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Descriptive Inorganic Chemistry

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  • Descriptive Inorganic Chemistry FIFTH EDITION

    Geoff Rayner-CanhamSir Wilfred Grenfell CollegeMemorial University

    Tina OvertonUniversity of Hull

    W. H. FREEMAN AND COMPANYNEW YORK

  • Publisher: Clancy Marshall

    Acquisitions Editors: Jessica Fiorillo/Kathryn Treadway

    Marketing Director: John Britch

    Media Editor: Dave Quinn

    Cover and Text Designer: Vicki Tomaselli

    Senior Project Editor: Mary Louise Byrd

    Illustrations: Network Graphics/Aptara

    Senior Illustration Coordinator: Bill Page

    Production Coordinator: Susan Wein

    Composition: Aptara

    Printing and Binding: World Color Versailles

    Library of Congress Control Number: 2009932448

    ISBN-13: 978-1-4292-2434-5

    ISBN-10: 1-4292-1814-2

    @2010, 2006, 2003, 2000 by W. H. Freeman and Company

    All rights reserved

    Printed in the United States of America

    First printing

    W. H. Freeman and Company

    41 Madison Avenue

    New York, NY 10010

    Houndmills, Basingstoke RG21 6XS, England

    www.whfreeman.com

  • CHAPTER 1 The Electronic Structure of the Atom: A Review 1

    CHAPTER 2 An Overview of the Periodic Table 19

    CHAPTER 3 Covalent Bonding 41

    CHAPTER 4 Metallic Bonding 81

    CHAPTER 5 Ionic Bonding 93

    CHAPTER 6 Inorganic Thermodynamics 113

    CHAPTER 7 Solvent Systems and Acid-Base Behavior 137

    CHAPTER 8 Oxidation and Reduction 167

    CHAPTER 9 Periodic Trends 191

    CHAPTER 10 Hydrogen 227

    CHAPTER 11 The Group 1 Elements: The Alkali Metals 245

    CHAPTER 12 The Group 2 Elements: The Alkaline Earth Metals 271

    CHAPTER 13 The Group 13 Elements 291

    CHAPTER 14 The Group 14 Elements 315

    CHAPTER 15 The Group 15 Elements: The Pnictogens 363

    CHAPTER 16 The Group 16 Elements: The Chalcogens 409

    CHAPTER 17 The Group 17 Elements: The Halogens 453

    CHAPTER 18 The Group 18 Elements: The Noble Gases 487

    CHAPTER 19 Transition Metal Complexes 499

    CHAPTER 20 Properties of the 3d Transition Metals 533

    CHAPTER 21 Properties of the 4d and 5d Transition Metals 579

    CHAPTER 22 The Group 12 Elements 599

    CHAPTER 23 Organometallic Chemistry 611

    On the Web www.whfreeman.com/descriptive5e

    CHAPTER 24 The Rare Earth and Actinoid Elements 651w

    Appendices A-1

    Index I-1

    Overview

    iii

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  • Contents

    What Is Descriptive Inorganic Chemistry? xiiiPreface xvAcknowledgments xixDedication xxi

    CHAPTER 1

    The Electronic Structure of the Atom:A Review 1Atomic Absorption Spectroscopy 21.1 The Schrdinger Wave Equation and Its

    Signi cance 31.2 Shapes of the Atomic Orbitals 51.3 The Polyelectronic Atom 91.4 Ion Electron Con gurations 141.5 Magnetic Properties of Atoms 151.6 Medicinal Inorganic Chemistry: An Introduction 16

    CHAPTER 2

    An Overview of the Periodic Table 192.1 Organization of the Modern

    Periodic Table 212.2 Existence of the Elements 232.3 Stability of the Elements and Their Isotopes 24The Origin of the Shell Model of the Nucleus 262.4 Classi cations of the Elements 272.5 Periodic Properties: Atomic Radius 292.6 Periodic Properties: Ionization Energy 332.7 Periodic Properties: Electron Af nity 35Alkali Metal Anions 372.8 The Elements of Life 37

    CHAPTER 3

    Covalent Bonding 413.1 Models of Covalent Bonding 423.2 Introduction to Molecular Orbitals 433.3 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 50

    3.6 A Brief Review of Lewis Structures 513.7 Partial Bond Order 533.8 Formal Charge 543.9 Valence-Shell Electron-Pair Repulsion Rules 543.10 The Valence-Bond Concept 593.11 Network Covalent Substances 613.12 Intermolecular Forces 63The Origins of the Electronegativity Concept 653.13 Molecular Symmetry 663.13 Symmetry and Vibrational Spectroscopy 72Transient SpeciesA New Direction for Inorganic Chemistry 743.15 Covalent Bonding and the Periodic Table 78

    CHAPTER 4

    Metallic Bonding 814.1 Metallic Bonding 814.2 Bonding Models 824.3 Structure of Metals 844.4 Unit Cells 864.5 Alloys 87Memory Metal: The Shape of Things to Come 884.6 Nanometal Particles 894.7 Magnetic Properties of Metals 90

    CHAPTER 5

    Ionic Bonding 935.1 The Ionic Model and the Size of Ions 935.2 Hydrated Salts 955.3 Polarization and Covalency 965.4 Ionic Crystal Structures 995.5 Crystal Structures Involving

    Polyatomic Ions 1055.6 The Bonding Continuum 106Concrete: An Old Material with a New Future 109

    v

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    CHAPTER 6

    Inorganic Thermodynamics 1136.1 Thermodynamics of the Formation

    of Compounds 1146.2 Formation of Ionic Compounds 1206.3 The Born-Haber Cycle 1226.4 Thermodynamics of the Solution

    Process for Ionic Compounds 1246.5 Formation of Covalent Compounds 127The Hydrogen Economy 1286.6 Thermodynamic versus Kinetic Factors 129

    CHAPTER 7

    Solvent Systems and Acid-Base Behavior 1377.1 Solvents 1387.2 Brnsted-Lowry Acids 142Antacids 1447.3 Brnsted-Lowry Bases 147Cyanide and Tropical Fish 1487.4 Trends in Acid-Base Behavior 148Superacids and Superbases 1507.5 Acid-Base Reactions of Oxides 1537.6 Lewis Theory 1557.7 Pearson Hard-Soft Acid-Base Concepts 1567.8 Applications of the HSAB Concept 1587.9 Biological Aspects 161

    CHAPTER 8

    Oxidation and Reduction 1678.1 Redox Terminology 1678.2 Oxidation Number Rules 1688.3 Determination of Oxidation Numbers

    from Electronegativities 1698.4 The Difference between Oxidation

    Number and Formal Charge 1718.5 Periodic Variations of Oxidation

    Numbers 1728.6 Redox Equations 173Chemosynthesis: Redox Chemistry on the Sea oor 1758.7 Quantitative Aspects of Half-Reactions 176

    8.8 Electrode Potentials as Thermodynamic Functions 177

    8.9 Latimer (Reduction Potential) Diagrams 1788.10 Frost (Oxidation State) Diagrams 1808.11 Pourbaix Diagrams 1828.12 Redox Synthesis 1848.13 Biological Aspects 185

    CHAPTER 9

    Periodic Trends 1919.1 Group Trends 1929.2 Periodic Trends in Bonding 1959.3 Isoelectronic Series in Covalent

    Compounds 1999.4 Trends in Acid-Base Properties 2019.5 The (n) Group and (n 10) Group

    Similarities 202Chemical Topology 2069.6 Isomorphism in Ionic Compounds 207New Materials: Beyond the Limitations of Geochemistry 2099.7 Diagonal Relationships 210Lithium and Mental Health 2119.8 The Knights Move Relationship 2129.9 The Early Actinoid Relationships 2159.10 The Lanthanoid Relationships 2169.11 Combo Elements 2179.12 Biological Aspects 221Thallium Poisoning: Two Case Histories 223

    CHAPTER 10

    Hydrogen 22710.1 Isotopes of Hydrogen 22810.2 Nuclear Magnetic Resonance 229Isotopes in Chemistry 23010.3 Properties of Hydrogen 231Searching the Depths of Space for the Trihydrogen Ion 23310.4 Hydrides 23310.5 Water and Hydrogen Bonding 237Water: The New Wonder Solvent 23810.6 Clathrates 239

    Contents

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    10.7 Biological Aspects of Hydrogen Bonding 241

    Is There Life Elsewhere in Our Solar System? 24210.8 Element Reaction Flowchart 242

    CHAPTER 11

    The Group 1 Elements: The Alkali Metals 24511.1 Group Trends 24611.2 Features of Alkali Metal Compounds 24711.3 Solubility of Alkali Metal Salts 249Mono Lake 25011.4 Lithium 25211.5 Sodium 25511.6 Potassium 25611.7 Oxides 25711.8 Hydroxides 25911.9 Sodium Chloride 261Salt Substitutes 26111.10 Potassium Chloride 26211.11 Sodium Carbonate 26211.12 Sodium Hydrogen Carbonate 26411.13 Ammonia Reaction 26411.14 Ammonium Ion as a Pseudo

    Alkali-Metal Ion 26511.15 Biological Aspects 26511.16 Element Reaction Flowcharts 266

    CHAPTER 12

    The Group 2 Elements: The Alkaline Earth Metals 27112.1 Group Trends 27112.2 Features of Alkaline Earth Metal

    Compounds 27212.3 Beryllium 27512.4 Magnesium 27612.5 Calcium and Barium 27812.6 Oxides 27912.7 Calcium Carbonate 280How Was Dolomite Formed? 28112.8 Cement 28212.9 Calcium Chloride 283

    Biomineralization: A New Interdisciplinary Frontier 28412.10 Calcium Sulfate 28412.11 Calcium Carbide 28512.12 Biological Aspects 28612.15 Element Reaction Flowcharts 287

