Hauptseite Organic Chemistry: Structure and Function
Es kann für Sie interessant sein Powered by Rec2Me
Am meisten angefragte Begriffe
Hello. I am Zarif Samarov I work as chemistry teacher
27 August 2020 (11:46)
Hello. I am SunYanchun I work as Organic synthesis researcher in China
05 September 2020 (16:54)
Hello. I am Galbadrakh Tsogbadrakh. I am work as chemistry teacher.
11 October 2020 (17:19)
Damn you niggas have ugly ass names lmao
19 May 2021 (07:47)
looks a very good book. let me read more of it before I say how great it is.
19 May 2021 (19:12)
About the Authors K. PETER C. VOLLHARDT was born in Madrid, raised in Buenos Aires and Munich, studied at the University of Munich, received his Ph.D. with Professor Peter Garratt at the University College, London, and was a postdoctoral fellow with Professor Bob Bergman (then) at the California Institute of Technology. He moved to Berkeley in 1974 when he began his efforts toward the development of organocobalt reagents in organic synthesis, the preparation of theoretically interesting hydrocarbons, the assembly of novel transition metal arrays with potential in catalysis, and the discovery of a parking space. Among other pleasant experiences, he was a Studienstiftler, Adolf Windaus medalist, Humboldt Senior Scientist, ACS Organometallic Awardee, Otto Bayer Prize Awardee, A. C. Cope Scholar, Japan Society for the Promotion of Science Prize Holder, and recipient of the Medal of the University Aix-Marseille and an Honorary Doctorate from The University of Rome Tor Vergata. He is an editor of Synlett. Among his more than 350 publications, he treasures especially this textbook in organic chemistry, translated into 13 languages. Peter is married to Marie-José Sat, a French artist, and they have three children, Maïa (b. 1982, Peter’s stepdaughter), whose splendid tattoo you can admire on p. 1067, Paloma (b. 1994), and Julien (b. 1997). NEIL E. SCHORE was born in Newark, New Jersey in 1948. His education took him through the public schools of the Bronx, New York, and Ridgefield, New Jersey, after which he completed a B.A. with honors in chemistry at the University of Pennsylvania in 1969. Moving back to New York, he worked with the late Professor Nicholas J. Turro at Columbia University, studying photochemical and photophysical processes of organic compounds for his Ph.D. thesis. He first met Peter Vollhardt when he and Peter were doing postdoctoral work in Professor Robert Bergman’s laboratory at Cal Tech in the 1970s. Since joining the U.C. Davis faculty in 1976, he has taught organic chemistry to some 20; ,000 nonchemistry majors, winning seven teaching awards, publishing over 100 papers in various areas related to organic chemistry, and refereeing several hundred local youth soccer games. He has also pioneered Study Abroad programs in Taiwan and the U.K. for chemistry students and is an Adjunct Professor in the Korea University International Summer Campus program. Neil is married to Carrie Erickson, a microbiologist at the U.C. Davis School of Veterinary Medicine. They have two children, Michael (b. 1981) and Stefanie (b. 1983), both of whom carried out experiments for this book. Grandson Roman (b. 2016) is a bit too young for that as yet. PETER VOLLHARDT University of California at Berkeley NEIL SCHORE University of California at Davis Vice President: Ben Roberts Program Manager: Beth Cole Development Editor: Randi Blatt Rossignol Marketing Manager: Maureen Rachford Marketing Assistant: Savannah DiMarco Director of Content: Kristen Ford Lead Content Developer: Lily Huang Assistant Editor: Allison Greco Director, Content Management Enhancement: Tracey Kuehn Managing Editor: Lisa Kinne Director of Design, Content Management: Diana Blume Photo Editor: Sheena Goldstein Photo Researcher: Lisa Passmore Design Director, Content Management: Diana Blume Media Project Manager: Daniel Comstock Senior Design Manager: Vicki Tomaselli Project Management and Composition: Aptara®, Inc. Illustrations: Network Graphics; Precision Graphics Senior Workflow Supervisor: Susan Wein Cover Image: Jürgen Müller/imageBROKER/Alamy Library of Congress Control Number: 2017950317 ISBN-13: 978-1-319-18896-2 (EPUB) © 2018, 2014, 2011, 2007 by W. H. Freeman and Company All rights reserved W. H. Freeman and Company One New York Plaza Suite 4500 New York, NY 10004-1562 www.macmillanlearning.com Brief Contents PREFACE: A User’s Guide to ORGANIC CHEMISTRY: Structure and Function 1 STRUCTURE AND BONDING IN ORGANIC MOLECULES 2 STRUCTURE AND REACTIVITY Acids and Bases, Polar and Nonpolar Molecules 3 REACTIONS OF ALKANES Bond-Dissociation Energies, Radical Halogenation, and Relative Reactivity 4 CYCLOALKANES 5 STEREOISOMERS 6 PROPERTIES AND REACTIONS OF HALOALKANES Bimolecular Nucleophilic Substitution 7 FURTHER REACTIONS OF HALOALKANES Unimolecular Substitution and Pathways of Elimination 8 HYDROXY FUNCTIONAL GROUP: ALCOHOLS Properties, Preparation, and Strategy of Synthesis 9 FURTHER REACTIONS OF ALCOHOLS AND THE CHEMISTRY OF ETHERS 10 USING NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY TO DEDUCE STRUCTURE 11 ALKENES: INFRARED SPECTROSCOPY AND MASS SPECTROMETRY 12 REACTIONS OF ALKENES 13 ALKYNES The Carbon–Carbon Triple Bond 14 DELOCALIZED Pi SYSTEMS Investigation by Ultraviolet and Visible Spectroscopy 15 BENZENE AND AROMATICITY Electrophilic Aromatic Substitution 16 ELECTROPHILIC ATTACK ON DERIVATIVES OF BENZENE Substituents Control Regioselectivity 17 ALDEHYDES AND KETONES The Carbonyl Group 18 ENOLS, ENOLATES, AND THE ALDOL CONDENSATION α,β-Unsaturated Aldehydes and Ketones 19 CARBOXYLIC ACIDS 20 CARBOXYLIC ACID DERIVATIVES 21 AMINES AND THEIR DERIVATIVES Functional Groups Containing Nitrogen 22 CHEMISTRY OF BENZENE SUBSTITUENTS Alkylbenzenes, Phenols, and Anilines 23 ESTER ENOLATES AND THE CLAISEN CONDENSATION Synthesis of β-Dicarbonyl Compounds; Acyl Anion Equivalents 24 CARBOHYDRATES Polyfunctional Compounds in Nature 25 HETEROCYCLES Heteroatoms in Cyclic Organic Compounds 26 AMINO ACIDS, PEPTIDES, PROTEINS, AND NUCLEIC ACIDS Nitrogen-Containing Polymers in Nature Answers to Exercises Index Contents PREFACE: A User’s Guide to ORGANIC CHEMISTRY: Structure and Function 1 STRUCTURE AND BONDING IN ORGANIC MOLECULES 1-1 1-2 1-3 1-4 1-5 1-6 1-7 1-8 1-9 1-10 2 The Scope of Organic Chemistry: An Overview Real Life: Nature 1-1 Urea: From Urine to Wöhler’s Synthesis to Industrial Fertilizer Coulomb Forces: A Simplified View of Bonding Ionic and Covalent Bonds: The Octet Rule Electron-Dot Model of Bonding: Lewis Structures Resonance Forms Atomic Orbitals: A Quantum Mechanical Description of Electrons Around the Nucleus Molecular Orbitals and Covalent Bonding Hybrid Orbitals: Bonding in Complex Molecules Structures and Formulas of Organic Molecules A General Strategy for Solving Problems in Organic Chemistry Worked Examples: Integrating the Concepts Important Concepts Problems STRUCTURE AND REACTIVITY Acids and Bases, Polar and Nonpolar Molecules 2-1 2-2 2-3 2-4 2-5 2-6 2-7 2-8 2-9 2-10 3 Kinetics and Thermodynamics of Simple Chemical Processes Keys to Success: Using Curved “Electron-Pushing” Arrows to Describe Chemical Reactions Acids and Bases Real Life: Medicine 2-1 Stomach Acid, Peptic Ulcers, Pharmacology, and Organic Chemistry Functional Groups: Centers of Reactivity Straight-Chain and Branched Alkanes Naming the Alkanes Structural and Physical Properties of Alkanes Real Life: Nature 2-2 “Sexual Swindle” by Means of Chemical Mimicry Rotation about Single Bonds: Conformations Rotation in Substituted Ethanes Worked Examples: Integrating the Concepts Important Concepts Problems REACTIONS OF ALKANES Bond-Dissociation Energies, Radical Halogenation, and Relative Reactivity 3-1 3-2 3-3 3-4 3-5 3-6 3-7 Strength of Alkane Bonds: Radicals Structure of Alkyl Radicals: Hyperconjugation Conversion of Petroleum: Pyrolysis Real Life: Sustainability 3-1 Sustainability and the Needs of the 21st Century: “Green” Chemistry Chlorination of Methane: The Radical Chain Mechanism Other Radical Halogenations of Methane Keys to Success: Using the “Known” Mechanism as a Model for the “Unknown” Chlorination of Higher Alkanes: Relative Reactivity and Selectivity 3-8 3-9 3-10 3-11 3-12 4 4-1 4-2 4-3 4-4 4-5 4-6 4-7 4-8 5 5-1 Selectivity in Radical Halogenation with Fluorine and Bromine Synthetic Radical Halogenation Real Life: Medicine 3-2 Chlorination, Chloral, and DDT: The Quest to Eradicate Malaria Synthetic Chlorine Compounds and the Stratospheric Ozone Layer Combustion and the Relative Stabilities of Alkanes Worked Examples: Integrating the Concepts Important Concepts Problems CYCLOALKANES Names and Physical Properties of Cycloalkanes Ring Strain and the Structure of Cycloalkanes Cyclohexane: A Strain-Free Cycloalkane Substituted Cyclohexanes Larger Cycloalkanes Polycyclic Alkanes Carbocyclic Products in Nature Real Life: Materials 4-1 Cyclohexane, Adamantane, and Diamandoids: Diamond “Molecules” Real Life: Medicine 4-2 Cholesterol: How Is It Bad and How Bad Is It? Real Life: Medicine 4-3 Controlling Fertility: From “the Pill” to RU-486 to Male Contraceptives Worked Examples: Integrating the Concepts Important Concepts Problems STEREOISOMERS Chiral Molecules Real Life: Nature 5-1 Chiral Substances in Nature 5-2 5-3 5-4 5-5 5-6 5-7 5-8 5-9 6 Optical Activity Absolute Configuration: R,S Sequence Rules Fischer Projections Molecules Incorporating Several Stereocenters: Diastereomers Real Life: Nature 5-2 Stereoisomers of Tartaric Acid Meso Compounds Stereochemistry in Chemical Reactions Real Life: Medicine 5-3 Chiral Drugs—Racemic or Enantiomerically Pure? Real Life: Medicine 5-4 Why Is Nature “Handed”? Resolution: Separation of Enantiomers Worked Examples: Integrating the Concepts Important Concepts Problems PROPERTIES AND REACTIONS OF HALOALKANES Bimolecular Nucleophilic Substitution 6-1 6-2 6-3 6-4 6-5 6-6 6-7 6-8 6-9 6-10 Physical Properties of Haloalkanes Real Life: Medicine 6-1 Fluorinated Pharmaceuticals Nucleophilic Substitution Reaction Mechanisms Involving Polar Functional Groups: Using “Electron-Pushing” Arrows A Closer Look at the Nucleophilic Substitution Mechanism: Kinetics Frontside or Backside Attack? Stereochemistry of the SN2 Reaction Consequences of Inversion in SN2 Reactions Structure and SN2 Reactivity: The Leaving Group Structure and SN2 Reactivity: The Nucleophile Keys to Success: Choosing Among Multiple Mechanistic Pathways Structure and SN2 Reactivity: The Substrate 6-11 6-12 7 The SN2 Reaction at a Glance Worked Examples: Integrating the Concepts Important Concepts Problems FURTHER REACTIONS OF HALOALKANES Unimolecular Substitution and Pathways of Elimination 7-1 7-2 7-3 7-4 7-5 7-6 7-7 7-8 7-9 7-10 8 Solvolysis of Tertiary and Secondary Haloalkanes Unimolecular Nucleophilic Substitution Stereochemical Consequences of SN1 Reactions Effects of Solvent, Leaving Group, and Nucleophile on Unimolecular Substitution Effect of the Alkyl Group on the SN1 Reaction: Carbocation Stability Real Life: Medicine 7-1 Unusually Stereoselective SN1 Displacement in Anticancer Drug Synthesis Unimolecular Elimination: E1 Bimolecular Elimination: E2 Keys to Success: Substitution versus Elimination—Structure Determines Function Summary of Reactivity of Haloalkanes Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems HYDROXY FUNCTIONAL GROUP: ALCOHOLS Properties, Preparation, and Strategy of Synthesis 8-1 8-2 8-3 8-4 8-5 8-6 8-7 8-8 8-9 9 Naming the Alcohols Structural and Physical Properties of Alcohols Alcohols as Acids and Bases Synthesis of Alcohols by Nucleophilic Substitution Synthesis of Alcohols: Oxidation–Reduction Relation Between Alcohols and Carbonyl Compounds Real Life: Medicine 8-1 Oxidation and Reduction in the Body Real Life: Medicine 8-2 Don’t Drink and Drive: The Breath Analyzer Test Organometallic Reagents: Sources of Nucleophilic Carbon for Alcohol Synthesis Organometallic Reagents in the Synthesis of Alcohols Keys to Success: An Introduction to Synthetic Strategy Real Life: Chemistry 8-3 What Magnesium Does Not Do, Copper Can: Alkylation of Organometallics Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems FURTHER REACTIONS OF ALCOHOLS AND THE CHEMISTRY OF ETHERS 9-1 9-2 9-3 9-4 9-5 9-6 9-7 9-8 Reactions of Alcohols with Base: Preparation of Alkoxides Reactions of Alcohols with Strong Acids: Alkyloxonium Ions in Substitution and Elimination Reactions of Alcohols Carbocation Rearrangements Esters from Alcohols and Haloalkane Synthesis Names and Physical Properties of Ethers Williamson Ether Synthesis Real Life: Nature 9-1 Chemiluminescence of 1,2Dioxacyclobutanes Synthesis of Ethers: Alcohols and Mineral Acids Reactions of Ethers Real Life: Medicine 9-2 Protecting Groups in the Synthesis