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

The Free High School Science Texts: Textbooks for High School Students Chemistry
FHSST Authors
The Free High School Science Texts:
Textbooks for High School Students
Studying the Sciences
Chemistry
Grades 10 - 12
Version 0
November 9, 2008
ii
Copyright 2007 “Free High School Science Texts”
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FHSST Core Team
Mark Horner ; Samuel Halliday ; Sarah Blyth ; Rory Adams ; Spencer Wheaton
FHSST Editors
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Whitfield
FHSST Contributors
Rory Adams ; Prashant Arora ; Richard Baxter ; Dr. Sarah Blyth ; Sebastian Bodenstein ;
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Daniels ; Sean Dobbs ; Fernando Durrell ; Dr. Dan Dwyer ; Frans van Eeden ; Giovanni
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Andrew Kubik ; Dr. Marco van Leeuwen ; Dr. Anton Machacek ; Dr. Komal Maheshwari ;
Kosma von Maltitz ; Nicole Masureik ; John Mathew ; JoEllen McBride ; Nikolai Meures ;
Riana Meyer ; Jenny Miller ; Abdul Mirza ; Asogan Moodaly ; Jothi Moodley ; Nolene Naidu ;
Tyrone Negus ; Thomas O’Donnell ; Dr. Markus Oldenburg ; Dr. Jaynie Padayachee ;
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iii
iv
Contents
I
II
Introduction
1
Matter and Materials
3
1 Classification of Matter - Grade 10
1.1
1.2
5
Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5
1.1.1
Heterogeneous mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . .
6
1.1.2
Homogeneous mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . .
6
1.1.3
Separating mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7
Pure Substances: Elements and Compounds . . . . . . . . . . . . . . . . . . . .
9
1.2.1
Elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
9
1.2.2
Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
9
1.3
Giving names and formulae to substances . . . . . . . . . . . . . . . . . . . . . 10
1.4
Metals, Semi-metals and Non-metals . . . . . . . . . . . . . . . . . . . . . . . . 13
1.4.1
Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13
1.4.2
Non-metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14
1.4.3
Semi-metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14
1.5
Electrical conductors, semi-conductors and insulators . . . . . . . . . . . . . . . 14
1.6
Thermal Conductors and Insulators . . . . . . . . . . . . . . . . . . . . . . . . . 15
1.7
Magnetic and Non-magnetic Materials . . . . . . . . . . . . . . . . . . . . . . . 17
1.8
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18
2 What are the objects around us made of? - Grade 10
21
2.1
Introduction: The atom as the building block of matter . . . . . . . . . . . . . . 21
2.2
Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21
2.2.1
Representing molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . 21
2.3
Intramolecular and intermolecular forces . . . . . . . . . . . . . . . . . . . . . . 25
2.4
The Kinetic Theory of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26
2.5
The Properties of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28
2.6
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31
3 The Atom - Grade 10
3.1
35
Models of the Atom . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35
3.1.1
The Plum Pudding Model . . . . . . . . . . . . . . . . . . . . . . . . . . 35
3.1.2
Rutherford’s model of the atom
v
. . . . . . . . . . . . . . . . . . . . . . 36
CONTENTS
3.1.3
3.2
3.3
CONTENTS
The Bohr Model . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37
How big is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.2.1
How heavy is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.2.2
How big is an atom? . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
Atomic structure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 38
3.3.1
The Electron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39
3.3.2
The Nucleus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39
3.4
Atomic number and atomic mass number . . . . . . . . . . . . . . . . . . . . . 40
3.5
Isotopes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42
3.6
3.7
3.8
3.9
3.5.1
What is an isotope? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42
3.5.2
Relative atomic mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . 45
Energy quantisation and electron configuration . . . . . . . . . . . . . . . . . . 46
3.6.1
The energy of electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . 46
3.6.2
Energy quantisation and line emission spectra . . . . . . . . . . . . . . . 47
3.6.3
Electron configuration . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47
3.6.4
Core and valence electrons . . . . . . . . . . . . . . . . . . . . . . . . . 51
3.6.5
The importance of understanding electron configuration . . . . . . . . . 51
Ionisation Energy and the Periodic Table . . . . . . . . . . . . . . . . . . . . . . 53
3.7.1
Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 53
3.7.2
Ionisation Energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 55
The Arrangement of Atoms in the Periodic Table . . . . . . . . . . . . . . . . . 56
3.8.1
Groups in the periodic table
. . . . . . . . . . . . . . . . . . . . . . . . 56
3.8.2
Periods in the periodic table . . . . . . . . . . . . . . . . . . . . . . . . 58
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59
4 Atomic Combinations - Grade 11
63
4.1
Why do atoms bond? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63
4.2
Energy and bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63
4.3
What happens when atoms bond? . . . . . . . . . . . . . . . . . . . . . . . . . 65
4.4
Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65
4.4.1
The nature of the covalent bond . . . . . . . . . . . . . . . . . . . . . . 65
4.5
Lewis notation and molecular structure . . . . . . . . . . . . . . . . . . . . . . . 69
4.6
Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72
4.7
4.8
4.6.1
Non-polar and polar covalent bonds . . . . . . . . . . . . . . . . . . . . 73
4.6.2
Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73
Ionic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74
4.7.1
The nature of the ionic bond . . . . . . . . . . . . . . . . . . . . . . . . 74
4.7.2
The crystal lattice structure of ionic compounds . . . . . . . . . . . . . . 76
4.7.3
Properties of Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . 76
Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 76
4.8.1
The nature of the metallic bond . . . . . . . . . . . . . . . . . . . . . . 76
4.8.2
The properties of metals . . . . . . . . . . . . . . . . . . . . . . . . . . 77
vi
CONTENTS
4.9
CONTENTS
Writing chemical formulae
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 78
4.9.1
The formulae of covalent compounds . . . . . . . . . . . . . . . . . . . . 78
4.9.2
The formulae of ionic compounds . . . . . . . . . . . . . . . . . . . . . 80
4.10 The Shape of Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 82
4.10.1 Valence Shell Electron Pair Repulsion (VSEPR) theory . . . . . . . . . . 82
4.10.2 Determining the shape of a molecule . . . . . . . . . . . . . . . . . . . . 82
4.11 Oxidation numbers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85
4.12 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 88
5 Intermolecular Forces - Grade 11
91
5.1
Types of Intermolecular Forces . . . . . . . . . . . . . . . . . . . . . . . . . . . 91
5.2
Understanding intermolecular forces . . . . . . . . . . . . . . . . . . . . . . . . 94
5.3
Intermolecular forces in liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . 96
5.4
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97
6 Solutions and solubility - Grade 11
101
6.1
Types of solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 101
6.2
Forces and solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102
6.3
Solubility . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103
6.4
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106
7 Atomic Nuclei - Grade 11
107
7.1
Nuclear structure and stability . . . . . . . . . . . . . . . . . . . . . . . . . . . 107
7.2
The Discovery of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 107
7.3
Radioactivity and Types of Radiation . . . . . . . . . . . . . . . . . . . . . . . . 108
7.4
7.3.1
Alpha (α) particles and alpha decay . . . . . . . . . . . . . . . . . . . . 109
7.3.2
Beta (β) particles and beta decay . . . . . . . . . . . . . . . . . . . . . 109
7.3.3
Gamma (γ) rays and gamma decay . . . . . . . . . . . . . . . . . . . . . 110
Sources of radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112
7.4.1
Natural background radiation . . . . . . . . . . . . . . . . . . . . . . . . 112
7.4.2
Man-made sources of radiation . . . . . . . . . . . . . . . . . . . . . . . 113
7.5
The ’half-life’ of an element . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113
7.6
The Dangers of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116
7.7
The Uses of Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117
7.8
Nuclear Fission
7.9
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118
7.8.1
The Atomic bomb - an abuse of nuclear fission . . . . . . . . . . . . . . 119
7.8.2
Nuclear power - harnessing energy . . . . . . . . . . . . . . . . . . . . . 120
Nuclear Fusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 120
7.10 Nucleosynthesis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121
7.10.1 Age of Nucleosynthesis (225 s - 103 s) . . . . . . . . . . . . . . . . . . . 121
7.10.2 Age of Ions (103 s - 1013 s) . . . . . . . . . . . . . . . . . . . . . . . . . 122
7.10.3 Age of Atoms (1013 s - 1015 s) . . . . . . . . . . . . . . . . . . . . . . . 122
7.10.4 Age of Stars and Galaxies (the universe today) . . . . . . . . . . . . . . 122
7.11 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122
vii
CONTENTS
CONTENTS
8 Thermal Properties and Ideal Gases - Grade 11
125
8.1
A review of the kinetic theory of matter . . . . . . . . . . . . . . . . . . . . . . 125
8.2
Boyle’s Law: Pressure and volume of an enclosed gas . . . . . . . . . . . . . . . 126
8.3
Charles’s Law: Volume and Temperature of an enclosed gas . . . . . . . . . . . 132
8.4
The relationship between temperature and pressure . . . . . . . . . . . . . . . . 136
8.5
The general gas equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 137
8.6
The ideal gas equation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 140
8.7
Molar volume of gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 145
8.8
Ideal gases and non-ideal gas behaviour . . . . . . . . . . . . . . . . . . . . . . 146
8.9
Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 147
9 Organic Molecules - Grade 12
151
9.1
What is organic chemistry? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151
9.2
Sources of carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151
9.3
Unique properties of carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4
Representing organic compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4.1
Molecular formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 152
9.4.2
Structural formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 153
9.4.3
Condensed structural formula . . . . . . . . . . . . . . . . . . . . . . . . 153
9.5
Isomerism in organic compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 154
9.6
Functional groups . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 155
9.7
The Hydrocarbons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 155
9.7.1
The Alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 158
9.7.2
Naming the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 159
9.7.3
Properties of the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . 163
9.7.4
Reactions of the alkanes . . . . . . . . . . . . . . . . . . . . . . . . . . 163
9.7.5
The alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166
9.7.6
Naming the alkenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166
9.7.7
The properties of the alkenes . . . . . . . . . . . . . . . . . . . . . . . . 169
9.7.8
Reactions of the alkenes
9.7.9
The Alkynes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171
. . . . . . . . . . . . . . . . . . . . . . . . . . 169
9.7.10 Naming the alkynes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171
9.8
9.9
The Alcohols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172
9.8.1
Naming the alcohols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 173
9.8.2
Physical and chemical properties of the alcohols . . . . . . . . . . . . . . 175
Carboxylic Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 176
9.9.1
Physical Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177
9.9.2
Derivatives of carboxylic acids: The esters . . . . . . . . . . . . . . . . . 178
9.10 The Amino Group . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178
9.11 The Carbonyl Group . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178
9.12 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 179
viii
CONTENTS
CONTENTS
10 Organic Macromolecules - Grade 12
185
10.1 Polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 185
10.2 How do polymers form? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 186
10.2.1 Addition polymerisation . . . . . . . . . . . . . . . . . . . . . . . . . . . 186
10.2.2 Condensation polymerisation . . . . . . . . . . . . . . . . . . . . . . . . 188
