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ADVANCED ORGANIC CHEMISTRY
Reactions, Mechanisms, and Structure
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BICENTENNIAL
The Wiley Bicentennial-Knowledge for generations
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ach generation has its unique needs and aspirations. When Charles Wiley first opened his small printing shop in lower Manhattan in 1807, it was a generation of boundless potential searching for an identity. And we were there, helping to define a new American literary tradition. Over half a century later, in the midst of the Second Industrial Revolution, it was a generation focused on building the future. Once again, we were there, supplying the critical scientific, technical, and engineering knowledge that helped frame the world. Throughout the 20th Century, and into the new millennium, nations began to reach out beyond their own borders and a new international community was bom. Wiley was there, expanding its operations around the world to enable a global exchange of ideas, opinions, and know-how.
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William J. Pesce Peter Booth Wiley
President and Chief Executive Officer Chairman of the Board
MARCH’S ADVANCED ORGANIC CHEMISTRY
REACTIONS, MECHANISMS,
AND STRUCTURE
SIXTH EDITION
Michael B. Smith
Professor of Chemistry
Jerry March
Professor of Chemistry
3ICENTENNIAL
BICENTENNIAL
WILEY-INTERSCIENCE A JOHN WILEY & SONS, INC., PUBLICATION
Copyright © 2007 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 is available .
Smith, Michael B., March, Jerry
March’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Sixth Edition
ISBN 13: 978-0-471-72091-1 ISBN 10: 0-471-72091-7
Printed in the United States of America
10 987654321
CONTENTS
PREFACE y
BIOGRAPHICAL NOTE xv
ABBREVIATIONS xvii
PART 1 1
1. Localized Chemical Bonding 3
2. Delocalized Chemical Bonding 32
3. Bonding Weaker than Covalent 106
4. Stereochemistry 136
5. Carbocations, Carbanions, Free Radicals,
Carbenes, and Nitrenes 234
6. Mechanisms and Methods of Determining Them 296
7. Irradiation Processes in Organic Chemistry 328
8. Acids and Bases 356
9. Effects of Structure and Medium on Reactivity 395
PART 2 417
10. Aliphatic Substitution: Nucleophilic and Organometallic 425
11. Aromatic Substitution, Electrophilic 657
12. Aliphatic, Alkenyl, and Alkynyl Substitution,
Electrophilic and Organometallic 752
13. Aromatic Substitution, Nucleophilic and Organometallic 853
14. Substitution Reactions: Free Radicals 934
15. Addition to Carbon-Carbon Multiple Bonds 999
xiii
XIV
CONTENTS
16. Addition to Carbon-Hetero Multiple Bonds 1251
17. Eliminations 1477
18. Rearrangements 1559
19. Oxidations and Reductions 1703
Appendix A The Literature of Organic Chemistry 1870
Appendix B Classification of Reactions by Type of
Compounds Synthesized 1911
Indexes
Author Index 1937
Subject Index 2190
PREFACE
Organic chemistry is a vibrant and growing scientific discipline that touches a vast number of scientific areas. This sixth edition of “March’s Advanced Organic Chemistry” has been thoroughly updated to reflect new areas of Organic chemistry, as well as new advances in well-known areas of Organic chemistry. Every topic retained from the fifth edition has been brought up to date. Changes include the addition of a few new sections, significant revision to sections that have seen explo¬ sive growth in that area of research, moving sections around within the book to bet¬ ter reflect logical and reasonable chemical classifications, and a significant rewrite of much of the book. More than 7000 new references have been added. As with the fifth edition, when older references were deleted and in cases where a series of papers by the same principal author were cited, all but the most recent were deleted. The older citations should be found within the more recent one or ones. The funda¬ mental structure of the sixth edition is essentially the same as that of all previous ones, although acyl substitution reactions have been moved from chapter 10 to chapter 16, and many oxidation or reduction reactions have been consolidated into chapter 19.
Like the first five editions, the sixth is intended to be a textbook for a course in advanced organic chemistry taken by students who have had the standard under¬ graduate organic and physical chemistry courses.
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 has completed a course based on this book should be able to approach the literature directly, with a sound knowledge of modern basic organic chemistry. Major special areas of organic chemistry: terpenes, carbohydrates, proteins, many organometallic reagents, combinatorial chemistry, polymerization and electroche¬ mical reactions, steroids, etc. have been treated lightly or ignored completely. I share the late Professor March’s opinion that these topics are best approached after the first year of graduate study, when the fundamentals have been mastered, either in advanced courses, or directly, by consulting the many excellent books and review articles available on these subjects. In addition, many of these topics are so vast, they are beyond the scope of this book.
The organization is based on reaction types, so the student can be shown that despite the large number of organic reactions, a relatively few principles suffice to explain nearly all of them. Accordingly, the reactions-mechanisms section of this book (Part 2) is divided into 10 chapters (10-19), each concerned with a dif¬ ferent type of reaction. In the first part of each chapter the appropriate basic
v
VI
PREFACE
mechanisms are discussed along with considerations of reactivity and orientation, while the second part consists of numbered sections devoted to individual reactions, where the scope and the mechanism of each reaction are discussed. Numbered sec¬ tions are used for the reactions. Since the methods for the preparation of individual classes of compounds (e.g., ketones, nitriles, etc.) are not treated all in one place, an index has been provided (Appendix B) by use of which all methods for the prepara¬ tion of a given type of compound will be found. For each reaction, a list of Organic Syntheses references is given where they have been reported. Thus for many reac¬ tions the student can consult actual examples in Organic Syntheses. It is important to note that the numbers for each reaction differ from one edition to the other, and many of the sections in the fifth edition do not correlate with the fourth. A correla¬ tion table is included at the end of this Preface that directly correlates the sections found in the 5th edition with the new ones in the 6th edition.
The structure of organic compounds is discussed in the first five chapters of Part 1 . This section provides a necessary background for understanding mechanisms and is also important in its own right. The discussion begins with chemical bonding and ends with a chapter on stereochemistry. There follow two chapters on reaction mechanisms in general, one for ordinary reactions and the other for photochemical reactions. Part 1 concludes with two more chapters that give further background to the study of mechanisms.
In addition to reactions, mechanisms, and structure, the student should have some familiarity with the literature of organic chemistry. A chapter devoted to this topic has been placed in Appendix A, though many teachers may wish to cover this material at the beginning of the course.
The IUPAC names for organic transformations are included, first introduced in the third edition. Since then the rules have been broadened to cover additional cases; hence more such names are given in this edition. Furthermore, IUPAC has now published a new system for designating reaction mechanisms (see p. 420), and some of the simpler designations are included.
In treating a subject as broad as the basic structures, reactions, and mechanisms of organic chemistry, it is obviously not possible to cover each topic in great depth. Nor would this be desirable even if possible. Nevertheless, students will often wish to pursue individual topics further. An effort has therefore been made to guide the reader to pertinent review articles and books published since about 1965. In this respect, this book is intended to be a guide to the secondary literature (since about 1965) of the areas it covers. Furthermore, in a graduate course, students should be encouraged to consult primary sources. To this end, more than 20,000 references to original papers have been included.
Although basically designed for a one-year course on the graduate level, this book can also be used in advanced undergraduate courses, but a one-year course in organic chemistry prior to this is essential, and a one year course in physical chemistry is strongly recommended. It can also be adapted, by the omission of a large part of its contents, to a one- semester course. Indeed, even for a one-year course, more is included than can be conveniently covered. Many individual sec¬ tions can be easily omitted without disturbing continuity.
PREFACE Vii
The reader will observe that this text contains much material that is included in first-year organic and physical chemistry courses, though in most cases it goes more deeply into each subject and, of course, provides references, which first-year texts do not. It has been my experience that students who have completed the first-year courses often have a hazy recollection of the material and greatly profit from a representation of the material if it is organized in a different way. It is hoped that the organization of the material on reactions and mechanisms will greatly aid the memory and the understanding. In any given course the teacher may want to omit some chapters because students already have an adequate knowledge of the material, or because there are other graduate courses that cover the areas more thoroughly. Chapters 1, 4, and 7 especially may fall into one of these categories.
This book is probably most valuable as a reasonably up-to-date reference work. Students preparing for qualifying examinations and practicing organic chemists will find that Part 2 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 on which bonds are broken and formed. Also valuable for reference pur¬ poses are the previously mentioned lists of reactions classified by type of compound prepared (Appendix B) and of all of the Organic Syntheses references to each reaction.
Anyone who writes a book such as this is faced with the question of which units to use, in cases where international rules mandate one system, but published papers use another. Two instances are the units used for energies and for bond distances. For energies, IUPAC mandates joules, and many journals do use this unit exclu¬ sively. However, organic chemists who publish in United States journals overwhel¬ mingly use calories and this situation shows no signs of changing in the near future. Since previous editions of this book have been used extensively both in this country and abroad, I have now adopted the practice of giving virtually all energy values in both calories and joules. The question of units for bond distances is easier to answer. Although IUPAC does not recommend Angstrom units, nearly all bond dis¬ tances published in the literature anywhere in the world, whether in organic or in crystallographic journals, are in these units, though a few papers do use picometers. Therefore, I continue to use only Angstrom units.
I would like to acknowledge the contributions of those chemists cited and thanked by Professor March in the first four editions. I especially thank George Majetich, Warren Hehre, and Amos B. Smith III for generous contributions to spe¬ cialized sections in the book as well as reviewing those sections. I also thank the many people who have contributed comments or have pointed out errors in the 5th edition that were invaluable to putting together the 6th edition. I thank Cambridge- Soft Inc. for providing ChemOffice, with ChemDraw , which was used to prepare all reactions and several structures in this book. I thank Dr. Warren Hehre and Wave- function, Inc. for providing MacSpartan, allowing the incorporation of Spartan 3D models for selected molecules and intermediates.
Special thanks are due to the Interscience division of John Wiley & Sons and to Dr. Darla Henderson without whose support the book would not have been com¬ pleted. Special thanks are also given to Shirley Thomas and Rebekah Amos at
PREFACE
viii
Wiley for their fine work as editors in turning the manuscript into the finished book. I also thank Ms. Jeannette Stiefel, for an excellent job of copy editing the manu¬ script. I gratefully acknowledge the work of the late Professor Jerry March, upon whose work this new edition is built, and who is responsible for the concept of this book and for carrying it through four very successful editions.
I encourage those who read and use the sixth edition to contact me directly with comments, errors, and with publications that might be appropriate for future edi¬ tions. I hope that this new edition will carry on the tradition that Professor March began with the first edition.
My Email address is
michael . smith @ uconn . edu
and my homepage is
http://orgchem.chem.uconn.edu/home/mbs-home.html
Finally, I want to thank my wife Sarah for her patience and understanding during the preparation of this manuscript. I also thank my son Steven for his support. With¬ out their support, this work would not have been possible.
