Atkins Inorganic Chem

Descriptive

Inorganic Chemistry

F O U R T H E D I T I O N

Geoff Rayner-CanhamSir Wilfred Grenfell College Memorial University

Tina OvertonUniversity of Hull

W. H. FREEMAN AND COMPANY

NEW YORK

The cover illustration shows an updated and modied version of the Hull periodic table, devised by Dr. P. G. Nelson at the University of Hull, U.K. The basic shape of this periodic table was suggested by T. Bayley in 1882. The idea of grading the connections and showing them by different kinds of line was proposed by I. D. Margary (1921). The Hull table takes these ideas and applies them to all the elements. For further details, see Education in Chemistry 24 (1987): 17173; 25 (1988): 3. A similar table was suggested by W. B. Jensen in the United States.

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iii

Overview

What Is Descriptive Inorganic Chemistry? . . . . . . . . . xiv

Preface . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xv

Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . xviii

Dedication . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xx

C H A P T E R 1 The Electronic Structure of the Atom: A Review . . . . . 1

C H A P T E R 2 An Overview of the Periodic Table . . . . . . . . . . . . . . . . 18

C H A P T E R 3 Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

C H A P T E R 4 Metallic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

C H A P T E R 5 Ionic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

C H A P T E R 6 Inorganic Thermodynamics . . . . . . . . . . . . . . . . . . . . . 108

C H A P T E R 7 Solvent Systems and Acid-Base Behavior . . . . . . . . . 129

C H A P T E R 8 Oxidation and Reduction . . . . . . . . . . . . . . . . . . . . . . . 159

C H A P T E R 9 Periodic Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 184

C H A P T E R 1 0 Hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 221

C H A P T E R 1 1 The Group 1 Elements: The Alkali Metals . . . . . . . . . 240

C H A P T E R 1 2 The Group 2 Elements: The Alkaline Earth Metals . . 263

C H A P T E R 1 3 The Group 13 Elements . . . . . . . . . . . . . . . . . . . . . . . . 282




C H A P T E R 1 7 The Group 17 Elements: The Halogens . . . . . . . . . . . 441

C H A P T E R 1 8 The Group 18 Elements: The Noble Gases . . . . . . . . 473

C H A P T E R 1 9 Introduction to Transition Metal Complexes . . . . . . . 484

C H A P T E R 2 0 Properties of Transition Metals . . . . . . . . . . . . . . . . . . 513


C H A P T E R 2 2 Organometallic Chemistry . . . . . . . . . . . . . . . . . . . . . . 577

C H A P T E R 2 3 The Rare Earth and Actinoid Elements . . . . . . . . . . 617w

Appendices . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A-1

Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . I-1

On the Web www.whfreeman.com/rayner4e

Contents

What Is Descriptive Inorganic Chemistry? xiv

Preface xv

Acknowledgments xviii

Dedication xx

C H A P T E R 1 The Electronic Structure of the Atom: A Review . . . . . . . . . 1

1.1 The Schrdinger Wave Equation and Its

Signicance . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

Atomic Absorption Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . 3

1.2 Shapes of the Atomic Orbitals . . . . . . . . . . . . . . . . . . . 6

1.3 The Polyelectronic Atom . . . . . . . . . . . . . . . . . . . . . . . 9

1.4 Ion Electron Congurations . . . . . . . . . . . . . . . . . . . . 13

1.5 Magnetic Properties of Atoms . . . . . . . . . . . . . . . . . . 14

C H A P T E R 2 An Overview of the Periodic Table . . . . . . . . . . . . . . . . . . . . 18

2.1 Organization of the Modern Periodic Table . . . . . . . 20

2.2 Existence of the Elements . . . . . . . . . . . . . . . . . . . . . 22

2.3 Stability of the Elements and Their Isotopes . . . . . . 23

The Origin of the Shell Model of the Nucleus . . . . . . . . . . . . 25

2.4 Classication of the Elements . . . . . . . . . . . . . . . . . . 27

Medicinal Inorganic Chemistry: An Introduction . . . . . . . . . 28

2.5 Periodic Properties: Atomic Radius . . . . . . . . . . . . . 30

2.6 Periodic Properties: Ionization Energy . . . . . . . . . . 33

2.7 Periodic Properties: Electron Afnity . . . . . . . . . . . . 35

Alkali Metal Anions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37

2.8 Biochemistry of the Elements . . . . . . . . . . . . . . . . . . 37

C H A P T E R 3 Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

3.1 Theories of Covalent Bonding . . . . . . . . . . . . . . . . . . 42

3.2 Introduction to Molecular Orbital Theory . . . . . . . . 43

3.3 Molecular Orbitals for Period 1 Diatomic

Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44

3.4 Molecular Orbitals for Period 2 Diatomic

Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46

3.5 Molecular Orbitals for Heteronuclear Diatomic

Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49

C o n t e n t s v

3.6 A Brief Review of the Lewis Theory . . . . . . . . . . . . . 51

3.7 Partial Bond Order . . . . . . . . . . . . . . . . . . . . . . . . . . . 53

3.8 Formal Charge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 53

3.9 Valence Shell Electron Pair Repulsion Theory . . . . 54

3.10 Valence-Bond Theory . . . . . . . . . . . . . . . . . . . . . . . . . 59

3.11 Network Covalent Substances . . . . . . . . . . . . . . . . . . 61

3.12 Intermolecular Forces . . . . . . . . . . . . . . . . . . . . . . . . 62

The Origins of the Electronegativity Concept . . . . . . . . . . . . 64

3.13 Molecular Symmetry . . . . . . . . . . . . . . . . . . . . . . . . . . 66

3.14 Symmetry and Vibrational Spectroscopy . . . . . . . . . 71

3.15 Radiation Trapping: The Greenhouse Effect . . . . . . 73

3.16 Covalent Bonding and the Periodic Table . . . . . . . . 75

C H A P T E R 4 Metallic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

4.1 Metallic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

4.2 Bonding Models . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 81

4.3 Structure of Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . 83

4.4 Unit Cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85

Memory Metal: The Shape of Things to Come . . . . . . . . . . . 86

4.5 Alloys . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 87

C H A P T E R 5 Ionic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

5.1 The Ionic Model and the Size of Ions . . . . . . . . . . . . 90

5.2 Polarization and Covalency . . . . . . . . . . . . . . . . . . . . 92

5.3 Hydrated Salts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95

5.4 Ionic Crystal Structures . . . . . . . . . . . . . . . . . . . . . . . . 95

5.5 The Bonding Continuum . . . . . . . . . . . . . . . . . . . . . 101

Concrete: An Old Material with a New Future . . . . . . . . . . 104

C H A P T E R 6 Inorganic Thermodynamics . . . . . . . . . . . . . . . . . . . . . . . . . 108

6.1 Thermodynamics of the Formation of

Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 109

6.2 Formation of Ionic Compounds . . . . . . . . . . . . . . . . 115

6.3 The Born-Haber Cycle . . . . . . . . . . . . . . . . . . . . . . . 116

6.4 Thermodynamics of the Solution Process for

Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 118

6.5 Formation of Covalent Compounds . . . . . . . . . . . . 121

The Hydrogen Economy . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122

6.6 Thermodynamic versus Kinetic Factors . . . . . . . . . 123

vi C o n t e n t s

C H A P T E R 7 Solvent Systems and Acid-Base Behavior . . . . . . . . . . . . . 129

7.1 Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 130

7.2 Brnsted-Lowry Acids . . . . . . . . . . . . . . . . . . . . . . . 134

Antacids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 138

7.3 Brnsted-Lowry Bases . . . . . . . . . . . . . . . . . . . . . . . 139

Cyanide and Tropical Fish . . . . . . . . . . . . . . . . . . . . . . . . . . 140

7.4 Trends in Acid-Base Behavior . . . . . . . . . . . . . . . . . 141

7.5 Acid-Base Reactions of Oxides . . . . . . . . . . . . . . . . 144

Superacids and Superbases . . . . . . . . . . . . . . . . . . . . . . . . . 146

7.6 Lewis Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 147

7.7 Pearson Hard-Soft Acid-Base Concepts . . . . . . . . . 148

7.8 Applications of the HSAB Concept . . . . . . . . . . . . 150

7.9 Biological Aspects . . . . . . . . . . . . . . . . . . . . . . . . . . . 153

C H A P T E R 8 Oxidation and Reduction . . . . . . . . . . . . . . . . . . . . . . . . . . . 159

8.1 Redox Terminology . . . . . . . . . . . . . . . . . . . . . . . . . . 159

8.2 Oxidation Number Rules . . . . . . . . . . . . . . . . . . . . . 160

8.3 Determination of Oxidation Numbers from

Electronegativities . . . . . . . . . . . . . . . . . . . . . . . . . 161

8.4 The Difference between Oxidation Number

and Formal Charge . . . . . . . . . . . . . . . . . . . . . . . . 163

8.5 Periodic Variations of Oxidation Numbers . . . . . . . 163

8.6 Redox Equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . 164

Chemosynthesis: Redox Chemistry on the Seaoor . . . . . . 167

8.7 Quantitative Aspects of Half-Reactions . . . . . . . . 167

8.8 Electrode Potentials as Thermodynamic Functions 169

8.9 Latimer (Reduction Potential) Diagrams . . . . . . . . 170

8.10 Frost (Oxidation State) Diagrams . . . . . . . . . . . . . . 172

8.11 Pourbaix Diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . 173

Technetium: The Most Important Radiopharmaceutical . . 176

8.12 Redox Synthesis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177


C H A P T E R 9 Periodic Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 184