10.3 The chemical properties of polymers . . . . . . . . . . . . . . . . . . . . . . . . 190
10.4 Types of polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191
10.5 Plastics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191
10.5.1 The uses of plastics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 192
10.5.2 Thermoplastics and thermosetting plastics . . . . . . . . . . . . . . . . . 194
10.5.3 Plastics and the environment . . . . . . . . . . . . . . . . . . . . . . . . 195
10.6 Biological Macromolecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 196
10.6.1 Carbohydrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 197
10.6.2 Proteins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 199
10.6.3 Nucleic Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 202
10.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 204
III
Chemical Change
209
11 Physical and Chemical Change - Grade 10
211
11.1 Physical changes in matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 211
11.2 Chemical Changes in Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212
11.2.1 Decomposition reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 213
11.2.2 Synthesis reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 214
11.3 Energy changes in chemical reactions . . . . . . . . . . . . . . . . . . . . . . . . 217
11.4 Conservation of atoms and mass in reactions . . . . . . . . . . . . . . . . . . . . 217
11.5 Law of constant composition . . . . . . . . . . . . . . . . . . . . . . . . . . . . 219
11.6 Volume relationships in gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . 219
11.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 220
12 Representing Chemical Change - Grade 10
223
12.1 Chemical symbols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 223
12.2 Writing chemical formulae
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224
12.3 Balancing chemical equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224
12.3.1 The law of conservation of mass . . . . . . . . . . . . . . . . . . . . . . 224
12.3.2 Steps to balance a chemical equation
. . . . . . . . . . . . . . . . . . . 226
12.4 State symbols and other information . . . . . . . . . . . . . . . . . . . . . . . . 230
12.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 232
13 Quantitative Aspects of Chemical Change - Grade 11
233
13.1 The Mole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 233
13.2 Molar Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 235
13.3 An equation to calculate moles and mass in chemical reactions . . . . . . . . . . 237
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CONTENTS
13.4 Molecules and compounds
CONTENTS
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 239
13.5 The Composition of Substances . . . . . . . . . . . . . . . . . . . . . . . . . . . 242
13.6 Molar Volumes of Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 246
13.7 Molar concentrations in liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . 247
13.8 Stoichiometric calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 249
13.9 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252
14 Energy Changes In Chemical Reactions - Grade 11
255
14.1 What causes the energy changes in chemical reactions? . . . . . . . . . . . . . . 255
14.2 Exothermic and endothermic reactions . . . . . . . . . . . . . . . . . . . . . . . 255
14.3 The heat of reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 257
14.4 Examples of endothermic and exothermic reactions . . . . . . . . . . . . . . . . 259
14.5 Spontaneous and non-spontaneous reactions . . . . . . . . . . . . . . . . . . . . 260
14.6 Activation energy and the activated complex . . . . . . . . . . . . . . . . . . . . 261
14.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 264
15 Types of Reactions - Grade 11
267
15.1 Acid-base reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.1 What are acids and bases? . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.2 Defining acids and bases . . . . . . . . . . . . . . . . . . . . . . . . . . 267
15.1.3 Conjugate acid-base pairs . . . . . . . . . . . . . . . . . . . . . . . . . . 269
15.1.4 Acid-base reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 270
15.1.5 Acid-carbonate reactions . . . . . . . . . . . . . . . . . . . . . . . . . . 274
15.2 Redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 276
15.2.1 Oxidation and reduction
. . . . . . . . . . . . . . . . . . . . . . . . . . 277
15.2.2 Redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 278
15.3 Addition, substitution and elimination reactions . . . . . . . . . . . . . . . . . . 280
15.3.1 Addition reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 280
15.3.2 Elimination reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 281
15.3.3 Substitution reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 282
15.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 283
16 Reaction Rates - Grade 12
287
16.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 287
16.2 Factors affecting reaction rates . . . . . . . . . . . . . . . . . . . . . . . . . . . 289
16.3 Reaction rates and collision theory . . . . . . . . . . . . . . . . . . . . . . . . . 293
16.4 Measuring Rates of Reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . 295
16.5 Mechanism of reaction and catalysis . . . . . . . . . . . . . . . . . . . . . . . . 297
16.6 Chemical equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 300
16.6.1 Open and closed systems . . . . . . . . . . . . . . . . . . . . . . . . . . 302
16.6.2 Reversible reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 302
16.6.3 Chemical equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . 303
16.7 The equilibrium constant . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 304
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CONTENTS
16.7.1 Calculating the equilibrium constant . . . . . . . . . . . . . . . . . . . . 305
16.7.2 The meaning of kc values . . . . . . . . . . . . . . . . . . . . . . . . . . 306
16.8 Le Chatelier’s principle . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 310
16.8.1 The effect of concentration on equilibrium . . . . . . . . . . . . . . . . . 310
16.8.2 The effect of temperature on equilibrium . . . . . . . . . . . . . . . . . . 310
16.8.3 The effect of pressure on equilibrium . . . . . . . . . . . . . . . . . . . . 312
16.9 Industrial applications . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 315
16.10Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 316
17 Electrochemical Reactions - Grade 12
319
17.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 319
17.2 The Galvanic Cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 320
17.2.1 Half-cell reactions in the Zn-Cu cell . . . . . . . . . . . . . . . . . . . . 321
17.2.2 Components of the Zn-Cu cell . . . . . . . . . . . . . . . . . . . . . . . 322
17.2.3 The Galvanic cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 323
17.2.4 Uses and applications of the galvanic cell . . . . . . . . . . . . . . . . . 324
17.3 The Electrolytic cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 325
17.3.1 The electrolysis of copper sulphate . . . . . . . . . . . . . . . . . . . . . 326
17.3.2 The electrolysis of water . . . . . . . . . . . . . . . . . . . . . . . . . . 327
17.3.3 A comparison of galvanic and electrolytic cells . . . . . . . . . . . . . . . 328
17.4 Standard Electrode Potentials . . . . . . . . . . . . . . . . . . . . . . . . . . . . 328
17.4.1 The different reactivities of metals . . . . . . . . . . . . . . . . . . . . . 329
17.4.2 Equilibrium reactions in half cells . . . . . . . . . . . . . . . . . . . . . . 329
17.4.3 Measuring electrode potential . . . . . . . . . . . . . . . . . . . . . . . . 330
17.4.4 The standard hydrogen electrode . . . . . . . . . . . . . . . . . . . . . . 330
17.4.5 Standard electrode potentials . . . . . . . . . . . . . . . . . . . . . . . . 333
17.4.6 Combining half cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 337
17.4.7 Uses of standard electrode potential . . . . . . . . . . . . . . . . . . . . 338
17.5 Balancing redox reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342
17.6 Applications of electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . 347
17.6.1 Electroplating . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 347
17.6.2 The production of chlorine . . . . . . . . . . . . . . . . . . . . . . . . . 348
17.6.3 Extraction of aluminium
. . . . . . . . . . . . . . . . . . . . . . . . . . 349
17.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 349
IV
Chemical Systems
353
18 The Water Cycle - Grade 10
355
18.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355
18.2 The importance of water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355
18.3 The movement of water through the water cycle . . . . . . . . . . . . . . . . . . 356
18.4 The microscopic structure of water . . . . . . . . . . . . . . . . . . . . . . . . . 359
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18.4.1 The polar nature of water . . . . . . . . . . . . . . . . . . . . . . . . . . 359
18.4.2 Hydrogen bonding in water molecules . . . . . . . . . . . . . . . . . . . 359
18.5 The unique properties of water . . . . . . . . . . . . . . . . . . . . . . . . . . . 360
18.6 Water conservation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 363
18.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 366
19 Global Cycles: The Nitrogen Cycle - Grade 10
369
19.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 369
19.2 Nitrogen fixation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 369
19.3 Nitrification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 371
19.4 Denitrification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372
19.5 Human Influences on the Nitrogen Cycle . . . . . . . . . . . . . . . . . . . . . . 372
19.6 The industrial fixation of nitrogen . . . . . . . . . . . . . . . . . . . . . . . . . 373
19.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 374
20 The Hydrosphere - Grade 10
377
20.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 377
20.2 Interactions of the hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . 377
20.3 Exploring the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 378
20.4 The Importance of the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . 379
20.5 Ions in aqueous solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 379
20.5.1 Dissociation in water . . . . . . . . . . . . . . . . . . . . . . . . . . . . 380
20.5.2 Ions and water hardness . . . . . . . . . . . . . . . . . . . . . . . . . . . 382
20.5.3 The pH scale . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 382
20.5.4 Acid rain . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 384
20.6 Electrolytes, ionisation and conductivity . . . . . . . . . . . . . . . . . . . . . . 386
20.6.1 Electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 386
20.6.2 Non-electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 387
20.6.3 Factors that affect the conductivity of water . . . . . . . . . . . . . . . . 387
20.7 Precipitation reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 389
20.8 Testing for common anions in solution . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.1 Test for a chloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.2 Test for a sulphate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 391
20.8.3 Test for a carbonate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 392
20.8.4 Test for bromides and iodides . . . . . . . . . . . . . . . . . . . . . . . . 392
20.9 Threats to the Hydrosphere . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 393
20.10Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 394
21 The Lithosphere - Grade 11
397
21.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 397
21.2 The chemistry of the earth’s crust . . . . . . . . . . . . . . . . . . . . . . . . . 398
21.3 A brief history of mineral use . . . . . . . . . . . . . . . . . . . . . . . . . . . . 399
21.4 Energy resources and their uses . . . . . . . . . . . . . . . . . . . . . . . . . . . 400
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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 17
Electrochemical Reactions - Grade
12
17.1
Introduction
Chapter 15 in Grade 11 discussed oxidation, reduction and redox reactions. Oxidation involves
a loss of electrons and reduction involves a gain of electrons. A redox reaction is a reaction
where both oxidation and reduction take place. What is common to all of these processes is that
they involve a transfer of electrons and a change in the oxidation state of the elements that are
involved.
Exercise: Oxidation and reduction
1. Define the terms oxidation and reduction.
2. In each of the following reactions say whether the iron in the reactants is
oxidised or reduced.
(a)
(b)
(c)
(d)
(e)
F e → F e2+ + 2e−
F e3+ + e− → F e2+
F e2 O3 → F e
F e2+ → F e3+ + e−
F e2 O3 + 2Al → Al2
3. In each of the following equations, say which elements in the reactants are
oxidised and which are reduced.
(a)
(b)
(c)
(d)
CuO(s) + H2 (g) → Cu(s) + H2 O(g)
2N O(g) + 2CO(g) → N2 (g) + 2CO2 (g)
M g(s) + F eSO4 (aq) → M gSO4 (aq) + F e(s)
Zn(s) + 2AgN O3 (aq) → 2Ag + Zn(N O3 )2 (aq)
4. Which one of the substances listed below acts as the oxidising agent in the
following reaction?
