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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CONTENTS
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 15
Types of Reactions - Grade 11
There are many different types of chemical reactions that can take place. In this chapter, we will
be looking at a few of the more common reaction types: acid-base and acid-carbonate reactions,
redox reactions and addition, elimination and substitution reactions.
15.1
Acid-base reactions
15.1.1
What are acids and bases?
In our daily lives, we encounter many examples of acids and bases. In the home, vinegar (acetic
acid), lemon juice (citric acid) and tartaric acid (the main acid found in wine) are common, while
hydrochloric acid, sulfuric acid and nitric acid are examples of acids that are more likely to be
found in laboratories and industry. Hydrochloric acid is also found in the gastric juices in the
stomach. Even fizzy drinks contain acid (carbonic acid), as do tea and wine (tannic acid)! Bases
that you may have heard of include sodium hydroxide (caustic soda), ammonium hydroxide and
ammonia. Some of these are found in household cleaning products. Acids and bases are also
important commercial products in the fertiliser, plastics and petroleum refining industries. Some
common acids and bases, and their chemical formulae, are shown in table 15.1.
Table 15.1: Some common acids and
Acid
Formula
Hydrochoric acid
HCl
Sulfuric acid
H2 SO4
Nitric acid
HNO3
Acetic (ethanoic) acid CH3 COOH
Carbonic acid
H2 CO3
Sulfurous acid
H2 SO3
Phosphoric acid
H3 PO4
bases and their chemical
Base
Sodium hydroxide
Potassium hydroxide
Sodium carbonate
Calcium hydroxide
Magnesium hydroxide
Ammonia
Sodium bicarbonate
formulae
Formula
NaOH
KOH
Na2 CO3
Ca(OH)2
Mg(OH)2
NH3
NaHCO3
Most acids share certain characteristics, and most bases also share similar characteristics. It
is important to be able to have a definition for acids and bases so that they can be correctly
identified in reactions.
15.1.2
Defining acids and bases
A number of definitions for acids and bases have developed over the years. One of the earliest
was the Arrhenius definition. Arrhenius (1887) noticed that water dissociates (splits up) into
hydronium (H3 O+ ) and hydroxide (OH− ) ions according to the following equation:
H2 O ⇔ H3 O+ + OH−
267
15.1
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
Arrhenius described an acid as a compound that increases the concentration of H3 O+ ions
in solution, and a base as a compound that increases the concentration of OH− ions in a
solution. Look at the following examples showing the dissociation of hydrochloric acid and
sodium hydroxide (a base) respectively:
1. HCl + H2 O → H3 O+ + Cl−
Hydrochloric acid in water increases the concentration of H3 O+ ions and is therefore an
acid.
2. NaOH + H2 O → Na+ + OH−
Sodium hydroxide in water increases the concentration of OH− ions and is therefore a
base.
However, this definition could only be used for acids and bases in water. It was important to
come up with a much broader definition for acids and bases.
It was Lowry and Bronsted (1923) who took the work of Arrhenius further to develop a broader
definition for acids and bases. The Bronsted-Lowry model defines acids and bases in terms of
their ability to donate or accept protons.
Definition: Acids and bases
According to the Bronsted-Lowry theory of acids and bases, an acid is a substance that
gives away protons (H+ ), and is therefore called a proton donor. A base is a substance
that takes up protons, and is therefore called a proton acceptor.
Below are some examples:
1. HCl(g) + NH3 (g) → NH4+ + Cl−
In order to decide which substance is a proton donor and which is a proton acceptor, we
need to look at what happens to each reactant. The reaction can be broken down as
follows:
HCl → Cl− + H+ and
NH3 + H+ → NH+
4
From these reactions, it is clear that HCl is a proton donor and is therefore an acid, and
that NH3 is a proton acceptor and is therefore a base.
2. CH3 COOH + H2 O → H3 O+ + CH3 COO−
The reaction can be broken down as follows:
CH3 COOH → CH3 COO− + H+ and
H2 O + H+ → H3 O+
In this reaction, CH3 COOH (acetic acid) is a proton donor and is therefore the acid. In
this case, water acts as a base because it accepts a proton to form H3 O+ .
−
3. NH3 + H2 O → NH+
4 + OH
The reaction can be broken down as follows:
268
For
more
information
on dissociation, refer to
chapter 20.
