Catherine E. Housecroft And Alan G. Sharpe 2nd Edition.pdf

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This book has established itself as a leading textbook in thesubject by offering a fresh and exciting approach to theteaching of modern inorganic chemistry. It gives a clearintroduction to key principles with strong coverage ofdescriptive chemistry of the elements. Special selectedtopics chapters are included, covering inorganic kineticsand mechanism, catalysis, solid state chemistry andbioinorganic chemistry.

A new full-colour text design and three-dimensionalillustrations bring inorganic chemistry to life. Topic boxeshave been used extensively throughout the book to relatethe chemistry described in the text to everyday life, thechemical industry, environmental issues and legislation, andnatural resources.

Teaching aids throughout the text have been carefullydesigned to help students learn effectively. The manyworked examples take students through each calculation orexercise step by step, and are followed by related self-studyexercises tackling similar problems with answers to helpdevelop their confidence. In addition, end-of-chapterproblems reinforce learning and develop subjectknowledge and skills. Definitions boxes and end-of-chapterchecklists provide excellent revision aids, while furtherreading suggestions, from topical articles to recentliterature papers, will encourage students to explore topicsin more depth.

Catherine E. Housecroft is Professor of Chemistry at theUniversity of Basel, Switzerland. She is the author of anumber of textbooks and has extensive teachingexperience in the UK, Switzerland, South Africa and theUSA. Alan G. Sharpe is a Fellow of Jesus College, Universityof Cambridge, UK and has had many years of experienceteaching inorganic chemistry to undergraduates

SECONDEDITION

INORGANICCHEMISTRY

INO

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AN

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CATHERINE E. HOUSECROFT AND ALAN G. SHARPE

SECONDEDITION

New to this edition• Many more self-study exercises have been

introduced throughout the book with theaim of making stronger connectionsbetween descriptive chemistry andunderlying principles.

• Additional ‘overview problems’ havebeen added to the end-of-chapterproblem sets.

• The descriptive chemistry has beenupdated, with many new results from theliterature being included.

• Chapter 4 – Bonding in polyatomicmolecules, has been rewritten withgreater emphasis on the use of grouptheory for the derivation of ligand grouporbitals and orbital symmetry labels.

• There is more coverage of supercriticalfluids and ‘green’ chemistry.

• The new full-colour text design enhancesthe presentation of the many molecularstructures and 3-D images.

Supporting this edition• Companion website featuring multiple-

choice questions and rotatable 3-Dmolecular structures, available atwww.pearsoned.co.uk/housecroft. For fullinformation including details of lecturermaterial see the Contents list inside thebook.

• A Solutions Manual, written by CatherineE. Housecroft, with detailed solutions toall end-of-chapter problems within thetext is available for purchase separatelyISBN 0131 39926 8.

INORGANICCHEMISTRYCATHERINE E. HOUSECROFT AND ALAN G. SHARPE

C ATHERINE E.HOU

SECROFTAN

D ALAN G.SHARPE

SECONDEDITION

For additional learning resources visit:www.pearsoned.co.uk/housecroft

Cover illustration by Gary Thompson

SECOND

EDITION

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Visit the Inorganic Chemistry, second edition Companion Website atwww.pearsoned.co.uk/housecroft to find valuable student learning materialincluding:

. Multiple choice questions to help test your learning

. Web-based problems for Chapter 3

. Rotatable 3D structures taken from the book

. Interactive Periodic Table

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First edition 2001Second edition 2005

# Pearson Education Limited 2001, 2005

The rights of Catherine E. Housecroft and Alan G. Sharpe to beidentified as the authors of this Work have been asserted by them inaccordance with the Copyright, Designs and Patents Act 1988.

All rights reserved. No part of this publication may be reproduced, storedin a retrieval system, or transmitted in any form or by any means,electronic, mechanical, photocopying, recording, or otherwise, withouteither the prior written permission of the publisher or a licence permittingrestricted copying in the United Kingdom issued by the CopyrightLicensing Agency Ltd, 90 Tottenham Court Road, London W1T 4LP.

All trademarks used herein are the property of their respective owners.The use of any trademark in this text does not vest in the author orpublisher any trademark ownership rights in such trademarks, nor doesthe use of such trademarks imply any affiliation with or endorsement ofthis book by such owners.

ISBN 0130-39913-2

British Library Cataloguing-in-Publication DataA catalogue record for this book is available from the British Library

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Typeset in 912 /12 pt Times by 60Printed by Ashford Colour Press Ltd., Gosport

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Contents

Preface to the second edition xxxiPreface to the first edition xxiii

1 Some basic concepts 1

1.1 Introduction 1

Inorganic chemistry: it is not an isolated branch of chemistry 1The aims of Chapter 1 1

1.2 Fundamental particles of an atom 1

1.3 Atomic number, mass number and isotopes 2

Nuclides, atomic number and mass number 2Relative atomic mass 2Isotopes 2

1.4 Successes in early quantum theory 3

Some important successes of classical quantum theory 4Bohr’s theory of the atomic spectrum of hydrogen 5

1.5 An introduction to wave mechanics 6

The wave-nature of electrons 6The uncertainty principle 6The Schrödinger wave equation 6

1.6 Atomic orbitals 9

The quantum numbers n, l and ml 9The radial part of the wavefunction, RðrÞ 10The radial distribution function, 4�r2RðrÞ2 11The angular part of the wavefunction, Að�; �Þ 12Orbital energies in a hydrogen-like species 13Size of orbitals 13The spin quantum number and the magnetic spin quantum number 15The ground state of the hydrogen atom 16

1.7 Many-electron atoms 16

The helium atom: two electrons 16Ground state electronic configurations: experimental data 16Penetration and shielding 17

1.8 The periodic table 17

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1.9 The aufbau principle 21

Ground state electronic configurations 21Valence and core electrons 22Diagrammatic representations of electronic configurations 22

1.10 Ionization energies and electron affinities 23

Ionization energies 23Electron affinities 25

1.11 Bonding models: an introduction 26

A historical overview 26Lewis structures 26

1.12 Homonuclear diatomic molecules: valence bond (VB) theory 27

Uses of the term homonuclear 27Covalent bond distance, covalent radius and van der Waals radius 27The valence bond (VB) model of bonding in H2 27The valence bond (VB) model applied to F2 , O2 and N2 28

1.13 Homonuclear diatomic molecules: molecular orbital (MO) theory 29

An overview of the MO model 29Molecular orbital theory applied to the bonding in H2 29The bonding in He2, Li2 and Be2 31The bonding in F2 and O2 32What happens if the s–p separation is small? 33

1.14 The octet rule 36

1.15 Electronegativity values 36

Pauling electronegativity values, �P 37Mulliken electronegativity values, �M 37Allred–Rochow electronegativity values, �AR 38Electronegativity: final remarks 38

1.16 Dipole moments 39

Polar diatomic molecules 39Molecular dipole moments 40

1.17 MO theory: heteronuclear diatomic molecules 41

Which orbital interactions should be considered? 41Hydrogen fluoride 42Carbon monoxide 42

1.18 Isoelectronic molecules 43

1.19 Molecular shape and the VSEPR model 43

Valence-shell electron-pair repulsion theory 43Structures derived from a trigonal bipyramid 47Limitations of VSEPR theory 48

1.20 Molecular shape: geometrical isomerism 48

Square planar species 48Octahedral species 48Trigonal bipyramidal species 49High coordination numbers 49Double bonds 49

vi Contents

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2 Nuclear properties 53

2.1 Introduction 53

2.2 Nuclear binding energy 53

Mass defect and binding energy 53The average binding energy per nucleon 54

2.3 Radioactivity 55

Nuclear emissions 55Nuclear transformations 55The kinetics of radioactive decay 56Units of radioactivity 57

2.4 Artificial isotopes 57

Bombardment of nuclei by high-energy a-particles and neutrons 57Bombardment of nuclei by ‘slow’ neutrons 57

2.5 Nuclear fission 58

The fission of uranium-235 58The production of energy by nuclear fission 60Nuclear reprocessing 61

2.6 Syntheses of transuranium elements 61

2.7 The separation of radioactive isotopes 62

Chemical separation 62The Szilard–Chalmers effect 62

2.8 Nuclear fusion 62

2.9 Applications of isotopes 63

Infrared (IR) spectroscopy 63Kinetic isotope effects 64Radiocarbon dating 64Analytical applications 65

2.10 Sources of 2H and 13C 65

Deuterium: electrolytic separation of isotopes 65Carbon-13: chemical enrichment 65

2.11 Multinuclear NMR spectroscopy in inorganic chemistry 67

Which nuclei are suitable for NMR spectroscopic studies? 68Chemical shift ranges 68Spin–spin coupling 69Stereochemically non-rigid species 72Exchange processes in solution 73

2.12 Mössbauer spectroscopy in inorganic chemistry 73

The technique of Mössbauer spectroscopy 73What can isomer shift data tell us? 75

Contents vii

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3 An introduction to molecular symmetry 79

3.1 Introduction 79

3.2 Symmetry operations and symmetry elements 79

Rotation about an n-fold axis of symmetry 80Reflection through a plane of symmetry (mirror plane) 80Reflection through a centre of symmetry (inversion centre) 82Rotation about an axis, followed by reflection through a plane perpendicularto this axis 82Identity operator 82

3.3 Successive operations 84

3.4 Point groups 85

C1 point group 85C1v point group 85D1h point group 85Td, Oh or Ih point groups 86Determining the point group of a molecule or molecular ion 86

3.5 Character tables: an introduction 89

3.6 Why do we need to recognize symmetry elements? 90

3.7 Infrared spectroscopy 90

How many vibrational modes are there for a given molecular species? 90Selection rule for an infrared active mode of vibration 91Linear (D1h or C1v) and bent (C2v) triatomic molecules 92XY3 molecules with D3h or C3v symmetry 92XY4 molecules with Td or D4h symmetry 93Observing IR spectroscopic absorptions: practical problems 94

