Marchs advanced organic chemistry 7e by michael b smith 1

HMO H€uckel molecular orbital1H NMR Proton nuclear magnetic resonance spectroscopy HOMO Highest occupied molecular orbital IUPAC International Union of Pure and Applied Chemistry LCAO Li

MARCH’S ADVANCED ORGANIC CHEMISTRY

MARCH’S ADVANCED

ORGANIC CHEMISTRY

REACTIONS, MECHANISMS, AND STRUCTURE

SEVENTH EDITION

Michael B Smith

Professor of Chemistry

Copyright # 2013 by John Wiley & Sons, Inc All rights reserved.

Published by John Wiley & Sons, Inc., Hoboken, New Jersey.

Published simultaneously in Canada.

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Library of Congress Cataloging-in-Publication Data:

Smith, Michael, 1946 Oct 17- March’s Advanced Organic Chemistry : Reactions, Mechanisms, and Structure – 7th Edition / Michael B Smith, Professor of Chemistry.

10 9 8 7 6 5 4 3 2 1

2.B Bond Energies and Distances in Compounds Containing

v

2.K Aromatic Systems with Electron Numbers other than Six 65

4.J.iii cis–trans Isomerism of Fused and Bridged Ring Systems 167

vi CONTENTS

4.N.ii Conformation in Six-Membered Rings 1804.N.iii Conformation in Six-Membered Rings Containing Heteroatoms 186

CONTENTS vii

7 Irradiation Processes in Organic Chemistry 289

7.A.vii The Fate of the Excited Molecule: Chemical Processes 301

9.C Quantitative Treatments of the Effect of Structure on Reactivity 352

viii CONTENTS

10.C The Neighboring-Group Mechanism 39110.C.i Neighboring-Group Participation byp and s Bonds:

10.E Nucleophilic Substitution at an Allylic Carbon: Allylic

10.F Nucleophilic Substitution at an Aliphatic Trigonal Carbon:

10.G.vii Ambident (Bidentant) Nucleophiles: Regioselectivity 446

11.B.i Orientation and Reactivity in Monosubstituted

11.B.iv Orientation in Benzene Rings with More Than One Substituent 583

11.D A Quantitative Treatment of Reactivity of the Electrophile: The Selectivity

12 Aliphatic, Alkenyl, and Alkynyl Substitution, Electrophilic

12.A.iii Electrophilic Substitution Accompanied by Double-Bond Shifts 657

14.A.iv Neighboring-Group Assistance in Free Radical Reactions 810

x CONTENTS

14.C Reactions 821

16.A.i Nucleophilic Substitution at an Aliphatic Trigonal Carbon: The

16.B.i Reactions in which Hydrogen or a Metallic Ion Adds to the

CONTENTS xi

17.C Stereochemistry of the Double Bond 1273

17.F.i Reactions in which CC and CC Bonds are Formed 1282

17.F.iii Reactions in which CN or CN Bonds are Formed 1310

to make room for new ones, and in cases where a series of papers by the same principalauthor were cited, all but the most recent were deleted The older citations are usuallyfound by referring to the more recent publication(s) Many of the figures relating tomolecular orbitals dated to the 1960s In all cases possible, they have been replaced bymolecular orbitals drawings using Spartan software from Wavefunction, Inc The funda-mental structure of the seventh edition is essentially the same as that of all previous ones.The goal, as in previous editions is to give equal weight to the three fundamental aspects

of the study of organic chemistry: reactions, mechanisms, and structure A student who hascompleted a course based on this book should be able to approach the literature directly,with a sound knowledge of modern organic chemistry Major special areas of organicchemistry: terpenes, carbohydrates, proteins, many organometallic reagents, combinatorialchemistry, polymerization and electrochemical reactions, steroids, and so on, have beentreated lightly or ignored completely The use of this book in the first year of graduate studyshould help master the fundamentals It is hoped that this book will lead a student to consultthe many excellent books and review articles cited for various topics in order to understandthe subject in more detail Indeed, many of these topics are so vast, they cannot beexplained completely in this book

The organization is based on reaction types, and a relatively few principles suffice toexplain nearly all of them despite the large number of organic reactions Accordingly, thereactions-mechanisms section of this book (Part II) is divided into 10 chapters (10–19),each concerned with a different type of reaction In the first part of each chapter, theappropriate basic mechanisms are discussed along with considerations of reactivity andorientation, while the second part consists of numbered sections devoted to individualreactions, where the scope and the mechanism of each reaction are discussed Numberedsections are used for the reactions and are set in boldface Since the methods for thepreparation of individual classes of compounds (ketones, nitriles, etc.) are not treated all inone place, an updated and revised index has been provided (Appendix B) by use of whichthe synthesis of a given type of compound will be found It is important to note that thenumbers for each reaction in the 7th edition aredifferent from editions 1–5 in many cases,but are the same as found in the 6th edition For this reason, a correlation table is included atthe end of this Preface that directly correlates the sections found in the 5th edition with thenew ones in both the 6th and 7th editions