9.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 185

9.2 Periodic Trends in Bonding . . . . . . . . . . . . . . . . . . . . 188

9.3 Isoelectronic Series in Covalent Compounds . . . . . 192

9.4 Trends in Acid-Base Properties . . . . . . . . . . . . . . . . 192

9.5 The (n) Group and (n + 10) Group Similarities . . . 194

Chemical Topology . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 198

9.6 Isomorphism in Ionic Compounds . . . . . . . . . . . . . . 200

C o n t e n t s vii

New Materials: Beyond the Limitations of Geochemistry . 202

9.7 Diagonal Relationships . . . . . . . . . . . . . . . . . . . . . . . 202

Lithium and Mental Health . . . . . . . . . . . . . . . . . . . . . . . . . . 204

9.8 The Knights Move Relationship . . . . . . . . . . . . . 205

9.9 The Early Actinoid Relationships . . . . . . . . . . . . . . 208

9.10 The Lanthanoid Relationships . . . . . . . . . . . . . . . . . 208

9.11 Combo Elements . . . . . . . . . . . . . . . . . . . . . . . . . . 210

9.12 Pseudo-elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . 213


Thallium Poisoning: Two Case Histories . . . . . . . . . . . . . . . 217

C H A P T E R 1 0 Hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 221

10.1 Isotopes of Hydrogen . . . . . . . . . . . . . . . . . . . . . . . . 222

10.2 Nuclear Magnetic Resonance . . . . . . . . . . . . . . . . . . 223

Isotopes in Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 224

10.3 Properties of Hydrogen . . . . . . . . . . . . . . . . . . . . . . . 225

Searching the Depths of Space for the Trihydrogen Ion . . . 227

10.4 Hydrides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 228

10.5 Water and Hydrogen Bonding . . . . . . . . . . . . . . . . . 231

Water: The New Wonder Solvent . . . . . . . . . . . . . . . . . . . . . . 232

10.6 Clathrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 233

10.7 Biological Aspects of Hydrogen Bonding . . . . . . . . 235

Is There Life Elsewhere in Our Solar System? . . . . . . . . . . 236

10.8 Element Reaction Flowchart . . . . . . . . . . . . . . . . . . 237

C H A P T E R 1 1 The Group 1 Elements: The Alkali Metals . . . . . . . . . . . . . 240

11.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 241

11.2 Features of Alkali Metal Compounds . . . . . . . . . . . 242

11.3 Solubility of Alkali Metal Salts . . . . . . . . . . . . . . . . 243

Mono Lake . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 244

11.4 Lithium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 247

11.5 Sodium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 249

11.6 Potassium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 251

11.7 Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252

11.8 Hydroxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 253

11.9 Sodium Chloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . 255

Salt Substitutes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 255

11.10 Potassium Chloride . . . . . . . . . . . . . . . . . . . . . . . . . . 256

11.11 Sodium Carbonate . . . . . . . . . . . . . . . . . . . . . . . . . . . 257

11.12 Sodium Hydrogen Carbonate . . . . . . . . . . . . . . . . . 258

11.13 Ammonia Reaction . . . . . . . . . . . . . . . . . . . . . . . . . . 258

viii C o n t e n t s



C H A P T E R 1 2 The Group 2 Elements: The Alkaline Earth Metals . . . . . 263

12.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 263

12.2 Features of Alkaline Earth Metal Compounds . . . 265

12.3 Solubility of Alkaline Earth Metal Salts . . . . . . . . . 265

12.4 Beryllium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 267

12.5 Magnesium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 268

12.6 Calcium and Barium . . . . . . . . . . . . . . . . . . . . . . . . . 270

12.7 Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 271

12.8 Calcium Carbonate . . . . . . . . . . . . . . . . . . . . . . . . . . 272

How Was Dolomite Formed? . . . . . . . . . . . . . . . . . . . . . . . . 273

12.9 Cement . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 274

12.10 Calcium Chloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . 275

12.11 Calcium Sulfate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 277

Biomineralization: A New Interdisciplinary Frontier . . . 276

12.12 Calcium Carbide . . . . . . . . . . . . . . . . . . . . . . . . . . . . 277



C H A P T E R 1 3 The Group 13 Elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . 282

13.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 282

13.2 Boron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 284

13.3 Borides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 285

Inorganic Fibers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 286

13.4 Boranes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 286

Boron Neutron Capture Therapy . . . . . . . . . . . . . . . . . . . . . 289

13.5 Boron Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 291

13.6 Aluminum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 292

13.7 Aluminum Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . 297

13.8 Aluminum Potassium Sulfate . . . . . . . . . . . . . . . . . . 298

13.9 Spinels . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 298




14.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 306

14.2 Carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 307

The Discovery of Buckminsterfullerene . . . . . . . . . . . . . . . . 310

14.3 Isotopes of Carbon . . . . . . . . . . . . . . . . . . . . . . . . . . 313

C o n t e n t s ix

14.4 The Extensive Chemistry of Carbon . . . . . . . . . . . . 314

14.5 Carbides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 314

Moissanite: The Diamond Substitute . . . . . . . . . . . . . . . . . . 316

14.6 Carbon Monoxide . . . . . . . . . . . . . . . . . . . . . . . . . . . 317

14.7 Carbon Dioxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 319

Carbon Dioxide, Supercritical Fluid . . . . . . . . . . . . . . . . . . 321

14.8 Carbonates and Hydrogen Carbonates . . . . . . . . . . 321

14.9 Carbon Suldes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 323

14.10 Carbon Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 324

14.11 Methane . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 326

14.12 Cyanides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 327

14.13 Silicon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 328

14.14 Silicon Dioxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 329

14.15 Silicates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 332

14.16 Aluminosilicates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 334

14.17 Silicones . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 337

Inorganic Polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 338

14.18 Tin and Lead . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 339

14.19 Tin and Lead Oxides . . . . . . . . . . . . . . . . . . . . . . . . . 340

14.20 Tin and Lead Halides . . . . . . . . . . . . . . . . . . . . . . . . 341

14.21 Tetraethyllead . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342

A Case History . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 343




15.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 352

15.2 Anomalous Nature of Nitrogen . . . . . . . . . . . . . . . . 353

15.3 Nitrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 355

Propellants and Explosives . . . . . . . . . . . . . . . . . . . . . . . . . . 356

15.4 Overview of Nitrogen Chemistry . . . . . . . . . . . . . . . 357

The First Dinitrogen Compound . . . . . . . . . . . . . . . . . . . . . 358

15.5 Nitrogen Hydrides . . . . . . . . . . . . . . . . . . . . . . . . . . . 359

Haber and Scientic Morality . . . . . . . . . . . . . . . . . . . . . . . . 361

15.6 Nitrogen Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 365

15.7 The Ammonium Ion . . . . . . . . . . . . . . . . . . . . . . . . . 366

15.8 Nitrogen Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 367

15.9 Nitrogen Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372

15.10 Nitrous Acid and Nitrites . . . . . . . . . . . . . . . . . . . . . 373

15.11 Nitric Acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 374

15.12 Nitrates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 375

x C o n t e n t s

15.13 Overview of Phosphorus Chemistry . . . . . . . . . . . . 377

15.14 Phosphorus and Its Allotropes . . . . . . . . . . . . . . . . . 378

Nauru, the Worlds Richest Island . . . . . . . . . . . . . . . . . . . . . 380

15.15 Phosphine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 381

15.16 Phosphorus Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . 382

15.17 Phosphorus Chlorides . . . . . . . . . . . . . . . . . . . . . . . . 382

15.18 Common Oxyacids of Phosphorus . . . . . . . . . . . . . . 384

15.19 Phosphates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 385


Paul Erhlich and His Magic Bullet . . . . . . . . . . . . . . . . . . 389



16.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 398

16.2 Anomalous Nature of Oxygen . . . . . . . . . . . . . . . . . 399

Oxygen Isotopes in Geology . . . . . . . . . . . . . . . . . . . . . . . . . 399

16.3 Oxygen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 400

16.4 Bonding in Covalent Oxygen Compounds . . . . . . . 406

16.5 Trends in Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . 407

16.6 Mixed Metal Oxides . . . . . . . . . . . . . . . . . . . . . . . . . 408

New Pigments through Perovskites . . . . . . . . . . . . . . . . . . . 409

16.7 Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 410

16.8 Hydrogen Peroxide . . . . . . . . . . . . . . . . . . . . . . . . . . 411

16.9 Hydroxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 412

16.10 The Hydroxyl Radical . . . . . . . . . . . . . . . . . . . . . . . . 414

16.11 An Overview of Sulfur Chemistry . . . . . . . . . . . . . . 414

16.12 Sulfur and Its Allotropes . . . . . . . . . . . . . . . . . . . . . . 415

Cosmochemistry: Io, the Sulfur-Rich Moon . . . . . . . . . . . . . 416

16.13 Hydrogen Sulde . . . . . . . . . . . . . . . . . . . . . . . . . . . . 418

16.14 Suldes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 419

Disulde Bonds and Hair . . . . . . . . . . . . . . . . . . . . . . . . . . . 420

16.15 Sulfur Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 421

16.16 Sultes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424

16.17 Sulfuric Acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 425

16.18 Sulfates and Hydrogen Sulfates . . . . . . . . . . . . . . . . 427

16.19 Other Oxy-Sulfur Anions . . . . . . . . . . . . . . . . . . . . . 428

16.20 Sulfur Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 430

16.21 Sulfur-Nitrogen Compounds . . . . . . . . . . . . . . . . . . 432

16.22 Selenium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 432



C o n t e n t s xi

C H A P T E R 1 7 The Group 17 Elements: The Halogens . . . . . . . . . . . . . . . 441

