Main Organic Chemistry

Organic Chemistry

In Organic Chemistry, 3rd Edition, Dr. David Klein builds on the phenomenal success of the first two editions, which presented his unique skills-based approach to learning organic chemistry. Dr. Klein’s skills-based approach includes all of the concepts typically covered in an organic chemistry textbook, and places special emphasis on skills development to support these concepts. This emphasis on skills development in unique SkillBuilder examples provides extensive opportunities for two-semester Organic Chemistry students to develop proficiency in the key skills necessary to succeed in organic chemistry.
Year: 2017
Edition: 3
Publisher: John Wiley & Sons
Language: english
Pages: 1344 / 1319
ISBN 13: 978-1-119-31615-2
File: PDF, 92.40 MB
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Approximate pKa Values for Commonly Encountered Structural Types

R1
H
H
H
H

R2

X

pKa

Br
Cl
F

–10
–9.0
–7.0
3.2

R

pKa

CF3
OH
Me
Ph

–14
–9.0
–1.2
–0.6

R

pKa

CF3
H
Me
t-Bu
OH

–0.25
3.8
4.8
5.0
6.4

R3

pKa

NO2
H
H
NO2
H
H
H
OMe

7.1
8.4
9.9
10.2

R1

R2

pKa

Me
OEt
OMe
OEt

Me
Me
OMe
OEt

9.0
11
13
13.3

R1

R2

R3

pKa

Me
Me
Me
H
CF3
CF3

Me
Me
H
H
H
CF3

Me
H
H
H
H
H

18.0
16.5
16.0
15.5
12.5
9.3

R

pKa

t-Bu
Et

24.5
25.0

H

−10

X

H

⊕

O

R2

R1

−5

O
R

S

H

OH

O

O

R1 ⊕ R2
O
OH (–1.3)

⊕N

0

⊝O

O

⊕

5

OH

CH3CO3H (8.2)
R1

R2

⊕

N

H

R3
R2

R1
H

H
H

15

R1
R2

C

O

H

H
H

O
H

RO
H

H

H

R
H

R

C

H

C

H

25
O

35
R1

R2

pKa

H
Et
Et

H
H
Et

38
38
40

R1

N

H
S
(35)
H 3C
C
DMSO H H
H

H

H (36)

R2

40
H

H
C

R1

R2

R3

pKa

Ph
CH=CH2
H
Me
Me
Me

H
H
H
H
Me
Me

H
H
H
H
H
Me

41
43
48
50
51
53

45
R1
R2

C

H

R3

50

H

pKa
–3.8
–3.6
–2.4
–2.2
–1.7

R1

R2

R3

pKa

H
Me
Me
Me
Et
Pr

H
H
Me
Me
Et
Pr

H
H
H
Me
Et
H

9.2
10.5
10.6
10.6
10.8
11.1

(15)

O

20

R2
Me
Et
H
H
H

H (15.7)

O

R3

R1
Me
Et
Et
Me
H

H (7.0)

S

O

–8.0
–7.3
–6.5
–6.2
–6.1

H (4.7)

N

N

⊕

10

R2

O

N
H

R1
R3

⊝

pKa

H
Me
OMe
Ph
OH

H (3.4)

N
OH

R2

F (3.2)

H

R

R1
Me
Me
Me
Me
Me

(44)

C
H

R

pKa

Ph
H
Me

16.0
17.0
19.2

R

pKa

Ph
H

23
25

This page intentionally left blank

O r g a n i c C h e m i s t ry
T h i r d Ed i t i o n

D av i d K l e i n
Jo h n s H o p k i n s U n i v e r s i t y

VICE PRESIDENT: SCIENCE Petra Recter
EXECUTIVE EDITOR Sladjana Bruno
SPONSORING EDITOR Joan Kalkut
EXECUTIVE MARKETING MANAGER Kristine Ruff
PRODUCT DESIGNER Sean Hickey
SENIOR DESIGNER Thomas Nery
SENIOR PHOTO EDITOR Billy Ray
EDITORIAL ASSISTANTS Esther Kamar, Mili Ali
SENIOR PRODUCTION EDITOR/MEDIA SPECIALIST Elizabeth Swain
Production Manager Sofia Buono
Cover/preface photo credits: flask 1 (lemons) Africa Studio/Shutterstock; flask 2 (cells) Lightspring/
Shutterstock; flask 3 (pills) photka/Shutterstock.
The book was set in 10/12 Garamond by codeMantra and printed and bound by Quad Graphics.
The cover was printed by Quad Graphics.
Copyright © 2017, 2015, 2012 John Wiley and Sons, Inc. All rights reserved.
No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form
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ISBN 978-1-119-31615-2
Printed in the United States of America 10 9 8 7 6 5 4 3 2 1
The inside back cover will contain printing identification and country of origin if omitted from this
page. In addition, if the ISBN on the back cover differs from the ISBN on this page, the one on the
back cover is correct.

Dedication
To my father and mother,
You have saved me (quite literally) on so many occasions, always steering me in the
right direction. I have always cherished your guidance, which has served as a compass
for me in all of my pursuits. You ­repeatedly urged me to work on this textbook (“write the
book!”, you would say so often), with full confidence that it would be appreciated by students around the world. I will forever rely on the life lessons that you have taught me and
the values that you have instilled in me. I love you.

To Larry,
By inspiring me to pursue a career in organic chemistry instruction, you served as the
spark for the creation of this book. You showed me that any subject can be fascinating (even
organic chemistry!) when presented by a masterful teacher. Your mentorship and friendship
have profoundly shaped the course of my life, and I hope that this book will always serve as
a source of pride and as a reminder of the impact you’ve had on your students.

To my wife, Vered,
This book would not have been possible without your partnership. As I worked for years
in my office, you shouldered all of our life responsibilities, including taking care of all of
the needs of our five amazing children. This book is our collective accomplishment and will
forever serve as a testament of your constant support that I have come to depend on for
everything in life. You are my rock, my partner, and my best friend. I love you.

Brief Contents
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27

A Review of General Chemistry: Electrons, Bonds, and Molecular Properties 1
Molecular Representations 49
Acids and Bases 93
Alkanes and Cycloalkanes 132
Stereoisomerism 181
Chemical Reactivity and Mechanisms 226
Alkyl Halides: Nucleophilic Substitution and Elimination Reactions 271
Addition Reactions of Alkenes 343
Alkynes 400
Radical Reactions 435
Synthesis 479
Alcohols and Phenols 505
Ethers and Epoxides; Thiols and Sulfides 556
Infrared Spectroscopy and Mass Spectrometry 602
Nuclear Magnetic Resonance Spectroscopy 649
Conjugated Pi Systems and Pericyclic Reactions 701
Aromatic Compounds 751
Aromatic Substitution Reactions 790
Aldehydes and Ketones 844
Carboxylic Acids and Their Derivatives 898
Alpha Carbon Chemistry: Enols and Enolates 954
Amines 1008
Introduction to Organometallic Compounds 1054
Carbohydrates 1107
Amino Acids, Peptides, and Proteins 1147
Lipids 1190
Synthetic Polymers 1227

Contents

2.7 Introduction to Resonance 63
2.8 Curved Arrows 65
2.9 Formal Charges in Resonance Structures 68

1

A Review of General Chemistry:
Electrons, Bonds, and Molecular
Properties 1

2.10 D
 rawing Resonance Structures via Pattern
Recognition 70
2.11 A
 ssessing the Relative Importance of Resonance
Structures 75
2.12 The Resonance Hybrid 79
2.13 Delocalized and Localized Lone Pairs 81
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

1.1 Introduction to Organic Chemistry 2
1.2 The Structural Theory of Matter 3
1.3 Electrons, Bonds, and Lewis Structures 4
1.4 Identifying Formal Charges 8
1.5 Induction and Polar Covalent Bonds 9
PRACTICALLY SPEAKING Electrostatic Potential
Maps 12
1.6 Atomic Orbitals 12
1.7 Valence Bond Theory 16
1.8 Molecular Orbital Theory 17
1.9 Hybridized Atomic Orbitals 18
1.10 Predicting Molecular Geometry: VSEPR Theory 24
1.11 Dipole Moments and Molecular Polarity 28
1.12 Intermolecular Forces and Physical Properties 32
PRACTICALLY SPEAKING Biomimicry and
Gecko Feet 35
MEDICALLY SPEAKING Drug-Receptor Interactions 38
1.13 Solubility 38
MEDICALLY SPEAKING Propofol: The Importance of
Drug Solubility 40
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

3

Acids and Bases 93
3.1 Introduction to Brønsted-Lowry Acids
and Bases 94
3.2 Flow of Electron Density: Curved-Arrow
Notation 94
MEDICALLY SPEAKING Antacids and Heartburn 96
3.3 Brønsted-Lowry Acidity: A Quantitative
Perspective 97
MEDICALLY SPEAKING Drug Distribution and pKa 103
3.4 Brønsted-Lowry Acidity: Qualitative
Perspective 104
3.5 Position of Equilibrium and Choice
of Reagents 116
3.6 Leveling Effect 119
3.7 Solvating Effects 120
3.8 Counterions 120
PRACTICALLY SPEAKING Baking Soda versus
Baking Powder 121
3.9 Lewis Acids and Bases 121

2

Molecular Representations 49
2.1 Molecular Representations 50
2.2 Bond-Line Structures 51
2.3 Identifying Functional Groups 55
MEDICALLY SPEAKING Marine Natural Products 57
2.4 Carbon Atoms with Formal Charges 58
2.5 Identifying Lone Pairs 58
2.6 Three-Dimensional Bond-Line Structures 61
MEDICALLY SPEAKING Identifying the
Pharmacophore 62

Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

4

Alkanes and Cycloalkanes 132
4.1 Introduction to Alkanes 133
4.2 Nomenclature of Alkanes 133
PRACTICALLY SPEAKING Pheromones:
Chemical Messengers 137
MEDICALLY SPEAKING Naming Drugs 145
4.3 Constitutional Isomers of Alkanes 146

v

vi   CONTENTS
4.4 Relative Stability of Isomeric Alkanes 147
4.5 Sources and Uses of Alkanes 148
PRACTICALLY SPEAKING An Introduction
to Polymers 150
4.6 Drawing Newman Projections 150
4.7 Conformational Analysis of Ethane
and Propane 152
4.8 Conformational Analysis of Butane 154
MEDICALLY SPEAKING Drugs and Their
Conformations 158
4.9 Cycloalkanes 158
MEDICALLY SPEAKING Cyclopropane as an
Inhalation Anesthetic 160
4.10 Conformations of Cyclohexane 161
4.11 Drawing Chair Conformations 162
4.12 Monosubstituted Cyclohexane 164
4.13 Disubstituted Cyclohexane 166
4.14 cis-trans Stereoisomerism 170
4.15 Polycyclic Systems 171
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

6

Chemical Reactivity and Mechanisms 226
6.1 Enthalpy 227
6.2 Entropy 230
6.3 Gibbs Free Energy 232
PRACTICALLY SPEAKING Explosives 233
PRACTICALLY SPEAKING Do Living Organisms Violate
the Second Law of Thermodynamics? 235
6.4 Equilibria 235
6.5 Kinetics 237
MEDICALLY SPEAKING Nitroglycerin: An Explosive
with Medicinal Properties 240
PRACTICALLY SPEAKING Beer Making 241
6.6 Reading Energy Diagrams 242
6.7 Nucleophiles and Electrophiles 245
6.8 Mechanisms and Arrow Pushing 248
6.9 Combining the Patterns of Arrow Pushing 253
6.10 Drawing Curved Arrows 255
6.11 Carbocation Rearrangements 257

5

6.12 Reversible and Irreversible Reaction Arrows 259
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

Stereoisomerism 181
5.1 Overview of Isomerism 182
5.2 Introduction to Stereoisomerism 183
PRACTICALLY SPEAKING The Sense of Smell 188
5.3 Designating Configuration Using the
Cahn-Ingold-Prelog System 188
MEDICALLY SPEAKING Chiral Drugs 193
5.4 Optical Activity 194
5.5 Stereoisomeric Relationships: Enantiomers and
Diastereomers 200
5.6 Symmetry and Chirality 203
5.7 Fischer Projections 207
5.8 Conformationally Mobile Systems 209
5.9 Chiral Compounds That Lack a
Chiral Center 210
5.10 Resolution of Enantiomers 211
5.11 E
 and Z Designations for Diastereomeric
Alkenes 213
MEDICALLY SPEAKING Phototherapy Treatment for
Neonatal Jaundice 215
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