    CHAPTER 13

    The Group 13 Elements 29113.1 Group Trends 29213.2 Boron 29313.3 Borides 294Inorganic Fibers 29513.4 Boranes 295Boron Neutron Capture Therapy 29813.5 Boron Halides 30013.6 Aluminum 30113.7 Aluminum Halides 30613.8 Aluminum Potassium Sulfate 30713.9 Spinels 30813.10 Aluminides 30913.11 Biological Aspects 30913.12 Element Reaction Flowcharts 311

    CHAPTER 14

    The Group 14 Elements 31514.1 Group Trends 31614.2 Contrasts in the Chemistry of Carbon

    and Silicon 31614.3 Carbon 318The Discovery of Buckminsterfullerene 32214.4 Isotopes of Carbon 32514.5 Carbides 326Moissanite: The Diamond Substitute 32714.6 Carbon Monoxide 32814.7 Carbon Dioxide 330Carbon Dioxide, Supercritical Fluid 33214.8 Carbonates and Hydrogen Carbonates 33314.9 Carbon Sul des 33514.10 Carbon Halides 33514.11 Methane 33814.12 Cyanides 339

    Contents

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    14.13 Silicon 33914.14 Silicon Dioxide 34114.15 Silicates 34314.16 Aluminosilicates 34514.17 Silicones 349Inorganic Polymers 35014.18 Tin and Lead 35114.19 Tin and Lead Oxides 35214.20 Tin and Lead Halides 35314.21 Tetraethyllead 354TEL: A Case History 35514.22 Biological Aspects 35614.23 Element Reaction Flowcharts 359

    CHAPTER 15

    The Group 15 Elements: The Pnictogens 36315.1 Group Trends 36415.2 Contrasts in the Chemistry

    of Nitrogen and Phosphorus 36515.3 Overview of Nitrogen Chemistry 368The First Dinitrogen Compound 36915.4 Nitrogen 369Propellants and Explosives 37015.5 Nitrogen Hydrides 371Haber and Scienti c Morality 37415.6 Nitrogen Ions 37715.7 The Ammonium Ion 37815.8 Nitrogen Oxides 37915.9 Nitrogen Halides 38415.10 Nitrous Acid and Nitrites 38515.11 Nitric Acid and Nitrates 38615.12 Overview of Phosphorus Chemistry 38915.13 Phosphorus 390Nauru, the Worlds Richest Island 39115.14 Phosphine 39315.15 Phosphorus Oxides 39315.16 Phosphorus Chlorides 39415.17 Phosphorus Oxo-Acids and Phosphates 39515.18 The Pnictides 39915.19 Biological Aspects 399Paul Erhlich and His Magic Bullet 40115.29 Element Reaction Flowcharts 402

    CHAPTER 16

    The Group 16 Elements: The Chalcogens 40916.1 Group Trends 41016.2 Contrasts in the Chemistry of

    Oxygen and Sulfur 41116.3 Oxygen 412Oxygen Isotopes in Geology 41216.4 Bonding in Covalent Oxygen

    Compounds 41816.5 Trends in Oxide Properties 41916.6 Mixed-Metal Oxides 421New Pigments through Perovskites 42216.7 Water 42216.8 Hydrogen Peroxide 42416.9 Hydroxides 42416.10 The Hydroxyl Radical 42616.11 Overview of Sulfur Chemistry 42616.12 Sulfur 427Cosmochemistry: Io, the Sulfur-Rich Moon 42816.13 Hydrogen Sul de 43116.14 Sul des 432Disul de Bonds and Hair 43216.15 Sulfur Oxides 43416.16 Sul tes 43716.17 Sulfuric Acid 43816.18 Sulfates and Hydrogen Sulfates 44016.19 Other Oxy-Sulfur Anions 44116.20 Sulfur Halides 44316.21 Sulfur-Nitrogen Compounds 44516.22 Selenium 44516.23 Biological Aspects 44616.24 Element Reaction Flowcharts 448

    CHAPTER 17

    The Group 17 Elements: The Halogens 45317.1 Group Trends 45417.2 Contrasts in the Chemistry of

    Fluorine and Chlorine 45517.3 Fluorine 458The Fluoridation of Water 45917.4 Hydrogen Fluoride and Hydro uoric Acid 46017.5 Overview of Chlorine Chemistry 462

    Contents

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    17.6 Chlorine 46317.7 Hydrochloric Acid 46417.8 Halides 46517.9 Chlorine Oxides 46917.10 Chlorine Oxyacids and Oxyanions 471Swimming Pool Chemistry 473

    The Discovery of the Perbromate Ion 47417.11 Interhalogen Compounds and

    Polyhalide Ions 47517.12 Cyanide Ion as a Pseudo-halide Ion 47717.13 Biological Aspects 47817.14 Element Reaction Flowcharts 481

    CHAPTER 18

    The Group 18 Elements: The Noble Gases 48718.1 Group Trends 48818.2 Unique Features of Helium 48918.3 Uses of the Noble Gases 48918.4 A Brief History of Noble Gas

    Compounds 491Is It Possible to Make Compounds of the Early Noble Gases? 49218.5 Xenon Fluorides 49218.6 Xenon Oxides 49418.7 Other Noble Gas Compounds 49518.8 Biological Aspects 49518.9 Element Reaction Flowchart 496

    CHAPTER 19

    Transition Metal Complexes 49919.1 Transition Metals 49919.2 Introduction to Transition Metal

    Complexes 50019.3 Stereochemistries 50219.4 Isomerism in Transition Metal

    Complexes 503Platinum Complexes and Cancer Treatment 50619.5 Naming Transition Metal Complexes 50719.6 An Overview of Bonding Theories

    of Transition Metal Compounds 51019.7 Crystal Field Theory 51119.8 Successes of Crystal Field Theory 517

    The Earth and Crystal Structures 52119.9 More on Electronic Spectra 52119.10 Ligand Field Theory 52319.11 Thermodynamic versus Kinetic Factors 52519.12 Synthesis of Coordination Compounds 52619.13 Coordination Complexes and the

    HSAB Concept 52719.14 Biological Aspects 529

    CHAPTER 20

    Properties of the 3d Transition Metals 53320.1 Overview of the 3d Transition Metals 53420.2 Group 4: Titanium 53620.3 Group 5: Vanadium 53720.4 Group 6: Chromium 53820.5 Group 7: Manganese 544Mining the Sea oor 54520.6 Group 8: Iron 54920.7 Group 9: Cobalt 55820.8 Group 10: Nickel 56220.9 Group 11: Copper 56320.10 Biological Aspects 56920.11 Element Reaction Flowcharts 572

    CHAPTER 21

    Properties of the 4d and 5d Transition Metals 57921.1 Comparison of the Transition Metals 58021.2 Features of the Heavy

    Transition Metals 58121.3 Group 4: Zirconium and Hafnium 58421.4 Group 5: Niobium and Tantalum 58521.5 Group 6: Molybdenum and Tungsten 58621.6 Group 7: Technetium and Rhenium 587Technetium: The Most Important Radiopharmaceutical 58821.7 The Platinum Group Metals 58921.8 Group 8: Ruthenium and Osmium 59021.9 Group 9: Rhodium and Iridium 59121.10 Group 10: Palladium and Platinum 59121.11 Group 11: Silver and Gold 59121.12 Biological Aspects 594

    Contents

  • xCHAPTER 22

    The Group 12 Elements 59922.1 Group Trends 60022.2 Zinc and Cadmium 60022.3 Mercury 60322.4 Biological Aspects 605Mercury Amalgam in Teeth 60722.5 Element Reaction Flowchart 608

    CHAPTER 23

    Organometallic Chemistry 61123.1 Introduction to Organometallic

    Compounds 61223.2 Naming Organometallic Compounds 61223.3 Counting Electrons 61323.4 Solvents for Organometallic Chemistry 61423.5 Main Group Organometallic

    Compounds 615Grignard Reagents 618

    The Death of Karen Wetterhahn 62323.6 Organometallic Compounds of the

    Transition Metals 62323.7 Transition Metal Carbonyls 62523.8 Synthesis and Properties of Simple

    Metal Carbonyls 63023.9 Reactions of Transition Metal Carbonyls 63223.10 Other Carbonyl Compounds 63323.11 Complexes with Phosphine Ligands 63423.12 Complexes with Alkyl, Alkene, and

    Alkyne Ligands 635Vitamin B12A Naturally Occurring Organometallic Compound 63823.13 Complexes with Allyl and 1,3-Butadiene

    Ligands 63923.14 Metallocenes 64023.15 Complexes with 6-Arene Ligands 64223.16 Complexes with Cycloheptatriene and

    Cyclooctatetraene Ligands 643

    23.17 Fluxionality 64323.18 Organometallic Compounds in

    Industrial Catalysis 644

    CHAPTER 24 ON THE WEB www.whfreeman.com/descriptive5e

    The Rare Earth and Actinoid Elements 651w24.1 The Group 3 Elements 653w24.2 The Lanthanoids 653wSuperconductivity 655w23.3 The Actinoids 656w24.4 Uranium 659wA Natural Fission Reactor 661w24.5 The Postactinoid Elements 662w

    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-16Appendix 4 Ionization Energies of Selected

    Metals A-18Appendix 5 Electron Af nities of Selected

    Nonmetals A-20Appendix 6 Selected Lattice Energies A-21Appendix 7 Selected Hydration Enthalpies A-22Appendix 8 Selected Ionic Radii A-23ON THE WEB www.whfreeman.com/descriptive5e

    Appendix 9 Standard Half-Cell Electrode Potentials of Selected Elements A-25w

    ON THE WEB www.whfreeman.com/descriptive5e

    Appendix 10 Electron Con guration of the Elements A-35w

    INDEX I-1

    Contents

  • 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 the subject, with the synthesis of new and novel materials, the properties are being linked with explanations for the formulas and structures of compounds together with an understanding of the chemical reactions they undergo. In addition, we are no longer looking at inorganic chemistry as an isolated subject but as a part of essential scienti c knowledge with applications throughout science and our lives. Because of a need for greater contextualization, we have added more features and more applications. In many colleges and universities, descriptive inorganic chemistry is offered as a sophomore or junior course. In this way, students come to know something of the fundamental properties of important and interesting elements and their compounds. Such knowledge is important for careers not only in pure or applied chemistry but also in pharmacy, medicine, geology, and environmental science. This course can then be followed by a junior or senior course that focuses on the theoretical principles and the use of spectroscopy to a greater depth than is covered in a descriptive text. In fact, 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 this has been true of inorganic chemistry. Up until the 1960s, it was very much pure descriptive, requiring exclusively memorization. In the 1970s and 1980s, upper-level texts focused exclusively on the theoretical principles. Now it is ap-parent that descriptive is very importantnot the traditional memorization of facts but the linking of facts, where possible, to underlying principles. Students need to have modern descriptive inorganic chemistry as part of their educa-tion. Thus, we must ensure that chemists are aware of the new descriptive inorganic chemistry.