17.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 442

17.2 Anomalous Nature of Fluorine . . . . . . . . . . . . . . . . 444

17.3 Fluorine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 445

The Fluoridation of Water . . . . . . . . . . . . . . . . . . . . . . . . . . . 446

17.4 Hydrogen Fluoride and Hydrouoric Acid . . . . . . 448

17.5 Chlorine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 449

17.6 Hydrochloric Acid . . . . . . . . . . . . . . . . . . . . . . . . . . . 451

17.7 Halides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 452

17.8 Chlorine Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 456

17.9 Chlorine Oxyacids and Oxyanions . . . . . . . . . . . . . 458

Swimming Pool Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . 460

The Discovery of the Perbromate Ion . . . . . . . . . . . . . . . . . 462

17.10 Interhalogen Compounds and Polyhalide Ions . . . 463



C H A P T E R 1 8 The Group 18 Elements: The Noble Gases . . . . . . . . . . . . 473

18.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 474

18.2 Unique Features of Helium . . . . . . . . . . . . . . . . . . . 475

18.3 Uses of the Noble Gases . . . . . . . . . . . . . . . . . . . . . . 475

18.4 A Brief History of Noble Gas Compounds . . . . . . 476

Is It Possible to Make Compounds of the Early

Noble Gases? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 477

18.5 Xenon Fluorides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 478

18.6 Xenon Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 479



C H A P T E R 1 9 Introduction to Transition Metal Complexes . . . . . . . . . . . 484

19.1 Transition Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . 484

19.2 Transition Metal Complexes . . . . . . . . . . . . . . . . . . . 485

19.3 Stereochemistries . . . . . . . . . . . . . . . . . . . . . . . . . . . . 487

19.4 Isomerism in Transition Metal Complexes . . . . . . . 488

Platinum Complexes and Cancer Treatment . . . . . . . . . . . . 491

19.5 Naming Transition Metal Complexes . . . . . . . . . . . 492

19.6 An Overview of Bonding Theories of

Transition Metal Compounds . . . . . . . . . . . . . . . . 494

19.7 Crystal Field Theory . . . . . . . . . . . . . . . . . . . . . . . . . 496

19.8 Successes of Crystal Field Theory . . . . . . . . . . . . . . 501

The Earth and Crystal Structures . . . . . . . . . . . . . . . . . . . . . 504

xii C o n t e n t s

19.9 More on Electronic Spectra . . . . . . . . . . . . . . . . . . . 504

19.10 Thermodynamic versus Kinetic Factors . . . . . . . . . 506

19.11 Synthesis of Coordination Compounds . . . . . . . . . . 506

19.12 Coordination Complexes and the HSAB

Concept . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 508


C H A P T E R 2 0 Properties of Transition Metals . . . . . . . . . . . . . . . . . . . . . . 513

20.1 Overview of the Transition Metals . . . . . . . . . . . . . . 514

20.2 Group 4: Titanium, Zirconium, and Hafnium . . . . . 517

20.3 Group 5: Vanadium, Niobium, and Tantalum . . . . . 518

20.4 Group 6: Chromium, Molybdenum, and Tungsten 519

20.5 Group 7: Manganese, Technetium, and Rhenium . 528

Mining the Seaoor . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 529

20.6 Group 8: Iron, Ruthenium, and Osmium . . . . . . . . 533

20.7 Group 9: Cobalt, Rhodium, and Iridium . . . . . . . . . 544

20.8 Group 10: Nickel, Palladium, and Platinum . . . . . . 547

20.9 Group 11: Copper, Silver, and Gold . . . . . . . . . . . . . 549



21.1 Group Trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 567

21.2 Zinc and Cadmium . . . . . . . . . . . . . . . . . . . . . . . . . . 567

21.3 Mercury . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 570

Mercury Amalgam in Teeth . . . . . . . . . . . . . . . . . . . . . . . . . . 574



C H A P T E R 2 2 Organometallic Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . 577

22.1 Naming Organometallic Compounds . . . . . . . . . . . 578

22.2 Counting Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . 579

22.3 Solvents for Organometallic Chemistry . . . . . . . . . 580

22.4 Main Group Organometallic Compounds . . . . . . . 581

Grignard Reagents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 583

The Death of Karen Wetterhahn . . . . . . . . . . . . . . . . . . . . . . 588

22.5 Organometallic Compounds of the Transition

Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 589

22.6 Transition Metal Carbonyls . . . . . . . . . . . . . . . . . . . 591

22.7 Synthesis and Properties of Simple Metal

Caubonyls . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 595

22.8 Reactions of Transition Metal Carbonyls . . . . . . . . 597

C o n t e n t s xiii

22.9 Other Carbonyl Compounds . . . . . . . . . . . . . . . . . . 599

22.10 Complexes with Phosphine Ligands . . . . . . . . . . . . 600

22.11 Complexes with Alkyl, Alkene, and Alkyne

Ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 601

The Preservation of Books . . . . . . . . . . . . . . . . . . . . . . . . . . 603

22.12 Complexes with Allyl and 1,3-Butadiene Ligands . 604

22.13 Metallocenes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 606

22.14 Complexes with 6-Arene Ligands . . . . . . . . . . . . . 607

22.15 Complexes with Cycloheptatriene and

Cyclooctatetraene Ligands . . . . . . . . . . . . . . . . . . 608

22.16 Fluxionality . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 609

22.17 Organometallic Compounds in Industrial

Catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 609

C H A P T E R 2 3 The Rare Earth and Actinoid Elements . . . . . . . . . . . . . 617w

23.1 Properties of the Rare Earth Elements . . . . . . . . 619w

Superconductivity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 621w

23.2 Properties of the Actinoids . . . . . . . . . . . . . . . . . . 622w

23.3 Extraction of Uranium . . . . . . . . . . . . . . . . . . . . . . 625w

A Natural Fission Reactor . . . . . . . . . . . . . . . . . . . . . . . . . 625w

23.4 Enriched and Depleted Uranium . . . . . . . . . . . . . 627w

23.5 The Postactinoid Elements . . . . . . . . . . . . . . . . . . 628w

APPENDICES

Appendix 1 Thermodynamic Properties of Some

Selected Inorganic Compounds.............. A-1

Appendix 2 Charge Densities of Selected Ions ........... A-13

Appendix 3 Selected Bond Energies ............................ A-16

Appendix 4 Ionization Energies of Selected Metals .. A-18

Appendix 5 Electron Afnities of Selected Nonmetals. A-20

Appendix 6 Selected Lattice Energies ......................... A-21

Appendix 7 Selected Hydration Enthalpies................. A-22

Appendix 8 Selected Ionic Radii ................................... A-23

Appendix 9 Standard Half-Cell Electrode Potentials

of Selected ElementsWeb .............. A-25w

Appendix 10 Electron Conguration of the

ElementsWeb................................... A-35w

Index .......................................................................... I-1




What Is Descriptive Inorganic Chemistry?

Descriptive inorganic chemistry was traditionally concerned with the prop-erties of the elements and their compounds. Now, in the renaissance of thesubject, the properties are being linked with explanations for the formulasand structures of compounds together with an understanding of the chemi-cal reactions they undergo. In addition, we are no longer looking at inor-ganic chemistry as an isolated subject but as a part of essential scientificknowledge with applications throughout science and our lives. And it isbecause of a need for greater contextualization that we have added morefeatures and more applications.

In many colleges and universities, descriptive inorganic chemistry isoffered as a sophomore or junior course. In this way, students come to knowsomething of the fundamental properties of important and interesting ele-ments and their compounds. Such knowledge is important for careers notonly in pure or applied chemistry but also in pharmacy, medicine, geology,environmental science, and so on. This course can then be followed by ajunior or senior course that focuses on the theoretical principles and the useof spectroscopy to a greater depth than is covered in a descriptive text. Infact, the theoretical course builds nicely on the descriptive background.Without the descriptive grounding, however, the theory becomes sterile,uninteresting, and irrelevant.

Education has often been a case of the swinging pendulum, and thishas been very true of inorganic chemistry. Up until the 1960s, it was verymuch pure descriptive, requiring exclusively memorization. In the 1970sand 1980s, upper-level texts focused exclusively on the theoretical princi-ples. Now it is apparent that descriptive is very importantnot the tradi-tional memorization of facts, but the linking of facts, where possible, tounderlying principles. Students need to have modern descriptive inorganicchemistry as part of their education. Thus, we must ensure that chemists areaware of the new descriptive inorganic chemistry.