3SO2 + Cr2 O72− + 2H + → 3SO42− + 2Cr3+ + H2 O
(a)
(b)
(c)
(d)
H+
Cr3+
SO2
Cr2 O2−
7
319
17.2
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
In Grade 11, an experiment was carried out to see what happened when zinc granules are added
to a solution of copper(II) sulphate. In the experiment, the Cu2+ ions from the copper(II)
sulphate solution were reduced to copper metal, which was then deposited in a layer on the zinc
granules. The zinc atoms were oxidised to form Zn2+ ions in the solution. The half reactions
are as follows:
Cu2+ (aq) + 2e− → Cu(s) (reduction half reaction)
Zn(s) → Zn2+ (aq) + 2e− (oxidation half reaction)
The overall redox reaction is:
Cu2+ (aq) + Zn → Cu(s) + Zn2+ (aq)
There was an increase in the temperature of the reaction when you carried out this experiment.
Is it possible that this heat energy could be converted into electrical energy? In other words, can
we use a chemical reaction where there is an exchange of electrons, to produce electricity? And
if this is possible, what would happen if an electrical current was supplied to cause some type
of chemical reaction to take place?
An electrochemical reaction is a chemical reaction that produces a voltage, and therefore a
flow of electrical current. An electrochemical reaction can also be the reverse of this process, in
other words if an electrical current causes a chemical reaction to take place.
Definition: Electrochemical reaction
If a chemical reaction is caused by an external voltage, or if a voltage is caused by a chemical
reaction, it is an electrochemical reaction.
Electrochemistry is the branch of chemistry that studies these electrochemical reactions. In
this chapter, we will be looking more closely at different types of electrochemical reactions, and
how these can be used in different ways.
17.2
The Galvanic Cell
Activity :: Experiment : Electrochemical reactions
Aim:
To investigate the reactions that take place in a zinc-copper cell
Apparatus:
zinc plate, copper plate, measuring balance, zinc sulphate (ZnSO4 ) solution (1
mol.dm−3 ), copper sulphate (CuSO4 ) solution (1 mol.dm−3 ), two 250 ml beakers,
U-tube, Na2 SO4 solution, cotton wool, ammeter, connecting wire.
Method:
1. Measure the mass of the copper and zinc plates and record your findings.
2. Pour about 200 ml of the zinc sulphate solution into a beaker and put the zinc
plate into it.
3. Pour about 200 ml of the copper sulphate solution into the second beaker and
place the copper plate into it.
4. Fill the U-tube with the Na2 SO4 solution and seal the ends of the tubes with
the cotton wool. This will stop the solution from flowing out when the U-tube
is turned upside down.
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17.2
5. Connect the zinc and copper plates to the ammeter and observe whether the
ammeter records a reading.
6. Place the U-tube so that one end is in the copper sulphate solution and the
other end is in the zinc sulphate solution. Is there a reading on the ammeter?
In which direction is the current flowing?
7. Take the ammeter away and connect the copper and zinc plates to each other
directly using copper wire. Leave to stand for about one day.
8. After a day, remove the two plates and rinse them first with distilled water,
then with alcohol and finally with ether. Dry the plates using a hair dryer.
9. Weigh the zinc and copper plates and record their mass. Has the mass of the
plates changed from the original measurements?
Note: A voltmeter can also be used in place of the ammeter. A voltmeter will
measure the potential difference across the cell.
electron flow
A
+
+
Cu
-
salt
bridge
CuSO4(aq)
Zn
ZnSO4(aq)
Results:
During the experiment, you should have noticed the following:
• When the U-tube containing the Na2 SO4 solution was absent, there was no
reading on the ammeter.
• When the U-tube was connected, a reading was recorded on the ammeter.
• After the plates had been connected directly to each other and left for a day,
there was a change in their mass. The mass of the zinc plate decreased, while
the mass of the copper plate increased.
• The direction of electron flow is from the zinc plate towards the copper plate.
Conclusions:
When a zinc sulphate solution containing a zinc plate is connected by a U-tube
to a copper sulphate solution containing a copper plate, reactions occur in both
solutions. The decrease in mass of the zinc plate suggests that the zinc metal has
been oxidised. The increase in mass of the copper plate suggests that reduction has
occurred here to produce more copper metal. This will be explained in detail below.
17.2.1
Half-cell reactions in the Zn-Cu cell
The experiment above demonstrated a zinc-copper cell. This was made up of a zinc half cell
and a copper half cell.
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Definition: Half cell
A half cell is a structure that consists of a conductive electrode surrounded by a conductive
electrolyte. For example, a zinc half cell could consist of a zinc metal plate (the electrode)
in a zinc sulphate solution (the electrolyte).
How do we explain what has just been observed in the zinc-copper cell?
• Copper plate
At the copper plate, there was an increase in mass. This means that Cu2+ ions from the
copper sulphate solution were deposited onto the plate as atoms of copper metal. The
half-reaction that takes place at the copper plate is:
Cu2+ + 2e− → Cu (Reduction half reaction)
Another shortened way to represent this copper half-cell is Cu2+ /Cu.
• Zinc plate
At the zinc plate, there was a decrease in mass. This means that some of the zinc goes
into solution as Z2+ ions. The electrons remain on the zinc plate, giving it a negative
charge. The half-reaction that takes place at the zinc plate is:
Zn → Zn2+ + 2e− (Oxidation half reaction)
The shortened way to represent the zinc half-cell is Zn/Zn2+ .
The overall reaction is:
Zn + Cu2+ + 2e− → Zn2+ + Cu + 2e− or, if we cancel the electrons:
Zn + Cu2+ → Zn2+ + Cu
For this electrochemical cell, the standard notation is:
Zn|Zn2+ ||Cu2+ |Cu
where
| =
|| =
a phase boundary (solid/aqueous)
the salt bridge
In the notation used above, the oxidation half-reaction at the anode is written on the left, and
the reduction half-reaction at the cathode is written on the right. In the Zn-Cu electrochemical
cell, the direction of current flow in the external circuit is from the zinc electrode (where there
has been a build up of electrons) to the copper electrode.
17.2.2
Components of the Zn-Cu cell
In the zinc-copper cell, the copper and zinc plates are called the electrodes. The electrode
where oxidation occurs is called the anode, and the electrode where reduction takes place is
called the cathode. In the zinc-copper cell, the zinc plate is the anode and the copper plate is
the cathode.
Definition: Electrode
An electrode is an electrical conductor that is used to make contact with a metallic part
of a circuit. The anode is the electrode where oxidation takes place. The cathode is the
electrode where reduction takes place.
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17.2
The zinc sulphate and copper sulphate solutions are called the electrolyte solutions.
Definition: Electrolyte
An electrolyte is a substance that contains free ions and which therefore behaves as an
electrical conductor.
The U-tube also plays a very important role in the cell. In the Zn/Zn2+ half-cell, there is a build
up of positive charge because of the release of electrons through oxidation. In the Cu2+ /Cu halfcell, there is a decrease in the positive charge because electrons are gained through reduction.
This causes a movement of SO2−
4 ions into the beaker where there are too many positive ions,
in order to neutralise the solution. Without this, the flow of electrons in the outer circuit stops
completely. The U-tube is called the salt bridge. The salt bridge acts as a transfer medium
that allows ions to flow through without allowing the different solutions to mix and react.
Definition: Salt bridge
A salt bridge, in electrochemistry, is a laboratory device that is used to connect the oxidation
and reduction half-cells of a galvanic cell.
17.2.3
The Galvanic cell
In the zinc-copper cell the important thing to notice is that the chemical reactions that take place
at the two electrodes cause an electric current to flow through the outer circuit. In this type
of cell, chemical energy is converted to electrical energy. These are called galvanic cells.
The zinc-copper cell is one example of a galvanic cell. A galvanic cell (which is also sometimes
referred to as a voltaic or electrochemical cell) consists of two metals that are connected by
a salt bridge between the individual half-cells. A galvanic cell generates electricity using the
reactions that take place at these two metals, each of which has a different reaction potential.
So what is meant by the ’reaction potential’ of a substance? Every metal has a different half
reaction and different dissolving rates. When two metals with different reaction potentials are
used in a galvanic cell, a potential difference is set up between the two electrodes, and the result
is a flow of current through the wire that connects the electrodes. In the zinc-copper cell, zinc
has a higher reaction potential than copper and therefore dissolves more readily into solution.
The metal ’dissolves’ when it loses electrons to form positive metal ions. These electrons are
then transferred through the connecting wire in the outer circuit.
Definition: Galvanic cell
A galvanic (voltaic) cell is an electrochemical cell that uses a chemical reaction between
two dissimilar electrodes dipped in an electrolyte, to generate an electric current.
teresting It was the Italian physician and anatomist Luigi Galvani who marked the birth
Interesting
Fact
Fact
of electrochemistry by making a link between chemical reactions and electricity.
In 1780, Galvani discovered that when two different metals (copper and zinc for
example) were connected together and then both touched to different parts of a
nerve of a frog leg at the same time, they made the leg contract. He called this
”animal electricity”. While many scientists accepted his ideas, another scientist,
Alessandro Volta, did not. In 1800, because of his professional disagreement over
the galvanic response that had been suggested by Luigi Galvani, Volta developed
the voltaic pile, which was very similar to the galvanic cell. It was the work of
these two men that paved the way for all electrical batteries.
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Worked Example 83: Understanding galvanic cells
Question: For the following cell:
Zn|Zn2+ ||Ag + |Ag
1. Give the anode and cathode half-reactions.
2. Write the overall equation for the chemical reaction.
3. Give the direction of the current in the external circuit.
Answer
Step 1 : Identify the oxidation and reduction reactions
In the standard notation format, the oxidation reaction is written on the left and the
reduction reaction on the right. So, in this cell, zinc is oxidised and silver ions are
reduced.
Step 2 : Write the two half reactions
Oxidation half-reaction:
Zn → Zn2+ + 2e−
Reduction half-reaction:
Ag + + e− → Ag
Step 3 : Combine the half-reactions to get the overall equation.
When you combine the two half-reactions, all the reactants must go on the left side
of the equation and the products must go on the right side of the equation. The
overall equation therefore becomes:
Zn + Ag+ + e− → Zn2+ + 2e− + Ag
Note that this equation is not balanced. This will be discussed later in the chapter.
Step 4 : Determine the direction of current flow
A build up of electrons occurs where oxidation takes place. This is at the zinc
electrode. Current will therefore flow from the zinc electrode to the silver electrode.
17.2.4
Uses and applications of the galvanic cell
The principles of the galvanic cell are used to make electrical batteries. In science and technology, a battery is a device that stores chemical energy and makes it available in an electrical
form. Batteries are made of electrochemical devices such as one or more galvanic cells, fuel
cells or flow cells. Batteries have many uses including in torches, electrical appliances (long-life
alkaline batteries), digital cameras (lithium battery), hearing aids (silver-oxide battery), digital
watches (mercury battery) and military applications (thermal battery). Refer to chapter 23 for
more information on batteries.
The galvanic cell can also be used for electroplating. Electroplating occurs when an electrically
conductive object is coated with a layer of metal using electrical current. Sometimes, electroplating is used to give a metal particular properties such as corrosion protection or wear resistance.