of 9-9 9-10 9-11 9-12 10 Testosterone Reactions of Oxacyclopropanes Real Life: Chemistry 9-3 Hydrolytic Kinetic Resolution of Oxacyclopropanes Sulfur Analogs of Alcohols and Ethers Physiological Properties and Uses of Alcohols and Ethers Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems USING NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY TO DEDUCE STRUCTURE 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 Physical and Chemical Tests Defining Spectroscopy Hydrogen Nuclear Magnetic Resonance Real Life: Spectroscopy 10-1 Recording an NMR Spectrum Using NMR Spectra to Analyze Molecular Structure: The Proton Chemical Shift Tests for Chemical Equivalence Real Life: Medicine 10-2 Magnetic Resonance Imaging (MRI) in Medicine Integration of NMR Signals Spin–Spin Splitting: The Effect of Nonequivalent Neighboring Hydrogens Spin–Spin Splitting: Some Complications Real Life: Spectroscopy 10-3 The Nonequivalence of Diastereotopic Hydrogens Carbon-13 Nuclear Magnetic Resonance Real Life: Spectroscopy 10-4 How to Determine Atom Connectivity in NMR Real Life: Medicine 10-5 Structural Characterization of Natural and “Unnatural” Products: An Antioxidant from Grape Seeds and a Fake Drug in Herbal Medicines 10-10 11 Worked Examples: Integrating the Concepts Important Concepts Problems ALKENES: INFRARED SPECTROSCOPY AND MASS SPECTROMETRY 11-1 11-2 11-3 11-4 11-5 11-6 11-7 11-8 11-9 11-10 11-11 11-12 12 Naming the Alkenes Structure and Bonding in Ethene: The Pi Bond Physical Properties of Alkenes Nuclear Magnetic Resonance of Alkenes Real Life: Medicine 11-1 NMR of Complex Molecules: The Powerfully Regulating Prostaglandins Catalytic Hydrogenation of Alkenes: Relative Stability of Double Bonds Preparation of Alkenes from Haloalkanes and Alkyl Sulfonates: Bimolecular Elimination Revisited Preparation of Alkenes by Dehydration of Alcohols Infrared Spectroscopy Measuring the Molecular Mass of Organic Compounds: Mass Spectrometry Real Life: Medicine 11-2 Detecting Performance-Enhancing Drugs Using Mass Spectrometry Fragmentation Patterns of Organic Molecules Degree of Unsaturation: Another Aid to Identifying Molecular Structure Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems REACTIONS OF ALKENES 12-1 12-2 12-3 12-4 12-5 12-6 12-7 12-8 12-9 12-10 12-11 12-12 12-13 12-14 12-15 12-16 12-17 12-18 13 Why Addition Reactions Proceed: Thermodynamic Feasibility Catalytic Hydrogenation Basic and Nucleophilic Character of the Pi Bond: Electrophilic Addition of Hydrogen Halides Alcohol Synthesis by Electrophilic Hydration: Thermodynamic Control Electrophilic Addition of Halogens to Alkenes The Generality of Electrophilic Addition Oxymercuration–Demercuration: A Special Electrophilic Addition Real Life: Medicine 12-1 Juvenile Hormone Analogs in the Battle Against Insect-Borne Diseases Hydroboration–Oxidation: A Stereospecific Anti-Markovnikov Hydration Diazomethane, Carbenes, and Cyclopropane Synthesis Oxacyclopropane (Epoxide) Synthesis: Epoxidation by Peroxycarboxylic Acids Vicinal Syn Dihydroxylation with Osmium Tetroxide Real Life: Medicine 12-2 Synthesis of Antitumor Drugs: Sharpless Enantioselective Oxacyclopropanation (Epoxidation) and Dihydroxylation Oxidative Cleavage: Ozonolysis Radical Additions: Anti-Markovnikov Product Formation Dimerization, Oligomerization, and Polymerization of Alkenes Synthesis of Polymers Ethene: An Important Industrial Feedstock Alkenes in Nature: Insect Pheromones Real Life: Medicine 12-3 Alkene Metathesis Transposes the Termini of Two Alkenes: Construction of Rings Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems ALKYNES The Carbon–Carbon Triple Bond 13-1 13-2 13-3 13-4 13-5 13-6 13-7 13-8 13-9 13-10 13-11 13-12 14 Naming the Alkynes Properties and Bonding in the Alkynes Spectroscopy of the Alkynes Preparation of Alkynes by Double Elimination Preparation of Alkynes from Alkynyl Anions Reduction of Alkynes: The Relative Reactivity of the Two Pi Bonds Electrophilic Addition Reactions of Alkynes Anti-Markovnikov Additions to Triple Bonds Chemistry of Alkenyl Halides Real Life: Synthesis 13-1 Metal-Catalyzed Stille, Suzuki, and Sonogashira Coupling Reactions Ethyne as an Industrial Starting Material Alkynes in Nature and in Medicine Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems DELOCALIZED Pi SYSTEMS Investigation by Ultraviolet and Visible Spectroscopy 14-1 14-2 14-3 14-4 14-5 14-6 Overlap of Three Adjacent p Orbitals: Electron Delocalization in the 2-Propenyl (Allyl) System Radical Allylic Halogenation Nucleophilic Substitution of Allylic Halides: SN1 and SN2 Allylic Organometallic Reagents: Useful Three-Carbon Nucleophiles Two Neighboring Double Bonds: Conjugated Dienes Electrophilic Attack on Conjugated Dienes: Kinetic and Thermodynamic Control 14-7 14-8 14-9 14-10 14-11 14-12 Delocalization Among More than Two Pi Bonds: Extended Conjugation and Benzene A Special Transformation of Conjugated Dienes: Diels-Alder Cycloaddition Real Life: Materials 14-1 Organic Polyenes Conduct Electricity Real Life: Sustainability 14-2 The Diels-Alder Reaction is “Green” Electrocyclic Reactions Polymerization of Conjugated Dienes: Rubber Electronic Spectra: Ultraviolet and Visible Spectroscopy Real Life: Spectroscopy 14-3 The Contributions of IR, MS, and UV to the Characterization of Viniferone Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems INTERLUDE A Summary of Organic Reaction Mechanisms 15 BENZENE AND AROMATICITY Electrophilic Aromatic Substitution 15-1 15-2 15-3 15-4 15-5 15-6 15-7 15-8 Naming the Benzenes Structure and Resonance Energy of Benzene: A First Look at Aromaticity Pi Molecular Orbitals of Benzene Spectral Characteristics of the Benzene Ring Real Life: Materials 15-1 Compounds Made of Pure Carbon: Graphite, Graphene, Diamond, and Fullerenes Polycyclic Aromatic Hydrocarbons Other Cyclic Polyenes: Hückel’s Rule Hückel’s Rule and Charged Molecules Synthesis of Benzene Derivatives: Electrophilic Aromatic 15-9 15-10 15-11 15-12 15-13 15-14 16 Substitution Halogenation of Benzene: The Need for a Catalyst Nitration and Sulfonation of Benzene Friedel-Crafts Alkylation Limitations of Friedel-Crafts Alkylations Friedel-Crafts Acylation Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems ELECTROPHILIC ATTACK ON DERIVATIVES OF BENZENE Substituents Control Regioselectivity 16-1 16-2 16-3 16-4 16-5 16-6 16-7 16-8 17 Activation or Deactivation by Substituents on a Benzene Ring Directing Electron-Donating Effects of Alkyl Groups Directing Effects of Substituents in Conjugation with the Benzene Ring Real Life: Materials 16-1 Explosive Nitroarenes: TNT and Picric Acid Electrophilic Attack on Disubstituted Benzenes Keys to Success: Synthetic Strategies Toward Substituted Benzenes Reactivity of Polycyclic Benzenoid Hydrocarbons Polycyclic Aromatic Hydrocarbons and Cancer Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems ALDEHYDES AND KETONES The Carbonyl Group 17-1 17-2 17-3 17-4 17-5 17-6 17-7 17-8 17-9 17-10 17-11 17-12 17-13 17-14 17-15 18 Naming the Aldehydes and Ketones Structure of the Carbonyl Group Spectroscopic Properties of Aldehydes and Ketones Preparation of Aldehydes and Ketones Reactivity of the Carbonyl Group: Mechanisms of Addition Addition of Water to Form Hydrates Addition of Alcohols to Form Hemiacetals and Acetals Acetals as Protecting Groups Nucleophilic Addition of Ammonia and Its Derivatives Real Life: Biochemistry 17-1 Imines Mediate the Biochemistry of Amino Acids Deoxygenation of the Carbonyl Group Addition of Hydrogen Cyanide to Give Cyanohydrins Addition of Phosphorus Ylides: The Wittig Reaction Oxidation by Peroxycarboxylic Acids: The Baeyer-Villiger Oxidation Oxidative Chemical Tests for Aldehydes Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems ENOLS, ENOLATES, AND THE ALDOL CONDENSATION α,β -Unsaturated Aldehydes and Ketones 18-1 18-2 18-3 18-4 18-5 Acidity of Aldehydes and Ketones: Enolate Ions Keto–Enol Equilibria Halogenation of Aldehydes and Ketones Alkylation of Aldehydes and Ketones Attack by Enolates on the Carbonyl Function: Aldol Condensation Real Life: Biology And Medicine 18-1 Stereoselective Aldol 18-6 18-7 18-8 18-9 18-10 18-11 18-12 19 19-1 19-2 19-3 19-4 19-5 19-6 19-7 19-8 19-9 19-10 19-11 19-12 19-13 Reactions in Nature and in the Laboratory: “Organocatalysis” Crossed Aldol Condensation Keys to Success: Competitive Reaction Pathways and the Intramolecular Aldol Condensation Real Life: Nature 18-2 Absorption of Photons by Unsaturated Aldehydes Enables Vision Properties of α,β -Unsaturated Aldehydes and Ketones Conjugate Additions to α,β -Unsaturated Aldehydes and Ketones 1,2- and 1,4-Additions of Organometallic Reagents Conjugate Additions of Enolate Ions: Michael Addition and Robinson Annulation Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems CARBOXYLIC ACIDS Naming the Carboxylic Acids Structural and Physical Properties of Carboxylic Acids Spectroscopy and Mass Spectrometry of Carboxylic Acids Acidic and Basic Character of Carboxylic Acids Carboxylic Acid Synthesis in Industry Methods for Introducing the Carboxy Functional Group Substitution at the Carboxy Carbon: The Addition–Elimination Mechanism Carboxylic Acid Derivatives: Acyl Halides and Anhydrides Carboxylic Acid Derivatives: Esters Carboxylic Acid Derivatives: Amides Reduction of Carboxylic Acids by Lithium Aluminum Hydride Bromination Next to the Carboxy Group: The Hell-VolhardZelinsky Reaction Biological Activity of Carboxylic Acids Real Life: Materials 19-1 Long-Chain Carboxylates and 19-14 20 20-1 20-2 20-3 20-4 20-5 20-6 20-7 20-8 20-9 21 Sulfonates Make Soaps and Detergents Real Life: Health 19-2 Artery-Clogging Trans Fatty Acids Phasing Out Real Life: Materials 19-3 Green Plastics, Fibers, and Energy from Biomass-Derived Hydroxyesters Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems CARBOXYLIC ACID DERIVATIVES Relative Reactivities, Structures, and Spectra of Carboxylic Acid Derivatives Chemistry of Acyl Halides Chemistry of Carboxylic Anhydrides Chemistry of Esters Esters in Nature: Waxes, Fats, Oils, and Lipids Real Life: Sustainability 20-1 Moving Away from Petroleum: Green Fuels from Vegetable Oil Amides: The Least Reactive Carboxylic Acid Derivatives Real Life: Medicine 20-2 Killing the Bugs that Kill the Drugs: Antibiotic Wars Amidates and Their Halogenation: The Hofmann Rearrangement Alkanenitriles: A Special Class of Carboxylic Acid Derivatives Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems AMINES AND THEIR DERIVATIVES Functional Groups Containing Nitrogen 21-1 21-2 21-3 21-4 21-5 21-6 21-7 21-8 21-9 21-10 21-11 22 Naming the Amines Real Life: Medicine 21-1 Physiologically Active Amines and Weight Control Structural and Physical Properties of Amines Spectroscopy of the Amine Group Acidity and Basicity of Amines Synthesis of Amines by Alkylation Synthesis of Amines by Reductive Amination Synthesis of Amines from Carboxylic Amides Reactions of Quaternary Ammonium Salts: Hofmann Elimination Mannich Reaction: Alkylation of Enols by Iminium Ions Nitrosation of Amines Real Life: Medicine 21-2 Sodium Nitrite as a Food Additive, N-Nitrosodialkanamines, and Cancer Real Life: Materials 21-3 Amines in Industry: Nylon, the “Miracle Fiber” Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems CHEMISTRY OF BENZENE SUBSTITUENTS Alkylbenzenes, Phenols, and Anilines 22-1 22-2 22-3 22-4 22-5 22-6 Reactivity at the Phenylmethyl (Benzyl) Carbon: Benzylic Resonance Stabilization Oxidations and Reductions of Substituted Benzenes Names and Properties of Phenols Real Life: Medicine 22-1 Two Phenols in the News: Bisphenol A and Resveratrol Preparation of Phenols: Nucleophilic Aromatic Substitution Alcohol Chemistry of Phenols Real Life: Medicine 22-2 Aspirin: The Miracle Drug Electrophilic Substitution of Phenols 22-7 22-8 22-9 22-10 22-11 22-12 23 An Electrocyclic Reaction of the Benzene Ring: The Claisen Rearrangement Oxidation of Phenols: Benzoquinones Real Life: Biology 22-3 Chemical Warfare in Nature: The Bombardier Beetle Oxidation-Reduction Processes in Nature Arenediazonium Salts Electrophilic Substitution with Arenediazonium Salts: Diazo Coupling Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems ESTER ENOLATES AND THE CLAISEN CONDENSATION Synthesis of β-Dicarbonyl Compounds; Acyl Anion Equivalents 23-1 23-2 23-3 23-4 23-5 β -Dicarbonyl Compounds: Claisen Condensations Real Life: Nature 23-1 Claisen Condensations Assemble Biological Molecules β -Dicarbonyl Compounds as Synthetic Intermediates β -Dicarbonyl Anion Chemistry: Michael Additions Acyl Anion Equivalents: Preparation of α -Hydroxyketones Real Life: Nature 23-2 Thiamine: A Natural, Metabolically Active Thiazolium Salt Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems 