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”
Permission is granted to copy, distribute and/or modify this document under the
terms of the GNU Free Documentation License, Version 1.2 or any later version
published by the Free Software Foundation; with no Invariant Sections, no FrontCover Texts, and no Back-Cover Texts. A copy of the license is included in the
section entitled “GNU Free Documentation License”.
STOP!!!!
Did you notice the FREEDOMS we’ve granted you?
Our copyright license is different! It grants freedoms
rather than just imposing restrictions like all those other
textbooks you probably own or use.
• We know people copy textbooks illegally but we would LOVE it if you copied
our’s - go ahead copy to your hearts content, legally!
• Publishers’ revenue is generated by controlling the market, we don’t want any
money, go ahead, distribute our books far and wide - we DARE you!
• Ever wanted to change your textbook? Of course you have! Go ahead, change
ours, make your own version, get your friends together, rip it apart and put
it back together the way you like it. That’s what we really want!
• Copy, modify, adapt, enhance, share, critique, adore, and contextualise. Do
it all, do it with your colleagues, your friends, or alone but get involved!
Together we can overcome the challenges our complex and diverse country
presents.
• So what is the catch? The only thing you can’t do is take this book, make
a few changes and then tell others that they can’t do the same with your
changes. It’s share and share-alike and we know you’ll agree that is only fair.
• These books were written by volunteers who want to help support education,
who want the facts to be freely available for teachers to copy, adapt and
re-use. Thousands of hours went into making them and they are a gift to
everyone in the education community.
FHSST Core Team
Mark Horner ; Samuel Halliday ; Sarah Blyth ; Rory Adams ; Spencer Wheaton
FHSST Editors
Jaynie Padayachee ; Joanne Boulle ; Diana Mulcahy ; Annette Nell ; René Toerien ; Donovan
Whitfield
FHSST Contributors
Rory Adams ; Prashant Arora ; Richard Baxter ; Dr. Sarah Blyth ; Sebastian Bodenstein ;
Graeme Broster ; Richard Case ; Brett Cocks ; Tim Crombie ; Dr. Anne Dabrowski ; Laura
Daniels ; Sean Dobbs ; Fernando Durrell ; Dr. Dan Dwyer ; Frans van Eeden ; Giovanni
Franzoni ; Ingrid von Glehn ; Tamara von Glehn ; Lindsay Glesener ; Dr. Vanessa Godfrey ; Dr.
Johan Gonzalez ; Hemant Gopal ; Umeshree Govender ; Heather Gray ; Lynn Greeff ; Dr. Tom
Gutierrez ; Brooke Haag ; Kate Hadley ; Dr. Sam Halliday ; Asheena Hanuman ; Neil Hart ;
Nicholas Hatcher ; Dr. Mark Horner ; Robert Hovden ; Mfandaidza Hove ; Jennifer Hsieh ;
Clare Johnson ; Luke Jordan ; Tana Joseph ; Dr. Jennifer Klay ; Lara Kruger ; Sihle Kubheka ;
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 ;
Nicolette Pekeur ; Sirika Pillay ; Jacques Plaut ; Andrea Prinsloo ; Joseph Raimondo ; Sanya
Rajani ; Prof. Sergey Rakityansky ; Alastair Ramlakan ; Razvan Remsing ; Max Richter ; Sean
Riddle ; Evan Robinson ; Dr. Andrew Rose ; Bianca Ruddy ; Katie Russell ; Duncan Scott ;
Helen Seals ; Ian Sherratt ; Roger Sieloff ; Bradley Smith ; Greg Solomon ; Mike Stringer ;
Shen Tian ; Robert Torregrosa ; Jimmy Tseng ; Helen Waugh ; Dr. Dawn Webber ; Michelle
Wen ; Dr. Alexander Wetzler ; Dr. Spencer Wheaton ; Vivian White ; Dr. Gerald Wigger ;
Harry Wiggins ; Wendy Williams ; Julie Wilson ; Andrew Wood ; Emma Wormauld ; Sahal
Yacoob ; Jean Youssef
Contributors and editors have made a sincere effort to produce an accurate and useful resource.
Should you have suggestions, find mistakes or be prepared to donate material for inclusion,
please don’t hesitate to contact us. We intend to work with all who are willing to help make
this a continuously evolving resource!
www.fhsst.org
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
ix
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
x
CONTENTS
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
xi
CONTENTS
CONTENTS
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
xii
CONTENTS
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
xiii
CONTENTS
CONTENTS
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 16
Reaction Rates - Grade 12
16.1
Introduction
Before we begin this section, it might be useful to think about some different types of reactions
and how quickly or slowly they occur.
Exercise: Thinking about reaction rates
Think about each of the following reactions:
• rusting of metals
• photosynthesis
• weathering of rocks (e.g. limestone rocks being weathered by water)
• combustion
1. For each of the reactions above, write a chemical equation for the reaction that
takes place.
2. How fast is each of these reactions? Rank them in order from the fastest to
the slowest.
3. How did you decide which reaction was the fastest and which was the slowest?
4. Try to think of some other examples of chemical reactions. How fast or slow
is each of these reactions, compared with those listed earlier?
In a chemical reaction, the substances that are undergoing the reaction are called the reactants,
while the substances that form as a result of the reaction are called the products. The reaction
rate describes how quickly or slowly the reaction takes place. So how do we know whether a
reaction is slow or fast? One way of knowing is to look either at how quickly the reactants are
used up during the reaction or at how quickly the product forms. For example, iron and sulfur
react according to the following equation:
F e + S → F eS
In this reaction, we can see the speed of the reaction by observing how long it takes before there
is no iron or sulfur left in the reaction vessel. In other words, the reactants have been used up.
Alternatively, one could see how quickly the iron sulfide product forms. Since iron sulfide looks
very different from either of its reactants, this is easy to do.
In another example:
287
16.1
CHAPTER 16. REACTION RATES - GRADE 12
2M g(s) + O2 → 2M gO(s)
In this case, the reaction rate depends on the speed at which the reactants (solid magnesium
and oxygen gas) are used up, or the speed at which the product (magnesium oxide) is formed.
Definition: Reaction rate
The rate of a reaction describes how quickly reactants are used up or how quickly products
are formed during a chemical reaction. The units used are: moles per second (mols/second
or mol.s−1 ).
The average rate of a reaction is expressed as the number of moles of reactant used up, divided
by the total reaction time, or as the number of moles of product formed, divided by the reaction
time. Using the magnesium reaction shown earlier:
Average reaction rate =
moles M g used
reaction time (s)
or
Average reaction rate =
moles O2 used
reaction time (s)
or
Average reaction rate =
moles M gO produced
reaction time (s)
Worked Example 76: Reaction rates
Question: The following reaction takes place:
4Li + O2 → 2Li2 O
After two minutes , 4 g of Lithium has been used up. Calculate the rate of the
reaction.
Answer
Step 1 : Calculate the number of moles of Lithium that are used up in the
reaction.
n=
4
m
=
= 0.58mols
M
6.94
Step 2 : Calculate the time (in seconds) for the reaction.
t = 2 × 60 = 120s
Step 3 : Calculate the rate of the reaction.
Rate of reaction =
0.58
moles of Lithium used
=
= 0.005
time
120
The rate of the reaction is 0.005 mol.s−1
288
CHAPTER 16. REACTION RATES - GRADE 12
16.2
Exercise: Reaction rates
1. A number of different reactions take place. The table below shows the number
of moles of reactant that are used up in a particular time for each reaction.
Reaction
1
2
3
4
5
Mols used up
2
5
1
3.2
5.9
Time
30 secs
2 mins
1.5 mins
1.5 mins
30 secs
Reaction rate
(a) Complete the table by calculating the rate of each reaction.
(b) Which is the fastest reaction?
(c) Which is the slowest reaction?
2. Two reactions occur simultaneously in separate reaction vessels. The reactions
are as follows:
M g + Cl2 → M gCl2
2N a + Cl2 → 2N aCl
After 1 minute, 2 g of MgCl2 have been produced in the first reaction.
(a) How many moles of MgCl2 are produced after 1 minute?
(b) Calculate the rate of the reaction, using the amount of product that is
produced.
(c) Assuming that the second reaction also proceeds at the same rate, calculate...
i. the number of moles of NaCl produced after 1 minute.
ii. the mass (in g) of sodium that is needed for this reaction to take place.
16.2
Factors affecting reaction rates
Several factors affect the rate of a reaction. It is important to know these factors so that reaction
rates can be controlled. This is particularly important when it comes to industrial reactions, so
that productivity can be maximised. The following are some of the factors that affect the rate
of a reaction.
1. Nature of reactants
Substances have different chemical properties and therefore react differently and at different
rates.
2. Concentration (or pressure in the case of gases)
As the concentration of the reactants increases, so does the reaction rate.
3. Temperature
If the temperature of the reaction increases, so does the rate of the reaction.
4. Catalyst
Adding a catalyst increases the reaction rate.
289
16.2
CHAPTER 16. REACTION RATES - GRADE 12
5. Surface area of solid reactants
Increasing the surface area of the reactants (e.g. if a solid reactant is finely broken up)
will increase the reaction rate.
Activity :: Experiment : The nature of reactants.
Aim:
To determine the effect of the nature of reactants on the rate of a reaction.
Apparatus:
Oxalic acid ((COOH)2 ), iron(II) sulphate (FeSO4 ), potassium permanganate
(KMnO4 ), concentrated sulfuric acid (H2 SO4 ), spatula, test tubes, medicine dropper, glass beaker and glass rod.
