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Organic Chemistry: Structure and Function

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Organic Chemistry: Structure and Function 8emaintains the classic framework with a logical organization that an organic molecule's structure will determine its function and strengthens a focus on helping students understand reactions, mechanisms, and synthetic analysis and their practical applications. The eighth edition presents a refined methodology, rooted in teaching expertise to promote student understanding and build problem solving skills. Paired with SaplingPlus, students will have access to an interactive and fully mobile ebook, interactive media features and well respected Sapling tutorial style problems--Where every problem emphasizes learning with hints, targeted feedback and detailed solutions as well as a unique pedagogically focused drawing tool.


Accessible Organization with a Classic Framework:

Framework emphasizes that the structure of an organic molecule determines how that molecule functions--either it's physical behavior or in a chemical reaction

NEWchapter-opening Learning Objectivesandchapter-ending Summariesprovide a framework for student learning by outlining the concepts appearing in each chapter.

Early Introduction to Electron pushing arrows-- introduced in early sections and then reinforced in margins throughout the text to help students follow what is happening in the reaction

Focus on Reaction Mechanisms

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.

NEWannotations on key mechanistic illustrationsrecall and reinforce the key tenets of mechanism.

Interlude: A Summary of Organic Reaction Mechanisms?(following Chapter 14) summarizes the key mechanisms that drive the majority of organic reactions.

Reaction Mechanism Summary Roadmaps?summarize the reactivity of each major functional group.

Problem Solving Skills and Strategy

Expanded WHIP problem-solving strategy is included with solved exercises within each chapter, providing a guide for how students should approach problem solving: What How Information Proceed.

NEWGuidelineshighlight the steps that students will use to learn new materials and solve important problems.

Worked Examples Integrating the Concepts section at the end of every chapter includes a series of sample exercises with solutions of the chapter's most important concepts.

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.

A wide range of end-of-chapter problems?include?Chapter Integration Problems?with step-by-step, worked out solutions,Team Problems?for collaborative learning, and?Pre-Professional?Problems?in the style of the?MCAT.?

Practical Applications and Visualization

Every chapter features discussions of biological, medical, and industrial applications of organic chemistry

Real Life Boxes- Real Chemistry By Practicing Chemists

Really? Margin entries highlighting unusual and surprising aspects of Organic concepts intended to stimulate students' engagement

Animations in Sapling help students visualize the most important structures and concepts in organic chemistry.

Graphical cues, animations, and models
to help students visualize reactions to promotes understanding versus rote memorization
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Hello. I am Zarif Samarov I work as chemistry teacher
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Hello. I am SunYanchun I work as Organic synthesis researcher in China
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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.

University of California at Berkeley

University of California at Davis

Vice President: Ben Roberts
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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

Brief Contents
PREFACE: A User’s Guide to ORGANIC CHEMISTRY: Structure and




Acids and Bases, Polar and Nonpolar Molecules


Bond-Dissociation Energies, Radical Halogenation,
and Relative






Bimolecular Nucleophilic Substitution


Unimolecular Substitution and Pathways of


Properties, Preparation, and Strategy of Synthesis










The Carbon–Carbon Triple Bond


Investigation by Ultraviolet and Visible


Electrophilic Aromatic Substitution


Substituents Control Regioselectivity


The Carbonyl Group


α,β-Unsaturated Aldehydes and Ketones






Functional Groups Containing Nitrogen


Alkylbenzenes, Phenols, and Anilines



Synthesis of β-Dicarbonyl Compounds; Acyl Anion


Polyfunctional Compounds in Nature


Heteroatoms in Cyclic Organic Compounds


Nitrogen-Containing Polymers in Nature

Answers to Exercises

PREFACE: A User’s Guide to ORGANIC CHEMISTRY: Structure and






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


Acids and Bases, Polar and Nonpolar Molecules




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
Rotation about Single Bonds: Conformations
Rotation in Substituted Ethanes
Worked Examples: Integrating the Concepts
Important Concepts

Bond-Dissociation Energies, Radical Halogenation,
and Relative Reactivity



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 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
Combustion and the Relative Stabilities of Alkanes
Worked Examples: Integrating the Concepts
Important Concepts

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

Chiral Molecules
Real Life: Nature 5-1 Chiral Substances in Nature




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

Bimolecular Nucleophilic Substitution


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:
Frontside or Backside Attack? Stereochemistry of the SN2
Consequences of Inversion in SN2
Structure and SN2
Reactivity: The Leaving Group
Structure and SN2
Reactivity: The Nucleophile
Keys to Success: Choosing Among Multiple Mechanistic
Structure and SN2
Reactivity: The Substrate