Michael B. Smith June, 2006
5th edition — > 6th edition
10-1 — >10-1
10-18 — >10-14
10-35 — > 16-68
10-2 — > 10-2
10-19 — > 10-15
10-36 — > 10-24
10-3 — ► 10-3
10-20 — > 10-16
10-37 — > 10-25
10-4 — ► 10-4
10-21 — 16-61
10-38 — > 10-26
10-5 — > 10-5
10-22 — 16-62
10-39 — > 16-69
10-6 — ► 10-6
10-23 — ► 16-63
10-40 — >10-27
10-7 — ► 10-7
10-24 — > 16-64
10-41 — > 10-28
10-8 — > 16-57
10-25 — 16-65
10-42 — > 10-29
10-9 — ► 16-58
10-26 — > 10-17
10-43 — > 10-30
10-10 — > 16-59
10-27 — > 10-18
10-44 — > 10-31
10-11 — > 16-60
10-28 — > 10-19
10-46 — > 10-32
10-12 — >10-8
10-29 — ► 16-66
10-47 — > 10-33
10-13 — > 10-9
10-30 — ► 16-67
10-48 — > 16-70
10-14 — > 10-10
10-31 — 10-20
10-49 — > 10-34
10-15 — >10-11
10-32 — > 10-21
10-50 — > 10-35
10-16 — >10-12
10-33 — > 10-22
10-51 — > 10-37
10-17 — > 10-13
10-34 — > 10-23
10-52 — > 10-38
PREFACE
IX
10-53 — -> 10-39
10-54 — -> 10-40
10-55 — ► 16-72
10-56 — > 16-73
10-57 — > 16-74
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10-59 — -> 16-76
10-60 — > 16-77
10-61 — > 10-41
10-62 — > 10-42
10-63 — -> 10-36
10-64 — -> 10-42
10-65 — ► 10-43 10-66 — > 10-44 10-67 — ► 10-45 10-68 — -> 10-46 10-69 — > 10-47 10-70 — > 10-48 10-71 — ► 10-49 10-72 — ► 10-50 10-73 — > 10-51 10-74 — -> 10-52 10-75 — > 10-53 10-76 — ► 10-54 10-77 — ► 16-79 10-78 — ► 16-80 10-79 — > 19-53 10-80 — > 19-57 10-81 — ► 19-54 10-82 — > 19-58 10-83 — > 19-66 10-84 — > 19-56 10-85 — -> 19-35 10-86 — -> 19-59 10-87 — ► 19-67 10-88 — ► 19-70 10-89 — -> 19-39 10-90 — > 19-40 10-91 — > 19-41 10-92 — ► 10-55 10-93 — ► 10-56 10-94 — ► 10-57 10-95 — > 10-58 10-96 — -> 10-59 10-98 — ► 10-61 10-99 — ► 10-63
10-100
— -> 10-60
10-101
— * 10-64
10-102
— ► 10-65
10-103
— ► 10-66
10-104
— * 10-67
10-105
10-68
10-106
— * 10-70
10-107
— ► 10-71
10-108
— ► 10-72
10-109
— ► 10-73
10-110
— * 10-74
10-111
10-75
10-112
— ► 10-76
10-113
— ► 10-77
10-114
— > 16-81
10-115
— * 16-82
10-116
— ► 16-83
10-117
— * 16-84
10-118
— ► 16-85
10-119
— ► 16-86
10-120
— ► 16-87
10-121
— ► 16-88
10-122
— * 16-89
10-123
— ► 16-90
10-124
— ► 16-100
10-125
— > 16-101
10-126
— > 16-102
10-127
— ► 16-103
10-128
— ► 16-104
10-129
— ► 16-105
11-1 —
>11-1
11-2 —
>11-2
11-3 — > 11-3
11-4 — >11-4
11-5 — >11-5
11-6 —
>11-6
11-7 — >11-7
11-8 —
>11-8
11-9
11-10 — >11-9
11-11 — >11-10
11-12 —
->11-11
11-13-
— > 11-15
H_14 — >11-17
11-15 — >11-18
11-19 — > 11-19
11-20 — > 11-20
11-21 — > 11-21
11-22 — > 11-12
11-23 — > 11-13
11-24 — > 11-14
11-25 — > 11-22
11-26 — > 11-23
11-27 — > 11-24
11-28 — > 11-25
11-29 — > 11-26
11-30 — > 11-27
11-31 — > 11-28 11-32 — > 11-29 11-33 — > 11-30 11-34 — > 11-31 11-35 — > 11-32 11-36 — > 11-33 11-37 — > 11-34 11-38 — > 11-35 11-39 — > 11-36 H_40 — > 11-37 11-41 — > 11-38 11-42 — > 11-39
11- 43 — > 11-40 H_44 — > 11-41
12- 1 — > 12-1 12-2 — > 12-2
12-3 — > 12-3
12-4 — > 12-4
12-5 — » 12-5 12-6 — -> 12-6
12-7 — > 12-7 12-8 — > 12-8
12-9 — > 12-10 12-10 — > 12-11 12-11 — > 12-12
12-12 — > 12-13
12-13 — > 12-14
12-14 — > 12-16
12-15 — > 12-18
12-16 — > 12-19
12-17 — > 12-20
12-18 — > 10-69
12-19 — > 12-21
X
PREFACE
12-20 — -+ 12-22
12-21 — -+ 12-23
12-22 — » 12-17
12-23 — > 12-24
12-24 — > 12-25
12-25 — -+ 12-26
12-26 — -+ 12-27
12-27 — > 12-30
12-28 — > 12-31
12-29 — > 12-32
12-30 — -+ 12-33
12-31 — > 12-34
12-32 — » 12-35
12-33 — > 12-36
12-34 — » 12-37
12-36 — > 12-38
12-37 — 12-39
12-38 — -+ 12-40
12-39 — > 12-41
12-40 — » 12-42
12-41 — > 12-43
12-42 — -+ 12-44
12-43 — -+ 12-45 12-44 — » 12-46 12-45 — > 12-47 12-46 — » 12-48 12-47 — > 13-19 12-48 — -+ 12-49 12-49 — » 12-50 12-50 — » 13-24 12-51 — > 12-51 12-52 — -+ 12-52
12- 53 — -+ 12-53
13- 1 — » 13-1
13-2 — > 13-2
13-3 — > 13-3
13-4 — > 13-4
13-5 — > 13-5
13-6 — » 13-6
13-7 — > 13-7
13-8 — » 19-55
13-10 — > 13-8
13-11 — > 13-9
13-12 — » 13-14
13-13 — > 13-15
13-14 — > 13-11
13-15 — > 13-17
13-16 — > 13-18
13-17 — > 13-20
13-18 — > 13-21
13-19 — > 13-22
13-20 — > 13-23
13-21 — > 13-30
13-22 — » 13-31
13-23 — -+ 13-32
13- 24 — > 13-33
14- 1 — > 14-1
14-2 — > 14-3
14-3 — 14-4
14-4 — > 19-14
14-5 — 14-5
14-6 — > 19-23
14-7 — > 14-6
14-8 — > 14-7
14-9 — 14-8
14-10 — > 14-9
14-11 — > 14-10
14-12 — > 12-9
14-13 — > 14-11
14-14 — > 14-12
14-15 — > 14-14
14-16 — > 14-16
14-17 — > 13-27
14-18 — > 13-26
14-19 — > 13-10
14-20 — > 12-15
14-21 — 14-17
14-22 — > 14-18
14-23 — > 14-19 14_24 — 19-69 14-25 — 14-20 14-26 — > 14-21 14-27 — > 14-22 14-28 — > 13-28 14-29 — > 13-25 14-30 — > 14-23 14-31 — -+ 14-24 14-32 — > 14-26
14-33 — 14-25
14_34 14_27
14-35 — > 14-28 14_36 — ^ 14-29 14-37 — -+ 14-30
14- 38 — > 14-31
14_39 14_32
15- 1 — >15-1
15-2 — >15-2
15-3 — > 15-3
15-4 — >15-4
15-5 15-5
15-6 — > 15-6
15-7 — ► 15-7
15-8 — > 15-8
15-9 — > 15-9
15-10 — >15-10
15-11 — >15-11
15-12 — >15-12
15-13 — >15-14
15-14 — >15-13
15-15 — >15-15
15-16 — > 15-16
15-17 — > 15-17
15-18 — > 15-18
15-19 — > 15-20
15-20 — > 15-23
15-21 — > 15-24
15-22 — >15-21
15-23 — > 15-22 15-24 — > 15-25 15-25 — > 15-27 15-26 — > 15-28 15-27 — > 15-32 15-28 — > 15-33 15-29 — > 15-36 15-30 — > 15-35 15-31 — >15-37 15-32 — > 15-34 15-33 — > 15-38 15-34 — > 15-19 15-35 — 15-29 15-36 — > 15-30 15-37 — > 15-39
PREFACE
XI
15-38
— -+ 15-41
16-18 — >
16-22
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15-39
— -+ 15-40
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16-14
16-63 — -+ 16-95
15-40
— > 15-42
16-20 —
16-15
16-64 — » 16-96
15-41
— ► 15-43
16-21 — >
16-16
16-65 — -+ 16-97
15-42
— -+ 15-44
16-22 — ■+
16-23
16-66 — > 16-98
15-43
— -+ 15-45
16-23 — *
19-36
16-67 — > 16-99
15-44
— -+ 15-46
16-24 — ■*
19-42
15-45
— ► 15-47
16-25 — ►
19-43
17-1 — » 17-1
15-46
— ► 15-48
16-26 — ►
19-44
17-2 — > 17-2
15-47
— ► 15-49
16-27 — ►
16-24
17-3 — > 17-4
15-48
— -+ 15-50
16-28 — *
16-25
17-4 — > 17-5
15-49
— -+ 15-62
16-29 — 4
16-26
17-5 — > 17-6
15-50
— » 15-51
16-30^
16-27
17-6 — » 17-7
15-51
— » 15-52
16-31 — »
16-28
17-7 — > 17-8
15-52
— ► 15-53
16-32 — ►
16-29
17-8 — > 17-9
15-53
— -+ 15-54
16-33 deleted - combined
17-9 — > 17-10
15-54
— -+ 15-55
with 10-1
15
17-10 — > 17-11
15-55
— -+ 15-56
16-34 — 4
16-30
17-11 — > 17-12
15-56
— ► 15-57
16-35 — ►
16-31
17-12 — > 17-13
15-57
— ► 15-58
16-36 — »
16-32
17-13 — > 17-14
15-58
— -+ 15-60
16-37 — 4
16-33
17-14 — > 17-15
15-59
— -+ 15-61
16-38 — ■*
16-34
17-15 — > 17-16
15-60
— -+ 15-59
16-39 — ■*
16-35
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— ► 15-63
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16-36
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15-62
— ► 15-64
16-41 — >
16-38
17-18 — > 17-19
15-63
— ► 15-65
16-42 — >
16-41
17-19 — > 17-3
15-64
— -+ 15-66
16-43 — 4
16-42
17-20 — > 17-20
16-44 — ■*
16-39
17-21 — > 17-21
16-1 -
— ► 16-1
16-45 — ►
16-40
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16-2-
— > 16-2
16-46 — ►
16-43
17-23 — > 17-23
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— > 16-3
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16-44
17-24 — > 17-24
16-4 — > 16-4
16-48 —
16-45
17-25 — > 17-25
16-5 — > 16-5
16-49 — *
16-50
17-26 deleted
16-6 — > 16-7
16-50 — 4
16-51
combined with 17-25
16-7-
— > 16-8
16-51 — >
16-52
17-27 — > 17-26
16-8-
— > 16-9
16-52 — ►
16-53
17-28 — > 17-27
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16-53 — ■*
16-54
17-29 — > 17-28
16-10
— > 16-11
16-54 — 4
16-55
17-30 — > 17-29
16-11
— > 16-12
16-55 — ■*
16-56
17-31 deleted
16-12
— > 16-13
16-56 — ►
16-91
combined with 17-30
16-13
— » 16-18
16-57 — ►
16-6
17-32 — > 17-30
16-14
— » 16-17
16-58 — ►
16-92
17-33 — > 17-31
16-15
— > 16-19
16-59 — 4
16-93
17_34 — > 17-32
16-16
— -+ 16-20
16-60 — *
16-94
17-35 — ► 17-33
16-17
— > 16-21
16-61 — ►
16-46
17-36 — » 17-34
Xii PREFACE
17-37 — > 17-35
18-34 — > 18-34
19-27 — > 19-29
17-38 — > 17-36
18-35 — > 18-35
19-28 — > 19-31
17-39 — > 17-37
18-36 — > 18-36
19-29 — >19-24
17-40 — » 17-38
18-37 — > 18-37
19-30 — > 19-32
18-38 — > 18-38
19-31 — > 19-33
18-1 — > 18-1
18-39 — > 18-39
19-32 — > 19-34
18-2 — > 18-2
18-40 — > 18-40
19-33 — -+ 19-61
18-3 — » 18-3
18-42 — >18-42
19-34 — > 19-37
18-4 — >18-4
18-43 — > 18-43
19-35 — > 19-64
18-5 — > 18-5
18-44 — >18-44
19-36 — > 19-62
18-6 — > 18-6
19-37 — -+ 19-63
18-7 — >18-7
19-1 — > 19-1
19-38 — > 19-38
18-8 — > 18-8
19-2 — >19-2
19-39 — > 19-65
18-9 — > 18-9
19-3 — > 19-3
19-40 deleted
18-10 — > 18-10 .