xiii

The structure of organic compounds is discussed in Chapters 1–5 (Part I) This sectionprovides a necessary background for understanding mechanisms and is also important in itsown right The discussion begins with chemical bonding (Chapt 1) and ends with a chapter

on stereochemistry (Chapt 4) Two chapters follow (Chapt 6–7) on reaction mechanisms ingeneral, one for ordinary reactions and the other for photochemical reactions Part 1concludes with two more chapters (Chapt 8 and 9) that give further background to the study

of mechanisms

The IUPAC names for many organic transformations are included, first introduced inthe 3rd edition Since then the rules have been broadened to cover additional cases; hencemore such names are given in this edition Furthermore, International Union of Pure andApplied Chemistry (IUPAC) has now published a system for designating reactionmechanisms, and some of the simpler designations are included

Appendix A is devoted to the literature of organic chemistry

In treating subjects as broad as structure, reactions, and mechanisms of organicchemistry, it is impossible to cover each topic in great depth, and this would not bedesirable even if possible This book is intended to point the reader to the primary literature

of the areas it covers To this end, there are >20,000 references to original papers.Secondary literature sources including reviews, books, and monographs have beenincluded as well Appendix A provides a brief introduction to using computer-basedsearch engines (e.g.,Reaxys1andSciFinder1)

Although basically designed as a reference text for a one-year course on the graduatelevel, this book can also be used in advanced undergraduate courses, but only aftercompletion of a one-year course in organic chemistry A one year course in both inorganicand physical chemistry would be most helpful It has been my experience that students whohave completed the first-year courses often have a hazy recollection of the material andgreatly profit from a representation of the material if it is easily accessible The material inthe first nine chapters, particularly Chapters 1, 2, 4, 6, and 8 may be helpful for reviewingsuch material when this book is used in connection with a course

This book is probably most valuable as a reasonably up-to-date reference work.Students preparing for qualifying examinations and practicing organic chemists willfind that Part II contains a survey of what is known about the mechanism and scope of

a large number of reactions, arranged in an orderly manner based on reaction type and onwhich bonds are broken and formed

For units of energy, IUPAC mandates joules, and many journals do use this unitexclusively However, organic chemists who publish in United States journals commonlyuse calories Virtually all energy values are presented in both calories and joules AlthoughIUPAC does not recommend angstrom units for bond distances, but rather picometers (pm),

a vast number of bond distances published in the literature are in angstrom units, and thisbook therefore uses angstrom units

I would like to acknowledge the contributions of those chemists cited and thanked byProfessor March in the first-four editions, and those I thanked in the 5th and 6th editions.This book would not be possible without their contributions For the 7th edition, I thankLou Allinger for pointing out the deficiencies in the hyperconjugation section, andgraciously helping me write the new section appearing in this new edition I thank WarrenHehre for his invaluable help in calculating and presenting the molecular orbital drawingsusing Spartan I also thank Adrian Shell (Elsevier) for facilitating the transfer of materialrelating to the programReaxys, discussed in Appendix A I thank the many people whohave contributed comments or have pointed out errors in the 6th edition that were

xiv PREFACE

invaluable to putting together the 7th edition I thank Warren Hehre and Sean Ohlinger ofWavefunction, Inc., Irvine, CA (www.wavefun.com ) for providing Spartan 10 Macint osh(v 1.0.1), allowing the incorporation of Spartan models for selected molecules andintermediates All structures and line drawings in this book were done using ChemDraw1Ultra 11.0.1 (350440), graciously provided by CambridgeSoft Corporation, Cambridge,

MA (www.cambridge soft.com )

Special thanks are due to the Interscience division of John Wiley & Sons and to JonathanRose Special thanks are also given to Kristen Parrish and Amanda Amanullah, at Wiley fortheir fine work as editors in turning the manuscript into the finished book as well asSanchari Sil of Thomson Digital I also thank Jeanette Stiefel for an excellent job of copyediting the manuscript

With gratitude, I acknowledge the work of Jerry March, upon whose work this newedition is built, and who is responsible for the concept of this book and for carrying itthrough four very successful editions I used Jerry’s book as a student and it is an honor tocontinue this tradition

I encourage those who read and use the 7th edition to contact me directly withcomments, errors, and with publications that might be appropriate for future editions

I hope that this new edition will due justice to the tradition that Professor March beganwith the first edition

Finally, I want to thank my wife Sarah for her patience and understanding during thepreparation of this manuscript I also thank my son Steven for his support Without theirsupport, this work would not have been possible

MICHAELB SMITHMay, 2012Correlation Table

5th edition! 7th edition Reactions

10-35! 16-6810-36! 10-2410-37! 10-2510-38! 10-2610-39! 16-6910-40! 10-2710-41! 10-2810-42! 10-2910-43! 10-3010-44! 10-3110-46! 10-3210-47! 10-3310-48! 16-7010-49! 10-3410-50! 10-3510-51! 10-3710-52! 10-38