7

Alkyl Halides: Nucleophilic Substitution
and Elimination Reactions 271
7.1 Introduction to Substitution and Elimination
Reactions 272
7.2 Nomenclature and Uses of Alkyl Halides 273
7.3 SN2 Reactions 276
MEDICALLY SPEAKING Pharmacology and Drug
Design 283
7.4 N
 ucleophilic Strength and Solvent Effects in
SN2 Reactions 285
7.5 SN2 Reactions in Biological Systems—Methylation 287
7.6 Introduction to E2 Reactions 289
7.7 Nomenclature and Stability of Alkenes 291
7.8 R
 egiochemical and Stereochemical Outcomes for E2
Reactions 295
7.9 Unimolecular Reactions: (SN1 and E1) 305
7.10 Kinetic Isotope Effects in Elimination Reactions 315

CONTENTS   vii
7.11 Predicting Products: Substitution vs. Elimination 317
7.12 S
 ubstitution and Elimination Reactions with Other
Substrates 323
7.13 Synthesis Strategies 327
MEDICALLY SPEAKING Radiolabeled Compounds in
Diagnostic Medicine 330
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

8

Addition Reactions of Alkenes 343
8.1 Introduction to Addition Reactions 344
8.2 Alkenes in Nature and in Industry 345

PRACTICALLY SPEAKING Conducting Organic
Polymers 404
9.2 Nomenclature of Alkynes 404
 cidity of Acetylene and Terminal
9.3 A
Alkynes 406
9.4 Preparation of Alkynes 409
9.5 Reduction of Alkynes 411
9.6 Hydrohalogenation of Alkynes 414
9.7 Hydration of Alkynes 416
9.8 Halogenation of Alkynes 422
9.9 Ozonolysis of Alkynes 422
9.10 Alkylation of Terminal Alkynes 423
9.11 Synthesis Strategies 425
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

PRACTICALLY SPEAKING Pheromones to Control
Insect Populations 345
8.3 A
 ddition vs. Elimination: A Thermodynamic
Perspective 346
8.4 Hydrohalogenation 348
PRACTICALLY SPEAKING Cationic Polymerization
and Polystyrene 355

10

Radical Reactions 435

8.5 Acid-Catalyzed Hydration 356

10.1 Radicals 436

8.6 Oxymercuration-Demercuration 360
8.7 Hydroboration-Oxidation 361

10.2 C
 ommon Patterns in Radical
Mechanisms 441

8.8 Catalytic Hydrogenation 367

10.3 Chlorination of Methane 444

PRACTICALLY SPEAKING Partially Hydrogenated
Fats and Oils 372

10.4 Thermodynamic Considerations
for Halogenation Reactions 448

8.9 Halogenation and Halohydrin Formation 373

10.5 Selectivity of Halogenation 450

8.10 Anti Dihydroxylation 377

10.6 Stereochemistry of Halogenation 453

8.11 Syn Dihydroxylation 380
8.12 Oxidative Cleavage 381
8.13 P
 redicting the Products of an Addition
Reaction 383
8.14 Synthesis Strategies 385
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

9

Alkynes 400
9.1 Introduction to Alkynes 401
MEDICALLY SPEAKING The Role of Molecular
Rigidity 403

10.7 Allylic Bromination 455
10.8 A
 tmospheric Chemistry and the
Ozone Layer 458
PRACTICALLY SPEAKING Fighting Fires with
Chemicals 460
10.9 Autooxidation and Antioxidants 461
MEDICALLY SPEAKING Why Is an Overdose of
Acetaminophen Fatal? 463
10.10 R
 adical Addition of HBr: Anti-Markovnikov
Addition 464
10.11 Radical Polymerization 468
10.12 R
 adical Processes in the Petrochemical
Industry 470
10.13 H
 alogenation as a Synthetic
Technique 470
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

viii   CONTENTS

11

Synthesis 479
11.1 One-Step Syntheses 480

12.12 Oxidation of Phenol 539
12.13 Synthesis Strategies 541
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

11.2 Functional Group Transformations 481
11.3 Reactions That Change the Carbon
Skeleton 484
MEDICALLY SPEAKING Vitamins 486
11.4 How to Approach a Synthesis Problem 487
MEDICALLY SPEAKING The Total Synthesis of
Vitamin B12 489
11.5 Retrosynthetic Analysis 491
PRACTICALLY SPEAKING Retrosynthetic
Analysis 496
11.6 Green Chemistry 496
11.7 Practical Tips for Increasing Proficiency 497
MEDICALLY SPEAKING Total Synthesis of
Taxol 498
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems
Challenge Problems

13

Ethers and Epoxides; Thiols and
Sulfides 556
13.1 Introduction to Ethers 557
13.2 Nomenclature of Ethers 557
13.3 Structure and Properties of Ethers 559
MEDICALLY SPEAKING Ethers as Inhalation
Anesthetics 560
13.4 Crown Ethers 561
MEDICALLY SPEAKING Polyether Antibiotics 563
13.5 Preparation of Ethers 563
13.6 Reactions of Ethers 566
13.7 Nomenclature of Epoxides 569

12

Alcohols and Phenols 505

MEDICALLY SPEAKING Epothilones as Novel
Anticancer Agents 570
13.8 Preparation of Epoxides 570
MEDICALLY SPEAKING Active Metabolites
and Drug Interactions 573
13.9 Enantioselective Epoxidation 573

12.1 Structure and Properties of Alcohols 506
MEDICALLY SPEAKING Chain Length as a Factor
in Drug Design 510
12.2 Acidity of Alcohols and Phenols 510
12.3 Preparation of Alcohols via Substitution or
Addition 514
12.4 Preparation of Alcohols via Reduction 515
12.5 Preparation of Diols 521
PRACTICALLY SPEAKING Antifreeze 522
12.6 Preparation of Alcohols via Grignard
Reagents 522

13.10 Ring-Opening Reactions of Epoxides 575
PRACTICALLY SPEAKING Ethylene Oxide as a Sterilizing
Agent for Sensitive Medical Equipment 578
MEDICALLY SPEAKING Cigarette Smoke
and Carcinogenic Epoxides 582
13.11 Thiols and Sulfides 583
13.12 Synthesis Strategies Involving Epoxides 586
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

12.7 Protection of Alcohols 526
12.8 Preparation of Phenols 527
12.9 Reactions of Alcohols: Substitution and
Elimination 528
PRACTICALLY SPEAKING Drug Metabolism 531
12.10 Reactions of Alcohols: Oxidation 533

14

Infrared Spectroscopy
and Mass Spectrometry 602

12.11 Biological Redox Reactions 537
PRACTICALLY SPEAKING Biological Oxidation
of Methanol and Ethanol 539

14.1 Introduction to Spectroscopy 603
PRACTICALLY SPEAKING Microwave Ovens 605

CONTENTS   ix
14.2 IR Spectroscopy 605
MEDICALLY SPEAKING IR Thermal Imaging for
Cancer Detection 606
14.3 Signal Characteristics: Wavenumber 607
14.4 Signal Characteristics: Intensity 612
PRACTICALLY SPEAKING IR Spectroscopy for Testing
Blood Alcohol Levels 614

15.11 Acquiring a 13C NMR Spectrum 685
15.12 Chemical Shifts in 13C NMR Spectroscopy 685
15.13 DEPT 13C NMR Spectroscopy 687
MEDICALLY SPEAKING Magnetic Resonance
Imaging (MRI) 690
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

14.5 Signal Characteristics: Shape 614
14.6 Analyzing an IR Spectrum 618
 sing IR Spectroscopy to Distinguish between
14.7 U
Two Compounds 623
14.8 Introduction to Mass Spectrometry 624
PRACTICALLY SPEAKING Mass Spectrometry
for Detecting Explosives 626
14.9 Analyzing the (M)+• Peak 627
14.10 Analyzing the (M+1)+• Peak 628
14.11 Analyzing the (M+2)+• Peak 630
14.12 Analyzing the Fragments 631
14.13 High-Resolution Mass Spectrometry 634
14.14 Gas Chromatography–Mass Spectrometry 636
14.15 Mass Spectrometry of Large Biomolecules 637
MEDICALLY SPEAKING Medical Applications of
Mass Spectrometry 637
14.16 Hydrogen Deficiency Index: Degrees of
Unsaturation 638
Review of Concepts & Vocabulary • SkillBuilder Review
Practice Problems • Integrated Problems • Challenge Problems

16

Conjugated Pi Systems
and Pericyclic Reactions 701
16.1 Classes of Dienes 702
16.2 Conjugated Dienes 703
16.3 Molecular Orbital Theory 705
16.4 Electrophilic Addition 709
16.5 Thermodynamic Control vs. Kinetic
Control 712
PRACTICALLY SPEAKING Natural and Synthetic
Rubbers 715
16.6 An Introduction to Pericyclic Reactions 716
16.7 Diels–Alder Reactions 717
16.8 MO Description of Cycloadditions 723
16.9 Electrocyclic Reactions 726
16.10 Sigmatropic Rearrangements 731
MEDICALLY SPEAKING The Photoinduced
Biosynthesis of Vitamin D 733

15

16.11 UV-Vis Spectroscopy 734

Nuclear Magnetic Resonance
Spectroscopy 649
15.1 Introduction to NMR Spectroscopy 650
15.2 Acquiring a 1H NMR Spectrum 652
15.3 Characteristics of a 1H NMR Spectrum 653

PRACTICALLY SPEAKING Sunscreens 738
16.12 Color 739
PRACTICALLY SPEAKING Bleach 740
16.13 Chemistry of Vision 740
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

15.4 Number of Signals 654
15.5 Chemical Shift 660
15.6 Integration 667
15.7 Multiplicity 670
15.8 D
 rawing the Expected 1H NMR Spectrum of a
Compound 678

17

Aromatic Compounds 751

1

15.9 Using H NMR Spectroscopy to Distinguish between
Compounds 679
MEDICALLY SPEAKING Detection of Impurities in
Heparin Sodium Using 1H NMR Spectroscopy 681
15.10 Analyzing a 1H NMR Spectrum 682

17.1 Introduction to Aromatic Compounds 752
PRACTICALLY SPEAKING What Is Coal? 753
17.2 Nomenclature of Benzene Derivatives 753
17.3 Structure of Benzene 756

x   CONTENTS
17.4 Stability of Benzene 757
PRACTICALLY SPEAKING Molecular Cages 761
17.5 Aromatic Compounds Other Than
Benzene 764
MEDICALLY SPEAKING The Development of
Nonsedating Antihistamines 769
17.6 Reactions at the Benzylic Position 771

19

Aldehydes and Ketones 844
JerryB7/Getty Images, Inc

19.1 Introduction to Aldehydes and Ketones 845
19.2 Nomenclature 846
19.3 Preparing Aldehydes and Ketones: A Review 848

17.7 Reduction of Benzene and Its
Derivatives 776

19.4 Introduction to Nucleophilic Addition Reactions 849

17.8 Spectroscopy of Aromatic Compounds 778

19.5 Oxygen Nucleophiles 852

PRACTICALLY SPEAKING Buckyballs and
Nanotubes 781

19.6 Nitrogen Nucleophiles 860

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

MEDICALLY SPEAKING Acetals as Prodrugs 858
PRACTICALLY SPEAKING Beta-Carotene and
Vision 864
19.7 Hydrolysis of Acetals, Imines, and Enamines 868
MEDICALLY SPEAKING Prodrugs 871
19.8 Sulfur Nucleophiles 871

18

Aromatic Substitution Reactions 790
18.1 Introduction to Electrophilic Aromatic
Substitution 791
18.2 Halogenation 791
MEDICALLY SPEAKING Halogenation
in Drug Design 794
18.3 Sulfonation 795
PRACTICALLY SPEAKING What Are Those Colors
in Fruity Pebbles? 796
18.4 Nitration 797
MEDICALLY SPEAKING The Discovery of
Prodrugs 799
18.5 Friedel–Crafts Alkylation 800
18.6 Friedel–Crafts Acylation 802
18.7 Activating Groups 804