19-4 — >19-4
incorporated into 10-85
18-11 — > 18-11
19-5 — * 19-5
19_41 19_45
18-12 — > 18-12
19-6 — > 19-6
19-42 — -+ 19-46
18-13 — > 18-13
19-7 — > 19-7
19_43 — > 19-47
18-14 — >18-14
19-8 — > 19-8
19-44 — >19-48
18-15 — > 18-15
19-9 — > 19-9
19_45 19_50
18-16 — > 18-16
19-10 — > 19-10
19-46 — > 19-51
18-17 — >18-17
19-11 — > 19-11
19-47 — >19-71
18-18 — > 18-18
19-12 — > 19-12
19_48 — 19-68
18-19 — > 18-19
19-13 — > 19-13
19_49 — > 19-72
18-20 — >18-20
19_14 — >19-17
19-50 — > 19-60
18-21 — >18-21
19-15 — > 19-15
19-51 — > 19-49
18-22 — > 18-22
19-16 — > 19-18
19-52 — >19-73
18-23 — > 18-23
19-17 deleted
19-53 — > 19-74
18-24 — >18-24
incorporated in 19-14
19-54 — > 19-75
18-25 — > 18-25
19-18 — > 19-19
19-55 — > 19-76
18-26 — >18-26
19-19 — >19-20
19-56 — > 19-77
18-27 — >18-27
19-20 — > 19-21
19-57 — > 19-78
18-28 — > 18-28
19-21 — >19-22
19-58 — > 19-79
18-29 — > 18-29
19-22 — 19-25
19-59 — > 19-80
18-30 — > 18-30
19-23 — > 19-27
19-60 — > 19-81
18-31 — > 18-31
19-24 — > 19-28
19-61 — -+ 19-82
18-32 — > 18-32
19-25 — -+ 19-30
19-62 — > 19-83
18-33 — > 18-33
19-26 — 19-26
19-63 — -+ 19-84
BIOGRAPHICAL NOTE
Professor Michael B. Smith was bom in Detroit, Michigan in 1946 and lived there until 1957. In 1957, he and his family moved to Madison Heights, Virginia, where he attended high school and then Ferrum Jr. College, where he graduated with an A. A in 1966. Professor Smith then transferred to Virginia Polytechnic Institute (Virginia Tech), and graduated with a B.S in chemistry in 1969. After working 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 Purdue University under the mentorship of Professor Joseph Wolinsky. Professor Smith graduated with a Ph.D. in Organic chemistry in 1977. He then spent one year as a faculty research associate at the Arizona State University, in the Cancer Research Institute directed by Professor George R. Pettit. Professor Smith spent a second year doing postdoctoral work at the Massachusetts Institute of Technology under the mentorship of Professor Sidney Hecht. In 1979 Professor Smith began his independent academic career, where he now holds the rank of full professor.
Professor smith is the author of approximately 70 independent research articles, and is the author of 14 published books. The books include the 5th edition of March’s Advanced Organic Chemistry (Wiley), volumes 6-11 of the Compendium of Organic Synthetic Methods (Wiley), Organic Chemistry a Two Semester Course (HarperCollins) into its 2nd edition, and Organic Synthesis (McGraw-Hill) through its 2nd edition. The 3rd edition of the Organic Synthesis book is due out in 2007, published by Wavefunction, Inc.
Professor Smith’s current research involves the synthesis and structural verifica¬ tion of several bioactive lipids obtained from the dental pathogen Porphyromonas gingivalis. Another area of research examines the chemical reactivity of conducting polymers such as poly(ethylenedioxy)thiophene (PEDOT). Such polymers are sup¬ posed to be chemically inert but, in fact, induce a variety of chemical reactions, including Friedel-Crafts alkylation of aromatic compounds with alcohols. Another area of research involves the development of a dye-conjugate designed to target and image tumors, as well as the total synthesis of anti-cancer phenanthridone alkaloids such as pancratistatin.
xv
ABBREVIATIONS
Ac
Acetyl
acac
Acetylacetonato
AIBN
Azoisobutyronitrile
aq.
Aqueous
9-Borabicyclo [3 .3.1] nonylboryl
y^H3
9-BBN
BER
BINAP
Bn
Bz
BOC
bpy (bipy)
Bu
CAM
CAN
c-
cat.
Cbz
Chirald
Cod
Cot
Cp
CSA
CTAB
9-Borabicyclo[3.3.1]nonane Borohydride exchange resin (2R,35'),2,2/-bis(diphenylphosphino)- 1 , l'-binapthyl Benzyl
Benzoyl O
tert- Butoxy carbonyl 2,2/-Bipyridyl n- Butyl
Carboxamidomethyl Ceric ammonium nitrate Cyclo- Catalytic Carbobenzyloxy
Ot-Bu
-ch2ch2ch2ch3
(NH)2Ce(N03)6
o
xs
OCH2Ph
(25',3^)-(+)-4-dimethylamino-l,2-diphenyl-3-methylbutan-2-ol 1,5-Cyclooctadiene (ligand)
1,3,5,7-Cyclooctatetraene (ligand)
Cyclopentadienyl Camphorsulfonic acid
Cetyltrimethylammonium bromide Ci6H33NMe3+Br
Cy (c-CeHn) °C
DABCO
dba
DBE
DBU
DBN
DCC
DCE
Cyclohexyl
Temperature in degrees Centigrade
1 .4- Diazobicyclo[2.2.2]octane Dibenzylidene acetone
1 .2- Dibromoethane l,8-Diazabicyclo[5.4.0]undec-7-ene
1.5- Diazabicyclo[4.3.0]non-5-ene
1 .3- Dicyclohexylcarbodiimide
1 ,2-Dichloroethane
BrCH2CH2Br
c-C6H13— N=C=N-c-C6H13
cich2ch2ci
xvii
ABBREVIATIONS
xviii
DDQ % de DEA DEAD Dibal-H Diphos (dppe) Diphos-4 (dppb) DMAP DMA DME
DMF
dmp
DMSO
dpm
dppb
dppe
dppf
dppp
dvb
e-
% ee
EE
Et
EDA
EDTA
FMN
fod
Fp
FVP
h
hv
1,5-HD
HMPA
HMPT
/Pr
IR
LICA (LIPCA)
LDA
LHMDS
LTMP
MABR
2.3- Dichloro-5,6-dicyano-l,4-benzoquinone % Diasteromeric excess Diethylamine Diethylazodicarboxylate Diisobutylaluminum hydride
1.2- bis(Diphenylphosphino)ethane
1.4- bis(Diphenylphosphino)butane 4-Dimethylaminopyridine Dimethylacetamide
1 .2- Dimethoxy ethane
HN(CH2CH3)2
Et02C-N=NC02Et
(Me2CHCH2)2AIH
Ph2PCH2CH2PPh2
Ph2P(CH2)4PPh2
MeOCH2CH2OMe
N,Nf -Dimethy lformamide
o
N(CH3)2
Ph2P(CH2)3PPh2
EtO(Me)HCO—
-CH2CH3
h2nch2ch2nh2
bis-[ 1 ,3-Di(p-methoxyphenyl)- 1 ,3-propanedionato]
Dimethyl sulfoxide Dipivaloylmethanato
1.4- bis(Diphenylphosphino)butane Ph2P(CH2)4PPh2
l,2-bis(Diphenylphosphino)ethane Ph2PCH2CH2CH2PPh2 bis(Diphenylphosphino)ferrocene 1 , 3 -bis(Dipheny lphosphino)propane Divinylbenzene Electrolysis
% Enantiomeric excess 1 -Ethoxy ethyl Ethyl
Ethylenediamine Ethylenediaminetetraacetic acid Flavin mononucleotide
/Ws-(6,6,7,7,8,8,8)-Heptafluoro-2,2-dimethyl-3,5-octanedionate Cyclopentadienyl-bis(carbonyl iron)
Flash vacuum pyrolysis Hour (hours)
Irradiation with light
1.5- Hexadienyl Hexamethylphosphoramide Hexamethylphorous triamide Isopropyl Infrared
Lithium cyclohexylisopropylamide Lithium diisopropylamide Lithium hexamethyl disilazide Lithium 2,2,6 , 6- tetramethy lpiperidide Methylaluminum bis(4-bromo-2,6-di-te/V-butylphenoxide)
(Me3N)3P=0 (Me3N)3P — CHMe2
LiN(/Pr)2
LiN(SiMe3)2
ABBREVIATIONS
XIX
MAD
bis(2,6-Di-^rt-butyl-4-methylphenoxy)methyl aluminum
mCPBA
meto-Chloroperoxybenzoic acid
Me
Methyl
-ch3
MEM
(3-Methoxyethoxymethyl
MeOCH2CH2OCH2—
Mes
Mesityl
2,4,6-tri-Me-C6H2
MOM
Methoxymethyl
MeOCH2—
Ms
Methanesulfonyl
ch3so2-
MS
Molecular sieves (3 A or 4 A)
MTM
Methylthiomethyl
ch3sch2-
NAD
Nicotinamide adenine dinucleotide
NADP
Sodium triphosphopyridine nucleotide
Napth
Naphthyl (C10H8)
NBD
Norbornadiene
NBS
A-Bromosuccinimide
NCS
A-Chlorosuccinimide
NIS
A-Iodosuccinimide
Ni(R)
Raney nickel
NMP
A-Methyl-2-pyrrolidinone
NY
New York
NMR
Nuclear magnetic resonance
Oxone
2 KHS05KHS04K2S04
®
Polymeric backbone
PCC
Pyridinium chlorochromate
PDC
Pyridinium dichromate
PEG
Polyethylene glycol
Ph
Phenyl
PhH
Benzene
PhMe
Toluene
Phth
Phthaloyl
pic
2-Pyridinecarboxylate
+o
Pip
Piperidyl
PMP
4-Methoxyphenyl
Pr
^-Propyl
-ch2ch2ch3
Py
Pyridine
o
quant.
Quantitative yield
Red-Al
[(Me0CH2CH20)2AlH2]Na
sBu
sec-Butyl
CH3CH2CH(CH3)
vBuLi
sec-Butyllithium
CH3CH2CH(Li)CH3
Siamyl
Diisoamyl
(CH3)2CHCH(CH3)-
TADDOL
oc,oc,oc V-Tetraaryl-4,5-dimethoxy- 1 ,3-dioxolane
TASF
£rA-(Diethylamino)sulfonium difluorotrimethyl silicate
TBAF
Tetrabutylammonium fluoride
«-Bu4N+F“
TBDMS
te/t-Butyldimethylsilyl
r-BuMesSi
TBHP
te/t-Butylhydroperoxide (7-BuOOH)
Me3COOH
XX
ABBREVIATIONS
£-Bu
tert- Butyl
-C(CH3)3
TBS
tert- Butyl dimethylsilyl
?-BuMe2Si
TEBA
Triethylbenzylammonium
Bn(CH3)3N+
TEMPO
Tetramethylpiperdinyloxy free radical
TFA
Trifluoroacetic acid
cf3cooh
TFAA
Trifluoroacetic anhydride
(CF3CO)20
Tf (OTf)
Triflate
-so2cf3(-oso2cf3)
THF
Tetrahydrofuran
THP
Tetrahydropyran
TMEDA
Tetramethylethylenediamine
Me2NCH2CH2NMe2
TMG
1 , 1 ,3,3-Tetramethylguanidine
TMS
Trimethylsilyl
— Si(CH3)3
TMP
2,2,6,6-Tetramethylpiperidine
TPAP
tetra-ft-Propylammonium perruthenate
Tol
Tolyl
4MeC6H4
Tr
Trityl
— CPh3
TRIS
Triisopropylphenylsulfonyl
Ts(Tos)
Tosyl = /?-Toluenesulfonyl
4-MeC6H4
UV
Ultraviolet
Xc
Chiral auxiliary
PART ONE
This book contains 19 chapters. Chapters 10-19, which make up Part 2, are directly concerned with organic reactions and their mechanisms. Chapters 1-9 may be thought of as an introduction to Part 2. The first five chapters deal with the structure of organic compounds. These chapters discuss the kinds of bonding important in organic chemistry, the three-dimensional structure of organic molecules, and the structure of species in which the valence of carbon is less than 4. Chapters 6-9 are concerned with other topics that help to form a background to Part 2: acids and bases, photochemistry, the relationship between structure and reactivity, and a general discussion of mechanisms and the means by which they are determined.
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. In Chapter 2, we will consider delocalized bonding , in which electrons are shared by more than two nuclei.
COVALENT BONDING1
Wave mechanics is based on the fundamental principle that electrons behave as waves (e.g., they can be diffracted) and that consequently a wave equation can be written for them, in the same sense that light waves, sound waves, and so on can be described by wave equations. The equation that serves as a mathematical model for electrons is known as the Schrodinger equation , which for a one-electron system is
52\| / 52\| / 52\| / 87i2m
bx2 by2 + bz2 + h 2
(E - V)v|/ = 0
where m is the mass of the electron, E is its total energy, V is its potential energy, and h is Planck’s constant. In physical terms, the function ^ expresses the square root of the probability of finding the electron at any position defined by the coordinates x, y, and z, where the origin is at the nucleus. For systems containing more than one electron, the equation is similar, but more complicated.