PREFACE xv

11-10! 11-911-11! 11-1011-12! 11-1111-13! 11-1511-14! 11-1711-15! 11-1811-16 -deleted11-17 -deleted

11-18 -deleted11-19! 11-1911-20! 11-2011-21! 11-2111-22! 11-1211-23! 11-1311-24! 11-1411-25! 11-2211-26! 11-2311-27! 11-2411-28! 11-2511-29! 11-2611-30! 11-2711-31! 11-2811-32! 11-2911-33! 11-3011-34! 11-3111-35! 11-3211-36! 11-3311-37! 11-3411-38! 11-3511-39! 11-3611-40! 11-3711-41! 11-3811-42! 11-3911-43! 11-4011-44! 11-4112-1! 12-112-2! 12-212-3! 12-312-4! 12-412-5! 12-512-6 ! 12-612-7 ! 12-712-8 ! 12-812-9 ! 12-1012-10! 12-1112-11! 12-1212-12! 12-1312-13! 12-1412-14! 12-1612-15! 12-1812-16! 12-1912-17! 12-2012-18! 10-6912-19! 12-21

xvi PREFACE

14-37! 14-3014-38! 14-3114-39! 14-3215-1! 15-115-2! 15-215-3! 15-315-4! 15-415-5! 15-515-6 ! 15-615-7 ! 15-715-8 ! 15-815-9 ! 15-915-10! 15-1015-11! 15-1115-12! 15-1215-13! 15-1415-14! 15-1315-15! 15-1515-16! 15-1615-17! 15-1715-18! 15-1815-19! 15-2015-20! 15-2315-21! 15-2415-22! 15-2115-23! 15-2215-24! 15-2515-25! 15-2715-26! 15-2815-27! 15-3215-28! 15-3315-29! 15-3615-30! 15-3515-31! 15-3715-32! 15-3415-33! 15-3815-34! 15-1915-35! 15-2915-36! 15-3015-37! 15-3915-38! 15-4115-39! 15-4015-40! 15-4215-41! 15-4315-42! 15-4415-43! 15-4515-44! 15-46PREFACE xvii

17-4! 17-517-5! 17-617-6! 17-717-7! 17-817-8! 17-917-9! 17-1017-10! 17-1117-11! 17-1217-12! 17-1317-13! 17-1417-14! 17-1517-15! 17-1617-16! 17-1717-17! 17-1817-18! 17-1917-19! 17-317-20! 17-2017-21! 17-2117-22! 17-2217-23! 17-2317-24! 17-2417-25! 17-2517-26deletedcombinedwith 17-2517-27! 17-2617-28! 17-2717-29! 17-2817-30! 17-2917-31deletedcombinedwith 17-3017-32! 17-3017-33! 17-3117-34! 17-3217-35! 17-3317-36! 17-3417-37! 17-3517-38! 17-3617-39! 17-3717-40! 17-3818-1 ! 18-118-2 ! 18-218-3 ! 18-318-4 ! 18-418-5 ! 18-5

in 19-1419-18! 19-1919-19! 19-2019-20! 19-2119-21! 19-2219-22! 19-2519-23! 19-2719-24! 19-2819-25! 19-3019-26! 19-2619-27! 19-2919-28! 19-3119-29! 19-2419-30! 19-3219-31! 19-33

19-32! 19-3419-33! 19-6119-34! 19-3719-35! 19-6419-36! 19-6219-37! 19-6319-38! 19-3819-39! 19-6519-40deletedincorporatedinto 10-8519-41! 19-4519-42! 19-4619-43! 19-4719-44! 19-4819-45! 19-5019-46! 19-5119-47! 19-7119-48! 19-6819-49! 19-7219-50! 19-6019-51! 19-4919-52! 19-7319-53! 19-7419-54! 19-7519-55! 19-7619-56! 19-7719-57! 19-7819-58! 19-7919-59! 19-8019-60! 19-8119-61! 19-8219-62! 19-8319-63! 19-84

PREFACE xix

BINAP (2R,3S)-2,20-bis-(diphenylphosphino)-1,10-binapthyl

CNDO Complete Neglect of Differential Overlap

xxi

cot 1,3,5-Cyclooctatrienyl (ligand)

O

HMO H€uckel molecular orbital

1H NMR Proton nuclear magnetic resonance spectroscopy

HOMO Highest occupied molecular orbital

IUPAC International Union of Pure and Applied Chemistry

LCAO Linear combination of atomic orbitals

LICA

(LIPCA)

LithiumN-isopropyl-N-cyclohexylamide

LTMP Lithium 2,2,6,6-tetramethylpiperidide

NOESY Nuclear overhauser effect spectroscopy

PIFA Phenyliodine (III)-bis-(trifluoroacetate)