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  • Preface

    Inorganic chemistry goes beyond academic interest: it is an im-portant part of our lives.

    Inorganic chemistry is interestingmore than thatit is exciting! So much of our twenty- rst-century science and technology rely on natural and syn-thetic materials, often inorganic compounds, many of which are new and novel. Inorganic chemistry is ubiquitous in our daily lives: household products, some pharmaceuticals, our transportationboth the vehicles themselves and the synthesis of the fuelsbattery technology, and medical treatments. There is the industrial aspect, the production of all the chemicals required to drive our economy, everything from steel to sulfuric acid to glass and cement. Environ-mental chemistry is largely a question of the inorganic chemistry of the atmo-sphere, water, and soil. Finally, there are the profound issues of the inorganic chemistry of our planet, the solar system, and the universe.

    This textbook is designed to focus on the properties of selected interesting, important, and unusual elements and compounds. However, to understand inorganic chemistry, it is crucial to tie this knowledge to the underlying chemi-cal principles and hence provide explanations for the existence and behavior of compounds. For this reason, almost half the chapters survey the relevant concepts of atomic theory, bonding, intermolecular forces, thermodynamics, acid-base behavior, and reduction-oxidation properties as a prelude to, and preparation for, the descriptive material.

    For this fth edition, the greatest change has been the expansion of coverage of the 4d and 5d transition metals to a whole chapter.

    The heavier transition metals have unique trends and patterns, and the new chapter highlights these. Having an additional chapter on transition met-als also better balances the coverage between the main group elements and the transition elements.

    Also, the fth edition has a second color. With the addition of a second color, gures are much easier to understand, and tables and text are easier to read.

    On a chapter-by-chapter basis, the signi cant improvements are as follows:

    Chapter 1: The Electronic Structure of the Atom: A ReviewThe Introduction and Section 1.3, The Polyelectronic Atom, have been revised.

    Chapter 3: Covalent BondingSection 3.11, Network Covalent Substances, has a new subsection: AmorphousSilicon.

    Chapter 4: Metallic BondingSection 4.6, Nanometal Particles, was added.Section 4.7, Magnetic Properties of Metals, was added.

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    Chapter 5: Ionic BondingSection 5.3, Polarization and Covalency, has a new subsection: The Ionic-Covalent Boundary.Section 5.4, Ionic Crystal Structures, has a new subsection: Quantum Dots.

    Chapter 9: Periodic TrendsSection 9.3, Isoelectronic Series in Covalent Compounds, has been revised and improved.Section 9.8, The Knights Move Relationship, has been revised and improved.

    Chapter 10: HydrogenSection 10.4, Hydrides, has a revised and expanded subsection: Ionic Hydrides.

    Chapter 11: The Group 1 ElementsSection 11.14, Ammonium Ion as a PseudoAlkali-Metal Ion, moved from Chapter 9.

    Chapter 13: The Group 13 ElementsSection 13.10, Aluminides, was added.

    Chapter 14: The Group 14 ElementsSection 14.2, Contrasts in the Chemistry of Carbon and Silicon, was added.Section 14.3, Carbon, has a new subsection: Graphene.Section 14.7, Carbon Dioxide, has a new subsection: Carbonia.

    Chapter 15: The Group 15 ElementsSection 15.2, Contrasts in the Chemistry of Nitrogen and Phosphorus, was added.Section 15.18, The Pnictides, was added.

    Chapter 16: The Group 16 ElementsSection 16.2, Contrasts in the Chemistry of Oxygen and Sulfur, was added.Section 16.14, Sul des, has a new subsection: Disul des.

    Chapter 17: The Group 17 ElementsSection 17.2, Contrasts in the Chemistry of Fluorine and Chlorine, was added.Section 17.12, Cyanide Ion as a Pseudo-halide Ion, moved from Chapter 9.

    Chapter 18: The Group 18 ElementsSection 18.7, Other Noble Gas Compounds, was added.

    Chapter 19: Transition Metal ComplexesSection 19.10, Ligand Field Theory, was added.

    Chapter 20: Properties of the 3d Transition MetalsSection 20.1, Overview of the 3d Transition Metals, was added.

    Chapter 21: Properties of the 4d and 5d Transition MetalsNEW CHAPTER added (for details, see the previous page).

    Chapter 24: The Rare Earth and Actinoid ElementsThis chapter has been signi cantly revised with the new subsections Scandium, Yttrium, and Thorium.

    Preface

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    ALSOVideo ClipsDescriptive inorganic chemistry by de nition is visual, so what better way to appreciate a chemical reaction than to make it visual? We now have a series of at least 60 Web-based video clips to bring some of the reactions to life. The text has a margin icon to indicate where a reaction is illustrated.

    Text Figures and TablesAll the illustrations and tables in the book are available as .jpg les for inclusion in PowerPoint presentations on the instructor side of the Web site at www.whfreeman.com/descriptive5e.

    Additional ResourcesA list of relevant SCIENTIFIC AMERICAN articles is found on the text Web site at www.whfreeman.com/descriptive5e. The text has a margin icon to indicate where a Scienti c American article is available.

    SupplementsThe Student Solutions Manual, ISBN: 1-4292-2434-7 contains the worked solutions to all the odd-numbered end-of-chapter problems.

    The Companion Web Site www.whfreeman.com/descriptive5eContains the following student-friendly materials: Chapter 24: The Rare Earth and Actinoid Elements, Appendices, Lab Experiments, Tables, and over 50 usefulvideos of elements and metals in reactions and oxidations.

    Instructors Resource CD-ROM, ISBN: 1-4292-2428-2Includes PowerPoint and videos as well as all text art and solutions to all prob-lems in the book.

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

    Preface

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  • Acknowledgments

    Many thanks must go to the team at W. H. Freeman and Company who have contributed their talents to the ve editions of this book. We offer our sincere gratitude to the editors of the fth edition, Jessica Fiorillo, Kathryn Treadway, and Mary Louise Byrd; of the fourth edition, Jessica Fiorillo, Jenness Crawford, and Mary Louise Byrd; of the third edition, Jessica Fiorillo and Guy Copes; of the second edition, Michelle Julet and Mary Louise Byrd; and a special thanks to Deborah Allen, who bravely commissioned the rst edition of the text. Each one of our fabulous editors has been a source of encouragement, support, and helpfulness. We wish to acknowledge the following reviewers of this edition, whose criticisms and comments were much appreciated: Theodore Betley at Harvard University; Dean Campbell at Bradley University; Maria Contel at Brooklyn College (CUNY); Gerry Davidson at St. Francis College; Maria Derosa at Carleton University; Stan Duraj at Cleveland State University; Dmitri Giarkios at Nova Southeastern University; Michael Jensen at Ohio UniversityMain Campus; David Marx at the University of Scranton; Joshua Moore at Tennessee State UniversityNashville; Stacy OReilly at Butler University; William Pen-nington at Clemson University; Daniel Rabinovich at the University of North Carolina at Charlotte; Hal Rogers at California State UniversityFullerton; Thomas Schmedake at the University of North Carolina at Charlotte; Bradley Smucker at Austin College; Sabrina Sobel at Hofstra University; Ronald Strange at Fairleigh Dickinson UniversityMadison; Mark Walters at New York University; Yixuan Wang at Albany State University; and Juchao Yan at Eastern New Mexico University; together with prereviewers: Londa Borer at California State UniversitySacramento; Joe Fritsch at Pepperdine Univer-sity; Rebecca Roesner at Illinois Wesleyan University, and Carmen Works at Sonoma College. We acknowledge with thanks the contributions of the reviewers of the fourth edition: Rachel Narehood Austin at Bates College; Leo A. Bares at the University of North CarolinaAsheville; Karen S. Brewer at Hamilton College; Robert M. Burns at Alma College; Do Chang at Averett University; Georges Dns at Concordia University; Daniel R. Derringer at Hollins University; Carl P. Fictorie at Dordt College; Margaret Kastner at Bucknell University; Michael Laing at the University of Natal, Durban; Richard H. Langley at Stephen F. Austin State University; Mark R. McClure at the University of North Carolina at Pembroke; Louis Mercier at Laurentian University; G. Merga at Andrews University; Stacy OReilly at Butler University; Larry D. Pedersen at College Misercordia; Robert D. Pike at the College of William and Mary; William Quintana at New Mexico State University; David F. Rieck at Salisbury 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 the State University of New York at Cortland.