3.8 Chiral molecules 95

4 Bonding in polyatomic molecules 100

4.1 Introduction 100

4.2 Valence bond theory: hybridization of atomic orbitals 100

What is orbital hybridization? 100sp Hybridization: a scheme for linear species 101sp2 Hybridization: a scheme for trigonal planar species 102sp3 Hybridization: a scheme for tetrahedral and related species 103Other hybridization schemes 104

4.3 Valence bond theory: multiple bonding in polyatomic molecules 105

C2H4 105HCN 105BF3 106

4.4 Molecular orbital theory: the ligand group orbital approach andapplication to triatomic molecules 107

Molecular orbital diagrams: moving from a diatomic to polyatomic species 107

viii Contents

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MO approach to the bonding in linear XH2: symmetry matching by inspection 107MO approach to bonding in linear XH2: working from molecular symmetry 109A bent triatomic: H2O 109

4.5 Molecular orbital theory applied to the polyatomic molecules BH3,NH3 and CH4 112

BH3 112NH3 113CH4 115A comparison of the MO and VB bonding models 116

4.6 Molecular orbital theory: bonding analyses soon become complicated 117

4.7 Molecular orbital theory: learning to use the theory objectively 119

�-Bonding in CO2 119[NO3]

� 120SF6 120Three-centre two-electron interactions 123A more advanced problem: B2H6 124

5 Structures and energetics of metallic and ionic solids 131


5.2 Packing of spheres 131

Cubic and hexagonal close-packing 131The unit cell: hexagonal and cubic close-packing 132Interstitial holes: hexagonal and cubic close-packing 133Non-close-packing: simple cubic and body-centred cubic arrays 134

5.3 The packing-of-spheres model applied to the structures of elements 134

Group 18 elements in the solid state 134H2 and F2 in the solid state 134Metallic elements in the solid state 134

5.4 Polymorphism in metals 136

Polymorphism: phase changes in the solid state 136Phase diagrams 136

5.5 Metallic radii 136

5.6 Melting points and standard enthalpies of atomization of metals 137

5.7 Alloys and intermetallic compounds 139

Substitutional alloys 139Interstitial alloys 139Intermetallic compounds 140

5.8 Bonding in metals and semiconductors 141

Electrical conductivity and resistivity 141Band theory of metals and insulators 141The Fermi level 142Band theory of semiconductors 143

Contents ix

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5.9 Semiconductors 143

Intrinsic semiconductors 143Extrinsic (n- and p-type) semiconductors 143

5.10 Sizes of ions 144

Ionic radii 144Periodic trends in ionic radii 145

5.11 Ionic lattices 146

The rock salt (NaCl) lattice 148The caesium chloride (CsCl) lattice 149The fluorite (CaF2) lattice 149The antifluorite lattice 149The zinc blende (ZnS) lattice: a diamond-type network 149The b-cristobalite (SiO2) lattice 150The wurtzite (ZnS) structure 151The rutile (TiO2) structure 151The CdI2 and CdCl2 lattices: layer structures 151The perovskite (CaTiO3) lattice: a double oxide 152

5.12 Crystal structures of semiconductors 152

5.13 Lattice energy: estimates from an electrostatic model 152

Coulombic attraction within an isolated ion-pair 152Coulombic interactions in an ionic lattice 153Born forces 153The Born–Landé equation 154Madelung constants 154Refinements to the Born–Landé equation 155Overview 155

5.14 Lattice energy: the Born–Haber cycle 155

5.15 Lattice energy: ‘calculated’ versus ‘experimental’ values 156

5.16 Applications of lattice energies 157

Estimation of electron affinities 157Fluoride affinities 157Estimation of standard enthalpies of formation and disproportionation 157The Kapustinskii equation 158

5.17 Defects in solid state lattices: an introduction 158

Schottky defect 158Frenkel defect 158Experimental observation of Schottky and Frenkel defects 159

6 Acids, bases and ions in aqueous solution 162


6.2 Properties of water 162

Structure and hydrogen bonding 162The self-ionization of water 163Water as a Brønsted acid or base 163

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6.3 Definitions and units in aqueous solution 165

Molarity and molality 165Standard state 165Activity 165

6.4 Some Brønsted acids and bases 166

Carboxylic acids: examples of mono-, di- and polybasic acids 166Inorganic acids 167Inorganic bases: hydroxides 167Inorganic bases: nitrogen bases 168

6.5 The energetics of acid dissociation in aqueous solution 169

Hydrogen halides 169H2S, H2Se and H2Te 170

6.6 Trends within a series of oxoacids EOn(OH)m 170

6.7 Aquated cations: formation and acidic properties 171

Water as a Lewis base 171Aquated cations as Brønsted acids 172

6.8 Amphoteric oxides and hydroxides 173

Amphoteric behaviour 173Periodic trends in amphoteric properties 173

6.9 Solubilities of ionic salts 174

Solubility and saturated solutions 174Sparingly soluble salts and solubility products 174The energetics of the dissolution of an ionic salt: �solG

o 175The energetics of the dissolution of an ionic salt: hydration of ions 176Solubilities: some concluding remarks 177

6.10 Common-ion effect 178

6.11 Coordination complexes: an introduction 178

Definitions and terminology 178Investigating coordination complex formation 179

6.12 Stability constants of coordination complexes 180

Determination of stability constants 182Trends in stepwise stability constants 182Thermodynamic considerations of complex formation: an introduction 182

6.13 Factors affecting the stabilities of complexes containing onlymonodentate ligands 186

Ionic size and charge 186Hard and soft metal centres and ligands 187

7 Reduction and oxidation 192


Oxidation and reduction 192Oxidation states 192Stock nomenclature 193

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7.2 Standard reduction potentials, Eo, and relationships between Eo,�Go and K 193

Half-cells and galvanic cells 193Defining and using standard reduction potentials, Eo 195Dependence of reduction potentials on cell conditions 197

7.3 The effect of complex formation or precipitation on Mzþ/M reductionpotentials 199

Half-cells involving silver halides 199Modifying the relative stabilities of different oxidation states of a metal 200

7.4 Disproportionation reactions 203

Disproportionation 203Stabilizing species against disproportionation 203

7.5 Potential diagrams 203

7.6 Frost–Ebsworth diagrams 205

Frost–Ebsworth diagrams and their relationship to potential diagrams 205Interpretation of Frost–Ebsworth diagrams 206

7.7 The relationships between standard reduction potentials and someother quantities 208

Factors influencing the magnitudes of standard reduction potentials 208Values of �fG

o for aqueous ions 209

7.8 Applications of redox reactions to the extraction of elements from theirores 210

Ellingham diagrams 210

8 Non-aqueous media 214


8.2 Relative permittivity 214

8.3 Energetics of ionic salt transfer from water to an organic solvent 215

8.4 Acid–base behaviour in non-aqueous solvents 216

Strengths of acids and bases 216Levelling and differentiating effects 217‘Acids’ in acidic solvents 217Acids and bases: a solvent-oriented definition 217

8.5 Self-ionizing and non-ionizing non-aqueous solvents 217

8.6 Liquid ammonia 218

Physical properties 218Self-ionization 218Reactions in liquid NH3 218Solutions of s-block metals in liquid NH3 219Redox reactions in liquid NH3 221

xii Contents

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8.7 Liquid hydrogen fluoride 221

Physical properties 221Acid–base behaviour in liquid HF 221Electrolysis in liquid HF 222

8.8 Sulfuric acid 222

Physical properties 222Acid–base behaviour in liquid H2SO4 223

8.9 Fluorosulfonic acid 223

Physical properties 223Superacids 224

8.10 Bromine trifluoride 224

Physical properties 224Behaviour of fluoride salts and molecular fluorides in BrF3 225Reactions in BrF3 225

8.11 Dinitrogen tetraoxide 225

Physical properties 225Reactions in N2O4 226

8.12 Ionic liquids 227

Molten salt solvent systems 227Ionic liquids at ambient temperatures 227Reactions in and applications of molten salt/ionic liquid media 229

8.13 Supercritical fluids 230

Properties of supercritical fluids and their uses as solvents 230Supercritical fluids as media for inorganic chemistry 232

9 Hydrogen 236

9.1 Hydrogen: the simplest atom 236

9.2 The Hþ and H� ions 236

The hydrogen ion (proton) 236The hydride ion 237

9.3 Isotopes of hydrogen 237

Protium and deuterium 237Deuterated compounds 237Tritium 238

9.4 Dihydrogen 238

Occurrence 238Physical properties 238Synthesis and uses 238Reactivity 242

9.5 Polar and non-polar E�H bonds 244

9.6 Hydrogen bonding 244

The hydrogen bond 244Trends in boiling points, melting points and enthalpies of vaporization forp-block binary hydrides 246

Contents xiii

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Infrared spectroscopy 246Solid state structures 247Hydrogen bonding in biological systems 250

9.7 Binary hydrides: classification and general properties 251

Classification 251Interstitial metal hydrides 251Saline hydrides 251Molecular hydrides and complexes derived from them 253Polymeric hydrides 254Intermediate hydrides 255

10 Group 1: the alkali metals 257


10.2 Occurrence, extraction and uses 257

Occurrence 257Extraction 257Major uses of the alkali metals and their compounds 259

10.3 Physical properties 259

General properties 259Atomic spectra and flame tests 260Radioactive isotopes 261NMR active nuclei 261

10.4 The metals 261

Appearance 261Reactivity 261

10.5 Halides 263

10.6 Oxides and hydroxides 264

Oxides, peroxides, superoxides, suboxides and ozonides 264Hydroxides 265

10.7 Salts of oxoacids: carbonates and hydrogencarbonates 265

10.8 Aqueous solution chemistry including macrocyclic complexes 267

Hydrated ions 267Complex ions 268

10.9 Non-aqueous coordination chemistry 271

11 The group 2 metals 275



Occurrence 275Extraction 276Major uses of the group 2 metals and their compounds 277

xiv Contents

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General properties 278Flame tests 279Radioactive isotopes 279