19.9 Hydrogen Nucleophiles 872
19.10 Carbon Nucleophiles 873
PRACTICALLY SPEAKING Organic Cyanide Compounds
in Nature 876
19.11 B
 aeyer–Villiger Oxidation of Aldehydes and
Ketones 881
19.12 Synthesis Strategies 882
19.13 Spectroscopic Analysis of Aldehydes and
Ketones 885
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

20

Carboxylic Acids
and Their Derivatives 898

18.8 Deactivating Groups 808

20.1 Introduction to Carboxylic Acids 899

18.9 Halogens: The Exception 810

20.2 Nomenclature of Carboxylic Acids 899

18.10 Determining the Directing Effects of a
Substituent 812

20.4 Preparation of Carboxylic Acids 904

18.11 Multiple Substituents 815

20.5 Reactions of Carboxylic Acids 905

18.12 Synthesis Strategies 821

20.6 Introduction to Carboxylic Acid Derivatives 906

18.13 Nucleophilic Aromatic Substitution 827

20.3 Structure and Properties of Carboxylic Acids 901

MEDICALLY SPEAKING Sedatives 908

18.14 Elimination-Addition 829

20.7 Reactivity of Carboxylic Acid Derivatives 910

18.15 Identifying the Mechanism of an Aromatic
Substitution Reaction 831

20.8 Preparation and Reactions of Acid Chlorides 917

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

20.9 Preparation and Reactions of Acid Anhydrides 922
MEDICALLY SPEAKING How Does Aspirin Work? 924
20.10 Preparation of Esters 925

CONTENTS   xi
20.11 Reactions of Esters 926
PRACTICALLY SPEAKING How Soap Is Made 927
MEDICALLY SPEAKING Esters as Prodrugs 928
20.12 Preparation and Reactions of Amides 931
PRACTICALLY SPEAKING Polyamides and Polyesters 932
MEDICALLY SPEAKING Beta-Lactam Antibiotics 934

22.5 Preparation of Amines via Substitution
Reactions 1020
22.6 Preparation of Amines via Reductive Amination 1023
22.7 Synthesis Strategies 1025
22.8 Acylation of Amines 1028
22.9 Hofmann Elimination 1029

20.13 Preparation and Reactions of Nitriles 935

22.10 Reactions of Amines with Nitrous Acid 1032

20.14 Synthesis Strategies 938

22.11 Reactions of Aryl Diazonium Ions 1034

20.15 S
 pectroscopy of Carboxylic Acids and Their
Derivatives 943

22.12 Nitrogen Heterocycles 1038

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

21

Alpha Carbon Chemistry:
Enols and Enolates 954
21.1 Introduction to Alpha Carbon Chemistry:
Enols and Enolates 955
21.2 Alpha Halogenation of Enols and Enolates 962
21.3 Aldol Reactions 966
PRACTICALLY SPEAKING Muscle Power 969

MEDICALLY SPEAKING H2-Receptor Antagonists
and the Development of Cimetidine 1039
22.13 Spectroscopy of Amines 1041
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

23

Introduction to Organometallic
Compounds 1054
23.1 General Properties of Organometallic
Compounds 1055
23.2 Organolithium and Organomagnesium
Compounds 1056

21.4 Claisen Condensations 976

23.3 Lithium Dialkyl Cuprates (Gilman Reagents) 1059

21.5 Alkylation of the Alpha Position 979

23.4 The Simmons–Smith Reaction and Carbenoids 1063

21.6 Conjugate Addition Reactions 986
MEDICALLY SPEAKING Glutathione Conjugation
and Biological Michael Reactions 988
21.7 Synthesis Strategies 992
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

22

Amines 1008
22.1 Introduction to Amines 1009
MEDICALLY SPEAKING Drug Metabolism Studies 1010

23.5 Stille Coupling 1066
23.6 Suzuki Coupling 1071
23.7 Negishi Coupling 1077
23.8 The Heck Reaction 1082
23.9 Alkene Metathesis 1087
PRACTICALLY SPEAKING Improving Biodiesel
via Alkene Metathesis 1092
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

24

Carbohydrates 1107

22.2 Nomenclature of Amines 1010
22.3 Properties of Amines 1013
MEDICALLY SPEAKING Fortunate Side Effects 1014
PRACTICALLY SPEAKING Chemical Warfare
Among Ants 1018
22.4 Preparation of Amines: A Review 1019

24.1 Introduction to Carbohydrates 1108
24.2 Classification of Monosaccharides 1108
24.3 Configuration of Aldoses 1111
24.4 Configuration of Ketoses 1112
24.5 Cyclic Structures of Monosaccharides 1114

xii   CONTENTS
24.6 Reactions of Monosaccharides 1121
24.7 Disaccharides 1128
MEDICALLY SPEAKING Lactose Intolerance 1131
PRACTICALLY SPEAKING Artificial Sweeteners 1132
24.8 Polysaccharides 1133
24.9 Amino Sugars 1134
24.10 N-Glycosides 1135
MEDICALLY SPEAKING Aminoglycoside Antibiotics 1136
MEDICALLY SPEAKING Erythromycin Biosynthesis 1139
Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

26.4 Reactions of Triglycerides 1196
PRACTICALLY SPEAKING Soaps Versus Synthetic
Detergents 1201
26.5 Phospholipids 1205
MEDICALLY SPEAKING Selectivity of Antifungal
Agents 1207
26.6 Steroids 1208
MEDICALLY SPEAKING Cholesterol
and Heart Disease 1211
MEDICALLY SPEAKING Anabolic Steroids
and Competitive Sports 1214
26.7 Prostaglandins 1214
MEDICALLY SPEAKING NSAIDs and COX-2 Inhibitors 1216
26.8 Terpenes 1217

25

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

25.1 Introduction to Amino Acids, Peptides, and
Proteins 1148

27

Amino Acids, Peptides, and Proteins 1147

25.2 Structure and Properties of Amino Acids 1149
PRACTICALLY SPEAKING Nutrition and Sources
of Amino Acids 1151
PRACTICALLY SPEAKING Forensic Chemistry
and Fingerprint Detection 1155
25.3 Amino Acid Synthesis 1156
25.4 Structure of Peptides 1160
MEDICALLY SPEAKING Polypeptide
Antibiotics 1165
25.5 Sequencing a Peptide 1166
25.6 Peptide Synthesis 1169
25.7 Protein Structure 1177
MEDICALLY SPEAKING Diseases Caused
by Misfolded Proteins 1180

Synthetic Polymers 1227
27.1 Introduction to Synthetic Polymers 1228
27.2 Nomenclature of Synthetic Polymers 1229
27.3 Copolymers 1230
27.4 Polymer Classification by Reaction Type 1231
27.5 Polymer Classification by Mode of Assembly 1239
27.6 Polymer Classification by Structure 1241
27.7 Polymer Classification by Properties 1244
PRACTICALLY SPEAKING Safety Glass and Car
Windshields 1245
27.8 Polymer Recycling 1246

25.8 Protein Function 1180

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

Review of Reactions • Review of Concepts & Vocabulary
SkillBuilder Review • Practice Problems
Integrated Problems • Challenge Problems

Appendix A: Nomenclature of Polyfunctional
Compounds A–1
Glossary G–1

26

Lipids 1190
26.1 Introduction to Lipids 1191
26.2 Waxes 1192
26.3 Triglycerides 1193

Credits CR–1
Index I–1

Preface
WHY I WROTE THIS BOOK

A SKILLS-BASED APPROACH

Students who perform poorly on organic chemistry exams often
report having invested countless hours studying. Why do many
students have difficulty preparing themselves for organic chemistry exams? Certainly, there are several contributing factors,
including inefficient study habits, but perhaps the most dominant factor is a fundamental disconnect between what students
learn in the lecture hall and the tasks expected of them during
an exam. To illustrate the disconnect, consider the following
­analogy.
Imagine that a prestigious university offers a course entitled
“Bike-Riding 101.” Throughout the course, physics and engineering professors explain many concepts and principles (for example,
how bicycles have been engineered to minimize air resistance).
Students invest significant time studying the information that was
presented, and on the last day of the course, the final exam consists
of riding a bike for a distance of 100 feet. A few students may
have innate talents and can accomplish the task without falling.
But most students will fall several times, slowly making it to the
finish line, bruised and hurt; and many students will not be able to
ride for even one second without falling. Why? Because there is a
disconnect between what the students learned and what they were
expected to do for their exam.
Many years ago, I noticed that a similar disconnect exists in
traditional organic chemistry instruction. That is, learning organic
chemistry is much like bicycle riding; just as the students in the
bike-riding analogy were expected to ride a bike after attending lectures, it is often expected that organic chemistry students will independently develop the necessary skills for solving problems. While
a few students have innate talents and are able to develop the necessary skills independently, most students require guidance. This
guidance was not consistently integrated within existing textbooks,
prompting me to write the first edition of my textbook, Organic
Chemistry. The main goal of my text was to employ a skills-based
approach to bridge the gap between theory (concepts) and practice (problem-solving skills). The second edition further supported
this goal by introducing hundreds of additional problems based
on the chemical literature, thereby exposing students to exciting
real-world examples of chemical research being conducted in real
laboratories. The phenomenal success of the first two editions has
been extremely gratifying because it provided strong evidence that
my skills-based approach is indeed effective at bridging the gap
described above.
I firmly believe that the scientific discipline of organic chemistry is NOT merely a compilation of principles, but rather, it is
a disciplined method of thought and analysis. Students must certainly understand the concepts and principles, but more importantly, students must learn to think like organic chemists . . . that
is, they must learn to become proficient at approaching new situations methodically, based on a repertoire of skills. That is the true
essence of organic chemistry.

To address the disconnect in organic chemistry instruction, I have
developed a skills-based approach to instruction. The textbook
includes all of the concepts typically covered in an organic chemistry textbook, complete with conceptual checkpoints that promote
mastery of the concepts, but special emphasis is placed on skills
development through SkillBuilders to support these concepts.
Each SkillBuilder contains three parts:
Learn the Skill: contains a solved problem that demonstrates a

particular skill.

Practice the Skill: includes numerous problems (similar to the

solved problem in Learn the Skill) that give students valuable
­opportunities to practice and master the skill.

Apply the Skill: contains one or two more problems in which
the student must apply the skill to solve real-world problems (as
reported in the chemical literature). These problems include conceptual, cumulative, and applied problems that encourage students
to think outside of the box. Sometimes problems that foreshadow
concepts introduced in later chapters are also included.

At the end of each SkillBuilder, a Need More Practice? reference suggests end-of-chapter problems that students can work to
practice the skill.
This emphasis upon skills development provides students with
a greater opportunity to develop proficiency in the key skills necessary to succeed in organic chemistry. Certainly, not all necessary
skills can be covered in a textbook. However, there are certain skills
that are fundamental to all other skills.
As an example, resonance structures are used repeatedly
throughout the course, and students must become masters of resonance structures early in the course. Therefore, a significant portion of Chapter 2 is devoted to pattern-recognition for drawing
resonance structures. Rather than just providing a list of rules and
then a few follow-up problems, the skills-based approach provides
students with a series of skills, each of which must be mastered in
sequence. Each skill is reinforced with numerous practice problems. The sequence of skills is designed to foster and develop proficiency in drawing resonance structures.
The skills-based approach to organic chemistry instruction
is a unique approach. Certainly, other textbooks contain tips for
problem solving, but no other textbook consistently presents skills
development as the primary vehicle for instruction.

WHAT’S NEW IN THIS EDITION
Peer review played a very strong role in the development of the
first and second editions of Organic Chemistry. Specifically, the first
edition manuscript was reviewed by nearly 500 professors and over
5,000 students, and the second edition manuscript was based on

xiii

xiv   PREFACE
comments received from 300 professors and 900 students. In preparing the third edition, peer review has played an equally prominent role. We have received a tremendous amount of input from
the market, including surveys, class tests, diary reviews, and phone
interviews. All of this input has been carefully culled and has been
instrumental in identifying the focus of the third edition.