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.1
H2 O → OH− + H+ and
NH3 + H+ → NH+
4
In this reaction, water donates a proton and is therefore an acid in this reaction. Ammonia
accepts the proton and is therefore the base. Notice that in the previous equation, water
acted as a base and that in this equation it acts as an acid. Water can act as both an
acid and a base depending on the reaction. This is also true of other substances. These
substances are called ampholytes and are said to be amphoteric.
Definition: Amphoteric
An amphoteric substance is one that can react as either an acid or base. Examples of
amphoteric substances include water, zinc oxide and beryllium hydroxide.
15.1.3
Conjugate acid-base pairs
Look at the reaction between hydrochloric acid and ammonia to form ammonium and chloride
ions:
−
HCl + NH3 ⇔ NH+
4 + Cl
Looking firstly at the forward reaction (i.e. the reaction that proceeds from left to right), the
changes that take place can be shown as follows:
HCl → Cl− + H+ and
NH3 + H+ → NH+
4
Looking at the reverse reaction (i.e. the reaction that proceeds from right to left), the changes
that take place are as follows:
+
NH+
4 → NH3 + H and
Cl− + H+ → HCl
In the forward reaction, HCl is a proton donor (acid) and NH3 is a proton acceptor (base).
In the reverse reaction, the chloride ion is the proton acceptor (base) and NH+
4 is the proton
donor (acid). A conjugate acid-base pair is two compounds in a reaction that change into
each other through the loss or gain of a proton. The conjugate acid-base pairs for the above
reaction are shown below.
conjugate pair
HCl + NH3
acid1 base2
−
NH+
4 + Cl
acid2 base1
conjugate pair
The reaction between ammonia and water can also be used as an example:
269
15.1
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
conjugate pair
−
NH+
4 + OH
acid2 base1
H2 O + NH3
acid1 base2
conjugate pair
Definition: Conjugate acid-base pair
The term refers to two compounds that transform into each other by the gain or loss of a
proton.
Exercise: Acids and bases
1. In the following reactions, identify (1) the acid and the base in the reactants
and (2) the salt in the product.
(a)
(b)
(c)
(d)
H2 SO4 + Ca(OH)2 → CaSO4 + 2H2 O
CuO + H2 SO4 → CuSO4 + H2 O
H2 O + C6 H5 OH → H3 O+ + C6 H5 O−
HBr + C5 H5 N → (C5 H5 NH+ )Br−
2. In each of the following reactions, label the conjugate acid-base pairs.
(a)
(b)
(c)
(d)
15.1.4
H2 SO4 + H2 O ⇔ H3 O+ + HSO−
4
−
NH+
+
F
⇔
HF
+
NH
3
4
H2 O + CH3 COO− ⇔ CH3 COOH + OH−
H2 SO4 + Cl− ⇔ HCl + HSO−
4
Acid-base reactions
When an acid and a base react, they neutralise each other to form a salt. If the base contains
hydroxide (OH− ) ions, then water will also be formed. The word salt is a general term which
applies to the products of all acid-base reactions. A salt is a product that is made up of the
cation from a base and the anion from an acid. When an acid reacts with a base, they neutralise
each other. In other words, the acid becomes less acidic and the base becomes less basic. Look
at the following examples:
1. Hydrochloric acid reacts with sodium hydroxide to form sodium chloride (the salt) and
water. Sodium chloride is made up of Na+ cations from the base (NaOH) and Cl− anions
from the acid (HCl).
HCl + NaOH → H2 O + NaCl
2. Hydrogen bromide reacts with potassium hydroxide to form potassium bromide (the salt)
and water. Potassium bromide is made up of K+ cations from the base (KOH) and Br−
anions from the acid (HBr).
HBr + KOH → H2 O + KBr
270
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.1
3. Hydrochloric acid reacts with sodium hydrocarbonate to form sodium chloride (the salt)
and hydrogen carbonate. Sodium chloride is made up of Na+ cations from the base
(NaHCO3 ) and Cl− anions from the acid (HCl).
HCl + NaHCO3 → H2 CO3 + NaCl
You should notice that in the first two examples, the base contained OH− ions, and therefore the
products were a salt and water. NaCl (table salt) and KBr are both salts. In the third example,
NaHCO3 also acts as a base, despite not having OH− ions. A salt is still formed as one of the
products, but no water is produced.
It is important to realise how important these neutralisation reactions are. Below are some
examples:
• Domestic uses
Calcium oxide (CaO) is put on soil that is too acid. Powdered limestone (CaCO3 ) can
also be used but its action is much slower and less effective. These substances can also
be used on a larger scale in farming and also in rivers.