11.4 The metals 279

Appearance 279Reactivity 279

11.5 Halides 280

Beryllium halides 280Halides of Mg, Ca, Sr and Ba 282

11.6 Oxides and hydroxides 283

Oxides and peroxides 283Hydroxides 285

11.7 Salts of oxoacids 286

11.8 Complex ions in aqueous solution 287

Aqua species of beryllium 287Aqua species of Mg2þ, Ca2þ, Sr2þ and Ba2þ 288Complexes with ligands other than water 288

11.9 Complexes with amido or alkoxy ligands 288

11.10 Diagonal relationships between Li and Mg, and between Be and Al 288

Lithium and magnesium 289Beryllium and aluminium 290

12 The group 13 elements 293



Occurrence 293Extraction 293Major uses of the group 13 elements and their compounds 295


Electronic configurations and oxidation states 296NMR active nuclei 299

12.4 The elements 299

Appearance 299Structures of the elements 300Reactivity 301

12.5 Simple hydrides 301

Neutral hydrides 301The ½MH4�� ions 305

12.6 Halides and complex halides 307

Boron halides: BX3 and B2X4 307Al(III), Ga(III), In(III) and Tl(III) halides and their complexes 309Lower oxidation state Al, Ga, In and Tl halides 311

Contents xv

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12.7 Oxides, oxoacids, oxoanions and hydroxides 313

Boron oxides, oxoacids and oxoanions 313Aluminium oxides, oxoacids, oxoanions and hydroxides 316Oxides of Ga, In and Tl 317

12.8 Compounds containing nitrogen 317

Nitrides 317Ternary boron nitrides 318Molecular species containing B–N or B–P bonds 319Molecular species containing group 13 metal–nitrogen bonds 321

12.9 Aluminium to thallium: salts of oxoacids, aqueous solution chemistryand complexes 322

Aluminium sulfate and alums 322Aqua ions 322Redox reactions in aqueous solution 322Coordination complexes of the M3þ ions 323

12.10 Metal borides 324

12.11 Electron-deficient borane and carbaborane clusters: an introduction 326

Boron hydrides 326




Occurrence 338Extraction and manufacture 339Uses 339


Ionization energies and cation formation 342Some energetic and bonding considerations 343NMR active nuclei 344Mössbauer spectroscopy 344

13.4 Allotropes of carbon 345

Graphite and diamond: structure and properties 345Graphite: intercalation compounds 345Fullerenes: synthesis and structure 348Fullerenes: reactivity 349Carbon nanotubes 353

13.5 Structural and chemical properties of silicon, germanium, tin and lead 353

Structures 353Chemical properties 353

13.6 Hydrides 354

Binary hydrides 354Halohydrides of silicon and germanium 356

13.7 Carbides, silicides, germides, stannides and plumbides 357

Carbides 357

xvi Contents

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Silicides 358Germides, stannides and plumbides 358

13.8 Halides and complex halides 361

Carbon halides 361Silicon halides 363Halides of germanium, tin and lead 364

13.9 Oxides, oxoacids and hydroxides 365

Oxides and oxoacids of carbon 365Silica, silicates and aluminosilicates 369Oxides, hydroxides and oxoacids of germanium, tin and lead 373

13.10 Silicones 376

13.11 Sulfides 377

13.12 Cyanogen, silicon nitride and tin nitride 379

Cyanogen and its derivatives 379Silicon nitride 380Tin(IV) nitride 381

13.13 Aqueous solution chemistry and salts of oxoacids of germanium,tin and lead 381




Occurrence 386Extraction 387Uses 387


Bonding considerations 390NMR active nuclei 391Radioactive isotopes 391


Nitrogen 392Phosphorus 392Arsenic, antimony and bismuth 393

14.5 Hydrides 394

Trihydrides, EH3 (E¼N, P, As, Sb and Bi) 394Hydrides E2H4 (E¼N, P, As) 397Chloramine and hydroxylamine 398Hydrogen azide and azide salts 399

14.6 Nitrides, phosphides, arsenides, antimonides and bismuthides 401

Nitrides 401Phosphides 402Arsenides, antimonides and bismuthides 402

Contents xvii

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14.7 Halides, oxohalides and complex halides 403

Nitrogen halides 403Oxofluorides and oxochlorides of nitrogen 405Phosphorus halides 406Phosphoryl trichloride, POCl3 408Arsenic and antimony halides 409Bismuth halides 411

14.8 Oxides of nitrogen 412

Dinitrogen monoxide, N2O 412Nitrogen monoxide, NO 412Dinitrogen trioxide, N2O3 413Dinitrogen tetraoxide, N2O4, and nitrogen dioxide, NO2 414Dinitrogen pentaoxide, N2O5 415

14.9 Oxoacids of nitrogen 415

Hyponitrous acid, H2N2O2 415Nitrous acid, HNO2 415Nitric acid, HNO3, and its derivatives 416

14.10 Oxides of phosphorus, arsenic, antimony and bismuth 417

Oxides of phosphorus 418Oxides of arsenic, antimony and bismuth 419

14.11 Oxoacids of phosphorus 419

Phosphinic acid, H3PO2 419Phosphonic acid, H3PO3 420Hypophosphoric acid, H4P2O6 420Phosphoric acid, H3PO4, and its derivatives 421

14.12 Oxoacids of arsenic, antimony and bismuth 422

14.13 Phosphazenes 424

14.14 Sulfides and selenides 426

Sulfides and selenides of phosphorus 426Arsenic, antimony and bismuth sulfides 428

14.15 Aqueous solution chemistry 428





15.3 Physical properties and bonding considerations 434

NMR active nuclei and isotopes as tracers 437


Dioxygen 437Ozone 438

xviii Contents

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Sulfur: allotropes 439Sulfur: reactivity 440Selenium and tellurium 441

15.5 Hydrides 442

Water, H2O 442Hydrogen peroxide, H2O2 442Hydrides H2E (E¼S, Se, Te) 445Polysulfanes 445

15.6 Metal sulfides, polysulfides, polyselenides and polytellurides 446

Sulfides 446Polysulfides 446Polyselenides and polytellurides 447

15.7 Halides, oxohalides and complex halides 448

Oxygen fluorides 448Sulfur fluorides and oxofluorides 448Sulfur chlorides and oxochlorides 450Halides of selenium and tellurium 451

15.8 Oxides 453

Oxides of sulfur 453Oxides of selenium and tellurium 456

15.9 Oxoacids and their salts 457

Dithionous acid, H2S2O4 457Sulfurous and disulfurous acids, H2SO3 and H2S2O5 457Dithionic acid, H2S2O6 458Sulfuric acid, H2SO4 459Fluoro- and chlorosulfonic acids, HSO3F and HSO3Cl 461Polyoxoacids with S�O�S units 461Peroxosulfuric acids, H2S2O8 and H2SO5 461Thiosulfuric acid, H2S2O3, and polythionates 461Oxoacids of selenium and tellurium 462

15.10 Compounds of sulfur and selenium with nitrogen 462

Sulfur–nitrogen compounds 462Tetraselenium tetranitride 464

15.11 Aqueous solution chemistry of sulfur, selenium and tellurium 464



Fluorine, chlorine, bromine and iodine 468Astatine 469



16.3 Physical properties and bonding considerations 471

NMR active nuclei and isotopes as tracers 473

Contents xix

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Difluorine 474Dichlorine, dibromine and diiodine 475Charge transfer complexes 475Clathrates 477

16.5 Hydrogen halides 477

16.6 Metal halides: structures and energetics 478

16.7 Interhalogen compounds and polyhalogen ions 479

Interhalogen compounds 479Bonding in ½XY2�� ions 482Polyhalogen cations 482Polyhalide anions 483

16.8 Oxides and oxofluorides of chlorine, bromine and iodine 483

Oxides 483Oxofluorides 484

16.9 Oxoacids and their salts 485

Hypofluorous acid, HOF 485Oxoacids of chlorine, bromine and iodine 485

16.10 Aqueous solution chemistry 488






NMR active nuclei 495

17.4 Compounds of xenon 496

Fluorides 496Chlorides 498Oxides 499Oxofluorides 499Other compounds of xenon 499

17.5 Compounds of krypton and radon 501

18 Organometallic compounds of s- and p-block elements 503


18.2 Group 1: alkali metal organometallics 504

xx Contents

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18.3 Group 2 organometallics 507

Beryllium 507Magnesium 509Calcium, strontium and barium 510

18.4 Group 13 511

Boron 511Aluminium 511Gallium, indium and thallium 514

18.5 Group 14 518

Silicon 518Germanium 520Tin 521Lead 524Coparallel and tilted C5-rings in group 14 metallocenes 526

18.6 Group 15 527

Bonding aspects and E¼E bond formation 527Arsenic, antimony and bismuth 527

18.7 Group 16 530

Selenium and tellurium 530

19 d-Block chemistry: general considerations 535

19.1 Topic overview 535

19.2 Ground state electronic configurations 535

d-Block metals versus transition elements 535Electronic configurations 536


19.4 The reactivity of the metals 538

19.5 Characteristic properties: a general perspective 538

Colour 538Paramagnetism 539Complex formation 539Variable oxidation states 539

19.6 Electroneutrality principle 539

19.7 Coordination numbers 541

The Kepert model 541Coordination number 2 543Coordination number 3 543Coordination number 4 543Coordination number 5 544Coordination number 6 544Coordination number 7 545Coordination number 8 546Coordination number 9 547Coordination numbers of 10 and above 547

Contents xxi

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19.8 Isomerism in d-block metal complexes 547

Structural isomerism: ionization isomers 548Structural isomerism: hydration isomers 548Structural isomerism: coordination isomerism 549Structural isomerism: linkage isomerism 549Structural isomerism: polymerization isomerism 549Stereoisomerism: geometrical isomers 549Stereoisomerism: optical isomers 549

20 d-Block chemistry: coordination complexes 555


High- and low-spin states 555

20.2 Bonding in d-block metal complexes: valence bond theory 555

Hybridization schemes 555Applying VB theory 556

20.3 Crystal field theory 557

The octahedral crystal field 558Crystal field stabilization energy: high- and low-spin octahedral complexes 560Jahn–Teller distortions 561The tetrahedral crystal field 562The square planar crystal field 562Other crystal fields 564Crystal field theory: uses and limitations 564

