 2Cl−
2. Reaction 2: Sodium sulphate reacts with copper(II) chloride
The ions that are present in solution are Cu2+ , Cl− , N a+ and SO42− .
The ions collide with each other and may recombine in different ways. The possible combinations of the ions are as follows:
Cu2+ + SO2−
4 → CuSO4
Cu2+ + 2Cl− → CuCl2
Na+ + Cl− → NaCl
Na+ + SO2−
4 → Na2 SO4
If we look up which of these salts are soluble and which are insoluble, we see that most
chlorides and most sulphates are soluble. This is why no precipitate forms in this second
reaction. Even when the ions recombine, they immediately separate and go back into
solution. The reaction that has taken place between the ions in solution is as follows:
2N a+ + SO42− + Cu2+ + 2Cl− → 2N a+ + SO42− + Cu2+ + 2Cl−
Table 20.1 shows some of the general rules about the solubility of different salts based on a
number of investigations:
390
CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.8
Table 20.1: General rules for the solubility of salts
Salt
Nitrates
Potassium, sodium and ammonium salts
Chlorides
Sulphates
Carbonates
20.8
Solubility
All are soluble
All are soluble
All are soluble except silver chloride, lead(II)chloride and mercury(II)chloride
All
are
soluble
except
lead(II)sulphate,
barium sulphate and calcium sulphate
All are insoluble except those of
potassium, sodium and ammonium
Testing for common anions in solution
It is also possible to carry out tests to determine which ions are present in a solution.
20.8.1
Test for a chloride
Prepare a solution of the unknown salt using distilled water and add a small amount of silver
nitrate solution. If a white precipitate forms, the salt is either a chloride or a carbonate.
−
Cl− + Ag+ + NO−
3 → AgCl + NO3 (AgCl is white precipitate)
−
+
CO2−
3 + 2Ag + 2NO3 → Ag2 CO3 + 2NO2 (Ag2 CO3 is white precipitate)
The next step is to treat the precipitate with a small amount of concentrated nitric acid. If
the precipitate remains unchanged, then the salt is a chloride. If carbon dioxide is formed, and
the precipitate disappears, the salt is a carbonate.
AgCl + HNO3 → (no reaction; precipitate is unchanged)
Ag2 CO3 + 2HNO3 → 2AgNO3 + H2 O + CO2 (precipitate disappears)
20.8.2
Test for a sulphate
Add a small amount of barium chloride solution to a solution of the test salt. If a white precipitate
forms, the salt is either a sulphate or a carbonate.
2+
SO2−
+ Cl− → BaSO4 + Cl− (BaSO4 is a white precipitate)
4 + Ba
2+
CO2−
+ Cl− → BaCO3 + Cl− (BaCO3 is a white precipitate)
3 + Ba
If the precipitate is treated with nitric acid, it is possible to distinguish whether the salt is a
sulphate or a carbonate (as in the test for a chloride).
BaSO4 + HNO3 → (no reaction; precipitate is unchanged)
BaCO3 + 2HNO3 → Ba(NO3 )2 + H2 O + CO2 (precipitate disappears)
391
20.8
20.8.3
CHAPTER 20. THE HYDROSPHERE - GRADE 10
Test for a carbonate
If a sample of the dry salt is treated with a small amount of acid, the production of carbon
dioxide is a positive test for a carbonate.
Acid + CO32− → CO2
If the gas is passed through limewater and the solution becomes milky, the gas is carbon dioxide.
Ca(OH)2 + CO2 → CaCO3 + H2 O (It is the insoluble CaCO3 precipitate that makes the
limewater go milky)
20.8.4
Test for bromides and iodides
As was the case with the chlorides, the bromides and iodides also form precipitates when they
are reacted with silver nitrate. Silver chloride is a white precipitate, but the silver bromide and
silver iodide precipitates are both pale yellow. To determine whether the precipitate is a bromide
or an iodide, we use chlorine water and carbon tetrachloride (CCl4 ).
Chlorine water frees bromine gas from the bromide, and colours the carbon tetrachloride a reddish brown.
Chlorine water frees iodine gas from an iodide, and colours the carbon tetrachloride is coloured
purple.
Exercise: Precipitation reactions and ions in solution
1. Silver nitrate (AgNO3 ) reacts with potassium chloride (KCl) and a white precipitate is formed.
(a) Write a balanced equation for the reaction that takes place.