H2 SO4
KMnO4
H2 SO4
KMnO4
Test tube 1
Iron (II) sulphate solution
Test tube 2
Oxalic acid solution
Method:
1. In the first test tube, prepare an iron (II) sulphate solution by dissolving about
two spatula points of iron (II) sulphate in 10 cm3 of water.
2. In the second test tube, prepare a solution of oxalic acid in the same way.
3. Prepare a solution of sulfuric acid by adding 1 cm3 of the concentrated acid
to about 4 cm3 of water. Remember always to add the acid to the water, and
never the other way around.
4. Add 2 cm3 of the sulfuric acid solution to the iron(II) and oxalic acid solutions
respectively.
5. Using the medicine dropper, add a few drops of potassium permanganate to
the two test tubes. Once you have done this, observe how quickly each solution
discolours the potassium permanganate solution.
Results:
• You should have seen that the oxalic acid solution discolours the potassium
permanganate much more slowly than the iron(II) sulphate.
2+
• It is the oxalate ions (C2 O2−
ions that cause the discolouration.
4 ) and the Fe
It is clear that the Fe2+ ions act much more quickly than the C2 O2−
4 ions. The
reason for this is that there are no covalent bonds to be broken in the ions
before the reaction can take place. In the case of the oxalate ions, covalent
bonds between carbon and oxygen atoms must be broken first.
Conclusions:
The nature of the reactants can affect the rate of a reaction.
290
CHAPTER 16. REACTION RATES - GRADE 12
16.2
teresting Oxalic acids are abundant in many plants. The leaves of the tea plant (Camellia
Interesting
Fact
Fact
sinensis) contain very high concentrations of oxalic acid relative to other plants.
Oxalic acid also occurs in small amounts in foods such as parsley, chocolate, nuts
and berries. Oxalic acid irritates the lining of the gut when it is eaten, and can
be fatal in very large doses.
Activity :: Experiment : Surface area and reaction rates.
Marble (CaCO3 ) reacts with hydrochloric acid (HCl) to form calcium chloride,
water and carbon dioxide gas according to the following equation:
CaCO3 + 2HCl → CaCl2 + H2 O + CO2
Aim:
To determine the effect of the surface area of reactants on the rate of the reaction.
Apparatus:
2 g marble chips, 2 g powdered marble, hydrochloric acid, beaker, two test tubes.
beaker containing dilute
hydrochloric acid
Test tube 1
marble chips
Test tube 2
powdered marble
Method:
1. Prepare a solution of hydrochloric acid in the beaker by adding 2 cm3 of the
concentrated solution to 20 cm3 of water.
2. Place the marble chips and powdered marble into separate test tubes.
3. Add 10 cm3 of the dilute hydrochloric acid to each of the test tubes and observe
the rate at which carbon dioxide gas is produced.
Results:
• Which reaction proceeds the fastest?
• Can you explain this?
Conclusions:
The reaction with powdered marble is the fastest. The smaller the pieces of
marble are, the greater the surface area for the reaction to take place. The greater
the surface area of the reactants, the faster the reaction rate will be.
291
16.2
CHAPTER 16. REACTION RATES - GRADE 12
Activity :: Experiment : Reactant concentration and reaction rate.
Aim:
To determine the effect of reactant concentration on reaction rate.
Apparatus:
Concentrated hydrochloric acid (HCl), magnesium ribbon, two beakers, two test
tubes, measuring cylinder.
Method:
1. Prepare a solution of dilute hydrochloric acid in one of the beakers by diluting
1 part concentrated acid with 10 parts water. For example, if you measure 1
cm3 of concentrated acid in a measuring cylinder and pour it into a beaker, you
will need to add 10 cm3 of water to the beaker as well. In the same way, if you
pour 2 cm3 of concentrated acid into a beaker, you will need to add 20 cm3 of
water. Both of these are 1:10 solutions. Pour 10 cm3 of the 1:10 solution into
a test tube and mark it ’A’. Remember to add the acid to the water, and not
the other way around.
2. Prepare a second solution of dilute hydrochloric acid by diluting 1 part concentrated acid with 20 parts water. Pour 10cm3 of this 1:20 solution into a second
test tube and mark it ’B’.
3. Take two pieces of magnesium ribbon of the same length. At the same time,
put one piece of magnesium ribbon into test tube A and the other into test
tube B, and observe closely what happens.
Mg ribbon
Mg ribbon
Test tube A
1:10 HCl solution
Test tube B
1:20 HCl solution
The equation for the reaction is:
2HCl + M g → M gCl2 + H2
Results:
• Which of the two solutions is more concentrated, the 1:10 or 1:20 hydrochloric
acid solution?
• In which of the test tubes is the reaction the fastest? Suggest a reason for this.
• How can you measure the rate of this reaction?
• What is the gas that is given off?
• Why was it important that the same length of magnesium ribbon was used for
each reaction?
Conclusions:
The 1:10 solution is more concentrated and this reaction therefore proceeds
faster. The greater the concentration of the reactants, the faster the rate of the
reaction. The rate of the reaction can be measured by the rate at which hydrogen
gas is produced.
292
CHAPTER 16. REACTION RATES - GRADE 12
16.3
Activity :: Group work : The effect of temperature on reaction rate
1. In groups of 4-6, design an experiment that will help you to see the effect of
temperature on the reaction time of 2 cm of magnesium ribbon and 20 ml of
vinegar. During your group discussion, you should think about the following:
• What equipment will you need?
• How will you conduct the experiment to make sure that you are able to
compare the results for different temperatures?
• How will you record your results?
• What safety precautions will you need to take when you carry out this
experiment?
2. Present your experiment ideas to the rest of the class, and give them a chance
to comment on what you have done.
3. Once you have received feedback, carry out the experiment and record your
results.
4. What can you conclude from your experiment?
16.3
Reaction rates and collision theory
It should be clear now that the rate of a reaction varies depending on a number of factors. But
how can we explain why reactions take place at different speeds under different conditions? Why,
for example, does an increase in the surface area of the reactants also increase the rate of the
reaction? One way to explain this is to use collision theory.
For a reaction to occur, the particles that are reacting must collide with one another. Only a
fraction of all the collisions that take place actually cause a chemical change. These are called
’successful’ collisions. When there is an increase in the concentration of reactants, the chance
that reactant particles will collide with each other also increases because there are more particles
in that space. In other words, the collision frequency of the reactants increases. The number of
successful collisions will therefore also increase, and so will the rate of the reaction. In the same
way, if the surface area of the reactants increases, there is also a greater chance that successful
collisions will occur.
Definition: Collision theory
Collision theory is a theory that explains how chemical reactions occur and why reaction
rates differ for different reactions. The theory assumes that for a reaction to occur the
reactant particles must collide, but that only a certain fraction of the total collisions, the
effective collisions, actually cause the reactant molecules to change into products. This is
because only a small number of the molecules have enough energy and the right orientation
at the moment of impact to break the existing bonds and form new bonds.
When the temperature of the reaction increases, the average kinetic energy of the reactant
particles increases and they will move around much more actively. They are therefore more likely
to collide with one another (Figure 16.1). Increasing the temperature also increases the number
of particles whose energy will be greater than the activation energy for the reaction (refer section
16.5).
293
CHAPTER 16. REACTION RATES - GRADE 12
A
A
B
B
A
B
B
A
B
B
16.3
B
A
A
B
A
A
B
B
B
A
A
A
B
Low Temperature
A
High Temperature
Figure 16.1: An increase in the temperature of a reaction increases the chances that the reactant
particles (A and B) will collide because the particles have more energy and move around more.
Exercise: Rates of reaction
Hydrochloric acid and calcium carbonate react according to the following equation:
CaCO3 + 2HCl → CaCl2 + H2 O + CO2
The volume of carbon dioxide that is produced during the reaction is measured
at different times. The results are shown in the table below.
Time (mins)
1
2
3
4
5
6
7
8
9
10
Volume of CO2 produced (cm3 )
14
26
36
44
50
58
65
70
74
77
Note: On a graph of production against time, it is the gradient of the graph that
shows the rate of the reaction.
Questions:
1. Use the data in the table to draw a graph showing the volume of gas that is
produced in the reaction, over a period of 10 minutes.
2. At which of the following times is the reaction fastest? Time = 1 minute; time
= 6 minutes or time = 8 minutes.
3. Suggest a reason why the reaction slows down over time.
4. Use the graph to estimate the volume of gas that will have been produced after
11 minutes.
5. After what time do you think the reaction will stop?
6. If the experiment was repeated using a more concentrated hydrochloric acid
solution...
(a) would the rate of the reaction increase or decrease from the one shown in
the graph?
(b) draw a rough line on the graph to show how you would expect the reaction
to proceed with a more concentrated HCl solution.
294
CHAPTER 16. REACTION RATES - GRADE 12
16.4
16.4
Measuring Rates of Reaction
How the rate of a reaction is measured will depend on what the reaction is, and what product
forms. Look back to the reactions that have been discussed so far. In each case, how was the
rate of the reaction measured? The following examples will give you some ideas about other
ways to measure the rate of a reaction:
• Reactions that produce hydrogen gas:
When a metal dissolves in an acid, hydrogen gas is produced. A lit splint can be used
to test for hydrogen. The ’pop’ sound shows that hydrogen is present. For example,
magnesium reacts with sulfuric acid to produce magnesium sulphate and hydrogen.