20.4 Molecular orbital theory: octahedral complexes 564

Complexes with no metal–ligand �-bonding 564Complexes with metal–ligand �-bonding 566

20.5 Ligand field theory 570

20.6 Electronic spectra 570

Spectral features 570Selection rules 571Electronic spectra of octahedral and tetrahedral complexes 574Microstates 576Tanabe–Sugano diagrams 577

20.7 Evidence for metal–ligand covalent bonding 578

The nephelauxetic effect 578ESR spectroscopy 579

20.8 Magnetic properties 579

Magnetic susceptibility and the spin-only formula 579Spin and orbital contributions to the magnetic moment 581The effects of temperature on �eff 583Spin crossover 584Ferromagnetism, antiferromagnetism and ferrimagnetism 584

20.9 Thermodynamic aspects: ligand field stabilization energies (LFSE) 585

Trends in LFSE 585Lattice energies and hydration energies of Mnþ ions 586Octahedral versus tetrahedral coordination: spinels 587

xxii Contents

Black plate (23,1)

20.10 Thermodynamic aspects: the Irving–Williams series 587

20.11 Thermodynamic aspects: oxidation states in aqueous solution 588

21 d-Block metal chemistry: the first row metals 593



21.3 Physical properties: an overview 597

21.4 Group 3: scandium 597

The metal 597Scandium(III) 598

21.5 Group 4: titanium 598

The metal 598Titanium(IV) 598Titanium(III) 601Low oxidation states 601

21.6 Group 5: vanadium 602

The metal 602Vanadium(V) 602Vanadium(IV) 604Vanadium(III) 605Vanadium(II) 605

21.7 Group 6: chromium 606

The metal 606Chromium(VI) 606Chromium(V) and chromium(IV) 607Chromium(III) 608Chromium(II) 609Chromium–chromium multiple bonds 610

21.8 Group 7: manganese 611

The metal 611Manganese(VII) 612Manganese(VI) 613Manganese(V) 613Manganese(IV) 613Manganese(III) 614Manganese(II) 616

21.9 Group 8: iron 617

The metal 617Iron(VI), iron(V) and iron(IV) 617Iron(III) 618Iron(II) 622

21.10 Group 9: cobalt 624

The metal 624Cobalt(IV) 624

Contents xxiii

Black plate (24,1)

Cobalt(III) 624Cobalt(II) 627

21.11 Group 10: nickel 630

The metal 630Nickel(IV) and nickel(III) 630Nickel(II) 631Nickel(I) 634

21.12 Group 11: copper 634

The metal 634Copper(IV) and (III) 634Copper(II) 635Copper(I) 637

21.13 Group 12: zinc 639

The metal 639Zinc(II) 640

22 d-Block metal chemistry: the second and third row metals 645




Effects of the lanthanoid contraction 649Coordination numbers 649NMR active nuclei 649

22.4 Group 3: yttrium 651

The metal 651Yttrium(III) 651

22.5 Group 4: zirconium and hafnium 652

The metals 652Zirconium(IV) and hafnium(IV) 652Lower oxidation states of zirconium and hafnium 652Zirconium clusters 653

22.6 Group 5: niobium and tantalum 654

The metals 654Niobium(V) and tantalum(V) 654Niobium(IV) and tantalum(IV) 656Lower oxidation state halides 656

22.7 Group 6: molybdenum and tungsten 658

The metals 658Molybdenum(VI) and tungsten(VI) 659Molybdenum(V) and tungsten(V) 662Molybdenum(IV) and tungsten(IV) 663Molybdenum(III) and tungsten(III) 663Molybdenum(II) and tungsten(II) 665

22.8 Group 7: technetium and rhenium 666

The metals 666

xxiv Contents

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High oxidation states of technetium and rhenium: M(VII), M(VI) and M(V) 667Technetium(IV) and rhenium(IV) 669Technetium(III) and rhenium(III) 669

22.9 Group 8: ruthenium and osmium 671

The metals 671High oxidation states of ruthenium and osmium: M(VIII), M(VII) and M(VI) 671Ruthenium(V), (IV) and osmium(V), (IV) 673Ruthenium(III) and osmium(III) 675Ruthenium(II) and osmium(II) 676Mixed-valence ruthenium complexes 678

22.10 Group 9: rhodium and iridium 679

The metals 679High oxidation states of rhodium and iridium: M(VI) and M(V) 679Rhodium(IV) and iridium (IV) 680Rhodium(III) and iridium(III) 680Rhodium(II) and iridium(II) 682Rhodium(I) and iridium(I) 683

22.11 Group 10: palladium and platinum 684

The metals 684The highest oxidation states: M(VI) and M(V) 684Palladium(IV) and platinum(IV) 684Palladium(III), platinum(III) and mixed-valence complexes 685Palladium(II) and platinum(II) 686

22.12 Group 11: silver and gold 689

The metals 689Gold(V) and silver(V) 690Gold(III) and silver(III) 690Gold(II) and silver(II) 691Gold(I) and silver(I) 692Gold(�I) and silver(�I) 694

22.13 Group 12: cadmium and mercury 694

The metals 694Cadmium(II) 695Mercury(II) 695Mercury(I) 696

23 Organometallic compounds of d-block elements 700


Hapticity of a ligand 700

23.2 Common types of ligand: bonding and spectroscopy 700

�-Bonded alkyl, aryl and related ligands 700Carbonyl ligands 701Hydride ligands 702Phosphine and related ligands 703�-Bonded organic ligands 704Dinitrogen 706Dihydrogen 707

23.3 The 18-electron rule 707

Contents xxv

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23.4 Metal carbonyls: synthesis, physical properties and structure 709

Synthesis and physical properties 710Structures 711

23.5 The isolobal principle and application of Wade’s rules 714

23.6 Total valence electron counts in d-block organometallic clusters 716

Single cage structures 717Condensed cages 718Limitations of total valence counting schemes 719

23.7 Types of organometallic reactions 719

Substitution of CO ligands 719Oxidative addition 719Alkyl and hydrogen migrations 720b-Hydrogen elimination 721a-Hydrogen abstraction 721Summary 722

23.8 Metal carbonyls: selected reactions 722

23.9 Metal carbonyl hydrides and halides 723

23.10 Alkyl, aryl, alkene and alkyne complexes 724

�-Bonded alkyl and aryl ligands 724Alkene ligands 725Alkyne ligands 726

23.11 Allyl and buta-1,3-diene complexes 727

Allyl and related ligands 727Buta-1,3-diene and related ligands 728

23.12 Carbene and carbyne complexes 729

23.13 Complexes containing Z5-cyclopentadienyl ligands 730Ferrocene and other metallocenes 731ðZ5-CpÞ2Fe2ðCOÞ4 and derivatives 732

23.14 Complexes containing Z6- and Z7-ligands 734Z6-Arene ligands 734Cycloheptatriene and derived ligands 735

23.15 Complexes containing the Z4-cyclobutadiene ligand 737

24 The f -block metals: lanthanoids and actinoids 741


24.2 f -Orbitals and oxidation states 742

24.3 Atom and ion sizes 743

The lanthanoid contraction 743Coordination numbers 743

xxvi Contents

Black plate (27,1)

24.4 Spectroscopic and magnetic properties 744

Electronic spectra and magnetic moments: lanthanoids 744Luminescence of lanthanoid complexes 746Electronic spectra and magnetic moments: actinoids 746

24.5 Sources of the lanthanoids and actinoids 747

Occurrence and separation of the lanthanoids 747The actinoids 748

24.6 Lanthanoid metals 748

24.7 Inorganic compounds and coordination complexes of the lanthanoids 749

Halides 749Hydroxides and oxides 750Complexes of Ln(III) 750

24.8 Organometallic complexes of the lanthanoids 751

�-Bonded complexes 751Cyclopentadienyl complexes 753Bis(arene) derivatives 755Complexes containing the Z8-cyclooctatetraenyl ligand 755

24.9 The actinoid metals 755

24.10 Inorganic compounds and coordination complexes of thorium,uranium and plutonium 756

Thorium 756Uranium 757Plutonium 758

24.11 Organometallic complexes of thorium and uranium 759

�-Bonded complexes 759Cyclopentadienyl derivatives 760Complexes containing the Z8-cyclooctatetraenyl ligand 761

25 d-Block metal complexes: reaction mechanisms 764


25.2 Ligand substitutions: some general points 764

Kinetically inert and labile complexes 764Stoichiometric equations say nothing about mechanism 764Types of substitution mechanism 765Activation parameters 765

25.3 Substitution in square planar complexes 766

Rate equations, mechanism and the trans-effect 766Ligand nucleophilicity 769

25.4 Substitution and racemization in octahedral complexes 769

Water exchange 770The Eigen–Wilkins mechanism 772Stereochemistry of substitution 774Base-catalysed hydrolysis 774Isomerization and racemization of octahedral complexes 776

Contents xxvii

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25.5 Electron-transfer processes 777

Inner-sphere mechanism 777Outer-sphere mechanism 779

26 Homogeneous and heterogeneous catalysis 786

26.1 Introduction and definitions 786

26.2 Catalysis: introductory concepts 786

Energy profiles for a reaction: catalysed versus non-catalysed 786Catalytic cycles 787Choosing a catalyst 788

26.3 Homogeneous catalysis: alkene (olefin) metathesis 789

26.4 Homogeneous catalysis: industrial applications 791

Alkene hydrogenation 791Monsanto acetic acid synthesis 793Tennessee–Eastman acetic anhydride process 794Hydroformylation (Oxo-process) 795Alkene oligomerization 797

26.5 Homogeneous catalyst development 797

Polymer-supported catalysts 797Biphasic catalysis 798d-Block organometallic clusters as homogeneous catalysts 799

26.6 Heterogeneous catalysis: surfaces and interactions with adsorbates 799

26.7 Heterogeneous catalysis: commercial applications 802

Alkene polymerization: Ziegler–Natta catalysis 802Fischer–Tropsch carbon chain growth 803Haber process 804Production of SO3 in the Contact process 805Catalytic converters 805Zeolites as catalysts for organic transformations: uses of ZSM-5 806