'The treatment of orbitals given here is necessarily simplified. For much fuller treatments of orbital theory as applied to organic chemistry, see Matthews, P.S.C. Quantum Chemistry of Atoms and Molecules , Cambridge University Press, Cambridge, 1986, Clark, T. A Handbook of Computational Chemistry , Wiley, NY, 1985; Albright, T.A.; Burdett, J.K.; Whangbo, M. Orbital Interactions in Chemistry, Wiley, NY, 1985; Mac Weeny, R.M. Coulson’s Valence, Oxford University Press, Oxford, 1980; Murrell, J.N.; Kettle, S.F.A; Tedder, J.M. The Chemical Bond, Wiley, NY, 1978; Dewar, M.J.S.; Dougherty. R.C. The PMO Theory of Organic Chemistry, Plenum, NY, 1975; Zimmerman, H.E. Quantum Mechanics for Organic Chemists, Academic Press, NY, 1975; Borden, W.T. Modem Molecular Orbital Theory for Organic Chemists, Prentice-Hall, Englewood Cliffs, NJ, 1975; Dewar, M.J.S. The Molecular Orbital Theory of Organic Chemistry, McGraw-Hill, NY, 1969; Liberies, A. Introduction to Molecular Orbital Theory, Holt, Rinehart, and Winston, NY, 1966.
March ’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Sixth Edition, by Michael B. Smith and Jerry March Copyright © 2007 John Wiley & Sons, Inc.
3
4 LOCALIZED CHEMICAL BONDING
Fig. 1.1. ( a ) The Is orbital. ( b ) The three 2 p orbitals.
The Schrodinger equation is a differential equation, which means that solutions of it are themselves equations, but the solutions are not differential equations. They are simple equations for which graphs can be drawn. Such graphs, which are three- dimensional (3D) pictures that show the electron density, are called orbitals or elec¬ tron clouds. 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 prob¬ ability of finding the electron is extremely small.2 Also note that in Fig. 1.1 some lobes of the orbitals are labeled + 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 function Tf When two parts of an orbital are separated by a node, always has 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 Schrodinger equation can be solved exactly only for one- electron systems, such as the hydrogen atom. If it could be solved exactly for mole¬ cules containing two or more electrons,3 we would have a precise picture of the shape of the orbitals available to each electron (especially for the important ground state) and the energy for each orbital. Since exact solutions are not available, drastic approximations must be made. There are two chief general methods of approxima¬ tion: the molecular-orbital method and the valence-bond method.
In the molecular-orbital method, bonding is considered to arise from the overlap of atomic orbitals. When any number of atomic orbitals overlap, they combine to
2When wave-mechanical calculations are made according to the Schrodinger 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. and 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 review of relativistic effects on chemical structures in general, see Pyykko, P. Chem. Rev. 1988, 88, 563.
3For a number of simple systems containing two or more electrons, such as the H2 molecule or the He atom, approximate solutions are available that are so accurate that for practical purposes they are as good as exact solutions. See, for example, 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.
CHAPTER 1
COVALENT BONDING 5
form an equal number of new orbitals, called molecular orbitals. Molecular orbitals differ from atomic orbitals in that they are clouds that surround the nuclei of two or more atoms, rather than just one atom. In localized bonding the number of atomic orbitals that overlap 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 energy. Orbitals of lower energy fill first. Since the two original atomic orbitals each held one electron, both of these electrons can now go into the new molecular bonding orbital, since any orbital can hold two electrons. The antibonding orbital remains empty in the ground state. The greater the overlap, the stronger the bond, although total overlap is prevented by repulsion between the nuclei. Figure 1.2 shows the bonding and antibonding orbitals that arise by the overlap of two Is electrons. Note that since the antibonding orbital has a node between the nuclei, there is practically no electron density in that area, so that this orbital cannot be expected to bond very well. Molecular orbitals formed by the over¬ lap of two atomic orbitals when the centers of electron density are on the axis com¬ mon to the two nuclei are called a {sigma) orbitals, and the bonds are called a bonds. Corresponding antibonding orbitals are designated & *. Sigma orbitals are formed not only by the overlap of two s orbitals, but also by the overlap of any of the kinds of atomic orbital (s, p, d, or/) whether the same or different, but the two lobes that overlap must have the same sign: a positive s orbital can form a bond only by overlapping with another positive s orbital or with a positive lobe of a p, d , or/ orbital. Any s orbital, no matter what kind of atomic orbitals it has arisen from, may be represented as approximately ellipsoidal in shape.
Orbitals are frequently designated by their symmetry properties. The a orbital of hydrogen is often written \) tg. The g stands for gerade. A gerade orbital is one in which the sign on the orbital does not change when it is inverted through its center of symmetry. The a* orbital is ungerade (designated \| ju). An ungerade orbital changes sign when inverted through its center of symmetry.
Fig. 1.2. Overlap of two Is orbitals gives rise to a a and a a orbital.
6 LOCALIZED CHEMICAL BONDING
In molecular-orbital calculations, a wave function is formulated that is a linear combination of the atomic orbitals that have overlapped (this method is often called the linear combination of atomic orbitals , or LCAO). Addition of the atomic orbitals gives the bonding molecular orbital:
v|/ = Ca'I'a + Cb'I'b O1)
The functions v| iA and v|/B are the functions for the atomic orbitals of atoms A and B, respectively, and cA and cB represent weighting factors. Subtraction is also a linear combination:
v|/ = cAv|/A - cbv|/b (1-2)
This gives rise to the antibonding molecular orbital.
In the valence-bond method, a wave equation is written for each of various pos¬ sible electronic structures that a molecule may have (each of these is called a cano¬ nical form), and the total v| / is obtained by summation of as many of these as seem plausible, each with its weighting factor:
i f = +c2\|/2 H - (1-3)
This resembles Eq. (1), but here each \| / represents a wave equation for an imagin¬ ary canonical form and each c is the amount contributed to the total picture by that form. For example, a wave function can be written for each of the following canonical forms of the hydrogen molecule:4
H-H H H+ +H H
Values for c in each method are obtained by solving the equation for various values of each c and choosing the solution of lowest energy. In practice, both methods give similar solutions for molecules that contain only localized electrons, and these are in agreement with the Lewis structures long familiar to the organic chemist. Delocalized systems are considered in Chapter 2.
MULTIPLE VALENCE
A univalent atom has only one orbital available for bonding. But atoms with a valence of 2 or more must form bonds by using at least two orbitals. An oxygen atom has two half-filled orbitals, 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 nuclei should form an angle of 90° with the oxygen nucleus, since the two available orbitals on oxygen are p orbitals, which are perpendicular. Similarly, we should expect that nitrogen, which has three mutually perpendicular p orbitals, would have bond angles of 90° when it forms three single bonds. However, these are not the observed bond angles. The bond
4In this book, a pair of electrons, whether in a bond or unshared, is represented by a straight line.
CHAPTER 1
HYBRIDIZATION 7
angles are,5 in water, 104o27', and in ammonia, 106°46/. For alcohols and ethers the angles are even larger (see p. 25). A discussion of this will be deferred to p. 25, but it is important to note that covalent compounds do have definite bond angles. Although the atoms are continuously vibrating, the mean position is the same for each molecule of a given compound.
HYBRIDIZATION
Consider the case of mercury. Its electronic structure is
[Xe core]4/145(/106^2
Although it has no half-filled orbitals, it has a valence of 2 and forms two covalent bonds. We can explain this by imagining that one of the 6s electrons is promoted to a vacant 6 p orbital to give the excited configuration
[Xe cor e\4fu5dw6sl6pl
In this state, the atom has two half-filled orbitals, but they are not equivalent. If bonding were to occur by the overlap of these orbitals with the orbitals of external atoms, the two bonds would not be equivalent. The bond formed from the 6p orbital would be more stable than the one formed from the 6s orbital, since a larger amount of overlap is possible with the former. A more stable situation is achieved when, in the course of bond formation, the 6s and 6 p orbitals combine to form two new orbi¬ tals that are equivalent; these are shown in Fig. 1.3.
Since these new orbitals are a mixture of the two original orbitals, they are called hybrid orbitals. Each is called an sp orbital, since a merger of an s and a p orbital was required to form it. The sp orbitals, each of which consists of a large lobe and a very small one, are atomic orbitals, although they arise only in the bonding process and do not represent a possible structure for the free atom. A mercury atom forms
Fig. 1.3. The two sp orbitals formed by mercury. 5Bent, H.A. Chem. Rev. 1961, 61, 275, p. 277.
8 LOCALIZED CHEMICAL BONDING
its two bonds by overlapping each of the large lobes shown in Fig. 1.3 with an orbi¬ tal from an external atom. This external orbital may be any of the atomic orbitals previously considered (s, p, d , or/) or it may be another hybrid orbital, although only lobes of the same sign can overlap. In any of these cases, the molecular orbital that arises is called a a orbital since it fits our previous definition of a a orbital.
In general, because of mutual repulsion, equivalent orbitals lie as far away from each other as possible, so the two sp orbitals form an angle of 180°. This means that HgCl2, for example, should be a linear molecule (in contrast to H20), and it is. This kind of hybridization is called digonal hybridization. An sp hybrid orbital 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 greater overlap. Although it would require energy to promote a 6s electron to the 6p state, the extra bond energy more than makes up the difference.
Many other kinds of hybridization are possible. Consider boron, which has the electronic configuration
\s12s12p'
yet has a valence of 3. Once again we may imagine promotion and hybridization:
1?2,V Is22sWM hybridiZati°"> 1 s2(sp2)3
In this case, there are three equivalent hybrid orbitals, each called sp2 ( trigonal hybridization). This method of designating hybrid orbitals is perhaps unfortunate since nonhybrid orbitals are designated by single letters, but it must be kept in mind that each of the three orbitals is called sp2. These orbitals are shown in Fig. 1.4. The three axes are all in one plane and point to the corners of an equilateral triangle. This accords with the known structure of BF3, a planar molecule with angles of 120°.
The case of carbon (in forming four single bonds) may be represented as
1^2 ,2
Hi vciti r»n
Fig. 1.4. The three sp2 and the four sp 3 orbitals.
CHAPTER 1
MULTIPLE BONDS 9
There are four equivalent orbitals, each called sp3, which point to the corners of a regular tetrahedron (Fig. 1.4). The bond angles of methane would thus be expected to be 109° 28', which is the angle for a regular tetrahedron.
Although the hybrid orbitals discussed in this section satisfactorily account for most of the physical and chemical properties of the molecules involved, it is necessary to point out that the sp 3 orbitals, for example, stem from only one possible approxi¬ mate solution of the Schrodinger equation. The s and the three p atomic orbitals can also be combined in many other equally valid ways. As we shall see on p. 13, the four C— H bonds of methane do not always behave as if they are equivalent.
MULTIPLE BONDS
If we consider the ethylene molecule in terms of the molecular-orbital concepts dis¬ cussed so far, we have each carbon using sp 2 orbitals to form bonds with the three atoms to which it is connected. These sp 2 orbitals arise from hybridization of the 2 sl, 2 p\, and 2 ply electrons of the promoted state shown on p. 8. We may consider that any carbon atom that is bonded to only three different atoms uses sp2 orbitals for this bonding. Each carbon of ethylene is thus bonded by three a bonds: one to each hydrogen and one to the other carbon. Each carbon therefore has another elec¬ tron in the 2 pz orbital that is perpendicular to the plane of the sp2 orbitals. The two parallel 2 pz orbitals can overlap sideways to generate two new orbitals, a bonding and an antibonding orbital (Fig. 1.5). Of course, in the ground state, both electrons go into the bonding orbital and the antibonding orbital remains vacant. Molecular orbitals formed by the overlap of atomic orbitals whose axes are parallel are called 7i orbitals if they are bonding and n if they are antibonding.
In this picture of ethylene, the two orbitals that make up the double bond are not equivalent.6 The a orbital is ellipsoidal and symmetrical about the C— C axis. The 7i orbital is in the shape of two ellipsoids, one above the plane and one below. The plane itself represents a node for the n orbital. In order for the p orbitals to maintain maximum overlap, they must be parallel. This means that free rotation is not pos¬ sible about the double bond, since the two p orbitals would have to reduce their overlap to allow one H— C— H plane to rotate with respect to the other. The six atoms of a double bond are therefore in a plane with angles that should be ^120°. Double bonds are shorter than the corresponding single bonds because maximum stability is obtained when the p orbitals overlap as much as possible. Double bonds between carbon and oxygen or nitrogen are similarly represented: they consist of one a and one n orbital.
In triple-bond compounds, carbon is connected to only two other atoms and hence uses sp hybridization, which means that the four atoms are in a straight
6The double bond can also be pictured as consisting of two equivalent orbitals, where the centers of electron density point away from the C— C axis. This is the bent-bond or banana-bond picture. Support for this view is found in Pauling. L. Theoretical Organic Chemistry, The Kekule Symposium, Butterworth, London, 1959, pp. 2-5; Palke, W.E. J. Am. Chem. Soc. 1986, 108, 6543. However, most of the literature of organic chemistry is written in terms of the o-n picture, and we will use it in this book.