24 CARBOHYDRATES Polyfunctional Compounds in Nature 24-1 24-2 24-3 24-4 24-5 24-6 24-7 24-8 24-9 24-10 24-11 24-12 24-13 25 Names and Structures of Carbohydrates Conformations and Cyclic Forms of Sugars Anomers of Simple Sugars: Mutarotation of Glucose Polyfunctional Chemistry of Sugars: Oxidation to Carboxylic Acids Oxidative Cleavage of Sugars Reduction of Monosaccharides to Alditols Carbonyl Condensations with Amine Derivatives Ester and Ether Formation: Glycosides Step-by-Step Buildup and Degradation of Sugars Real Life: Nature 24-1 Biological Sugar Synthesis Relative Configurations of the Aldoses: An Exercise in Structure Determination Complex Sugars in Nature: Disaccharides Real Life: Food Chemistry 24-2 Manipulating Our Sweet Tooth Polysaccharides and Other Sugars in Nature Real Life: Medicine 24-3 Sialic Acid, “Bird Flu,” and Rational Drug Design Worked Example: Integrating the Concepts New Reactions Important Concepts Problems HETEROCYCLES Heteroatoms in Cyclic Organic Compounds 25-1 25-2 Naming the Heterocycles Nonaromatic Heterocycles Real Life: Medicine 25-1 Smoking, Nicotine, Cancer, and Medicinal Chemistry 25-3 25-4 25-5 25-6 25-7 25-8 25-9 26 Structures and Properties of Aromatic Heterocyclopentadienes Reactions of the Aromatic Heterocyclopentadienes Structure and Preparation of Pyridine: An Azabenzene Reactions of Pyridine Real Life: Biochemistry 25-2 Lessons from Redox-Active Pyridinium Salts in Nature: Nicotinamide Adenine Dinucleotide, Dihydropyridines, and Synthesis Quinoline and Isoquinoline: The Benzopyridines Real Life: Biology 25-3 Folic Acid, Vitamin D, Cholesterol, and the Color of Your Skin Alkaloids: Physiologically Potent Nitrogen Heterocycles in Nature Real Life: Nature 25-4 Nature is Not Always Green: Natural Pesticides Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems AMINO ACIDS, PEPTIDES, PROTEINS, AND NUCLEIC ACIDS Nitrogen-Containing Polymers in Nature 26-1 26-2 26-3 26-4 26-5 26-6 Structure and Properties of Amino Acids Real Life: Medicine 26-1 Arginine and Nitric Oxide in Biochemistry and Medicine Synthesis of Amino Acids: A Combination of Amine and Carboxylic Acid Chemistry Synthesis of Enantiomerically Pure Amino Acids Real Life: Chemistry 26-2 Enantioselective Synthesis of Optically Pure Amino Acids: Phase-Transfer Catalysis Peptides and Proteins: Amino Acid Oligomers and Polymers Determination of Primary Structure: Amino Acid Sequencing Synthesis of Polypeptides: A Challenge in the Application of Protecting Groups 26-7 26-8 26-9 26-10 26-11 26-12 Merrifield Solid-Phase Peptide Synthesis Polypeptides in Nature: Oxygen Transport by the Proteins Myoglobin and Hemoglobin Biosynthesis of Proteins: Nucleic Acids Real Life: Medicine 26-3 Synthetic Nucleic Acid Bases and Nucleosides in Medicine Protein Synthesis Through RNA DNA Sequencing and Synthesis: Cornerstones of Gene Technology Real Life: Forensics 26-4 DNA Fingerprinting Worked Examples: Integrating the Concepts New Reactions Important Concepts Problems Answers to Exercises Index Preface A User’s Guide to ORGANIC CHEMISTRY: Structure and Function In the eighth edition of Organic Chemistry: Structure and Function, we maintain the logical framework for understanding contemporary organic chemistry that is a hallmark of the text, while also strengthening the focus on helping students understand reactions, mechanisms, and synthetic analysis, as well as their practical applications. The classic framework of the text emphasizes that the structure of an organic molecule determines how that molecule functions, be it with respect to its physical behavior or in a chemical reaction. Extensive revisions to the eighth edition build on this framework by presenting a refined methodology, rooted in teaching expertise, to promote student understanding, build problem-solving skills, and provide applications of organic chemistry in the life and material sciences. Organic Chemistry: Structure and Function is also offered in a cost-effective SaplingPlus format as an interactive and fully mobile e-Book supported by interactive media features and the well-respected Sapling Learning tutorial style problems. The key innovations of the eighth edition as well as the hallmark features of the text are illustrated in the following pages. Accessible Organization within a Classic Framework Our framework emphasizes that the structure (electronic and spatial) of an organic molecule determines how that molecule functions—both its physical behavior and in a chemical reaction. Emphasizing this connection in early chapters allows students to build a true grasp of reaction mechanisms, encouraging understanding over memorization. NEW chapter-opening Learning Objectives and chapter-ending Summaries provide a framework for students and set the goals for the chapter. Electron-pushing arrows are introduced in early sections and then reinforced throughout the text to help students follow the movement of electrons and atoms in the reaction. Focus on Reaction Mechanisms Animations of key reactions allow students to visualize mechanisms and their dynamics by providing a moving image. NEW PowerPoint slides for the classroom include embedded animations along with questions appropriate for in-class work or open response systems. NEW annotations on key mechanistic illustrations recall and reinforce the key tenets of a mechanism. Keys to Success sections teach and reinforce conceptual understanding of mechanisms. A few examples include: Chapter 3, Section 3-6: Using the “Known” Mechanism as a Model for the “Unknown” Chapter 6, Section 6-9: Choosing Among Multiple Mechanistic Pathways Chapter 7, Section 7-8: Substitution versus Elimination—Structure Determines Function Chapter 8, Section 8-8: An Introduction to Synthetic Strategy Interlude: A Summary of Organic Reaction Mechanisms (following Chapter 14) reviews the key mechanisms that drive the majority of organic reactions, encouraging understanding over memorization. Reaction Summary Road Maps provide one-page overviews of the reactivity of each major functional group. The preparation maps indicate the possible origins of a functionality—that is, the precursor functional groups. The reaction maps show what each functional group does. Section numbers direct students to the corresponding discussion in the text. Problem Solving Skills and Strategy Expanded WHIP problem-solving sections are included with solved exercises within each chapter, providing strategies for how students should approach problem solving: What does the problem ask and what information is provided? How to begin? Information needed? Proceed, step-by-step, and do not skip any steps. In-text problems allow students to practice applications of new material, guide them to the solutions, and teach them how to recognize the fundamental types of questions they are likely to encounter. NEW Guidelines provide a blueprint of the steps that will help students apply concepts and organize their approach to problem solving. One example is Guidelines: IUPAC Rules for Naming Straight-Chain Alkanes (p. 80). NEW. Numerous marginal “Remember:” entries highlighting common pitfalls that students encounter, particularly in the formulation of mechanisms. ChemCasts replicate the face-to-face experience of watching an instructor work a problem. Using a virtual whiteboard, the Organic ChemCast tutors show students the steps involved in solving key Worked Examples, while explaining the concepts along the way. The Worked Examples featured in the ChemCasts were chosen with the input of organic chemistry students. End-of-chapter problems include a range of difficulty levels and a variety of practical applications: Worked Examples: Integrating the Concepts with worked-out, stepby-step solutions to problems involving several concepts from within chapters and from among several chapters Team Problems to encourage discussion and collaborative learning among students Pre-Professional Problems offering practice for students in solving problems similar to those they will encounter on the MCAT, especially those that test for scientific inquiry, reasoning, and research. Practical Applications and Visualization Every chapter features discussions of biological, medical, and industrial applications of organic chemistry, many of them new to this edition. Really? margin entries highlight unusual and surprising aspects of organic chemistry concepts intended to stimulate students’ engagement. Real Life boxes describe real chemistry by practicing chemists. Medicinal chemistry fundamentals in over 70 entries introduce drug design, absorption, metabolism, mode of action, and medicinal terminology. Animations help students visualize the most important structures and concepts in organic chemistry. New numerous potential-energy diagrams have been added to increase the visual understanding of reaction energetics. NEW AND UPDATED TOPICS As with all new editions, each chapter has been carefully reviewed, revised, and expanded. New Relevant portions of the 2013 IUPAC recommendations for naming organic molecules have been implemented. For cyclic hydrocarbons the ring is now the stem, regardless of the length of an alkyl substituent (unsaturated or not). For acyclic unsaturated hydrocarbons, the longest chain rules, regardless of the presence or location of double or triple bonds. Thioethers are named alkylthioalkanes, in analogy to alkoxyalkanes for ethers. New Over 220 new in-chapter problems have been added to aid students in practicing and applying new concepts as they learn new material. New entries, updates, and improvements include: Expanded and improved coverage of acid-base chemistry (Chapter 2) Improved presentation of the factors determining bond strengths (Chapters 2, 3, 6, 11, and 20) Improved presentation of leaving group ability (Chapter 6) Numerous new examples of reactions and expanded presentation of mechanisms (Chapters 6–26) Updated coverage of the ozone layer (Chapter 3) Improved and expanded discussion of the substitution and elimination reactions of alkyloxonium ions (Chapter 9) Updated mechanism of the Wittig reaction (Chapter 17) Expanded coverage of enolates (Chapter 18) Expanded mechanistic coverage of carboxylic acid derivatives (Chapters 19 and 20) Expanded coverage of amine synthesis (Chapter 21) Expanded mechanistic coverage of nitrosation (Chapter 21) Expanded coverage of oxidations and reductions of substituted benzenes: the Birch reduction (Chapter 22) Completely revised and updated “Top Ten” Drug List (Chapter 25) Revised and expanded coverage of heteroaromatic compounds (Chapter 25) Resources in Every Problem Counts This comprehensive and robust online platform combines innovative high quality teaching and learning features with Sapling Learning’s acclaimed online chemistry homework. Sapling Online Homework Proven effective in raising students’ comprehension and problem-solving skills and recently updated with hundreds of additional questions. A special effort has been made to increase the number of spectroscopy, synthesis, and mechanism problems. Sapling Learning’s innovative online homework offers: Immediate, Individualized Feedback Each student gets the guidance they need when they need it. Efficient Course Management Automatic grading, tracking, and analytics helps instructors save time and tailor assignments to student needs. Industry-Leading Peer-to-Peer Support Each instructor is paired with a fully trained PhD or master’s level colleague ready to help with everything from quizzes and assignments to syllabus planning and tech support. SaplingPlus for Vollhardt and Schore, Organic Chemistry, Eighth Edition, features: New! Interactive e-Book. Mobile accessible with native apps on Mac, Windows, iOS, Android, Chrome, and Kindle. Allows students to read offline, have the book read aloud to them, highlight, take notes, and search for keywords. Sapling Learning Problems. Tutorial style problems. Every problem emphasizes learning with hints, targeted feedback, and detailed solutions, as well as a unique pedagogically focused drawing tool. Sapling Learning for Vollhardt and Schore, Organic Chemistry includes problems taken from the text’s end-of-chapter sections, particularly questions focused on medical applications. Newly added spectroscopy, mechanistic, and synthesis problems. Animated Mechanisms allow students to visualize mechanisms by providing a moving image. NEW PowerPoint slides for the classroom include embedded animations as well as questions appropriate for in-class work or open response systems. Animations in Sapling Animations of processes (in addition to the animated mechanisms) help students visualize the most important structures and concepts