M g(s) + H2 SO4 → M gSO4 + H2
• Reactions that produce carbon dioxide:
When a carbonate dissolves in an acid, carbon dioxide gas is produced. When carbon
dioxide is passes through limewater, it turns the limewater milky. This is the test for the
presence of carbon dioxide. For example, calcium carbonate reacts with hydrochloric acid
to produce calcium chloride, water and carbon dioxide.
CaCO3 (s) + 2HCl(aq) → CaCl2 (aq) + H2 O(l) + CO2 (g)
• Reactions that produce gases such as oxygen or carbon dioxide:
Hydrogen peroxide decomposes to produce oxygen. The volume of oxygen produced can
be measured using the gas syringe method (figure 16.2). The gas collects in the syringe,
pushing out against the plunger. The volume of gas that has been produced can be read
from the markings on the syringe. For example, hydrogen peroxide decomposes in the
presence of a manganese(IV) oxide catalyst to produce oxygen and water.
2H2 O2 (aq) → 2H2 O(l) + O2 (g)
Gas Syringe System
Gas
[glassType=erlen,niveauLiquide1=40,tubeCoude]
Reactants
Figure 16.2: Gas Syringe Method
• Precipitate reactions:
In reactions where a precipitate is formed, the amount of precipitate formed in a period of
time can be used as a measure of the reaction rate. For example, when sodium thiosulphate
reacts with an acid, a yellow precipitate of sulfur is formed. The reaction is as follows:
N a2 S2 O3 (aq) + 2HCl(aq) → 2N aCl(aq) + SO2 (aq) + H2 O(l) + S(s)
295
16.4
CHAPTER 16. REACTION RATES - GRADE 12
One way to estimate the rate of this reaction is to carry out the investigation in a conical
flask and to place a piece of paper with a black cross underneath the bottom of the flask.
At the beginning of the reaction, the cross will be clearly visible when you look into the
flask (figure 16.3). However, as the reaction progresses and more precipitate is formed,
the cross will gradually become less clear and will eventually disappear altogether. Noting
the time that it takes for this to happen will give an idea of the reaction rate. Note that
it is not possible to collect the SO2 gas that is produced in the reaction, because it is very
soluble in water.
[glassType=erlen,niveauLiquide1=40]
Figure 16.3: At the beginning of the reaction beteen sodium thiosulphate and hydrochloric acid,
when no precipitate has been formed, the cross at the bottom of the conical flask can be clearly
seen.
• Changes in mass:
The rate of a reaction that produces a gas can also be measured by calculating the mass
loss as the gas is formed and escapes from the reaction flask. This method can be used for
reactions that produce carbon dioxide or oxygen, but are not very accurate for reactions
that give off hydrogen because the mass is too low for accuracy. Measuring changes in
mass may also be suitable for other types of reactions.
Activity :: Experiment : Measuring reaction rates
Aim:
To measure the effect of concentration on the rate of a reaction.
Apparatus:
• 300 cm3 of sodium thiosulphate (Na2 S2 O3 ) solution. Prepare a solution of
sodium thiosulphate by adding 12 g of Na2 S2 O3 to 300 cm3 of water. This is
solution ’A’.
• 300 cm3 of water
• 100 cm3 of 1:10 dilute hydrochloric acid. This is solution ’B’.
• Six 100 cm3 glass beakers
• Measuring cylinders
• Paper and marking pen
• Stopwatch or timer
Method:
One way to measure the rate of this reaction is to place a piece of paper with a
cross underneath the reaction beaker to see how quickly the cross is made invisible
by the formation of the sulfur precipitate.
296
CHAPTER 16. REACTION RATES - GRADE 12
16.5
1. Set up six beakers on a flat surface and mark them from 1 to 6. Under each
beaker you will need to place a piece of paper with a large black cross.
2. Pour 60 cm3 solution A into the first beaker and add 20 cm3 of water
3. Use the measuring cylinder to measure 10 cm3 HCl. Now add this HCl to the
solution that is already in the first beaker (NB: Make sure that you always clean
out the measuring cylinder you have used before using it for another chemical).
4. Using a stopwatch with seconds, record the time it takes for the precipitate
that forms to block out the cross.
5. Now measure 50 cm3 of solution A into the second beaker and add 30 cm3 of
water. To this second beaker, add 10 cm3 HCl, time the reaction and record
the results as you did before.
6. Continue the experiment by diluting solution A as shown below.
Beaker
1
2
3
4
5
6
Solution
(cm3 )
60
50
40
30
20
10
A
Water (cm3 )
20
30
40
50
60
70
Solution
(cm3 )
10
10
10
10
10
10
B
Time
(s)
The equation for the reaction between sodium thiosulphate and hydrochloric acid
is:
N a2 S2 O3 (aq) + 2HCl(aq) → 2N aCl(aq) + SO2 (aq) + H2 O(l) + S(s)
Results:
• Calculate the reaction rate in each beaker. This can be done using the following
equation:
Rate of reaction =
1
time
• Represent your results on a graph. Concentration will be on the x-axis and
reaction rate on the y-axis. Note that the original volume of Na2 S2 O3 can be
used as a measure of concentration.
• Why was it important to keep the volume of HCl constant?
• Describe the relationship between concentration and reaction rate.
Conclusions:
The rate of the reaction is fastest when the concentration of the reactants was
the highest.
16.5
Mechanism of reaction and catalysis
Earlier it was mentioned that it is the collision of particles that causes reactions to occur and
that only some of these collisions are ’successful’. This is because the reactant particles have a
wide range of kinetic energy, and only a small fraction of the particles will have enough energy
to actually break bonds so that a chemical reaction can take place. The minimum energy that
is needed for a reaction to take place is called the activation energy. For more information on
the energy of reactions, refer to chapter 14.
297
16.5
CHAPTER 16. REACTION RATES - GRADE 12
Definition: Activation energy
The energy that is needed to break the bonds in reactant molecules so that a chemical
reaction can proceed.
Probability of particle with
that KE
Even at a fixed temperature, the energy of the particles varies, meaning that only some of them
will have enough energy to be part of the chemical reaction, depending on the activation energy
for that reaction. This is shown in figure 16.4. Increasing the reaction temperature has the effect
of increasing the number of particles with enough energy to take part in the reaction, and so the
reaction rate increases.
The distribution of particle kinetic
energies at a fixed temperature.
Average KE
Kinetic Energy of Particle (KE)
Figure 16.4: The distribution of particle kinetic energies at a fixed temperature
A catalyst functions slightly differently. The function of a catalyst is to lower the activation
energy so that more particles now have enough energy to react. The catalyst itself is not changed
during the reaction, but simply provides an alternative pathway for the reaction, so that it needs
less energy. Some metals e.g. platinum, copper and iron can act as catalysts in certain reactions.
In our own human bodies, enzymes are catalysts that help to speed up biological reactions. Catalysts generally react with one or more of the reactants to form a chemical intermediate which
then reacts to form the final product. The chemical intermediate is sometimes called the activated complex.
The following is an example of how a reaction that involves a catalyst might proceed. C represents the catalyst, A and B are reactants and D is the product of the reaction of A and B.
Step 1: A + C → AC
Step 2: B + AC → ABC
Step 3: ABC → CD
Step 4: CD → C + D
In the above, ABC represents the intermediate chemical. Although the catalyst (C) is consumed
by reaction 1, it is later produced again by reaction 4, so that the overall reaction is as follows:
A+B+C→D+C
You can see from this that the catalyst is released at the end of the reaction, completely unchanged.
Definition: Catalyst
A catalyst speeds up a chemical reaction, without being altered in any way. It increases the
reaction rate by lowering the activation energy for a reaction.
298
CHAPTER 16. REACTION RATES - GRADE 12
16.5
Energy diagrams are useful to illustrate the effect of a catalyst on reaction rates. Catalysts
decrease the activation energy required for a reaction to proceed (shown by the smaller ’hump’
on the energy diagram in figure 16.5), and therefore increase the reaction rate.
activated complex
Potential energy
activation
energy
products
activation energy
with a catalyst
reactants
Time
Figure 16.5: The effect of a catalyst on the activation energy of a reaction
Activity :: Experiment : Catalysts and reaction rates
Aim:
To determine the effect of a catalyst on the rate of a reaction
Apparatus:
Zinc granules, 0.1 M hydrochloric acid, copper pieces, one test tube and a glass
beaker.
Method:
1. Place a few of the zinc granules in the test tube.
2. Measure the mass of a few pieces of copper and keep them separate from the
rest of the copper.
3. Add about 20 cm3 of HCl to the test tube. You will see that a gas is released.
Take note of how quickly or slowly this gas is released. Write a balanced
equation for the chemical reaction that takes place.
4. Now add the copper pieces to the same test tube. What happens to the rate
at which the gas is produced?
5. Carefully remove the copper pieces from the test tube (do not get HCl on your
hands), rinse them in water and alcohol and then weigh them again. Has the
mass of the copper changed since the start of the experiment?
Results:
During the reaction, the gas that is released is hydrogen. The rate at which the
hydrogen is produced increases when the copper pieces (the catalyst) are added.
The mass of the copper does not change during the reaction.
Conclusions:
The copper acts as a catalyst during the reaction. It speeds up the rate of the
reaction, but is not changed in any way itself.