26.8 Heterogeneous catalysis: organometallic cluster models 807

27 Some aspects of solid state chemistry 813


27.2 Defects in solid state lattices 813

Types of defect: stoichiometric and non-stoichiometric compounds 813Colour centres (F-centres) 814Thermodynamic effects of crystal defects 814

27.3 Electrical conductivity in ionic solids 815

Sodium and lithium ion conductors 815d-Block metal(II) oxides 816

xxviii Contents

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27.4 Superconductivity 817

Superconductors: early examples and basic theory 817High-temperature superconductors 817Superconducting properties of MgB2 819Applications of superconductors 819

27.5 Ceramic materials: colour pigments 819

White pigments (opacifiers) 820Adding colour 820

27.6 Chemical vapour deposition (CVD) 820

High-purity silicon for semiconductors 821a-Boron nitride 821Silicon nitride and carbide 821III–V Semiconductors 822Metal deposition 823Ceramic coatings 824Perovskites and cuprate superconductors 824

27.7 Inorganic fibres 826

Boron fibres 826Carbon fibres 826Silicon carbide fibres 827Alumina fibres 827

28 The trace metals of life 830


Amino acids, peptides and proteins: some terminology 830

28.2 Metal storage and transport: Fe, Cu, Zn and V 832

Iron storage and transport 832Metallothioneins: transporting some toxic metals 835

28.3 Dealing with O2 837

Haemoglobin and myoglobin 837Haemocyanin 839Haemerythrin 841Cytochromes P-450 843

28.4 Biological redox processes 843

Blue copper proteins 844The mitochondrial electron-transfer chain 845Iron–sulfur proteins 847Cytochromes 851

28.5 The Zn2+ ion: Nature’s Lewis acid 854

Carbonic anhydrase II 854Carboxypeptidase A 855Carboxypeptidase G2 858Cobalt-for-zinc ion substitution 859

Contents xxix

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Appendices 863

1 Greek letters with pronunciations 864

2 Abbreviations and symbols for quantities and units 865

3 Selected character tables 869

4 The electromagnetic spectrum 873

5 Naturally occurring isotopes and their abundances 875

6 Van der Waals, metallic, covalent and ionic radii for the s-, p- andfirst row d-block elements 877

7 Pauling electronegativity values (�P) for selected elements of theperiodic table 879

8 Ground state electronic configurations of the elements andionization energies for the first five ionizations 880

9 Electron affinities 883

10 Standard enthalpies of atomization (�aHo) of the elements at 298K 884

11 Selected standard reduction potentials (298K) 885

Answers to non-descriptive problems 888

Index 905

xxx Contents

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Preface to the second edition

The second edition of Inorganic Chemistry is a natural progression from the first editionpublished in 2001. In this last text, we stated that our aim was to provide a singlevolume that gives a critical introduction to modern inorganic chemistry. Our approachto inorganic chemistry continues as before: we provide a foundation of physical inorganicprinciples and theory followed by descriptive chemistry of the elements, and a number of‘special topics’ that can, if desired, be used for modular teaching. Boxed material has beenused extensively to relate the chemistry described in the text to everyday life, the chemicalindustry, environmental issues and legislation, and natural resources.In going from the first to second editions, the most obvious change has been a move

from two to full colour. This has given us the opportunity to enhance the presentationsof many of the molecular structures and 3D images. In terms of content, the descriptivechemistry has been updated, with many new results from the literature being included.Some exciting advances have taken place in the past two to three years spanning smallmolecule chemistry (for example, the chemistry of [N5]

þ), solid state chemistry (e.g.the first examples of spinel nitrides) and bioinorganic systems (a landmark discovery isthat of a central, 6-coordinate atom, probably nitrogen, at the centre of the FeMo-cofactor in nitrogenase). Other changes to the book have their origins in feedbackfrom people using the text. Chapters 3 and 4 have been modified; in particular, therole of group theory in determining ligand group orbitals and orbital symmetry labelshas been more thoroughly explored. However, we do not feel that a book, the primepurpose of which is to bring chemistry to a student audience, should evolve into atheoretical text. For this reason, we have refrained from an in-depth treatment ofgroup theory. Throughout the book, we have used the popular ‘worked examples’ and‘self-study exercises’ as a means of helping students to grasp principles and concepts.Many more self-study exercises have been introduced throughout the book, with theaim of making stronger connections between descriptive chemistry and underlyingprinciples. Additional ‘overview problems’ have been added to the end-of-chapterproblem sets; in Chapter 3, a set of new problems has been designed to work inconjunction with rotatable structures on the accompanying website (www.pearsoned.co.uk/housecroft).Supplementary data accompanying this text include a Solutions Manual written by

Catherine E. Housecroft. The accompanying website includes features for both studentsand lecturers and can be accessed from www.pearsoned.co.uk/housecroft.The 3D-molecular structures the book have been drawn using atomic coordinates

accessed from the Cambridge Crystallographic Data Base and implemented through theETH in Zürich, or from the Protein Data Bank (http://www/rcsb.org/pdb).We are very grateful to many lecturers who have passed on their comments and

criticisms of the first edition of Inorganic Chemistry. Some of these remain anonymousto us and can be thanked only as ‘the review panel set up by Pearson Education.’ Inaddition to those colleagues whom we acknowledged in the preface to the first edition,we are grateful to Professors Duncan Bruce, Edwin Constable, Ronald Gillespie,Robert Hancock, Laura Hughes, Todd Marder, Christian Reber, David Tudela andKarl Wieghardt, and Drs Andrew Hughes and Mark Thornton-Pett who provided uswith a range of thought-provoking comments. We are, of course, indebted to the teamat Pearson Education who have supported the writing project and have taken themanuscript and graphics files through to their final form and provided their expertisefor the development of the accompanying website. Special thanks go to Bridget Allen,

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Kevin Ancient, Melanie Beard, Pauline Gillett, Simon Lake, Mary Lince, Paul Nash,Abigail Woodman and Ros Woodward.Having another inorganic chemist on-call in the house during the preparation of the

book has been more than beneficial: one of us owes much to her husband, EdwinConstable, for his critical comments. His insistence that a PC should replace the long-serving series of Macs has proved a bonus for the production of artwork. Finally, twobeloved feline companions have once again taken an active role (not always helpful) inthe preparation of this text – Philby and Isis have a unique ability to make sure theyare the centre of attention, no matter how many deadlines have to be met.

Catherine E. Housecroft (Basel)Alan G. Sharpe (Cambridge)

March 2004

Online resourcesVisit www.pearsoned.co.uk/housecroft to find valuable online resources

Companion Website for students

. Multiple choice questions to help test your learning

. Web-based problems for Chapter 3


. Interactive Periodic Table

For instructors

. Guide for lecturers


. PowerPoint slides

Also: The Companion Website provides the following features:

. Search tool to help locate specific items of content

. E-mail results and profile tools to send results of quizzes to instructors

. Online help and support to assist with website usage and troubleshooting

For more information please contact your local Pearson Education sales representative

or visit www.pearsoned.co.uk/housecroft

xxxii Preface to the second edition

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Preface to the first edition

Inorganic Chemistry has developed from the three editions of Alan Sharpe’s InorganicChemistry and builds upon the success of this text. The aim of the two books is thesame: to provide a single volume that gives a critical introduction to modern inorganicchemistry. However, in making the transition, the book has undergone a complete over-haul, not only in a complete rewriting of the text, but also in the general format, pedago-gical features and illustrations. These changes give Inorganic Chemistry a more modernfeel while retaining the original characteristic approach to the discussions, in particularof general principles of inorganic chemistry. Inorganic Chemistry provides students withnumerous fully-worked examples of calculations, extensive end-of-chapter problems,and ‘boxed’ material relating to chemical and theoretical background, chemical resources,the effects of chemicals on the environment and applications of inorganic chemicals. Thebook contains chapters on physical inorganic chemistry and descriptive chemistry of theelements. Descriptive chapters build upon the foundations laid in the earlier chapters.The material is presented in a logical order but navigation through the text is aided bycomprehensive cross-references. The book is completed by four ‘topic’ chapters coveringinorganic kinetics, catalysis, aspects of the solid state and bioinorganic chemistry. Eachchapter in the book ends with a summary and a checklist of new chemical terms. Thereading lists contain suggestions both for books and articles in the current literature.Additional information about websites of interest to readers of this book can be accessedvia: http://www.booksites.net/housecroftThe content of all descriptive chemistry chapters contains up-to-date information and

takes into account the results of the latest research; in particular, the chapters on organo-metallic chemistry of the s- and p-block and d-block elements reflect a surge in researchinterest in this area of chemistry. Another major development from Alan Sharpe’s originaltext has been to extend the discussion of molecular orbital theory, with an aim not only ofintroducing the topic but also showing how an objective (and cautious) approach canprovide insight into particular bonding features of molecular species. Greater emphasison the use of multinuclear NMR spectroscopy has been included; case studies introduceI > 1

2 nuclei and the observation of satellite peaks and applications of NMR spectroscopyare discussed where appropriate throughout the text. Appendices are included and are afeature of the book; they provide tables of physical data, selected character tables, anda list of abbreviations.Answers to non-descriptive problems are included in Inorganic Chemistry, but a

separate Solutions Manual has been written by Catherine Housecroft, and this givesdetailed answers or essay plans for all end of chapter problems.Most of the 3D-structural diagrams in the book have been drawn using Chem3D Pro.

with coordinates accessed from the Cambridge Crystallographic Data Base and imple-mented through the ETH in Zürich. The protein structures in Chapter 28 have beendrawn using Rasmol with data from the Protein Data Bank (http://www/rcsb.org/pdb).Suggestions passed on by readers of Alan Sharpe’s Inorganic Chemistry have helped us

to identify ‘holes’ and, in particular, we thank Professor Derek Corbridge. We gratefullyacknowledge comments made on the manuscript by members of the panel of reviewers(from the UK, the Netherlands and the US) set up by Pearson Education. A number ofcolleagues have read chapters of the manuscript and their suggestions and criticismshave been invaluable: special thanks go to Professors Steve Chapman, Edwin Constable,Michael Davies and Georg Süss-Fink, and Dr Malcolm Gerloch. We should also liketo thank Dr Paul Bowyer for information on sulfur dioxide in wine production, and

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Dr Bo Sundman for providing data for the iron phase-diagram. A text of this type cannotbecome reality without dedicated work from the publisher: from among those at PearsonEducation who have seen this project develop from infancy and provided us with support,particular thanks go to Lynn Brandon, Pauline Gillett, Julie Knight, Paul Nash, AlexSeabrook and Ros Woodward, and to Bridget Allen and Kevin Ancient for tireless anddedicated work on the design and artwork.One of us must express sincere thanks to her husband, Edwin Constable, for endless

discussions and critique. Thanks again to two very special feline companions, Philbyand Isis, who have sat, slept and played by the Macintosh through every minute of thewriting of this edition – they are not always patient, but their love and affection is anintegral part of writing.