in organic chemistry. ChemCasts replicate the face-to-face experience of watching an instructor work a problem. Using a virtual whiteboard, the Organic ChemCast tutors show students the steps involved in solving key Worked Examples, while explaining the concepts along the way. The Worked Examples featured in the ChemCasts were chosen with the input of organic chemistry students. Molecular Modeling Problems provide an additional bank of detailed problems intended for use with electronic structure software. Lecture PowerPoints, adapted from the lecture slides used in Peter Vollhardt’s course, serve as a framework for lecture outlines. Text art in PNG and PPT includes all figures from the text, optimized for use in large lecture halls. Additional Resources Study Guide and Solutions Manual, by Neil Schore, University of California at Davis ISBN: 978-1-319-19574-8 Written by Organic Chemistry coauthor Neil Schore, this invaluable manual includes chapter introductions that highlight new material, chapter outlines, detailed comments for each chapter section, a glossary, and solutions to the end-of-chapter problems, presented in a way that shows students how to reason their way to the answer. Molecular Model Set A modeling set offers a simple, practical way for students to see, manipulate, and investigate molecular behavior. Polyhedra mimic atoms, pegs serve as bonds, and oval discs become orbitals. Maruzen Company Basic Molecular Modeling Kit: ISBN: 978-1-319-12052-8. Darling Modeling Kit #3: ISBN: 978-1-319-08374-8 Acknowledgments We are grateful to the following professors who reviewed the manuscript for the eighth edition. Jung-Mo Ahn, University of Texas at Dallas Kim Albizati, University of California, San Diego Taro Amagata, San Francisco State University Shawn Amorde, Austin Community College Donal Aue, University of California, Santa Barbara David Baker, Delta College Koushik Banerjee, Georgia College and State University Francis Barrios, Bellarmine University Mikael Bergdahl, San Diego State University Thomas Bertolini, University of Southern California Kelvin Billingsley, San Francisco State University Richard Broene, Bowdoin College Corey Causey, University North Florida Steven Chung, Bowling Green State University Edward Clennan, University of Wyoming Oana Cojocaru, Tennessee Technological University Perry Corbin, Ashland University Lisa Crow, Southern Nazarene University Michael Danahy, Bowdoin College Patrick Donoghue, Appalachian State University Steven Farmer, Sonoma State University Balazs Hargittai, Saint Francis University Bruce Hathaway, LeTourneau University Sheng-Lin (Kevin) Huang, Azusa Pacific University John Jewett, University of Arizona Bob Kane, Baylor University Jeremy Klosterman, Bowling Green State University Brian Love, East Carolina University Philip Lukeman, St. John’s University Jordan Mader, Shepherd University Matt McIntosh, University of Arkansas Cheryl Moy, University of North Carolina Joseph Mullins, Le Moyne College Shaun Murphree, Allegheny College Jacqueline Nikles, University of Alabama at Birmingham Herman Odens, Southern Adventist University Jon Parquette, Ohio State University Bhavna Rawal, Houston Community College Kevin Shaughnessy, University of Alabama Nicholas Shaw, Appalachian State University Supriya Sihi, Houston Community College Melinda Stephens, Geneva College John Tovar, Johns Hopkins University Elizabeth Waters, University of North Carolina at Wilmington Haim Weizman, University of California, San Diego Patrick Willoughby, Ripon College We remain grateful to the professors who contributed in so many ways to the development of previous editions of Organic Chemistry. Marc Anderson, San Francisco State University George Bandik, University of Pittsburgh Anne Baranger, University of California, Berkeley Kevin Bartlett, Seattle Pacific University Scott Borella, University of North Carolina—Charlotte Stefan Bossmann, Kansas State University Alan Brown, Florida Institute of Technology Paul Carlier, Virginia Tech University Robert Carlson, University of Kansas Toby Chapman, University of Pittsburgh Robert Coleman, Ohio State University William Collins, Fort Lewis College Robert Corcoran, University of Wyoming Stephen Dimagno, University of Nebraska, Lincoln Rudi Fasan, University of Rochester James Fletcher, Creighton University Sara Fitzgerald, Bridgewater College Joseph Fox, University of Delaware Terrence Gavin, Iona College Joshua Goodman, University of Rochester Christopher Hadad, Ohio State University Ronald Halterman, University of Oklahoma Michelle Hamm, University of Richmond Kimi Hatton, George Mason University Sean Hightower, University of North Dakota Shawn Hitchcock, Illinois State University Stephen Hixson, University of Massachusetts, Amherst Danielle Jacobs, Rider University Ismail Kady, East Tennessee State University Rizalia Klausmeyer, Baylor University Krishna Kumar, Tufts University Julie Larson, Bemidji State University Scott Lewis, James Madison University Carl Lovely, University of Texas at Arlington Claudia Lucero, California State University—Sacramento Sarah Luesse, Southern Illinois University—Edwardsville John Macdonald, Worcester Polytechnical Institute Lisa Ann McElwee-White, University of Florida Linda Munchausen, Southeastern Louisiana State University Richard Nagorski, Illinois State University Liberty Pelter, Purdue University—Calumet Jason Pontrello, Brandeis University MaryAnn Robak, University of California, Berkeley Joseph Rugutt, Missouri State University—West Plains Kirk Schanze, University of Florida Pauline Schwartz, University of New Haven Trent Selby, Mississippi College Gloria Silva, Carnegie Mellon University Dennis Smith, Clemson University Leslie Sommerville, Fort Lewis College Jose Soria, Emory University Michael Squillacote, Auburn University Mark Steinmetz, Marquette University Jennifer Swift, Georgetown University James Thompson, Alabama A&M University Carl Wagner, Arizona State University James Wilson, University of Miami Alexander Wurthmann, University of Vermont Neal Zondlo, University of Delaware Eugene Zubarev, Rice University We are also grateful to the following professors who reviewed the manuscript for the sixth edition: Michael Barbush, Baker University Debbie J. Beard, Mississippi State University Robert Boikess, Rutgers University Cindy C. Browder, Northern Arizona University Kevin M. Bucholtz, Mercer University Kevin C. Cannon, Penn State Abington J. Michael Chong, University of Waterloo Jason Cross, Temple University Alison Flynn, Ottawa University Roberto R. Gil, Carnegie Mellon University Sukwon Hong, University of Florida Jeffrey Hugdahl, Mercer University Colleen Kelley, Pima Community College Vanessa McCaffrey, Albion College Keith T. Mead, Mississippi State University James A. Miranda, Sacramento State University David A. Modarelli, University of Akron Thomas W. Ott, Oakland University Hasan Palandoken, Western Kentucky University Gloria Silva, Carnegie Mellon University Barry B. Snider, Brandeis University David A. Spiegel, Yale University Paul G. Williard, Brown University Shmuel Zbaida, Rutgers University Eugene Zubarev, Rice University Peter Vollhardt thanks his colleagues at UC Berkeley, in particular Professors John Arnold, Anne Baranger, Bob Bergman, Ron Cohen, Felix Fischer, Matt Francis, John Hartwig, Darleane Hoffman, Tom Maimone, Richmond Sarpong, Rich Saykally, Andrew Streitwieser, and Dean Toste, for suggestions, updates, general discussions, and stimulus. He would also like to thank his administrative assistant, Bonnie Kirk, for helping with the logistics of producing and handling manuscript and galleys. Neil Schore thanks Dr. Melekeh Nasiri and Professor Mark Mascal for their ongoing comments and suggestions, and the numerous undergraduates at UC Davis who eagerly pointed out errors, omissions, and sections that could be improved or clarified. Our thanks go to the many people who helped with this edition. Beth Cole, acquisitions editor, and Randi Rossignol, development editor, at Macmillan Learning, guided this edition from concept to completion. Marketing Manager Maureen Rachford helped refine the story of the eighth edition and support our adopters. The team at Sapling Learning, led by Lily Huang and Stacy Benson, managed the media content with great skill and knowledge. The Sapling team included Sarah Egner, Rene Flores, Alexandra Gordon, Chris Knarr, Robley Light, Cheryl McCutchan, Heather Southerland, Thomas Turner, and Andrew Waldeck. Allison Greco, assistant editor, helped coordinate our efforts. Also many thanks to Sheena Goldstein, our photo editor, Vicki Tomaselli, our designer, and Susan Wein, Senior Workflow Supervisor, for their fine work and attention to the smallest detail. Thanks also to Dennis Free and Sherrill Redd at Aptara for their unlimited patience. Chapter 1 Structure and Bonding in Organic Molecules Tetrahedral carbon, the essence of organic chemistry, exists as a lattice of six-membered rings in diamonds. In 2003, a family of molecules called diamandoids was isolated from petroleum. Diamandoids are subunits of diamond in which the excised pieces are capped off with hydrogen atoms. An example is the beautifully crystalline pentamantane (molecular model on top right and picture on the left. which consists of five “cages” of the diamond lattice. The top right of the picture shows the carbon frame of pentamantane stripped of its hydrogens and its superposition on the lattice of diamond. Learning Objectives Relate your knowledge of beginning general chemistry to organic molecules: ionic and covalent bonding, shape, the octet rule, and Lewis structures Recognize the importance of Coulomb’s Law in organic chemistry Recognize the importance of the spreading out of electron density Relate the valence electron count to the stabilization of the elements through bond formation Learn to write resonance forms for structures that exhibit delocalization Review the orbital picture of electrons around the nucleus Apply hybridization to describe bonding in simple organic systems, such as methane Illustrate the drawing of three-dimensional structures of organic molecules How do chemicals regulate your body? Why did your muscles ache this morning after last night’s long jog? What is in the pill you took to get rid of that headache you got after studying all night? What happens to the gasoline you pour into the gas tank of your car? What is the molecular composition of the things you wear? What is the difference between a cotton shirt and one made of silk? What is the origin of the odor of garlic? You will find the answers to these questions, and many others that you may have asked yourself, in this book on organic chemistry. Chemistry is the study of the structure of molecules and the rules that govern their interactions. As such, it interfaces closely with the fields of biology, physics, and mathematics. What, then, is organic chemistry? What distinguishes it from other chemical disciplines, such as physical, inorganic, or nuclear chemistry? A common definition provides a partial answer: Organic chemistry is the chemistry of carbon and its compounds. These compounds are called organic molecules. Organic molecules constitute the chemical building blocks of life. Fats, sugars, proteins, and the nucleic acids are compounds in which the principal component is carbon. So are countless substances that we take for granted in everyday use. Virtually all the clothes that we wear are made of organic molecules—some of natural fibers, such as cotton and silk; others artificial, such as polyester. Toothbrushes, toothpaste, soaps, shampoos, deodorants, perfumes—all contain organic compounds, as do furniture, carpets, the plastic in light fixtures and cooking utensils, paintings, food, and countless other items. Consequently, organic chemical industries are among the largest in the world, including petroleum refining and processing, agrochemicals, plastics, pharmaceuticals, paints and coatings, and the food conglomerates. Organic substances such as gasoline, medicines, pesticides, and polymers have improved the quality of our lives. Yet the uncontrolled disposal of organic chemicals has polluted the environment, causing deterioration of animal and plant life as well as injury and disease to humans. If we are to create useful molecules—and learn to control their effects—we need a knowledge of their properties and an understanding of their behavior. We must be able to apply the principles of organic chemistry. This chapter explains how the basic ideas of chemical structure and bonding apply to organic molecules. Most of it is a review of topics that you covered in your general chemistry courses, including molecular bonds, Lewis structures and resonance, atomic and molecular orbitals, and the geometry around bonded atoms. 