299
16.6
CHAPTER 16. REACTION RATES - GRADE 12
Exercise: Reaction rates
1. For each of the following, say whether the statement is true or false. If it is
false, re-write the statement correctly.
(a) A catalyst increases the energy of reactant molecules so that a chemical
reaction can take place.
(b) Increasing the temperature of a reaction has the effect of increasing the
number of reactant particles that have more energy that the activation
energy.
(c) A catalyst does not become part of the final product in a chemical reaction.
2. 5 g of zinc granules are added to 400 cm3 of 0.5 mol.dm−3 hydrochloric acid.
To investigate the rate of the reaction, the change in the mass of the flask
containing the zinc and the acid was measured by placing the flask on a direct
reading balance. The reading on the balance shows that there is a decrease
in mass during the reaction. The reaction which takes place is given by the
following equation:
Zn(s) + 2HCl(aq) → ZnCl2 (aq) + H2 (g)
(a) Why is there a decrease in mass during the reaction?
(b) The experiment is repeated, this time using 5 g of powdered zinc instead
of granulated zinc. How will this influence the rate of the reaction?
(c) The experiment is repeated once more, this time using 5 g of granulated
zinc and 600 cm3 of 0.5 mol.dm−3 hydrochloric acid. How does the rate
of this reaction compare to the original reaction rate?
(d) What effect would a catalyst have on the rate of this reaction?
(IEB Paper 2 2003)
3. Enzymes are catalysts. Conduct your own research to find the names of common enzymes in the human body and which chemical reactions they play a role
in.
4. 5 g of calcium carbonate powder reacts with 20 cm3 of a 0.1 mol.dm−3 solution
of hydrochloric acid. The gas that is produced at a temperature of 250 C is
collected in a gas syringe.
(a) Write a balanced chemical equation for this reaction.
(b) The rate of the reaction is determined by measuring the volume of gas
thas is produced in the first minute of the reaction. How would the rate
of the reaction be affected if:
i. a lump of calcium carbonate of the same mass is used
ii. 40 cm3 of 0.1 mol.dm−3 hydrochloric acid is used
16.6
Chemical equilibrium
Having looked at factors that affect the rate of a reaction, we now need to ask some important
questions. Does a reaction always proceed in the same direction or can it be reversible? In other
words, is it always true that a reaction proceeds from reactants to products, or is it possible that
sometimes, the reaction will reverse and the products will be changed back into the reactants?
And does a reaction always run its full course so that all the reactants are used up, or can a
reaction reach a point where reactants are still present, but there does not seem to be any further
change taking place in the reaction? The following demonstration might help to explain this.
300
CHAPTER 16. REACTION RATES - GRADE 12
16.6
Activity :: Demonstration : Liquid-vapour phase equilibrium
Apparatus and materials:
2 beakers; water; bell jar
Method:
1. Half fill two beakers with water and mark the level of the water in each case.
2. Cover one of the beakers with a bell jar.
3. Leave the beakers and, over the course of a day or two, observe how the water
level in the two beakers changes. What do you notice? Note: You could speed
up this demonstration by placing the two beakers over a bunsen burner to heat
the water. In this case, it may be easier to cover the second beaker with a glass
cover.
Observations:
You should notice that in the beaker that is uncovered, the water level drops
quickly because of evaporation. In the beaker that is covered, there is an initial drop
in the water level, but after a while evaporation appears to stop and the water level
in this beaker is higher than that in the one that is open. Note that the diagram
below shows the situation ate time=0.
bell jar
= condensation
= evaporation
Discussion:
In the first beaker, liquid water becomes water vapour as a result of evaporation
and the water level drops. In the second beaker, evaporation also takes place.
However, in this case, the vapour comes into contact with the surface of the bell
jar and it cools and condenses to form liquid water again. This water is returned to
the beaker. Once condensation has begun, the rate at which water is lost from the
beaker will start to decrease. At some point, the rate of evaporation will be equal to
the rate of condensation above the beaker, and there will be no change in the water
level in the beaker. This can be represented as follows:
liquid ⇔ vapour
In this example, the reaction (in this case, a change in the phase of water) can
proceed in either direction. In one direction there is a change in phase from liquid to
vapour. But the reverse can also take place, when vapour condenses to form water
again.
In a closed system it is possible for reactions to be reversible, such as in the demonstration
above. In a closed system, it is also possible for a chemical reaction to reach equilibrium. We
will discuss these concepts in more detail.
301
16.6
CHAPTER 16. REACTION RATES - GRADE 12
16.6.1
Open and closed systems
An open system is one in which matter or energy can flow into or out of the system. In the
liquid-vapour demonstration we used, the first beaker was an example of an open system because
the beaker could be heated or cooled (a change in energy ), and water vapour (the matter ) could
evaporate from the beaker.
A closed system is one in which energy can enter or leave, but matter cannot. The second
beaker covered by the bell jar is an example of a closed system. The beaker can still be heated or
cooled, but water vapour cannot leave the system because the bell jar is a barrier. Condensation
changes the vapour to liquid and returns it to the beaker. In other words, there is no loss of
matter from the system.
Definition: Open and closed systems
An open system is one whose borders allow the movement of energy and matter into and
out of the system. A closed system is one in which only energy can be exchanged, but not
matter.
16.6.2
Reversible reactions
Some reactions can take place in two directions. In one direction the reactants combine to form
the products. This is called the forward reaction. In the other, the products react to form
reactants again. This is called the reverse reaction. A special double-headed arrow is used to
show this type of reversible reaction:
XY + Z ⇔ X + Y Z
So, in the following reversible reaction:
H2 (g) + I2 (g) ⇔ 2HI(g)
The forward reaction is H2 (g) + I2 (g) → 2HI(g). The reverse reaction is 2HI(g) → H2 (g) + I2 (g).
Definition: A reversible reaction
A reversible reaction is a chemical reaction that can proceed in both the forward and reverse
directions. In other words, the reactant and product of one reaction may reverse roles.
Activity :: Demonstration : The reversibility of chemical reactions
Apparatus and materials:
Lime water (Ca(OH)2 ); calcium carbonate (CaCO3 ); hydrochloric acid; 2 test
tubes with rubber stoppers; delivery tube; retort stand and clamp; bunsen burner.
Method and observations:
1. Half-fill a test tube with clear lime water (Ca(OH)2 ).
2. In another test tube, place a few pieces of calcium carbonate (CaCO3 ) and
cover the pieces with dilute hydrochloric acid. Seal the test tube with a rubber
stopper and delivery tube.
3. Place the other end of the delivery tube into the test tube containing the lime
water so that the carbon dioxide that is produced from the reaction between calcium carbonate and hydrochloric acid passes through the lime water. Observe
what happens to the appearance of the lime water.
The equation for the reaction that takes place is:
302
CHAPTER 16. REACTION RATES - GRADE 12
16.6
Ca(OH)2 + CO2 → CaCO3 + H2 O
CaCO3 is insoluble and it turns the limewater milky.
4. Allow the reaction to proceed for a while so that carbon dioxide continues to
pass through the limewater. What do you notice? The equation for the reaction
that takes place is:
CaCO3 (s) + H2 O + CO2 → Ca(HCO3 )2
In this reaction, calcium carbonate becomes one of the reactants to produce
hydrogen carbonate (Ca(HCO3 )2 ) and so the solution becomes clear again.
5. Heat the solution in the test tube over a bunsen burner. What do you observe?
You should see bubbles of carbon dioxide appear and the limewater turns milky
again. The reaction that has taken place is:
Ca(HCO3 )2 → CaCO3 (s) + H2 O + CO2
delivery tube
rubber stopper
rubber stopper
[glassType=tube,bouchon=true,niveauLiquide1=30]
[glassType=tube,bouchon=true,niveauLiquide1=60]
calcium carbonate &
hydrochloric acid
limewater
Discussion:
• If you look at the last two equations you will see that the one is the reverse of
the other. In other words, this is a reversible reaction and can be written as
follows:
CaCO3 (s) + H2 O + CO2 ⇔ Ca(HCO3 )2
• Is the forward reaction endothermic or exothermic? Is the reverse reaction
endothermic or exothermic? You should have noticed that the reverse reaction only took place when the solution was heated. Sometimes, changing the
temperature of a reaction can change its direction.
16.6.3
Chemical equilibrium
Using the same reversible reaction that we used in an earlier example:
H2 (g) + I2 (g) ⇔ 2HI(g)
The forward reaction is:
H2 + I2 → 2HI
303
16.7
CHAPTER 16. REACTION RATES - GRADE 12
The reverse reaction is:
2HI → H2 + I2
When the rate of the forward reaction and the reverse reaction are equal, the system is said to
be in equilbrium. Figure 16.6 shows this. Initially (time = 0), the rate of the forward reaction
is high and the rate of the reverse reaction is low. As the reaction proceeds, the rate of the
forward reaction decreases and the rate of the reverse reaction increases, until both occur at the
same rate. This is called equilibrium.
Rate of Reaction
Although it is not always possible to observe any macroscopic changes, this does not mean
that the reaction has stopped. The forward and reverse reactions continue to take place and
so microscopic changes still occur in the system. This state is called dynamic equilibrium. In
the liquid-vapour phase equilibrium demonstration, dynamic equilibrium was reached when there
was no observable change in the level of the water in the second beaker even though evaporation
and condensation continued to take place.
