Quantum Chemistry: A Unified Approach

by David B Cook

Overview

This book is a presentation of a qualitative theory of chemical bonding stressing the physical processes which occur on bond formation. It differs from most (if not all) other books in that it does not seek to “rationalize” the phenomena of bonding by a series of mnemonic rules. A principal feature is a unified and consistent treatment across all types of bonding in organic, physical and inorganic chemistry.

Product Details

ISBN-13:	9781848162655
Publisher:	Imperial College Press
Publication date:	11/16/2008
Pages:	276
Product dimensions:	6.20(w) x 9.10(h) x 0.80(d)

Table of Contents

Preface xi

Acknowledgments xiii

1 How Science Deals with Complex Problems 1

1.1 Introduction: Levels in Science 2

1.2 What are Molecules Made Of? 4

1.3 Interactions Between Atoms 6

1.4 The Simplest examples: H[subscript 2] and LiH 8

1.4.1 The Hydrogen Molecule 8

1.4.2 The Lithium Hydride Molecule 10

1.4.2.1 What About the Other Li Electrons? 14

1.4.2.2 What About the Nuclear Repulsions? 14

1.4.3 Comments on H[subscript 2] and LiH 15

1.5 How to Proceed? 16

Appendix A How to Interpret 3D Contours 19

A.1 Thinking in 3D 19

A.2 The Electron Distribution of the Lithium 2s Electron 22

A.2.1 How does this relate to the Text-Book "Orbitals" 25

A.2.2 What if the Distribution is not Spherical? 25

Appendix B Must We Use Quantum Theory? 29

B.1 Connections to Laws of Nature 29

B.2 Stable Molecules 30

B.3 The Equipartition of Energy 31

B.4 Quantum Summary 33

2 What We Know about Atoms and Molecules 35

2.1 Atomic Electronic Structure 35

2.1.1 The Hydrogen Atom 36

2.1.2 Many-electron Atoms 38

2.1.3 The Pauli Principle 40

2.1.3.1 Statement of the Pauli Principle 40

2.1.4 Current Summary for Atoms 41

2.2 Empirical Chemistry 41

Appendix C The Interpretation of Orbitals 45

C.1 What is an Orbital? 45

C.2 Orbitals: Atomic and Molecular 48

3 A Strategy for Electronic Structure 49

3.1 Review 49

3.2 Lithium Hydride Again 52

3.2.1 Polarisation and Hybrid AOs 53

3.2.2 Molecular Orbitals 55

3.2.2.1 Quick Summary 59

Appendix D Is Hybridisation a Real Process? 61

4 The Pauli Principle and Orbitals 63

4.1 A Difficulty with Helium 64

4.2 When are Orbitals Mutually Exclusive? 66

4.3 Does This Work for AOs? 69

4.4 TheHelium Molecule-Again 72

4.5 The Role of Atomic Orbitals in Valence Theory 75

4.6 Current Summary for LiH and "He[subscript 2]" 76

5 A Model Polyatomic: Methane 79

5.1 The Methane Molecule: CH[subscript 4] 79

5.2 The Electronic Structure of Methane 81

5.3 The Shape of the Methane Molecule 83

5.4 What About the Pauli Principle? 84

5.4.1 Preliminary Summary for Methane 85

5.5 The Chemist's Description of Methane 86

5.5.1 How to Use these Structures: the Valence Bond Method 88

5.6 Summary for Methane 91

6 Lone Pairs of Electrons 93

6.1 Why are Not All Electrons Involved in Bonding? 93

6.2 What is a Lone Pair? 96

6.2.1 The Ammonia Molecule 97

6.2.2 The Water Molecule 101

6.3 The Shapes of Simple Molecules 102

6.3.1 The Water Molecule-Again 103

6.4 "Reactions" of Lone pairs 105

6.5 A Working Summary 106

7 Organic Molecules with Multiple Bonds 107

7.1 Double and Triple Bonds 108

7.2 The Possibilities 110