Red-Al [(MeOCH2CH2O)2AlH2]Na

(Sia)2BH Disiamylboranesec-Isoamyl

TEMPO 2,2,6,6-Tetramethylpiperidinyloxy free radical

TOSMIC Toluenesulfonylmethyl isocyanide

xxiv COMMON ABBREVIATIONS

BIOGRAPHICAL STATEMENT

Professor Michael B Smith was born in Detroit, Michigan in 1946 In 1957, he and hisfamily moved to Madison Heights, Virginia After graduation from Amherst County highschool, he entered Ferrum Jr College and graduated with an A.A Professor Smithtransferred to Virginia Polytechnic Institute (Virginia Tech), where he did undergraduateresearch with Professor Harold Bell, and graduated with a B.S in chemistry in 1969 Afterworking as an analytical chemist at the Newport News Shipbuilding and Dry Dock Co.(Tenneco) in Newport News, Virginia for three years, he began graduate studies at PurdueUniversity under the mentorship of Professor Joseph Wolinsky Professor Smith graduatedwith a Ph.D in Organic chemistry in 1977 He spent one year as a faculty research associate

at the Arizona State University in the Cancer Research Institute, directed by ProfessorGeorge R Pettit, and a second year doing postdoctoral work at the Massachusetts Institute

of Technology under the mentorship of Professor Sidney Hecht In 1979 he began hisindependent academic career, where he now holds the rank of full professor

Professor Smith is the author of approximately 90 independent research articles, and 20published books The books include the 5th and 6th edition of March’s Advanced OrganicChemistry (Wiley), volumes 6–12 of the Compendium of Organic Synthetic Methods(Wiley), Organic Chemistry a Two Semester Course (HarperCollins) into its 2nd edition,and Organic Synthesis (Elsevier) in its 3rd edition A new undergraduate organic chemistrybook, Organic Chemistry: An Acid-Base Approach, was published in 2011 by the CRC Press.Professor Smith’s current research involves the synthesis and structural verification oflipids obtained from the dental pathogen Porphyromonas gingivalis, which show inflam-matory activity, induce bone degeneration and are involved in triggering multiple sclerosis

A main area of research is the synthesis of fluorescent dye-heterocyclic conjugates thattarget hypoxic cancerous tumors, allowing non-invasive fluorescence imaging in the near

IR The synthesis of anti-cancer alkaloids is also ongoing

xxv

PART I

INTRODUCTION

This book contains 19 chapters Chapters 1–9 may be thought of as an introduction toPart II The first-five chapters deal with the structure of organic compounds These chaptersdiscuss the kinds of bonding important in organic chemistry, the fundamental principles ofconformation and stereochemistry of organic molecules, and reactive intermediates inorganic chemistry Chapters 6–9 are concerned with general principles of mechanism inorganic chemistry, including acids and bases, photochemistry, sonochemistry and micro-wave irradiation, and finally the relationship between structure and reactivity

Chapters 10–19, which make up Part II, are directly concerned with the nature and thescope of organic reactions and their mechanisms

March’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Seventh Edition.

Michael B Smith.

Ó 2013 John Wiley & Sons, Inc Published 2013 by John Wiley & Sons, Inc.

1

CHAPTER 1

Localized Chemical Bonding

Localized chemical bonding may be defined as bonding in which the electrons are shared

by two and only two nuclei Such bonding is the essential feature associated with thestructure of organic molecules.1 Chapter 2 will discuss delocalized bonding, in whichelectrons are shared by more than two nuclei

1.A COVALENT BONDING2

Wave mechanics is based on the fundamental principle that electrons behave as waves(e.g., they can be diffracted) Consequently, a wave equation can be written for electrons, inthe same sense that light waves, sound waves, and so on, can be described by waveequations The equation that serves as a mathematical model for electrons is known as theSchr€odinger equation, which for a one-electron system is

The Schr€odinger equation is a differential equation, so solutions to it are themselvesequations, but the solutions are not differential equations They are just simple equationsfor which graphs can be drawn Such graphs are essentially three-dimensional (3D)

March’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Seventh Edition.

3

pictures that show the electron density, and these pictures are called orbitals or electronclouds Most students are familiar with the shapes of the s and p atomic orbitals (Fig 1.1).Note that each p orbital has a node: A region in space where the probability of finding theelectron is extremely small.3Also note that in Fig 1.1 some lobes of the orbitals are labeled

þ and others  These signs do not refer to positive or negative charges, since both lobes of

an electron cloud must be negatively charged They are the signs of the wave functionC.When a node separates two parts of an orbital, a point of zero electron density,C alwayshas opposite signs on the two sides of the node According to the Pauli exclusion principle,

no more than two electrons can be present in any orbital, and they must have opposite spins.Unfortunately, the Schr€odinger equation can be solved exactly only for one-electronsystems (e.g., the hydrogen atom) If it could be solved exactly for molecules containingtwo or more electrons,4 a precise picture of the shape of the orbitals available to eachelectron (especially for the important ground state) would become available, as well as theenergy for each orbital Since exact solutions are not available, drastic approximationsmust be made There are two chief general methods of approximation: the molecularorbital (MO) method and the valence bond method