• Biological uses
Acids in the stomach (e.g. hydrochloric acid) play an important role in helping to digest
food. However, when a person has a stomach ulcer, or when there is too much acid in the
stomach, these acids can cause a lot of pain. Antacids are taken to neutralise the acids so
that they don’t burn as much. Antacids are bases which neutralise the acid. Examples of
antacids are aluminium hydroxide, magnesium hydroxide (’milk of magnesia’) and sodium
bicarbonate (’bicarbonate of soda’). Antacids can also be used to relieve heartburn.
• Industrial uses
Alkaline calcium hydroxide (limewater) can be used to absorb harmful SO2 gas that is
released from power stations and from the burning of fossil fuels.
teresting Bee stings are acidic and have a pH between 5 and 5.5. They can be soothed
Interesting
Fact
Fact
by using substances such as calomine lotion, which is a mild alkali based on zinc
oxide. Bicarbonate of soda can also be used. Both alkalis help to neutralise the
acidic bee sting and relieve some of the itchiness!
271
15.1
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
Important: Acid-base titrations
The neutralisation reaction between an acid and a base can be very useful. If an acidic
solution of known concentration (a standard solution) is added to an alkaline solution until
the solution is exactly neutralised (i.e. it has neither acidic nor basic properties), it is
possible to calculate the exact concentration of the unknown solution. It is possible to do
this because, at the exact point where the solution is neutralised, chemically equivalent
amounts of acid and base have reacted with each other. This type of calculation is called
volumetric analysis. The process where an acid solution and a basic solution are added
to each other for this purpose, is called a titration, and the point of neutralisation is called
the end point of the reaction. So how exactly can a titration be carried out to determine
an unknown concentration? Look at the following steps to help you to understand the
process.
Step 1:
A measured volume of the solution with unknown concentration is put into a flask.
Step 2:
A suitable indicator is added to this solution (bromothymol blue and phenolpthalein are
common indicators).
Step 3:
A volume of the standard solution is put into a burette and is slowly added to the solution
in the flask, drop by drop.
Step 4:
At some point, adding one more drop will change the colour of the unknown solution.
For example, if the solution is basic and bromothymol blue is being used as the indicator in the titration, the bromothymol blue would originally have coloured the solution
blue. At the end point of the reaction, adding one more drop of acid will change the
colour of the basic solution from blue to yellow. Yellow shows that the solution is now acidic.
Step 5:
Record the volume of standard solution that has been added up to this point.
Step 6:
Use the information you have gathered to calculate the exact concentration of the unknown
solution. A worked example is shown below.
Important: Titration calculations
When you are busy with these calculations, you will need to remember the following:
1dm3 = 1 litre = 1000ml = 1000cm3, therefore dividing cm3 by 1000 will give you an answer in
dm3 .
Some other terms and equations which will be useful to remember are shown below:
• Molarity is a term used to describe the concentration of a solution, and is measured in
mol.dm−3 . The symbol for molarity is M. Refer to chapter 13 for more information on
molarity.
• Moles = molarity (mol.dm−3 ) x volume (dm3 )
• Molarity (mol.dm−3 ) = mol / volume
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CHAPTER 15. TYPES OF REACTIONS - GRADE 11
Worked Example 74: Titration calculation
Question: Given the equation:
NaOH + HCl → NaCl + H2 O
25cm3 of a sodium hydroxide solution was pipetted into a conical flask and titrated
with 0.2M hydrochloric acid. Using a suitable indicator, it was found that 15cm3
of acid was needed to neutralise the alkali. Calculate the molarity of the sodium
hydroxide.
Answer
Step 1 : Write down all the information you know about the reaction, and
make sure that the equation is balanced.
NaOH: V = 25 cm3
HCl: V = 15 cm3 ; C = 0.2 M
The equation is already balanced.
Step 2 : Calculate the number of moles of HCl that react according to this
equation.
n
V
Therefore, n(HCl) = M × V (make sure that all the units are correct!)
M=
M = 0.2mol.dm−3
V = 15cm3 = 0.015dm3
Therefore
n(HCl) = 0.2 × 0.015 = 0.003
There are 0.003 moles of HCl that react
Step 3 : Calculate the number of moles of sodium hydroxide in the reaction
Look at the equation for the reaction. For every mole of HCl there is one mole
of NaOH that is involved in the reaction. Therefore, if 0.003 moles of HCl react,
we can conclude that the same quantity of NaOH is needed for the reaction. The
number of moles of NaOH in the reaction is 0.003.