Preface

Inorganic chemistry goes beyond academic interest: it is an important partof our lives.

Inorganic chemistry is interestingmore than thatit is exciting! So muchof our twenty-first-century science and technology relies on natural andsynthetic materials, often inorganic compounds. Inorganic chemistry isubiquitous in our daily lives: household products, some pharmaceuticals,our transportationboth the vehicles themselves and the synthesis of thefuelsbattery technology, and medical treatments. There is the industrialaspect, the production of all the chemicals that are required to drive oureconomy, everything from steel to sulfuric acid to glass and cement. Envi-ronmental chemistry is largely a question of the inorganic chemistry of theatmosphere, water, and soil. Finally, there are the profound issues of theinorganic chemistry of our planet, the solar system, and the universe.

This text is designed to focus on the properties of selected interesting,important, and unusual elements and compounds. However, to understandinorganic chemistry, it is crucial to tie this knowledge to the underlyingchemical principles and hence provide explanations for the existence andbehavior of compounds. For this reason, almost half the chapters survey therelevant concepts of atomic theory, bonding, intermolecular forces, thermo-dynamics, acid-base behavior, and reduction-oxidation properties as a prel-ude to, and preparation for, the descriptive material.

For this fourth edition, the major improvements are as follows:

Chapter 3: Covalent BondingBy popular request, we have added THREE NEW SECTIONS related to sym-metry: Molecular Symmetry, Symmetry and Vibrational Spectroscopy, andRadiation Trapping: The Greenhouse Effect.

There is now a CONSOLIDATED SECTION, Intermolecular Forces.

Chapter 5: Ionic BondingThe latter part of the chapter has been reworked. In particular, there is aREORGANIZED SECTION, Ionic Crystal Structures.The REVISED SECTION, The Bonding Continuum, includes discussion ofboth the bond triangle and the bond tetrahedron.

Chapter 6: Inorganic ThermodynamicsThis chapter begins with a REMODELED SECTION, Thermodynamics of theFormation of Compounds, to make it a smoother introduction to the topic.

xvi P r e f a c e

Chapter 7: Solvent Systems and Acid-Base BehaviorThis renamed chapter begins with a NEW SECTION, Solvents, on the typesand properties of solvents for inorganic reactions, and we excised materialthat excessively overlapped with general chemistry courses.A brief discussion of Lux-Flood theory was also added to the chapter.

Chapter 8: Oxidation and ReductionA NEW SECTION, Redox Synthesis, was added.

Chapter 10: HydrogenAn EXPANDED SECTION, Clathrates, shows the increasing importance of gashydrates.

Chapters 1116Each of these chapters has REORGANIZED CONTENT so that there are fewersections, at the same time providing a better ow (for example, Chapter 15has been reduced from 28 sections to 21 sections).

Chapter 13: The Group 13 ElementsAn EXPANDED SECTION, Boranes, gives more comprehensive coverage.

Chapters 1417In our discussions of these four groups, we have put much greater emphasison the use of bond energy data to offer explanations of element and com-pound behavior.

Chapter 20: Properties of Transition MetalsThe RESTRUCTURED SECTIONS Group 6: Chromium, Molybdenum, andTungsten and Group 7: Manganese,Technetium, and Rhenium provide a bet-ter ow and linkage between topics.

Chapter 23: The Rare Earth and Actinoid ElementsThe UPDATED SECTION, The Postactinoid Elements, reects recent elementdiscoveries.

Video ClipsDescriptive inorganic chemistry by denition is visual. What better way toappreciate a chemical reaction than to make it visual? So, with this edition,we have A SERIES OF AT LEAST 50 WEB-BASED VIDEO CLIPS to bring someof the reactions to life. The text has the margin icon to indicate where areaction is illustrated.

Key IdeasEach chapter now concludes with a checklist of KEY IDEAS in thatchapter.

VI

DEO CLIP

P r e f a c e xvii

FeaturesWe have added one ADDITIONAL FEATURE, Chemical Topology, whichshows the use of computer technology to nd new linkages among chemicalelements. Other features have been updated.

Text Figures and TablesAll of the text figures and tables are available as .jpg files for inclusionin PowerPoint presentations on the instructor side of the Web site atwww.whfreeman.com/rayner4e.

Additional ResourcesA list of READABLE, RELEVANT LITERATURE SOURCES is found on the textWeb site at www.whfreeman.com/rayner4e.

SupplementsThe Student Solutions Manual, 0-7167-6177-7, includes the worked solu-tions to all of the odd-numbered problems found in the text.

Web Site www.whfreeman.com/rayner4e, features New visual demonstrations of inorganic chemistry (videos) All art from the text (instructor side) Lab experiments Additional Web resources and study tools Appendices Video clips

Instructors Resource CD-ROM, 0-7167-6176-9Includes all text art, PowerPoint, videos, and solutions to all problems fromthe text.

This book was written to pass on to another generation our fascination withdescriptive inorganic chemistry. Thus, the comments of readers, both stu-dents and instructors, will be sincerely appreciated. Any suggestions foradded or updated additional readings are also welcome. Our current e-mailaddresses are [email protected] and [email protected].

Acknowledgments

Many thanks must go to the team at W. H. Freeman who have contributedtheir talents to the four editions of this book.We offer our sincere gratitudeto our editors of the fourth edition, Jessica Fiorillo and Jenness Crawford(and to Mary Louise Byrd, who rejoined the project); of the third edition,Jessica Fiorillo and Guy Copes; of the second edition, Michelle Julet andMary Louise Byrd; and a special thanks to Deborah Allen, who bravelycommissioned the rst edition of the text. Each one of our fabulous editorshas been a source of encouragement, support, and helpfulness.

We wish to acknowledge the following reviewers of this edition, whosecriticisms and comments were much appreciated: Rachel Narehood Austinat Bates College; Leo A. Bares at the University of North CarolinaAsheville; Karen S. Brewer at Hamilton College; Robert M. Burns at AlmaCollege; Do Chang at Averett University; Georges Dns at ConcordiaUniversity; Daniel R. Derringer at Hollins University; Carl P. Fictorie atDordt College; Margaret Kastner at Bucknell University; Michael Laing atthe University of Natal, Durban; Richard H. Langley at Stephen F. AustinState University; Mark R. McClure at the University of North Carolina atPembroke; Louis Mercier at Laurentian University; G. Merga at AndrewsUniversity; Stacy OReilly at Butler University; Larry D. Pedersen atCollege Misercordia; Robert D. Pike at the College of William and Mary;William Quintana at New Mexico State University; David F. Rieck at Salis-bury University; John Selegue at the University of Kentucky; Melissa M.Strait at Alma College; Daniel J. Williams at Kennesaw State University;Juchao Yan at Eastern New Mexico University; and Arden P. Zipp at theState University of New York at Cortland.

We acknowledge with thanks the contributions of the reviewers of thethird edition: Franois Caron at Laurentian University; Thomas D. Getmanat Northern Michigan University; Janet R. Morrow at the State Universityof New York at Buffalo; Robert D. Pike at the College of William and Mary;Michael B. Wells at Cambell University; and particularly Joe Takats of theUniversity of Alberta for his comprehensive critique of the second edition.

And the contributions of the reviewers of the second edition: F. C. Hentzat North Carolina State University; Michael D. Johnson at New MexicoState University; Richard B. Kaner at the University of California, LosAngeles; Richard H. Langley at Stephen F. Austin State University; JamesM. Mayer at the University of Washington; Jon Melton at Messiah College;Joseph S. Merola at Virginia Technical Institute; David Phillips at WabashCollege; John R. Pladziewicz at the University of Wisconsin, Eau Claire;Daniel Rabinovich at the University of North Carolina at Charlotte; DavidF. Reich at Salisbury State University;Todd K.Trout at Mercyhurst College;

A c k n o w l e d g m e n t s xix

Steve Watton at the Virginia Commonwealth University; and John S. Woodat the University of Massachusetts, Amherst.

Likewise, the reviewers of the first edition: E. Joseph Billo at BostonCollege; David Finster at Wittenberg University; Stephen J. Hawkes atOregon State University; Martin Hocking at the University of Victoria;Vake Marganian at Bridgewater State College; Edward Mottel at theRose-Hulman Institute of Technology; and Alex Whitla at Mount AllisonUniversity.

As a personal acknowledgment, Geoff Rayner-Canham wishes to espe-cially thank three teachers and mentors who had a major influence on hiscareer: Briant Bourne, Harvey Grammar School; Margaret Goodgame,Imperial College, London University; and Derek Sutton, Simon Fraser Uni-versity. And he expresses his eternal gratitude to his spouse, Marelene, forher support and encouragement.

Tina Overton would like to thank her colleague Phil King for his invalu-able suggestions for improvements and his assistance with the illustrations.Thanks must also go to her family, Dave, John, and Lucy, for their patienceduring the months when this project lled all her waking hours.