In the MO method, bonding is considered to arise from the overlap of atomic orbitals.When any number of atomic orbitals overlap, they combine to form an equal number of neworbitals, called molecular orbitals Molecular orbitals differ from atomic orbitals in that anelectron cloud effectively surrounds the nuclei of two or more atoms, rather than just oneatom In other words, the electrons are shared by two atoms rather than being localized on oneatom In localized bonding for a single covalent bond, the number of atomic orbitals thatoverlap is two (each containing one electron), so that two molecular orbitals are generated.One of these, called a bonding orbital, has a lower energy than the original atomic orbitals(otherwise a bond would not form), and the other, called an antibonding orbital, has a higher

•

z y

– –

FIG 1.1 (a) The 1s orbital (b) The three 2p orbitals

3 When wave mechanical calculations are made according to the Schr €odinger equation, the probability of finding the electron in a node is zero, but this treatment ignores relativistic considerations When such considerations are applied, Dirac has shown that nodes do have a very small electron density: Powell, R.E J Chem Educ 1968, 45,

558 See also, Ellison, F.O.; Hollingsworth, C.A J Chem Educ 1976, 53, 767; McKelvey, D.R J Chem Educ.

1983, 60, 112; Nelson, P.G J Chem Educ 1990, 67, 643 For a general review of relativistic effects on chemical structures, see Pyykk €o, P Chem Rev 1988, 88, 563.

4 See Roothaan, C.C.J.; Weiss, A.W Rev Mod Phys 1960, 32, 194; Kolos, W.; Roothaan, C.C.J Rev Mod Phys.

1960, 32, 219 For a review, see Clark, R.G.; Stewart, E.T Q Rev Chem Soc 1970, 24, 95.

energy Orbitals of lower energy fill first Since the two original atomic orbitals each held oneelectron, both of these electrons will reside in the new molecular bonding orbital, which islower in energy Remember that any orbital can hold two electrons The higher energyantibonding orbital remains empty in the ground state.

The strength of a bond is determined by the amount of electron density that residesbetween the two nuclei The greater the overlap of the orbitals, the stronger the bond, buttotal overlap is prevented by repulsion between the nuclei Figure 1.2 shows the bondingand antibonding orbitals that arise by the overlap of two 1s electrons Note that since theantibonding orbital has a node between the nuclei, there is practically no electron density inthat area, so that this orbital cannot be expected to bond very well When the centers ofelectron density are on the axis common to the two nuclei, the molecular orbitals formed bythe overlap of two atomic orbitals are calleds (sigma) orbitals, and the bonds are called sbonds The corresponding antibonding orbitals are designateds Sigma orbitals may be

formed by the overlap of any of the atomic orbital (s, p, d, or f ) whether the same ordifferent, not only by the overlap of two s orbitals However, the two lobes that overlapmust have the same sign: A positive s orbital can form a bond only by overlapping withanother positive s orbital or with a positive lobe of a p, d, or f orbital Anys molecularorbital may be represented as approximately ellipsoidal in shape

Orbitals are frequently designated by their symmetry properties The s orbital ofhydrogen is often writtencg The g stands for gerade A gerade orbital is one in which thesign on the orbital does not change when it is inverted through its center of symmetry The

The functionscAandcBare the functions for the atomic orbitals of atoms A and B, respectively,and cAand cBrepresent weighting factors Subtraction is also a linear combination:

c ¼ cAcA cBcB (1-2)This gives rise to the antibonding molecular orbital

In the valence bond method, a wave equation is written for each of various possibleelectronic structures that a molecule may have (each of these is called a canonical form),and the totalc is obtained by summation of as many of these as seem plausible, each withits weighting factor:

c ¼ c1c1þ c2c2þ    (1-3)This resembles Eq (1-1), but here each c represents a wave equation for an imaginarycanonical form and each c is the amount contributed to the total picture by that form Forexample, a wave function can be written for each of the following canonical forms of thehydrogen molecule5:

HH H: Hþ þH H:

Values for c in each method are obtained by solving the equation for various values ofeach c, and choosing the solution of lowest energy In practice, both methods give similarsolutions for molecules that contain only localized electrons, and these are in agreementwith the Lewis structures long familiar to the organic chemist Delocalized systems areconsidered in Chapter 2 Note that orbital functions can indeed be reconstructed frommeasured data using several different approaches Often, however, the results are still lessaccurate than those achieved with purely theoretical methods.6

1.B MULTIPLE VALENCE

A univalent atom has only one orbital available for bonding But atoms with a valence of 2 ormore must form bonds by using at least two orbitals An oxygen atom has two half-filledorbitals, giving it a valence of 2 It forms single bonds by the overlap of these with the orbitals

of two other atoms According to the principle of maximum overlap, the other two nucleishould form an angle of 90 with the oxygen nucleus, since the two available orbitals on

oxygen are p orbitals, which are perpendicular If this is correct, nitrogen, which has threemutually perpendicular p orbitals, would have bond angles of 90when it forms three single

bonds However, these are not the observed bond angles The bond angles in water are,7

104270, and in ammonia, 106460 For alcohols and ethers, the angles are even larger (see

Sec 1.K) A discussion of this will be deferred to Section 1.K, but it is important to note thatcovalent compounds do have definite bond angles Although the atoms are continuouslyvibrating, the mean position is the same for each molecule of a given compound

5 In this book, a pair of electrons in a bond is represented by two dots.

6 Schwarz, W.H.E Angew Chem Int Ed 2006, 45, 1508 For the ball-in-box model, see Pierrefixe, S.C.A.H.; Guerra, C.F.; Bickelhaupt, F.M Chem Eur J 2008, 14, 819; Pierrefixe, S.C.A.H.; Bickelhaupt, F.M J Phys Chem A 2008, 112, 12816.