1 0 LOCALIZED CHEMICAL BONDING
Fig. 1.5. Overlapping p orbitals form a n and a n orbital. The a orbitals are shown in the upper figure. They are still there in the states represented by the diagrams below, but have been removed from the picture for clarity.
line (Fig. 1.6). 7 Each carbon has two p orbitals remaining, with one electron in each. These orbitals are perpendicular to each other and to the C— C axis. They overlap in the manner shown in Fig. 1.7 to form two n orbitals. A triple bond is thus composed of one a and two n orbitals. Triple bonds between carbon and nitro¬ gen can be represented in a similar manner.
Double and triple bonds are important only for the first-row elements carbon, nitrogen, and oxygen.8 For second-row elements multiple bonds are rare and
H
C C
Fig. 1.6. The a electrons of acetylene.
H
7For reviews of triple bonds, 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.
8This statement applies to the representative elements. Multiple bonding is also important for some transition elements. For a review of metal-metal multiple bonds, see Cotton, F.A. J. Chem. Educ. 1983, 60, 713.
CHAPTER 1
MULTIPLE BONDS 11
Fig. 1.7. Overlap of p orbitals in a triple bond for clarity, the a orbitals have been removed from the drawing on the left, although they are shown on the right.
compounds containing them are generally less stable9 because these elements tend to form weaker n bonds than do the first-row elements.10 The only ones of any importance at all are C=S bonds, and C=S compounds are generally much less stable than the corresponding C=0 compounds (however, see pn-dn bonding, p. $$$). Stable compounds with Si=C and Si=Si bonds are rare, but examples have been reported,11 including a pair of cis and trans Si=Si isomers.12
9For a review of double bonds between carbon and elements other than C, N, S, or O, see Jutzi, P. Angew. Chem. Int. Ed. 1975, 14, 232. For reviews of multiple bonds involving silicon and germanium, see Barrau, J.; Escudie, J.; Satge, J. Chem. Rev. 1990, 90, 283 (Ge only); Raabe, G.; Michl, J., in Patai, S. and Rappoport, Z. The Chemistry of Organic Silicon Compounds, part 2, Wiley: NY, 1989, pp. 1015-1142; Chem. Rev. 1985, 85, 419 (Si only); Wiberg, N. J. Organomet. Chem. 1984, 273, 141 (Si only); Gusel’nikov, L.E.; Nametkin, N.S. Chem. Rev. 1979, 79, 529 (Si only). For reviews of C=P and C+P bonds, see Regitz, 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. For reviews of other second-row double bonds, see West, R. Angew. Chem. Int. Ed. 1987, 26, 1201 (Si=Si bonds); Brook, A.G.; Baines, K.M. Adv. Organometal. Chem. 1986, 25, 1 (Si=C bonds); Kutney, G.W.; Turnbull, K. Chem. Rev. 1982, 82, 333 (S=S bonds). For reviews of multiple bonds between heavier elements, see Cowley, A.H.; Norman, N.C. Prog. Inorg. Chem. 1986, 34, 1; Cowley, A.H. Polyhedron 1984, 3, 389; Acc. Chem. Res. 1984, 17, 386. For a theoretical study of multiple bonds to silicon, see Gordon, M.S. Mol. Struct. Energ. 1986, 1, 101. 10For 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.
nFor Si=C bonds, see Brook, A.G.; Nyburg, S.C.; Abdesaken, F.; Gutekunst, B.; Gutekunst, G.; Kallury, R.K.M.R.; Poon, Y.C.; Chang, Y.; Wong-Ng, W. J. Am. Chem. Soc. 1982, 104, 5667; Schaefer III, H.F. Acc. Chem. Res. 1982, 15, 283; Wiberg, N.; Wagner, G.; Riede, J.; Muller, G. Organometallics 1987, 6, 32. For Si=Si bonds, see West, R.; Fink, M.J.; Michl, J. Science 1981, 214, 1343; Boudjouk, R; Han, B.; Anderson, K.R. J. Am. Chem. Soc. 1982, 104, 4992; 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. For an Si=N bond, see Wiberg, N.; Schurz, K.; Reber, G.; Muller, G. J. Chem. Soc. Chem. Commun. 1986, 591.
12Michalczyk, M.J.; West, R.; Michl, J. J. Am. Chem. Soc. 1984, 106, 821, Organometallics 1985, 4, 826.
1 2 LOCALIZED CHEMICAL BONDING
PHOTOELECTRON SPECTROSCOPY
Although the four bonds of methane are equivalent according to most physical and chemical methods of detection (e.g., neither the nuclear magnetic resonances (NMR) nor the infrared (IR) spectrum of methane contains peaks that can be attributed to different kinds of C— H bonds), there is one physical technique that shows that the eight valence electrons of methane can be differentiated. In this technique, called photoelectron spectroscopy ,13 a 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, and the difference between the energy of the radiation used and that of the ejected electron is the ionization potential of that elec¬ tron. A molecule that contains several electrons of differing energies can lose any one of them as long as its ionization potential is less than the energy of the radiation used (a single molecule loses only one electron; the loss of two electrons by any individual molecule almost never occurs). A photoelectron spectrum therefore consists of a series of bands, each corresponding to an orbital of a different energy. The spectrum gives a direct experimental picture of all the orbitals present, in order of their energies, provided that radiation of sufficiently high energy is used.14 Broad
Orbital 5
Fig. 1.8. Photoelectron spectrum of N2.15
13Only the briefest description of this subject is given here. For monographs, 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. Photoelectron 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; Carlson, T.A. Annu. Rev. Phys. Chem. 1975, 26, 211; Baker, C.; Brundle, C.R.; Thompson, M. Chem. Soc. Rev. 1972, 1, 355; Bock, H.; Mollere, P.D. J. Chem. Educ. 1974, 51, 506; 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 the electron spectroscopies, see Porter, H.Q.; Turner, D.W. Pure Appl. Chem. 1987, 59, 1343. 14The 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.
CHAPTER 1
PHOTOELECTRON SPECTROSCOPY 13
2 pxl tyy1
L _L
2py] 2 pzl
Nitrogen atom Nitrogen molecule
:NEN
Nitrogen atom
Fig. 1.9. Electronic structure of N2 (inner- shell electrons omitted).
bands usually correspond to strongly bonding electrons and narrow bands to weakly bonding or nonbonding electrons. A typical spectrum is that of N2, shown in Fig. 1.8. 15 The N2 molecule has the electronic structure shown in Fig. 1.9. The two 2s orbitals of the nitrogen atoms combine to give the two orbitals marked 1 (bonding) and 2 (antibonding), while the six 2 p orbitals combine to give six orbi¬ tals, three of which (marked 3, 4, and 5) are bonding. The three antibonding orbitals (not indicated in Fig. 1.9) are unoccupied. Electrons ejected from orbital 1 are not found in Fig. 1.8 because the ionization potential of these electrons is greater than the energy of the light used (they can be seen when higher energy light is used). The broad band in Fig. 1.8 (the individual peaks within this band are caused by different vibrational levels; see Chapter 7) corresponds to the four electrons in the degenerate orbitals 3 and 4. The triple bond of N2 is therefore composed of these two orbitals and orbital 1. The bands corresponding to orbitals 2 and 5 are narrow; hence these orbitals contribute little to the bonding and may be regarded as the two unshared pairs of N=N. Note that this result is contrary to that expected from a naive con¬ sideration of orbital roverlaps, where it would be expected that the two unshared pairs would be those of orbitals 1 and 2, resulting from the overlap of the filled 2s orbitals, and that the triple bond would be composed of orbitals 3, 4, and 5, resulting from overlap of the p orbitals. This example is one illustration of the value of photoelectron spectroscopy.
The photoelectron spectrum of methane16 shows two bands,17 at ^23 and 14 eV, and not the single band we would expect from the equivalency of the four C— H
15From Brundle, C.R.; Robin, M.B., in Nachod, F.C.; Zuckerman, JJ. Determination of Organic Structures by Physical Methods, Vol. 3, Academic Press, NY, 1971, p. 18.
16Brundle, C.R.; Robin, M.B.; Basch, H. J. Chem. Phys. 1970, 53, 2196; Baker, A.D.; Betteridge, D.; Kemp, N.R.; Kirby, R.E. J. Mol Struct. 1971, 8, 75; Potts, A.W.; Price, W.C. Proc. R. Soc. London, Ser A 1972, 326, 165.
17 A third band, at 290 eV, caused by the Is electrons of carbon, can also found if radiation of sufficiently high energy is used.
1 4 LOCALIZED CHEMICAL BONDING
bonds. The reason is that ordinary sp 3 hybridization is not adequate to explain phe¬ nomena involving ionized molecules (e.g., the CH4 radical ion, which is left behind when an electron is ejected from methane). For these phenomena it is necessary to use other combinations of atomic orbitals (see p. 9). The band at 23 eV comes from two electrons in a low-energy level (called the ax level), which can be regarded as arising from a combination of the 2s orbital of carbon with an appropriate combi¬ nation of hydrogen Is orbitals. The band at 14 eV comes from six electrons in a triply degenerate level (the t2 level), arising from a combination of the three 2 p orbitals of carbon with other combinations of Is hydrogen orbitals. As was men¬ tioned above, most physical and chemical processes cannot distinguish these levels, but photoelectron spectroscopy can. The photoelectron spectra of many other organic molecules are known as well,18 including monocyclic alkenes, in which bands <10 eV are due to Ti-orbital ionization and those >10 eV originate from ionization of s-orbitals only.19
ELECTRONIC STRUCTURES OF MOLECULES
For each molecule, ion, or free radical that has only localized electrons, it is pos¬ sible to draw an electronic formula, called a Lewis structure , that shows the location of these electrons. Only the valence electrons are shown. Valence electrons may be found in covalent bonds connecting two atoms or they may be unshared.20 The stu¬ dent must be able to draw these structures correctly, since the position of electrons changes in the course of a reaction, and it is necessary to know where the electrons are initially before one can follow where they are going. To this end, the following rules operate:
1. The total number of valence electrons in the molecule (or ion or free radical) must be the sum of all outer-shell electrons “contributed” to the molecule by each atom plus the negative charge or minus the positive charge, for the case of ions. Thus, for H2S04, there are 2 (one for each hydrogen) + 6 (for the sulfur) + 24 (6 for each oxygen) = 32; while for S04_, the number is also 32, since each atom “contributes” 6 plus 2 for the negative charge.
2. Once the number of valence electrons has been ascertained, it is necessary to determine which of them are found in covalent bonds and which are unshared. Unshared electrons (either a single electron or a pair) form part of the outer shell of just one atom, but electrons in a covalent bond are part of the outer shell of both atoms of the bond. First-row atoms (B, C, N, O, F) can have a maximum of eight valence electrons , and usually have this number, although some cases are known where a first-row atom has only six or seven.
18See Robinson, J.W., Practical Handbook of Spectroscopy, CRC Press, Boca Raton, FL, 1991, p. 178. 19Novak, I.; Potts, A.W. Tetrahedron 1997, 53, 14713.
20It has been argued that although the Lewis picture of two electrons making up a covalent bond may work well for organic compounds, it cannot be successfully applied to the majority of inorganic compounds: Jprgensen, C.K. Top. Curr. Chem. 1984, 124, 1.
CHAPTER 1
ELECTRONIC STRUCTURES OF MOLECULES 1 5
Where there is a choice between a structure that has six or seven electrons around a first-row atom and one in which all such atoms have an octet, it is the latter that generally has the lower energy and that consequently exists. For example, ethylene is
C=C
and not
H H
® c— C:0 \
H
or
H
H H
•c-c*
/ \
H H
There are a few exceptions. In the case of the molecule 02, the structure :6-o: has a lower energy than :o = o:- Although first-row atoms are limited to 8 valence electrons, this is not so for second-row atoms, which can accommodate 10 or even 12 because they can use their empty d orbitals for this purpose.21 For example, PC15 and SF6 are stable compounds. In SF6, one s and one p electron from the ground state 3s23p4 of the sulfur are promoted to empty d orbitals, and the six orbitals hybridize to give six sp3d2 orbitals, which point to the corners of a regular octahedron.
3. It is customary to show the formal charge on each atom. For this purpose, an atom is considered to “own” all unshared electrons, but only one -half of the electrons in covalent bonds. The sum of electrons that thus “belong” to an atom is compared with the number “contributed” by the atom. An excess belonging to the atom results in a negative charge, and a deficiency results in a positive charge. The total of the formal charges on all atoms equals the charge on the whole molecule or ion. Note that the counting procedure is not the same for determining formal charge as for determining the number of valence electrons. For both purposes, an atom “owns” all unshared electrons, but for outer- shell purposes it “owns” both the electrons of the covalent bond, while for formal-charge purposes it “owns” only one-half of these electrons.