At other times, it can be for aesthetic reasons for example in the production of jewellery. This
will be discussed in more detail later in this chapter.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.3
Exercise: Galvanic cells
1. The following half-reactions take place in an electrochemical cell:
Fe → Fe3+ + 3e−
Fe2+ + 2e− → Fe
(a)
(b)
(c)
(d)
(e)
Which is the oxidation half-reaction?
Which is the reduction half-reaction?
Name one oxidising agent.
Name one reducing agent.
Use standard notation to represent this electrochemical cell.
2. For the following cell:
M g|M g 2+||M n2+ |M n
(a)
(b)
(c)
(d)
(e)
(f)
(g)
Give the cathode half-reaction.
Give the anode half-reaction.
Give the overall equation for the electrochemical cell.
What metals could be used for the electrodes in this electrochemical cell.
Suggest two electrolytes for this electrochemical cell.
In which direction will the current flow?
Draw a simple sketch of the complete cell.
3. For the following cell:
Sn|Sn2+ ||Ag + |Ag
(a)
(b)
(c)
(d)
17.3
Give the cathode half-reaction.
Give the anode half-reaction.
Give the overall equation for the electrochemical cell.
Draw a simple sketch of the complete cell.
The Electrolytic cell
In section 17.2, we saw that a chemical reaction that involves a transfer of electrons, can be used
to produce an electric current. In this section, we are going to see whether the ’reverse’ process
applies. In other words, is it possible to use an electric current to force a particular chemical
reaction to occur, which would otherwise not take place? The answer is ’yes’, and the type of
cell that is used to do this, is called an electrolytic cell.
Definition: Electrolytic cell
An electrolytic cell is a type of cell that uses electricity to drive a non-spontaneous reaction.
An electrolytic cell is activated by applying an electrical potential across the anode and cathode
to force an internal chemical reaction between the ions that are in the electrolyte solution. This
process is called electrolysis.
Definition: Electrolysis
In chemistry and manufacturing, electrolysis is a method of separating bonded elements and
compounds by passing an electric current through them.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Activity :: Demonstration : The movement of coloured ions
A piece of filter paper is soaked in an ammonia-ammonium chloride solution and
placed on a microscope slide. The filter paper is then connected to a supply of
electric current using crocodile clips and connecting wire as shown in the diagram
below. A line of copper chromate solution is placed in the centre of the filter paper.
The colour of this solution is initially green-brown.
negative ions
copper chromate (green brown)
+
positive ions
+
-
-
+ -
+ start of reaction
after 20 minutes
The current is then switched on and allowed to run for about 20 minutes. After
this time, the central coloured band disappears and is replaced by two bands, one
yellow and the other blue, which seem to have separated out from the first band of
copper chromate.
Explanation:
• The cell that is used to supply an electric current sets up a potential difference
across the circuit, so that one of the electrodes is positive and the other is
negative.
• The chromate (CrO2−
4 ) ions in the copper chromate solution are attracted
to the positive electrode, while the Cu2+ ions are attracted to the negative
electrode.
Conclusion:
The movement of ions occurs because the electric current in the outer circuit
sets up a potential difference between the two electrodes.
Similar principles apply in the electrolytic cell, where substances that are made of ions can be
broken down into simpler substances through electrolysis.
17.3.1
The electrolysis of copper sulphate
There are a number of examples of electrolysis. The electrolysis of copper sulphate is just one.
Activity :: Demonstration : The electrolysis of copper sulphate
Two copper electrodes are placed in a solution of blue copper sulphate and are
connected to a source of electrical current as shown in the diagram below. The
current is turned on and the reaction is left for a period of time.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
+
+
–
–
positive anode
negative cathode
copper electrode
copper electrode
SO2−
4
17.3
Cu2+
CuSO4 solution
Observations:
• The initial blue colour of the solution remains unchanged.
• It appears that copper has been deposited on one of the electrodes but dissolved
from the other.
Explanation:
• At the negative cathode, positively charged Cu2+ ions are attracted to the
negatively charged electrode. These ions gain electrons and are reduced to
form copper metal, which is deposited on the electrode. The half-reaction that
takes place is as follows:
Cu2+ (aq) + 2e− → Cu(s) (reduction half reaction)
• At the positive anode, copper metal is oxidised to form Cu2+ ions. This is
why it appears that some of the copper has dissolved from the electrode. The
half-reaction that takes place is as follows:
Cu(s) → Cu2+ (aq) + 2e− (oxidation half reaction)
• The amount of copper that is deposited at one electrode is approximately the
same as the amount of copper that is dissolved from the other. The number
of Cu2+ ions in the solution therefore remains almost the same and the blue
colour of the solution is unchanged.
Conclusion:
In this demonstration, an electric current was used to split CuSO4 into its component ions, Cu2+ and SO2−
4 . This process is called electrolysis.
17.3.2
The electrolysis of water
Water can also undergo electrolysis to form hydrogen gas and oxygen gas according to the
following reaction:
2H2 O(l) → 2H2 (g) + O2 (g)
This reaction is very important because hydrogen gas has the potential to be used as an energy source. The electrolytic cell for this reaction consists of two electrodes (normally platinum
metal), submerged in an electrolyte and connected to a source of electric current.
The reduction half-reaction that takes place at the cathode is as follows:
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
2H2 O(l) + 2e− → H2 (g) + 2OH − (aq)
The oxidation half-reaction that takes place at the anode is as follows:
2H2 O(l) → O2 (g) + 4H + (aq) + 4e−
17.3.3
A comparison of galvanic and electrolytic cells
It should be much clearer now that there are a number of differences between a galvanic and an
electrolytic cell. Some of these differences have been summarised in table 17.1.
Item
Metals used for electrode
Charge of the anode
Charge of the cathode
The electrolyte solution/s
Energy changes
Applications
Galvanic cell
Two metals with different
reaction potentials are used
as electrodes
negative
positive
The electrolyte solutions
are kept separate from one
another, and are connected
only by a salt bridge
Chemical potential energy
from chemical reactions is
converted to electrical energy
Run batteries, electroplating
Electrolytic cell
The same metal can be
used for both the cathode
and the anode
positive
negative
The cathode and anode are
in the same electrolyte
An external supply of electrical energy causes a chemical reaction to occur
Electrolysis e.g. of water,
NaCl
Table 17.1: A comparison of galvanic and electrolytic cells
Exercise: Electrolyis
1. An electrolytic cell consists of two electrodes in a silver chloride (AgCl) solution,
connected to a source of current. A current is passed through the solution and
Ag+ ions are reduced to a silver metal deposit on one of the electrodes.
(a) Give the equation for the reduction half-reaction.
(b) Give the equation for the oxidation half-reacion.
2. Electrolysis takes place in a solution of molten lead bromide (PbBr) to produce
lead atoms.
(a) Draw a simple diagram of the electrolytic cell.
(b) Give equations for the half-reactions that take place at the anode and
cathode, and include these in the diagram.
(c) On your diagram, show the direction in which current flows.
17.4
Standard Electrode Potentials
If a voltmeter is connected in the circuit of an electrochemical cell, a reading is obtained. In
other words, there is a potential difference between the two half cells. In this section, we are
going to look at this in more detail to try to understand more about the electrode potentials
of each of the electrodes in the cell. We are going to break this section down so that you build
up your understanding gradually. Make sure that you understand each subsection fully before
moving on, otherwise it might get confusing!
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.4.1
17.4
The different reactivities of metals
All metals have different reactivities. When metals react, they give away electrons and form
positive ions. But some metals do this more easily than others. Look at the following two half
reactions:
Zn → Zn2+ + 2e−
Cu → Cu2+ + 2e−
Of these two metals, zinc is more reactive and is more likely to give away electrons to form Zn2+
ions in solution, than is copper.
17.4.2
Equilibrium reactions in half cells
Let’s think back to the Zn-Cu electrochemical cell. This cell is made up of two half cells and
the reactions that take place at each of the electrodes are as follows:
Zn → Zn2+ + 2e−
Cu2+ + 2e− → Cu
At the zinc electrode, the zinc metal loses electrons and forms Zn2+ ions. The electrons are
concentrated on the zinc metal while the Zn2+ ions are in solution. But some of the ions will be
attracted back to the negatively charged metal, will gain their electrons again and will form zinc
metal. A dynamic equilibrium is set up between the zinc metal and the Zn2+ ions in solution
when the rate at which ions are leaving the metal is equal to the rate at which they are joining
it again. The situation looks something like the diagram in figure 17.1.
--2+
-
- --
2+
2+
zinc metal
concentration of electrons on metal surface
2+
Zn2+ ions in solution
2+
Figure 17.1: Zinc loses electrons to form positive ions in solution. The electrons accumulate on
the metal surface.
The equilibrium reaction is represented like this:
Zn2+ (aq) + 2e− ⇔ Zn(s)
(NOTE: By convention, the ions are written on the left hand side of the equation)
In the zinc half cell, the equilibrium lies far to the left because the zinc loses electrons easily
to form Zn2+ ions. We can also say that the zinc is oxidised and that it is a strong reducing agent.
At the copper electrode, a similar process takes place. The difference though is that copper is
not as reactive as zinc and so it does not form ions as easily. This means that the build up of
electrons on the copper electrode is less (figure 17.2).
The equilibrium reaction is shown like this:
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17.4
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
-
-
-2+
copper metal
concentration of electrons on metal surface
Cu2+ ions in solution
2+
Figure 17.2: Zinc loses electrons to form positive ions in solution. The electrons accumulate on
the metal surface.
Cu2+ (aq) + 2e− ⇔ Cu(s)
The equation lies far to the right because most of the copper is present as copper metal rather
than as Cu2+ ions. In this half reaction, the Cu2+ ions are reduced.
17.4.3
Measuring electrode potential
If we put the two half cells together, a potential difference is set up in two places in the Zn-Cu
cell:
1. There is a potential difference between the metal and the solution surrounding it because
one is more negative than the other.
2. There is a potential difference between the Zn and Cu electrodes because one is more
negative than the other.
It is the potential difference (recorded as a voltage) between the two electrodes that causes
electrons, and therefore current, to flow from the more negative electrode to the less negative
electrode.
The problem though is that we cannot measure the potential difference (voltage) between a
metal and its surrounding solution in the cell. To do this, we would need to connect a voltmeter
to both the metal and the solution, which is not possible. This means we cannot measure the
exact electrode potential (Eo V) of a particular metal. The electrode potential describes the
ability of a metal to give up electrons. And if the exact electrode potential of each of the
electrodes involved can’t be measured, then it is difficult to calculate the potential difference
between them. But what we can do is to try to describe the electrode potential of a metal
relative to another substance. We need to use a standard reference electrode for this.
17.4.4
The standard hydrogen electrode
Before we look at the standard hydrogen electrode, it may be useful to have some more understanding of the ideas behind a ’reference electrode’. Refer to the Tip box on ’Understanding the
ideas behind a reference electrode’ before you read further.
Important: Understanding the ideas behind a reference electrode
Adapted from www.chemguide.co.uk
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.4
Let’s say that you have a device that you can use to measure heights from some distance away.