(b) What is the name of the insoluble salt that forms?
(c) Which of the salts in this reaction are soluble?
2. Barium chloride reacts with sulfuric acid to produce barium sulphate and hydrochloric acid.
(a) Write a balanced equation for the reaction that takes place.
(b) Does a precipitate form during the reaction?
(c) Describe a test that could be used to test for the presence of barium
sulphate in the products.
3. A test tube contains a clear, colourless salt solution. A few drops of silver
nitrate solution are added to the solution and a pale yellow precipitate forms.
Which one of the following salts was dissolved in the original solution?
(a)
(b)
(c)
(d)
NaI
KCl
K2 CO3
Na2 SO4
(IEB Paper 2, 2005)
392
CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.9
20.9
Threats to the Hydrosphere
It should be clear by now that the hydrosphere plays an extremely important role in the survival of
life on Earth, and that the unique properties of water allow various important chemical processes
to take place which would otherwise not be possible. Unfortunately for us however, there are
a number of factors that threaten our hydrosphere, and most of these threats are because of
human activities. We are going to focus on two of these issues: overuse and pollution and look
at ways in which these problems can possibly be overcome.
1. Overuse of water
We mentioned earlier that only a very small percentage of the hydrosphere’s water is
available as freshwater. However, despite this, humans continue to use more and more
water to the point where water consumption is fast approaching the amount of water
that is available. The situation is a serious one, particularly in countries such as South
Africa which are naturally dry, and where water resources are limited. It is estimated that
between 2020 and 2040, water supplies in South Africa will no longer be able to meet the
growing demand for water in this country. This is partly due to population growth, but
also because of the increasing needs of industries as they expand and develop. For each
of us, this should be a very scary thought. Try to imagine a day without water...difficult
isn’t it? Water is so much a part of our lives, that we are hardly aware of the huge part
that it plays in our daily lives.
Activity :: Discussion : Creative water conservation
As populations grow, so do the demands that are placed on dwindling water
resources. While many people argue that building dams helps to solve this watershortage problem, the reality is that dams are only a temporary solution, and
that they often end up doing far more ecological damage than good. The only
sustainable solution is to reduce the demand for water, so that water supplies
are sufficient to meet this. The more important question then is how to do this.
Discussion:
Divide the class into groups, so that there are about five people in each.
Each group is going to represent a different sector within society. Your teacher
will tell you which sector you belong to from the following: Farming, industry,
city management or civil society (i.e. you will represent the ordinary ’man on
the street’). In your groups, discuss the following questions as they relate to the
group of people you represent: (Remember to take notes during your discussions,
and nominate a spokesperson to give feedback to the rest of the class on behalf
of your group)
• What steps could be taken by your group to conserve water?
• Why do you think these steps are not being taken?
• What incentives do you think could be introduced to encourage this group
to conserve water more efficiently?
2. Pollution
Pollution of the hydrosphere is also a major problem. When we think of pollution, we
sometimes only think of things like plastic, bottles, oil and so on. But any chemical that
is present in the hydrosphere in an amount that is not what it should be is a pollutant.
Animals and plants that live in the hydrosphere are specially adapted to surviving within a
certain range of conditions. If these conditions are changed (e.g. through pollution), these
organisms may not be able to survive. Pollution then, can affect entire aquatic ecosystems.
The most common forms of pollution in the hydrosphere are waste products from humans
and from industries, nutrient pollution e.g. fertiliser runoff which causes eutrophication
(this will be discussed in a later section) and toxic trace elements such as aluminium,
mercury and copper to name a few. Most of these elements come from mines or from
industries.
393
20.10
CHAPTER 20. THE HYDROSPHERE - GRADE 10
It is important to realise that our hydrosphere exists in a delicate balance with other systems,
and that disturbing this balance can have serious consequences for life on this planet.
Activity :: Group Project : School Action Project
There is a lot that can be done within a school to save water. As a class,
discuss what actions could be taken by your class to make people more aware of how
important it is to conserve water.
20.10
Summary
• The hydrosphere includes all the water that is on Earth. Sources of water include freshwater (e.g. rivers, lakes), saltwater (e.g. oceans), groundwater (e.g. boreholes) and water
vapour. Ice (e.g. glaciers) is also part of the hydrosphere.