Examples of electronic structures are (as mentioned in Ref. 4, an electron pair, whether unshared or in a bond, is represented by a straight line):
H-O
•O-H
H
H-C*
H
ch3 + l - h3c — n-o:
ch3
h3c
h3c
F-B
ch3
A coordinate-covalent bond, represented by an arrow, is one in which both elec¬ trons come from the same atom; that is, the bond can be regarded as being formed by the overlap of an orbital containing two electrons with an empty one. Thus tri- methylamine oxide would be represented
CH3
H3c— N~-0:“
ch3
21For a review concerning sulfur compounds with a valence shell larger than eight, see Salmond, W.G. Q. Rev. Chem. Soc. 1968, 22, 235.
1 6 LOCALIZED CHEMICAL BONDING
For a coordinate-covalent bond the rule concerning formal charge is amended, so that both electrons count for the donor and neither for the recipient. Thus the nitro¬ gen and oxygen atoms of trimethylamine oxide bear no formal charges. However, it is apparent that the electronic picture is exactly the same as the picture of trimethy¬ lamine oxide given just above, and we have our choice of drawing an arrowhead or a charge separation. Some compounds, for example, amine oxides, must be drawn one way or the other. It seems simpler to use charge separation, since this spares us from having to consider as a “different” method of bonding a way that is really the same as ordinary covalent bonding once the bond has formed.
ELECTRONEGATIVITY
The electron cloud that bonds two atoms is not symmetrical (with respect to the plane that is the perpendicular bisector of the bond) except when the two atoms are the same and have the same substituents. The cloud is necessarily distorted toward one side of the bond or the other, depending on which atom (nucleus plus electrons) maintains the greater attraction for the cloud. This attraction is called electronegativity ;22 and it is greatest for atoms in the upper-right corner of the periodic table and lowest for atoms in the lower-left comer. Thus a bond between fluorine and chlorine is distorted so that there is a higher probability of finding the electrons near the fluorine than near the chlorine. This gives the fluorine a partial negative charge and the chlorine a partial positive charge.
A number of attempts have been made to set up quantitative tables of electro¬ negativity that indicate the direction and extent of electron-cloud distortion for a bond between any pair of atoms. The most popular of these scales, devised by Paul¬ ing, is based on bond energies (see p. 27) of diatomic molecules. It is rationalized that if the electron distribution were symmetrical in a molecule A— B, the bond energy would be the mean of the energies of A— A and B— B, since in these cases the cloud must be undistorted. If the actual bond energy of A— B is higher than this (and it usually is), it is the result of the partial charges, since the charges attract each other and make a stronger bond, which requires more energy to break. It is neces¬ sary to assign a value to one element arbitrarily (F = 4.0). Then the electronegativ¬ ity of another is obtained from the difference between the actual energy of A— B and the mean of A— A and B— B (this difference is called A) by the formula
x^-x‘=\fSi
where vA and xB are the electronegativities of the known and unknown atoms and 23.06 is an arbitrary constant. Part of the scale derived from this treatment is shown in Table 1.1.
22For a collection of articles on this topic, see Sen, K.D.; Jprgensen, C.K. Electronegativity (Vol. 6 of Structure and Bonding)-, Springer: NY, 1987. For a review, see Batsanov, S.S. Russ. Chem. Rev. 1968, 37, 332.
CHAPTER 1
ELECTRONEGATIVITY 1 7
TABLE 1.1. Electronegativities of Some Atoms on the Pauling23 and Sanderson24 Scales
Element
Pauling
Sanderson
Element
Pauling
Sanderson
F
4.0
4.000
H
2.1
2.592
O
3.5
3.654
P
2.1
2.515
Cl
3.0
3.475
B
2.0
2.275
N
3.0
3.194
Si
1.8
2.138
Br
2.8
3.219
Mg
1.2
1.318
S
2.5
2.957
Na
0.9
0.835
I
2.5
2.778
Cs
0.7
0.220
C
2.5
2.746
Other treatments25 have led to scales that are based on different principles, for example, the average of the ionization potential and the electron affinity,26 the aver¬ age one-electron energy of valence- shell electrons in ground- state free atoms,27 or the “compactness” of an atom’s electron cloud.24 In some of these treatments electronegativities can be calculated for different valence states, for different hybri¬ dizations (e.g., sp carbon atoms are more electronegative than sp2, which are still more electronegative than s/?3),28and even differently for primary, secondary, and tertiary carbon atoms. Also, electronegativities can be calculated for groups rather than atoms (Table 1.2).29
Electronegativity information can be obtained from NMR spectra. In the absence of a magnetically anisotropic group30the chemical shift of a XH or a 13C nucleus is approximately proportional to the electron density around it and hence to the elec¬ tronegativity of the atom or group to which it is attached. The greater the electro¬ negativity of the atom or group, the lower the electron density around the proton, and the further downfield the chemical shift. An example of the use of this correla¬ tion is found in the variation of chemical shift of the ring protons in the series
23Taken from Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, NY, p. 93, except for the value for Na, which is from Sanderson, R.T. J. Am. Chem. Soc. 1983, 105, 2259; J. Chem. Educ. 1988, 65, 112, 223.
24See Sanderson, R.T. J. Am. Chem. Soc. 1983, 105, 2259; J. Chem. Educ. 1988, 65, 112, 223.
25For several sets of electronegativity values, see Huheey, J.E. Inorganic Chemistry, 3rd ed., Harper and Row: NY, 1983, pp. 146-148; Mullay, J., in Sen, K.D.; Jprgensen, C.K. Electronegativity (Vol. 6 of Structure and Bonding ), Springer, NY, 1987, p. 9.
26Mulliken, R.S. J. Chem. Phys. 1934, 2, 782; Iczkowski, R.P; Margrave, J.L. J. Am. Chem. Soc. 1961, 83, 3547; Hinze, J.; Jaffe, H.H. J. Am. Chem. Soc. 1962, 84, 540; Rienstra-Kiracofe, J.C.; Tschumper, G.S.; Schaefer III, H.F.; Nandi, S.; Ellison, G.B. Chem. Rev. 2002, 102, 231.
27 Allen, L.C. J. Am. Chem. Soc. 1989, 111, 9003.
28Walsh, A.D. Discuss. Faraday Soc. 1947, 2, 18; Bergmann, D.; Hinze, J., in Sen, K.D.; Jprgensen, C.K. Electronegativity (Vol. 6 of Structure and Bonding), Springer, NY, 1987, pp. 146-190.
29Inamoto, N.; Masuda, S. Chem. Lett. 1982, 1003. For a review of group electronegativities, see Wells, PR. Prog. Phys. Org. Chem. 1968, 6, 111. See also Bratsch, S.G. J. Chem. Educ., 1988, 65, 223; Mullay, J. J. Am. Chem. Soc. 1985, 107, 7271; Zefirov, N.S.; Kirpichenok, M.A.; Izmailov, F.F.; Trofimov, M.I. Dokl. Chem. 1987, 296, 440; Boyd, R.J.; Edgecombe, K.E. J. Am. Chem. Soc. 1988, 110, 4182.
30A magnetically anisotropic group is one that is not equally magnetized along all three axes. The most common such groups are benzene rings (see p. 55) and triple bonds.
1 8 LOCALIZED CHEMICAL BONDING
TABLE 1.2. Some Group Electronegativites Relative to H = 2.176.29
ch3
2.472
cci3
2.666
CH3CH2
2.482
C6H5
2.717
CH2C1
2.538
cf3
2.985
CBr3
2.561
C=N
3.208
CHC12
2.602
no2
3.421
toluene, ethylbenzene, isopropylbenzene, terf-butylbenzene (there is a magnetically anisotropic group here, but its effect should be constant throughout the series). It is found that the electron density surrounding the ring protons decreases31 in the order given.32 However, this type of correlation is by no means perfect, since all the mea¬ surements are being made in a powerful field, which itself may affect the electron density distribution. Coupling constants between the two protons of a system -chch-x have atso been found to depend on the electronegativity of X.33
When the difference in electronegativities is great, the orbital may be so far over to one side that it barely covers the other nucleus. This is an ionic bond , which is seen to arise naturally out of the previous discussion, leaving us with basically only one type of bond in organic molecules. Most bonds can be considered intermediate between ionic and covalent. We speak of percent ionic character of a bond, which indicates the extent of electron-cloud distortion. There is a continuous gradation from ionic to covalent bonds.
DIPOLE MOMENT
The dipole moment is a property of the molecule that results from charge separations like those discussed above. However, it is not possible to measure the dipole moment of an individual bond within a molecule; we can measure only the total moment of the molecule, which is the vectorial sum of the individual bond moments.34 These individual moments are roughly the same from molecule to molecule,35 but this con¬ stancy is by no means universal. Thus, from the dipole moments of toluene and nitro¬ benzene (Fig. 1.10)36 we should expect the moment of p-nitrotoluene to be ~4.36D.
31This order is opposite to that expected from the field effect (p. 19). It is an example of the Baker-Nathan order (p. 96).
32Moodie, R.B.; Connor, T.M.; Stewart, R. Can. J. Chem. 1960, 38, 626.
33Williamson, K.L. J. Am. Chem. Soc. 1963, 85, 516; Laszlo, P.; Schleyer, P.v.R. J. Am. Chem. Soc. 1963, 85, 2709; Niwa, J. Bull. Chem. Soc. Jpn. 1967, 40, 2192.
34For methods of determining dipole moments and discussions of their applications, see Exner, O. Dipole Moments in Organic Chemistry, Georg Thieme Publishers: Stuttgart, 1975. For tables of dipole moments, see McClellan, A.L. Tables of Experimental Dipole Moments, Vol. 1 ; W.H. Freeman: San Francisco, 1963’, Vol. 2, Rahara Enterprises: El Cerrito, CA, 1974.
35For example, see Koudelka, J.; Exner, O. Collect. Czech. Chem. Commun. 1985, 50, 188, 200.
36The values for toluene, nitrobenzene, and p-nitro toluene are from MacClellan, A.L., Tables of Experimental Dipole Moments, Vol. 1, W.H. Freeman, San Francisco, 1963’, Vol. 2, Rahara Enterprises, El Cerrito, CA, 1974. The values for phenol and p-cresol were determined by Goode, E.V.; Ibbitson, D.A. J. Chem. Soc. 1960, 4265.
CHAPTER 1
INDUCTIVE AND FIELD EFFECTS 1 9
Fig. 1.10. Some dipole moments, in debye units, measured in benzene. In the 3D model, the arrow indicates the direction of the dipole moment for the molecule, pointing to the negative part of the molecule.36
The actual value 4.39 D is reasonable. However, the moment of 77-cresol (1.57 D) is quite far from the predicted value of 1.11 D. In some cases, molecules may have sub¬ stantial individual bond moments but no total moments at all because the individual moments are canceled out by the overall symmetry of the molecule. Some examples are CC14, trans- 1 ,2-dibromoethene, and p-dinitrobenzene.
Because of the small difference between the electronegativities of carbon and hydrogen, alkanes have very small dipole moments, so small that they are difficult to measure. For example, the dipole moment of isobutane is 0.132 D37 and that of propane is 0.085 D.38 Of course, methane and ethane, because of their symmetry, have no dipole moments.39 Few organic molecules have dipole moments >7 D.
INDUCTIVE AND FIELD EFFECTS
The C— C bond in ethane has no polarity because it connects two equivalent atoms. However, the C— C bond in chloroethane is polarized by the presence of the electrone¬ gative chlorine atom. This polarization is actually the sum of two effects. In the first of these, the C-l atom, having been deprived of some of its electron density by the
S+ 5+ 5-
1CH3 — 2CH2 > » Cl
37Maryott, A.A.; Birnbaum, G. J. Chem. Phys. 1956, 24, 1022; Lide Jr., D.R.; Mann, D.E. J. Chem. Phys. 1958, 29, 914.
38Muenter, J.S.; Laurie, V.W. J. Chem. Phys. 1966, 45, 855.
39Actually, symmetrical tetrahedral molecules like methane do have extremely small dipole moments, caused by centrifugal distortion effects; these moments are so small that they can be ignored for all practical purposes. For CH4 p is ~ 5.4 x 10-6 D: Ozier, I. Phys. Rev. Lett. 1971, 27, 1329; Rosenberg, A.; Ozier, I.; Kudian, A.K. J. Chem. Phys. 1972, 57, 568.