7

Bent, H.A Chem Rev 1961, 61, 275, 277.

1.C HYBRIDIZATION

Consider the case of mercury Its electronic structure is

½Xe core4f145d106s2Although it has no half-filled orbitals, it has a valence of 2 and forms two covalent bonds.This result can be explained by imagining that one of the 6s electrons is promoted to avacant 6p orbital to give the excited configuration

½Xe core4f145d106s16p1

In this state, the atom has two half-filled orbitals, but they are not equivalent If bondingwere to occur by the overlap of these orbitals with the orbitals of external atoms, the twobonds would not be equivalent The bond formed from the 6p orbital would be more stablethan the one formed from the 6s orbital, since a larger amount of overlap is possible withthe former A more stable situation is achieved when, in the course of bond formation, the6s and 6p orbitals combine to form two new orbitals that are equivalent; these are shown inFig 1.3

The new molecular orbitals are a mixture of the two original orbitals, so they are calledhybrid orbitals Each orbital is a merger of an s and p orbital and is called an sp orbital The

sp orbitals, each of which consists of a large lobe and a very small one, arise only in thebonding process and do not represent a possible structure for the free atom A mercuryatom forms its two bonds by overlapping each of the large lobes shown in Fig 1.3 with anorbital from an external atom The orbital of this external atom may be any of the atomicorbitals previously considered (s, p, d, or f), or it may be another hybrid orbital Note thatonly lobes of the same sign can overlap In any of these cases, the molecular orbital thatarises is called as orbital since it fits the previous definition of a s orbital

In general, equivalent orbitals lie as far away from each other as possible because ofmutual repulsion, so two sp orbitals form an angle of 180 In other words, an atom that

forms only twos bonds uses two sp orbitals so HgCl2, for example, should be a linearmolecule, and it is This kind of hybridization is called digonal hybridization An sp hybridorbital forms a stronger covalent bond than either an s or a p orbital because it extends out

in space in the direction of the other atom’s orbital farther than the s or the p and permits

z

y x

+ –

+ –

FIG 1.3 The two sp orbitals formed by mercury

greater overlap Compare HgCl2with water (OH2) It is known that the shape of HgCl2islinear, but water is angular This fact suggests that the hybrid orbitals utilized by oxygen inwater is different from those used by mercury in HgCl2.

Many other kinds of hybridization are possible Consider boron, which has the electronicconfiguration 1s22s22p1yet has a valence of 3 To begin, boron has only three valenceelectrons available to form bonds, hence the valence of three Any hybridization model musttake this into account As before, imagine promotion of an electron and hybridization:

1s22s22p1!promotion 1s22s12p1x2p1y!hybridization 1s2ð2s2Þ3

In this case, there are three equivalent hybrid orbitals, each called sp2(trigonal tion) This method of designating hybrid orbitals is perhaps unfortunate since nonhybridorbitals are designated by single letters, but keep in mind that each of the three orbitals iscalled sp2 The key is to understand that the atom forms twos bonds for sp hybridization andthrees bonds for sp2hybridization The sp2hybrid orbitals just noted are shown in Fig 1.4.The three axes are all in one plane and point to the corners of an equilateral triangle Thisaccords with the known structure of BF3, a planar molecule with angles of 120.

hybridiza-Another type of hybrid orbital is possible, formed by atoms that can form fours bonds.Carbon is an important atom that can form four single bonds (four s bonds) Imaginepromotion of an electron and hybridization that leads to

1s22s22p1x2p1y!promotion 1s22s12p1x2p1y2p1z!hybridization 1s2ð2sp3Þ4

There are four equivalent molecular orbitals, each called sp3, and electron repulsion leads

to a shape in which the orbitals point to the corners of a regular tetrahedron (Fig 1.4) Atypical molecule is methane (CH4) and assuming that carbon forms four bonds with sp3hybrid orbitals, the bond angles of methane would thus be expected to be 109280, which is

the angle for a regular tetrahedron In reality, electrons are not “promoted” in atomicorbitals, but atomic orbitals are different from molecular orbitals (e.g., those found inmethane) The model of promoting an electron is a mathematical device to describemolecular orbitals using the atomic orbitals