New Features in the Third Edition
• A new chapter on organometallic reactions covers modern synthetic techniques, including Stille coupling, Suzuki coupling,
Negishi coupling, the Heck reaction, and alkene metathesis.
• Substitution and elimination reactions have been combined
into one chapter. This chapter (Chapter 7) also features a
new section covering the preparation and reactions of alkyl
tosylates, as well as a new section covering kinetic isotope
effects. In addition, a new section introducing r­ etrosynthesis
has been added to the end of the chapter, so that synthesis
and retrosynthesis are now introduced much earlier.
• For most SkillBuilders throughout the text, the Apply the
Skill problem(s) have been replaced with moderate-level,
literature-based problems. There are at least 150 of these
new problems, which will expose students to exciting realworld examples of chemical research being conducted in
real ­laboratories. Students will see that organic chemistry is
a vibrant field of study, with endless possibilities for exploration and research that can benefit the world in concrete ways.
• Throughout the text, the distribution of problems has been
improved by reducing the number of easy problems, and
increasing the number of moderate-level, literature-based
problems.
• Each chapter now includes a problem set that mimics the
style of the ACS Organic Chemistry Exam.
• The section covering oxidation of alcohols (in Chapter 12,
and then again in Chapter 19) has been enhanced to include
modern oxidation methods, such as Swern and DMP-based
oxidations.
• Coverage of Wittig reactions has been updated to include
stereochemical outcomes and the Horner–Wadsworth–
Emmons variation.
• Section 2.11 has been revised (Assessing the relative importance of resonance structures). The rules have been completely rewritten to focus on the importance of octets and
locations of charges. The improved rules will provide students with a deeper conceptual understanding.
• In Chapter 2, a new section covers the skills necessary for
drawing a resonance hybrid.
• At the end of Chapter 5 (Stereoisomerism), a new section
introduces chiral compounds that lack chiral centers, including chiral allenes and chiral biphenyls.
• A new section in Chapter 11 (Synthesis) introduces “green
chemistry” (atom economy, toxicology issues, etc.).
• Coverage of E-Z nomenclature has been moved earlier. It
now appears in Chapter 5, which covers stereoisomerism.

TEXT ORGANIZATION
The sequence of chapters and topics in Organic Chemistry, 3e does
not differ markedly from that of other organic chemistry textbooks.
Indeed, the topics are presented in the traditional order, based on
functional groups (alkenes, alkynes, alcohols, ethers, aldehydes and
ketones, carboxylic acid derivatives, etc.). Despite this traditional
order, a strong emphasis is placed on mechanisms, with a focus on
pattern recognition to illustrate the similarities between reactions
that would otherwise appear unrelated. No shortcuts were taken in
any of the mechanisms, and all steps are clearly illustrated, including all proton transfer steps.
Two chapters (6 and 11) are devoted almost entirely to
skill development and are generally not found in other textbooks. Chapter 6, Chemical Reactivity and Mechanisms, emphasizes skills that are necessary for drawing mechanisms, while
Chapter 11, Synthesis, prepares the students for proposing syntheses. These two chapters are strategically positioned within
the traditional order described above and can be assigned to the
students for independent study. That is, these two chapters do
not need to be covered during precious lecture hours, but can
be, if so desired.
The traditional order allows instructors to adopt the skillsbased approach without having to change their lecture notes or
methods. For this reason, the spectroscopy chapters (Chapters
14 and 15) were written to be stand-alone and portable, so that
instructors can cover these chapters in any order desired. In fact,
five of the chapters (Chapters 2, 3, 7, 12, and 13) that precede
the spectroscopy chapters include end-of-chapter spectroscopy
problems, for those students who covered spectroscopy earlier.
Spectroscopy coverage also appears in subsequent functional
group chapters, specifically Chapter 17 (Aromatic Compounds),
Chapter 19 (Aldehydes and Ketones), Chapter 20 (Carboxylic
Acids and Their Derivatives), Chapter 22 (Amines), Chapter 24
(Carbohydrates), and Chapter 25 (Amino Acids, Peptides, and
Proteins).

THE WileyPLUS ADVANTAGE
WileyPLUS is a research-based online environment for effective
teaching and learning. WileyPLUS is packed with interactive study
tools and resources, including the complete online textbook.

New to WileyPLUS for Organic Chemistry, 3e
WileyPLUS for Organic Chemistry, 3e highlights David Klein’s
innovative pedagogy and teaching style:
• NEW Author-created question assignments
• NEW solved problem videos by David Klein for all new
Apply the Skill Problems
• NEW Author-curated course includes reading materials,
embedded resources, practice, and problems that have been
chosen specifically by the author
• NEW embedded Interactive exercises: over 300 interactive exercises designed to engage students with the content

PREFACE   xv

WileyPLUS for Organic Chemistry, 3e is now supported by an
adaptive learning module called ORION. Based on cognitive science, ORION provides students with a personal, adaptive learning
experience so they can build proficiency in concepts and use their
study time effectively. WileyPLUS with ORION helps students
learn by learning about them.
WileyPLUS with ORION is great as:
• An adaptive pre-lecture tool that assesses your students’ conceptual knowledge so they come to class better prepared.
• A personalized study guide that helps students understand
both strengths and areas where they need to invest more time,
especially in preparation for quizzes and exams.

•
•
•
•

Concept Review Exercises
SkillBuilder Review Exercises
Reaction Review Exercises
A list of new reagents for each chapter, with a description of
their function.
• A list of “Common Mistakes to Avoid” in every chapter.
Molecular Visions™ Model Kit To support the learning of
organic chemistry concepts and allow students the tactile experience of manipulating physical models, we offer a molecular modeling kit from the Darling Company. The model kit can be bundled
with the textbook or purchased stand alone.

ADDITIONAL INSTRUCTOR
RESOURCES

CONTRIBUTORS TO ORGANIC
CHEMISTRY, 3E

Testbank Prepared by Christine Hermann, Radford University.

I owe special thanks to my contributors for their collaboration,
hard work, and creativity. Many of the new, literature-based,
SkillBuilder problems were written by Laurie Starkey, California
State Polytechnic University, Pomona; Tiffany Gierasch, University
of Maryland, Baltimore County, Seth Elsheimer, University of
Central Florida; and James Mackay, Elizabethtown College. Sections
2.11 and 19.10 were rewritten by Laurie Starkey, and Section
2.12 was written by Tiffany Gierasch. Many of the new Medically
Speaking and Practically Speaking applications throughout the
text were written by Ron Swisher, Oregon Institute of Technology.

PowerPoint Lecture Slides with Answer Slides Prepared by

Adam Keller, Columbus State Community College.
PowerPoint Art Slides Prepared by Kevin Minbiole, Villanova
University.
Personal Response System (“Clicker”) Questions Prepared
by Dalila Kovacs, Grand Valley State University and Randy
Winchester, Grand Valley State University.

STUDENT RESOURCES
(ISBN
9781118700815) Authored by David Klein. The third edition
of the Student Study Guide and Solutions Manual to accompany
Organic Chemistry, 3e contains:
• More detailed explanations within the solutions for every
problem.

Student

Study

Guide

and

Solutions

Manual

ACKNOWLEDGMENTS
The feedback received from both faculty and students supported
the creation, development, and execution of each edition of
Organic Chemistry. I wish to extend sincere thanks to my colleagues
(and their students) who have graciously devoted their time to offer
valuable comments that helped shape this textbook.

ThiRD Edition Reviewers: Class Test Participants,
Focus Group Participants, and Accuracy Checkers
Reviewers
A l aba m a  Rita Collier, Gadsden State
Community College; Anne Gorden, Auburn
University; Eta Isiorho, Auburn University;
Donna Perygin, Jacksonville State University;
Kevin Shaughnessy, The University of Alabama;
Cynthia Tidwell, University of Montevallo;
Stephen Woski, The University of Alabama

Cindy Browder, Northern
Arizona University; Smitha Pillai, Arizona State
University

Cory Antonakos, Diablo
Valley College; Stephen Corlett, Laney College;
Kay Dutz, Mt. San Antonio College; Jason
Hein, University of California, Merced; Carl
Hoeger, University of California, San Diego;
Peggy Kline, Santa Monica College; Megan
McClory, Stanford University; Wayne Pitcher,
Chabot College; Ming Tang, University of
California, Riverside; John Toivonen, Santa
Monica College; William Trego, Laney College;
Erik Woodbury, De Anza College

Martin Campbell, Henderson
State University; Kevin Stewart, Harding
University

C o l o r ad o David Anderson, University
of Colorado, Colorado Springs; Alex Leontyev,
Adams State University

A r i z o n a 

Ark ansas

C a l i f o r n i a 

Del awa re

of Delaware

Bruce Hietbrink, University

F l o r i da  Eric Ballard, University of

Tampa; Edie Banner, University of South
Florida, Sarasota; Adam Braunschweig,
University of Miami; Deborah Bromfield
Lee, Florida Southern College; David Brown,
Florida Gulf Coast University; Mapi Cuevas,
Santa Fe College; Andrew Frazer, University of
Central Florida; Salvatore Profeta, Florida State
University; Bobby Roberson, Pensacola State
College; Christine Theodore, The University of
Tampa

xvi   PREFACE
David Boatright, University
of West Georgia; David Goode, Mercer
University; Shainaz Landge, Georgia Southern
University; David Pursell, Georgia Gwinnett
College; Caroline Sheppard, Clayton State
University; Joseph Sloop, Georgia Gwinnett
College; Michele Smith, Georgia Southwestern
State University; Nina Weldy, Kennesaw State
University

G e o r g i a 

Steve Gentemann, Southwestern
Illinois College; Valerie Keller, University of
Chicago; Jennifer Van Wyk, Southwestern
Illinois College

Illinois

Adam Azman, Butler University;
Jason Dunham, Ball State University; Ryan
Jeske, Ball State University; LuAnne McNulty,
Butler University; Cathrine Reck, Indiana
University

I n d i a n a 

I o w a 

John Gitua, Drake University

O h i o James Beil, Lorain County
Community College
K e n t u c k y Rebecca Brown, West
Kentucky Community and Technical College;
Tanea Reed, Eastern Kentucky University;
Ashley Steelman, University of Kentucky
L o u i s i a n a 

University
Ma i n e

Scott Grayson, Tulane

Richard Broene, Bowdin College

Benjamin Norris, Frostburg
State University; Mark Perks, University of
Maryland, Baltimore County; Emerald Wilson,
Prince George’s Community College

Ma r y l a n d 

Jeremy Andreatta,
Worcester State University; Rich Gurney,

Ma s s ac h u s e t t s

Simmons College; Robert Stolow, Tufts
University

M i c h i ga n Michael Fuertes, Monroe
County Community College; James Kiddle,

Western Michigan University; Jill Morris,
Grand Valley State University; Anja Mueller,
Central Michigan University; Michael Rathke,
Michigan State University
M i s s o u r i Gautam Bhattacharyya,
Missouri State University; Brian Ganley,
University of Missouri, Columbia; Reni Joseph,
St. Louis Community College; Anne Moody,
Truman State University; Vidyullata Waghulde,
St. Louis Community College, Meramec
M o n ta n a 

State University
N e b r a s k a 

University

Kristian Schlick, Montana
James Fletcher, Creighton

N e w Y o r k Martin Di Grandi, Fordham
University; Pamela Kerrigan, College of Mount
Saint Vincent; Ruben Savizky, Cooper Union;
Lucas Tucker, Siena College; Stephen Zawacki,
Erie Community College - North
N o r t h C a r o l i n a  Nicole Bennett,
Appalachian State University; Lindsay

Comstock, Wake Forest University; Stacey
Johnson, Western Piedmont Community College;
Angela King, Wake Forest University

N o r t h D a k o ta  Dennis Viernes,

University of Mary

O h i o Judit Beagle, University of Dayton;
James Beil, Lorain County Community College;
Christopher Callam, The Ohio State University;
Adam Keller, Columbus State Community
College; Noel Paul, The Ohio State University;
Joel Shulman, University of Cincinnati; Sharon
Stickley, Columbus State Community College;
Daniel Turner, University of Dayton
P e n n s y l v a n i a  Qi Chen, Slippery Rock
University; Dian He, Holy Family University;
Steven Kennedy, Millersville University of
Pennsylvania; George Lengyel, Slippery Rock
University; James MacKay, Elizabethtown