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    And the contributions of the reviewers of the third edition: Franois Caron at Laurentian University; Thomas D. Getman at Northern Michigan Univer-sity; Janet R. Morrow at the State University of New York at Buffalo; Robert D. Pike at the College of William and Mary; Michael B. Wells at Cambell Uni-versity; and particularly Joe Takats of the University of Alberta for his compre-hensive critique of the second edition. And the contributions of the reviewers of the second edition: F. C. Hentz at North Carolina State University; Michael D. Johnson at New Mexico State University; Richard B. Kaner at the University of California, Los Angeles; Richard H. Langley at Stephen F. Austin State University; James M. Mayer at the University of Washington; Jon Melton at Messiah College; Joseph S. Merola at Virginia Technical Institute; David Phillips at Wabash College; John R. Pladziewicz at the University of Wisconsin, Eau Claire; Daniel Rabinovich at the University of North Carolina at Charlotte; David F. Reich at Salisbury State University; Todd K. Trout at Mercyhurst College; Steve Watton at the Virginia Commonwealth University; and John S. Wood at the University of Massachusetts, Amherst. Likewise, the reviewers of the rst edition: E. Joseph Billo at Boston Col-lege; David Finster at Wittenberg University; Stephen J. Hawkes at Oregon State University; Martin Hocking at the University of Victoria; Vake Marganian at Bridgewater State College; Edward Mottel at the Rose-Hulman Institute of Technology; and Alex Whitla at Mount Allison University. As a personal acknowledgment, Geoff Rayner-Canham wishes to especial-ly thank three teachers and mentors who had a major in uence on his career: Briant Bourne, Harvey Grammar School; Margaret Goodgame, Imperial Col-lege, London University; and Derek Sutton, Simon Fraser University. And he expresses his eternal gratitude to his spouse, Marelene, for her support and encouragement. Tina Overton would like to thank her colleague Phil King for his invaluable suggestions for improvements and his assistance with the illustrations. Thanks must also go to her family, Dave, John, and Lucy, for their patience during the months when this project lled all her waking hours.

    Acknowledgments

  • Dedication

    Chemistry is a human endeavor. New discoveries are the result of the work of enthusiastic people and groups of people who want to explore the molecular world. We hope that you, the reader, will come to share our own fascination with inorganic chemistry. We have chosen to dedicate this book to two scientists who, for very different reasons, never did receive the ultimate accolade of a Nobel Prize.

    Henry Moseley (18871915)Although Mendeleev is identi ed as the discoverer of the peri-odic table, his version was based on an increase in atomic mass. In some cases, the order of elements had to be reversed to match properties with location. It was a British scientist, Henry Moseley, who put the periodic table on a much rmer footing by discov-ering that, on bombardment with electrons, each element emit-ted X-rays of characteristic wavelengths. The wavelengths tted a formula related by an integer number unique to each element. We know that number to be the number of protons. With the es-tablishment of the atomic number of an element, 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 scienti c talents of the twentieth century died at the age of 27. The famous American scientist Robert Milliken commented: Had the European War had no other result than the snuf ng out of this young life, that alone would make it one of the most hideous and most irreparable crimes in history. Unfortunately, Nobel Prizes are only awarded to living scientists. In 1924, the discovery of element 43 was claimed, and it was named mose-leyum; however, the claim was disproved by the very method that Moseley had pioneered.

    xix

  • xx

    Lise Meitner (18781968)In the 1930s, scientists were bombarding atoms of heavy elements such as uranium with subatomic particles to try to make new ele-ments and extend the periodic table. Austrian scientist Lise Meit-ner had shared leadership with Otto Hahn of the German research team working on the synthesis of new elements; the team thought they had discovered nine new elements. Shortly after the claimed discovery, Meitner was forced to ee Germany because of her Jewish ancestry, and she settled in Sweden. Hahn reported to 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 drop of water. No wonder the element formed behaved like barium: it was barium! Thus was born the concept of nuclear ssion. She informed Hahn of her proposal. When Hahn wrote the research paper on the work, he barely mentioned the

    vital contribution of Meitner and Frisch. As a result, Hahn and his colleague Fritz Strassmann received the Nobel Prize. Meitners ash of genius was ignored. Only recently has Meitner received the acclaim she deserved 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-Twentieth Century. 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.

    Dedication

  • Isaac Newton was the original model for the absentminded professor. Supposedly, he always timed the boiled egg he ate at breakfast; one morning, his maid found him standing by the pot of boiling water, hold-ing an egg in his hand and gazing intently at the watch in the bottom of the pot! Nevertheless, it was Newton who initiated the study of the electronic structure of the atom in about 1700, when he noticed that the passage of sunlight through a prism produced a continuous visible spectrum. Much later, in 1860, Robert Bunsen (of burner fame) inves-tigated the light emissions from ames and gases. Bunsen observed that the emission spectra, rather than being continuous, were series of colored lines (line spectra).

    The proposal that electrons existed in concentric shells had its origin in the research of two overlooked pioneers: Johann Jakob Balmer, a Swiss mathematician, and Johannes Robert Rydberg, a Swedish physicist. After an undistinguished career in mathematics, in 1885, at the age of 60, Balmer studied the visible emission lines of the hydrogen atom and found that there was a mathematical relationship between the wave-lengths. Following from Balmers work, in 1888, Rydberg deduced a more general relationship:

    1l

    5 RHa 1n2f

    21

    n2ib

    where l is the wavelength of the emission line, RH is a constant, later known as the Rydberg constant, and nf and ni are integers. For the visible lines seen by Balmer and Rydberg, nf had a value of 2. The Rydberg formula received further support in 1906, when Theodore Lyman found a series of lines in the far-ultraviolet spectrum of hydrogen,

    1.1 The Schrdinger Wave Equation and Its Signi cance

    Atomic Absorption Spectroscopy

    1.2 Shapes of the Atomic Orbitals

    1.3 The Polyelectronic Atom

    1.4 Ion Electron Con gurations

    1.5 Magnetic Properties of Atoms

    1.6 Medicinal Inorganic Chemistry: An Introduction

    The Electronic Structure of the Atom: A Review

    CHAPTER 1

    1

    To understand the behavior of inorganic compounds, we need to study

    the nature of chemical bonding. Bonding, in turn, relates to the behavior

    of electrons in the constituent atoms. Our study of inorganic chemistry,

    therefore, starts with a review of the models of the atom and a survey of

    the probability models applications to the electron con gurations of atoms

    and ions.

  • CHAPTER 1 The Electronic Structure of the Atom: A Review2

    corresponding to the Rydberg formula with nf 5 1. Then in 1908, Friedrich Paschen discovered a series of far-infrared hydrogen lines, tting the equation with nf 5 3.

    In 1913, Niels Bohr, a Danish physicist, became aware of Balmers and Rydbergs experimental work and of the Rydberg formula. At that time, he was trying to combine Ernest Rutherfords planetary model for electrons in an atom with Max Plancks quantum theory of energy exchanges. Bohr contended that an electron orbiting an atomic nucleus could only do so at certain xed distances and that whenever the electron moved from a higher to a lower orbit, the atom emitted characteristic electromagnetic radiation.

    Rydberg had deduced his equation from experimental observations of atomic hydrogen emission spectra. Bohr was able to derive the same equation from quantum theory, showing that his theoretical work meshed with reality. From this result, the Rutherford-Bohr model of the atom of concentric elec-tron shells was devised, mirroring the recurring patterns in the periodic table of the elements (Figure 1.1). Thus the whole concept of electron energy levels can be traced back to Rydberg. In recognition of Rydbergs contribution, excited atoms with very high values of the principal quantum number, n, are called Rydberg atoms.

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

    n 3

    n 2

    n 1

    ZeE hv

    FIGURE 1.1 The Rutherford-Bohr electron-shell model of the atom, showing the n 5 1, 2, and 3 energy levels.

    Aglowing body, such as the Sun, is expected to emit a continuous spectrum of electromagnetic radiation. However, in the early nineteenth century, a German sci-entist, Josef von Fraunhofer, noticed that the visible spec-trum from the Sun actually contained a number of dark bands. Later investigators realized that the bands were the result of the absorption of particular wavelengths by cooler atoms in the atmosphere above the surface of the Sun. The electrons of these atoms were in the ground state, and they were absorbing radiation at wavelengths corresponding to the energies needed to excite them to higher energy states. A study of these negative spectra led to the discovery of helium. Such spectral studies are still of great importance in cosmochemistrythe study of the chemical 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 presence of elements

    at very low concentrations. Each element has a particu-lar absorption spectrum corresponding to the various separations of (differences between) the energy levels in its atoms. When light from an atomic emission source is passed through a vaporized sample of an element, the particular wavelengths corresponding to the various en-ergy separations will be absorbed. We nd that the higher the concentration of the atoms, the greater the proportion of the light that will be absorbed. This linear relationship between light absorption and concentration is known as Beers law. The sensitivity of this method is extremely high, and concentrations of parts per million are easy to determine; some elements can be detected at the parts per billion level. Atomic absorption spectroscopy has now become a routine analytical tool in chemistry, metal-lurgy, geology, medicine, forensic science, and many other elds of scienceand it simply requires the movement of electrons from one energy level to another.

    Atomic Absorption Spectroscopy

  • 3

    1.1 The Schrdinger Wave Equation and Its Signi cance

    The more sophisticated quantum mechanical model of atomic structure was derived from the work of Louis de Broglie. De Broglie showed that, just as elec-tromagnetic waves could be treated as streams of particles (photons), moving particles could exhibit wavelike properties. Thus, it was equally valid to picture electrons either as particles or as waves. Using this wave-particle duality, Erwin Schrdinger developed a partial differential equation to represent the behavior of an electron around an atomic nucleus. One form of this equation, given here for a one-electron atom, shows the relationship between the wave function of the electron, C, and E and V, the total and potential energies of the system, re-spectively. The second differential terms relate to the wave function along each of the Cartesian coordinates x, y, and z, while m is the mass of an electron, and h is Plancks constant.

    020x2

    1020y2

    1020z2

    18p2m

    h21E 2 V2 5 0

    The derivation of this equation and the method of solving it are in the realm of physics and physical chemistry, but the solution itself is of great importance to inorganic chemists. We should always keep in mind, however, that the wave equation is simply a mathematical formula. We attach meanings to the solution simply because most people need concrete images to think about subatomic phenomena. The conceptual models that we create in our macroscopic world cannot hope to reproduce the subatomic reality.