1-1 The Scope of Organic Chemistry: An Overview A goal of organic chemistry is to relate the structure of a molecule to the reactions that it can undergo. We can then study the steps by which each type of reaction takes place, and we can learn to create new molecules by applying those processes. Thus, it makes sense to classify organic molecules according to the subunits and bonds that determine their chemical reactivity: These determinants are groups of atoms called functional groups. The study of the various functional groups and their respective reactions provides the structure of this book. Functional groups determine the reactivity of organic molecules We begin with the alkanes, composed of only carbon and hydrogen atoms (“hydrocarbons”) connected by single bonds. They lack any functional groups and as such constitute the basic scaffold of organic molecules. As with each class of compounds, we present the systematic rules for naming alkanes, describe their structures, and examine their physical properties (Chapter 2). An example of an alkane is ethane. Its structural mobility is the starting point for a review of thermodynamics and kinetics. This review is then followed by a discussion of the strength of alkane bonds, which can be broken by heat, light, or chemical reagents. We illustrate these processes with the chlorination of alkanes (Chapter 3). A Chlorination Reaction Next we look at cyclic alkanes (Chapter 4), which contain carbon atoms in a ring. This arrangement can lead to new properties and changes in reactivity. The recognition of a new type of isomerism in cycloalkanes bearing two or more substituents—either on the same side or on opposite sides of the ring plane—sets the stage for a general discussion of stereoisomerism. Stereoisomerism is exhibited by compounds with the same connectivity but differing in the relative positioning of their component atoms in space (Chapter 5). We shall then study the haloalkanes, our first example of compounds containing a functional group—the carbon–halogen bond. The haloalkanes participate in two types of organic reactions: substitution and elimination (Chapters 6 and 7). In a substitution reaction, one halogen atom may be replaced by another; in an elimination process, adjacent atoms may be removed from a molecule to generate a double bond. Like the haloalkanes, each of the major classes of organic compounds is characterized by a particular functional group. For example, the carbon– carbon triple bond is the functional group of alkynes (Chapter 13); the smallest alkyne, acetylene, is the chemical burned in a welder’s torch. A carbon-oxygen double bond is characteristic of aldehydes and ketones (Chapter 17); formaldehyde and acetone are major industrial commodities. The amines (Chapter 21), which include drugs such as nasal decongestants and amphetamines, contain nitrogen in their functional group; methylamine is a starting material in many syntheses of medicinal compounds. We shall study the tools for identifying these molecular subunits, especially the various forms of spectroscopy (Chapters 10, 11, and 14). Organic chemists rely on an array of spectroscopic methods to elucidate the structures of unknown compounds. All of these methods depend on the absorption of electromagnetic radiation at specific wavelengths and the correlation of this information with structural features. Subsequently, we shall encounter organic molecules that are especially crucial in biology and industry. Many of these, such as the carbohydrates (Chapter 24) and amino acids (Chapter 26), contain multiple functional groups. However, in every class of organic compounds, the principle remains the same: The structure of the molecule determines the reactions that it can undergo. Synthesis is the making of new molecules Carbon compounds are called “organic” because it was originally thought that they could be produced only from living organisms. In 1828, Friedrich Wöhler1 proved this idea to be false when he converted the inorganic salt lead cyanate into urea, an organic product of protein metabolism in mammals (Real Life 1-1). Wöhler’s Synthesis of Urea Synthesis, or the making of molecules, is a very important part of organic chemistry (Chapter 8). Since Wöhler’s time, many millions of organic substances have been synthesized from simpler materials, both organic and inorganic.2 These substances include many that also occur in nature, such as the penicillin antibiotics, as well as entirely new compounds. Some, such as cubane, have given chemists the opportunity to study special kinds of bonding and reactivity. Others, such as the artificial sweetener saccharin, have become a part of everyday life. Typically, the goal of synthesis is to construct complex organic chemicals from simpler, more readily available ones. To be able to convert one molecule into another, chemists must know organic reactions. They must also know the physical factors that govern such processes, such as temperature, pressure, solvent, and molecular structure. This knowledge is equally valuable in analyzing reactions in living systems. As we study the chemistry of each functional group, we shall develop the tools both for planning effective syntheses and for predicting the processes that take place in nature. But how? The answer lies in looking at reactions step by step. An organic molecular architect at work. REAL LIFE: NATURE 1-1 Urea: From Urine to Wöhler’s Synthesis to Industrial Fertilizer Urination is the main process by which we excrete nitrogen from our bodies. Urine is produced by the kidneys and then stored in the bladder, which begins to contract when its volume exceeds about 200 mL. The average human excretes about 1.5 L of urine daily, and a major component is urea, about 20 g per liter. In an attempt to probe the origins of kidney stones, early (al)chemists, in the 18th century, attempted to isolate the components of urine by crystallization, but they were stymied by the cocrystallization with the also present sodium chloride. William Prout,3 an English chemist and physician, is credited with the preparation of pure urea in 1817 and the determination of its accurate elemental analysis as CH4N2O. Prout was an avid proponent of the then revolutionary thinking that disease has a molecular basis and could be understood as such. This view clashed with that of the so-called vitalists, who believed that the functions of a living organism are controlled by a “vital principle” and cannot be explained by chemistry (or physics). Into this argument entered Wöhler, an inorganic chemist, who attempted to make ammonium cyanate, NH4+OCN− (also CH4N2O ), from lead cyanate and ammonia in 1828, but who obtained the same compound that Prout had characterized as urea. To one of his mentors, Wöhler wrote, “I can make urea without a kidney, or even a living creature.” In his landmark paper, “On the Artificial Formation of Urea,” he commented on his synthesis as a “remarkable fact, as it is an example of the artificial generation of an organic material from inorganic materials.” He also alluded to the significance of the finding that a compound with an identical elemental composition as ammonium cyanate can have such completely different chemical properties, a forerunner to the recognition of isomeric compounds. Wöhler’s synthesis of urea forced his contemporary vitalists to accept the notion that simple organic compounds could be made in the laboratory. As you shall see in this book, over the ensuing decades, synthesis has yielded much more complex molecules than urea, some of them endowed with self-replicating and other “lifelike” properties, such that the boundaries between what is lifeless and what is alive are dwindling. Apart from its function in the body, urea’s high nitrogen content makes it an ideal fertilizer. It is also a raw material in the manufacture of plastics and glues, an ingredient of some toiletry products and fire extinguishers, and an alternative to rock salt for deicing roads. It is produced industrially from ammonia and carbon dioxide to the tune of 200 million tons per year worldwide. The effect of nitrogen fertilizer on wheat growth: treated on the left; untreated on the right. Reactions are the vocabulary and mechanisms are the grammar of organic chemistry When we introduce a chemical reaction, we will first show just the starting compounds, or reactants (also called substrates), and the products. In the chlorination process mentioned earlier, the substrates—methane, CH4, and chlorine, Cl2 —may undergo a reaction to give chloromethane, CH3Cl, and hydrogen chloride, HCl. We described the overall transformation as CH4+Cl2→CH3Cl+HCl. However, even a simple reaction such as this one may proceed through a complex sequence of steps. The reactants could have first formed one or more unobserved substances—call these X—that rapidly changed into the observed products. These underlying details of the reaction constitute the reaction mechanism. In our example, the mechanism consists of two major parts: CH4+Cl2→X followed by X→CH3Cl+HCl. Each part is crucial in determining whether the overall reaction will proceed. Substances X in our chlorination reaction are examples of reaction intermediates, species formed on the pathway between reactants and products. We shall learn the mechanism of this chlorination process and the nature of the reaction intermediates in Chapter 3. How can we determine reaction mechanisms? The strict answer to this question is, we cannot. All we can do is amass circumstantial evidence that is consistent with (or points to) a certain sequence of molecular events that connect starting materials and products (“the postulated mechanism”). To do so, we exploit the fact that organic molecules are no more than collections of bonded atoms. We can, therefore, study how, when, and how fast bonds break and form, in which way they do so in three dimensions, and how changes in substrate structure affect the outcome of reactions. Thus, although we cannot strictly prove a mechanism, we can certainly rule out many (or even all) reasonable alternatives and propose a most likely pathway. In a way, the “learning” and “using” of organic chemistry is much like learning and using a language. You need the vocabulary (i.e., the reactions) to be able to use the right words, but you also need the grammar (i.e., the mechanisms) to be able to converse intelligently. Neither one on its own gives complete knowledge and understanding, but together they form a powerful means of communication, rationalization, and predictive analysis. Before we begin our study of the principles of organic chemistry, let us review some of the elementary principles of bonding. We shall find these concepts useful in understanding and predicting the chemical reactivity and the physical properties of organic molecules. 1-2 Coulomb Forces: A Simplified View of Bonding The bonds between atoms hold a molecule together. But what causes bonding? Two atoms form a bond only if their interaction is energetically favorable, that is, if energy—heat, for example—is released when the bond is formed. Conversely, breaking that bond requires the input of the same amount of energy. The two main causes of the energy release associated with bonding are based on Coulomb’s law of electric charge: 1. Opposite charges attract each other (electrons are attracted to protons). 