You want to use this to find out how tall a particular person is. Unfortunately, you can’t see
their feet because they are standing in long grass. Although you can’t measure their absolute
height, what you can do is to measure their height relative to the post next to them. Let’s say
that person A for example is 15 cm shorter than the height of the post. You could repeat this
for a number of other people (B and C). Person B is 30 cm shorter than the post and person C
is 10 cm taller than the post.
A
B
C
You could summarise your findings as follows:
Person
A
B
C
Height relative to post (cm)
-15
-30
+10
Although you don’t know any of their absolute heights, you can rank them in order, and do some
very simple sums to work out exactly how much taller one is than another. For example, person
C is 25 cm taller than A and 40 cm taller than B.
As mentioned earlier, it is difficult to measure the absolute electrode potential of a particular
substance, but we can use a reference electrode (similar to the ’post’ in the Tip box example)
that we use to calculate relative electrode potentials for these substances. The reference elctrode
that is used is the standard hydrogen electrode (figure 17.3).
Definition: Standard hydrogen electrode
The standard hydrogen electrode is a redox electrode which forms the basis of the scale of
oxidation-reduction potentials. The actual electrode potential of the hydrogen electrode is
estimated to be 4.44 0.02 V at 250 C, but its standard electrode potential is said to be zero
at all temperatures so that it can be used as for comparison with other electrodes. The
hydrogen electrode is based on the following redox half cell:
2H+ (aq) + 2e− → H2 (g)
A standard hydrogen electrode consists of a platinum electrode in a solution containing H+ ions.
The solution (e.g. H2 SO4 ) that contains the H+ ions has a concentration of 1 mol.dm−3 . As
the hydrogen gas bubbles over the platinum electrode, an equilibrium is set up between hydrogen
molecules and hydrogen ions in solution. The reaction is as follows:
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17.4
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
b
H2 gas
(1 x atmosphereric pressure)
c
b
Pt
cb
b
c
c
b
[H3 O+ ]
(1 mol.dm−3 )
25 ◦ C
c
b
Figure 17.3: The standard hydrogen electrode
2H + (aq) + 2e− ⇔ H2 (g)
The position of this equilibrium can change if you change some of the conditions (e.g. concentration, temperature). It is therefore important that the conditions for the standard hydrogen
electrode are standardised as follows: pressure = 100 kPa (1atm); temperature = 298 K (250 C)
and concentration = 1 mol.dm−3 .
In order to use the hydrogen electrode, it needs to be attached to the electrode system that
you are investigating. For example, if you are trying to determine the electrode potential of
copper, you will need to connect the copper half cell to the hydrogen electrode; if you are trying
to determine the electrode potential of zinc, you will need to connect the zinc half cell to the
hydrogen electrode and so on. Let’s look at the examples of zinc and copper in more detail.
1. Zinc
Zinc has a greater tendency than hydrogen to form ions, so if the standard hydrogen
electrode is connected to the zinc half cell, the zinc will be relatively more negative because
the electrons that are released when zinc is oxidised will accumulate on the metal. The
equilibria on the two electrodes are as follows:
Zn2+ (aq) + 2e− ⇔ Zn(s)
2H + (aq) + 2e− ⇔ H2 (g)
In the zinc half-reaction, the equilibrium lies far to the left and in the hydrogen halfreaction, the equilibrium lies far to the right. A simplified representation of the cell is
shown in figure 17.4.
The voltmeter measures the potential difference between the charge on these electrodes. In
this case, the voltmeter would read 0.76 and would show that Zn is the negative electrode
(i.e. it has a relatively higher number of electrons).
2. Copper
Copper has a lower tendency than hydrogen to form ions, so if the standard hydrogen
electrode is connected to the copper half cell, the hydrogen will be relatively more negative.
The equilibria on the two electrodes are as follows:
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.4
V
-
H electrode
(less negative) -
-----
-
Zn electrode with electrons
Figure 17.4: When zinc is connected to the standard hydrogen electrode, relatively few electrons
build up on the platinum (hydrogen) electrode. There are lots of electrons on the zinc electrode.
Cu2+ (aq) + 2e− ⇔ Cu(s)
2H + (aq) + 2e− ⇔ H2 (g)
In the copper half-reaction, the equilibrium lies far to the right and in the hydrogen halfreaction, the equilibrium lies far to the left. A simplified representation of the cell is shown
in figure 17.5.
V
H electrode
-----
-
-
Cu electrode
-
Figure 17.5: When copper is connected to the standard hydrogen electrode, relatively few electrons build up on the copper electrode. There are lots of electrons on the hydrogen electrode.
The voltmeter measures the potential difference between the charge on these electrodes. In
this case, the voltmeter would read 0.34 and would show that Cu is the positive electrode
(i.e. it has a relatively lower number of electrons).
17.4.5
Standard electrode potentials
The voltages recorded earlier when zinc and copper were connected to a standard hydrogen
electrode are in fact the standard electrode potentials for these two metals. It is important
to remember that these are not absolute values, but are potentials that have been measured
relative to the potential of hydrogen if the standard hydrogen electrode is taken to be zero.
Important: Conventions and voltage sign
By convention, the hydrogen electrode is written on the left hand side of the cell. The sign of
the voltage tells you the sign of the metal electrode.
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17.4
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
In the examples we used earlier, zinc’s electrode potential is actually -0.76 and copper is +0.34.
So, if a metal has a negative standard electrode potential, it means it forms ions easily. The
more negative the value, the easier it is for that metal to form ions. If a metal has a positive
standard electrode potential, it means it does not form ions easily. This will be explained in more
detail below.
Luckily for us, we do not have to calculate the standard electrode potential for every metal. This
has been done already and the results are recorded in a table of standard electrode potentials
(table 17.2).
A few examples from the table are shown in table 17.3. These will be used to explain some of
the trends in the table of electrode potentials.
Refer to table 17.3 and notice the following trends:
• Metals at the top of series (e.g. Li) have more negative values. This means they ionise
easily, in other words, they release electrons easily. These metals are easily oxidised and
are therefore good reducing agents.
• Metal ions at the bottom of the table are good at picking up electrons. They are easily
reduced and are therefore good oxidising agents.
• The reducing ability (i.e. the ability to act as a reducing agent) of the metals in the table
increases as you move up in the table.
• The oxidising ability of metals increases as you move down in the table.
Worked Example 84: Using the table of Standard Electrode Potentials
Question:
The following half-reactions take place in an electrochemical cell:
Cu2+ + 2e− ⇔ Cu
Ag− + e− ⇔ Ag
1. Which of these reactions will be the oxidation half-reaction in the cell?
2. Which of these reactions will be the reduction half-reaction in the cell?
Answer
Step 5 : Determine the electrode potential for each metal
From the table of standard electrode potentials, the electrode potential for the copper half-reaction is +0.34 V. The electrode potential for the silver half-reaction is
+0.80 V.
Step 6 : Use the electrode potential values to determine which metal is
oxidised and which is reduced
Both values are positive, but silver has a higher positive electrode potential than
copper. This means that silver does not form ions easily, in other words, silver is
more likely to be reduced. Copper is more likely to be oxidised and to form ions more
easily than silver. Copper is the oxidation half-reaction and silver is the reduction
half-reaction.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Half-Reaction
Li+ + e− ⇋ Li
K + + e− ⇋ K
Ba2+ + 2e− ⇋ Ba
Ca2+ + 2e− ⇋ Ca
N a+ + e − ⇋ N a
M g 2+ + 2e− ⇋ M g
M n2+ + 2e− ⇋ M n
2H2O + 2e− ⇋ H2 (g) + 2OH −
Zn2+ + 2e− ⇋ Zn
Cr2+ + 2e− ⇋ Cr
F e2+ + 2e− ⇋ F e
Cr3+ + 3e− ⇋ Cr
Cd2+ + 2e− ⇋ Cd
Co2+ + 2e− ⇋ Co
N i2+ + 2e− ⇋ N i
Sn2+ + 2e− ⇋ Sn
P b2+ + 2e− ⇋ P b
F e3+ + 3e− ⇋ F e
2H + + 2e− ⇋ H2 (g)
S + 2H + + 2e− ⇋ H2 S(g)
Sn4+ + 2e− ⇋ Sn2+
Cu2+ + e− ⇋ Cu+
SO42+ + 4H + + 2e− ⇋ SO2 (g) + 2H2 O
Cu2+ + 2e− ⇋ Cu
2H2 O + O2 + 4e− ⇋ 4OH −
Cu+ + e− ⇋ Cu
I2 + 2e− ⇋ 2I −
O2 (g) + 2H + + 2e− ⇋ H2 O2
F e3+ + e− ⇋ F e2+
N O3− + 2H + + e− ⇋ N O2 (g) + H2 O
Hg 2+ + 2e− ⇋ Hg(l)
Ag + + e− ⇋ Ag
N O3− + 4H + + 3e− ⇋ N O(g) + 2H2 O
Br2 + 2e− ⇋ 2Br−
O2 (g) + 4H + + 4e− ⇋ 2H2 O
M nO2 + 4H + + 2e− ⇋ M n2+ + 2H2 O
Cr2 O72− + 14H + + 6e− ⇋ 2Cr3+ + 7H2 O
Cl2 + 2e− ⇋ 2Cl−
Au3+ + 3e− ⇋ Au
M nO4− + 8H + + 5e− ⇋ M n2+ + 4H2 O
Co3+ + e− ⇋ Co2+
F2 + 2e− ⇋ 2F −
17.4
E0V
-3.04
-2.92
-2.90
-2.87
-2.71
-2.37
-1.18
-0.83
-0.76
-0.74
-0.44
-0.41
-0.40
-0.28
-0.25
-0.14
-0.13
-0.04
0.00
0.14
0.15
0.16
0.17
0.34
0.40
0.52
0.54
0.68
0.77
0.78
0.78
0.80
0.96
1.06
1.23
1.28
1.33
1.36
1.50
1.52
1.82
2.87
Table 17.2: Standard Electrode Potentials
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Half-Reaction
Li+ + e− ⇋ Li
Zn2+ + 2e− ⇋ Zn
F e3+ + 3e− ⇋ F e
2H + + 2e− ⇋ H2 (g)
Cu2+ + 2e− ⇋ Cu
Hg 2+ + 2e− ⇋ Hg(l)
Ag + + e− ⇋ Ag
E0V
-3.04
-0.76
-0.04
0.00
0.34
0.78
0.80
Table 17.3: A few examples from the table of standard electrode potentials
Important: Learning to understand the question in a problem.
Before you tackle this problem, make sure you understand exactly what the question is
asking. If magnesium is able to displace silver from a solution of silver nitrate, this means
that magnesium metal will form magnesium ions and the silver ions will become silver metal.
In other words, there will now be silver metal and a solution of magnesium nitrate. This
will only happen if magnesium has a greater tendency than silver to form ions. In other
words, what the question is actually asking is whether magnesium or silver forms ions more
easily.
Worked Example 85: Using the table of Standard Electrode Potentials
Question: Is magnesium able to displace silver from a solution of silver nitrate?
Answer
Step 1 : Determine the half-reactions that would take place if magnesium
were to displace silver nitrate.