Step 4 : Calculate the molarity of the sodium hydroxide
First convert the volume into dm3 . V = 0.025 dm3 . Then continue with the
calculation.
M=
0.003
n
=
= 0.12
V
0.025
The molarity of the NaOH solution is 0.12 mol.dm3 or 0.12 M
Worked Example 75: Titration calculation
Question: 4.9 g of sulfuric acid is dissolved in water and the final solution has a
volume of 220 cm3 . Using titration, it was found that 20 cm3 of this solution was
273
15.1
15.1
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
able to completely neutralise 10 cm3 of a sodium hydroxide solution. Calculate the
concentration of the sodium hydroxide in mol.dm−3 .
Answer
Step 1 : Write a balanced equation for the titration reaction.
H2 SO4 + 2NaOH → Na2 SO4 + 2H2 O
Step 2 : Calculate the molarity of the sulfuric acid solution.
M = n/V
V = 220 cm3 = 0.22 dm3
n=
4.9g
m
= 0.05mols
=
M
98g.mol−1
Therefore,
M=
0.05
= 0.23mol.dm−3
0.22
Step 3 : Calculate the moles of sulfuric acid that were used in the neutralisation reaction.
Remember that only 20 cm3 of the sulfuric acid solution is used.
M = n/V, therefore n = M × V
n = 0.23 × 0.02 = 0.0046mol
Step 4 : Calculate the number of moles of sodium hydroxide that were
neutralised.
According to the balanced chemical equation, the mole ratio of H2 SO4 to NaOH is
1:2. Therefore, the number of moles of NaOH that are neutralised is 0.0046 × 2 =
0.0092 mols.
Step 5 : Calculate the concentration of the sodium hydroxide solution.
M=
15.1.5
0.0092
n
=
= 0.92M
V
0.01
Acid-carbonate reactions
Activity :: Demonstration : The reaction of acids with carbonates
Apparatus and materials:
Small amounts of sodium carbonate and calcium carbonate (both in powder
form); hydrochloric acid and sulfuric acid; retort stand; two test tubes; two rubber
stoppers for the test tubes; a delivery tube; lime water. The demonstration should
be set up as shown below.
274
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.1
delivery tube
rubber stopper
rubber stopper
[glassType=tube,bouchon=true,niveauLiquide1=30]
[glassType=tube,bouchon=true,niveauLiquide1=60]
sodium carbonate &
hydrochloric acid
limewater
Method:
1. Pour limewater into one of the test tubes and seal with a rubber stopper.
2. Pour a small amount of hydrochloric acid into the remaining test tube.
3. Add a small amount of sodium carbonate to the acid and seal the test tube
with the rubber stopper.
4. Connect the two test tubes with a delivery tube.
5. Observe what happens to the colour of the limewater.
6. Repeat the above steps, this time using sulfuric acid and calcium carbonate.
Observations:
The clear lime water turns milky meaning that carbon dioxide has been produced.
When an acid reacts with a carbonate a salt, carbon dioxide and water are formed. Look at the
following examples:
• Nitric acid reacts with sodium carbonate to form sodium nitrate, carbon dioxide and water.
2HNO3 + Na2 CO3 → 2NaNO3 + CO2 + H2 O
• Sulfuric acid reacts with calcium carbonate to form calcium sulphate, carbon dioxide and
water.
H2 SO4 + CaCO3 → CaSO4 + CO2 + H2 O
• Hydrochloric acid reacts with calcium carbonate to form calcium chloride, carbon dioxide
and water.
2HCl + CaCO3 → CaCl2 + CO2 + H2 O
Exercise: Acids and bases
1. The compound NaHCO3 is commonly known as baking soda. A recipe requires
1.6 g of baking soda, mixed with other ingredients, to bake a cake.
275
15.2
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
(a) Calculate the number of moles of NaHCO3 used to bake the cake.
(b) How many atoms of oxygen are there in the 1.6 g of baking soda?
During the baking process, baking soda reacts with an acid to produce
carbon dioxide and water, as shown by the reaction equation below:
+
HCO−
3 (aq) + H (aq) → CO2 (g) + H2 O(l)
(c) Identify the reactant which acts as the Bronsted-Lowry base in this reaction. Give a reason for your answer.
(d) Use the above equation to explain why the cake rises during this baking
process.