2. Like charges repel each other (electrons spread out in space). Bonds are made by simultaneous coulombic attraction and electron exchange Each atom consists of a nucleus, containing electrically neutral particles, or neutrons, and positively charged protons. Surrounding the nucleus are negatively charged electrons, equal in number to the protons so that the net charge is zero. As two atoms approach each other, the positively charged nucleus of the first atom attracts the electrons of the second atom; similarly, the nucleus of the second atom attracts the electrons of the first atom. As a result, the nuclei are held together by the electrons located between them. This sort of bonding is described by Coulomb’s4 law: Opposite charges attract each other with a force inversely proportional to the square of the distance between the centers of the charges. Charge separation is rectified by Coulomb’s law, appropriately in the heart of Paris. Coulomb’s Law Attracting force=constant×(+)charge×(-)chargedistance2 This attractive force causes energy to be released as the neutral atoms are brought together. The resulting decrease in energy is called the bond strength. When the atoms reach a certain closeness, no more energy is released. The distance between the two nuclei at this point is called the bond length (Figure 1-1). Bringing the atoms closer together than this distance results in a sharp increase in energy. Why? As stated above, just as opposite charges attract, like charges repel. If the atoms are too close, the electron–electron and nuclear–nuclear repulsions become stronger than the attractive forces. When the nuclei are the appropriate bond length apart, the electrons are spread out around both nuclei, and attractive and repulsive forces balance for maximum bonding. The energy content of the two-atom system is then at a minimum, the most stable situation (Figure 1-2). Figure 1-1 The changes in energy, E, that result when two atoms are brought into close proximity. At the separation defined as bond length, maximum bonding is achieved. Figure 1-2 Covalent bonding. Attractive (solid-line) and repulsive (dashed-line) forces in the bonding between two atoms. The large spheres represent areas in space in which the electrons are found around the nucleus. The small circled plus sign denotes the nucleus. An alternative to this type of bonding results from the complete transfer of an electron from one atom to the other. The result is two charged ions: one positively charged, a cation, and one negatively charged, an anion (Figure 1-3). Again, the bonding is based on coulombic attraction, this time between two ions. Figure 1-3 Ionic bonding. An alternative mode of bonding results from the complete transfer of an electron from atom 1 to atom 2, thereby generating two ions whose opposite charges attract each other. The coulombic bonding models of attracting and repelling charges shown in Figures 1-2 and 1-3 are highly simplified views of the interactions that take place in the bonding of atoms. Nevertheless, even these simple models explain many of the properties of organic molecules. In the sections to come, we shall examine increasingly more sophisticated views of bonding. 1-3 Ionic and Covalent Bonds: The Octet Rule We have seen that attraction between negatively and positively charged particles is a basis for bonding. How does this concept work in real molecules? Two extreme types of bonding explain the interactions between atoms in organic molecules: 1. A covalent bond is formed by the sharing of electrons (as shown in Figure 1-2). 2. An ionic bond is based on the electrostatic attraction of two ions with opposite charges (as shown in Figure 1-3). We shall see that many atoms bind to carbon in a way that is intermediate between these extremes: Some ionic bonds have covalent character and some covalent bonds are partly ionic (polarized). What are the factors that account for the two types of bonds? To answer this question, let us return to the atoms and their compositions. We start by looking at the periodic table and at how the electronic makeup of the elements changes as the atomic number increases. The periodic table underlies the octet rule The partial periodic table depicted in Table 1-1 includes those elements most widely found in organic molecules: carbon (C), hydrogen (H), oxygen (O), nitrogen (N), sulfur (S), chlorine (Cl), bromine (Br), and iodine (I). Certain reagents, indispensable for synthesis and commonly used, contain elements such as lithium (Li), magnesium (Mg), boron (B), and phosphorus (P). (If you are not familiar with these elements, refer to Table 1-1 or the periodic table on the inside cover.) Table 1.1 Partial Periodic Table Period First Second Third Fourth Fifth H1 Li2,1 Na2,8,1 K2,8,8,1 Be2,2 Mg2,8,2 B2,3 Al2,8,3 C2,4 Si2,8,4 N2,5 P2,8,5 Note: The superscripts indicate the number of electrons in each principal shell of the atom. The elements in the periodic table are listed according to nuclear charge (number of protons), which equals the number of electrons. The nuclear charge increases by one with each element listed. The electrons occupy energy levels, or “shells,” each with a fixed capacity. For example, the first shell has room for two electrons; the second, eight; and the third, 18. Helium, with two electrons in its shell, and the other noble gases, with eight electrons (called octets) in their outermost shells, are especially stable. These elements show very little chemical reactivity. All other elements (including carbon) lack octets in their outermost electron shells. Noble Gases Atoms tend to form molecules in such a way as to reach an octet in the outer electron shell and attain a noble-gas configuration. Carbon Atom In the next two sections, we describe two extreme ways in which this goal may be accomplished: by the formation of pure ionic or pure covalent bonds. Exercise 1-1 (a) Redraw Figure 1-1 for a weaker bond than the one depicted. (b) Write the elements in Table 1-1 from memory. In pure ionic bonds, electron octets are formed by transfer of electrons Sodium (Na), a reactive metal, interacts with chlorine, a reactive gas, in a violent manner to produce a stable substance: sodium chloride. Similarly, sodium reacts with fluorine (F), bromine, or iodine to give the respective salts. Other alkali metals, such as lithium and potassium (K), undergo the same reactions. These transformations succeed because both reaction partners attain noble-gas character by the transfer of outer-shell electrons, called valence electrons, from the alkali metals on the left side of the periodic table to the halogens on the right. Let us see how this works for the ionic bond in sodium chloride. Why is the interaction energetically favorable? First, it takes energy to remove an electron from an atom. This energy is the ionization potential (IP) of the atom. For sodium gas, the ionization energy amounts to 119kcalmol−1. 5 Conversely, energy may be released when an electron attaches itself to an atom. For chlorine, this energy, called its electron affinity (EA), is −83kcalmol−1. These two processes result in the transfer of an electron from sodium to chlorine. Together, they require a net energy input of 119−83=36kcal mol−1. Why, then, do the atoms readily form NaCl? The reason is their electrostatic attraction, which pulls them together in an ionic bond. At the most favorable interatomic distance [about 2.8 Å (angstroms; 1 Å = 10−10 m (one ten-billionth of a meter or 0.1 nm)] in the gas phase], this attraction releases (see Figure 1-1) about 120kcalmol−1(502kJmol−1). This energy release is enough to make the reaction of sodium with chlorine energetically highly favorable [+36−120= −84kcal mol−1(−351kJ mol−1)]. Formation of Ionic Bonds by Electron Transfer More than one electron may be donated (or accepted) to achieve noblegas electronic configurations. Magnesium, for example, has two valence electrons. Donation to an appropriate acceptor produces the corresponding doubly charged cation, Mg2+, with the electronic structure of neon. In this way, the ionic bonds of typical salts are formed. A representation of how charge (re)distributes itself in molecules is given by electrostatic potential maps. These computer-generated maps not only show a form of the molecule’s “electron cloud,” they also use color to depict deviations from charge neutrality. Excess electron density—for example, a negative charge—is shown in colors shaded toward red; conversely, diminishing electron density—ultimately, a positive charge—is shown in colors shaded toward blue. Charge-neutral regions are indicated by green. The reaction of a sodium atom with a chlorine atom to produce Na+Cl− is pictured this way in the margin. In the product, Na+ is blue, Cl− is red. A more convenient way of depicting valence electrons is by means of dots around the symbol for the element. In this case, the letters represent the nucleus including all the electrons in the inner shells, together called the core configuration. Valence Electrons as Electron Dots Electron-Dot Picture of Salts The hydrogen atom is unique because it may either lose an electron to become a bare nucleus, the proton, or accept an electron to form the hydride ion, [H, i.e., H:]−, which possesses the helium configuration. Indeed, the hydrides of lithium, sodium, and potassium (Li+H−,Na+H−, and K+H−) are commonly used reagents. Exercise 1-2 Draw electron-dot pictures for ionic LiBr, and MgS. Na2O,BeF2,AlCl3, In covalent bonds, electrons are shared to achieve octet configurations Formation of ionic bonds between two identical elements is difficult because the electron transfer is usually very unfavorable. For example, in H2, formation of H+H− would require an energy input of nearly 300 kcal mol−1 (1255 kJ mol−1). For the same reason, none of the halogens, F2,Cl2,Br2, and I2, has an ionic bond. The high IP of hydrogen also prevents the bonds in the hydrogen halides from being ionic. For elements nearer the center of the periodic table, the formation of ionic bonds is unfeasible, because it becomes more and more difficult to donate or accept enough electrons to attain the noble-gas configuration. Such is the case for carbon, which would have to shed four electrons to reach the helium electronic structure or add four electrons for a neon-like arrangement. The large amount of charge that would develop makes these processes very energetically unfavorable. Instead, covalent bonding takes place: The elements share electrons so that each atom attains a noble-gas configuration. Typical products of such sharing are H2 and HCl. In HCl, the chlorine atom assumes an octet structure by sharing one of its valence electrons with that of hydrogen. Similarly, the chlorine molecule, Cl2, is diatomic because both component atoms gain octets by sharing two electrons. Such bonds are called covalent single bonds. Electron-Dot Picture of Covalent Single Bonds Because carbon has four valence electrons, it must acquire a share of four electrons to gain the neon configuration, as in methane. Nitrogen has five valence electrons and needs three to share, as in ammonia; and oxygen, with six valence electrons, requires only two to share, as in water. It is possible for one atom to supply both of the electrons required for covalent bonding. This occurs upon addition of a proton to ammonia, thereby forming NH4+, or to water, thereby forming H3O+. Besides two-electron (single) bonds, atoms may form four-electron (double) and six-electron (triple) bonds to gain noble-gas configurations. Atoms that share more than one electron pair are found in ethene and ethyne. *In labels of molecules, systematic names (introduced in Section 2-6) will be given first, followed in parentheses by so-called common names that are still used frequently. The drawings above, with pairs of electron dots representing bonds, are also called Lewis7 structures. We shall develop the general rules for formulating such structures in Section 1-4. Exercise 1-3 Draw electron-dot structures for F2,C_F4,C_H2Cl2,P_H3,BrI,HO−,H2N_−, and H3C_−. (Where applicable, the underlined element is at the center of the molecule.) Make sure that all atoms have noble-gas electron configurations. In most organic bonds, the electrons are not shared equally: polar covalent bonds The preceding two sections presented two extreme ways in which atoms attain noble-gas configurations by entering into bonding: pure ionic and pure covalent. In reality, most bonds are of a nature that lies between these two extremes: polar covalent. As a result, the ionic bonds in most salts have some covalent character; conversely, the covalent bonds to carbon have some ionic or polar character. Recall (Section 1-2) that both sharing of electrons and coulombic attraction contribute to the stability of a bond. How polar are polar covalent bonds, and what is the direction of the polarity? We can answer these questions by considering the periodic table and keeping in mind that the positive nuclear charge of the elements increases from left to right. Therefore, the elements on the left of the periodic table are often called electropositive, electron donating, or “electron pushing,” because their electrons are held by the nucleus less tightly than are those of elements to the right. These elements at the right of the periodic table are described as electronegative, electron accepting, or “electron pulling.” Table 1-2 lists the relative electronegativities of some elements. On this scale, fluorine, the most electronegative of them all, is assigned the value 4. You will note that the values for electronegativity decrease steadily going down a column in the periodic table, for example, from fluorine to iodine. This observation is a consequence of Coulomb’s Law: as the atoms get larger, the electrons surrounding them are located further and further away from