The half-reactions are as follows:
M g 2+ + 2e− ⇔ M g
Ag + + e− ⇔ Ag
Step 2 : Use the table of electrode potentials to see which metal forms
ions more easily.
Looking at the electrode potentials for the magnesium and silver reactions:
For the magnesium half-reaction: Eo V = -2.37
For the silver half-reaction: Eo V = 0.80
This means that magnesium is more easily oxidised than silver and the equilibrium in this half-reaction lies to the left. The oxidation reaction will occur
spontaneously in magnesium. Silver is more easily reduced and the equilibrium
lies to the right in this half-reaction. It can be concluded that magnesium will
displace silver from a silver nitrate solution so that there is silver metal and
magnesium ions in the solution.
Exercise: Table of Standard Electrode Potentials
1. In your own words, explain what is meant by the ’electrode potential’ of a
metal.
2. Give the standard electrode potential for each of the following metals:
(a) magnesium
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.4
(b) lead
(c) nickel
3. Refer to the electrode potentials in table 17.3.
(a)
(b)
(c)
(d)
Which of the metals is most likely to be oxidised?
Which metal is most likely to be reduced?
Which metal is the strongest reducing agent?
In the copper half-reaction, does the equilibrium position for the reaction
lie to the left or to the right? Explain your answer.
(e) In the mercury half-reaction, does the equilibrium position for the reaction
lie to the left or to the right? Explain your answer.
(f) If silver was added to a solution of copper sulphate, would it displace the
copper from the copper sulphate solution? Explain your answer.
4. Use the table of standard electrode potentials to put the following in order from
the strongest oxidising agent to the weakest oxidising agent.
•
•
•
•
Cu2+
MnO−
4
Br2
Zn2+
5. Look at the following half-reactions.
•
•
•
•
Ca2+ + 2e− → Ca
Cl2 + 2e− → 2Cl
F e3+ + 3e− → F e
I2 + 2e− → 2I −
(a) Which substance is the strongest oxidising agent?
(b) Which substance is the strongest reducing agent?
6. Which one of the substances listed below acts as the oxidising agent in the
following reaction?
2−
+
3+
3SO2 + Cr2 O2−
+ H2 O
7 + 2H → 3SO4 + 2Cr
(a)
(b)
(c)
(d)
H+
Cr3+
SO2
Cr2 O2−
7
(IEB Paper 2, 2004)
7. If zinc is added to a solution of magnesium sulphate, will the zinc displace the
magnesium from the solution? Give a detailed explanation for your answer.
17.4.6
Combining half cells
Let’s stay with the example of the zinc and copper half cells. If we combine these cells as we
did earlier in the chapter (section 17.2), the following two equilibria exist:
Zn2+ + 2e− ⇔ Zn(E 0 = −0.76V )
Cu2+ + 2e− ⇔ Cu(E 0 = +0.34V )
We know from demonstrations, and also by looking at the sign of the electrode potential, that
when these two half cells are combined, zinc will be the oxidation half-reaction and copper will be
the reduction half-reaction. A voltmeter connected to this cell will show that the zinc electrode
is more negative than the copper electrode. The reading on the meter will show the potential
difference between the two half cells. This is known as the electromotive force (emf) of the
cell.
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Definition: Electromotive Force (emf)
The emf of a cell is defined as the maximum potential difference between two electrodes or
half cells in a voltaic cell. emf is the electrical driving force of the cell reaction. In other
words, the higher the emf, the stronger the reaction.
Definition: Standard emf (E0cell )
Standard emf is the emf of a voltaic cell operating under standard conditions (i.e. 100 kPa,
concentration = 1 mol.dm−3 and temperature = 298 K). The symbol 0 denotes standard
conditions.
When we want to represent this cell, it is shown as follows:
Zn|Zn2+ (1mol.dm−3 )||Cu2+ (1mol.dm−3 )|Cu
The anode half cell (where oxidation takes place) is always written on the left. The cathode
half cell (where reduction takes place) is always written on the right.
It is important to note that the potential difference across a cell is related to the extent to which
the spontaneous cell reaction has reached equilibrium. In other words, as the reaction proceeds
and the concentration of reactants decreases and the concentration of products increases, the
reaction approaches equilibrium. When equilibrium is reached, the emf of the cell is zero and
the cell is said to be ’flat’. There is no longer a potential difference between the two half cells,
and therefore no more current will flow.
17.4.7
Uses of standard electrode potential
Standard electrode potentials have a number of different uses.
Calculating the emf of an electrochemical cell
To calculate the emf of a cell, you can use any one of the following equations:
E0(cell) = E0 (right) - E0 (left) (’right’ refers to the electrode that is written on the right in
standard cell notation. ’Left’ refers to the half-reaction written on the left in this notation)
E0(cell) = E0 (reduction half reaction) - E0 (oxidation half reaction)
E0(cell) = E0 (oxidising agent) - E0 (reducing agent)
E0(cell) = E0 (cathode) - E0 (anode)
So, for the Zn-Cu cell,
E0(cell) = 0.34 - (-0.76)
= 0.34 + 0.76
= 1.1 V
Worked Example 86: Calculating the emf of a cell
Question: The following reaction takes place:
Cu(s) + Ag + (aq) → Cu2+ (aq) + Ag(s)
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
1. Represent the cell using standard notation.
2. Calculate the cell potential (emf) of the electrochemical cell.
Answer
Step 1 : Write equations for the two half reactions involved
Cu2+ + 2e− ⇔ Cu (Eo V = 0.16V)
Ag + + e− ⇔ Ag (Eo V = 0.80V)
Step 2 : Determine which reaction takes place at the cathode and
which is the anode reaction
Both half-reactions have positive electrode potentials, but the silver half-reaction
has a higher positive value. In other words, silver does not form ions easily, and
this must be the reduction half-reaction. Copper is the oxidation half-reaction.
Copper is oxidised, therefore this is the anode reaction. Silver is reduced and so
this is the cathode reaction.
Step 3 : Represent the cell using standard notation
Cu|Cu2+ (1mol.dm−3 )||Ag + (1mol.dm−3 )|Ag
Step 4 : Calculate the cell potential
E0(cell) = E0 (cathode) - E0 (anode)
= +0.80 - (+0.34)
= +0.46 V
Worked Example 87: Calculating the emf of a cell
Question: Calculate the cell potential of the electrochemical cell in which the following reaction takes place, and represent the cell using standard notation.
M g(s) + 2H + (aq) → M g2+(aq) + H2 (g)
Answer
Step 1 : Write equations for the two half reactions involved
M g 2+ + 2e− ⇔ M g (Eo V = -2.37)
2H + + 2e− ⇔ H2 (Eo V = 0.00)
Step 2 : Determine which reaction takes place at the cathode and
which is the anode reaction
From the overall equation, it is clear that magnesium is oxidised and hydrogen
ions are reduced in this reaction. Magnesium is therefore the anode reaction and
hydrogen is the cathode reaction.
Step 3 : Represent the cell using standard notation
M g|M g 2+ ||H + |H2
Step 4 : Calculate the cell potential
E0(cell) = E0 (cathode) - E0 (anode)
= 0.00 - (-2.37)
= +2.37 V
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17.4
17.4
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Predicting whether a reaction will take place spontaneously
Look at the following example to help you to understand how to predict whether a reaction will
take place spontaneously or not.
In the reaction,
P b2+ (aq) + 2Br− (aq) → Br2 (l) + P b(s)
the two half reactions are as follows:
P b2+ + 2e− ⇔ P b (-0.13 V)
Br2 + 2e− ⇔ 2Br− (+1.06 V)
Important: Half cell reactions
You will see that the half reactions are written as they appear in the table of standard electrode
potentials. It may be useful to highlight the reacting substance in each half reaction. In this
case, the reactants are Pb2+ and Br− ions.
Look at the electrode potential for the first half reaction. The negative value shows that lead
loses electrons easily, in other words it is easily oxidised. The reaction would normally proceed
from right to left (i.e. the equilibrium lies to the left), but in the original equation, the opposite
is happening. It is the Pb2+ ions that are being reduced to lead. This part of the reaction is
therefore not spontaneous. The positive electrode potential value for the bromine half-reaction
shows that bromine is more easily reduced, in other words the equilibrium lies to the right. The
spontaneous reaction proceeds from left to right. This is not what is happening in the original
equation and therefore this is also not spontaneous. Overall it is clear then that the reaction will
not proceed spontaneously.
Worked Example 88: Predicting whether a reaction is spontaneous
Question: Will copper react with dilute sulfuric acid (H2 SO4 )? You are given the
following half reactions:
Cu2+ (aq) + 2e− ⇔ Cu(s) (E0 = +0.34 V)
2H + (aq) + 2e− ⇔ H2 (g) (E0 = 0 V)
Answer
Step 5 : For each reaction, look at the electrode potentials and decide in
which direction the equilibrium lies
In the first half reaction, the positive electrode potential means that copper does
not lose electrons easily, in other words it is more easily reduced and the equilibrium
position lies to the right. Another way of saying this is that the spontaneous reaction
is the one that proceeds from left to right, when copper ions are reduced to copper
metal.
In the second half reaction, the spontaneous reaction is from right to left.
Step 6 : Compare the equilibrium positions to the original reaction
What you should notice is that in the original reaction, the reactants are copper
(Cu) and sulfuric acid (2H+ ). During the reaction, the copper is oxidised and the
hydrogen ions are reduced. But from an earlier step, we know that neither of these
half reactions will proceed spontaneously in the direction indicated by the original
reaction. The reaction is therefore not spontaneous.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.4
Important:
A second method for predicting whether a reaction is spontaneous
Another way of predicting whether a reaction occurs spontaneously, is to look at the sign of the
emf value for the cell. If the emf is positive then the reaction is spontaneous. If the emf is
negative, then the reaction is not spontaneous.
Balancing redox reactions
We will look at this in more detail in the next section.
Exercise: Predicting whether a reaction will take place spontaneously
1. Predict whether the following reaction will take place spontaneously or not.
Show all your working.
2Ag(s) + Cu2+ (aq) → Cu(s) + 2Ag + (aq)
2. Zinc metal reacts with an acid, H+ (aq) to produce hydrogen gas.
(a) Write an equation for the reaction, using the table of electrode potentials.
(b) Predict whether the reaction will take place spontaneously. Show your
working.
3. Four beakers are set up, each of which contains one of the following solutions:
(a)
(b)
(c)
(d)
Mg(NO3 )2
Ba(NO3 )2
Cu(NO3 )2
Al(NO3 )2
Iron is added to each of the beakers. In which beaker will a spontaneous
reaction take place?
4. Which one of the following solutions can be stored in an aluminium container?
(a)
(b)
(c)
(d)
Cu(SO)4
Zn(SO)4
NaCl
Pb(NO3 )2
Exercise: Electrochemical cells and standard electrode potentials
1. An electrochemical cell is made up of a copper electrode in contact with a
copper nitrate solution and an electrode made of an unknown metal M in
contact with a solution of MNO3 . A salt bridge containing a KNO3 solution
joins the two half cells. A voltmeter is connected across the electrodes. Under
standard conditions the reading on the voltmeter is 0.46V.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
V
Cu
Salt bridge (KNO3 )
Cu(NO3 )2 (aq)
M
MNO3 (aq)
The reaction in the copper half cell is given by:
Cu → Cu2+ + 2e−
(a) Write down the standard conditions which apply to this electrochemical
cell.