This book could not have been created without the incredible
efforts of the following people at John Wiley & Sons, Inc. Photo
Editor Billy Ray helped identify exciting photos. Tom Nery conceived of a visually refreshing and compelling interior design and
cover. Senior Production Editor Elizabeth Swain kept this book
on schedule and was vital to ensuring such a high-quality product.
Joan Kalkut, Sponsoring Editor, was invaluable in the creation of
each edition of this book. Her tireless efforts, together with her
day-to-day guidance and insight, made this project possible. Sean
Hickey, Product Designer, conceived of and built a compelling
WileyPLUS course. Executive Marketing Manager Kristine Ruff
enthusiastically created an exciting message for this book. Mallory
Fryc, Associate Development Editor, managed the review and

College; Kevin Minbiole, Villanova University;
Ernie Trujillo, Wilkes University
R h o d e I s l a n d  Andrew Karatjas,
Johnson and Wales University
S o u t h C a r o l i n a  Tim Barker,

College of Charleston

Tennessee

College

Charity Brannen, Baptist

T e x a s Ashley Ayers, Tarrant County
College, SE Campus; John J. Coniglio, Tarrant
County College; Frank Foss, University of
Texas, Arlington; Martha Gilchrist, Tarrant
County College; Kenn Harding, Texas A&M
University; Drew Murphy, Northeast Texas
Community College; Phillip Pelphrey, Texas
Wesleyan University; Claudia Taenzler,
University of Texas, Dallas; Sammer Tearli,
Collin College; Greg Wilson, Dallas Baptist
University
U ta h

University

Mackay Steffensen, Southern Utah

Kerry Breno, Whitworth
University; Jeffrey Engle, Tacoma Community
College; Trisha Russell, Whitworth University

Wa s h i n g t o n

W i s c o n s i n David Brownholland,
Carthage College; Brian Esselman, University of
Wisconsin State
C a n ada  Mike Chong, University of
Waterloo; Isabelle Dionne, Dawson College;
Bryan Hill, Brandon University; Philip Hultin,
University of Manitoba; Anne Johnson, Ryerson
University; Jimmy Lowe, British Columbia
Institue of Technology; Isabel Molina, Algoma
University; Scott Murphy, University of Regina;
John Sorensen, University of Manitoba; Jackie
Stewart, University of British Columbia;
Christopher Wilds, Concordia University;
Vincent Ziffle, First Nations University of
Canada

s­ upplements process. Publisher Petra Recter provided strong vision
and guidance in bringing this book to market. Sladjana Bruno,
Executive Editor, continued the vision and supported the launch
to market.
Despite my best efforts, as well as the best efforts of the reviewers, accuracy checkers, and class testers, errors may still exist. I take
full responsibility for any such errors and would encourage those
using my textbook to contact me with any errors that you may
find.
David R. Klein, Ph.D.
Johns Hopkins University
klein@jhu.edu

A Review of
General Chemistry
ELECTRONS, BONDS, AND MOLECULAR PROPERTIES

1
1.1 Introduction to Organic Chemistry
1.2 The Structural Theory of Matter
1.3 Electrons, Bonds, and Lewis Structures
1.4 Identifying Formal Charges

Did you ever wonder . . .
what causes lightning?

B

1.5 Induction and Polar Covalent Bonds
1.6 Atomic Orbitals
1.7 Valence Bond Theory
1.8 Molecular Orbital Theory

elieve it or not, the answer to this question is still the subject of debate (that’s right … scientists have not yet figured out
everything, contrary to popular belief  ). There are various theories
that attempt to explain what causes the buildup of electric charge in
clouds. One thing is clear, though—lightning involves a flow of electrons. By studying the nature of electrons and how electrons flow, it
is possible to control where lightning will strike. A tall building can
be protected by installing a lightning rod (a tall metal column at the
top of the building) that attracts any nearby lightning bolt, thereby
preventing a direct strike on the building itself. The lightning rod on
the top of the Empire State Building is struck over a hundred times
each year.
Just as scientists have discovered how to direct electrons in a
bolt of lightning, chemists have also discovered how to direct electrons in chemical reactions. We will soon see that
although organic chemistry is literally defined
as the study of compounds containing carbon atoms, its true essence
is ­actually the study of electrons,
not atoms. Rather than thinking
of reactions in terms of the motion
of atoms, we must recognize that
continued >

1.9 Hybridized Atomic Orbitals
1.10 Predicting Molecular Geometry:
VESPR Theory
1.11 Dipole Moments and Molecular Polarity
1.12 Intermolecular Forces and
Physical Properties
1.13 Solubility

2   CHAPTER

1   A Review of General Chemistry

reactions occur as a result of the motion of electrons. For example, in the following reaction the
curved arrows represent the motion, or flow, of electrons. This flow of electrons causes the
chemical change shown:
HO

⊝

H

H

+

H

C

HO

C

H

+

⊝

H

H

Throughout this course, we will learn how, when, and why electrons flow during
reactions. We will learn about the barriers that prevent electrons from flowing, and
we will learn how to overcome those barriers. In short, we will study the behavioral
­patterns of electrons, enabling us to predict, and even control, the outcomes of chemical
­reactions.
This chapter reviews some relevant concepts from your general chemistry course that
should be familiar to you. Specifically, we will focus on the central role of electrons in forming bonds and influencing molecular properties.

1.1 Introduction to Organic Chemistry
In the early nineteenth century, scientists classified all known compounds into two categories: Organic
compounds were derived from living organisms (plants and animals), while inorganic compounds were
derived from nonliving sources (minerals and gases). This distinction was fueled by the observation
that organic compounds seemed to possess different properties than inorganic compounds. Organic
compounds were often difficult to isolate and purify, and upon heating, they decomposed more readily than inorganic compounds. To explain these curious observations, many scientists subscribed to
a belief that compounds obtained from living sources possessed a special “vital force” that inorganic
compounds lacked. This notion, called vitalism, stipulated that it should be impossible to convert
inorganic compounds into organic compounds without the introduction of an outside vital force.
Vitalism was dealt a serious blow in 1828 when German chemist Friedrich Wöhler demonstrated the
conversion of ammonium cyanate (a known inorganic salt) into urea, a known organic compound
found in urine:
O
NH4OCN

BY THE WAY
There are some
carbon‑containing
compounds that are
traditionally excluded
from organic classification.
For example, ammonium
cyanate (seen on this
page) is still classified as
inorganic, despite the
presence of a carbon
atom. Other exceptions
include sodium carbonate
(Na2CO3) and potassium
cyanide (KCN), both of
which are also considered
to be inorganic compounds.
We will not encounter
many more exceptions.

Ammonium cyanate
(Inorganic)

Heat

H2N

C

NH2

Urea
(Organic)

Over the decades that followed, other examples were found, and the concept of vitalism was
gradually rejected. The downfall of vitalism shattered the original distinction between organic and
inorganic compounds, and a new definition emerged. Specifically, organic compounds became
defined as those compounds containing carbon atoms, while inorganic compounds generally were
defined as those compounds lacking carbon atoms.
Organic chemistry occupies a central role in the world around us, as we are surrounded by
organic compounds. The food that we eat and the clothes that we wear are comprised of organic
compounds. Our ability to smell odors or see colors results from the behavior of organic compounds.
Pharmaceuticals, pesticides, paints, adhesives, and plastics are all made from organic compounds. In
fact, our bodies are constructed mostly from organic compounds (DNA, RNA, proteins, etc.) whose
behavior and function are determined by the guiding principles of organic chemistry. The responses
of our bodies to pharmaceuticals are the results of reactions guided by the principles of organic
chemistry. A deep understanding of those principles enables the design of new drugs that fight disease
and improve the overall quality of life and longevity. Accordingly, it is not surprising that organic
chemistry is required knowledge for anyone entering the health professions.

  3

1.2    The Structural Theory of Matter

1.2 The Structural Theory of Matter
In the mid-nineteenth century three individuals, working independently, laid the conceptual foundations for the structural theory of matter. August Kekulé, Archibald Scott Couper, and Alexander
M. Butlerov each suggested that substances are defined by a specific arrangement of atoms. As an
example, consider the following two compounds:
H
H

C

H
O

C

H

H

H

H

Dimethyl ether
Boiling point = –23°C

H

H

C

C

H

H

O

H

Ethanol
Boiling point = 78.4°C

These compounds have the same molecular formula (C2H6O), yet they differ from each other
in the way the atoms are connected—that is, they differ in their constitution. As a result, they
are called constitutional isomers. Constitutional isomers have different physical properties and
different names. The first compound is a colorless gas used as an aerosol spray propellant, while
the second compound is a clear liquid, commonly referred to as “alcohol,” found in alcoholic
beverages.
According to the structural theory of matter, each element will generally form a predictable
number of bonds. For example, carbon generally forms four bonds and is therefore said to be
­tetravalent. Nitrogen generally forms three bonds and is therefore trivalent. Oxygen forms two
bonds and is divalent, while hydrogen and the halogens form one bond and are monovalent
(Figure 1.1).

Figure 1.1
Valencies of some common
elements encountered
in organic chemistry.

Tetravalent

Trivalent

Divalent

C

N

O

Monovalent

H

X

(where X = F, Cl, Br, or )

Carbon generally
forms four bonds.

Nitrogen generally
forms three bonds.

Oxygen generally
forms two bonds.

Hydrogen and halogens
generally form one bond.

SKILLBUILDER
1.1

drawing constitutional isomers of small molecules

LEARN the skill

Draw all constitutional isomers that have the molecular formula C3H8O.

Solution
STEP 1
Determine the valency of
each atom that appears
in the molecular formula.
STEP 2
Connect the atoms of
highest valency, and
place the monovalent
atoms at the periphery.

Begin by determining the valency of each atom that appears in the molecular formula.
Carbon is tetravalent, hydrogen is monovalent, and oxygen is divalent. The atoms with the
highest valency are connected first. So, in this case, we draw our first isomer by connecting
the three carbon atoms, as well as the oxygen atom, as shown below. The drawing is com‑
pleted when the monovalent atoms (H) are placed at the periphery:

C

C

C

O

C

C

C

O

H

H

H

H

C

C

C

H

H

H

O

H

4   CHAPTER

1   A Review of General Chemistry
This isomer (called 1-propanol) can be drawn in many different ways, some of which are
shown here:
H

H

H

O

H

C

C

C

H

H

H

H

3

2

1

H

H

H

H

C

C

C

H

H

H

3

1-Propanol

STEP 3
Consider other ways to
connect the atoms.

2

1

O

H

H

1-Propanol

H

H

H

C

C

C

H

H

H

O

H

3

2

1

O

H

H

H

H

C

C

C

H

H

H

1

1-Propanol

2

3

H

1-Propanol

All of these drawings represent the same isomer. If we number the carbon atoms (C1, C2,
and C3), with C1 being the carbon atom connected to oxygen, then all of the drawings
above show the same connectivity: a three-carbon chain with an oxygen atom attached at
one end of the chain.
Thus far, we have drawn just one isomer that has the molecular formula C3H8O. Other
constitutional isomers can be drawn if we consider other possible ways of connecting the three
carbon atoms and the oxygen atom. For example, the oxygen atom can be connected to C2
(rather than C1), giving a compound called 2-propanol (shown below). Alternatively, the oxy‑
gen atom can be inserted between two carbon atoms, giving a compound called ethyl methyl
ether (also shown below). For each isomer, two of the many acceptable drawings are shown:
H

H
H

H

O

H

C

C

C

1

H

2

H

H

3

H

H

C

1

H

H

H

C

3

C

2

H
O

H

H

H

H

H

C

C

H

H

H
O

C

H

H
H

H

H

C

O

H

H

C

C

H

H

H

Ethyl methyl ether

2-Propanol

If we continue to search for alternate ways of connecting the three carbon atoms and the
oxygen atom, we will not find any other ways of connecting them. So in summary, there are
a total of three constitutional isomers with the molecular formula C3H8O, shown here:
H
H

H

H

H

C

C

C

H

H

H

O

H

Oxygen is connected to C1

H

H

O

H

C

C

C

H

H

H

H

Oxygen is connected to C2

H

H

H

C

C

H

H

H
O

C

H

H

Oxygen is between two carbon atoms

Additional skills (not yet discussed) are required to draw constitutional isomers of com‑
pounds containing a ring, a double bond, or a triple bond. Those skills will be developed in
Section 14.16.