    It was contended that the real meaning of the equation could be found from the square of the wave function, C2, which represents the probability of nding the electron at any point in the region surrounding the nucleus. There are a number of solutions to a wave equation. Each solution describes a different orbital and, hence, a different probability distribution for an elec-tron in that orbital. Each of these orbitals is uniquely de ned by a 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 theory, a fourth quantum number had to be de ned to explain the results of an experi-ment in 1922. In this experiment, Otto Stern and Walther Gerlach found that passing a beam of silver atoms through a magnetic eld caused about half the atoms to be de ected in one direction and the other half in the opposite direc-tion. Other investigators proposed that the observation was the result of two different electronic spin orientations. The atoms possessing an electron with one spin were de ected one way, and the atoms whose electron had the oppo-site spin were de ected in the opposite direction. This spin quantum number was assigned the symbol ms.

    The possible values of the quantum numbers are de ned as follows:

    n, the principal quantum number, can have all positive integer values from 1 to q.

    1.1 The Schrdinger Wave Equation and Its Signi cance

  • CHAPTER 1 The Electronic Structure of the Atom: A Review4

    l, the angular momentum quantum number, can have all integer values from n 2 1 to 0.

    ml, the magnetic quantum number, can have all integer values from 1lthrough 0 to 2l.

    ms, the spin quantum number, can have values of 112 and 2

    12.

    When the value of the principal quantum number is 1, there is only one possible set of quantum numbers n, l, and ml (1, 0, 0), whereas for a principal quantum number of 2, there are four sets of quantum numbers (2, 0, 0; 2, 1, 1; 2, 1, 0; 2, 1, 11). This situation is shown diagrammatically in Figure 1.2. To identify the electron orbital that corresponds to each set of quantum numbers, we use the value of the principal quantum number n, followed by a letter for the angular momentum quantum number l. Thus, when n 5 1, there is only the 1s orbital.

    When n 5 2, there is one 2s orbital and three 2p orbitals (corresponding to the ml values of 11, 0, and 1). The letters s, p, d, and f are derived from categories of the spectral lines: sharp, principal, diffuse, and fundamental. The correspondences are shown in Table 1.1.

    When the principal quantum number n 5 3, there are nine sets of quantum numbers (Figure 1.3). These sets correspond to one 3s, three 3p, and ve 3dorbitals. A similar diagram for the principal quantum number n 5 4 would show 16 sets of quantum numbers, corresponding to one 4s, three 4p, ve 4d,

    FIGURE 1.2 The possible sets of quantum numbers for n 5 1 and n 5 2.

    n

    l

    ml

    1s 2s 2p

    1 100

    0 1

    1

    0

    0

    2

    TABLE 1.1 Correspondence between angular momentum number l and orbital designation

    l Value Orbital designation

    0 s

    1 p

    2 d

    3 f

  • 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 display the complete solution for each orbital. As a realistic alternative, we break the wave equation into two parts: a radial part and an angular part.

    Each of the three quantum numbers derived from the wave equation rep-resents 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 the orbital.

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

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

    It is the value of the principal quantum number and, to a lesser extent the angular momentum quantum number, which determines the energy of the electron. Although the electron may not literally be spinning, it behaves as if it was, and it has the magnetic properties expected for a spinning particle.

    An orbital diagram is used to indicate the probability of nding an electron at any point in space. We de ne a location where an electron is most probably

    TABLE 1.2 Correspondence between angular momentum number l and number of orbitals

    l Value Number of orbitals

    0 1

    1 3

    2 5

    3 7

    FIGURE 1.3 The possible sets of quantum numbers for n 5 3.n

    l

    ml

    0 1

    3

    0 11 2 1 1 20 0

    2

    3s 3p 3d

    and seven 4f orbitals (Table 1.2). Theoretically, we can go on and on, but as we will see, the f orbitals represent the limit of orbital types among the elements of the periodic table for atoms in their electronic ground states.

    5

  • CHAPTER 1 The Electronic Structure of the Atom: A Review6

    found as an area of high electron density. Conversely, locations with a low prob-ability are called areas of low electron density.

    The s OrbitalsThe s orbitals are spherically symmetric about the atomic nucleus. As the prin-cipal quantum number increases, the electron tends to be found farther from the nucleus. To express this idea in a different way, we say that, as the principal quantum number increases, the orbital becomes more diffuse. A unique fea-ture of electron behavior in an s orbital is that there is a nite probability of nding the electron close to, and even within, the nucleus. This penetration by s orbital electrons plays a role in atomic radii (see Chapter 2) and as a means of studying nuclear structure.

    Same-scale representations of the shapes (angular functions) of the 1s and 2s orbitals of an atom are compared in Figure 1.4. 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 at the center of the spheres. These spheres represent the region in which there is a 99 percent probability of nding an electron. The total prob-ability cannot be represented, for the probability of nding an electron drops to zero only at an in nite distance from the nucleus.

    The probability of nding the electron within an orbital will always be posi-tive (since the probability is derived from the square of the wave function and squaring a negative makes a positive). However, when we discuss the bonding of atoms, we nd that the sign related to the original wave function has impor-tance. For this reason, it is conventional to superimpose the sign of the wave function on the representation of each atomic orbital. For an s orbital, the sign is positive.

    In addition to the considerable difference in size between the 1s and the 2sorbitals, the 2s orbital has, at a certain distance from the nucleus, a spherical surface 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 principal quantum number increases by 1, the number of nodal surfaces also increases by 1. We can visualize nodal surfaces more clearly by plotting a graph of the ra-dial density distribution function as a function of distance from the nucleus for any direction. Figure 1.5 shows plots for the 1s, 2s, and 3s orbitals. These plots show that the electron tends to be farther from the nucleus as the principal quantum number increases. The areas under all three curves are the same.

    FIGURE 1.4 Representations of the shapes and comparative sizes of the 1s and 2s orbitals.

  • 7

    Electrons in an s orbital are different from those in p, d, or f orbitals in two signi cant ways. First, only the s orbital has an electron density that varies in the same way in every direction out from the atomic nucleus. Second, there is a nite probability that an electron in an s orbital is at the nucleus of the atom. Every other orbital has a node at the nucleus.

    The p OrbitalsUnlike the s orbitals, the p orbitals consist of two separate volumes of space (lobes), with the nucleus located between the two lobes. Because there are three p orbitals, we assign each orbital a direction according to Cartesian co-ordinates: we have px, py, and pz. Figure 1.6 shows representations of the three 2p orbitals. At right angles to the axis of higher probability, there is a nodal plane through the nucleus. For example, the 2pz orbital has a nodal surface in the xy plane. In terms of wave function sign, one lobe is positive and the other negative.

    FIGURE 1.5 The variation of the radial density distribution function with distance from the nucleus for electrons in the 1s, 2s, and 3s orbitals of a hydrogen atom.

    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

    1.2 Shapes of the Atomic Orbitals

    FIGURE 1.6 Representations of the shapes of the 2px, 2py, and 2pzorbitals.

  • CHAPTER 1 The Electronic Structure of the Atom: A Review8

    If we compare graphs of electron density as a function of atomic radius for the 2s orbital and a 2p orbital (the latter plotted along the axis of higher prob-ability), we nd that the 2s orbital has a much greater electron density close to the nucleus than does the 2p orbital (Figure 1.7). Conversely, the second maximum of the 2s orbital is farther out than the single maximum of the 2porbital. However, the mean distance of maximum probability is the same for both orbitals.

    Like the s orbitals, the p orbitals develop additional nodal surfaces within the orbital structure as the principal quantum number increases. Thus, a 3porbital does not look exactly like a 2p orbital since it has an additional nodal surface. However, the detailed differences in orbital shapes for a particular angular momentum quantum number are of little relevance in the context of basic inorganic chemistry.

    The d OrbitalsThe ve d orbitals have more complex shapes. Three of them are located between the Cartesian axes, and the other two are oriented along the axes. In all cases, the nucleus is located at the intersection of the axes. Three orbitals each have four lobes that are located between pairs of axes (Figure 1.8). These orbitals are identi ed as dxy, dxz, and dyz. The other two d orbitals, dz2 and dx22y2, are shown in Figure 1.9. The dz2 orbital looks somewhat similar to a pz orbital (see Figure 1.6), except that it has an additional doughnut-shaped ring of high electron density in the xy plane. The dx22y2 orbital is identical to the dxy orbital but has been rotated through 458.

    FIGURE 1.7 The variation of the radial density distribution function with distance from the nucleus for electrons in the 2s and 2p orbitals of a hydrogen 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

    FIGURE 1.8 Representations of the shapes of the 3dxy, 3dxz, and 3dyz orbitals.

  • 9

    The f OrbitalsThe f orbitals are even more complex than the d orbitals. There are seven f orbitals, four of which have eight lobes. The other three look like the dz2orbital but have two doughnut-shaped rings instead of one. These orbitals are rarely 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 among the 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 the ground state, they occupy the orbitals of lowest energy, thereby minimizing the atoms total electronic energy. Thus, the con guration of an atom can be described simply by adding electrons one by one until the total number required for the element has been reached.

    Before starting to construct electron con gurations, we need to take into account a second rule: the Pauli exclusion principle. According to this rule, no two electrons in an atom may possess identical sets of the four quantum num-bers. Thus, there can be only one orbital of each three-quantum-number set per atom and each orbital can hold only two electrons, one with ms5 1

    12 and

    the other with ms5 212.

    Filling the s OrbitalsThe simplest con guration is that of the hydrogen atom. According to the Aufbau principle, the single electron will be located in the 1s orbital. This con- guration is the ground state of the hydrogen atom. Adding energy would raise the electron to one of the many higher energy states. These con gurations are referred to as excited states. In the diagram of the ground state of the hydro-gen atom (Figure 1.10), a half-headed arrow is used to indicate the direction of electron spin. The electron con guration is written as 1s1, with the superscript 1 indicating the number of electrons in that orbital.