(b) Identify the metal M. Show calculations.
(c) Use the standard electrode potentials to write down equations for the:
i. cathode half-reaction
ii. anode half-reaction
iii. overall cell reaction
(d) What is the purpose of the salt bridge?
(e) Explain why a KCl solution would not be suitable for use in the salt bridge
in this cell.
(IEB Paper 2, 2004)
2. Calculate the emf for each of the following standard electrochemical cells:
(a)
M g|M g 2+ ||H + |H2
(b)
F e|F e3+ ||F e2+ |F e
(c)
Cr|Cr2+||Cu2+ |Cu
(d)
P b|P b2+ ||Hg 2+ |Hg
3. Given the following two half-reactions:
• F e3+ (aq) + e− ⇔ F e2+ (aq)
• M nO4− (aq) + 8H + (aq) + 5e− ⇔ M n2+ (aq) + 4H2 O(l)
(a) Give the standard electrode potential for each half-reaction.
(b) Which reaction takes place at the cathode and which reaction takes place
at the anode?
(c) Represent the electrochemical cell using standard notation.
(d) Calculate the emf of the cell
17.5
Balancing redox reactions
Half reactions can be used to balance redox reactions. We are going to use some worked examples
to help explain the method.
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Worked Example 89: Balancing redox reactions
Question: Magnesium reduces copper (II) oxide to copper. In the process, magnesium is oxidised to magnesium ions. Write a balanced equation for this reaction.
Answer
Step 1 : Write down the unbalanced oxidation half reaction.
M g → M g 2+
Step 2 : Balance the number of atoms on both sides of the equation.
You are allowed to add hydrogen ions (H+ ) and water molecules if the reaction takes
place in an acid medium. If the reaction takes place in a basic medium, you can add
either hydroxide ions (OH− ) or water molecules. In this case, there is one magnesium atom on the left and one on the right, so no additional atoms need to be added.
Step 3 : Once the atoms are balanced, check that the charges balance.
Charges can be balanced by adding electrons to either side. The charge on the left
of the equation is 0, but the charge on the right is +2. Therefore, two electrons
must be added to the right hand side so that the charges balance. The half reaction
is now:
M g → M g 2+ + 2e−
Step 4 : Repeat the above steps, but this time using the reduction half
reaction.
The reduction half reaction is:
Cu2+ → Cu
The atoms balance but the charges don’t. Two electrons must be added to the right
hand side.
Cu2+ + 2e− → Cu
Step 5 : Multiply each half reaction by a suitable number so that the number
of electrons released in the oxidation half reaction is made equal to the
number of electrons that are accepted in the reduction half reaction.
No multiplication is needed because there are two electrons on either side.
Step 6 : Combine the two half reactions to get a final equation for the overall
reaction.
M g + Cu2+ + 2e− → M g 2+ + Cu + 2e− (The electrons on either side cancel
and you get...)
M g + Cu2+ → M g 2+ + Cu
Step 7 : Do a final check to make sure that the equation is balanced
In this case, it is.
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17.5
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Worked Example 90: Balancing redox reactions
Question: Chlorine gas oxidises Fe(II) ions to Fe(III) ions. In the process, chlorine
is reduced to chloride ions. Write a balanced equation for this reaction.
Answer
Step 1 : Write down the oxidation half reaction.
F e2+ → F e3+
Step 2 : Balance the number of atoms on both sides of the equation.
There is one iron atom on the left and one on the right, so no additional atoms need
to be added.
Step 3 : Once the atoms are balanced, check that the charges balance.
The charge on the left of the equation is +2, but the charge on the right is +3.
Therefore, one electron must be added to the right hand side so that the charges
balance. The half reaction is now:
F e2+ → F e3+ + e−
Step 4 : Repeat the above steps, but this time using the reduction half
reaction.
The reduction half reaction is:
Cl2 → Cl−
The atoms don’t balance, so we need to multiply the right hand side by two to fix
this. Two electrons must be added to the left hand side to balance the charges.
Cl2 + 2e− → 2Cl−
Step 5 : Multiply each half reaction by a suitable number so that the number
of electrons released in the oxidation half reaction is made equal to the
number of electrons that are accepted in the reduction half reaction.
We need to multiply the oxidation half reaction by two so that the number of electrons
on either side are balanced. This gives:
2F e2+ → 2F e3+ + 2e−
Step 6 : Combine the two half reactions to get a final equation for the overall
reaction.
2F e2+ + Cl2 → 2F e3+ + 2Cl−
Step 7 : Do a final check to make sure that the equation is balanced
The equation is balanced.
Worked Example 91: Balancing redox reactions in an acid medium
Question: The following reaction takes place in an acid medium:
Cr2 O72− + H2 S → Cr3+ + S
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CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
Write a balanced equation for this reaction.
Answer
Step 1 : Write down the oxidation half reaction.
Cr2 O72− → Cr3+
Step 2 : Balance the number of atoms on both sides of the equation.
We need to multiply the right side by two so that the number of Cr atoms will
balance. To balance the oxygen atoms, we will need to add water molecules to the
right hand side.
Cr2 O72− → 2Cr3+ + 7H2 O
Now the oxygen atoms balance but the hydrogens don’t. Because the reaction takes
place in an acid medium, we can add hydrogen ions to the left side.
Cr2 O72− + 14H + → 2Cr3+ + 7H2 O
Step 3 : Once the atoms are balanced, check that the charges balance.
The charge on the left of the equation is (-2+14) = +12, but the charge on the
right is +6. Therefore, six electrons must be added to the left hand side so that the
charges balance. The half reaction is now:
Cr2 O72− + 14H + + 6e− → 2Cr3+ + 7H2 O
Step 4 : Repeat the above steps, but this time using the reduction half
reaction.
The reduction half reaction after the charges have been balanced is:
S 2− → S + 2e−
Step 5 : Multiply each half reaction by a suitable number so that the number
of electrons released in the oxidation half reaction is made equal to the
number of electrons that are accepted in the reduction half reaction.
We need to multiply the reduction half reaction by three so that the number of
electrons on either side are balanced. This gives:
3S 2− → 3S + 6e−
Step 6 : Combine the two half reactions to get a final equation for the overall
reaction.
Cr2 O72− + 14H + + 3S 2− → 3S + 2Cr3+ + 7H2 O
Step 7 : Do a final check to make sure that the equation is balanced
Worked Example 92: Balancing redox reactions in an alkaline medium
Question: If ammonia solution is added to a solution that contains cobalt(II) ions, a
complex ion is formed, called the hexaaminecobalt(II) ion (Co(NH3 )2+
6 ). In a chemical reaction with hydrogen peroxide solution, hexaaminecobalt ions are oxidised by
hydrogen peroxide solution to the hexaaminecobalt(III) ion Co(NH3 )3+
6 . Write a
balanced equation for this reaction.
Answer
Step 1 : Write down the oxidation half reaction
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17.5
17.5
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
3+
Co(N H3 )2+
6 → Co(N H3 )6
Step 2 : Balance the number of atoms on both sides of the equation.
The number of atoms are the same on both sides.
Step 3 : Once the atoms are balanced, check that the charges balance.
The charge on the left of the equation is +2, but the charge on the right is +3.
One elctron must be added to the right hand side to balance the charges in the
equation.The half reaction is now:
3+
−
Co(N H3 )2+
6 → Co(N H3 )6 + e
Step 4 : Repeat the above steps, but this time using the reduction half
reaction.
Although you don’t actually know what product is formed when hydrogen peroxide
is reduced, the most logical product is OH− . The reduction half reaction is:
H2 O2 → OH −
After the atoms and charges have been balanced, the final equation for the reduction
half reaction is:
H2 O2 + 2e− → 2OH −
Step 5 : Multiply each half reaction by a suitable number so that the number
of electrons released in the oxidation half reaction is made equal to the
number of electrons that are accepted in the reduction half reaction.
We need to multiply the oxidation half reaction by two so that the number of electrons
on both sides are balanced. This gives:
3+
−
2Co(N H3 )2+
6 → 2Co(N H3 )6 + 2e
Step 6 : Combine the two half reactions to get a final equation for the overall
reaction.
3+
−
2Co(N H3 )2+
6 + H2 O2 → 2Co(N H3 )6 + 2OH
Step 7 : Do a final check to make sure that the equation is balanced
Exercise: Balancing redox reactions
1. Balance the following equations.
(a) HN O3 + P bS → P bSO4 + N O + H2 O
(b) N aI + F e2 (SO4 )3 → I2 + F eSO4 + N a2 SO4
2. Manganate(VII) ions (MnO−
4 ) oxidise hydrogen peroxide (H2 O2 ) to oxygen
gas. The reaction is done in an acid medium. During the reaction, the manganate(VII) ions are reduced to manganese(II) ions (Mn2+ ). Write a balanced
equation for the reaction.
3. Chlorine gas is prepared in the laboratory by adding concentrated hydrochloric
acid to manganese dioxide powder. The mixture is carefully heated.
(a) Write down a balanced equation for the reaction which takes place.
(b) Using standard electrode potentials, show by calculations why this mixture
needs to be heated.
(c) Besides chlorine gas which is formed during the reaction, hydrogen chloride
gas is given off when the conentrated hydrochloric acid is heated. Explain
why the hydrogen chloride gas is removed from the gas mixture when the
gas is bubbled through water.
(IEB Paper 2, 2004)
346
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
17.6
4. The following equation can be deduced from the table of standard electrode
potentials:
+
3+
2Cr2 O2−
(aq) + 3O2 (g) + 8H2 O(l) (E0 =
7 (aq) + 16H (aq) → 4Cr
+0.10V)
This equation implies that an acidified solution of aqueous potassium dichromate (orange) should react to form Cr3+ (green). Yet aqueous laboratory
solutions of potassium dichromate remain orange for years. Which ONE of the
following best explains this?
(a)
(b)
(c)
(d)
Laboratory solutions of aqueous potassium dichromate are not acidified
The E0 value for this reaction is only +0.10V
The activation energy is too low
The reaction is non-spontaneous
(IEB Paper 2, 2002)
5. Sulfur dioxide gas can be prepared in the laboratory by heating a mixture of
copper turnings and concentrated sulfuric acid in a suitable flask.
(a) Derive a balanced ionic equation for this reaction using the half-reactions
that take place.
(b) Give the E0 value for the overall reaction.
(c) Explain why it is necessary to heat the reaction mixture.
(d) The sulfur dioxide gas is now bubbled through an aqueous solution of
potassium dichromate. Describe and explain what changes occur during
this process.
(IEB Paper 2, 2002)
17.6
Applications of electrochemistry
Electrochemistry has a number of different uses, particularly in industry. We are going to look
at a few examples.