(DoE Grade 11 Paper 2, 2007)
2. Label the acid-base conjugate pairs in the following equation:
2−
+
HCO−
3 + H2 O ⇔ CO3 + H3 O
3. A certain antacid tablet contains 22.0 g of baking soda (NaHCO3 ). It is used to
neutralise the excess hydrochloric acid in the stomach. The balanced equation
for the reaction is:
NaHCO3 + HCl → NaCl + H2 O + CO2
The hydrochloric acid in the stomach has a concentration of 1.0 mol.dm−3 .
Calculate the volume of the hydrochloric acid that can be neutralised by the
antacid tablet.
(DoE Grade 11 Paper 2, 2007)
4. A learner is asked to prepare a standard solution of the weak acid, oxalic acid
(COOH)2 2H2 O for use in a titration. The volume of the solution must be 500
cm3 and the concentration 0.2 mol.dm−3 .
(a) Calculate the mass of oxalic acid which the learner has to dissolve to make
up the required standard solution. The leaner titrates this 0.2 mol.dm−3
oxalic acid solution against a solution of sodium hydroxide. He finds that
40 cm3 of the oxalic acid solution exactlt neutralises 35 cm3 of the sodium
hydroxide solution.
(b) Calculate the concentration of the sodium hydroxide solution.
5. A learner finds some sulfuric acid solution in a bottle labelled ’dilute sulfuric
acid’. He wants to determine the concentration of the sulphuric acid solution.
To do this, he decides to titrate the sulphuric acid against a standard potassium
hydroxide (KOH) solution.
(a) What is a standard solution?
(b) Calculate the mass of KOH which he must use to make 300 cm3 of a 0.2
mol.dm−3 KOH solution.
(c) Calculate the pH of the 0.2 mol.dm−3 KOH solution (assume standard
temperature).
(d) Write a balanced chemical equation for the reaction between H2 SO4 and
KOH.
(e) During the titration he finds that 15 cm3 of the KOH solution neutralises
20 cm3 of the H2 SO4 solution. Calculate the concentration of the H2 SO4
solution.
(IEB Paper 2, 2003)
15.2
Redox reactions
A second type of reaction is the redox reaction, in which both oxidation and reduction take
place.
276
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.2.1
15.2
Oxidation and reduction
If you look back to chapter 4, you will remember that we discussed how, during a chemical
reaction, an exchange of electrons takes place between the elements that are involved. Using
oxidation numbers is one way of tracking what is happening to these electrons in a reaction.
Refer back to section 4.11 if you can’t remember the rules that are used to give an oxidation
number to an element. Below are some examples to refresh your memory before we carry on
with this section!
Examples:
1. CO2
Each oxygen atom has an oxidation number of -2. This means that the charge on two
oxygen atoms is -4. We know that the molecule of CO2 is neutral, therefore the carbon
atom must have an oxidation number of +4.
2. KMnO4
Overall, this molecule has a neutral charge, meaning that the sum of the oxidation numbers of the elements in the molecule must equal zero. Potassium (K) has an oxidation
number of +1, while oxygen (O) has an oxidation number of -2. If we exclude the atom
of manganese (Mn), then the sum of the oxidation numbers equals +1+(-2x4)= -7. The
atom of manganese must therefore have an oxidation number of +7 in order to make the
molecule neutral.
By looking at how the oxidation number of an element changes during a reaction, we can easily
see whether that element is being oxidised or reduced.
Definition: Oxidation and reduction
Oxidation is the loses of an electron by a molecule, atom or ion. Reduction is the gain of
an electron by a molecule, atom or ion.
Example:
Mg + Cl2 → MgCl2
As a reactant, magnesium has an oxidation number of zero, but as part of the product magnesium
chloride, the element has an oxidation number of +2. Magnesium has lost two electrons and
has therefore been oxidised. This can be written as a half-reaction. The half-reaction for this
change is:
Mg → Mg2+ + 2e−
As a reactant, chlorine has an oxidation number of zero, but as part of the product magnesium
chloride, the element has an oxidation number of -1. Each chlorine atom has gained an electron
and the element has therefore been reduced. The half-reaction for this change is:
Cl2 + 2e− → 2Cl−
Definition: Half-reaction
A half reaction is either the oxidation or reduction reaction part of a redox reaction. A
half reaction is obtained by considering the change in oxidation states of the individual
substances that are involved in the redox reaction.