Catherine E. HousecroftAlan G. Sharpe

June 2000

The publishers are grateful to the following for permission to reproduce copyright maerial:

Professor B. N. Figgis for Figure 20.20 from Figgis, B. N. (1966) Introduction to LigandFields, New York: Interscience.

In some instances we have been unable to trace the owners of copyright material, and wewould appreciate any information that would enable us to do so.

xxxiv Preface to the first edition

Black plate (35,1)

The rock salt (NaCl) lattice

In salts of formula MX, the coordination numbers of M and

X must be equal.

Rock salt (or halite, NaCl) occurs naturally as cubic crystals,

which, when pure, are colourless or white. Figure 5.15

shows two representations of the unit cell (see Section 5.2)

of NaCl. Figure 5.15a illustrates the way in which the

ions occupy the space available; the larger Cl� ions

(rCl� ¼ 181pm) define an fcc arrangement with the Naþ ions(rNaþ ¼ 102 pm) occupying the octahedral holes. Thisdescription relates the structure of the ionic lattice to the

close-packing-of-spheres model. Such a description is often

employed, but is not satisfactory for salts such as KF; while

this adopts an NaCl lattice, the Kþ and F� ions are almost

the same size (rKþ ¼ 138, rF� ¼ 133pm) (see Box 5.4).Although Figure 5.15a is relatively realistic, it hides most of

the structural details of the unit cell and is difficult to reproduce

when drawing the unit cell. The more open representation

shown in Figure 5.15b tends to be more useful.

The complete NaCl lattice is built up by placing unit cells

next to one another so that ions residing in the corner, edge

or face sites (Figure 5.15b) are shared between adjacent unit

cells. Bearing this in mind, Figure 5.15b shows that eachNaþ

and Cl� ion is 6-coordinate in the crystal lattice, while within

a single unit cell, the octahedral environment is defined

completely only for the central Naþ ion.

Figure 5.15b is not a unique representation of a unit cell of

the NaCl lattice. It is equally valid to draw a unit cell with

Naþ ions in the corner sites; such a cell has a Cl� ion in

the unique central site. This shows that the Naþ ions are

also in an fcc arrangement, and the NaCl lattice could there-

fore be described in terms of two interpenetrating fcc lattices,

one consisting of Naþ ions and one of Cl� ions.

Among the many compounds that crystallize with the

NaCl lattice are NaF, NaBr, NaI, NaH, halides of Li, K

and Rb, CsF, AgF, AgCl, AgBr, MgO, CaO, SrO, BaO,

MnO, CoO, NiO, MgS, CaS, SrS and BaS.

Worked example 5.2 Compound stoichiometry from a

unit cell

Show that the structure of the unit cell for sodium chloride

(Figure 5.15b) is consistent with the formula NaCl.

In Figure 5.15b, 14 Cl� ions and 13 Naþ ions are shown.

However, all but one of the ions are shared between two or

more unit cells.

There are four types of site:

. unique central position (the ion belongs entirely to theunit cell shown);

. face site (the ion is shared between two unit cells);

. edge sites (the ion is shared between four unit cells);

. corner site (the ion is shared between eight unit cells).

The total number of Naþ and Cl� ions belonging to the unit

cell is calculated as follows:

Site Number of Naþ Number of Cl�

Central 1 0Face 0 ð6� 12Þ ¼ 3Edge ð12� 14Þ ¼ 3 0Corner 0 ð8� 18Þ ¼ 1TOTAL 4 4

The ratio of Naþ :Cl� ions is 4 :4 ¼ 1 :1This ratio is consistent with the formula NaCl.

Fig. 5.15 Two representations of the unit cell of NaCl: (a) shows a space-filling representation, and (b) shows a ‘ball-and-stick’ representation which reveals the coordination environments of the ions. The Cl� ions are shown in green and the Naþ

ions in purple; since both types of ion are in equivalent environments, a unit cell with Naþ ions in the corner sites is also valid.There are four types of site in the unit cell: central (not labelled), face, edge and corner positions.

148 Chapter 5 . Structures and energetics of metallic and ionic solids

Self-study exercises

1. Show that the structure of the unit cell for caesium chloride(Figure 5.16) is consistent with the formula CsCl.

2. MgO adopts an NaCl lattice. How many Mg2þ and O2� ionsare present per unit cell? [Ans. 4 of each]

3. The unit cell of AgCl (NaCl type lattice) can be drawn with Agþ

ions at the corners of the cell, or Cl� at the corners. Confirmthat the number of Agþ and Cl� ions per unit cell remainsthe same whichever arrangement is considered.

The caesium chloride (CsCl) lattice

In the CsCl lattice, each ion is surrounded by eight others of

opposite charge. A single unit cell (Figure 5.16a) makes the

connectivity obvious only for the central ion. However, by

extending the lattice, one sees that it is constructed of

interpenetrating cubes (Figure 5.16b), and the coordination

number of each ion is seen. Because the Csþ and Cl� ions

are in the same environments, it is valid to draw a unit cell

either with Csþ or Cl� at the corners of the cube. Note the

relationship between the structure of the unit cell and bcc

packing.

The CsCl structure is relatively uncommon but is also

adopted by CsBr, CsI, TlCl and TlBr. At 298K, NH4Cl

and NH4Br possess CsCl lattices; [NH4]þ is treated as a

spherical ion (Figure 5.17), an approximation that can be

made for a number of simple ions in the solid state due to

their rotating or lying in random orientations about a fixed

point. Above 457 and 411K respectively, NH4Cl and

NH4Br adopt NaCl lattices.

The fluorite (CaF2) lattice

In salts of formula MX2, the coordination number of X mustbe half that of M.

Calcium fluoride occurs naturally as the mineral fluorite

(fluorspar). Figure 5.18a shows a unit cell of CaF2. Each

cation is 8-coordinate and each anion 4-coordinate; six of

the Ca2þ ions are shared between two unit cells and the

8-coordinate environment can be appreciated by envisaging

two adjacent unit cells. (Exercise: How does the coordination

number of 8 for the remaining Ca2þ ions arise?) Other

compounds that adopt this lattice type include group 2 metal

fluorides, BaCl2, and the dioxides of the f -block metals.

The antifluorite lattice

If the cation and anion sites in Figure 5.18a are exchanged,

the coordination number of the anion becomes twice that

of the cation, and it follows that the compound formula is

M2X. This arrangement corresponds to the antifluorite

structure, and is adopted by the group 1 metal oxides and

sulfides of type M2O and M2S; Cs2O is an exception.

The zinc blende (ZnS) lattice: a diamond-type network

Figure 5.18b shows the structure of zinc blende (ZnS). A

comparison of this with Figure 5.18a reveals a relationship

between the structures of zinc blende and CaF2; in going

from Figure 5.18a to 5.18b, half of the anions are removed

and the ratio of cation:anion changes from 1 :2 to 1 :1.

An alternative description is that of a diamond-type

network. Figure 5.19a gives a representation of the structure

of diamond; each C atom is tetrahedrally sited and the

structure is very rigid. This structure type is also adopted

by Si, Ge and a-Sn (grey tin). Figure 5.19b (with atomlabels that relate it to Figure 5.19a) shows a view of the

diamond network that is comparable with the unit cell of

zinc blende in Figure 5.18b. In zinc blende, every other site

in the diamond-type array is occupied by either a zinc or a

sulfur centre. The fact that we are comparing the structure

of an apparently ionic compound (ZnS) with that of a

covalently bonded species should not cause concern. As we

have already mentioned, the hard sphere ionic model is a

convenient approximation but does not allow for the fact

Fig. 5.16 (a) The unit cell of CsCl; Csþ ions are shownin yellow and Cl� in green, but the unit cell could also

be drawn with the Csþ ion in the central site. The unit cell isdefined by the yellow lines. (b) One way to describe the CsCllattice is in terms of interpenetrating cubic units of Csþ andCl� ions.

Fig. 5.17 The [NH4]þ ion can be treated as a sphere in

descriptions of solid state lattices; some other ions (e.g.[BF4]

�, [PF6]�) can be treated similarly.

Chapter 5 . Ionic lattices 149

emerald and aquamarine. Magnesium and calcium are the

eighth and fifth most abundant elements, respectively, in

the Earth’s crust, and Mg, the third most abundant in the

sea. The elements Mg, Ca, Sr and Ba are widely distributed

in minerals and as dissolved salts in seawater; some

important minerals are dolomite (CaCO3�MgCO3), mag-nesite (MgCO3), olivine ((Mg,Fe)2SiO4), carnallite (KCl�MgCl2�6H2O), CaCO3 (in the forms of chalk, limestoneand marble), gypsum (CaSO4�2H2O), celestite (SrSO4),strontianite (SrCO3) and barytes (BaSO4). The natural

abundances of Be, Sr and Ba are far less than those of Mg

and Ca (Figure 11.1).

Extraction

Of the group 2 metals, only Mg is manufactured on a large

scale (see Box 11.1). Dolomite is thermally decomposed to

a mixture of MgO and CaO, and MgO is reduced by

ferrosilicon in Ni vessels (equation 11.1); Mg is removed

by distillation in vacuo.