    1.3 The Polyelectronic Atom

    1s

    FIGURE 1.10 Electron con guration of a hydrogen atom.

    FIGURE 1.9 Representations of the shapes of the 3dx22y2 and 3dz2 orbitals.

  • CHAPTER 1 The Electronic Structure of the Atom: A Review10

    With a two-electron atom (helium), there is a choice: the second electron could go in the 1s orbital (Figure 1.11a) or the next higher energy orbital, the 2sorbital (Figure 1.11b). Although it might seem obvious that the second electron would enter the 1s orbital, it is not so simple. If the second electron entered the 1s orbital, it would be occupying the same volume of space as the electron al-ready in that orbital. The very strong electrostatic repulsions, the pairing energy, would discourage the occupancy of the same orbital. However, by occupying an orbital with a high probability closer to the nucleus, the second electron will experience a much greater nuclear attraction. The nuclear attraction is greater than the inter-electron repulsion. Hence, the actual con guration will be 1s2,although it must be emphasized that electrons pair up in the same orbital only when pairing is the lower energy option.

    In the lithium atom the 1s orbital is lled by two electrons, and the third electron must be in the next higher energy orbital, the 2s orbital. Thus, lithium has the con guration of 1s22s1. Because the energy separation of an s and its corresponding p orbitals is always greater than the pairing energy in a poly-electronic atom, the electron con guration of beryllium will be 1s22s2 rather than 1s22s12p1.

    Filling the p OrbitalsBoron marks the beginning of the lling of the 2p orbitals. A boron atom has an electron con guration of 1s22s22p1. Because the p orbitals are degenerate (that is, they all have the same energy), it is impossible to decide which one of the three orbitals contains the electron.

    Carbon is the second ground-state atom with electrons in the p orbitals. Its electron con guration provides another challenge. There are three pos-sible arrangements of the two 2p electrons (Figure 1.12): (a) both electrons in one orbital, (b) two electrons with parallel spins in different orbitals, and (c) two electrons with opposed spins in different orbitals. On the basis of elec-tron repulsions, the rst possibility (a) can be rejected immediately. The deci-sion between the other two possibilities is less obvious and requires a deeper knowledge of quantum theory. In fact, if the two electrons have parallel spins, there is a zero probability of their occupying the same space. However, if the spins are opposed, there is a nite possibility that the two electrons will occupy the same region in space, thereby resulting in some repulsion and a higher energy state. Hence, the parallel spin situation (b) will have the lowest energy. This preference for unpaired electrons with parallel spins has been formalized in Hunds rule: When lling a set of degenerate orbitals, the num-ber of unpaired electrons will be maximized and these electrons will have parallel spins.

    After the completion of the 2p electron set at neon (1s22s22p6), the 3s and 3porbitals start to ll. Rather than write the full electron con gurations, a short-ened form can be used. In this notation, the inner electrons are represented by the noble gas symbol having that con guration. Thus, magnesium, whose full electron con guration would be written as 1s22s22p63s2, can be represented as having a neon noble gas core, and its con guration is written as [Ne]3s2. An

    FIGURE 1.11 Two possible electron con gurations for helium.

    (a) (b)

    2s2s

    1s1s

    2p

    (a)

    2p

    (b)

    2p

    (c)

    FIGURE 1.12 Possible 2pelectron con gurations for carbon.

  • 11

    advantage of the noble gas core representation is that it emphasizes the outer-most (valence) electrons, and it is these electrons that are involved in chemical bonding. Then lling the 3p orbitals brings us to argon.

    Filling the d OrbitalsIt is at this point that the 3d and 4s orbitals start to ll. The simple orbital energy level concept breaks down because the energy levels of the 4s and 3dorbitals are very close. What becomes most important is not the minimum energy for a single electron but the con guration that results in the least number of inter-electron repulsions for all the electrons. For potassium, this is [Ar]4s1; for calcium, [Ar]4s2.

    In general, the lowest overall energy for each transition metal is obtained by lling the s orbitals rst; the remaining electrons then occupy the d orbitals. Although there are minor uctuations in con gurations throughout the d-blockand 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

    Figure 1.13 shows the elements organized by order of orbital lling.This order is shown as an energy-level diagram in Figure 1.14 (page 13).

    The orbitals ll in this order because the energy differences between the s, p, d,and f orbitals of the same principal quantum number become so great beyond n 5 2 that they overlap with the orbitals of the following principal quantum numbers. It is important to note that Figure 1.14 shows the lling order, not the order for any particular element. For example, for elements beyond zinc, electrons in the 3d orbitals are far lower in energy than those in the 4s orbitals. Thus, at this point, the 3d orbitals have become inner orbitals and have no role in chemical bonding. Hence, their precise ordering is unimportant.

    Although these are the generalized rules, to illustrate how this delicate balance changes with increasing numbers of protons and electrons, the outer electrons in each of the Group 3 to Group 12 elements are listed here. These con gurations are not important in themselves, but they do show how close the ns and (n 1)d electrons are in energy.

    1.3 The Polyelectronic Atom

    Atom Con guration Atom Con guration Atom Con guration

    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

  • d-Block

    La

    Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

    Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb

    f-Block

    Sc Ti V Cr Mn Fe Co ZnNi

    Lr Rf Db Sg Bh Hs Mt Ds Rg Uub

    Lu Hf Ta W Re Os Pt Au Hg

    Y Zr Nb Mo Tc Ru Pd Ag Cd

    Cu

    Ir

    Rh

    p-Block

    s-Block

    Li

    Na

    CaK

    Rb

    Cs

    Fr Ra

    Ba

    Sr

    Mg

    Be

    H He

    KrBrSeAsGeGa

    Al

    NeFONCB

    ArClSPSi

    UuoUuhUupUuqUut

    Tl Pb Bi Rn

    Xe

    AtPo

    In Sn ITeSb

    FIGURE 1.13 In this version of the periodic table, the chemical elements are organized in order of orbital lling.

    12

  • 13

    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

    1

    FIGURE 1.14 Representation of the comparative energies of the atomic orbitals for lling order purposes.

    For certain elements, the lowest energy is obtained by shifting one or both of the s electrons to d orbitals. Looking at the rst series in isolation would lead to the conclusion that there is some preference for a half-full or full set of d orbitals by chromium and copper. However, it is more accurate to say that the inter-electron repulsion between the two s electrons is suf cient in several cases to result in an s1 con guration.

    For the elements from lanthanum (La) to ytterbium (Yb), the situation is even more uid because the 6s, 5d, and 4f orbitals all have similar energies. For example, lanthanum has a con guration of [Xe]6s25d1, whereas the next ele-ment, cerium, has a con guration of [Xe]6s24f 2. The most interesting electron con guration in this row is that of gadolinium, [Xe]6s25d14f 7, rather than the predicted [Xe]6s24f 8. This con guration provides more evidence of the impor-tance of inter-electron repulsion in the determination of electron con guration

    1.3 The Polyelectronic Atom

  • CHAPTER 1 The Electronic Structure of the Atom: A Review14

    when adjacent orbitals have similar energies. Similar complexities occur among the elements from actinium (Ac) to nobelium (No), in which the 7s, 6d, and 5forbitals have similar energies.

    1.4 Ion Electron Con gurations

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

    Atom Electron con guration Ion Electron con guration

    Na [Ne]3s1 Na1 [Ne]

    Mg [Ne]3s2 Mg21 [Ne]

    Al [Ne]3s23p1 Al31 [Ne]

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

    Atom Electron con guration Ion Electron con guration

    N [He]2s22p3 N32 [Ne]

    O [He]2s22p4 O22 [Ne]

    F [He]2s22p5 F2 [Ne]

    Some of the later main group metals form two ions with different charges. For example, lead forms Pb21 and (rarely) Pb41. The 21 charge can be explained by the loss of the 6p electrons only (the inert pair effect, which we discuss in Chapter 9, Section 9.8), whereas the 41 charge results from loss of both 6sand 6p electrons:

    Atom Electron con guration Ion Electron con guration

    Pb [Xe]6s24f 145d106p2 Pb21 [Xe]6s24f 145d10

    Pb41 [Xe]4f 145d10

    Notice that the electrons of the higher principal quantum number are lost rst. This rule is found to be true for all the elements. For the transition metals, the selectrons are always lost rst when a metal cation is formed. In other words, for the transition metal cations, the 3d orbitals are always lower in energy than the 4s orbitals, and a charge of 21, representing the loss of the two s electrons, is common for the transition metals and the Group 12 metals. For example, zinc always forms an ion of 21 charge:

    Atom Electron con guration Ion Electron con guration

    Zn [Ar]4s23d10 Zn21 [Ar]3d10

  • 15

    Iron forms ions with charges of 21 and 31 and, as shown here, it is tempting to ascribe the formation of the 31 ion to a process in which inter-electron repul-sion forces out the only paired d electron:

    1.5 Magnetic Properties of Atoms

    Atom Electron con guration Ion Electron con guration

    Fe [Ar]4s23d6 Fe21 [Ar]3d6

    Fe31 [Ar]3d5

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

    Atom Electron con guration Ion Electron con guration

    Ni [Ar]4s23d8 Ni21 [Ar]3d8

    Pd [Kr]5s04d10 Pd21, Pd41 [Kr]4d8, [Kr]4d6

    Pt [Xe]6s15d9 Pt21, Pt41 [Xe]5d8, [Xe]5d6

    1.5 Magnetic Properties of Atoms

    In the discussions of electron con guration, we saw that some atoms possess unpaired electrons. The presence of unpaired electrons in the atoms of an ele-ment can be determined easily from the elements magnetic properties. If atoms containing only spin-paired electrons are placed in a magnetic eld, they are weakly repelled by the eld. This phenomenon is called diamagnetism. Con-versely, atoms containing one or more unpaired electrons are attracted by the magnetic eld. This behavior of unpaired electrons is named paramagnetism.The attraction of each unpaired electron is many times stronger than the repul-sion of all the spin-paired electrons in that atom.