their respective nuclei and hence less and less attracted by it. Table 1.2 Electronegativities of Selected Elements H 2.2 Li 1.0 Na 0.9 K 0.8 Be 1.6 Mg 1.3 B 2.0 Al 1.6 C 2.6 Si 1.9 N 3.0 P 2.2 O 3.4 S 2.6 F 4.0 Cl 3.2 Br 3.0 I 2.7 Note: Values established by L. Pauling and updated by A. L. Allred (see Journal of Inorganic and Nuclear Chemistry, 1961, 17, 215). Consideration of Table 1-2 readily explains why the most ionic (least covalent) bonds occur between elements at the two extremes (e.g., the alkali metal salts, such as sodium chloride). On the other hand, the purest covalent bonds are formed between atoms of equal electronegativity (i.e., identical elements, as in H2,N2,O2,F2, and so on) or in carbon–carbon bonds. However, most covalent bonds are between atoms of differing electronegativity, resulting in their polarization. The polarization of a bond is the consequence of a shift of the center of electron density in the bond toward the more electronegative atom. It is indicated in a very qualitative manner (using the Greek letter delta, δ ) by designating a partial positive charge, δ+, and partial negative charge, δ−, to the respective less or more electronegative atom. The larger the difference in electronegativity, the bigger is the charge separation. As a rule of thumb, electronegativity differences of 0.3 to 2.0 units indicate polar covalent bonds; lesser values are typical of essentially “pure” covalent bonds, larger values of “pure” ionic ones. The separation of opposite charges is called an electric dipole, symbolized by an arrow crossed at its tail and pointing from positive to negative. A polarized bond can impart polarity to a molecule as a whole, as in HF, ICl, and CH3F. Molecules Can Have Polar Bonds but No Net Polarization Polar Bonds In symmetrical structures, the polarizations of the individual bonds may cancel, thus leading to molecules with no net polarization, such as CO2 and CCl4 . To know whether a molecule is polar, we have to know its shape, because the net polarity is the vector sum of the bond dipoles. The electrostatic potential maps of p. 11 clearly illustrate the polarization in CO2 and CCl4, showing the respective carbon atoms shaded relatively blue, the attached, more electronegative atoms relatively red. Moreover, you can recognize how the shape of each molecule renders it nonpolar as a whole. There are two cautions in viewing electrostatic potential maps: (1) The scale on which the color differentials are rendered may vary. For example, a much more sensitive scale is used for the molecules in the p. 11, in which the charges are only partial, than for NaCl on p. 8, in which the atoms assume full charges. Hence, it may be misleading to compare the electrostatic potential maps of one set of molecules with those of another, electronically very different group. Most organic structures shown in this book will be on a comparative scale, unless mentioned otherwise. (2) Because of the way in which the potential at each point is calculated, it will contain contributions from all nuclei and electrons in the vicinity. As a consequence, the color of the spatial regions around individual nuclei is not uniform. Electrostatic Potential Maps Color Scale Valence electron repulsion controls the shapes of molecules Molecules adopt shapes in which electron repulsion (including both bonding and nonbonding electrons) is minimized. In diatomic species such as H2 or LiH, there is only one bonding electron pair and one possible arrangement of the two atoms. However, beryllium fluoride, BeF2, is a triatomic species. Will it be bent or linear? Electron repulsion is at a minimum in a linear structure, because the bonding and nonbonding 8 electrons are placed as far from each other as possible, at 180°. Linearity is also expected for other derivatives of beryllium, as well as of other elements in the same column of the periodic table. In boron trichloride, the three valence electrons of boron allow it to form covalent bonds with three chlorine atoms. Electron repulsion enforces a regular trigonal arrangement—that is, the three halogens are at the corners of an equilateral triangle, the center of which is occupied by boron, and the bonding (and nonbonding) electron pairs of the respective chlorine atoms are at maximum distance from each other, that is, 120°. Other derivatives of boron, and the analogous compounds with other elements in the same column of the periodic table, are again expected to adopt trigonal structures. Applying this principle to carbon, we can see that methane, CH4, has to be tetrahedral. Placing the four hydrogens at the vertices of a tetrahedron minimizes the electron repulsion of the corresponding bonding electron pairs. This method for determining molecular shape by minimizing electron repulsion is called the valence-shell electron-pair repulsion (VSEPR) method. Note that we often draw molecules such as BCl3 and CH4 as if they were flat and had 90° angles. This depiction is for ease of drawing only. Do not confuse such two-dimensional drawings with the true threedimensional molecular shapes (trigonal for BCl3 and tetrahedral for CH4 ). Exercise 1-4 Show the bond polarization in H2O,SCO,SO,IBr,CH4,CHCl3,CH2Cl2, and CH3Cl by using dipole arrows to indicate separation of charge. (In the last four examples, place the carbon in the center of the molecule.) Exercise 1-5 Ammonia, :NH3, is not trigonal but pyramidal, with bond angles of 107.3°. Water, H2 , is not linear but bent (104.5°). Why? (Hint: Consider the effect of the nonbonding electron pairs.) In Summary There are two extreme types of bonding, ionic and covalent. Both derive favorable energetics from Coulomb forces and the attainment of noble-gas electronic structures. Most bonds are better described as something between the two types: the polar covalent (or covalent ionic) bonds. Polarity in bonds may give rise to polar molecules. The outcome depends on the shape of the molecule, which is determined in a simple manner by arrangement of its bonds and nonbonding electrons to minimize electron repulsion. 1-4 Electron-Dot Model of Bonding: Lewis Structures Lewis structures are important for predicting geometry and polarity (hence reactivity) of organic compounds, and we shall use them for that purpose throughout this book. In this section, we provide rules for writing such structures correctly and for keeping track of valence electrons. Guidelines: Drawing Lewis Structures The procedure for drawing correct electron-dot structures is straightforward, as long as the following rules are observed. Rule 1. Draw the (given or desired) molecular skeleton. As an example, consider methane. The molecule has four hydrogen atoms bonded to one central carbon atom. HHCHHHHCHHCorrectIncorrect Rule 2. Count the number of available valence electrons. Add up all the valence electrons of the component atoms. Special care has to be taken with charged structures (anions or cations), in which case the appropriate number of electrons has to be added or subtracted to account for extra charges. Rule 3. (The octet rule) Depict all covalent bonds by two shared electrons, giving as many atoms as possible a surrounding electron octet, except for H, which requires a duet. Make sure that the number of electrons used is exactly the number counted according to rule 2. Elements at the right in the periodic table may contain pairs of valence electrons not used for bonding, called lone electron pairs or just lone pairs. Consider, for example, hydrogen bromide. The shared electron pair supplies the hydrogen atom with a duet, the bromine with an octet, because the bromine carries three lone electron pairs. Conversely, in methane, the four C–H bonds satisfy the requirement of the hydrogens and, at the same time, furnish the octet for carbon. Examples of correct and incorrect Lewis structures for HBr are shown below. Correct Lewis Structure Incorrect Lewis Structures Frequently, the number of valence electrons is not sufficient to satisfy the octet rule only with single bonds. In this event, double bonds (two shared electron pairs) and even triple bonds (three shared pairs) are necessary to obtain octets. An example is the nitrogen molecule, N2, which has ten valence electrons. An N–N single bond would leave both atoms with electron sextets, and a double bond provides only one nitrogen atom with an octet. It is the molecule with a triple bond that satisfies both atoms. You may find a simple procedure useful that gives you the total number of bonds needed in a molecule to give every atom an octet (or duet). Thus, after you have counted the supply of available electrons (rule 2), add up the total “electron demand,” that is, two electrons for each hydrogen atom and eight for each other element atom. Then subtract supply from “demand” and divide by 2. For N2, demand is 16 electrons, supply is 10, and hence the number of bonds is 3. Further examples of molecules with double and triple bonds are shown below. Correct Lewis Structures In practice, another simple sequence may help. First, connect all mutually bonded atoms in your structure by single bonds (i.e., shared electron pairs); second, if there are any electrons left, distribute them as lone electron pairs to maximize the number of octets; and finally, if some of the atoms lack octet structures, change as many lone electron pairs into shared electron pairs as required to complete the octet shells (see also Solved Exercises 1-7, 1-23, and 1-24). Rule 4. Assign (formal) charges to atoms in the molecule. Each lone pair contributes two electrons to the valence electron count of an atom in a molecule, and each bonding (shared) pair contributes one. An atom is charged if this total is different from the outer-shell electron count in the free, nonbonded atom. Thus, we have the formula Formal charge=(number of outer-shellelectrons on thefree, neutral atom)(number of unsharedelectrons on the atomin the molecule)-12(number of bondingelectrons surrounding theatom in the molecule) or simply Formal charge=number of valence electrons−number of lone pair electrons −12number of bonding electrons The reason for the term formal is that, in molecules, charge is not localized on one atom but is distributed to varying degrees over its surroundings. As an example, which atom bears the positive charge in the hydronium ion? Each hydrogen has a valence electron count of 1 from the shared pair in its bond to oxygen. Because this value is the same as the electron count in the free atom, the (formal) charge on each hydrogen is zero. The electron count on the oxygen in the hydronium ion is 2 (the lone pair) + 3 (half of 6 bonding electrons) = 5. This value is one short of the number of outer-shell electrons in the free atom, thus giving the oxygen a charge of +1. Hence, the positive charge is assigned to oxygen. Another example is the nitrosyl cation, NO+. The molecule bears a lone pair on nitrogen, in addition to the triple bond connecting the nitrogen to the oxygen atom. This gives nitrogen five valence electrons, a value that matches the count in the free atom; therefore, the nitrogen atom has no charge. The same number of valence electrons (5) is found on oxygen. Because the free oxygen atom requires six valence electrons to be neutral, the oxygen in NO+ possesses the +1 charge. Other examples are shown below. Sometimes the octet rule leads to charges on atoms even in neutral molecules. The Lewis structure is then said to be charge separated. An example is carbon monoxide, CO. Some compounds containing nitrogen–oxygen bonds, such as nitric acid, HNO3, also exhibit this property. REALLY? In the evolution of his ideas on the chemical bond, Gilbert Lewis at first drew “cubical atoms,” in which the electrons were positioned at the eight corners of a cube: Drawings of cubical atoms by G. N. Lewis, 1902. [J. F. Kennedy Library, California State University, Los Angeles] Exercise 1-6 Draw Lewis structures for the following molecules: HI,CH3CH2CH3,CH3OH,HSSH,SiO2(OSiO),O2,CS2(SCS). The octet rule does not always hold The octet rule strictly holds only for the elements of the second row and then only if there is a sufficient number of valence electrons to satisfy it. There are three exceptions to be considered. Exception 1. You will have noticed that all our examples of “correct” Lewis structures contain an even number of electrons; that is, all are distributed as bonding or lone pairs. This distribution is not possible in species having an odd number of electrons, such as nitrogen oxide (NO) and neutral methyl (methyl radical, ⋅CH3; see Section 3-1). Exception 2. Some compounds of the early second-row elements, such as BeH2 and BH3, have a deficiency