Dedication

Chemistry is a human endeavor. New discoveries are the result of the workof enthusiastic individuals and groups of individuals who want to explorethe molecular world. We hope that you, the reader, will come to share ourown fascination with inorganic chemistry. We have chosen to dedicate thisbook to two people who, for very different reasons, never did receive theultimate accolade of a Nobel Prize.

Henry Moseley (18871915)

Although Mendeleev is identified as the discoverer of the periodic table,his version was based on an increase in atomic mass. In some cases, theorder of elements had to be reversed to match properties with location. Itwas a British scientist, Henry Moseley, who put the periodic table on amuch firmer footingby discovering that,on bombardmentwith electrons, eachelement emitted X-rays of characteris-tic wavelengths. Thewavelengths fitted aformula related byan integer numberunique to each ele-ment. We know thatnumber to be thenumber of protons.With the establish-ment of the atomicnumber of an ele-ment, chemists at last knew the fundamental organization of the table.Sadly, Moseley was killed at the battle of Gallipoli in World War I. Thus,one of the brightest scientific talents of the twentieth century died at theage of 27. The famous American scientist Robert Milliken commented:Had the European War had no other result than the snuffing out of thisyoung life, that alone would make it one of the most hideous and mostirreparable crimes in history. Unfortunately, Nobel Prizes are onlyawarded to living scientists. In 1924, the discovery of element 43 wasclaimed, and it was named moseleyum; however, the claim was disprovedby the very method that Moseley had pioneered.

D e d i c a t i o n xxi

Lise Meitner (18781968)

In the 1930s, scientists werebombarding atoms of heavyelements such as uranium withsubatomic particles to try tomake new elements andextend the periodic table. Aus-trian scientist Lise Meitnerhad shared leadership withOtto Hahn of the Germanresearch team working on thesynthesis of new elements.They thought they had discov-ered nine new elements.Shortly after the claimed dis-covery, Meitner was forced toflee Germany because of herJewish ancestry, and she set-tled in Sweden. Hahn reportedto her that one of the new elements behaved chemically just like barium.During a famous walk in the snow with her nephew, physicist Otto Frisch,Meitner realized that an atomic nucleus could break in two just like a dropof water. No wonder the element formed behaved like barium: it was bar-ium! Thus was born the concept of nuclear fission. She informed Hahn ofher proposal. When Hahn wrote the research paper on the work, he barelymentioned the vital contribution of Meitner and Frisch. As a result, Hahnand his colleague Fritz Strassmann received the Noble Prize. Meitners ashof genius was ignored. Only recently has Meitner received the acclaim shedeserved by the naming of an element after her, element 109, meitnerium.

Additional ReadingHeibron, J. L. H. G. J. Moseley. University of California Press, Berkeley,1974.Rayner-Canham, M. F., and G. W. Rayner-Canham. Women in Chemistry:Their Changing Roles from Alchemical Times to the Mid-TwentiethCentury. Chemical Heritage Foundation, Philadelphia, 1998.Sime, R. L. Lise Meitner: A Life in Physics. University of California Press,Berkeley, 1996.Weeks, M. E., and H. M. Leicester. Discovery of the Elements, 7th ed.Journal of Chemical Education, Easton, PA, 1968.

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1C H A P T E R 1

The Electronic Structure of the Atom:A Review

To understand the behavior of inorganic compounds, we need to study the nature of chemical bonding. Bonding, in turn, relates to the behaviorof electrons in the constituent atoms. Our study of inorganic chemistry,therefore, starts with a review of the probability model of the atom and asurvey of the models applications to the electron congurations of atomsand ions.

It is amazing that Isaac Newton discovered anything at all, for he was theoriginal model for the absentminded professor. Supposedly, he alwaystimed the boiled egg he ate at breakfast; one morning, his maid found himstanding by the pot of boiling water, holding an egg in his hand and gazingintently at the watch in the bottom of the pot! Nevertheless, he initiated thestudy of the electronic structure of the atom in about 1700, when he noticedthat the passage of sunlight through a prism produced a continuous visiblespectrum (Figure 1.1).

Much later, in 1860, Robert Bunsen (of burner fame) investigated thelight emissions from flames and gases. Bunsen observed that the emissionspectra, rather than being continuous, were series of colored lines (linespectra). He noted that each chemical element produced a unique and char-acteristic spectrum (Figure 1.2). Other investigators subsequently showedthat there were, in fact, several sets of spectral lines for the hydrogen atom:one set in the ultraviolet region, one in the visible region, and a number ofsets in the infrared part of the electromagnetic spectrum (Figure 1.3).

The explanation of the spectral lines was one of the triumphs of the Bohrmodel of the atom. In 1913, Niels Bohr proposed that the electrons couldoccupy only certain energy levels that corresponded to various circularorbits around the nucleus. He identied each of these levels with an integer,

1.1 The Schrdinger WaveEquation and Its Signicance

Atomic AbsorptionSpectroscopy

1.2 Shapes of the AtomicOrbitals

1.3 The Polyelectronic Atom

1.4 Ion Electron Congurations

1.5 Magnetic Properties ofAtoms

FIGURE 1.1 A prism splits whitelight into the wavelengths of thevisible spectrum.

Sun Slit

Prism

RedOrangeYellowGreenBlueIndigoViolet

Continuousspectrum

2 C H A P T E R 1 The Electronic Structure of the Atom: A Review

which he called a quantum number.The value of this parameter could rangefrom 1 to . He argued that, as energy was absorbed by an atom from aflame or electrical discharge, electrons moved from one quantum level toone or more higher energy levels. The electrons sooner or later returned tolower quantum levels, closer to the nucleus, and when they did, light wasemitted. The wavelength of the emitted light directly corresponded to theenergy of separation of the initial and nal quantum levels. When the elec-trons occupied the lowest-possible energy level, they were said to be in theground state. If one or more electrons absorbed enough energy to moveaway from the nucleus, then they were said to be in an excited state.

The energy of an electron in each level could be found from the relationship

where E is the electron energy, n is the quantum number, and RH is theRydberg constant for hydrogen. The energy of the light emitted could becalculated from the difference in energies of the initial and final energylevels. The wavelength of the light could then be found from the relation-ships and where h is Plancks constant, is the frequency,c is the velocity of light, and is the wavelength of the emitted light.

However, the Bohr model had a number of aws. For example, the spec-tra of multi-electron atoms had far more lines than the simple Bohr modelpredicted. Nor could the Bohr model explain the splitting of the spectrallines in a magnetic field (a phenomenon known as the Zeeman effect).Within a short time, a radically different model, the quantum mechanicalmodel, was proposed to account for these observations.

c ,E h

E RHa 1n2b

Ultra-

violet

Visible

Infrared

200 nm

500 nm

1000 nm2000 nm

100 nm

FIGURE 1.3 Emission spectrum ofhydrogen.

Slit

PrismLinespectrum

Sample

Burner

FIGURE 1.2 A line spectrum isproduced when an element is heatedin a ame.

1.1 The Schrdinger Wave Equation and Its Signicance 3

1.1 The Schrdinger Wave Equation and Its Signicance

The more sophisticated quantum mechanical model of atomic structure wasderived from the work of Louis de Broglie. De Broglie showed that, just aselectromagnetic waves could be treated as streams of particles (photons),moving particles could exhibit wavelike properties. Thus, it was equallyvalid to picture electrons either as particles or as waves. Using this wave-particle duality, Erwin Schrdinger developed a partial differential equa-tion to represent the behavior of an electron around an atomic nucleus. Thisequation, given here for a one-electron atom, shows the relationship betweenthe wave function of the electron, and E and V, the total and potentialenergies of the system, respectively.The second differential terms representthe wave function along each of the Cartesian coordinates x, y, and z; m isthe mass of an electron, and h is Plancks constant.

00x2

00y2

00z2

82m

b21E V2 0

,

Atomic Absorption Spectroscopy

A glowing body, such as the Sun, is expected to emit a continuous spectrum ofelectromagnetic radiation. However, in the early nineteenth century, a Germanscientist, Josef von Fraunhofer, noticed that the visible spectrum from the Sunactually contained a number of dark bands. Later investigators realized that thebands were the result of the absorption of particular wavelengths by cooleratoms in the atmosphere above the surface of the Sun. The electrons of theseatoms were in the ground state, and they were absorbing radiation at wavelengthscorresponding to the energies needed to excite them to higher energy states. Astudy of these negative spectra led to the discovery of helium. Such spectralstudies are still of great importance in cosmochemistrythe study of thechemical composition of stars.

In 1955, two groups of scientists, one in Australia and the other in Holland,nally realized that the absorption method could be used to detect the presenceof elements at very low concentrations. Each element has a particular absorp-tion spectrum corresponding to the various separations of (differencesbetween) the energy levels in its atoms. When light from a powerful source ispassed through a vaporized sample of an element, the particular wavelengthscorresponding to the various energy separations will be absorbed. We nd thatthe higher the concentration of the atoms, the greater the proportion of thelight that will be absorbed. This linear relationship between light absorption andconcentration is known as Beers law. The sensitivity of this method is extremelyhigh, and concentrations of parts per million are easy to determine; some ele-ments can be detected at the parts per billion level. Atomic absorption spec-troscopy has now become a routine analytical tool in chemistry, metallurgy,geology, medicine, forensic science, and many other elds of scienceand itsimply requires the movement of electrons from one energy level to another.