    To explain paramagnetism in simple terms, we can visualize the electron as a particle spinning on its axis and generating a magnetic moment, just as an electric current owing through a wire does. This permanent magnetic mo-ment results in an attraction into the stronger part of the eld. When electrons have their spins paired, the magnetic moments cancel each other. As a result, the paired electrons are weakly repelled by the lines of force of the magnetic eld. In paramagnetic materials, application of a magnetic eld aligns some of the normally randomly oriented electron spins with the applied magnetic eld (Figure 1.15a and b). It is this alignment that results in the attraction of the material into the magnetic eld. We will encounter this phenomenon again in our discussions of covalent bonding and the bonding in transition metal compounds.

  • CHAPTER 1 The Electronic Structure of the Atom: A Review16

    1.6 Medicinal Inorganic Chemistry: An Introduction

    Inorganic chemistry affects our lives directly in two ways. First, as we discuss at the end of later chapters, many chemical elements are required for the func-tioning of living organisms. Second, inorganic elements and compounds have been used as medicines since earliest times. Periodically in this text, we give examples of the use of inorganic compounds as medicinal substances, but it is useful to provide an overview.

    Many inorganic compounds have been used as medicines through the ages. A fashionable habit in European countries was to drink the waters at spa cities. In some cases, the springs were mineral-rich; for example, the water in Vichy, France (now available bottled), is rich in magnesium ion, which acts as a potent laxative. That water, therefore, should only be drunk in small quantities. The solid salt, magnesium sulfate heptahydrate, MgSO4.7H2O, has the same effect. It was named Epsom salts after the town in England where it was rst discovered. During the nineteenth century, one British hospital was using 2.5 tonnes per year on its patients!

    Some cultures practice geophagy, the eating of soilusually clay. Clays are a complex class of minerals, as we discuss in Chapter 14. One form of clay is kaolina substance that is known for its absorptive abilities. Several types of tablets to combat stomach upsets employ kaolin, which, it is believed, can surface-absorb toxins produced by ingested harmful bacteria. Other clays and soils can supply trace nutrients. However, persistent clay eating is not advised since the clay can line the stomach and prevent nutrient adsorption. Also, many natural clays contain high concentrations of harmful elements, such as lead.

    Inorganic medicinal chemistry can appear in the most unusual contexts. For example, religious statues made from the mineral realgar, diarsenic disul de (As2S2), were popular among devotees of the Chinese Taoist religion. Han-dling the statues was believed to restore health. In this particular case, chemis-try rather than faith might have contributed, for many people in tropical areas suffer from internal parasites and handling the statues would result in arsenic absorption through the skin, enough to kill the parasites but not enough to kill the devotee.

    FIGURE 1.15 The behavior of paramagnetic materials without (a) and with (b) an applied magnetic eld.

    (a)(b)

    H

  • 17

    In the following chapters, we mention a few of the many modern medicinal applications of inorganic compounds:

    Antacids (Chapter 7)

    Lithium in the treatment of bipolar disorder (Chapter 9)

    Boron neutron capture therapy (Chapter 13)

    Platinum complexes as anticancer agents (Chapter 19)

    Technetium as a radiopharmaceutical (Chapter 21)

    Gold in the treatment of rheumatoid arthritis (Chapter 21)

    KEY IDEAS

    The properties of an electron in an atom can be de ned in terms of four quantum numbers.

    There are a variety of shapes of orbitals (s, p, d, f ) as de ned by the angular momentum quantum number.

    Electrons in the ground state of an atom ll the orbitals of lowest energy.

    For the transition metals, the energies of the ns and (n 1)d orbitals are very similar.

    In the formation of a cation, the electrons in the orbitals of highest principal quantum number are lost rst.

    Paramagnetic behavior in a magnetic eld indicates the presence of unpaired electrons.

    EXERCISES

    1.1 De ne the following terms: (a) nodal surface; (b) Pauli exclusion principle; (c) paramagnetic.

    1.2 De ne the following terms: (a) orbital; (b) degenerate; (c) Hunds rule.

    1.3 Construct a quantum number tree for the principal quantum number n 5 4 similar to that depicted for n 5 3 in Figure 1.3.

    1.4 Determine the lowest value of n for which ml can (theoretically) have a value of 14.

    1.5 Identify the orbital that has n 5 5 and l 5 1.

    1.6 Identify the orbital that has n 5 6 and l 5 0.

    1.7 How does the quantum number n relate to the properties of an orbital?

    1.8 How does the quantum number l relate to the properties of an orbital?

    1.9 Explain concisely why carbon has two electrons in different p orbitals with parallel spins rather than the other possible arrangements.

    1.10 Explain concisely why beryllium has a ground-state electron con guration of 1s22s2 rather than 1s22s12p1.

    1.11 Write noble gas core ground-state electron con g-urations for atoms of (a) sodium; (b) nickel; (c) copper.

    1.12 Write noble gas core ground-state electron con g-urations for atoms of (a) calcium; (b) chromium; (c) lead.

    1.13 Write noble gas core ground-state electron con gu-rations for ions of (a) potassium; (b) scandium 31; (c) cop-per 21.

    1.14 Write noble gas core ground-state electron con gu-rations for ions of (a) chlorine; (b) cobalt 21; (c) manga-nese 41.

    1.15 Predict the common charges of the ions of thallium. Explain your reasoning in terms of electron con gurations.

    1.16 Predict the common charges of the ions of tin. Explain your reasoning in terms of electron con gurations.

    1.17 Predict the common charge of the silver ion. Explain your reasoning in terms of electron con gurations.

    1.18 Predict the highest possible charge of a zirconium ion. Explain your reasoning in terms of electron con gurations.

    1.19 Use diagrams similar to Figure 1.12 to determine the number of unpaired electrons in atoms of (a) oxygen; (b) magnesium; (c) chromium.

    Exercises

  • CHAPTER 1 The Electronic Structure of the Atom: A Review18

    BEYOND THE BASICS

    1.23 The next set of orbitals after the f orbitals are the gorbitals. How many g orbitals would there be? What would be the lowest principal quantum number n that would pos-sess g orbitals? Deduce the atomic number of the rst ele-ment at which g orbitals would begin to be lled on the basis of the patterns of the d and f orbitals.

    1.24 An alternative to the Schrdinger wave equation is the Dirac wave equation. Using online sources, research the Dirac wave equation and contrast it with the Schrdinger wave equation.

    1.25 Use an advanced inorganic chemistry text as a source of information on the f orbitals. What are their com-mon features? How do they differ among themselves?

    1.26 In Section 1.3, gadolinium is mentioned as having an electron con guration that deviates from the lanthanoid pattern. Which element in the actinoids should show a simi-lar deviation? What would be its electron con guration?

    1.27 In Figure 1.13, the elements are organized logically according to the order of orbital lling. Identify two disad-vantages of organizing the elements in this way.

    1.28 A philosophical question: Does an orbital exist even if it does not contain an electron? Discuss.

    1.20 Use diagrams similar to Figure 1.12 to determine the number of unpaired electrons in atoms of (a) nitrogen; (b) silicon; (c) iron.

    1.21 Write the electron con guration expected for ele-ment 113 and the con gurations for the two cations that it is most likely to form.

    1.22 Which of the following species are hydrogen-like? (a) He1; (b) He; (c) Li1; (d) Li21.

    ADDITIONAL RESOURCES

    For answers to odd-numbered questions: www.whfreeman.com/descriptive5e

    For accompanying video clips: www.whfreeman.com/descriptive5e

  • The search for patterns among the chemical elements really started with the work of the German chemist Johann Dbereiner. In 1829, he noticed that there were similarities in properties among various groups of three elements, such as calcium, strontium, and barium. He named these groups triads. Then in 1865, John Newlands, a British sugar re ner, realized that, when the elements were placed in order of increasing atomic weights, a cycle of properties repeated with every eight elements. Newlands called this pattern the law of octaves. At the time, scientists had started to look for a unity of the physical laws that would explain everything, so to correlate element organization with the musical scale seemed natural. Unfortunately, his proposal was laughed at by most chemists of the time.

    It was in 1869 that the Russian chemist Dmitri Mendeleev (pro-nounced Men-de-l-ev) independently devised the same concept (without linking it to music) and made the crucial advance of using the law as a predictive tool. It attracted little attention until four months later, when Lothar Meyer, a German chemist, published his own report on the periodic relationship. Meyer did acknowledge that Mendeleev had the same idea rst.

    Mendeleev and Meyer would hardly recognize the contemporary periodic table. In Mendeleevs proposal, the elements known at the time were organized in an eight-column format in order of increasing atomic mass. He claimed that each eighth element had similar proper-ties. Groups I to VII each contained two subgroups, and Group VIII contained four subgroups. The organization of one of his designs is shown in Figure 2.1. To ensure that the patterns in the properties of ele-ments t the table, it was necessary to leave spaces. Mendeleev assumed that these spaces corresponded to unknown elements. He argued that the properties of the missing elements could be predicted on the basis of the chemistry of each elements neighbors in the same group. For

    2.1 Organization of the Modern Periodic Table

    2.2 Existence of the Elements

    2.3 Stability of the Elements and Their Isotopes

    The Origin of the Shell Model of the Nucleus

    2.4 Classi cations of the Elements

    2.5 Periodic Properties: Atomic Radius

    2.6 Periodic Properties: Ionization Energy

    2.7 Periodic Properties: Electron Af nity

    Alkali Metal Anions

    2.8 The Elements of Life

    An Overview of the Periodic Table

    CHAPTER 2

    19

    The periodic table is the framework on which much of our understanding

    of inorganic chemistry is based. In this chapter, we provide the essential

    information that you will need for the more detailed discussions of the

    individual groups in later chapters.