Practice the skill 1.1 Draw all constitutional isomers with the following molecular formula.
(a) C3H7Cl

Apply the skill

(b) C4H10

(c) C5H12

(d) C4H10O

(e) C3H6Cl2

1.2 Chlorofluorocarbons (CFCs) are gases that were once widely used as refrigerants and
propellants. When it was discovered that these molecules contributed to the depletion of
the ozone layer, their use was banned, but CFCs continue to be detected as contaminants
in the environment.1 Draw all of the constitutional isomers of CFCs that have the molecular
formula C2Cl3F3.

need more PRACTICE? Try Problems 1.35, 1.46, 1.47, 1.54

1.3 Electrons, Bonds, and Lewis Structures
What Are Bonds?
As mentioned, atoms are connected to each other by bonds. That is, bonds are the “glue” that hold
atoms together. But what is this mysterious glue and how does it work? In order to answer this question, we must focus our attention on electrons.
The existence of the electron was first proposed in 1874 by George Johnstone Stoney (National
University of Ireland), who attempted to explain electrochemistry by suggesting the existence

1.3   Electrons, Bonds, and Lewis Structures

  5

of a particle bearing a unit of charge. Stoney coined the term electron to describe this particle.
In 1897, J. J. Thomson (Cambridge University) demonstrated evidence supporting the existence of
Stoney’s mysterious electron and is credited with discovering the electron. In 1916, Gilbert Lewis
(University of California, Berkeley) defined a covalent bond as the result of two atoms sharing a pair
of electrons. As a simple example, consider the formation of a bond between two hydrogen atoms:
+

H

Energy

0

–436 kJ/mol

H H
0.74 Å

H

H

Figure 1.2
An energy diagram showing
the energy as a function of the
internuclear distance between
two hydrogen atoms.

BY THE WAY

1 Å = 10−10 meters.

H

H

△H = –436 kJ/mol

H

Each hydrogen atom has one electron. When these electrons are shared to form a bond, there is a
decrease in energy, indicated by the negative value of ΔH. The energy diagram in Figure 1.2 plots
the energy of the two hydrogen atoms as a function of the distance between them. Focus on
the right side of the diagram, which represents the hydrogen atoms separated
by a large distance. Moving toward the left on the diagram, the hydrogen
atoms approach each other, and there are several forces that must
Internuclear distance
be taken into account: (1) the force of repulsion between the two
H +
H negatively charged electrons, (2) the force of repulsion between
the two positively charged nuclei, and (3) the forces of attraction
H
H
between the positively charged nuclei and the negatively charged electrons. As the hydrogen atoms get closer to each other, all of these forces get
H
H
stronger. Under these circumstances, the electrons are capable of moving in such
a way so as to minimize the repulsive forces between them while maximizing their attractive forces with the nuclei. This provides for a net force of attraction, which lowers the energy of
the system. As the hydrogen atoms move still closer together, the energy continues to be lowered
until the nuclei achieve a separation (internuclear distance) of 0.74 angstroms (Å). At that point,
the force of repulsion between the nuclei begins to overwhelm the forces of attraction, causing
the energy of the system to increase if the atoms are brought any closer together. The lowest point
on the curve represents the lowest energy (most stable) state. This state determines both the bond
length (0.74 Å) and the bond strength (436 kJ/mol).

Drawing the Lewis Structure of an Atom
Armed with the idea that a bond represents a pair of shared electrons, Lewis then devised a method
for drawing structures. In his drawings, called Lewis structures, the electrons take ­center stage. We
will begin by drawing individual atoms, and then we will draw Lewis structures for small molecules.
First, we must review a few simple features of atomic structure:
• The nucleus of an atom is comprised of protons and neutrons. Each proton has a charge of
+1, and each neutron is electrically neutral.
• For a neutral atom, the number of protons is balanced by an equal number of electrons,
which have a charge of −1 and exist in shells. The first shell, which is closest to the nucleus,
can contain two electrons, and the second shell can contain up to eight electrons.
• The electrons in the outermost shell of an atom are called the valence electrons. The number of
valence electrons in an atom is identified by its group number in the periodic table (Figure 1.3).
1A
2A

Li

Be

B

C

N

O

F

Ne

Na Mg

Al

Si

P

S

Cl

Ar

Ga Ge As Se Br

Kr

K
Figure 1.3
A periodic table showing
group numbers.

8A

H

Ca

Rb Sr
Cs Ba

3A 4A 5A 6A 7A He

Transition
Metal
Elements

n Sn Sb Te
Tl

Pb

Bi

Po

Xe
At Rn

The Lewis dot structure of an individual atom indicates the number of valence electrons, which
are placed as dots around the periodic symbol of the atom (C for carbon, O for oxygen, etc.). The
­placement of these dots is illustrated in the following SkillBuilder.

6   CHAPTER

1   A Review of General Chemistry

SKILLBUILDER
1.2

drawing the lewis dot structure of an atom

LEARN the skill

Draw the Lewis dot structure of (a) a boron atom and (b) a nitrogen atom.

Solution
STEP 1
Determine the number
of valence electrons.
STEP 2
Place one valence
electron by itself on each
side of the atom.

STEP 3
If the atom has more
than four valence
electrons, the remaining
electrons are paired with
the electrons already
drawn.

(a)	In a Lewis dot structure, only valence electrons are drawn, so we must first determine
the number of valence electrons. Boron belongs to group 3A on the periodic table, and
it therefore has three valence electrons. The periodic symbol for boron (B) is drawn, and
each electron is placed by itself (unpaired) around the B, like this:
B

(b)	Nitrogen belongs to group 5A on the periodic table, and it therefore has five valence
electrons. The periodic symbol for nitrogen (N) is drawn, and each electron is placed by
itself (unpaired) on a side of the N until all four sides are occupied:
N

	Any remaining electrons must be paired up with the electrons already drawn. In the case
of nitrogen, there is only one more electron to place, so we pair it up with one of the four
unpaired electrons (it doesn’t matter which one we choose):
N

Practice the skill 1.3 Draw a Lewis dot structure for each of the following atoms:
(a) Carbon

(b) Oxygen

(c) Fluorine

(d) Hydrogen

(e) Bromine

(f ) Sulfur

(g) Chlorine

(h) Iodine

1.4 Compare the Lewis dot structure of nitrogen and phosphorus and explain why you
might expect these two atoms to exhibit similar bonding properties.
1.5 Name one element that you would expect to exhibit bonding properties similar to
boron. Explain.
1.6 Draw a Lewis structure of a carbon atom that is missing one valence electron (and
therefore bears a positive charge). Which second-row element does this carbon atom resem‑
ble in terms of the number of valence electrons?

Apply the skill

1.7 Lithium salts have been used for decades to treat mental illnesses, including depres‑
sion and bipolar disorder. Although the treatment is effective, researchers are still trying to
determine how lithium salts behave as mood stabilizers.2
(a) Draw a Lewis structure of an uncharged lithium atom, Li.
(b) Lithium salts contain a lithium atom that is missing one valence electron (and therefore
bears a positive charge). Draw a Lewis structure of the lithium cation.

Drawing the Lewis Structure of a Small Molecule
The Lewis dot structures of individual atoms are combined to
H
produce Lewis dot structures of small molecules. These drawings
H
are constructed based on the observation that atoms tend to bond
HC H
H C H
H
in such a way so as to achieve the electron configuration of a
H
noble gas. For example, hydrogen will form one bond to achieve
the electron configuration of helium (two valence electrons), while second-row elements (C, N, O,
and F) will form the necessary number of bonds so as to achieve the electron configuration of neon
(eight valence electrons).

1.3   Electrons, Bonds, and Lewis Structures

  7

This observation, called the octet rule, explains why carbon is tetravalent. As just shown, it can
achieve an octet of electrons by using each of its four valence electrons to form a bond. The octet
rule also explains why nitrogen is trivalent. Specifically, it has five
H N H
HN H
valence electrons and requires three bonds in order to achieve an
H
octet of electrons. Notice that the nitrogen atom contains one pair
H
of unshared, or nonbonding, electrons, called a lone pair.
In the next chapter, we will discuss the octet rule in more detail; in particular, we will explore when
it can be violated and when it cannot be violated. For now, let’s practice drawing Lewis structures.

SKILLBUILDER
1.3

drawing the lewis structure of a small molecule

LEARN the skill

Draw the Lewis structure of CH2O.

Solution
There are four discrete steps when drawing a Lewis structure: First determine the number of
valence electrons for each atom.
STEP 1
Draw all individual
atoms.
STEP 2
Connect atoms that
form more than one
bond.

C

H

O

Then, connect any atoms that form more than one bond. Hydrogen atoms only form
one bond each, so we will save those for last. In this case, we connect the C and the O.
C O

Next, connect all hydrogen atoms. We place the hydrogen atoms next to carbon,
because carbon has more unpaired electrons than oxygen.

STEP 3
Connect the
hydrogen atoms.

STEP 4
Pair any unpaired
electrons so that each
atom achieves an
octet.

H

H C O
H

Finally, check to see if each atom (except hydrogen) has an octet. In fact, ­neither the carbon
nor the oxygen has an octet, so in a situation like this, the unpaired electrons are shared as
a double bond between carbon and oxygen.
H C O
H

H C O
H

Now all atoms have achieved an octet. When drawing Lewis structures, remember that
you cannot simply add more electrons to the drawing. For each atom to achieve an octet,
the existing electrons must be shared. The total number of valence electrons should be
correct when you are finished. In this example, there was one carbon atom, two hydrogen
atoms, and one oxygen atom, giving a total of 12 valence electrons (4 + 2 + 6). The drawing
above MUST have 12 valence electrons, no more and no less.

Practice the skill 1.8 Draw a Lewis structure for each of the following compounds:
(a) C2H6  (b) C2H4  (c) C2H2  (d) C3H8  (e) C3H6  (f ) CH3OH
1.9 Borane (BH3) is very unstable and quite reactive. Draw a Lewis structure of borane and
explain the source of the instability.
1.10 There are four constitutional isomers with the molecular formula C3H9N. Draw a Lewis
structure for each isomer and determine the number of lone pairs on the nitrogen atom in
each case.

Apply the skill

1.11 Smoking tobacco with a water pipe, or hookah, is often perceived as being less
­dangerous than smoking cigarettes, but hookah smoke has been found to contain the same

8   CHAPTER

1   A Review of General Chemistry
variety of toxins and carcinogens (cancer-causing compounds) as cigarette smoke.3 Draw a
Lewis structure for each of the following dangerous compounds found in tobacco smoke:
(a) HCN (hydrogen cyanide)

(b) CH2CHCHCH2 (1,3-butadiene)

need more PRACTICE? Try Problem 1.39

1.4 Identifying Formal Charges
A formal charge is associated with any atom that does not exhibit the appropriate number of valence
electrons. When such an atom is present in a Lewis structure, the formal charge must be drawn.
Identifying a formal charge requires two discrete tasks:
1. Determine the appropriate number of valence electrons for an atom.
2. Determine whether the atom exhibits the appropriate number of electrons.
The first task can be accomplished by inspecting the periodic table. As mentioned earlier, the
group number indicates the appropriate number of valence electrons for each atom. For example,
carbon is in group 4A and therefore has four valence electrons. Oxygen is in group 6A and has six
valence electrons.
O
After identifying the appropriate number of electrons for each atom in a Lewis strucH
C H
ture, the next task is to determine if any of the atoms exhibit an unexpected number of
electrons. For example, consider the following structure.
H

Each line represents two shared electrons (a bond). For our purposes, we must split each
bond apart equally, and then count the number of electrons on each atom.

O
H

C

H

H

Each hydrogen atom has one valence electron, as expected. The carbon atom also has
the appropriate number of valence electrons (four), but the oxygen atom does not. The
oxygen atom in this structure exhibits seven valence electrons, but it should only have six.
In this case, the oxygen atom has one extra electron, and it must therefore bear a negative
formal charge, which is indicated like this.

⊝

O
H

C

H

H

SKILLBUILDER
1.4

calculating formal charge

LEARN the skill

Consider the nitrogen atom in the structure below and determine if it has a formal charge:
H
N

H

H

H

Solution
STEP 1
Determine the
appropriate number
of valence electrons.
STEP 2
Determine the actual
number of valence
electrons in this case.