The SN2
Reaction at a Glance
Worked Examples: Integrating the Concepts
Important Concepts

Unimolecular Substitution and Pathways of




Solvolysis of Tertiary and Secondary Haloalkanes
Unimolecular Nucleophilic Substitution
Stereochemical Consequences of SN1
Effects of Solvent, Leaving Group, and Nucleophile on
Unimolecular Substitution
Effect of the Alkyl Group on the SN1
Reaction: Carbocation
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

Properties, Preparation, and Strategy of Synthesis





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




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




Reactions of Oxacyclopropanes
Real Life: Chemistry 9-3 Hydrolytic Kinetic Resolution of
Sulfur Analogs of Alcohols and Ethers
Physiological Properties and Uses of Alcohols and Ethers
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts





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
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



Worked Examples: Integrating the Concepts
Important Concepts






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
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
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts







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
Electrophilic Addition of Halogens to Alkenes
The Generality of Electrophilic Addition
Oxymercuration–Demercuration: A Special Electrophilic
Real Life: Medicine 12-1 Juvenile Hormone Analogs in the
Battle Against Insect-Borne Diseases
Hydroboration–Oxidation: A Stereospecific Anti-Markovnikov
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


The Carbon–Carbon Triple Bond



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
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

Investigation by Ultraviolet and Visible


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
Two Neighboring Double Bonds: Conjugated Dienes
Electrophilic Attack on Conjugated Dienes: Kinetic and
Thermodynamic Control




Delocalization Among More than Two Pi Bonds: Extended
Conjugation and Benzene
A Special Transformation of Conjugated Dienes: Diels-Alder
Real Life: Materials 14-1 Organic Polyenes Conduct
Real Life: Sustainability 14-2 The Diels-Alder Reaction is
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

A Summary of Organic Reaction Mechanisms


Electrophilic Aromatic Substitution



Naming the Benzenes
Structure and Resonance Energy of Benzene: A First Look at
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



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

Substituents Control Regioselectivity




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
Reactivity of Polycyclic Benzenoid Hydrocarbons
Polycyclic Aromatic Hydrocarbons and Cancer
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts


The Carbonyl Group



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
Oxidative Chemical Tests for Aldehydes
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts

α,β -Unsaturated Aldehydes and Ketones


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
Real Life: Biology And Medicine 18-1 Stereoselective Aldol




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
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

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
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






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

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
Alkanenitriles: A Special Class of Carboxylic Acid Derivatives
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts

Functional Groups Containing Nitrogen





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
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

Alkylbenzenes, Phenols, and Anilines



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




An Electrocyclic Reaction of the Benzene Ring: The Claisen
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
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts

Synthesis of β-Dicarbonyl Compounds; Acyl Anion




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


Polyfunctional Compounds in Nature





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
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
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

Heteroatoms in Cyclic Organic Compounds


Naming the Heterocycles
Nonaromatic Heterocycles
Real Life: Medicine 25-1 Smoking, Nicotine, Cancer, and
Medicinal Chemistry






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
Real Life: Nature 25-4 Nature is Not Always Green: Natural
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts

Nitrogen-Containing Polymers in Nature




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




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
Real Life: Forensics 26-4 DNA Fingerprinting
Worked Examples: Integrating the Concepts
New Reactions
Important Concepts

Answers to Exercises


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

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
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
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

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
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.

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
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

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
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:

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
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Patrick Donoghue, Appalachian State University
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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
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Nicholas Shaw, Appalachian State University
Supriya Sihi, Houston Community College
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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

Chapter 1 Structure and Bonding in Organic

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
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

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

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
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−
), 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

The effect of nitrogen fertilizer on wheat growth: treated on the left; untreated on the

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,
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
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
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







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 119⁢kcal⁢mol−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 −83⁢kcal⁢mol−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=36⁢kcal 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 120⁢kcal⁢mol−1⁢(502⁢kJ⁢mol−1).
This energy release is enough to make the
reaction of sodium with chlorine energetically highly favorable [+36−120=
−84⁢kcal mol−1⁢(−351⁢kJ 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,
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
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.


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_−.
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








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

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
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
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
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

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
Water, H2 , is not linear but bent (104.5°).
(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

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.

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

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
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
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
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.

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:

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
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
with each other spontaneously to give ethane, CH3–CH3,
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

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
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.
To solve such a problem, it is best to follow the preceding Guidelines for
Drawing Lewis Structures.

Rule 1: The molecular skeleton is given as unbranched, as shown.
Rule 2: Count the number of valence electrons:
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
electrons for the other three atoms, for a total of 26
electrons. Thus, we need (26−20)/2=3⁢bonds.
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):
(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
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
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