H2 +I2 →2HI
equilibrium
2HI→H2 +I2
Time
Figure 16.6: The change in rate of forward and reverse reactions in a closed system
There are, however, a number of factors that can change the chemical equilibrium of a reaction. Changing the concentration, the temperature or the pressure of a reaction can affect
equilibrium. These factors will be discussed in more detail later in this chapter.
Definition: Chemical equilibrium
Chemical equilibrium is the state of a chemical reaction, where the concentrations of the
reactants and products have no net change over time. Usually, this state results when the
forward chemical reactions proceed at the same rate as their reverse reactions.
16.7
The equilibrium constant
Definition: Equilibrium constant
The equilibrium constant (Kc ), relates to a chemical reaction at equilibrium. It can be
calculated if the equilibrium concentration of each reactant and product in a reaction at
equilibrium is known.
304
CHAPTER 16. REACTION RATES - GRADE 12
16.7.1
16.7
Calculating the equilibrium constant
Consider the following generalised reaction which takes place in a closed container at a constant
temperature:
A+B ⇔C +D
We know from section 16.2 that the rate of the forward reaction is directly proportional to the
concentration of the reactants. In other words, as the concentration of the reactants increases,
so does the rate of the forward reaction. This can be shown using the following equation:
Rate of forward reaction ∝ [A][B]
or
Rate of forward reaction = k1 [A][B]
Similarly, the rate of the reverse reaction is directly proportional to the concentration of the
products. This can be shown using the following equation:
Rate of reverse reaction ∝ [C][D]
or
Rate of reverse reaction = k2 [C][D]
At equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. This
can be shown using the following equation:
k1 [A][B] = k2 [C][D]
or
k1
[C][D]
=
k2
[A][B]
or, if the constants k1 and k2 are simplified to a single constant, the equation becomes:
kc =
[C][D]
[A][B]
A more general form of the equation for a reaction at chemical equilibrium is:
aA + bB ⇔ cC + dD
where A and B are reactants, C and D are products and a, b, c, and d are the coefficients of
the respective reactants and products. A more general formula for calculating the equilibrium
constant is therefore:
kc =
[C]c [D]d
[A]a [B]b
It is important to note that if a reactant or a product in a chemical reaction is in either the
liquid or solid phase, the concentration stays constant during the reaction. Therefore, these
values can be left out of the equation to calculate kc . For example, in the following reaction:
C(s) + H2 O(g) ⇔ CO(g) + H2 (g)
305
16.7
CHAPTER 16. REACTION RATES - GRADE 12
kc =
[CO][H2 ]
[H2 O]
Important:
1. The constant kc is affected by temperature and so, if the values of k c are being
compared for different reactions, it is important that all the reactions have taken
place at the same temperature.
2. kc values do not have units. If you look at the equation, the units all cancel each
other out.
16.7.2
The meaning of kc values
The formula for kc has the concentration of the products in the numerator and the concentration
of reactants in the denominator. So a high kc value means that the concentration of products
is high and the reaction has a high yield. We can also say that the equilibrium lies far to the
right. The opposite is true for a low kc value. A low kc value means that, at equilibrium, there
are more reactants than products and therefore the yield is low. The equilibrium for the reaction
lies far to the left.
Important: Calculations made easy
When you are busy with calculations that involve the equilibrium constant, the following tips
may help:
1. Make sure that you always read the question carefully to be sure of what you are being asked
to calculate. If the equilibrium constant is involved, make sure that the concentrations you
use are the concentrations at equilibrium, and not the concentrations or quantities that
are present at some other time in the reaction.
2. When you are doing more complicated calculations, it sometimes helps to draw up a table
like the one below and fill in the mole values that you know or those you can calculate.
This will give you a clear picture of what is happening in the reaction and will make sure
that you use the right values in your calculations.
Reactant 1
Reactant 2
Start of reaction
Used up
Produced
Equilibrium
Worked Example 77: Calculating kc
Question: For the reaction:
SO2 (g) + N O2 (g) → N O(g) + SO3 (g)
306
Product 1
CHAPTER 16. REACTION RATES - GRADE 12
the concentration of the reagents is as follows:
[SO3 ] = 0.2 mol.dm−3
[NO2 ] = 0.1 mol.dm−3
[NO] = 0.4 mol.dm−3
[SO2 ] = 0.2 mol.dm−3
Calculate the value of kc .
Answer
Step 1 : Write the equation for kc
kc =
[N O][SO3 ]
[SO2 ][N O2 ]
Step 2 : Fill in the values you know for this equation and calculate kc
kc =
(0.4 × 0.2)
=4
(0.2 × 0.1)
Worked Example 78: Calculating reagent concentration
Question: For the reaction:
S(s) + O2 (g) ⇔ SO2 (g)
1. Write an equation for the equilibrium constant.
2. Calculate the equilibrium concentration of O2 if Kc=6 and [SO2 ] = 3mol.dm−3
at equilibrium.
Answer
Step 1 : Write the equation for kc
kc =
[SO2 ]
[O2 ]
(Sulfur is left out of the equation because it is a solid and its concentration stays
constant during the reaction)
Step 2 : Re-arrange the equation so that oxygen is on its own on one side
of the equation
[O2 ] =
[SO2 ]
kc
Step 3 : Fill in the values you know and calculate [O2 ]
[O2 ] =
3mol.dm−3
= 0.5mol.dm−3
6
Worked Example 79: Equilibrium calculations
Question: Initially 1.4 moles of NH3 (g) is introduced into a sealed 2.0 dm−3 reaction
vessel. The ammonia decomposes when the temperature is increased to 600K and
reaches equilibrium as follows:
307
16.7
16.7
CHAPTER 16. REACTION RATES - GRADE 12
2N H3 (g) ⇔ N2 (g) + 3H2 (g)
When the equilibrium mixture is analysed, the concentration of NH3 (g) is 0.3 mol.dm−3
1. Calculate the concentration of N2 (g) and H2 (g) in the equilibrium mixture.
2. Calculate the equilibrium constant for the reaction at 900 K.
Answer
Step 1 : Calculate the number of moles of NH3 at equilibrium.
c=
n
V
Therefore,
n = c × V = 0.3 × 2 = 0.6mol
Step 2 : Calculate the number of moles of ammonia that react (are ’used
up’) in the reaction.
Moles used up = 1.4 - 0.6 = 0.8 moles
Step 3 : Calculate the number of moles of product that are formed.
Remember to use the mole ratio of reactants to products to do this. In this case,
the ratio of NH3 :N2 :H2 = 2:1:3. Therefore, if 0.8 moles of ammonia are used up in
the reaction, then 0.4 moles of nitrogen are produced and 1.2 moles of hydrogen are
produced.
Step 4 : Complete the following table
Start of reaction
Used up
Produced
Equilibrium
NH3
1.4
0.8
0
0.6
N2
0
0
0.4
0.4
H2
0
0
1.2
1.2
Step 5 : Using the values in the table, calculate [N2 ] and [H2 ]
[N2 ] =
0.4
n
=
= 0.2 mol.dm−3
V
2
[H2 ] =
n
1.2
=
= 0.6 mol.dm−3
V
2
Step 6 : Calculate kc
kc =
(0.6)3 (0.2)
[H2 ]3 [N2 ]
=
= 0.48
[N H3 ]2
(0.3)2
Worked Example 80: Calculating kc
Question: Hydrogen and iodine gas react according to the following equation:
H2 (g) + I2 (g) ⇔ 2HI(g)
When 0.496 mol H2 and 0.181 mol I2 are heated at 450oC in a 1 dm3 container, the
equilibrium mixture is found to contain 0.00749 mol I2 . Calculate the equilibrium
constant for the reaction at 450o C.
308
CHAPTER 16. REACTION RATES - GRADE 12
16.7
Answer
Step 1 : Calculate the number of moles of iodine used in the reaction.
Moles of iodine used = 0.181 - 0.00749 = 0.1735 mol
Step 2 : Calculate the number of moles of hydrogen that are used up in the
reaction.
The mole ratio of hydrogen:iodine = 1:1, therefore 0.1735 moles of hydrogen must
also be used up in the reaction.
Step 3 : Calculate the number of moles of hydrogen iodide that are produced.
The mole ratio of H2 :I2 :HI = 1:1:2, therefore the number of moles of HI produced
is 0.1735 × 2 = 0.347 mol.
So far, the table can be filled in as follows:
Start of reaction
Used up
Produced
Equilibrium
H2 (g)
0.496
0.1735
0
0.3225
I2
0.181
0.1735
0
0.0075
2HI
0
0
0.347
0.347
Step 4 : Calculate the concentration of each of the reactants and products
at equilibrium.
n
V
Therefore the equilibrium concentrations are as follows:
[H2 ] = 0.3225 mol.dm−3
[I2 ] = 0.0075 mol.dm−3
[HI] = 0.347 mol.dm−3
c=
Step 5 : Calculate kc
kc =
[HI]
0.347
=
= 143.47
[H2 ][I2 ]
0.3225 × 0.0075
Exercise: The equilibrium constant
1. Write the equilibrium constant expression, Kc for the following reactions:
(a) 2NO(g) + Cl2 (g) ⇔ 2NOCl
(b) H2 (g) + I2 (g) ⇔ 2HI(g)
2. The following reaction takes place:
Fe3+ (aq) + 4Cl− ⇔ FeCl−
4 (aq)
Kc for the reaction is 7.5 × 10−2 mol.dm−3 . At equilibrium, the concentration
−4
of FeCl−
mol.dm−3 and the concentration of free iron (Fe3+ )
4 is 0.95 × 10
−3
is 0.2 mol.dm . Calculate the concentration of chloride ions at equilibrium.