277
15.2
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
Important: Oxidation and reduction made easy!
An easy way to think about oxidation and reduction is to remember:
’OILRIG’ - Oxidation Is Loss of electrons, Reduction Is Gain of electrons.
An element that is oxidised is called a reducing agent, while an element that is reduced is
called an oxidising agent.
15.2.2
Redox reactions
Definition: Redox reaction
A redox reaction is one involving oxidation and reduction, where there is always a change
in the oxidation numbers of the elements involved.
Activity :: Demonstration : Redox reactions
Materials:
A few granules of zinc; 15 ml copper (II) sulphate solution (blue colour), glass
beaker.
zinc granules
copper sulphate
solution
Method:
Add the zinc granules to the copper sulphate solution and observe what happens.
What happens to the zinc granules? What happens to the colour of the solution?
Results:
• Zinc becomes covered in a layer that looks like copper.
• The blue copper sulphate solution becomes clearer.
Cu2+ ions from the CuSO4 solution are reduced to form copper metal. This is
what you saw on the zinc crystals. The reduction of the copper ions (in other words,
their removal from the copper sulphate solution), also explains the change in colour
of the solution. The equation for this reaction is:
Cu2+ + 2e− → Cu
Zinc is oxidised to form Zn2+ ions which are clear in the solution. The equation
for this reaction is:
Zn → Zn2+ + 2e−
The overall reaction is:
Cu2+ (aq) + Zn(s) → Cu(s) + Zn2+ (aq)
Conclusion:
A redox reaction has taken place. Cu2+ ions are reduced and the zinc is oxidised.
278
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.2
Below are some further examples of redox reactions:
• H2 + F2 → 2HF can be re-written as two half-reactions:
H2 → 2H+ + 2e− (oxidation) and
F2 + 2e− → 2F− (reduction)
• Cl2 + 2KI → 2KCl + I2 or Cl2 + 2I− → 2Cl− + I2 , can be written as two half-reactions:
Cl2 + 2e− → 2Cl− (reduction) and
2I− → I2 + 2e− (oxidation)
In Grade 12, you will go on to look at electrochemical reactions, and the role that electron
transfer plays in this type of reaction.
Exercise: Redox Reactions
1. Look at the following reaction:
2H2 O2 (l) → 2H2 O(l) + O2 (g)
(a) What is the oxidation number of the oxygen atom in each of the following
compounds?
i. H2 O2
ii. H2 O
iii. O2
(b) Does the hydrogen peroxide (H2 O2 ) act as an oxidising agent or a reducing
agent or both, in the above reaction? Give a reason for your answer.
2. Consider the following chemical equations:
1. Fe(s) → Fe2+ (aq) + 2e−
2. 4H+ (aq) + O2 (g) + 4e− → 2H2 O(l)
Which one of the following statements is correct?
(a)
(b)
(c)
(d)
Fe
Fe
Fe
Fe
is
is
is
is
oxidised and H+ is reduced
reduced and O2 is oxidised
oxidised and O2 is reduced
reduced and H+ is oxidised
(DoE Grade 11 Paper 2, 2007)
3. Which one of the following reactions is a redox reaction?
(a)
(b)
(c)
(d)
HCl + NaOH → NaCl + H2 O
AgNO3 + NaI → AgI + NaNO3
2FeCl3 + 2H2 O + SO2 → H2 SO4 + 2HCl + 2FeCl2
BaCl2 + MgSO4 → MgCl2 + BaSO4
279
15.3
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.3
Addition, substitution and elimination reactions
15.3.1
Addition reactions
An addition reaction occurs when two or more reactants combine to form a final product. This
product will contain all the atoms that were present in the reactants. The following is a general
equation for this type of reaction:
A+B → C
Notice that C is the final product with no A or B remaining as a residue.
The following are some examples.
1. The reaction between ethene and bromine to form 1,2-dibromoethane (figure 15.1).
C2 H4 + Br2 → C2 H4 Br2
H
H
C
H
+
C
Br
Br
H
H
H
H
C
C
Br
Br
H
Figure 15.1: The reaction between ethene and bromine is an example of an addition reaction
2. Polymerisation reactions
In industry, making polymers is very important. A polymer is made up of lots of smaller
units called monomers. When these monomers are added together, they form a polymer. Examples of polymers are polyvinylchloride (PVC) and polystyrene. PVC is often
used to make piping, while polystyrene is an important packaging and insulating material.