We begin by determining the appropriate number of valence electrons for a nitrogen atom.
Nitrogen is in group 5A of the periodic table, and it should therefore have five valence
­electrons.
Next, we count how many valence electrons are exhibited by the nitrogen atom in this par‑
ticular example.
H
H

N
H

H

  9

1.5   Induction and Polar Covalent Bonds

In this case, the nitrogen atom exhibits only four valence electrons. It is missing one electron,
so it must bear a positive charge, which is shown like this:

STEP 3
Assign a formal
charge.

H
H

⊕

N

H

H

Practice the skill 1.12 Identify any formal charges in the structures below:
H
H

(a)

Al

H

H
H

(b)

H

H

H

O

H

(c)

C

H
N

H

(f )

C
H

H

H

(g)

H

(d)

H

C

C

Cl

O

Cl

H

Cl

H

Cl

(h)

H

Al

H

C

Cl

H
H

C

H

O

(i)

(e)

H
H

H

O

N

C

C

H

H

C

H

H

O

1.13 Draw a structure for each of the following ions; in each case, indicate which atom
possesses the formal charge:
(a) BH4−    (b) NH2−    (c) C2H5+

Apply the skill

1.14 If you are having trouble paying attention during a long
lecture, your levels of ­acetylcholine (a neurotransmitter) may
be to blame.4 Identify any formal charges in ­acetylcholine.

H

H

O

C

C

O

H

need more PRACTICE? Try Problem 1.41

H

H
H
H
H C
H

C

C

H

H
H C
H
H
H

N

C

H

Acetylcholine

1.5 Induction and Polar Covalent Bonds
Chemists classify bonds into three categories: (1) covalent, (2) polar covalent, and (3) ionic. These
categories emerge from the electronegativity values of the atoms sharing a bond. Electronegativity is
a measure of the ability of an atom to attract electrons. Table 1.1 gives the electronegativity values for
elements commonly encountered in organic chemistry.
TABLE

1.1

ELECTRONEGATIVITY VALUES OF SOME COMMON ELEMENTS

Increasing electronegativity

H

2.1

Li

Be
1.5

2.0

2.5

3.0

3.5

4.0

Na

Mg

Al

Si

P

S

Cl

1.0
0.9

1.2

B

1.5

C

N

1.8

2.1

O

F

2.5

3.0

K

Increasing
electronegativity

Br

0.8

2.8

When two atoms form a bond, one critical consideration allows us to classify the bond:
What is the difference in the electronegativity values of the two atoms? Below are some rough
guidelines:
If the difference in electronegativity is less than 0.5, the electrons are considered to be
equally shared between the two atoms, resulting in a covalent bond. Examples include C−C and
C−H:
C

C

C

H

10   CHAPTER

1   A Review of General Chemistry

The C−C bond is clearly covalent, because there is no difference in electronegativity between the
two atoms forming the bond. Even a C−H bond is considered to be covalent, because the difference
in electronegativity between C and H is less than 0.5.
If the difference in electronegativity is between 0.5 and 1.7, the electrons are not shared equally
between the atoms, resulting in a polar covalent bond. For example, consider a bond between carbon and oxygen (C−O). Oxygen is significantly more electronegative (3.5) than carbon (2.5), and
therefore oxygen will more strongly attract the electrons of the bond. The withdrawal of electrons
toward oxygen is called induction, which is often indicated with an arrow like this.
C

O

Induction causes the formation of partial positive and partial negative charges, symbolized by
the Greek symbol delta (δ). The partial charges that result from induction will be very important in
upcoming chapters.
δ+

C

δ–

O

If the difference in electronegativity is greater than 1.7, the electrons are not shared at all. For
example, consider the bond between sodium and oxygen in sodium ­hydroxide (NaOH).
⊝

⊕

Na

OH

The difference in electronegativity between O and Na is so great that both electrons of the bond are
possessed solely by the oxygen atom, rendering the oxygen negatively charged and the sodium positively charged. The bond between oxygen and sodium, called an ionic bond, is the result of the force
of attraction between the two oppositely charged ions.
The cutoff numbers (0.5 and 1.7) should be thought of as rough guidelines. Rather than viewing
them as absolute, we must view the various types of bonds as belonging to a spectrum without clear
cutoffs (Figure 1.4).
Covalent
Figure 1.4
The nature of various bonds
commonly encountered
in organic chemistry.

C

C

C

Polar covalent
H

N

H

C

O

Li

Ionic
C

Li

Small difference
in electronegativity

N

NaCl

Large difference
in electronegativity

This spectrum has two extremes: covalent bonds on the left and ionic bonds on the right. Between
these two extremes are the polar covalent bonds. Some bonds fit clearly into one category, such as
C−C bonds (covalent), C−O bonds (polar covalent), or NaCl bonds (ionic). However, there are
many cases that are not so clear-cut. For example, a C−Li bond has a difference in electronegativity
of 1.5, and this bond is often drawn either as polar covalent or as ionic. Both drawings are acceptable:
C

Li

or

C

⊝ ⊕

Li

Another reason to avoid absolute cutoff numbers when comparing electronegativity values is that
the electronegativity values shown above are obtained via one particular method developed by Linus
Pauling. However, there are at least seven other methods for calculating electronegativity values, each of
which provides slightly different values. Strict adherence to the Pauling scale would suggest that C−Br
and C−I bonds are covalent, but these bonds will be treated as polar covalent throughout this course.

SKILLBUILDER
1.5

locating partial charges resulting from induction

LEARN the skill

Consider the structure of methanol. Identify all polar covalent bonds
and show any partial charges that result from inductive effects:

H
H

C

O

H
Methanol

H

  11

1.5   Induction and Polar Covalent Bonds

Solution
First identify all polar covalent bonds. The C−H bonds are considered to be covalent
because the electronegativity values for C and H are fairly close. It is true that carbon is more
electronegative than hydrogen, and therefore, there is a small inductive effect for each C−H
bond. However, we will generally consider this effect to be negligible for C−H bonds.
The C−O bond and the O−H bond are both polar covalent bonds:
STEP 1
Identify all polar
covalent bonds.

H
H

C
H

O

H

Polar covalent

Now determine the direction of the inductive effects. Oxygen is more electronegative than
C or H, so the inductive effects are shown like this:
STEP 2
Determine the
direction of each
dipole.

H
H

C

O

H

H

STEP 3
Indicate the location
of partial charges.

These inductive effects dictate the locations of the partial charges:
H
H

δ+ δ–

C

O

δ+

H

H

Practice the skill 1.15 For each of the following compounds, identify any polar covalent bonds by drawing
δ+ and δ− symbols in the appropriate locations:
H
H

(a)

C

O

H

H

H

C

C

H

H

H
O

C

H

F
H

H

H

(b)

C

H

Cl

(c)

H

C

Mg

Br

H

H

H

(d)

Apply the skill

O

H

O

H

C

C

C

H

H

O

H

H

H

(e)

H

O

H

C

C

C

H

Cl
Cl

H

(f )

H

1.16 The regions of δ+ in a compound are the regions most
likely to be attacked by an anion, such as hydroxide (HO−). In the
compound shown, identify the two carbon atoms that are most
likely to be attacked by a hydroxide ion.

H

C

Cl

Cl

H

O

H

H

H

C

C

C

C

C

H

H

H

H

Cl

1.17 Plastics and synthetic fibers are examples of the many materials made from repea­ting
subunits of carbon-containing molecules called polymers. Although most synthetic polymers
are prepared from fossil fuel sources, many researchers are exploring
H H
ways to make polymers from renewable sources instead. One example is
O
C
H
the synthesis of an epoxy resin polymer using a by-product from cashew
Cl
C
C
nut processing, another compound isolated from corn cobs, and epichlo‑
H
H
rohydrin, shown here.5 Identify any polar covalent bonds in epichlorohy‑
Epichlorohydrin
drin by drawing δ+ and δ− symbols in the appropriate ­locations.
need more PRACTICE? Try Problems 1.37, 1.38, 1.48, 1.57

12   CHAPTER

1   A Review of General Chemistry

Practically Speaking Electrostatic Potential Maps
Partial charges can be visualized with three-dimensional,
rainbow-like images called electrostatic potential maps. As an
example, consider the following electrostatic potential map of
chloromethane:
Most negative
(δ−)

δ−
δ−

Cl
H

C

δ+

H

H

Chloromethane

δ+

Electrostatic
potential map
of chloromethane

Most positive
(δ+)
Color scale

In the image, a color scale is used to represent areas of δ− and
δ+. As indicated, red represents a region that is δ−, while blue

represents a region that is δ+. In reality, electrostatic potential
maps are rarely used by practicing organic chemists when they
communicate with each other; however, these illustrations can
often be helpful to students who are learning organic chemistry.
Electrostatic potential maps are generated by performing a
series of calculations. Specifically, an imaginary point positive
charge is positioned at various locations, and for each location,
we calculate the potential energy associated with the attraction
between the point positive charge and the surrounding
electrons. A large attraction indicates a position of δ−, while a
small attraction indicates a position of δ+. The results are then
illustrated using colors, as shown.
A comparison of any two electrostatic potential maps is only
valid if both maps were prepared using the same color scale.
Throughout this book, care has been taken to use the same
color scale whenever two maps are directly compared to each
other. However, it will not be useful to compare two maps from
different pages of this book (or any other book), as the exact
color scales are likely to be different.

1.6 Atomic Orbitals
Quantum Mechanics
By the 1920s, vitalism had been discarded. Chemists were aware of constitutional isomerism and
had developed the structural theory of matter. The electron had been discovered and identified as the
source of bonding, and Lewis structures were used to keep track of shared and unshared electrons.
But the understanding of electrons was about to change dramatically.
In 1924, French physicist Louis de Broglie suggested that electrons, heretofore considered as
particles, also exhibited wavelike properties. Based on this assertion, a new theory of matter was born.
In 1926, Erwin Schrödinger, Werner Heisenberg, and Paul Dirac independently proposed a mathematical description of the electron that incorporated its wavelike properties. This new theory, called
wave mechanics, or quantum mechanics, radically changed the way we viewed the nature of matter
and laid the foundation for our current understanding of electrons and bonds.
Quantum mechanics is deeply rooted in mathematics and represents an entire subject by itself.
The mathematics involved is beyond the scope of our course, and we will not discuss it here. However,
in order to understand the nature of electrons, it is critical to understand a few simple highlights from
quantum mechanics:
• An equation is constructed to describe the total energy of a hydrogen atom (i.e., one proton
plus one electron). This equation, called the wave equation, takes into account the wavelike
behavior of an electron that is in the electric field of a proton.
• The wave equation is then solved to give a series of solutions called wavefunctions. The Greek
symbol psi (ψ) is used to denote each wavefunction (ψ1, ψ2, ψ3, etc.). Each of these wavefunctions corresponds to an allowed energy level for the electron. This result is incredibly important because it suggests that an electron, when contained in an atom, can only exist at discrete
energy levels (ψ1, ψ2, ψ3, etc.). In other words, the energy of the electron is quantized.
• Each wavefunction is a function of spatial location. It provides information that allows us to
assign a numerical value for each location in three-dimensional space relative to the nucleus.
The square of that value (ψ2 for any particular location) has a special meaning. It indicates the
probability of finding the electron in that location. Therefore, a three-dimensional plot of ψ2
will generate an image of an atomic orbital (Figure 1.5).

  13

1.6   Atomic Orbitals

y

y
z

y
z

x

y
z

x

z

x

x

Figure 1.5
Illustrations of an s orbital
and three p orbitals.