20 LOCALIZED CHEMICAL BONDING
greater electronegativity of Cl, is partially compensated by drawing the C— C electrons closer to itself, resulting in a polarization of this bond and a slightly positive charge on the C-2 atom. This polarization of one bond caused by the polarization of an adjacent bond is called the inductive effect. The effect is greatest for adjacent bonds but may also be felt farther away; thus the polarization of the C— C bond causes a (slight) polarization of the three methyl C— H bonds. The other effect operates not through bonds, but directly through space or solvent molecules, and is called the field effect.40 It is often very difficult to separate the two kinds of effect, but it has been done in a number of cases, generally by taking advantage of the fact that the field effect depends on the geometry of the molecule but the inductive effect depends only on the nature of the bonds. For example, in isomers 1 and 241 the inductive effect of the chlo¬ rine atoms on the position of the electrons in the COOH group (and hence on the
1
p Ka = 6.07
2
pKa = 5.67
acidity, see Chapter 8) should be the same since the same bonds intervene; but the field effect is different because the chlorines are closer in space to the COOH in 1 than they are in 2. Thus a comparison of the acidity of 1 and 2 should reveal whether a field effect is truly operating. The evidence obtained from such experiments is overwhelm¬ ing that field effects are much more important than inductive effects.42 In most cases, the two types of effect are considered together; in this book, we will not attempt to separate them, but will use the name field effect to refer to their combined action 43 Functional groups can be classified as electron- withdrawing (— I) or electron- donating (+/) groups relative to hydrogen. This means, for example, that N02, a —I group, will draw electrons to itself more than a hydrogen atom would if it
40Roberts, J.D.; Moreland, Jr., W.T. J. Am. Chem. Soc. 1953, 75, 2167.
41This example is from Grubbs, E.J.; Fitzgerald, R.; Phillips, R.E.; Petty, R. Tetrahedron 1971, 27, 935. 42For example, see Dewar, M.J.S.; Grisdale, P.J. J. Am. Chem. Soc. 1962, 84, 3548; Stock, L.M. J. Chem. Educ., 1972, 49, 400; Golden, R.; Stock, L.M. J. Am. Chem. Soc. 1972, 94, 3080; Liotta, C.; Fisher, W.F.; Greene Jr., G.H.; Joyner, B.L. J. Am. Chem. Soc. 1972, 94, 4891; Wilcox, C.F.; Leung, C. J. Am. Chem. Soc. 1968, 90, 336; Butler, A.R. J. Chem. Soc. B 1970, 867; Rees, J.H.; Ridd, J.H.; Ricci, A. J. Chem. Soc. Perkin Trans. 2 1976, 294; Topsom, R.D. J. Am. Chem. Soc. 1981, 103, 39; Grob, C.A.; Kaiser, A.; Schweizer, T. Helv. Chim. Acta 1977, 60, 391; Reynolds, W.F. J. Chem. Soc. Perkin Trans. 2 1980 , 985, Prog. Phys. Org. Chem. 1983, 14, 165-203; Adcock, W.; Butt, G.; Kok, G.B.; Marriott, S.; Topsom, R.D. J. Org. Chem. 1985, 50, 2551; Schneider, H.; Becker, N. J. Phys. Org. Chem. 1989, 2, 214; Bowden, K.; Ghadir, K.D.F. J. Chem. Soc. Perkin Trans. 2 1990, 1333. Inductive effects may be important in certain systems. See, for example, Exner, O.; Fiedler, P. Collect. Czech. Chem. Commun. 1980, 45, 1251; Li, Y.; Schuster, G.B. J. Org. Chem. 1987, 52, 3975.
43There has been some question as to whether it is even meaningful to maintain the distinction between the two types of effect: see Grob, C.A. Helv. Chim. Acta 1985, 68, 882; Lenoir, D.; Frank, R.M. Chem. Ber. 1985, 118, 753; Sacher, E. Tetrahedron Lett. 1986, 27, 4683.
CHAPTER 1
INDUCTIVE AND FIELD EFFECTS 21
TABLE 1.3. Field Effects of Various Groups Relative to Hydrogena
+/
-/
cr
nr3+
COOH
OR
COO"
SR+
F
COR
cr3
NH+
Cl
SH
chr2
no2
Br
SR
ch2r
so2r
I
OH
ch3
CN
OAr
Pi
U
III
u
D
S02Ar
COOR
Ar
c=cr2
"The groups are listed approximately in order of decreasing strength for both —I and +/ groups.
occupied the same position in the molecule.
o2n * — ch2 ^ — Ph
H - CH2 - Ph
Thus, in a-nitrotoluene, the electrons in the N— C bond are farther away from the car¬ bon atom than the electrons in the H— C bond of toluene. Similarly, the electrons of the C— Ph bond are farther away from the ring in a-nitrotoluene than they are in toluene. Field effects are always comparison effects. We compare the —I or +1 effect of one group with another (usually hydrogen). It is commonly said that, compared with hydro¬ gen, the N02 group is electron- withdrawing and the O- group electron-donating or elec¬ tron releasing. However, there is no actual donation or withdrawal of electrons, though these terms are convenient to use; there is merely a difference in the position of electrons due to the difference in electronegativity between H and N02 or between H and 0".
Table 1.3 lists a number of the most common —I and +/ groups.44 It can be seen that compared with hydrogen, most groups are electron withdrawing. The only elec- trondonating groups are groups with a formal negative charge (but not even all these), atoms of low electronegativity (Si 45 Mg, etc., and perhaps alkyl groups). Alkyl groups46 were formerly regarded as electron donating, but many examples of behavior have been found that can be interpreted only by the conclusion that alkyl groups are electron withdrawing compared with hydrogen.47 In accord with this is the value of 2.472 for the group electronegativity of CH3 (Table 1.2) compared with 2.176 for H. We will see that when an alkyl group is attached to an unsaturated or trivalent carbon (or other atom), its behavior is best explained by assuming it is +/ (see, e.g., pp. 239, 251, 388, 669), but when it is connected to a saturated atom, the results are not as clear,
44See also Ceppi, E.; Eckhardt, W.; Grob, C.A. Tetrahedron Lett. 1973, 3627 .
45For a review of field and other effects of silicon-containing groups, see Bassindale, A.R.; Taylor. P.G., in Patai, S.; Rappoport, Z. The Chemistry of Organic Silicon Compounds, pt. 2, Wiley, NY, 1989, pp. 893-963. 46For a review of the field effects of alkyl groups, see Levitt, L.S.; Widing, H.F. Prog. Phys. Org. Chem. 1976, 12, 119.
47See Sebastian, J.F. J. Chem. Educ. 1971, 48, 97.
22 LOCALIZED CHEMICAL BONDING
and alkyl groups seem to be +/ in some cases and —I in others48 (see also p. 391). Similarly, it is clear that the field-effect order of alkyl groups attached to unsaturated systems is tertiary > secondary > primary > CH3, but this order is not always main¬ tained when the groups are attached to saturated systems. Deuterium is electron- donating with respect to hydrogen.49 Other things being equal, atoms with sp bonding generally have a greater electron-withdrawing power than those with sp 2 bonding, which in turn have more electron-withdrawing power than those with sp3 bonding.50 This accounts for the fact that aryl, vinylic, and alkynyl groups are —I. Field effects always decrease with increasing distance, and in most cases (except when a very powerful +/ or — I group is involved), cause very little difference in a bond four bonds away or more. There is evidence that field effects can be affected by the solvent.51
For discussions of field effects on acid and base strength and on reactivity, see Chapters 8 and 9, respectively.
BOND DISTANCES52
The distances between atoms in a molecule are characteristic properties of the mole¬ cule and can give us information if we compare the same bond in different molecules. The chief methods of determining bond distances and angles are X-ray diffraction (only for solids), electron diffraction (only for gases), and spectroscopic methods, especially microwave spectroscopy. The distance between the atoms of a bond is not constant, since the molecule is always vibrating; the measurements obtained are therefore average values, so that different methods give different results.53 However, this must be taken into account only when fine distinctions are made.
Measurements vary in accuracy, but indications are that similar bonds have fairly constant lengths from one molecule to the next, though exceptions are known.54 The variation is generally less than 1%. Table 1.4 shows
48See, for example, Schleyer, P. von.R.; Woodworth, C.W. J. Am. Chem. Soc. 1968, 90, 6528; Wahl Jr., G.H.; Peterson Jr., M.R. J. Am. Chem. Soc. 1970, 92, 7238. The situation may be even more complicated. See, for example, Minot, C.; Eisenstein, O.; Hiberty, PC.; Anh, N.T. Bull. Soc. Chim. Fr. 1980, 11-119. 49Streitwieser Jr., A.; Klein, H.S. J. Am. Chem. Soc. 1963, 85, 2759.
50Bent, H.A. Chem. Rev. 1961 , 61, 275, p. 281.
51See Laurence, C.; Berthelot, M.; Lucon, M.; Helbert, M.; Morris, D.G.; Gal, J. J. Chem. Soc. Perkin Trans. 2 1984, 705.
52For tables of bond distances and angles, see Allen, F.H.; Kennard, O.; Watson, D.G.; Brammer, L.; Orpen, A.G.; Taylor, R. J. Chem. Soc. Perkin Trans. 2 1987, S1-S19 (follows p. 1914); Tables of Interatomic Distances and Configurations in Molecules and Ions Chem. Soc. Spec. Publ. No. 11, 1958’, Interatomic Distances Supplement Chem. Soc. Spec. Publ. No. 18, 1965’, Harmony, M.D. Faurie, V.W.; Kuczkowski, R.F.; Schwendeman, R.H.; Ramsay, D.A.; Lovas, F.J.; Lafferty, W.J.; Maki, A.G. J. Phys. Chem. Ref. Data 1979, 8, 619-721. For a review of molecular shapes and energies for many small organic molecules, radicals, and cations calculated by molecular-orbital methods, see Lathan, W.A.; Curtiss, L.A.; Hehre, W.J.; Lisle, J.B.; Pople, J.A. Prog. Phys. Org. Chem. 1974, 11, 175. For a discussion of substituent effects on bond distances, see Topsom, R.D. Prog. Phys. Org. Chem. 1987, 16, 85.
53Burkert, U.; Allinger, N.L. Molecular Mechanics’, ACS Monograph 111, American Chemical Society, Washington, 1982, pp. 6-9; Whiffen, D.H. Chem. Ber. 1971, 7, 57-61; Stals, J. Rev. Pure Appl. Chem. 1970, 20, 1, pp. 2-5.
54Schleyer, P.v.R.; Bremer, M. Angew. Chem. Int. Ed. 1989, 28, 1226.
CHAPTER 1
BOND DISTANCES 23
TABLE 1.4. Bond Lengths between sp3 Carbons in Some Compounds
C— C bond in
Reference
Bond length, A
Diamond
55
1.544
c2h6
56
1.5324 ±0.0011
c2h5ci
57
1.5495 ±0.0005
C3H8
58
1.532 ±0.003
Cyclohexane
59
1.540 ±0.015
tert-Buty\ chloride
60
1.532
w-Butane to n -heptane
61
1.531 - 1.534
Isobutane
62
1.535 ±0.001
distances for single bonds between two sp 3 carbons. However, an analysis of C— OR bond distances in >2000 ethers and carboxylic esters (all with sp 3 carbon) shows that this distance increases with increasing electron withdrawal in the R group and as the C changes from primary to secondary to tertiary.63 For these compounds, mean bond lengths of the various types ranged from 1.418 to 1.475 A. Certain sub¬ stituents can also influence bond length. The presence of a silyl substituent (3- to a C— O (ester) linkage can lengthen the C— O, thereby weakening it.64 This is believed to result from a- a interactions in which the C— Si a -bonding orbital acts as the donor and the C— O a* orbitals acts as the receptor.
Although a typical carbon— carbon single bond has a bond length of ~ 1.54 A, certain molecules are known that have significantly longer bond lengths.65 Calculations
55Lonsdale, K. Phil. Trans. R. Soc. London 1947 , A240, 219.
56Bartell, L.S.; Higginbotham, H.K. J. Chem. Phys. 1965, 42, 851.
57Wagner, R.S.; Dailey, B.P. J. Chem. Phys. 1957, 26, 1588.
58Iijima, T. Bull. Chem. Soc. Jpn. 1972, 45, 1291.
59Tables of Interatomic Distances, Ref. 52.
60Momany, F.A.; Bonham, R.A.; Druelinger, M.L. J. Am. Chem. Soc. 1963, 85, 3075; also see, Lide, Jr., D.R.; Jen, M. J. Chem. Phys. 1963, 38, 1504.
61Bonham, R.A.; Bartell, L.S.; Kohl, D.A. J. Am. Chem. Soc. 1959, 81, 4765.
62Hilderbrandt, R.L.; Wieser, J.D. J. Mol. Struct. 1973, 15, 27.
63 Allen, F.H.; Kirby, A.J. J. Am. Chem. Soc. 1984, 106, 6197; Jones, P.G.; Kirby, A.J. J. Am. Chem. Soc. 1984, 106, 6207.