7.3 Ethene and Methanal 112

7.4 The Double Bond in Ethene and Methanal 113

7.4.1 Sigma ([sigma]) and Pi ([pi]) Notation in Planar Molecules 115

7.5 The [sigma] and [pi] Orbitals in C[subscript 2]H[subscript 4] and CH[subscript 2]O 117

7.5.1 Ethene Contours 117

7.5.2 Methanal Contours 120

7.5.3 Relative Energies of the Two Bonds 122

7.6 Reactivity of a Double Bond 123

7.7 Multiple Bonds in General 124

8 Molecular Symmetry 125

8.1 The Question of Symmetry 125

8.2 Symmetry: Generalisation 128

8.3 Case Studies: H[subscript 2]O and Benzene 129

8.3.1 The H[subscript 2]O Molecule 129

8.3.2 The Benzene [sigma] system 132

8.4 Bond MOs and Symmetry MOs 135

8.5 A Cautionary Note 136

9 Diatomics with Multiple Bonds 139

9.1 Motivation 140

9.2 The Nitrogen Molecule: N[subscript 2] 140

9.2.1 Energies of the N[subscript 2] MOs 143

9.2.2 Symmetry and the N[subscript 2] Molecule 146

9.3 The Carbon Monoxide Molecule: CO 148

9.4 Other Homonuclear Diatomics 150

9.4.1 The Oxygen Molecule: O[subscript 2] 151

9.5 Lessons from Diatomics 153

10 Dative Bonds 155

10.1 Introduction: Familiar Reactions 155

10.1.1 "Solvation" 156

10.1.2 A Reactive Lone Pair: the CO Molecule 159

10.1.3 CO and Transition-metal Atoms 161

10.2 The Dative Bond: Summary 161

11 Delocalised Electronic Substructures: Aromaticity 163

11.1 The Benzene Molecule 163

11.2 Delocalised Electrons 167

11.3 Environment-insensitive [pi] Substructures? 172

11.4 Nomenclature and Summary 175

12 Organic and Inorganic Chemistry 177

12.1 Commentary on Results 177

12.2 Nitric Acid and Related Molecules 178

12.2.1 The Nitrate Ion NO[subscript 3 superscript -] 183

12.3 Carbonic acid and Carbonates 185

12.4 Sulphuric Acid and Sulphates 185

13 Further Down the Periodic Table 187

13.1 The Effect of Increasing Atomic Number 188

13.2 The Possible Demise of Lone Pairs 189

13.3 A Particular Case: Sulphur 190

13.4 The General Case: "Hypervalence" 192

13.4.1 Single or double bonds? 194

13.4.2 The Steric Effect 195

13.5 How to Describe These Bonds? 196

13.5.1 A Comparison: 16 valence electrons 198

13.6 An Updated Summary 203

14 Reconsidering Empirical Rules 205

14.1 Limitations of the Octet Rule 205

14.2 The Basis of the Octet Rule 206

14.3 Population Analysis 209

14.4 Resonance and Resonance Hybrids 212

14.5 Oxidation Number 213

14.6 Summary for Number Rules 215

15 Mavericks and other Lawbreakers 219

15.1 Exceptions to the Rules 220

15.2 Boron Hydrides and Bridges 221

15.2.1 The Expected Compound: BH[subscript 3] 222

15.2.2 The Compounds Which Are Found 223

15.2.3 Bridged, Three-Centre, Bonds 223

15.3 Other Three-Centre Bonds? 226

15.4 Metals and Crystals 229

15.4.1 Metals 230

15.4.2 Crystals 231

15.5 The Hydrogen Bond 233

15.6 Lawbreakers? 234

16 The Transition Elements 235

16.1 The Background 236

16.2 Transition Metals: effects of "d" electrons 237

16.3 "Screening" in the Electronic Structure of Atoms 238

16.4 History and Apology 241

16.4.1 The "Crystal" Model 241

16.4.2 The Molecular Orbital Model 242

16.4.3 The "Chemical" Model 245

16.4.4 Apology 246

16.5 Comments 247

17 Omissions and Conclusions 251

17.1 Omissions 251

17.1.1 Intermolecular Forces 252

17.1.2 Chemical Reactions 252

17.2 Conclusions 255