The hybrid orbitals discussed in this section stem from only one possible approximatesolution of the Schr€odinger equation The s and the three p atomic orbitals used to form sp3

orbitals, for example, can be combined in other equally valid ways As will be seen inSection 1.E, the four CH bonds of methane do not always behave as if they are

z

y x

z

y x

120°

120°

120°

FIG 1.4 The three sp2

and the four sp3orbitals

equivalent Bickelhaupt6has proposed an alternative approach to the bonding in carbonsuggesting that the maximum coordination number of carbn cannot exceed four because it

is too small to allow more than four substituents approach and form the appropriate bonds

high electron density

high electron density

high electron density

C R

R

C R

orbitals whose axes are parallel are calledp orbitals if they are bonding and pif they are

antibonding

In this picture of ethylene, there are two bonds connecting the adjacent carbon atoms,but the two orbitals that make up the double bond are not equivalent.8This means that thetwo bonds are different one from the other Thes orbital is ellipsoidal and symmetricalabout the CC axis, and this is the familiar s bond The p orbital is in the shape of twoellipsoids, one above the plane and one below, and forms the second bond, ap bond Theplane itself represents a node for the p orbital In order for the p orbitals to maintainmaximum overlap, they must be parallel Since both as bond and the p bond connect thetwo carbon atoms, free rotation is not possible about the double bond In other words, thetwo p orbitals would have to reduce their overlap to allow one HCH plane to rotatewith respect to the other (i.e., thep bond would have to disappear) With two sp2

hybridcarbon atoms in ethylene, the six atoms associated with the double bond (H2CCH2) are in

a plane with angles that should be120 Double bonds are shorter than the corresponding

single bonds because maximum stability is obtained when the p orbitals overlap as much aspossible (see Sec 1.J) Double bonds between carbon and oxygen (CO) or nitrogen(CN) similarly consist of one s and one p orbital.

When carbon is connected to another carbon atom by a triple bond, as in acetylene(HCCH), each carbon is connected to only two other atoms by a s bond, and hence uses

sp hybridization This fact requires that the four atoms of acetylene (2H and 2C) are in astraight line (Fig 1.6).9Each carbon has two p orbitals remaining, with one electron ineach These orbitals are perpendicular to each other and also to the CC axis Theyoverlap in the manner shown in Fig 1.7 to form two p orbitals A triple bond is thuscomposed of ones and two p orbitals Triple bonds between carbon and nitrogen can berepresented in a similar manner, CN.

For most organic molecules, double and triple bonds typically involve the first-rowelements carbon, nitrogen, and oxygen.10 Second-row elements tend to form weaker pbonds than do the first-row elements,11so multiple bonds are less common and compoundscontaining them are generally less stable.12Compounds with CS bonds are known, forexample, and CS compounds are generally much less stable than the correspondingCO compounds (however, see pp–dp bonding in Sec 2.H) Stable compounds with

C

FIG 1.6 The s orbitals of acetylene

8 For an alternative representation, see Pauling, L Theoretical Organic Chemistry, The Kekule Symposium, Butterworth, London, 1959, pp 2–5; Palke, W.E J Am Chem Soc 1986, 108, 6543.

9

See Simonetta, M.; Gavezzotti, A., in Patai, S The Chemistry of the Carbon–Carbon Triple Bond, Wiley, NY,

1978, pp 1–56; Dale, J., in Viehe, H.G Acetylenes, Marcel Dekker, NY, 1969, pp 3–96.

10 For a review of metal–metal multiple bonds, see Cotton, F.A J Chem Educ 1983, 60, 713.

11 For discussions, see Schmidt, M.W.; Truong, P.N.; Gordon, M.S J Am Chem Soc 1987, 109, 5217; Schleyer,

P von R.; Kost, D J Am Chem Soc 1988, 110, 2105.

12 For double bonds between carbon and elements other than C, N, S, or O, see Jutzi, P Angew Chem Int Ed.

1975, 14, 232; Raabe, G.; Michl, J Chem Rev 1985, 85, 419 (Si only); Wiberg, N J Organomet Chem 1984,

273, 141 (Si only); Gordon, M.S Mol Struct Energ 1986, 1, 101 For reviews of C P and CP bonds, seeRegitz, M Chem Rev 1990, 90, 191; Appel, R.; Knoll, F Adv Inorg Chem 1989, 33, 259; Markovski, L.N.; Romanenko, V.D Tetrahedron 1989, 45, 6019.

SiC and SiSi bonds are rare, but examples have been reported,13including a pair of cisand trans SiSi isomers.14

There is at least one report of a so-called two-electron, four-center CC bond for thedimer of tetracyanoethylene.15 While such multi-center bonding is not formally anexample of the multiple bonding described in this section, it constitutes a differenttype of bonding when compared to the simple CC bonds described earlier

1.E PHOTOELECTRON SPECTROSCOPY

Based on the hybridization model, methane is expected to have four equivalents bonds.Indeed, the four bonds of methane are equivalent according to most physical and chemicalmethods of detection The nuclear magnetic resonance (NMR) and the infrared (IR)spectrum of methane show no peaks that can be attributed to different kinds of CHbonds However, there is one physical technique showing that the eight valence electrons ofmethane can be differentiated In this technique, called photoelectron spectroscopy(PES),16a molecule or free atom is bombarded with vacuum ultraviolet (UV) radiation,causing an electron to be ejected The energy of the ejected electron can be measured, andthe difference between the energy of the radiation used and that of the ejected electron is