3. Ethanoic acid (CH3 COOH) reacts with ethanol (CH3 CH2 OH) to produce ethyl
ethanoate and water. The reaction is:
CH3 COOH + CH3 CH2 OH → CH3 COOCH2 CH3 + H2 O
At the beginning of the reaction, there are 0.5 mols of ethanoic acid and 0.5
mols of ethanol. At equilibrium, 0.3 mols of ethanoic acid was left unreacted.
The volume of the reaction container is 2 dm3 . Calculate the value of Kc .
309
16.8
16.8
CHAPTER 16. REACTION RATES - GRADE 12
Le Chatelier’s principle
A number of factors can influence the equilibrium of a reaction. These are:
1. concentration
2. temperature
3. pressure
Le Chatelier’s Principle helps to predict what a change in temperature, concentration or
pressure will have on the position of the equilibrium in a chemical reaction. This is very important,
particularly in industrial applications, where yields must be accurately predicted and maximised.
Definition: Le Chatelier’s Principle
If a chemical system at equilibrium experiences a change in concentration, temperature or
total pressure the equilibrium will shift in order to minimise that change.
16.8.1
The effect of concentration on equilibrium
If the concentration of a substance is increased, the equilibrium will shift so that this concentration decreases. So for example, if the concentration of a reactant was increased, the equilibrium
would shift in the direction of the reaction that uses up the reactants, so that the reactant concentration decreases and equilibrium is restored. In the reaction between nitrogen and hydrogen
to produce ammonia:
N2 (g) + 3H2 (g) ⇔ 2N H3 (g)
• If the nitrogen or hydrogen concentration was increased, Le Chatelier’s principle predicts
that equilibrium will shift to favour the forward reaction so that the excess nitrogen and
hydrogen are used up to produce ammonia. Equilibrium shifts to the right.
• If the nitrogen or hydrogen concentration was decreased, the reverse reaction would be
favoured so that some of the ammonia would change back to nitrogen and hydrogen to
restore equilibrium.
• The same would be true if the concentration of the product (NH3 ) was changed. If [NH3 ]
decreases, the forward reaction is favoured and if [NH3 ] increases, the reverse reaction is
favoured.
16.8.2
The effect of temperature on equilibrium
If the temperature of a reaction mixture is increased, the equilibrium will shift to decrease the
temperature. So it will favour the reaction which will use up heat energy, in other words the
endothermic reaction. The opposite is true if the temperature is decreased. In this case, the
reaction that produces heat energy will be favoured, in other words, the exothermic reaction.
The reaction shown below is exothermic (shown by the negative value for ∆ H). This means
that the forward reaction, where nitrogen and hydrogen react to form ammonia, gives off heat.
In the reverse reaction, where ammonia is broken down into hydrogen and nitrogen gas, heat is
used up and so this reaction is endothermic.
e.g. N2 (g) + 3H2 (g) ⇔ 2N H3 (g) and ∆H = −92kJ
An increase in temperature favours the reaction that is endothermic (the reverse reaction) because it uses up energy. If the temperature is increased, then the yield of ammonia (NH3 )
310
CHAPTER 16. REACTION RATES - GRADE 12
16.8
decreases.
A decrease in temperature favours the reaction that is exothermic (the forward reaction) because
it produces energy. Therefore, if the temperature is decreased, then the yield of NH3 increases.
Activity :: Experiment : Le Chatelier’s Principle
Aim:
To determine the effect of a change in concentration and temperature on chemical
equilibrium
Apparatus:
0.2 M CoCl2 solution, concentrated HCl, water, test tube, bunsen burner
Method:
1. Put 4-5 drops of 0.2M CoCl2 solution into a test tube.
2. Add 20-25 drops of concentrated HCl.
3. Add 10-12 drops of water.
4. Heat the solution for 1-2 minutes.
5. Cool the solution for 1 minute under a tap.
6. Observe and record the colour changes that take place during the reaction.
The equation for the reaction that takes place is:
−
e.g. CoCl42− + 6H2 O ⇔ Co(H2 O)2+
6 + 4Cl
{z
}
|
{z
}
|
blue
pink
Results:
Complete your observations in the table below, showing the colour changes that
take place, and also indicating whether the concentration of each of the ions in
solution increases or decreases.
Initial
colour
Final
colour
[Co2+ ]
[Cl− ]
[CoCl2−
4 ]
Add Cl−
Add H2 O
Increase
temp.
Decrease
temp.
Conclusions:
Use your knowledge of equilibrium principles to explain the changes that you
recorded in the table above. Draw a conclusion about the effect of a change in
concentration of either the reactants or products on the equilibrium position. Also
draw a conclusion about the effect of a change in temperature on the equilibrium
position.
311
16.8
CHAPTER 16. REACTION RATES - GRADE 12
16.8.3
The effect of pressure on equilibrium
In the case of gases, we refer to pressure instead of concentration. Similar principles apply as
those that were described before for concentration. When the pressure of a system increases,
there are more particles in a particular space. The equilibrium will shift in a direction that reduces
the number of gas particles so that the pressure is also reduced. To predict what will happen in a
reaction, we need to look at the number of moles of gas that are in the reactants and products.
Look at the example below:
e.g. 2SO2 (g) + O2 (g) ⇔ 2SO3 (g)
In this reaction, two moles of product are formed for every three moles of reactants. If we
increase the pressure on the closed system, the equilibrium will shift to the right because the
forward reaction reduces the number of moles of gas that are present. This means that the
yield of SO3 will increase. The opposite will apply if the pressure on the system decreases. the
equilibrium will shift to the left, and the concentration of SO2 and O2 will increase.
Important: The following rules will help in predicting the changes that take place in
equilibrium reactions:
1. If the forward reaction that forms the product is endothermic, then an increase in
temperature will favour this reaction and the yield of product will increase. Lowering
the temperature will decrease the product yield.
2. If the forward reaction that forms the product is exothermic, then a decrease in
temperature will favour this reaction and the product yield will increase. Increasing
the temperature will decrease the product yield.
3. Increasing the pressure favours the side of the equilibrium with the least number of
gas molecules. This is shown in the balanced symbol equation. This rule applies in
reactions with one or more gaseous reactants or products.
4. Decreasing the pressure favours the side of the equilibrium with the most number of
gas molecules. This rule applies in reactions with one or more gaseous reactants or
products.
5. If the concentration of a reactant (on the left) is increased, then some of it must
change to the products (on the right) for equilibrium to be maintained. The equilibrium position will shift to the right.
6. If the concentration of a reactant (on the left) is decreased, then some of the products
(on the right) must change back to reactants for equilibrium to be maintained. The
equilibrium position will shift to the left.
7. A catalyst does not affect the equilibrium position of a reaction. It only influences
the rate of the reaction, in other words, how quickly equilibrium is reached.
Worked Example 81: Reaction Rates 1
Question: 2N O2 (g) ⇔ 2N O(g) + O2 (g) and ∆H > 0 How will the rate of the
reverse reaction be affected by:
1. a decrease in temperature?
2. the addition of a catalyst?
3. the addition of more NO gas?
Answer
312
CHAPTER 16. REACTION RATES - GRADE 12
16.8
1. The rate of the forward reaction will increase since it is the forward reaction that
is exothermix and therefore produces energy to balance the loss of energy from
the decrease in temperature. The rate of the reverse reaction will decrease.
2. The rate of the reverse and the forward reaction will increase.
3. The rate of the reverse reaction will increase so that the extra NO gas is
converted into NO2 gas.
Worked Example 82: Reaction Rates 2
Question:
1. Write a balanced equation for the exothermic reaction between Zn(s) and HCl.
2. Name 3 ways to increase the reaction rate between hydrochloric acid and zinc
metal.
Answer
1. Zn(s) + 2HCl(aq) ⇔ ZnCl2 (aq) + H2 (g)
2. A catalyst could be added, the zinc solid could be ground into a fine powder
to increase its surface area, the HCl concentration could be increased or the
reaction temperature could be increased.
Exercise: Reaction rates and equilibrium
1. The following reaction reaches equilibrium in a closed container:
CaCO3 (s) ⇔ CaO(s) + CO2 (g)
The pressure of the system is increased by decreasing the volume of the container. How will the number of moles and the concentration of the CO2 (g)
have changed when a new equilibrium is reached at the same temperature?
A
B
C
D
moles of CO2
decreased
increased
decreased
decreased
[CO2 ]
decreased
increased
stays the same
increased
(IEB Paper 2, 2003)
2. The following reaction has reached equilibrium in a closed container:
C(s) + H2 O(g) ⇔ CO(g) + H2 (g) ∆H ¿ 0
The pressure of the system is then decreased by increasing the volume of the
container. How will the concentration of the H2 (g) and the value of Kc be
affected when the new equilibrium is established? Assume that the temperature
of the system remains unchanged.
A
B
C
D
[H2 ]
increases
increases
unchanged
decreases
313
Kc
increases
unchanged
unchanged
unchanged
16.8
CHAPTER 16. REACTION RATES - GRADE 12
(IEB Paper 2, 2004)
3. During a classroom experiment copper metal reacts with concentrated nitric
acid to produce NO2 gas, which is collected in a gas syringe. When enough gas
has collected in the syringe, the delivery tube is clamped so that no gas can
escape. The brown NO2 gas collected reaches an equilibrium with colourless
N2 O4 gas as represented by the following equation:
2N O2 (g) ⇔ N2 O4 (g)
Once this equilibrium has been established, there are 0.01 moles of NO2 gas
and 0.03 moles of N2 O4 gas present in the syringe.