  • CHAPTER 2 An Overview of the Periodic Table 20

    example, the missing element between silicon and tin, called eka-silicon (Es) by Mendeleev, should have properties intermediate between those of silicon and tin. Table 2.1 compares Mendeleevs predictions with the properties of germa-nium, discovered 15 years later.

    However, the Mendeleev periodic table had three major problems:

    1. If the order of increasing atomic mass was consistently followed, ele-ments did not always t in the group that had the matching properties. Thus, the order of nickel and cobalt had to be reversed, as did that of iodine and tellurium.

    2. Elements were being discovered, such as holmium and samarium, for which no space could be found. This dif culty was a particular embarrassment.

    3. Elements in the same group were sometimes quite different in their chemical reactivity. This discrepancy was particularly true of the rst group, which con-tained the very reactive alkali metals and the very unreactive coinage metals (copper, silver, and gold).

    As we now know, there was another aw: to establish a group of elements, at least one element has to be known already. Because none of the noble gases was known at that time, no space was left for them. Conversely, some spaces in Mendeleevs table were completely erroneous. This was because he tried to t the elements into repeating rows (periods) of eight. Now, of course, we know that the periods are not consistently eight members long but instead increase regularly: successive rows have 2, 8, 8, 18, 18, 32, and 32 elements.

    H

    Na

    Cu?

    Ag?

    Au?

    Li

    K

    Rb

    Cs

    Be

    Ca

    Sr

    Ba

    Mg

    Zn

    Cd

    Hg

    B

    Yt

    Er

    Al

    In

    Tl

    C

    Ti

    Zr

    Ce

    La

    Th

    Si

    Sn

    Pb

    N

    V

    Nb

    Ta

    P

    As

    Sb

    Bi

    O

    Cr

    Mo

    W

    U

    S

    Se

    Te

    F

    MnCl

    Br

    I

    Fe Ni Co Cu

    Os Pt Ir Au

    I II III IV V VI VII VIII

    Ru Pd Rh Ag

    FIGURE 2.1 The organization of one of Mendeleevs designs for the periodic table.

    TABLE 2.1 Comparison of Mendeleevs predictions for eka-silicon and the actual properties of germanium

    Element Atomic weight Density (g?cm23) Oxide formula Chloride formula

    Eka-silicon 72 5.5 EsO2 EsCl4Germanium 72.3 5.47 GeO2 GeCl4

  • 21

    The crucial transition to our modern ideas was provided by Henry Moseley, a British physicist, as we discuss in the dedication of this text. Placing the elements in order of the atomic number that he derived from spectroscopic measurements removed the irregularities of the table that was based on atomic masses, and it de ned exactly the spaces in the table where elements still needed to be found.

    2.1 Organization of the Modern Periodic Table

    In the modern periodic table, the elements are placed in order of increasing atomic number (the number of protons). There have been numerous designs of the table over the years, but the two most common are the long form and the short form. The long form (Figure 2.2) shows all the elements in numerical order.

    2.1 Organization of the Modern Periodic Table

    SrRb

    BaCs

    Lr RaFr Sg BhDbRf Hs Mt Ds Rg Uub Uut Uuq Uup Uuh Uuo

    Fe CoCr MnVTiSc

    Li

    Na

    K Ca

    Be

    Mg

    Mo TcNbZrY Ru Rh Sb TeSnInCdAgPd

    Bi PoPbTl At Rn

    Br Kr

    I Xe

    As SeGeGaZnNi Cu

    P SSiAl

    F Ne

    Cl Ar

    N OCB

    HeH

    Nd PmPrCeLa Sm Eu Tm YbErHoDyTbGd

    Ac U NpPaTh Pu Am CfBkCm Md NoFmEs

    Lu W ReTaHf Os Ir HgAuPt

    FIGURE 2.2 The long form of the periodic table.

    Main2s3s4s5s6s7s

    4f5f

    LanthanoidsActinoids

    3d4d5d6d

    5p6p7p

    2p3p4p

    Main

    Transition

    1s1s FIGURE 2.3 Electron orbital lling sequence in the periodic table.

    SciAmThe start of a new period always corresponds to the introduction of the rst electron into the s orbital of a new principal quantum number. In a par-ticular period, the principal quantum number of the p orbitals is the same as that of the s orbitals, whereas the d orbitals are one less and the f orbitals are two less. The number of elements in each period corresponds to the number of electrons required to ll those orbitals (Figure 2.3).

    Each group contains elements of similar electron con guration. For exam-ple, all Group 1 elements have an outer electron that is ns1, where n is the prin-cipal quantum number. Although elements in a group have similar properties, it is important to realize that every element is unique. Thus, although nitrogen and phosphorus are sequential elements in the same group, nitrogen gas is very unreactive and phosphorus is so reactive that it spontaneously reacts with the oxygen in the air.

  • CHAPTER 2 An Overview of the Periodic Table 22

    Because the long form of the periodic table is a very elongated diagram and because the elements from lanthanum to ytterbium and from actinium to nobelium show similar chemical behavior, the short form displays these two sets of elements in rows beneath the remainder of the table and the resulting space is closed up. Figure 2.4 shows this more compact, short form of the table. Chemists disagree about the choice of elements to be placed in Group 3 for Periods 6 and 7. Some consider lanthanum (element 57) and actinium (ele-ment 89) to be the correct choices, while others believe lutetium (element 71) and lawrencium (element 103) belong there. The electron con gurations of the contenders are shown below:

    Atom Electron con guration Atom Electron con guration

    La [Xe]6s25d1 Lu [Xe]6s24f145d1

    Ac [Rn]7s26d1 Lr [Rn]7s25f146d1

    Both arguments have their merits. The La-Ac supporters argue that the con g-urations match those of the elements above: scandium ([Ar]4s23d1) and yttrium ([Kr]5s24d1). The Lu-Lr supporters point out that all the other transition metals of Period 6 have 4f14 in their con gurations and those of Period 7 have 5f14 in their con gurations, making the Lu-Lr pair more consistent.

    According to the recommendations of the International Union of Pure and Applied Chemistry (IUPAC), the main and transition groups of elements are numbered from 1 to 18. This system replaces the old system of using a mixture of Roman numerals and letters, a notation that caused confusion because of differences in numbering between North America and the rest of the world. For example, in North America, IIIB referred to the group containing scandium, whereas in the rest of the world this designation was used for the group start-ing with boron. Numerical designations are not used for the series of elements from lanthanum (La) to ytterbium (Yb) and from actinium (Ac) to nobelium

    Mo TcNbZrYSrRb

    BaCs

    Fr

    W ReTaHf

    Sg BhDbRf

    Os Ir

    Hs Mt Ds Rg Uub Uut Uuq UupUuh Uuo

    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

    HeH1 2

    3 4 5 6 7 8 9 10 11 12

    13 14 15 16 17

    18

    *

    ** Lr

    Lu

    Np PuPa UTh NoFm MdEsCfBkAm Cm

    Lanthanoids

    Actinoids Ac

    PmPr NdCeLa Sm YbEr TmHoDyTbEu Gd*

    **

    Ra

    FIGURE 2.4 Short form of the periodic table displaying the group numbers. Main group elements are blue.

  • 23

    (No), because there is much more resemblance in properties within each of those rows of elements than vertically in groups.

    Groups 1 and 2 and 13 through 18 represent the main group elements, and these groups correspond to the lling of the s and p orbitals. Groups 4 through 11, corresponding to the lling of the d orbitals, are classi ed as the transition metals. The discussion of the main groups will take up the majority of space in this text because it is these elements that cover the widest range of chemi-cal and physical properties. The elements of Group 12, although sometimes included among the transition metals, have a very different chemistry from that series; hence, Group 12 will be considered separately. Several of the main groups have been given speci c names: alkali metals (Group 1), alkaline earth metals (Group 2), pnictogens (a lesser-used term for Group 15), chalcogens (a lesser-used term for Group 16), halogens (Group 17), and noble gases (Group 18). The elements in Group 11 are sometimes called the coinage metals.

    The elements corresponding to the lling of the 4f orbitals are called the lanthanoids, and those corresponding to the lling of the 5f orbitals are called the actinoids. They used to be named the lanthanides and actinides, but the -ide ending more correctly means a negative ion, such as oxide or chloride. For a few years, IUPAC was suggesting the names lanthanons and actinons, but because the ending -on is preferred for nonmetals (and the lanthanoids and actinoids are all metallic elements), the -oid ending is now recommended. The chemistry of the elements of Group 3, scandium (Sc), yttrium (Y), and lutetium (Lu), more closely resembles that of the lanthanoids than that of the transition metals. For this reason, these three elements are usually discussed together with the lanthanoid elements, lanthanum to ytterbium (see Chapter 24). There is, in fact, a collective name for the Group 3 and lanthanoid elements: the rare earth elements.

    Although the elements in the periodic table are arranged according to elec-tron structure, we make an exception for helium (1s2). Rather than placing it with the other ns2 con guration elements, the alkaline earth metals, it is placed with the other noble gases (of con guration ns2np6) because of chemical similarities (see Figure 2.3). Hydrogen is even more of a problem. Although some versions of the periodic table show it as a member of Group 1 or Group 17 or both, its chemistry is unlike either that of the alkali metals or the halogens. For this reason, it is placed on its own in the tables in this text to indicate its uniqueness.

    2.2 Existence of the Elements

    To understand why there are so many elements and to explain the pattern of the abundances of the elements, we must look at the most widely accepted theory of the origin of the universe. This is the big bang theory, which assumes that the uni-verse started from a single point. About one second after the universe came into existence, the temperature had dropped to about 1010 K, enabling protons and neutrons to exist. During the next three minutes, hydrogen-1, hydrogen-2, helium-3, helium-4, beryllium-7, and lithium-7 nuclei formed. (The num