2MgOþ 2CaOþ FeSi��"1450K

2Mgþ Ca2SiO4 þ Fe ð11:1Þ

Extraction of Mg by electrolysis of fused MgCl2 is also

important and is applied to the extraction of the metal

from seawater. The first step is precipitation (see Table

6.4) of Mg(OH)2 by addition of Ca(OH)2 (slaked lime),

produced from CaCO3 (available as various calcareous

deposits, see Figure 10.5). Neutralization with hydrochloric

acid (equation 11.2) and evaporation of water gives

MgCl2�xH2O, which, after heating at 990K, yields the

Fig. 11.1 Relative abundances in the Earth’s crust of thealkaline earth metals (excluding Ra); the data are plotted on alogarithmic scale. The units of abundance are ppm.

RESOURCES, ENVIRONMENTAL AND BIOLOGICAL

Box 11.1 Recycling of materials: magnesium

Recycling of materials became increasingly important duringthe last decades of the twentieth century, and continues tohave a significant influence on chemical industries. A largefraction of the total Mg consumed is in the form of Al/Mg

alloys (see Figure 11.2), and recycling of Al cans necessarilymeans recovery of Mg. The graph below shows the variationin total consumption of primary Mg in the US from 1960 to2000, and the increasing trend towards recovering the metal.

[Data from US Geological Survey]

276 Chapter 11 . The group 2 metals

Problems

11.1 (a) Write down, in order, the names and symbols of the

metals in group 2; check your answer by reference to thefirst page of this chapter. Which metals are classed asalkaline earth metals? (b) Give a general notation thatshows the ground state electronic configuration of each

metal.

11.2 Using data in Table 6.4, determine the relative solubilities

of Ca(OH)2 and Mg(OH)2 and explain the relevance ofyour answer to the extraction of magnesium fromseawater.

11.3 (a) Write an equation to show how Mg reacts with N2when heated. (b) Suggest how the product reacts with

water.

11.4 The structure of magnesium carbide, MgC2, is of the

NaCl type, elongated along one axis. (a) Explain how thiselongation arises. (b) What do you infer from the fact thatthere is no similar elongation in NaCN which also

crystallizes with a NaCl lattice?

11.5 Write balanced equations for the following reactions:

(a) the thermal decomposition of [NH4]2[BeF4];(b) the reaction between NaCl and BeCl2;(c) the dissolution of BeF2 in water;

11.6 (a) Suggest a likely structure for the dimer of BeCl2,present in the vapour phase below 1020K. Whathybridization scheme is appropriate for the Be centres?

(b) BeCl2 dissolves in diethyl ether to form monomericBeCl2�2Et2O; suggest a structure for this compound andgive a description of the bonding.

11.7 MgF2 has a TiO2 lattice. (a) Sketch a unit cell of MgF2,and (b) confirm the stoichiometry of MgF2 using the solid

state structure.

11.8 Discuss the trends in data in Table 11.4.

11.9 (a) How do anhydrous CaCl2 and CaH2 function asdrying agents? (b) Compare the solid state structures and

properties of BeCl2 and CaCl2.

11.10 How would you attempt to estimate the following?(a) �rH

o for the solid state reaction:

MgCl2 þMg��" 2MgCl

(b) �rHo for the reaction:

CaCO3ðcalciteÞ ��"CaCO3ðaragoniteÞ

11.11 (a) Identify the conjugate acid–base pairs in reaction11.23. (b) Suggest how BaO2 will react with water.

11.12 (a) Determine �rHo for the reactions of SrO and BaO

with water, given that values of �fHo(298K) for SrO(s),

BaO(s), Sr(OH)2(s), Ba(OH)2(s) and H2O(l) are �592.0,�553.5, �959.0, �944.7 and �285.5 kJmol�1respectively. (b) Compare the values of�rH

o with that for

the reaction of CaO with water (equation 11.5), andcomment on factors contributing to the trend in values.

11.13 (a) What qualitative test is used for CO2? (b) Whatreaction takes place, and (c) what is observed in a positive

test?

11.14 Discuss the data presented in Table 11.5; other relevantdata are available in this book.

11.15 Write a short account that justifies the so-called diagonal

relationship between Li and Mg.

11.16 Suggest why MgO is more soluble in aqueous MgCl2solution than in pure water.

Overview problems

11.17 Suggest explanations for the following observations.(a) The energy released when a mole of crystalline BaO is

formed from its constituent ions is less than that

released when a mole of MgO forms from its ions.(Note: Each compound possesses an NaCl lattice.)

(b) Despite being a covalent solid, BeF2 is very soluble in

water.(c) At 298K, Be adopts an hcp lattice; above 1523K, the

coordination number of a Be atom in elemental

beryllium is 8.

11.18 Comment on the following statements.(a) Na2S adopts a solid state structure that is related to

that of CaF2.(b) [C3]

4�, CO2 and [CN2]2� are isoelectronic species.

(c) Be(OH)2 is virtually insoluble in water, but is solublein aqueous solutions containing excess hydroxideions.

(d) MgO is used as a refractory material.Table 11.4 Data for problem 11.8.

Metal, M �fHo / kJmol�1

MF2 MCl2 MBr2 MI2

Mg �1113 �642 �517 �360Ca �1214 �795 �674 �535Sr �1213 �828 �715 �567Ba �1200 �860 �754 �602

Table 11.5 Data for problem 11.14: log K for the formation ofthe complexes [M(crypt-222)]nþ.

Mnþ Naþ Kþ Rbþ Mg2þ Ca2þ Sr2þ Ba2þ

log K 4.2 5.9 4.9 2.0 4.1 13.0 >15

Chapter 11 . Problems 291

Key definitions are highlighted in the text.

Worked examples are giventhroughout the text.

Topic boxes relate inorganic chemistry to real-lifeexamples of environmental and biological resources,illustrate applications, and provide information onchemical and theoretical background.

Self-study exercises allow students to testtheir understanding of what they have read.

Web icons indicate that a 3D rotatable graphic is available on thecompanion website (see p. xxxvi).

End-of chapter problems, including a set of‘overview’ problems, each covering a broadrange of material from the chapter.

Guided tour of the book xxxv

Black plate (36,1)

Multiple choice questions with immediate online results

Chime-viewable rotatable molecules

xxxvi Guided tour of the student website

Black plate (1,1)

Some basic concepts

1.1 Introduction

Inorganic chemistry: it is not an isolatedbranch of chemistry

If organic chemistry is considered to be the ‘chemistry of

carbon’, then inorganic chemistry is the chemistry of all

elements except carbon. In its broadest sense, this is true,

but of course there are overlaps between branches of

chemistry. A topical example is the chemistry of the fuller-

enes (see Section 13.4 ) including C60 (see Figure 13.5) and

C70; this was the subject of the award of the 1996 Nobel

Prize in Chemistry to Professors Sir Harry Kroto, Richard

Smalley and Robert Curl. An understanding of such

molecules and related species called nanotubes involves

studies by organic, inorganic and physical chemists as well

as by physicists and materials scientists.

Inorganic chemistry is not simply the study of elements

and compounds; it is also the study of physical principles.

For example, in order to understand why some compounds

are soluble in a given solvent and others are not, we apply

laws of thermodynamics. If our aim is to propose details of

a reaction mechanism, then a knowledge of reaction kinetics

is needed. Overlap between physical and inorganic chemistry

is also significant in the study of molecular structure. In the

solid state, X-ray diffraction methods are routinely used to

obtain pictures of the spatial arrangements of atoms in a

molecule or molecular ion. To interpret the behaviour of

molecules in solution, we use physical techniques such as

nuclear magnetic resonance (NMR) spectroscopy; the

equivalence or not of particular nuclei on a spectroscopic

timescale may indicate whether a molecule is static or under-

going a dynamic process (see Section 2.11). In this text, we

describe the results of such experiments but we will not, in

general, discuss underlying theories; several texts which

cover experimental details of such techniques are listed at

the end of Chapter 1.

The aims of Chapter 1

In this chapter, we outline some concepts fundamental to an

understanding of inorganic chemistry. We have assumed

that readers are to some extent familiar with most of these

concepts and our aim is to give a point of reference for

review purposes.

1.2 Fundamental particles of an atom

An atom is the smallest unit quantity of an element that iscapable of existence, either alone or in chemical combination

with other atoms of the same or another element. Thefundamental particles of which atoms are composed are theproton, electron and neutron.

Chapter

1TOPICS

& Fundamental particles

& Atomic number, mass number and isotopes

& An overview of quantum theory

& Orbitals of the hydrogen atom and quantumnumbers

& The multi-electron atom, the aufbau principleand electronic configurations

& The periodic table

& Ionization energies and electron affinities

& Lewis structures

& Valence bond theory

& Fundamentals of molecular orbital theory

& The octet rule

& Electronegativity

& Dipole moments

& MO theory: heteronuclear diatomic molecules

& Isoelectronic molecules

& Molecular shape and the VSEPR model

& Geometrical isomerism

Black plate (2,1)

A neutron and a proton have approximately the same mass

and, relative to these, an electron has negligible mass (Table

1.1). The charge on a proton is positive and of equal mag-

nitude, but opposite sign, to that on a negatively charged

electron; a neutron has no charge. In an atom of any

element, there are equal numbers of protons and electrons

and so an atom is neutral. The nucleus of an atom consists

of protons and (with the exception of protium; see Section

9.3) neutrons, and is positively charged; the nucleus of pro-

tium consists of a single proton. The electrons occupy a

region of space around the nucleus. Nearly all the mass of

an atom is concentrated in the nucleus, but the volume of

the nucleus is only a tiny fraction of that of the atom; the

radius of the nucleus is about 10�15 m while the atom

itself is about 105 times larger than this. It follows that

the density of the nucleus is enormous, more than 1012

times than of the metal Pb.

Although chemists tend to consider the electron, proton

and neutron as the fundamental (or elementary) particles of

an atom, particle physicists would disagree, since their

research shows the presence of yet smaller particles.