• The hydrosphere interacts with other global systems, including the atmosphere, lithosphere and biosphere.
• The hydrosphere has a number of important functions. Water is a part of all living cells,
it provides a habitat for many living organisms, it helps to regulate climate, and it is used
by humans for domestic, industrial and other use.
• The polar nature of water means that ionic compounds dissociate easily in aqueous
solution into their component ions.
• Ions in solution play a number of roles. In the human body for example, ions help to
regulate the internal environment (e.g. controlling muscle function, regulating blood pH).
Ions in solution also determine water hardness and pH.
• Water hardness is a measure of the mineral content of water. Hard water has a high
mineral concentration and generally also a high concentration of metal ions e.g. calcium
and magnesium. The opposite is true for soft water.
• pH is a measure of the concentration of hydrogen ions in solution. The formula used to
calculate pH is as follows:
pH = -log[H3O+ ] or pH = -log[H+ ]
A solution with a pH less than 7 is considered acidic and more than 7 is considered basic
(or alkaline). A neutral solution has a pH of 7.
• Gases such as CO2 , NO2 and SO2 dissolve in water to form weak acid solutions. Rain is
naturally acidic because of the high concentrations of carbon dioxide in the atmosphere.
Human activities such as burning fossil fuels, increase the concentration of these gases in
the atmosphere, resulting in acid rain.
• Conductivity is a measure of a solution’s ability to conduct an electric current.
• An electrolyte is a substance that contains free ions, and is therefore able to conduct an
electric current. Electrolytes can be divided into strong and weak electrolytes, based on
the extent to which the substance ionises in solution.
• A non-electrolyte cannot conduct an electric current because it dooes not contain free
ions.
• The type of substance, the concentration of ions and the temperature of the solution,
affect its conductivity.
394
CHAPTER 20. THE HYDROSPHERE - GRADE 10
20.10
• A precipitate is formed when ions in solution react with each other to form an insoluble
product. Solubility ’rules’ help to identify the precipitate that has been formed.
• A number of tests can be used to identify whether certain anions are present in a solution.
• Despite the importance of the hydrosphere, a number of factors threaten it. These include
overuse of water, and pollution.
Exercise: Summary Exercise
1. Give one word for each of the following descriptions:
(a)
(b)
(c)
(d)
the change in phase of water from a gas to a liquid
a charged atom
a term used to describe the mineral content of water
a gas that forms sulfuric acid when it reacts with water
2. 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)
Column A
1. A polar molecule
2. molecular solution
3. Mineral that increases water hardness
4. Substance that increases the hydrogen ion concentration
5. A strong electrolyte
6. A white precipitate
7. A non-conductor of electricity
Column B
A. H2 SO4
B. CaCO3
C. NaOH
D. salt water
E. calcium
F. carbon dioxide
G. potassium nitrate
H. sugar water
I. O2
3. For each of the following questions, choose the one correct answer from the
list provided.
(a) Which one of the following substances does not conduct electricity in the
solid phase but is an electrical conductor when molten?
i. Cu
ii. PbBr2
iii. H2 O
iv. I2
(IEB Paper 2, 2003)
(b) The following substances are dissolved in water. Which one of the solutions
is basic?
i. sodium nitrate
ii. calcium sulphate
iii. ammonium chloride
iv. potassium carbonate
(IEB Paper 2, 2005)
4. The concentration of hydronium and hydroxyl ions in a typical sample of seawater are 10−8 and 10−6 respectively.
(a) Is the seawater acidic or basic?
(b) Calculate the pH of this seawater.
5. Three test tubes (X, Y and Z) each contain a solution of an unknown potassium
salt. The following observations were made during a practical investigation to
identify the solutions in the test tubes:
A: A white precipitate formed when silver nitrate (AgNO3 ) was added to test
tube Z.
395
20.10
CHAPTER 20. THE HYDROSPHERE - GRADE 10
B: A white precipitate formed in test tubes X and Y when barium chloride
(BaCl2 ) was added.
C: The precipitate in test tube X dissolved in hydrochloric acid (HCl) and a gas
was released.
D: The precipitate in test tube Y was insoluble in hydrochloric acid.
(a) Use the above information to identify the solutions in each of the test tubes
X, Y and Z.
(b) Write a chemical equation for the reaction that took place in test tube X
before hydrochloric acid was added.
(DoE Exemplar Paper 2 2007)
396
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