of valence electrons. Because compounds falling under exceptions 1 and 2 do not have octet configurations, they are unusually reactive and transform readily in reactions that lead to octet structures. For example, two molecules of ·CH3 react with each other spontaneously to give ethane, CH3–CH3, and BH3, reacts with hydride, H−, to give borohydride, BH4−. Exception 3. While the preceding exceptions indicate that we can have molecules that contain atoms having less than eight electrons in their vicinity (or for hydrogen, less than two, as in H+ ), for the second row elements, we cannot exceed the octet count. However, beyond the second row, the simple Lewis model is not strictly applied, and elements may be surrounded by more than eight valence electrons, a feature referred to as valence-shell expansion. For example, phosphorus and sulfur (as relatives of nitrogen and oxygen) are trivalent and divalent, respectively, and we can readily formulate Lewis octet structures for their derivatives. But they also form stable compounds of higher valency, among them the familiar phosphoric and sulfuric acids. Some examples of octet and expanded-octet molecules containing these elements are shown below. An explanation for this apparent violation of the octet rule is found in a more sophisticated description of atomic structure by quantum mechanics (Section 1-6). However, you will notice that, even in these cases, you can construct dipolar forms in which the Lewis octet rule is preserved (see Section 1-5). Indeed, structural and computational data show that these formulations contribute to a varying degree to the resonance picture of such molecules. Covalent bonds can be depicted as straight lines Electron-dot structures can be cumbersome, particularly for larger molecules. It is simpler to represent covalent single bonds by single straight lines; double bonds are represented by two lines and triple bonds by three. Lone electron pairs can either be shown as dots or simply omitted. The use of such notation was first suggested by the German chemist August Kekulé,9 long before electrons were discovered; structures of this type are often called Kekulé structures. Straight-Line Notation for the Covalent Bond Solved Exercise 1-7 Working with the Concepts: Drawing Lewis Structures Draw the Lewis structure of HClO2(HOClO), including the assignment of any charges to atoms. Strategy: To solve such a problem, it is best to follow the preceding Guidelines for Drawing Lewis Structures. Solution: Rule 1: The molecular skeleton is given as unbranched, as shown. Rule 2: Count the number of valence electrons: H=1,2O=12,Cl=7,total=20 Rule 3: How many bonds (shared electron pairs) do we need? The supply of electrons is 20; the electron requirement is 2 for H and 3×8=24 electrons for the other three atoms, for a total of 26 electrons. Thus, we need (26−20)/2=3bonds. To distribute all valence electrons according to the octet rule, we first connect all atoms by two-electron bonds, H:O:Cl:O, using up 6 electrons. Second, we distribute the remaining 14 electrons to provide octets for all nonhydrogen atoms (arbitrarily) starting at the left oxygen. This process requires in turn 4, 4, and 6 electrons, resulting in octet structures without needing additional electron sharing: Rule 4: We determine any formal charges by noting any discrepancies between the “effective” valence electron count around each atom in the molecule we have found and its outer-shell count when isolated. For H in HOClO, the valence electron count is 1, which is the same as in the H atom, so it is neutral in the molecule. For the neighboring oxygen, the two values are again the same, 6. For Cl, the effective electron count is 6, but the neutral atom requires 7. Therefore, Cl bears a positive charge. For the terminal O, the electron counts are 7 (in the molecule) and 6 (neutral atom), giving it a negative charge. The final result is 1-8 Try It Yourself Draw Lewis structures of the following molecules, including the assignment of any charges to atoms (the order in which the atoms are attached is given in parentheses, when it may not be obvious from the form ula as it is commonly written): SO,F2O(FOF),BF3NH3(F3BNH3),CH3OH2+(H3COH2+),Cl2C═O,CN −,C22− (Caution: To draw Lewis structures correctly, it is essential that you know the number of valence electrons that belong to each atom. If you do not know this number, look it up before you begin. If a structure is charged, you must adjust the total number of valence electrons accordingly. For example, a species with a charge of −1 must have one electron more than the total number of valence electrons contributed by the constituent atoms.) In Summary Lewis structures describe bonding by the use of electron dots or straight lines. Whenever possible, they are drawn so as to give hydrogen an electron duet and other atoms an electron octet. Charges are assigned to each atom by evaluating its electron count. 1-5 Resonance Forms In organic chemistry, we also encounter molecules for which there are several correct Lewis structures. The carbonate ion has several correct Lewis structures Let us consider the carbonate ion, CO32−. Following our rules, we can easily draw a Lewis structure (A) in which every atom is surrounded by an octet. The two negative charges are located on the bottom two oxygen atoms; the third oxygen is neutral, connected to the central carbon by a double bond and bearing two lone pairs. But why choose the bottom two oxygen atoms as the charge carriers? There is no reason at all—it is a completely arbitrary choice. We could equally well have drawn structure B or C to describe the carbonate ion. The three correct Lewis pictures are called resonance forms. Resonance Forms of the Carbonate Ion The individual resonance forms are connected by double-headed arrows and are placed within one set of square brackets. They have the characteristic property of being interconvertible by electron-pair movement only, indicated by red arrows, the nuclear positions in the molecule remaining unchanged. Note that, to turn A into B and then into C, we have to shift two electron pairs in each case. Such movement of electrons can be depicted by curved arrows, a procedure informally called “electron pushing.” The use of curved arrows to depict electron-pair movement is a useful technique that will prevent us from making the common mistake of changing the total number of electrons when we draw resonance forms. It is also advantageous in keeping track of electrons when formulating mechanisms (Sections 2-2 and 6-3). But what is its true structure? Does the carbonate ion have one uncharged oxygen atom bound to carbon through a double bond and two other oxygen atoms bound through a single bond each, both bearing a negative charge, as suggested by the Lewis structures? Or, to put it differently, are A, B, and C equilibrating isomers? The answer is no. If that were true, the carbon-oxygen bonds would be of different lengths, because double bonds are normally shorter than single bonds. But the carbonate ion is perfectly symmetrical and contains a trigonal central carbon, all C–O bonds being of equal length—between the length of a double and that of a single bond. The negative charge is evenly distributed over all three oxygens: It is said to be delocalized, in accord with the tendency of electrons to “spread out in space” (Section 1-2). In other words, none of the individual Lewis representations of this molecule is correct on its own. Rather, the true structure is a composite of A, B, and C. The resulting picture is called a resonance hybrid. Because A, B, and C are equivalent (i.e., each is composed of the same number of atoms, bonds, and electron pairs), they contribute equally to the true structure of the molecule, but none of them by itself accurately represents it. Dotted-Line Notation of Carbonate as a Resonance Hybrid Because it minimizes coulombic repulsion, delocalization by resonance has a stabilizing effect: The carbonate ion is considerably more stable than would be expected for a doubly negatively charged organic molecule. The word resonance may imply to you that the molecule vibrates or equilibrates from one form to another. This inference is incorrect. The molecule never looks like any of the individual resonance forms; it has only one structure, the resonance hybrid. Unlike substances in ordinary chemical equilibria, resonance forms are not real, although each makes a partial contribution to reality. It is for this reason that the special convention of doubleheaded arrows and square brackets is used. The trigonal symmetry of carbonate is clearly evident in its electrostatic potential map shown on p. 18. An alternative convention used to describe resonance hybrids such as carbonate is to represent the bonds as a combination of solid and dotted lines. The 23− sign here indicates that a partial charge (23 of a negative charge) resides on each oxygen atom. The equivalence of all three carbonoxygen bonds and all three oxygens is clearly indicated by this convention. Other examples of resonance hybrids of octet Lewis structures are the acetate ion and the 2-propenyl (allyl) anion. Resonance is also possible for nonoctet molecules. For example, the 2propenyl (allyl) cation is stabilized by resonance. When drawing resonance forms, keep in mind that (1) pushing one electron pair toward one atom and away from another results in a movement of charge—the atom at the beginning of the arrow takes on a plus charge, that at the end, a minus charge; (2) the relative positions of all the atoms stay unchanged—only electrons are moved; (3) equivalent resonance forms contribute equally to the resonance hybrid; (4) the arrows connecting resonance forms are double headed (↔); and (5) we never exceed the octet count for elements in the second row. The recognition and formulation of resonance forms is important in predicting reactivity. For example, reaction of carbonate with acid can occur at any two of the three oxygens to give carbonic acid, H2CO3 (which is actually in equilibrium with CO2 and H2O ). Similarly, acetate ion is protonated at either oxygen to form acetic acid (see on p. 19). Analogously, the 2-propenyl anion is protonated at either terminus to furnish propene, and the corresponding cation reacts with hydroxide at either end to give the corresponding alcohol (see below). Exercise 1-9 (a) Consider molecules A–D. Does the arrow pushing in each structure lead to an acceptable resonance form? If so, draw it and explain your answer. (b) Draw two resonance forms for nitrite ion, NO2−. What can you say about the geometry of this molecule (linear or bent)? (Hint: Consider the effect of electron repulsion exerted by the lone pair on nitrogen.) (c) The possibility of valence-shell expansion increases the number of feasible resonance forms, and it is often difficult to decide on one that is “best.” One criterion that is used is whether the Lewis structure predicts bond lengths and bond angles with reasonable accuracy. Draw Lewis octet and valence-shellexpanded resonance forms for SO2(OSO). Considering the Lewis structure for SO (Exercise 1-8), its experimental bond length of 1.48Å, and the measured S–O distance in SO2 of 1.43Å, which one of the various structures would you consider “best”? You may find it easier to picture resonance by thinking about combining colors to produce a new one. For example, mixing yellow—one resonance form—and blue—a second resonance form—results in the color green: the resonance hybrid. Not all resonance forms are equivalent The molecules described above all have equivalent resonance forms. However, many molecules are described by resonance forms that are not equivalent. An example is the enolate ion. The two resonance forms differ in the locations of both the double bond and the charge. The Two Nonequivalent Resonance Forms of the Enolate Ion Although both forms are contributors to the true structure of the ion, we shall see that one contributes more than the other. The question is, which one? If we extend our consideration of nonequivalent resonance forms to include those containing atoms without electron octets, the question becomes more general. [Octet ↔ Nonoctet] Resonance Forms Such an extension requires that we relax our definitions of “correct” and “incorrect” Lewis structures and broadly regard all resonance forms as potential contributors to the true picture of a molecule. The task is then to recognize which resonance form is the most important one. In other words, which one is the major resonance contributor? Here are some guidelines. Guidelines: Drawing Resonance Structures Guideline 1. Structures with a maximum of oct