1.3 Atomic number, mass number andisotopes

Nuclides, atomic number and mass number

A nuclide is a particular type of atom and possesses a charac-

teristic atomic number,Z, which is equal to the number of pro-

tons in the nucleus; because the atom is electrically neutral, Z

also equals the number of electrons. Themass number, A, of a

nuclide is the number of protons and neutrons in the nucleus.

A shorthand method of showing the atomic number and mass

number of a nuclide along with its symbol, E, is:

EAZ

Mass number

Atomic numberNe20

10e.g.Element symbol

Atomic number ¼ Z ¼ number of protons in the nucleus =number of electrons

Mass number ¼ A ¼ number of protonsþ number ofneutrons

Number of neutrons ¼ A� Z

Relative atomic mass

Since the electrons are of minute mass, the mass of an atom

essentially depends upon the number of protons and neu-

trons in the nucleus. As Table 1.1 shows, the mass of a

single atom is a very small, non-integral number, and for

convenience we adopt a system of relative atomic masses.

We define the atomic mass unit as 1/12th of the mass of a126C atom so that it has the value 1:660� 10�27 kg. Relativeatomic masses (Ar) are thus all stated relative to126C¼ 12.0000. The masses of the proton and neutron canbe considered to be �1 u where u is the atomic mass unit(1 u � 1:660� 10�27 kg).

Isotopes

Nuclides of the same element possess the same number of

protons and electrons but may have different mass numbers;

the number of protons and electrons defines the element but

the number of neutrons may vary. Nuclides of a particular

element that differ in the number of neutrons and, therefore,

their mass number, are called isotopes (see Appendix 5).

Isotopes of some elements occur naturally while others

may be produced artificially.

Elements that occur naturally with only one nuclide are

monotopic and include phosphorus, 3115P, and fluorine,199F.

Elements that exist as mixtures of isotopes include C (126C

and 136C) and O (168O,

178O and

188O). Since the atomic

number is constant for a given element, isotopes are often

distinguished only by stating the atomic masses, e.g. 12C

and 13C.

Worked example 1.1 Relative atomic mass

Calculate the value of Ar for naturally occurring chlorine

if the distribution of isotopes is 75.77% 3517Cl and 24.23%3717Cl. Accurate masses for

35Cl and 37Cl are 34.97 and 36.97.

The relative atomic mass of chlorine is the weighted mean

of the mass numbers of the two isotopes:

Relative atomic mass,

Ar ¼�75:77

100� 34:97

�þ�24:23

100� 36:97

�¼ 35:45

Table 1.1 Properties of the proton, electron and neutron.

Proton Electron Neutron

Charge /C þ1:602� 10�19 �1:602� 10�19 0Charge number (relative charge) 1 �1 0Rest mass / kg 1:673� 10�27 9:109� 10�31 1:675� 10�27Relative mass 1837 1 1839

2 Chapter 1 . Some basic concepts

Black plate (3,1)

Self-study exercises

1. If Ar for Cl is 35.45, what is the ratio of35Cl:37Cl present in a

sample of Cl atoms containing naturally occurring Cl?[Ans. 3.17:1 ]

2. Calculate the value of Ar for naturally occurring Cu if the

distribution of isotopes is 69.2% 63Cu and 30.8% 65Cu;accurate masses are 62.93 and 64.93. [Ans. 63.5 ]

3. Why in question 2 is it adequate to write 63Cu rather than6329Cu?

4. Calculate Ar for naturally occurring Mg if the isotope distribu-

tion is 78.99% 24Mg, 10.00% 25Mg and 11.01% 26Mg;accurate masses are 23.99, 24.99 and 25.98. [Ans. 24.31]

Isotopes can be separated bymass spectrometry and Figure

1.1a shows the isotopic distribution in naturally occurring

Ru. Compare this plot (in which the most abundant

isotope is set to 100) with the values listed in Appendix 5.

Figure 1.1b shows a mass spectrometric trace for molecular

S8, the structure of which is shown in Figure 1.1c; five

peaks are observed due to combinations of the isotopes of

sulfur. (See problem 1.3 at the end of this chapter.)

Isotopes of an element have the same atomic number, Z, butdifferent atomic masses.

1.4 Successes in early quantum theory

We saw in Section 1.2 that electrons in an atom occupy a

region of space around the nucleus. The importance of

electrons in determining the properties of atoms, ions

and molecules, including the bonding between or within

them, means that we must have an understanding of the

CHEMICAL AND THEORETICAL BACKGROUND

Box 1.1 Isotopes and allotropes

Do not confuse isotope and allotrope! Sulfur exhibits both iso-

topes and allotropes. Isotopes of sulfur (with percentagenaturally occurring abundances) are 3216S (95.02%),

3316S

(0.75%), 3416S (4.21%),3616S (0.02%).

Allotropes of an element are different structural modifications

of that element. Allotropes of sulfur include cyclic structures, e.g.

S6 (see below) and S8 (Figure 1.1c), and Sx-chains of variouslengths (polycatenasulfur).

Further examples of isotopes and allotropes appear

throughout the book.

Fig. 1.1 Mass spectrometric traces for (a) atomic Ru and (b) molecular S8; the mass:charge ratio is m/z and in these tracesz ¼ 1. (c) The molecular structure of S8.

Chapter 1 . Successes in early quantum theory 3

Black plate (4,1)

electronic structures of each species. No adequate discussion

of electronic structure is possible without reference to

quantum theory and wave mechanics. In this and the next

few sections, we review some of the crucial concepts. The

treatment is mainly qualitative, and for greater detail and

more rigorous derivations of mathematical relationships,

the references at the end of Chapter 1 should be consulted.

The development of quantum theory took place in two

stages. In the older theories (1900–1925), the electron was

treated as a particle, and the achievements of greatest signif-

icance to inorganic chemistry were the interpretation of

atomic spectra and assignment of electronic configurations.

In more recent models, the electron is treated as a wave

(hence the name wave mechanics) and the main successes in

chemistry are the elucidation of the basis of stereochemistry

and methods for calculating the properties of molecules

(exact only for species involving light atoms).

Since all the results obtained by using the older quantum

theory may also be obtained from wave mechanics, it may

seem unnecessary to refer to the former; indeed, sophisticated

treatments of theoretical chemistry seldom do. However,

most chemists often find it easier andmore convenient to con-

sider the electron as a particle rather than a wave.

Some important successes of classicalquantum theory

Historical discussions of the developments of quantum

theory are dealt with adequately elsewhere, and so we

focus only on some key points of classical quantum theory

(in which the electron is considered to be a particle).

At low temperatures, the radiation emitted by a hot body

is mainly of low energy and occurs in the infrared, but as the

temperature increases, the radiation becomes successively

dull red, bright red and white. Attempts to account for this

observation failed until, in 1901, Planck suggested that

energy could be absorbed or emitted only in quanta of mag-

nitude �E related to the frequency of the radiation, �, by

equation 1.1. The proportionality constant is h, the Planck

constant (h ¼ 6:626� 10�34 J s).

�E ¼ h� Units: E in J; � in s�1 or Hz ð1:1Þc ¼ �� Units: � in m; � in s�1 or Hz ð1:2Þ

The hertz, Hz, is the SI unit of frequency.

Since the frequency of radiation is related to the wave-

length, �, by equation 1.2, in which c is the speed of light

in a vacuum (c ¼ 2:998� 108 m s�1), we can rewrite equation1.1 in the form of equation 1.3 and relate the energy of

radiation to its wavelength.

�E ¼ hc�

ð1:3Þ

On the basis of this relationship, Planck derived a relative

intensity/wavelength/temperature relationship which was in

good agreement with experimental data. This derivation is

not straightforward and we shall not reproduce it here.

One of the most important applications of early quantum

theory was the interpretation of the atomic spectrum of

hydrogen on the basis of the Rutherford–Bohr model of

the atom. When an electric discharge is passed through a

sample of dihydrogen, the H2 molecules dissociate into

atoms, and the electron in a particular excited H atom

may be promoted to one of many high energy levels.

These states are transient and the electron falls back to a

lower energy state, emitting energy as it does so. The

consequence is the observation of spectral lines in the

emission spectrum of hydrogen; the spectrum (a small

part of which is shown in Figure 1.2) consists of groups

of discrete lines corresponding to electronic transitions,

each of discrete energy. As long ago as 1885, Balmer

pointed out that the wavelengths of the spectral lines

observed in the visible region of the atomic spectrum of

hydrogen obeyed equation 1.4, in which R is the Rydberg

constant for hydrogen, �� is the wavenumber in cm�1, and

n is an integer 3, 4, 5 . . . This series of spectral lines is

known as the Balmer series.

Wavenumber ¼ reciprocal of wavelength; convenient (non-SI) units are ‘reciprocal centimetres’, cm�1

�� ¼ 1�¼ R

�1

22� 1n2

�ð1:4Þ

R ¼ Rydberg constant for hydrogen

¼ 1:097� 107 m�1 ¼ 1:097� 105 cm�1

Other series of spectral lines occur in the ultraviolet (Lyman

series) and infrared (Paschen, Brackett and Pfund series). All

Fig. 1.2 Part of the emission spectrum of atomic hydrogen. Groups of lines have particular names, e.g. Balmer and Lyman series.

4 Chapter 1 . Some basic concepts

Black plate (5,1)

lines in all the series obey the general expression given in

equation 1.5 where n’ > n. For the Lyman series, n ¼ 1, forthe Balmer series, n ¼ 2, and for the Paschen, Brackett andPfund series, n ¼ 3, 4 and 5 respectively. Figure 1.3 showssome of the allowed transitions of the Lyman and Balmer

series in the emission spectrum of atomic H. Note the use

of the word allowed; the transitions must obey selection

rules, to which we return in Section 20.6.

�� ¼ 1�¼ R

�1

n2� 1n’2

�ð1:5Þ

Bohr’s theory of the atomic spectrum ofhydrogen

In 1913, Niels Bohr combined elements of quantum theory

and classical physics in a treatment of the hydrogen atom.

He stated two postulates for an electron in an atom:

. Stationary states exist in which the energ

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