The derivation of this equation and the method of solving it are in therealm of physics and physical chemistry, but the solution itself is of greatimportance to inorganic chemists.We should always keep in mind, however,that the wave equation is simply a mathematical formula. We attach mean-ings to the solution simply because most people need concrete images tothink about subatomic phenomena.The images that we create correspondingto our macroscopic world can only vaguely resemble the subatomic reality.

Schrdinger argued that the real meaning of the equation could befound from the square of the wave function, which represents the prob-ability of finding the electron at any point in the region surrounding thenucleus. There are a number of solutions to a wave equation. Each solutiondescribes a different orbital and, hence, a different probability distributionfor an electron in that orbital. Each of these orbitals is uniquely dened bya set of three integers: n, l, and ml. Like the integers in the Bohr model,these integers are also called quantum numbers.

In addition to the three quantum numbers derived from the original the-ory, a fourth quantum number had to be dened to explain the results of alater experiment. In this experiment, it was found that passing a beam ofhydrogen atoms through a magnetic eld caused about half the atoms to bedeflected in one direction and the other half in the opposite direction.Other investigators proposed that the observation was the result of two dif-ferent electronic spin orientations. The atoms possessing an electron withone spin were deflected one way, and the atoms whose electron had theopposite spin were deflected in the opposite direction. This spin quantumnumber was assigned the symbol ms.

The possible values of the quantum numbers are dened as follows:

n, the principal quantum number, can have all positive integer valuesfrom 1 to .

l, the angular momentum quantum number, can have all integer valuesfrom n 1 to 0.

ml, the magnetic quantum number, can have all integer values from 1through 0 to 1.

ms, the spin quantum number, can have values of 12 and 12.

When the value of the principal quantum number is 1, there is only onepossible set of quantum numbers n, l, and ml (1, 0, 0), whereas for a principalquantum number of 2, there are four sets of quantum numbers (2, 0, 0; 2, 1,1; 2, 1, 0; 2, 1, 1). This situation is shown diagrammatically in Figure 1.4.To identify the electron orbital that corresponds to each set of quantumnumbers, we use the value of the principal quantum number n, followed bya letter for the angular momentum quantum number l. Thus, when n 1,there is only the 1s orbital.

When n 2, there is one 2s orbital and three 2p orbitals (correspondingto the ml values of 1, 0, and 1). The letters s, p, d, and f are derived fromcategories of the spectral lines: sharp, principal, diffuse, and fundamental.The correspondences are shown in Table 1.1.

2,

TABLE 1.1 Correspondencebetween angular momentumnumber l and orbital designation

l Value Orbital designation

0 s

1 p

2 d

3 f

1.1 The Schrdinger Wave Equation and Its Signicance 5

When the principal quantum number n 3, there are nine sets of quan-tum numbers (Figure 1.5).These sets correspond to one 3s, three 3p, and ve3d orbitals.A similar diagram for the principal quantum number n 4 wouldshow 16 sets of quantum numbers, corresponding to one 4s, three 4p, ve 4d,and seven 4f orbitals (Table 1.2). Theoretically, we can go on and on, but aswe shall see, the f orbitals represent the limit of orbital types among the ele-ments of the periodic table for atoms in their electronic ground states.

The Schrdinger wave equation is usually presented as the definitiverepresentation of the electrons of an atom, but it is not. As we discuss inChapter 2, Section 2.5, the equation fails to take into account the fact thatsome of the electrons in the more massive elements are traveling atextremely high velocities. As a result, the electron masses are inuenced byrelativistic effects. Although the Schrdinger equation can be modified toaccount for the problem, in 1928, the English physicist P.A.M. Dirac devel-oped a better wave equation that integrates relativity factors. The Diracequation provides four quantum numbers directly, although only the princi-pal quantum number, n, has the same signicance in both Schrdinger andDirac equations. Even the shapes of the orbitals derived from the Dirac equa-tion are different from those derived from the Schrdinger equation. How-ever, because this is a descriptive chemistry text, we emphasize the features ofthe simpler and more commonly used Schrdinger-derived orbitals.

n

l

ml

0 1

3

0 11 2 1 1 20 0

2

3s 3p 3dFIGURE 1.5 The possible sets ofquantum numbers for n 3.

n

l

ml

1s 2s 2p

1 100

0 1

1

0

0

2

FIGURE 1.4 The possible sets ofquantum numbers for n 1 and n 2.

TABLE 1.2 Correspondencebetween angular momentumnumber l and number of orbitals

l Value Number of orbitals

0 1

1 3

2 5

3 7


1.2 Shapes of the Atomic Orbitals

Representing the solutions to a wave equation on paper is not an easy task.In fact, we would need four-dimensional graph paper (if it existed) to dis-play the complete solution for each orbital. As a realistic alternative, webreak the wave equation into two parts: a radial part and an angular part.

Each of the three quantum numbers derived from the wave equationrepresents a different aspect of the orbital:

The principal quantum number n indicates the size of the orbital.

The angular momentum quantum number l represents the shape of theorbital.

The magnetic quantum number ml represents the spatial direction ofthe orbital.

The spin quantum number ms has little physical meaning; it merelyallows two electrons to occupy the same orbital.

The value of the principal quantum number and, to a lesser extent, that ofthe angular momentum quantum number determine the energy of the elec-tron. Although the electron may not literally be spinning, it behaves as if itwere, and it has the magnetic properties expected for a spinning particle.

An orbital diagram is used to indicate the probability of nding an elec-tron at any instant at any location.An alternative viewpoint is to consider thelocations of an electron over a lengthy period of time. We dene a locationwhere an electron seems to spend most of its time as an area of high electrondensity. Conversely, locations rarely visited by an electron are called areas oflow electron density.

The s orbitals are spherically symmetric about the atomic nucleus. As theprincipal quantum number increases, the electron tends to be found fartherfrom the nucleus. To express this idea in a different way, we say that, as the

principal quantum number increases, the orbital becomesmore diffuse. A unique feature of electron behavior in ans orbital is that there is a nite probability of nding the elec-tron close to, and even within, the nucleus.This penetration bys orbital electrons plays a role in atomic radii (see Chapter 2)and as a means of studying atomic structure. In fact, thetechnique of Mssbauer spectroscopy involves the study ofthe effect of changes in s orbital density on nuclear energies.

Same-scale representations of the shapes (angular func-tions) of the 1s and 2s orbitals of an atom are compared inFigure 1.6. The volume of a 2s orbital is about four times

greater than that of a 1s orbital. In both cases, the tiny nucleus is located atthe center of the spheres. These spheres represent the region in which thereis a 99 percent probability of nding an electron. The total probability can-not be represented, for the probability of nding an electron drops to zeroonly at an innite distance from the nucleus.

The Orbitalss

z

y

x

y

x

z

(1s) (2s)

FIGURE 1.6 Representations of theshapes of the 1s and 2s orbitals.[Adapted from D. A. McQuarrie andP. A. Rock, General Chemistry, 2nded. (New York: W. H. Freeman, 1991),p. 322.]

1.2 Shapes of the Atomic Orbitals 7

The probability of finding the electron within an orbital will always bepositive (as the probability is derived from the square of the wave functionand squaring a negative makes a positive). However, when we discuss thebonding of atoms, we nd that the sign related to the original wave functionhas importance. For this reason, it is conventional to superimpose the signof the wave function on the representation of each atomic orbital. For an sorbital, the sign is positive.

In addition to the enormous difference in size between the 1s and the 2sorbitals, the 2s orbital has, at a certain distance from the nucleus, a sphericalsurface on which the electron density is zero.A surface on which the probabil-ity of nding an electron is zero is called a nodal surface. When the principalquantum number increases by 1, the number of nodal surfaces also increasesby 1. We can visualize nodal surfaces more clearly by plotting a graph of theradial density distribution function as a function of distance from the nucleusfor any direction. Figure 1.7 shows plots for the 1s, 2s, and 3s orbitals. Theseplots show that the electron tends to be farther from the nucleus as the princi-pal quantum number increases.The areas under all three curves are the same.

Electrons in an s orbital are different from those in p, d, or f orbitals intwo significant ways. First, only the s orbital has an electron density thatvaries in the same way in every direction out from the atomic nucleus. Sec-ond, there is a finite probability that an electron in an s orbital is at thenucleus of the atom. Every other orbital has a node at the nucleus.

Unlike the s orbitals, the p orbitals are not spherically symmetric. In fact, thep orbitals consist of two separate volumes of space (lobes), with the nucleuslocated between the two lobes. Because there are three p orbitals, we assigneach orbital a direction according to Cartesian coordinates: we have px, py,and pz. Figure 1.8 shows representations of the three 2p orbitals. At rightangles to the axis of higher probability, there is a nodal plane through the

2s

Distance (nm)0.2 0.4 0.6 0.8 0.2 0.4 0.6 0.8 1.0 1.2

Distance (nm)

Prob

abili

ty 1s

0.2 0.4Distance (nm)

Prob

abili

ty

Prob

abili

ty 3s

FIGURE 1.7 The variation of theradial density distribution functionwith distance from the nucleus forelectrons in the 1s, 2s, and 3s orbitalsof a hydrogen atom.