Polystyrene is made up of lots of styrene monomers which are joined through addition
reactions (figure 15.2). ’Polymerisation’ refers to the addition reactions that eventually
help to form the polystyrene polymer.
CH2
CH
CH2
CH
CH2
CH
CH2
CH
polymerisation
etc
Figure 15.2: The polymerisation of a styrene monomer to form a polystyrene polymer
280
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
15.3
3. The hydrogenation of vegetable oils to form margarine is another example of an addition
reaction. Hydrogenation involves adding hydrogen (H2 ) to an alkene. An alkene is an
organic compound composed of carbon and hydrogen. It contains a double bond between
two of the carbon atoms. If this bond is broken, it means that more hydrogen atoms can
attach themselves to the carbon atoms. During hydrogenation, this double bond is broken,
and more hydrogen atoms are added to the molecule. The reaction that takes place is
shown below. Note that the ’R’ represents any side-chain. A side-chain is simply any
combination of atoms that are attached to the central part of the molecule.
RCHCH2 + H2 → RCH2 CH3
4. The production of the alcohol ethanol from ethene. Ethanol (CH3 CH2 OH) can be made
from alkenes such as ethene (C2 H4 ), through a hydration reaction like the one below. A
hydration reaction is one where water is added to the reactants.
C2 H4 + H2 O → CH3 CH2 OH
A catalyst is needed for this reaction to take place. The catalyst that is most commonly
used is phosphoric acid.
15.3.2
Elimination reactions
An elimination reaction occurs when a reactant is broken up into two products. The general
form of the equation is as follows:
A→B+C
The examples below will help to explain this:
1. The dehydration of an alcohol is one example. Two hydrogen atoms and one oxygen
atom are eliminated and a molecule of water is formed as a second product in the reaction,
along with an alkene.
CH3 CH2 OH → CH2 CH2 + H2 O
H
H
H
C
C
H
OH
H
H
H
C
+
C
H
H
H
H
2. The elimination of potassium bromide from a bromoalkane.
CH3 CH2 Br + KOH → CH2 CH2 + KBr + H2 O
H
H
Br
C
C
H
H
H + KOH
281
H
H
C
C + KBr
H
H
O
+ H2 O
H
15.3
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
3. Ethane cracking is an important industrial process used by SASOL and other petrochemical
industries. Hydrogen is eliminated from ethane (C2 H6 ) to produce an alkene called ethene
(C2 H4 ). Ethene is then used to produce other products such as polyethylene. You will
learn more about these compounds in chapter 23. The equation for the cracking of ethane
looks like this:
C2 H6 → C2 H4 + H2
15.3.3
Substitution reactions
A substitution reaction occurs when an exchange of reactants takes place. The initial reactants
are transformed or ’swopped around’ to give a final product. A simple example of a reaction like
this is shown below:
AB + CD → AC + BD
Some simple examples of substitution reactions are shown below:
CH4 + Cl2 → CH3 Cl + HCl
In this example, a chorine atom and a hydrogen atom are exchanged to create a new product.
2−
−
Cu(H2 O)2+
4 + 4Cl ⇔ Cu(Cl)4 + 4H2 O
Exercise: Addition, substitution and elimination reactions
1. Refer to the diagram below and then answer the questions that follow:
H
Cl
H
C
C
(i)
H
H
H + KOH
H
H
C
C + KCl
H
H
+ H2 O
(a) Is this reaction an example of substitution, elimination or addition?
(b) Give a reason for your answer above.
(c) What type of compound is the reactant marked (i)?
2. The following diagram shows the reactants in an addition reaction.
H
H
C
C + HCl
H
H
(a) Draw the final product in this reaction.
(b) What is the chemical formula of the product?
3. The following reaction takes place:
282
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
H
H
OH
C
C
H
H
H
H2 SO4
H
H
C
C + H2 O
H
H
15.4
Is this reaction a substitution, addition or dehydration reaction? Give a reason
for your answer.
4. Consider the following reaction:
Ca(OH)2 (s) + 2NH4 Cl(s) → CaCl2 (s) + 2NH3 (g) + 2H2 O(g)
Which one of the following best describes the type of reaction which takes
place?
(a) Redox reaction
(b) Acid-base reaction
(c) Dehydration reaction
15.4
Summary
• There are many different types of reactions that can take place. These include acid-base,
acid-carbonate, redox, addition, substitution and elimination reactions.