17.6.1
Electroplating
Electroplating is the process of using electrical current to coat an electrically conductive object
with a thin layer of metal. Mostly, this application is used to deposit a layer of metal that has
some desired property (e.g. abrasion and wear resistance, corrosion protection, improvement of
aesthetic qualities etc.) onto a surface that doesn’t have that property. Electro-refining (also
sometimes called electrowinning is electroplating on a large scale. Electrochemical reactions are
used to deposit pure metals from their ores. One example is the electrorefining of copper.
Copper plays a major role in the electrical reticulation industry as it is very conductive and is
used in electric cables. One of the problems though is that copper must be pure if it is to be
an effective current carrier. One of the methods used to purify copper, is electro-winning. The
copper electro-winning process is as follows:
1. Bars of crude (impure) copper containing other metallic impurities is placed on the anodes.
2. The cathodes are made up of pure copper with few impurities.
3. The electrolyte is a solution of aqueous CuSO4 and H2 SO4 .
347
17.6
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
4. When current passes through the cell, electrolysis takes place. The impure copper anode
dissolves to form Cu2+ ions in solution. These positive ions are attracted to the negative
cathode, where reduction takes place to produce pure copper metal. The reactions that
take place are as follows:
At the anode:
Cu(s) → Cu2+ (aq) + 2e−
At the cathode:
Cu+2 (aq) + 2e− → Cu(s)
(> 99%purity)
5. The other metal impurities (Zn, Au, Ag, Fe and Pb) do not dissolve and form a solid
sludge at the bottom of the tank or remain in solution in the electrolyte.
+
–
+
–
negative cathode
positive anode
impure copper electrode
pure copper electrode
Cu
2+
Figure 17.6: A simplified diagram to illustrate what happens during the electrowinning of copper
17.6.2
The production of chlorine
Electrolysis can also be used to produce chlorine gas from brine/seawater (NaCl). This is sometimes referred to as the ’Chlor-alkali’ process. The reactions that take place are as follows:
At the anode the reaction is:
2Cl− → Cl2 (g) + 2e−
whereas at the cathode, the following happens:
2N a+ + 2H2 O + 2e− → 2N a+ + 2OH − + H2
The overall reaction is:
2N a+ + 2H2 O + 2Cl− → 2N a+ + 2OH − + H2 + Cl2
Chlorine is a very important chemical. It is used as a bleaching agent, a disinfectant, in solvents,
pharmaceuticals, dyes and even plastics such as polyvinlychloride (PVC).
348
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
+
17.7
–
+
–
positive anode
negative cathode
electrode
electrode
Na+
Cl−
NaCl solution
Figure 17.7: The electrolysis of sodium chloride
17.6.3
Extraction of aluminium
Aluminum metal is a commonly used metal in industry where its properties of being both light
and strong can be utilized. It is also used in the manufacture of products such as aeroplanes
and motor cars. The metal is present in deposits of bauxite which is a mixture of silicas, iron
oxides and hydrated alumina (Al2 O3 x H2 O).
Electrolysis can be used to extract aluminum from bauxite. The process described below produces
99% pure aluminum:
1. Aluminum is melted along with cryolite (N a3 AlF6 ) which acts as the electrolyte. Cryolite
helps to lower the melting point and dissolve the ore.
2. The anode carbon rods provide sites for the oxidation of O2− and F − ions. Oxygen and
flourine gas are given off at the anodes and also lead to anode consumption.
3. At the cathode cell lining, the Al3+ ions are reduced and metal aluminum deposits on the
lining.
4. The AlF63− electrolyte is stable and remains in its molten state.
The basic electrolytic reactions involved are as follows: At the cathode:
Al+3 + 3e−
→ Al(s)
(99%purity)
At the anode:
2O2−
→ O2 (g) + 4e−
The overall reaction is as follows:
2Al2 O3
→ 4Al + 3O2
The only problem with this process is that the reaction is endothermic and large amounts of
electricity are needed to drive the reaction. The process is therefore very expensive.
17.7
Summary
• An electrochemical reaction is one where either a chemical reaction produces an external
voltage, or where an external voltage causes a chemical reaction to take place.
349
17.7
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
• In a galvanic cell a chemical reaction produces a current in the external circuit. An
example is the zinc-copper cell.
• A galvanic cell has a number of components. It consists of two electrodes, each of
which is placed in a separate beaker in an electrolyte solution. The two electrolytes are
connected by a salt bridge. The electrodes are connected two each other by an external
circuit wire.
• One of the electrodes is the anode, where oxidation takes place. The cathode is the
electrode where reduction takes place.
• In a galvanic cell, the build up of electrons at the anode sets up a potential difference
between the two electrodes, and this causes a current to flow in the external circuit.
• A galvanic cell is therefore an electrochemical cell that uses a chemical reaction between
two dissimilar electrodes dipped in an electrolyte to generate an electric current.
• The standard notation for a galvanic cell such as the zinc-copper cell is as follows:
Zn|Zn2+ ||Cu2+ |Cu
where
| =
a phase boundary (solid/aqueous)
|| =
the salt bridge
• The galvanic cell is used in batteries and in electroplating.
• An electrolytic cell is an electrochemical cell that uses electricity to drive a non-spontaneous
reaction. In an electrolytic cell, electrolysis occurs, which is a process of separating elements and compounds using an electric current.
• One example of an electrolytic cell is the electrolysis of copper sulphate to produce copper
and sulphate ions.
• Different metals have different reaction potentials. The reaction potential of metals (in
other words, their ability to ionise), is recorded in a standard table of electrode potential.
The more negative the value, the greater the tendency of the metal to be oxidised. The
more positive the value, the greater the tendency of the metal to be reduced.
• The values on the standard table of electrode potentials are measured relative to the
standard hydrogen electrode.
• The emf of a cell can be calculated using one of the following equations:
E0(cell) = E0 (right) - E0 (left)
E0(cell) = E0 (reduction half reaction) - E0 (oxidation half reaction)
E0(cell) = E0 (oxidising agent) - E0 (reducing agent)
E0(cell) = E0 (cathode) - E0 (anode)
• It is possible to predict whether a reaction is spontaneous or not, either by looking at the
sign of the cell’s emf or by comparing the electrode potentials of the two half cells.
• It is possible to balance redox equations using the half-reactions that take place.
• There are a number of important applications of electrochemistry. These include electroplating, the production of chlorine and the extraction of aluminium.
Exercise: Summary exercise
350
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
1. For each of the following, say whether the statement is true or false. If it is
false, re-write the statement correctly.
(a) The anode in an electrolytic cell has a negative charge.
(b) The reaction 2KClO3 → 2KCl + 3O2 is an example of a redox reaction.
(c) Lead is a stronger oxidising agent than nickel.
2. For each of the following questions, choose the one correct answer.
(a) Which one of the following reactions is a redox reaction?
i. HCl + N aOH → N aCl + H2 O
ii. AgN O3 + N aI → AgI + N aN O3
iii. 2F eCl3 + 2H2 O + SO2 → H2 SO4 + 2HCl + 2F eCl2
iv. BaCl2 + M gSO4 → M gCl2 + BaSO4
(IEB Paper 2, 2003)
(b) Consider the reaction represented by the following equation:
−
−
Br2(l) + 2Iaq
→ 2Braq
+ I2(s)
Which one of the following statements about this reaction is correct?
i. bromine is oxidised
ii. bromine acts as a reducing agent
iii. the iodide ions are oxidised
iv. iodine acts as a reducing agent
(IEB Paper 2, 2002)
(c) The following equations represent two hypothetical half-reactions:
X2 + 2e− ⇔ 2X − (+1.09 V) and
Y + + e− ⇔ Y (-2.80 V)
Which one of the following substances from these half-reactions has the
greatest tendency to donate electrons?
i. X−
ii. X2
iii. Y
iv. Y+
(d) Which one of the following redox reactions will not occur spontaneously
at room temperature?
i. M n + Cu2+ → M n2+ + Cu
ii. Zn + SO42− + 4H + → Zn2+ + SO2 + 2H2 O
iii. F e3+ + 3N O2 + 3H2 O → F e + 3N O3− + 6H +
iv. 5H2 S + 2M nO4− + 6H + → 5S + 2M n2+ + 8H2 O
(e) Which statement is CORRECT for a Zn-Cu galvanic cell that operates
under standard conditions?
i. The concentration of the Zn2+ ions in the zinc half-cell gradually decreases.
ii. The concentration of the Cu2+ ions in the copper half-cell gradually
increases.
iii. Negative ions migrate from the zinc half-cell to the copper half-cell.
iv. The intensity of the colour of the electrolyte in the copper half-cell
gradually decreases.
(DoE Exemplar Paper 2, 2008)
3. In order to investigate the rate at which a reaction proceeds, a learner places a
beaker containing concentrated nitric acid on a sensitive balance. A few pieces
of copper metal are dropped into the nitric acid.
(a) Use the relevant half-reactions from the table of Standard Reduction Potentials to derive the balanced nett ionic equation for the reaction that
takes place in the beaker.
(b) What chemical property of nitric acid is illustrated by this reaction?
(c) List three observations that this learner would make during the investigation.
351
17.7
17.7
CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12
(IEB Paper 2, 2005)
4. The following reaction takes place in an electrochemical cell:
Cu(s) + 2AgN O3 (aq) → Cu(N O3 )2 (aq) + 2Ag(s)
(a) Give an equation for the oxidation half reaction.
(b) Which metal is used as the anode?
(c) Determine the emf of the cell under standard conditions.
(IEB Paper 2, 2003)
5. The nickel-cadmium (NiCad) battery is small and light and is made in a sealed
unit. It is used in portable appliances such as calculators and electric razors.
The following two half reactions occur when electrical energy is produced by
the cell.
Half reaction 1: Cd(s) + 2OH− (aq) → Cd(OH)2 (s) + 2e−
Half reaction 2: NiO(OH)(s) + H2 O(l) + e− → Ni(OH)2 (s) + OH− (aq)
(a) Which half reaction (1 or 2) occurs at the anode? Give a reason for your
answer.
(b) Which substance is oxidised?
(c) Derive a balanced ionic equation for the overall cell reaction for the discharging process.
(d) Use your result above to state in which direction the cell reaction will
proceed (forward or reverse) when the cell is being charged.
(IEB Paper 2, 2001)
6. An electrochemical cell is constructed by placing a lead rod in a porous pot
containing a solution of lead nitrate (see sketch). The porous pot is then placed
in a large aluminium container filled with a solution of aluminium sulphate. The
lead rod is then connected to the aluminium container by a copper wire and
voltmeter as shown.
V
copper wire
lead rod
porous pot
Al2 (SO4 )3 (aq)
Pb(NO3 )2
(aq)
aluminium container
(a) Define the term reduction.
(b) In which direction do electrons flow in the copper wire? (Al to Pb or Pb
to Al)
(c) Write balanced equations for the reactions that take place at...
i. the anode
ii. the cathode
(d) Write a balanced nett ionic equation for the reaction which takes place in
this cell.
(e) What are the two functions of the porous pot?
(f) Calculate the emf of this cell under standard conditions.
(IEB Paper 2, 2005)
352
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