The Orbitalsp


z z z

yy

y

x

x

x

dxy dyz dxz

FIGURE 1.10 Representations of theshapes of the 3dxy, 3dxz, and 3dyzorbitals. [Adapted from L. Jones andP. Atkins, Chemistry: Molecules,Matter, and Change, 3rd ed. (NewYork: W. H. Freeman, 1997), p. 232.]

nucleus. For example, the 2pz orbital has a nodal surface in the xy plane. Interms of wave function sign, one lobe is positive and the other negative.

If we compare graphs of electron density as a function ofatomic radius for the 2s orbital and a 2p orbital (the latterplotted along the axis of higher probability), we nd that the2s orbital has a much greater electron density close to thenucleus than does the 2p orbital (Figure 1.9). Conversely, thesecond maximum of the 2s orbital is farther out than the sin-gle maximum of the 2p orbital. However, the mean distanceof maximum probability is the same for both orbitals.

Like the s orbitals, the p orbitals develop additional nodalsurfaces within the orbital structure as the principal quantum

number increases. Thus, a 3p orbital does not look exactly like a 2p orbitalsince it has an additional nodal surface. However, the detailed differences inorbital shapes for a particular angular momentum quantum number are oflittle relevance in the context of basic inorganic chemistry.

The five d orbitals have more complex shapes. Three of them are locatedbetween the Cartesian axes, and the other two are oriented along the axes. Inall cases, the nucleus is located at the intersection of the axes. Three orbitalseach have four lobes that are located between pairs of axes (Figure 1.10).These orbitals are identied as dxy, dxz, and dyz. The other two d orbitals,and are shown in Figure 1.11. The orbital looks somewhat similarto a pz orbital (see Figure 1.8), except that it has an additional doughnut-shaped ring of high electron density in the xy plane. The orbital isidentical to the dxy orbital but has been rotated through 45.

dx2y2

dz2dx2y2,dz2

The Orbitalsd

FIGURE 1.8 Representations of theshapes of the 2x, 2py, and 2pz orbitals.[Adapted from L. Jones and P. Atkins,Chemistry: Molecules, Matter, andChange, 3rd ed. (New York: W. H.Freeman, 1997), p. 231.] px py pz

z

y

x

FIGURE 1.9 The variation of theradial density distribution functionwith distance from the nucleus forelectrons in the 2s and 2p orbitals of ahydrogen atom.

2s

Distance (nm)

Prob

abili

ty

0.2 0.4 0.6 0.8

2p

Distance (nm)

Prob

abili

ty

0.2 0.4 0.6 0.8

VI

DEO CLIP

VI

DEO CLIP

1.3 The Polyelectronic Atom 9

The f orbitals are even more complex than the d orbitals. There are seven forbitals, four of which have eight lobes. The other three look like the orbital but have two doughnut-shaped rings instead of one.These orbitals arerarely involved in bonding, so we do not need to consider them in any detail.

1.3 The Polyelectronic Atom

In our model of the polyelectronic atom, the electrons are distributed amongthe orbitals of the atom according to the Aufbau (building-up) principle.This simple idea proposes that, when the electrons of an atom are all in theground state, they occupy the orbitals of lowest energy, thereby minimizingthe atoms total electronic energy.Thus, the conguration of an atom can bedescribed simply by adding electrons one by one until the total numberrequired for the element has been reached.

Before starting to construct electron congurations, we need to take intoaccount a second rule: the Pauli exclusion principle. According to this rule,no two electrons in an atom may possess identical sets of the four quantumnumbers. Thus, there can be only one orbital of each three-quantum-number set per atom and each orbital can hold only two electrons, one withms 12 and the other with ms 12.

The simplest conguration is that of the hydrogen atom. According to theAufbau principle, the single electron will be located in the 1s orbital.This con-guration is the ground state of the hydrogen atom. Adding energy wouldraise the electron to one of the many higher energy states. These congura-tions are referred to as excited states. In the diagram of the ground state ofthe hydrogen atom (Figure 1.12), a half-headed arrow is used to indicate thedirection of electron spin. The electron conguration is written as 1s1, withthe superscript 1 indicating the number of electrons in that orbital.

With a two-electron atom (helium), there is a choice: the second electroncould go in the 1s orbital (Figure 1.13a) or the next higher energy orbital,the 2s orbital (Figure 1.13b). Although it might seem obvious that the sec-ond electron would enter the 1s orbital, it is not so simple. If the secondelectron entered the 1s orbital, it would be occupying the same volume ofspace as the electron already in that orbital. The very strong electrostaticrepulsions would discourage the occupancy of the same orbital. For helium,the pairing energy, the energy needed to overcome the interelectronicrepulsive forces, is about However, by occupying an orbital3 MJ#mol1.

dz2

FIGURE 1.11 Representations of theshapes of the and orbitals.[Adapted from L. Jones and P. Atkins,Chemistry: Molecules, Matter, andChange, 3rd ed. (New York: W. H.Freeman, 1997), p. 232.]

3dz23dx2y2

Filling the Orbitalss

z z

y

y

xx

dx2 _ y2 dz2

The Orbitalsf

(a) (b)

2s2s

1s1s

FIGURE 1.13 Two possible electroncongurations for helium.

1s

FIGURE 1.12 Electron congurationof a hydrogen atom.


with a high probability closer to the nucleus, the second electron willexperience a much greater nuclear attraction. The nuclear attraction isgreater than the interelectron repulsion. Hence, the actual configurationwill be 1s2, although it must be emphasized that electrons pair up in thesame orbital only when pairing is the lower energy option.

In the lithium atom the 1s orbital is lled by two electrons, and the thirdelectron must be in the next higher energy orbital, the 2s orbital. Thus,lithium has the conguration of 1s22s1. Because the energy separation of ans and its corresponding p orbitals is always greater than the pairing energyin a polyelectronic atom, the electron configuration of beryllium will be1s22s2 rather than 1s22s12p1.

Boron marks the beginning of the filling of the 2p orbitals. A boron atomhas an electron conguration of 1s22s22p1.As the p orbitals are degenerate(that is, they all have the same energy), it is impossible to decide which oneof the three orbitals contains the electron.

Carbon is the second ground-state atom with electrons in the p orbitals. Itselectron conguration provides another challenge. There are three possiblearrangements of the two 2p electrons (Figure 1.14): (a) both electrons in oneorbital, (b) two electrons with parallel spins in different orbitals, and (c) twoelectrons with opposed spins in different orbitals. On the basis of electronrepulsions, the rst possibility (a) can be rejected immediately. The decisionbetween the other two possibilities is less obvious and requires a deeper knowl-edge of quantum theory. In fact, if the two electrons have parallel spins, there isa zero probability of their occupying the same space. However, if the spins areopposed, there is a nite possibility that the two electrons will occupy the sameregion in space, thereby resulting in some repulsion and a higher energy state.Hence, the parallel spin situation (b) will have the lowest energy. Thispreference for unpaired electrons with parallel spins has been formalized inHunds rule: When lling a set of degenerate orbitals, the number of unpairedelectrons will be maximized and these electrons will have parallel spins.

After the completion of the 2p electron set at neon (1s22s22p6), the 3s and3p orbitals start to ll. Rather than write the full electron congurations, ashortened form can be used. In this notation, the inner electrons are repre-sented by the noble gas symbol having that conguration. Thus, magnesium,whose full electron conguration would be written as 1s22s22p63s2, can berepresented as having a neon noble gas core, and its conguration is writtenas [Ne]3s2. An advantage of the noble gas core representation is that itemphasizes the outermost (valence) electrons, and it is these electrons thatare involved in chemical bonding.At this point, we have nished our analysisof the electron conguration of the two short periods (rows) of the periodictable (Figure 1.15): Period 2, lithium to neon, and Period 3, sodium to argon.

Once the 3p orbitals are lled (argon), the 3d and 4s orbitals start to ll. It ishere that the simple orbital energy level concept breaks down, because theenergy levels of the 4s and 3d orbitals are very close. What becomes mostimportant is not the minimum energy for a single electron but the congura-tion that results in the least number of interelectron repulsions for all the

Filling the Orbitalsp

2p

(a)

2p

(b)

2p

(c)

FIGURE 1.14 Possible 2p electroncongurations for carbon.

Filling the Orbitalsd

1.3 The Polyelectronic Atom 11

electrons. For potassium, this is [Ar]4s1; for calcium, [Ar]4s2.To illustrate howthis delicate balance changes with increasing numbers of protons and elec-trons, the outer electrons in each of the Group 3 to Group 12 elements arelisted here.These congurations are not important in themselves, but they doshow how close the ns and (n 1)d electrons are in energy.

