Principal Organic Chemistry
Organic ChemistryDavid R. Klein
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.
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This page intentionally left blank 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 or by any means, electronic, mechanical, photocopying, recording, scanning, or otherwise, except as permitted under Sections 107 or 108 of the 1976 United States Copyright Act, without either the prior written permission of the Publisher or authorization through payment of the appropriate per-copy fee to the Copyright Clearance Center, 222 Rosewood Drive, Danvers, MA 01923, (978) 750-8400, fax (978) 646-8600. Requests to the Publisher for permission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 111 River Street, Hoboken, NJ 07030-5774, (201) 748-6011, fax (201) 748-6008. Evaluation copies are provided to qualiﬁed academics and professionals for review purposes only, for use in their courses during the next academic year. These copies are licensed and may not be sold or transferred to a third party. Upon completion of the review period, please return the evaluation copy to Wiley. Return instructions and a free of charge return shipping label are available at www. wiley.com/go/returnlabel. Outside of the United States, please contact your local representative. 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. 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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. 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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 firstname.lastname@example.org 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, with each bond being formed from two overlapping atomic orbitals. In contrast, MO theory treats the bonding electrons as being associated with the entire molecule. The molecule has many molecular orbitals, each of which can be occupied by two electrons. Figure 1.16 illustrates one of the many molecular 18 CHAPTER 1 A Review of General Chemistry orbitals of CH3Br. This molecular orbital is capable of accommodating up to two electrons. Red and blue regions indicate the different phases, as described in Section 1.6. As we saw with molecular hydrogen, not all molecular orbitals will be occupied. The bonding electrons will occupy the lower energy molecular orbitals (such as the one shown in Figure 1.16), while the higher energy molecular orbitals remain unoccupied. For every molecule, two of its molecular orbitals will be of particular interest: (1) the highest energy orbital from among the occupied orbitals is called the highest occupied molecular orbital, or