• The Arrhenius definition of acids and bases defines an acid as a substance that increases
the concentration of hydrogen ions (H+ or H3 O+ ) in a solution. A base is a substance that
increases the concentration of hydroxide ions (OH− ) in a solution. However this definition
only applies to substances that are in water.
• The Bronsted-Lowry definition is a much broader one. An acid is a substance that
donates protons and a base is a substance that accepts protons.
• In different reactions, certain substances can act as both an acid and a base. These
substances are called ampholytes and are said to be amphoteric. Water is an example
of an amphoteric substance.
• A conjugate acid-base pair refers to two compounds in a reaction that change into other
through the loss or gain of a proton.
• When an acid and a base react, they form a salt and water. The salt is made up of a cation
from the base and an anion from the acid. An example of a salt is sodium chloride (NaCl),
which is the product of the reaction between sodium hydroxide (NaOH) and hydrochloric
acid (HCl).
• The reaction between an acid and a base is a neutralisation reaction.
• Titrations are reactions between an acid and a base that are used to calculate the concentration of one of the reacting substances. The concentration of the other reacting
substance must be known.
• In an acid-carbonate reaction, an acid and a carbonate react to form a salt, carbon
dioxide and water.
• A redox reaction is one where there is always a change in the oxidation numbers of the
elements that are involved in the reaction.
• Oxidation is the loss of electrons and reduction is the gain of electrons.
283
15.4
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
• When two or more reactants combine to form a product that contains all the atoms
that were in the reactants, then this is an addition reaction. Examples of addition
reactions include the reaction between ethene and bromine, polymerisation reactions and
hydrogenation reactions.
• A reaction where the reactant is broken down into one or more product, is called an elimination reaction. Alcohol dehydration and ethane cracking are examples of elimination
reactions.
• A substitution reaction is one where the reactants are transformed or swopped around
to form the final product.
Exercise: Summary Exercise
1. Give one word/term for each of the following descriptions:
(a) A chemical reaction during which electrons are transferred
(b) The addition of hydrogen across a double bond
(c) The removal of hydrogen and halogen atoms from an alkane to form an
elkene
2. For each of the following, say whether the statement is true or false. If the
statement is false, re-write the statement correctly.
(a) The conjugate base of NH+
4 is NH3 .
(b) The reactions C + O2 → CO2 and 2KClO3 → 2KCl + 3O2 are examples
of redox reactions.
3. For each of the following questions, choose the one correct statement from the
list provided.
A The following chemical equation represents the formation of the hydronium
ion:
H+ (aq) + H2 O(l) → H3 O+ (aq)
In this reaction, water acts as a Lewis base because it...
i. accepts protons
ii. donates protons
iii. accepts electrons
iv. donates electrons
B What is the concentration (in mol.dm−3 ) of H3 O+ ions in a NaOH solution
which has a pH of 12 at 250 C?
i. 1 × 1012
ii. 1 × 102
iii. 1 × 10−2
iv. 1 × 10−12
(IEB Paper 2, 2005)
C When chlorine water is added to a solution of potassium bromide, bromine
is produced. Which one of the following statements concerning this reaction is correct?
i. Br− is oxidised
ii. Cl2 is oxidised
iii. Br− is the oxidising agent
iv. Cl− is the oxidising agent
(IEB Paper 2, 2005)
4. The stomach secretes gastric juice, which contains hydrochloric acid. The
gastric juice helps with digestion. Sometimes there is an overproduction of
acid, leading to heartburn or indigestion. Antacids, such as milk of magnesia,
can be taken to neutralise the excess acid. Milk of magnesia is only slightly
soluble in water and has the chemical formula Mg(OH)2 .
284
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
a Write a balanced chemical equation to show how the antacid reacts with
the acid.
b The directions on the bottle recommend that children under the age of 12
years take one teaspoon of milk of magnesia, whereas adults can take two
teaspoons of the antacid. Briefly explain why the dosages are different.
c Why is it not advisable to take an overdose of the antacid in the stomach?
Refer to the hydrochloric acid concentration in the stomach in your answer.
In an experiment, 25.0 cm3 of a standard solution of sodium carbonate of
concentration 0.1 mol.dm−3 was used to neutralise 35.0 cm3 of a solution
of hydrochloric acid.
d Write a balanced chemical equation for the reaction.
e Calculate the concentration of the acid.
(DoE Grade 11 Exemplar, 2007)
285
15.4
15.4
CHAPTER 15. TYPES OF REACTIONS - GRADE 11
286
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