FIG 1.7 (a) The electron density map of acetylene Note the concentration of electron density along

a line between the nuclei of each atom, consistent with overlap ofs orbitals in a triple bond.(b) Electron potential map of acetylene showing the concentration of electron density between thecarbon atoms, consistent with two orthogonalp bonds

13 For SiC bonds, see Fink, M.J.; DeYoung, D.J.; West, R.; Michl, J J Am Chem Soc 1983, 105, 1070; Fink,M.J.; Michalczyk, M.J.; Haller, K.J.; West, R.; Michl, J Organometallics 1984, 3, 793; West, R Pure Appl Chem.

1984, 56, 163; Masamune, S.; Eriyama, Y.; Kawase, T Angew Chem Int Ed 1987, 26, 584; Shepherd, B.D.; Campana, C.F.; West, R Heteroat Chem 1990, 1, 1.

14

Michalczyk, M.J.; West, R.; Michl, J J Am Chem Soc 1984, 106, 821, Organometallics 1985, 4, 826 15

Miller, J.S.; Novoa, J.J Acc Chem Res 2007, 40, 189.

16 See Ballard, R.E Photoelectron Spectroscopy and Molecular Orbital Theory, Wiley, NY, 1978; Rabalais, J.W Principles of Ultraviolet Photoelectron Spectroscopy, Wiley, NY, 1977; Baker, A.D.; Betteridge, D Photo- electron Spectroscopy, Pergamon, Elmsford, NY, 1972; Turner, D.W.; Baker, A.D.; Baker, C.; Brundle, C.R High Resolution Molecular Photoelectron Spectroscopy, Wiley, NY, 1970 For reviews, see Westwood, N.P.C Chem Soc Rev 1989, 18, 317; Baker, C.; Brundle, C.R.; Thompson, M Chem Soc Rev 1972, 1, 355; Bock, H.; Ramsey, B.G Angew Chem Int Ed 1973, 12, 734; Turner, D.W Adv Phys Org Chem 1966, 4, 31 For the IUPAC descriptive classification of various electron spectroscopy techniques, see Porter, H.Q.; Turner, D.W Pure Appl Chem 1987, 59, 1343.

the ionization potential of that electron A molecule that contains several electrons ofdiffering energies can lose any one of them as long as its ionization potential is less than theenergy of the radiation used A single molecule loses only one electron; the loss of twoelectrons by any individual molecule almost never occurs Since electrons reside inorbitals, a photoelectron spectrum consists of a series of bands, each corresponding to

an orbital of a different energy The spectrum gives a direct experimental picture of allorbitals that are present, and they are ejected in ascending order of their energies, providedthat radiation of sufficiently high energy is used.17 Broad bands usually correspond tostrongly bonding electrons and narrow bands to weakly bonding or nonbonding electrons.Using PES, it is possible to probe the validity of the hybridization model for bonding.Dinitrogen (N2) is a typical diatomic molecule and is shown in Fig 1.8.18The N2moleculehas the electronic structure shown in Fig 1.9: The two 2s orbitals of the nitrogen atomscombine to give the two orbitals marked 1 (bonding) and 2 (antibonding), while the six 2porbitals combine to give six orbitals, three of which (marked 3, 4, and 5) are bonding Thethree antibonding orbitals (not indicated in Fig 1.9) are unoccupied Electrons ejectedfrom orbital 1 are not found in Fig 1.8 because the ionization potential of these electrons isgreater than the energy of the light used (they can be seen when higher energy light is used).The broad band in Fig 1.8 corresponds to the four electrons in the degenerate orbitals 3and 4 The individual peaks within this band are caused by different vibrational levels (seeChap 7) The triple bond of N2is therefore composed of these two orbitals and orbital 1.The bands corresponding to orbitals 2 and 5 are narrow; hence these orbitals contributelittle to the bonding and may be regarded as the two unshared pairs of €N  €N Note that thisresult is contrary to that expected from a naive consideration of orbital overlaps, where it

FIG 1.8 Photoelectron spectrum of N2.18 [Reprinted with permission from Brundle, C.R.; Robin,M.B in Nachod, F.C.; Zuckerman, J.J Determination of Organic Structures by Physical Methods,Vol 1, Academic Press, NY,1971, p 18 Copyright # 1971, with permission from Elsevier Science.With permission of C Richard Brundle, 2012.]

17 The correlation is not perfect, but the limitations do not seriously detract from the usefulness of the method The technique is not limited to vacuum UV radiation Higher energy radiation can also be used.

18 From Brundle, C.R.; Robin, M.B., in Nachod, F.C.; Zuckerman, J.J Determination of Organic Structures by Physical Methods, Vol 3, Academic Press, NY, 1971, p 18.