(a) A learner, noticing that the colour of the gas mixture in the syringe is no
longer changing, comments that all chemical reactions in the syringe must
have stopped. Is this assumption correct? Explain.
(b) The gas in the syringe is cooled. The volume of the gas is kept constant
during the cooling process. Will the gas be lighter or darker at the lower
temperature? Explain your answer.
(c) The volume of the syringe is now reduced to 75 cm3 by pushing the plunger
in and holding it in the new position. There are 0.032 moles of N2 O4
gas present once the equilibrium has been re-established at the reduced
volume (75 cm3 ). Calculate the value of the equilibrium constant for this
equilibrium.
(IEB Paper 2, 2004)
4. Consider the following reaction, which takes place in a closed container:
A(s) + B(g) → AB(g) ∆H < 0
If you wanted to increase the rate of the reaction, which of the following would
you do?
(a) decrease the concentration of B
(b) decrease the temperature of A
(c) grind A into a fine powder
(d) decrease the pressure
(IEB Paper 2, 2002)
5. Gases X and Y are pumped into a 2 dm3 container. When the container is
sealed, 4 moles of gas X and 4 moles of gas Y are present. The following
equilibrium is established:
2X(g) + 3Y(g) ⇔ X2 Y3
The graph below shows the number of moles of gas X and gas X2 Y3 that are
present from the time the container is sealed.
4
number
of
moles
0,5
30
70
100
time (s)
314
CHAPTER 16. REACTION RATES - GRADE 12
16.9
(a) How many moles of gas X2 Y3 are formed by the time the reaction reaches
equilibrium at 30 seconds?
(b) Calculate the value of the equilibrium constant at t = 50 s.
(c) At 70 s the temperature is increased. Is the forward reaction endothermic
or exothermic? Explain in terms of Le Chatelier’s Principle.
(d) How will this increase in temperature affect the value of the equilibrium
constant?
16.9
Industrial applications
The Haber process is a good example of an industrial process which uses the equilibrium
principles that have been discussed. The equation for the process is as follows:
N2 (g) + 3H2 (g) ⇔ 2N H3 (g) + energy
Since the reaction is exothermic, the forward reaction is favoured at low temperatures, and
the reverse reaction at high temperatures. If the purpose of the Haber process is to produce
ammonia, then the temperature must be maintained at a level that is low enough to ensure that
the reaction continues in the forward direction.
The forward reaction is also favoured by high pressures because there are four moles of reactant
for every two moles of product formed.
The k value for this reaction will be calculated as follows:
k=
[N H3 ]2
[N2 ][H2 ]3
Exercise: Applying equilibrium principles
Look at the values of k calculated for the Haber process reaction at different
temperatures, and then answer the questions that follow:
T oC
25
200
300
400
500
k
6.4
4.4
4.3
1.6
1.5
x
x
x
x
x
102
10−1
10−3
10−4
10−5
1. What happens to the value of k as the temperature increases?
2. Which reaction is being favoured when the temperature is 300 degrees celsius?
3. According to this table, which temperature would be best if you wanted to
produce as much ammonia as possible? Explain.
315
16.10
CHAPTER 16. REACTION RATES - GRADE 12
16.10
Summary
• The rate of a reaction describes how quickly reactants are used up, or how quickly
products form. The units used are moles per second.
• A number of factors can affect the rate of a reaction. These include the nature of the
reactants, the concentration of reactants, temperature of the reaction, the presence or
absence of a catalyst and the surface area of the reactants.
• Collision theory provides one way of explaining why each of these factors can affect the
rate of a reaction. For example, higher temperatures mean increased reaction rates because
the reactant particles have more energy and are more likely to collide successfully with each
other.
• Different methods can be used to measure the rate of a reaction. The method used
will depend on the nature of the product. Reactions that produce gases can be measured
by collecting the gas in a syringe. Reactions that produce a precipitate are also easy to
measure because the precipitate is easily visible.
• For any reaction to occur, a minimum amount of energy is needed so that bonds in the
reactants can break, and new bonds can form in the products. The minimum energy that
is required is called the activation energy of a reaction.
• In reactions where the particles do not have enough energy to overcome this activation
energy, one of two methods can be used to facilitate a reaction to take place: increase the
temperature of the reaction or add a catalyst.
• Increasing the temperature of a reaction means that the average energy of the reactant particles increases and they are more likely to have enough energy to overcome the
activation energy.
• A catalyst is used to lower the activation energy so that the reaction is more likely to
take place. A catalyst does this by providing an alternative, lower energy pathway, for the
reaction.
• A catalyst therefore speeds up a reaction but does not become part of the reaction in
any way.
• Chemical equilibrium is the state of a reaction, where the concentrations of the reactants
and the products have no net change over time. Usually this occurs when the rate of the
forward reaction is the same as the rate of the reverse reaction.
• The equilibrium constant relates to reactions at equilibrium, and can be calculated using
the following equation:
kc =
[C]c [D]d
[A]a [B]b
where A and B are reactants, C and D are products and a, b, c, and d are the coefficients
of the respective reactants and products.
• A high kc value means that the concentration of products at equilibrium is high and the
reaction has a high yield. A low kc value means that the concentration of products at
equilibrium is low and the reaction has a low yield.
• Le Chatelier’s Principle states that if a chemical system at equilibrium experiences a
change in concentration, temperature or total pressure the equilibrium will shift in order
to minimise that change. For example, if the pressure of a gaseous system at eqilibrium
was increased, the equilibrium would shift to favour the reaction that produces the lowest
quantity of the gas. If the temperature of the same system was to increase, the equilibrium
would shift to favour the endothermic reaction. Similar principles apply for changes in
concentration of the reactants or products in a reaction.
• The principles of equilibrium are very important in industrial applications such as the
Haber process, so that productivity can be maximised.
316
CHAPTER 16. REACTION RATES - GRADE 12
Exercise: Summary Exercise
1. For each of the following questions, choose the one correct answer from the
list provided.
(a) Consider the following reaction that has reached equilibrium after some
time in a sealed 1 dm3 flask:
P Cl5 (g) ⇔ P Cl3 (g) + Cl2 (g); ∆H is positive
Which one of the following reaction conditions applied to the system would
decrease the rate of the reverse reaction?
i.
ii.
iii.
iv.
increase the pressure
increase the reaction temperature
continually remove Cl2 (g) from the flask
addition of a suitable catalyst
(IEB Paper 2, 2001)
(b) The following equilibrium constant expression is given for a particular reaction:
Kc = [H2 O]4 [CO2 ]3 /[C3 H8 ][O2 ]5
For which one of the following reactions is the above expression of Kc is
correct?
i.
ii.
iii.
iv.
C3 H8 (g) + 5O2 (g) ⇔ 4H2 O(g) + 3CO2 (g)
4H2 O(g) + 3CO2 (g) ⇔ C3 H8 (g) + 5O2 (g)
2C3 H8 (g) + 7O2 (g) ⇔ 6CO(g) + 8H2 O(g)
C3 H8 (g) + 5O2 (g) ⇔ 4H2 O(l) + 3CO2 (g)
(IEB Paper 2, 2001)
2. 10 g of magnesium ribbon reacts with a 0.15 mol.dm−3 solution of hydrochloric
acid at a temperature of 250 C.
(a) Write a balanced chemical equation for the reaction.
(b) State two ways of increasing the rate of production of H2 (g).
(c) A table of the results is given below:
Time elapsed (min) Vol of H2 (g) (cm3 )
0
0
0.5
17
1.0
25
1.5
30
2.0
33
2.5
35
3.0
35
i. Plot a graph of volume versus time for these results.
ii. Explain the shape of the graph during the following two time intervals:
t = 0 to t = 2.0 min and then t = 2.5 and t = 3.0 min by referring
to the volume of H2 (g) produced.
(IEB Paper 2, 2001)
3. Cobalt chloride crystals are dissolved in a beaker containing ethanol and then
a few drops of water are added. After a period of time, the reaction reaches
equilibrium as follows:
−
CoCl42− (blue) +6H2 O ⇔ Co(H2 O)2+
6 (pink) +4Cl
The solution, which is now just blue, is poured into three test tubes. State,
in each case, what colour changes will be observed (if any) if the following are
added in turn to each test tube:
(a) 1 cm3 of distilled water
(b) A few crystals of sodium chloride
317
16.10
16.10
CHAPTER 16. REACTION RATES - GRADE 12
(c) The addition of dilute hydrochloric acid to the third test tube causes the
solution to turn pink. Explain why this occurs.
(IEB Paper 2, 2001)
318
APPENDIX A. GNU FREE DOCUMENTATION LICENSE
you must enclose the copies in covers that carry, clearly and legibly, all these Cover Texts: FrontCover Texts on the front cover, and Back-Cover Texts on the back cover. Both covers must also
clearly and legibly identify you as the publisher of these copies. The front cover must present the
full title with all words of the title equally prominent and visible. You may add other material on
the covers in addition. Copying with changes limited to the covers, as long as they preserve the
title of the Document and satisfy these conditions, can be treated as verbatim copying in other
respects.
If the required texts for either cover are too voluminous to fit legibly, you should put the first
ones listed (as many as fit reasonably) on the actual cover, and continue the rest onto adjacent
pages.