Electron Density and Atomic Orbitals
An orbital is a region of space that can be occupied by an electron. But care must be taken when trying to visualize this. There is a statement from the previous section that must be clarified because it
is potentially misleading: “ψ2 represents the probability of finding an electron in a particular location.”
This statement seems to treat an electron as if it were a particle flying around within a specific region
of space. But remember that an electron is not purely a particle—it has wavelike properties as well.
Therefore, we must construct a mental image that captures both of these properties. That is not easy to
do, but the following analogy might help. We will treat an occupied orbital as if it is a cloud—similar to
a cloud in the sky. No analogy is perfect, and there are certainly features of clouds that are very different
from orbitals. However, focusing on some of these differences between electron clouds (occupied orbitals) and real clouds makes it possible to construct a better mental model of an electron in an orbital:
• Clouds in the sky can come in any shape or size. However, electron clouds have specific shapes
and sizes (as defined by the orbitals).
• A cloud in the sky is comprised of billions of individual water molecules. An electron cloud
is not comprised of billions of particles. We must think of an electron cloud as a single entity,
even though it can be thicker in some places and thinner in other places. This concept is critical and will be used extensively throughout the course in explaining reactions.
• A cloud in the sky has edges, and it is possible to define a region of space that contains 100%
of the cloud. In contrast, an electron cloud does not have defined edges. We frequently use
the term electron density, which is associated with the probability of finding an electron in
a particular region of space. The “shape” of an orbital refers to a region of space that contains
90–95% of the electron density. Beyond this region, the remaining 5–10% of the electron
density tapers off but never ends. In fact, if we want to consider the region of space that contains 100% of the electron density, we must consider the entire universe.
In summary, we must think of an orbital as a region of space that can be occupied by electron
density. An occupied orbital must be treated as a cloud of electron density. This region of space is called
an atomic orbital (AO), because it is a region of space defined with respect to the nucleus of a single
atom. Examples of atomic orbitals are the s, p, d, and f orbitals that were discussed in your general
chemistry textbook.

Phases of Atomic Orbitals
Our discussion of electrons and orbitals has been based on the premise that electrons have wavelike
properties. As a result, it will be necessary to explore some of the characteristics of simple waves in
order to understand some of the characteristics of orbitals.
Consider a wave that moves across the surface of a lake (Figure 1.6). The wavefunction (ψ) mathematically describes the wave, and the value of the wavefunction is dependent on location. Locations

Average level
of lake
Figure 1.6
Phases of a wave moving
across the surface of a lake.

ψ is (+)

ψ is (+)
ψ is (–)
Node

ψ=0

ψ is (–)

14   CHAPTER

1   A Review of General Chemistry

ψ is (+)

Node

ψ is (–)

Figure 1.7
The phases of a p orbital.

above the average level of the lake have a positive value for ψ (indicated in red), and locations below the
average level of the lake have a negative value for ψ (indicated in blue). Locations where the value of ψ is
zero are called nodes.
Similarly, orbitals can have regions where the value of ψ is positive, negative, or zero. For example, consider a p orbital (Figure 1.7). Notice that the p orbital has two lobes: The top lobe is a region
of space where the values of ψ are positive, while the bottom lobe is a region where the values of ψ are
negative. Between the two lobes is a location where ψ = 0. This location represents a node.
Be careful not to confuse the sign of ψ (+ or −) with electrical charge. A positive value for ψ
does not imply a positive charge. The value of ψ (+ or −) is a mathematical convention that refers
to the phase of the wave (just like in the lake). Although ψ can have positive or negative values, nevertheless ψ2 (which describes the electron density as a function of location) will always be a positive
number. At a node, where ψ = 0, the electron density (ψ2) will also be zero. This means that there is
no electron density located at a node.
From this point forward, we will draw the lobes of an orbital with colors (red and blue) to indicate the phase of ψ for each region of space.

Filling Atomic Orbitals with Electrons
The energy of an electron depends on the type of orbital that it occupies. Most of the organic compounds that we will encounter will be composed of first- and second-row elements (H, C, N, and O).
These elements utilize the 1s orbital, the 2s orbital, and the three 2p orbitals. Our discussions will
­therefore focus primarily on these orbitals (Figure 1.8). Electrons are lowest in energy when they occupy
a 1s orbital, because the 1s orbital is closest to the nucleus and it has no nodes (the more nodes that an
orbital has, the greater its energy). The 2s orbital has one node and is farther away from the nucleus; it is
therefore higher in energy than the 1s orbital. After the 2s orbital, there are three 2p orbitals that are all
equivalent in energy to one another. Orbitals with the same energy level are called degenerate orbitals.
y

y
z

z

x
Figure 1.8
Illustrations of s orbitals
and three p orbitals.

1s

y

z

z

x

2s

y

y

x

z

x

2py

2px

x

2pz

As we move across the periodic table, starting with hydrogen, each element has one more
e­ lectron than the element before it (Figure 1.9). The order in which the orbitals are filled by electrons
is determined by just three simple principles:
1. The Aufbau principle. The lowest energy orbital is filled first.
2. The Pauli exclusion principle. Each orbital can accommodate a maximum of two electrons that
have opposite spin. To understand what “spin” means, we can imagine an electron spinning in
space (although this is an oversimplified explanation of the term “spin”). For reasons that are
beyond the scope of this course, electrons only have two possible spin states (designated by ⇃ or ↾).
In order for the orbital to accommodate two electrons, the electrons must have opposite spin states.

2p

Figure 1.9
Energy diagrams showing
the electron configurations
for H, He, Li, and Be.

Energy
1s
Hydrogen

1s
Helium

2p

2s

2s

1s
Lithium

1s
Beryllium

  15

1.6   Atomic Orbitals

3. Hund’s rule. When dealing with degenerate orbitals, such as p orbitals, one electron is placed
in each degenerate orbital first, before electrons are paired up.
The application of the first two principles can be seen in the electron configurations shown in Figure
1.9 (H, He, Li, and Be). The application of the third principle can be seen in the electron configurations for the remaining second-row elements (Figure 1.10).

Energy

2p

2p

2p

2p

2p

2s

2s

2s

2s

2s

1s

1s
Carbon

1s
Nitrogen

1s
Oxygen

1s
Fluorine

Boron

2p
2s
1s
Neon

Figure 1.10
Energy diagrams showing the electron configurations for B, C, N, O, F, and Ne.

SKILLBUILDER
1.6

identifying electron configurations

LEARN the skill

Identify the electron configuration of a nitrogen atom.

Solution
STEP 1
Place the valence
electrons in atomic
orbitals using the
Aufbau principle,
the Pauli exclusion
principle, and
Hund’s rule.

The electron configuration indicates which atomic orbitals are occupied by electrons.
Nitrogen has a total of seven electrons. These electrons occupy atomic orbitals of increasing
energy, with a maximum of two electrons in each orbital:
2p
2s
1s
Nitrogen

STEP 2
Identify the number
of valence electrons in
each atomic orbital.

Two electrons occupy the 1s orbital, two electrons occupy the 2s orbital, and three electrons
occupy the 2p orbitals. This is summarized using the following notation:

1s22s22p3

Practice the skill 1.18 Identify the electron configuration for each of the following atoms:
(a) Carbon  (b) Oxygen  (c) Boron  (d) Fluorine  (e) Sodium  (f ) Aluminum
1.19

Apply the skill

Identify the electron configuration for each of the following ions:

(a) A carbon atom with a negative charge

(c) A nitrogen atom with a positive charge

(b) A carbon atom with a positive charge

(d) An oxygen atom with a negative charge

1.20 Silicon is the second most abundant element in the Earth's crust, and its compounds
can be as ordinary as beach sand. However, silicon also plays an indispensable role in modern
devices such as computers, cell phones, semiconductors, and solar panels. A recent technol‑
ogy incorporates silicon in nanometer-sized particles called quantum dots that act as lumines‑
cent labels for pancreatic cancer cells.6 Identify the electron configuration of a silicon atom.

need more PRACTICE? Try Problem 1.44

16   CHAPTER

1   A Review of General Chemistry

1.7 Valence Bond Theory
With the understanding that electrons occupy regions of space called orbitals, we can now turn our attention to a deeper understanding of covalent bonds. Specifically, a covalent bond is formed from the overlap
of atomic orbitals. There are two commonly used theories for describing the nature of atomic orbital
overlap: valence bond theory and molecular orbital (MO) theory. The valence bond approach is more
simplistic in its treatment of bonds, and therefore we will begin our discussion with valence bond theory.
If we are going to treat electrons as waves, then we must quickly review what happens when
two waves interact with each other. Two waves that approach each other can interfere in one of two
possible ways—constructively or destructively. Similarly, when atomic orbitals overlap, they can
interfere either constructively (Figure 1.11) or destructively (Figure 1.12).
An electron
is like a wave
Figure 1.11
Constructive interference
resulting from the interaction
of two electrons.

An electron
is like a wave

Bring these
waves closer
together...

...and the
waves reinforce
each other

Constructive
interference

Internuclear
distance

Internuclear distance

Constructive interference produces a wave with larger amplitude. In contrast, destructive interference results in waves canceling each other, which produces a node (Figure 1.12).
Bring these
waves closer
together...

Figure 1.12
Destructive interference
resulting from the interaction
of two electrons.

...and the
waves cancel
each other

A node

Destructive
interference

According to valence bond theory, a bond is simply the sharing of electron density between two
atoms as a result of the constructive interference of their atomic orbitals. Consider, for example, the
bond that is formed between the two hydrogen atoms in molecular hydrogen (H2). This bond is
formed from the overlap of the 1s orbitals of each hydrogen atom (Figure 1.13).
The electron density of this bond is primarily located on the bond axis (the line that can be
drawn between the two hydrogen atoms). This type of bond is called a sigma (σ) bond and is characterized by circular symmetry with respect to the bond axis. To visualize what this means, imagine
a plane that is drawn perpendicular to the bond axis. This plane will carve out a circle (Figure 1.14).
This is the defining feature of σ bonds and will be true of all purely single bonds. Therefore, all single
bonds are σ bonds.

Circular
cross section

+

Figure 1.13
The overlap of the 1s atomic orbitals of two hydrogen
atoms, forming molecular hydrogen (H2).

Figure 1.14
An illustration of a sigma bond,
showing the circular symmetry
with respect to the bond axis.

1.8    Molecular Orbital Theory

  17

1.8 Molecular Orbital Theory
In most situations, valence bond theory will be sufficient for our purposes. However, there will be
cases in the upcoming chapters where valence bond theory will be inadequate to describe the observations. In such cases, we will utilize molecular orbital theory, a more sophisticated approach to viewing
the nature of bonds.
Molecular orbital (MO) theory uses mathematics as a tool to explore the consequences of atomic
orbital overlap. The mathematical method is called the linear combination of atomic orbitals (LCAO).
According to this theory, atomic orbitals are mathematically combined to produce new orbitals, called
molecular orbitals.
It is important to understand the distinction between atomic orbitals and molecular orbitals. Both
types of orbitals are used to accommodate electrons, but an atomic orbital is a region of space associated
with an individual atom, while a molecular orbital is associated with an entire molecule. That is, the
molecule is considered to be a single entity held together by many electron clouds, some of which can
actually span the entire length of the molecule. These molecular orbitals are filled with electrons in a
particular order in much the same way that atomic orbitals are filled. Specifically, electrons first occupy
the lowest energy orbitals, with a maximum of two electrons per orbital. In order to visualize what it
means for an orbital to be associated with an entire molecule, we will explore two molecules: molecular
hydrogen (H2) and bromomethane (CH3Br).
Consider the bond formed between the two hydrogen atoms in molecular hydrogen. This
bond is the result of the overlap of two atomic orbitals (s orbitals), each of which is occupied by one
electron. According to MO theory, when two atomic orbitals overlap, they cease to exist. Instead,
they are replaced by two molecular orbitals, each of which is associated with the entire molecule
(Figure 1.15).
Node
Antibonding MO

Energy

1s

1s

Bonding MO
Figure 1.15
An energy diagram showing the relative energy levels of bonding and antibonding
molecular orbitals.

Figure 1.16
A low-energy molecular
orbital of CH3Br. Red and
blue regions indicate the
different phases, as described
in Section 1.6. Notice that this
molecular orbital is associated
with the entire molecule,
rather than being associated
with two specific atoms.

In the energy diagram shown in Figure 1.15, the individual atomic orbitals are represented
on the right and left, with each atomic orbital having one electron. These atomic orbitals are
combined mathematically (using the LCAO method) to produce two molecular orbitals. The
lower energy molecular orbital, or bonding MO, is the result of constructive interference of
the original two atomic orbitals. The higher energy molecular orbital, or antibonding MO, is
the result of destructive interference. Notice that the antibonding MO has one node, which
explains why it is higher in energy. Both electrons occupy the bonding MO in order to achieve a
lower energy state. This lowering in energy is the essence of the bond. For an H−H bond, the
lowering in energy is equivalent to 436 kJ/mol. This energy corresponds with the bond strength
of an H−H bond (as shown in Figure 1.2).
Now let’s consider a molecule such as CH3Br, which contains more than just one bond.
Valence bond theory continues to view each bond separately, wi