Atom Conguration Atom Conguration Atom Conguration

Sc 4s23d1 Y 5s24d1 Lu 6s25d1

Ti 4s23d2 Zr 5s24d2 Hf 6s25d2

V 4s23d3 Nb 5s14d4 Ta 6s25d3

Cr 4s13d5 Mo 5s14d5 W 6s25d4

Mn 4s23d5 Tc 5s24d5 Re 6s25d5

Fe 4s23d6 Ru 5s14d7 Os 6s25d6

Co 4s23d7 Rh 5s14d8 Ir 6s25d7

Ni 4s23d8 Pd 5s04d10 Pt 6s15d9

Cu 4s13d10 Ag 5s14d10 Au 6s15d10

Zn 4s23d10 Cd 5s24d10 Hg 6s25d10

In general, the lowest overall energy for each transition metal is obtainedby filling the s orbitals first; the remaining electrons then occupy the dorbitals. However, for certain elements, the lowest energy is obtained byshifting one or both of the s electrons to d orbitals. Looking at the rst seriesin isolation would lead to the conclusion that there is some preference for ahalf-full or full set of d orbitals by chromium and copper. Although palla-dium and silver in the second transition series both favor the 4d10 congura-tion, it is more accurate to say that the interelectron repulsion between thetwo s electrons is sufcient in several cases to result in an s1 conguration.

Mo TcNbZrYSrRb

LuBaCs

LrRaFr

W ReTaHf

Sg BhDbRf

Os Ir

Hs Mt Ds Rg Uub Uut Uuq Uup Uuh

Fe Co

Ru Rh

Cr MnVTiSc

Li

Na

K Ca

Be

Mg

Sb TeSnInCdAgPd

HgAuPt Bi PoPbTl At Rn

Br Kr

I Xe

As SeGeGaZnNi Cu

P SSiAl

F Ne

Cl Ar

N OCB

He

Np PuPa UTh NoFm MdEsCfBkAm Cm

H1 2

3 4 5 6 7 8 9 10 11 12

13 14 15 16 17

18

Main groupelements

Main groupelements

Transition metals

Lanthanoids

Actinoids Ac

PmPr NdCeLa Sm YbEr TmHoDyTbEu Gd

1

2

3

4

5

6

7

FIGURE 1.15 Essential features ofthe periodic table.The numbers acrossthe top designate groups (columns) ofelements.The numbers down the right-hand side designate periods (rows).


For the elements from lanthanum (La) to ytterbium (Yb), the situation iseven more uid because the 6s, 5d, and 4f orbitals all have similar energies.For example, lanthanum has a conguration of [Xe]6s25d1, whereas the nextelement, cerium, has a configuration of [Xe]6s24f 2. The most interestingelectron configuration in this row is that of gadolinium, [Xe]6s25d14f 7,rather than the predicted [Xe]6s24f 8. This conguration provides more evi-dence of the importance of interelectron repulsion in the determination ofelectron conguration when adjacent orbitals have similar energies. Similarcomplexities occur among the elements from actinium (Ac) to nobelium(No), in which the 7s, 6d, and 5f orbitals have similar energies.

Although there are minor uctuations in congurations throughout thed-block and f-block elements, the order of lling is quite consistent:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

This order is also shown in Figure 1.16.The orbitals ll in this order becausethe energy differences between the s, p, d, and f orbitals of the same princi-pal quantum number become so great beyond n 2 that they overlap with

Period

Ener

gy

1 1s

2s2p

3p

4p

5p

6p

7p6d

5d

4d

4f

5f

3d

3s

4s

5s

6s

7s

7

6

5

4

3

2

1FIGURE 1.16 Representation of thecomparative energies of the atomicorbitals for lling order purposes.

Though on Earth essentially all atomsare in their ground state, it is not truein interstellar space. Electron transi-tions for the hydrogen atom from ashigh as the n 253 to n 252 levelhave been observed from this regionof the universe. The Bohr radius forsuch a hydrogen atom would beabout 0.34 mm, making the atom solarge it could (theoretically) be seen.

1.4 Ion Electron Congurations 13

the orbitals of the following principal quantum numbers. It is important tonote that Figure 1.16 shows the lling order, not the order for any particularelement. For example, for elements beyond zinc, electrons in the 3d orbitalsare far lower in energy than those in the 4s orbitals. Thus, at this point, the3d orbitals have become inner orbitals and have no role in chemicalbonding. Hence, their precise ordering is unimportant.

1.4 Ion Electron Congurations

For the early main group elements, the common ion electron congurationscan be predicted quite readily.Thus, metals tend to lose all the electrons in theouter orbital set. This situation is illustrated for the isoelectronic series (sameelectron conguration) of sodium, magnesium, and aluminum cations:

Atom Electron conguration Ion Electron conguration

Na [Ne]3s1 Na [Ne]

Mg [Ne]3s2 Mg2 [Ne]

Al [Ne]3s23p1 Al3 [Ne]

Nonmetals gain electrons to complete the outer orbital set. This situation isshown for nitrogen, oxygen, and uorine anions:


N [He]2s22p3 N3 [Ne]

O [He]2s22p4 O2 [Ne]

F [He]2s22p5 F [Ne]

Some of the later main group metals form two ions with different charges.For example, lead forms Pb2 and (rarely) Pb4. The 2 charge can beexplained by the loss of the 6p electrons only, whereas the 4 charge resultsfrom loss of both 6s and 6p electrons:


Pb [Xe]6s24f145d106p2 Pb2 [Xe]6s24f 145d10

Pb4 [Xe]4f145d10

Notice that the electrons of the higher principal quantum number are lostrst.This rule is found to be true for all the elements. For the transition met-als, the s electrons are always lost first when a metal cation is formed. Inother words, for the transition metal cations, the 3d orbitals are always


lower in energy than the 4s orbitals, and a charge of 2, representing theloss of the two s electrons, is common for the transition metals and theGroup 12 metals. For example, zinc always forms an ion of 2 charge:


Zn [Ar]4s23d10 Zn2 [Ar]3d10

Iron forms ions with charges of 2 and 3 and, as shown here, it is tempt-ing to ascribe the formation of the 3 ion to a process in which interelec-tron repulsion forces out the only paired d electron:


Fe [Ar]4s23d6 Fe2 [Ar]3d6

Fe3 [Ar]3d5

It is dangerous, however, to read too much into the electron congurationsof atoms as a means of predicting the ion charges. The series of nickel, pal-ladium, and platinum illustrate this point: they have different congurationsas atoms, yet their common ionic charges and corresponding ion electroncongurations are similar:


Ni [Ar]4s23d8 Ni2 [Ar]3d8

Pd [Kr]5s04d10 Pd2, Pd4 [Kr]4d8, [Kr]4d6

Pt [Xe]6s15d9 Pt2, Pt4 [Xe]5d8, [Xe]5d6

1.5 Magnetic Properties of Atoms

In the discussions of electron conguration, we saw that some atoms possessunpaired electrons. The presence of unpaired electrons in the atoms of anelement can be determined easily from the elements magnetic properties. Ifatoms containing only spin-paired electrons are placed in a magnetic eld,they are weakly repelled by the eld. This phenomenon is called diamagne-tism. Conversely, atoms containing one or more unpaired electrons areattracted by the magnetic eld.This behavior of unpaired electrons is namedparamagnetism. The attraction of each unpaired electron is many timesstronger than the repulsion of all the spin-paired electrons in that atom.

To explain paramagnetism in simple terms, we can visualize the electronas a particle spinning on its axis and generating a magnetic moment, just asan electric current flowing through a wire does. This permanent magnetic

1.5 Magnetic Properties of Atoms 15

moment results in an attraction into the stronger part of the field. Whenelectrons have their spins paired, the magnetic moments cancel each other.As a result, the paired electrons are weakly repelled by the lines of force ofthe magnetic field. In paramagnetic materials, application of a magneticeld aligns some of the normally randomly oriented electron spins with theapplied magnetic eld (Figure 1.17a and b). It is this alignment that resultsin the attraction of the material into the magnetic eld. We will encounterthis phenomenon again in our discussions of covalent bonding and thebonding in transition metal compounds.

There is a third relatively common form of magnetic behaviorferromagnetism. In ferromagnetic materials, the unpaired electrons are par-allel aligned with their neighbors even in the absence of a magnetic field.These groups of mutually aligned spins are known as magnetic domains.Application of a magnetic field causes all these domains to align with themagnetic field (Figure 1.17c and d). This alignment is much stronger thanthat of paramagnetism, and it can be permanent.

Ferromagnetism is found in elements (and some of their compounds)that have unpaired electrons in their d or f orbitals. Electrons in theseorbitals must weakly interact with those in neighboring atoms for the effectto occur. The phenomenon only occurs in the later 3d- and 4f-block ele-ments. When a ferromagnetic material is heated, the atomic vibrationscause a breakdown of the magnetic domains until, at the Curie temperature,the material reverts to the weaker paramagnetic behavior. Only four metalsexhibit ferromagnetism and have a Curie transition above 0C: iron, cobalt,nickel, and gadolinium.

A fourth class of magnetic behavior is antiferromagnetism. Antiferro-magnetism is similar to ferromagnetism except that the weak interactionsbetween neighboring atoms result in an antiparallel alignment. Thus, theattraction into a magnetic field is weaker than the paramagnetic effectwould predict but only up to the Nel temperature, at which the antiferro-magnetic material reverts to paramagnetic behavior.

(a)(b)

(d)(c)

H

FIGURE 1.17 The behav

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Atkins Inorganic Chem

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