If you publish or distribute Opaque copies of the Document numbering more than 100, you must
either include a machine-readable Transparent copy along with each Opaque copy, or state in or
with each Opaque copy a computer-network location from which the general network-using public
has access to download using public-standard network protocols a complete Transparent copy of
the Document, free of added material. If you use the latter option, you must take reasonably
prudent steps, when you begin distribution of Opaque copies in quantity, to ensure that this
Transparent copy will remain thus accessible at the stated location until at least one year after
the last time you distribute an Opaque copy (directly or through your agents or retailers) of that
edition to the public.
It is requested, but not required, that you contact the authors of the Document well before
redistributing any large number of copies, to give them a chance to provide you with an updated
version of the Document.
MODIFICATIONS
You may copy and distribute a Modified Version of the Document under the conditions of
sections A and A above, provided that you release the Modified Version under precisely this
License, with the Modified Version filling the role of the Document, thus licensing distribution
and modification of the Modified Version to whoever possesses a copy of it. In addition, you
must do these things in the Modified Version:
1. Use in the Title Page (and on the covers, if any) a title distinct from that of the Document,
and from those of previous versions (which should, if there were any, be listed in the History
section of the Document). You may use the same title as a previous version if the original
publisher of that version gives permission.
2. List on the Title Page, as authors, one or more persons or entities responsible for authorship
of the modifications in the Modified Version, together with at least five of the principal
authors of the Document (all of its principal authors, if it has fewer than five), unless they
release you from this requirement.
3. State on the Title page the name of the publisher of the Modified Version, as the publisher.
4. Preserve all the copyright notices of the Document.
5. Add an appropriate copyright notice for your modifications adjacent to the other copyright
notices.
6. Include, immediately after the copyright notices, a license notice giving the public permission to use the Modified Version under the terms of this License, in the form shown in the
Addendum below.
7. Preserve in that license notice the full lists of Invariant Sections and required Cover Texts
given in the Document’s license notice.
8. Include an unaltered copy of this License.
9. Preserve the section Entitled “History”, Preserve its Title, and add to it an item stating
at least the title, year, new authors, and publisher of the Modified Version as given on the
Title Page. If there is no section Entitled “History” in the Document, create one stating
the title, year, authors, and publisher of the Document as given on its Title Page, then
add an item describing the Modified Version as stated in the previous sentence.
469
APPENDIX A. GNU FREE DOCUMENTATION LICENSE
10. Preserve the network location, if any, given in the Document for public access to a Transparent copy of the Document, and likewise the network locations given in the Document
for previous versions it was based on. These may be placed in the “History” section. You
may omit a network location for a work that was published at least four years before the
Document itself, or if the original publisher of the version it refers to gives permission.
11. For any section Entitled “Acknowledgements” or “Dedications”, Preserve the Title of the
section, and preserve in the section all the substance and tone of each of the contributor
acknowledgements and/or dedications given therein.
12. Preserve all the Invariant Sections of the Document, unaltered in their text and in their
titles. Section numbers or the equivalent are not considered part of the section titles.
13. Delete any section Entitled “Endorsements”. Such a section may not be included in the
Modified Version.
14. Do not re-title any existing section to be Entitled “Endorsements” or to conflict in title
with any Invariant Section.
15. Preserve any Warranty Disclaimers.
If the Modified Version includes new front-matter sections or appendices that qualify as Secondary
Sections and contain no material copied from the Document, you may at your option designate
some or all of these sections as invariant. To do this, add their titles to the list of Invariant
Sections in the Modified Version’s license notice. These titles must be distinct from any other
section titles.
You may add a section Entitled “Endorsements”, provided it contains nothing but endorsements
of your Modified Version by various parties–for example, statements of peer review or that the
text has been approved by an organisation as the authoritative definition of a standard.
You may add a passage of up to five words as a Front-Cover Text, and a passage of up to 25
words as a Back-Cover Text, to the end of the list of Cover Texts in the Modified Version. Only
one passage of Front-Cover Text and one of Back-Cover Text may be added by (or through
arrangements made by) any one entity. If the Document already includes a cover text for the
same cover, previously added by you or by arrangement made by the same entity you are acting
on behalf of, you may not add another; but you may replace the old one, on explicit permission
from the previous publisher that added the old one.
The author(s) and publisher(s) of the Document do not by this License give permission to use
their names for publicity for or to assert or imply endorsement of any Modified Version.
COMBINING DOCUMENTS
You may combine the Document with other documents released under this License, under the
terms defined in section A above for modified versions, provided that you include in the combination all of the Invariant Sections of all of the original documents, unmodified, and list them
all as Invariant Sections of your combined work in its license notice, and that you preserve all
their Warranty Disclaimers.
The combined work need only contain one copy of this License, and multiple identical Invariant
Sections may be replaced with a single copy. If there are multiple Invariant Sections with the
same name but different contents, make the title of each such section unique by adding at the
end of it, in parentheses, the name of the original author or publisher of that section if known,
or else a unique number. Make the same adjustment to the section titles in the list of Invariant
Sections in the license notice of the combined work.
In the combination, you must combine any sections Entitled “History” in the various original
documents, forming one section Entitled “History”; likewise combine any sections Entitled “Acknowledgements”, and any sections Entitled “Dedications”. You must delete all sections Entitled
“Endorsements”.
470
APPENDIX A. GNU FREE DOCUMENTATION LICENSE
COLLECTIONS OF DOCUMENTS
You may make a collection consisting of the Document and other documents released under
this License, and replace the individual copies of this License in the various documents with a
single copy that is included in the collection, provided that you follow the rules of this License
for verbatim copying of each of the documents in all other respects.
You may extract a single document from such a collection, and distribute it individually under
this License, provided you insert a copy of this License into the extracted document, and follow
this License in all other respects regarding verbatim copying of that document.
AGGREGATION WITH INDEPENDENT WORKS
A compilation of the Document or its derivatives with other separate and independent documents
or works, in or on a volume of a storage or distribution medium, is called an “aggregate” if the
copyright resulting from the compilation is not used to limit the legal rights of the compilation’s
users beyond what the individual works permit. When the Document is included an aggregate,
this License does not apply to the other works in the aggregate which are not themselves derivative
works of the Document.
If the Cover Text requirement of section A is applicable to these copies of the Document, then if
the Document is less than one half of the entire aggregate, the Document’s Cover Texts may be
placed on covers that bracket the Document within the aggregate, or the electronic equivalent
of covers if the Document is in electronic form. Otherwise they must appear on printed covers
that bracket the whole aggregate.
TRANSLATION
Translation is considered a kind of modification, so you may distribute translations of the Document under the terms of section A. Replacing Invariant Sections with translations requires
special permission from their copyright holders, but you may include translations of some or
all Invariant Sections in addition to the original versions of these Invariant Sections. You may
include a translation of this License, and all the license notices in the Document, and any Warranty Disclaimers, provided that you also include the original English version of this License and
the original versions of those notices and disclaimers. In case of a disagreement between the
translation and the original version of this License or a notice or disclaimer, the original version
will prevail.
If a section in the Document is Entitled “Acknowledgements”, “Dedications”, or “History”, the
requirement (section A) to Preserve its Title (section A) will typically require changing the actual
title.
TERMINATION
You may not copy, modify, sub-license, or distribute the Document except as expressly provided
for under this License. Any other attempt to copy, modify, sub-license or distribute the Document
is void, and will automatically terminate your rights under this License. However, parties who
have received copies, or rights, from you under this License will not have their licenses terminated
so long as such parties remain in full compliance.
FUTURE REVISIONS OF THIS LICENSE
The Free Software Foundation may publish new, revised versions of the GNU Free Documentation
License from time to time. Such new versions will be similar in spirit to the present version, but
may differ in detail to address new problems or concerns. See http://www.gnu.org/copyleft/.
471
APPENDIX A. GNU FREE DOCUMENTATION LICENSE
Each version of the License is given a distinguishing version number. If the Document specifies
that a particular numbered version of this License “or any later version” applies to it, you have the
option of following the terms and conditions either of that specified version or of any later version
that has been published (not as a draft) by the Free Software Foundation. If the Document does
not specify a version number of this License, you may choose any version ever published (not as
a draft) by the Free Software Foundation.
ADDENDUM: How to use this License for your documents
To use this License in a document you have written, include a copy of the License in the document
and put the following copyright and license notices just after the title page:
c YEAR YOUR NAME. Permission is granted to copy, distribute and/or
Copyright modify this document under the terms of the GNU Free Documentation License,
Version 1.2 or any later version published by the Free Software Foundation; with no
Invariant Sections, no Front-Cover Texts, and no Back-Cover Texts. A copy of the
license is included in the section entitled “GNU Free Documentation License”.
If you have Invariant Sections, Front-Cover Texts and Back-Cover Texts, replace the “with...Texts.”
line with this:
with the Invariant Sections being LIST THEIR TITLES, with the Front-Cover Texts being LIST,
and with the Back-Cover Texts being LIST.
If you have Invariant Sections without Cover Texts, or some other combination of the three,
merge those two alternatives to suit the situation.
If your document contains nontrivial examples of program code, we recommend releasing these
examples in parallel under your choice of free software license, such as the GNU General Public
License, to permit their use in free software.
472
Was this manual useful for you? yes no
Thank you for your participation!

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Related manuals

Download PDF

advertisement