64White, J.M.; Robertson, G.B. J. Org. Chem. 1992, 57, 4638.
65Kaupp, G.; Boy, J Angew. Chem. Int. Ed. 1997, 36, 48.
24 LOCALIZED CHEMICAL BONDING
have been done for unstable molecules that showed them to have long bond lengths, and an analysis of the X-ray structure for the photoisomer of [2.2]-tetraben- zoparacyclophane (see Chapter 2) showed a C— C bond length of 1.77 A.66,65 Long bond lengths have been observed in stable molecules such as benzocyclobutane derivatives.67 A bond length of 1.729 A was reliably measured in l,l-di-terf-butyl-2, 2-diphenyl-3,8-dichlorocyclobutan[&] naphthalene, 3. 68 X-ray analysis of several of these derivations confirmed the presence of long C— C bonds, with 4 having a con¬ firmed bond length of 1.734 A.69
Bond distances for some important bond types are given in Table 1.5. 70 As can be seen in this table, carbon bonds are shortened by increasing s character.
TABLE 1.5. Bond distances"
Bond Type
Length, A
Typical Compounds
C-C
3 3
sp -sp
1.53
3 2
sp -sp
1.51
Acetaldehyde, toluene, propene
sp3- sp
1.47
Acetonitrile, propyne
2 2 sp -sp
1.48
Butadiene, glyoxal, biphenyl
sp2 - sp
1.43
Acrylonitrile, vinylacetylene
sp-sp
1.38
Cyanoacetylene, butadiyne
C=C
2 2
sp -sp
1.32
Ethylene
sp2-sp
1.31
Ketene, allenes
sp-sp 71
1.28
Butatriene, carbon suboxide
C=C72
sp-sp
1.18
Acetylene
C-H73
sp3- H
1.09
Methane
sp2- H
1.08
Benzene, ethylene
sp-Yl14
1.08
HCN, acetylene
66Ehrenberg, M. Acta Crystallogr. 1966, 20, 182.
67Toda, F.; Tanaka, K.; Stein, Z.; Goldberg, I Acta Crystallogr., Sect. C 1996, 52, 177.
68Toda, F.; Tanaka, K.; Watanabe, M.; Taura, K.; Miyahara, I.; Nakai, T.; Hirotsu, K. J. Org. Chem. 1999, 64, 3102.
69Tanaka, K.; Takamoto, N.; Tezuka, Y.; Kato, M.; Toda, F. Tetrahedron 2001, 57, 3761.
70Except where noted, values are from Allen, F.H.; Kennard, O.; Watson, D.G.; Brammer, L.; Orpen, A.G.; Taylor, R. J. Chem. Soc. Perkin Trans. 2 1987, S1-S19 (follows p. 1914). In this source, values are given to three significant figures.
71Costain, C.C.; Stoicheff, B.P. J. Chem. Phys. 1959, 30, 111.
72For a full discussion of alkyne bond distances, see Simonetta, M.; Gavezzotti, A, in Patai, S. The Chemistry of the Carbon-Carbon Triple Bond, Wiley, NY, 1978.
73For an accurate method of C— H bond distance determination, see Henry, B.R. Acc. Chem. Res. 1987, 20, 429.
74Bartell, L.S.; Roth, E.A.; Hollowell, C.D.; Kuchitsu, K.; Young, Jr., J.E. J. Chem. Phys. 1965, 42, 2683.
CHAPTER 1
BOND ANGLES 25
TABLE 1.5. (< continued )
Bond Type
Length, A
Typical Compounds
C-O
sp3- O
1.43
Dimethyl ether, ethanol
sp2- O
1.34
Formic acid
o
II
u
sp2- O
1.21
Formaldehyde, formic acid
sp- O59
1.16
co2
C-N
sp3- N
1.47
Methylamine
sp2- N
1.38
Formamide
C=N
sp2- N
1.28
Oximes, imines
C=N
sp- N
1.14
HCN
C-S
sp3- S
1.82
Methanethiol
sp2- S
1.75
Diphenyl sulfide
sp- S
1.68
CH3SCN
c=s
sp- S
1.67
CS2
C— halogen75
F
Cl
Br I
v/?3-halogen
1.40
1.79
1.97 2.16
v/?2-halogen
1.34
1.73
1.88 2.10
vp-halogen
1.2776
1.63
1.7977 1.9977
"The values given are average lengths and do not necessarily apply exactly to the compounds mentioned.70
This is most often explained by the fact that, as the percentage of s character in a hybrid orbital increases, the orbital becomes more like an s orbital and hence is held more tightly by the nucleus than an orbital with less s character. However, other explanations have also been offered (see p. 39), and the matter is not completely settled.
Indications are that a C— D bond is slightly shorter than a corresponding C— H bond. Thus, electron-diffraction measurements of C2H6 and C2D6 showed a C— H bond dis- tance of 1.1122 ± 0.0012 A and a C-D distance of 1.1071 ± 0.0012 A.56
BOND ANGLES
It might be expected that the bond angles of sp3 carbon would always be the tetra¬ hedral angle 109°28', but this is so only where the four groups are identical, as in
75For reviews of carbon-halogen bonds, see Trotter, J., in Patai, S. The Chemistry of the Carbon-Halogen Bond, pt. 1, Wiley, NY, 1973, pp. 49-62; Mikhailov, B.M. Russ. Chem. Rev. 1971, 40, 983.
76Lide, Jr., D.R. Tetrahedron 1962, 17, 125.
77Rajput, A.S.; Chandra, S. Bull Chem. Soc. Jpn. 1966, 39, 1854.
26 LOCALIZED CHEMICAL BONDING
methane, neopentane, or carbon tetrachloride. In most cases, the angles deviate a little from the pure tetrahedral value. For example, the C— C— Br angle in 2-bromo- propane is 114. 2°. 78 Similarly, slight variations are generally found from the ideal values of 120 and 180° for sp2 and sp carbon, respectively. These deviations occur because of slightly different hybridizations, that is, a carbon bonded to four other atoms hybridizes one s and three p orbitals, but the four hybrid orbitals thus formed are generally not exactly equivalent, nor does each contain exactly 25 % s and 75% p character. Because the four atoms have (in the most general case) different elec¬ tronegativities, each makes its own demand for electrons from the carbon atom.79 The carbon atom supplies more p character when it is bonded to more electronega¬ tive atoms, so that in chloromethane, for example, the bond to chlorine has some¬ what more than 75% p character, which of course requires that the other three bonds have somewhat less, since there are only three p orbitals (and one s) to be divided among the four hybrid orbitals.80 Of course, in strained molecules, the bond angles may be greatly distorted from the ideal values (see p. 216).
For oxygen and nitrogen, angles of 90° are predicted from p2 bonding. However, as we have seen (p. 6), the angles of water and ammonia are much larger than this, as are the angles of other oxygen and nitrogen compounds (Table 1.6); in fact, they are much closer to the tetrahedral angle of 109° 28' than to 90°. These facts have
TABLE 1.6. Oxygen, Sulfur, and Nitrogen Bond Angles in Some Compounds
Angle
Value
Compound
Reference
H-O-H
104°27'
Water
5
C-O-H
107-109°
Methanol
59
C-O-C
m°43'
Dimethyl ether
81
C-O-C
124° ± 5°
Diphenyl ether
82
H-S-H
92.1°
H2S
82
C-S-H
99.4°
Methanethiol
82
C-S-C
99.1°
Dimethyl sulfide
83
H-N-H
106°46'
Ammonia
5
H-N-H
106°
Methylamine
84
C-N-H
112°
Methylamine
83
C-N-C
108.7°
Trimethylamine
85
78Schwendeman, R.H.; Tobiason, F.L. J. Chem. Phys. 1965, 43, 201.
79For a review of this concept, see Bingel, W.A.; Ltittke, W. Angew. Chem. Int. Ed. 1981, 20, 899. 80This assumption has been challenged: see Pomerantz, M.; Liebman, J.F. Tetrahedron Lett. 1975, 2385. 81Blukis, V.; Kasai, P.H.; Myers, RJ. J. Chem. Phys. 1963, 38, 2753.
82 Abrahams, S.C. Q. Rev. Chem. Soc. 1956, 10, 407.
83Iijima, T.; Tsuchiya, S.; Kimura, M. Bull. Chem. Soc. Jpn. 1977, 50, 2564.
84Lide, Jr., D.R. J. Chem. Phys. 1957, 27, 343.
85Lide, Jr., D.R.; Mann, D.E. J. Chem. Phys. 1958, 28, 572.
CHAPTER 1
BOND ENERGIES 27
led to the suggestion that in these compounds oxygen and nitrogen use sp 3 bonding, that is, instead of forming bonds by the overlap of two (or three) p orbitals with Is orbi¬ tals of the hydrogen atoms, they hybridize their 2s and 2 p orbitals to form four sp3 orbi¬ tals and then use only two (or three) of these for bonding with hydrogen, the others remaining occupied by unshared pairs (also called lone pairs). If this description is valid, and it is generally accepted by most chemists today,86 it becomes necessary to explain why the angles of these two compounds are in fact not 109° 28' but a few degrees smaller. One explanation that has been offered is that the unshared pair actually has a greater steric requirement than a pair in a bond, since there is no second nucleus to draw away some of the electron density and the bonds are thus crowded together. How¬ ever, most evidence is that unshared pairs have smaller steric requirements than bonds87 and the explanation most commonly accepted is that the hybridization is not pure sp3. As we have seen above, an atom supplies more p character when it is bonded to more electronegative atoms. An unshared pair may be considered to be an “atom” of the lowest possible electronegativity, since there is no attracting power at all. Conse¬ quently, the unshared pairs have more s and the bonds more p character than pure sp3 orbitals, making the bonds somewhat more like p2 bonds and reducing the angle. As seen in Table 1.6, oxygen, nitrogen, and sulfur angles generally increase with decreas¬ ing electronegativity of the substituents. Note that the explanation given above cannot explain why some of these angles are greater than the tetrahedral angle.
BOND ENERGIES8889
There are two kinds of bond energy. The energy necessary to cleave a bond to give the constituent radicals is called the dissociation energy D. For example, D for H20 — > HO + H is 118 kcal mol-1 (494/mol). However, this is not taken as the energy of the O— H bond in water, since D for H— O — > H + O is 100 kcal mol-1 (418 kJ mol-1). The average of these two values, 109 kcal mol-1 (456 kJ mol-1), is taken as the bond energy E. In diatomic molecules, of course, D = E.
86An older theory holds that the bonding is indeed p2, and that the increased angles come from repulsion of the hydrogen or carbon atoms. See Laing, M., J. Chem. Educ. 1987, 64, 124.
87See, for example, Pumphrey, N.W.J.; Robinson, M.J.T. Chem. Ind. ( London ) 1963, 1903; Allinger, N.L.; Carpenter, J.G.D.; Karkowski, F.M. Tetrahedron Lett. 1964, 3345; Jones, R.A.Y.; Katritzky, A.R.; Richards, A.C.; Wyatt, R.J.; Bishop, R.J.; Sutton, L.E. J. Chem. Soc. B 1970, 127; Blackburne, I.D.; Katritzky, A.R.; Takeuchi, Y. J. Am. Chem. Soc. 1974, 96, 682; Acc. Chem. Res. 1975, 8, 300; Aaron, H.S.; Ferguson, C.P. J. Am. Chem. Soc. 1976, 98, 7013; Anet, F.A.L.; Yavari, I. J. Am. Chem. Soc. 1977, 99, 2794; Vierhapper, F.W.; Eliel, E.L. J. Org. Chem. 1979, 44, 1081; Gust, D.; Fagan, M.W. J. Org. Chem. 1980, 45, 2511. For other views, see Lambert, J.B.; Featherman, S.I. Chem. Rev. 1975, 75, 611; Crowley, P.J.; Morris, G.A.; Robinson, M.J.T. Tetrahedron Lett. 1976, 3575; Breuker, K.; Kos, N.J.; van der Plas, H.C.; van Veldhuizen, B. J. Org. Chem. 1982, 47, 963.
88Blanksby, S.J.; Ellison, G.B. Acc. Chem. Res. 2003, 36, 255.
89For reviews including methods of determination, see Wayner, D.D.M.; Griller, D. Adv. Free Radical Chem. ( Greenwich , Conn.) 1990, 1, 159; Kerr, J.A. Chem. Rev. 1966, 66, 465; Benson, S.W. J. Chem. Educ. 1965, 42, 520; Wiberg, K.B., in Nachod, F.C.; Zuckerman, J.J. Determination of Organic Structures by Physical Methods, Vol. 3, Academic Press, NY, 1971, pp. 207-245.
28 LOCALIZED CHEMICAL BONDING
kcal
kJ
+ 3.5 02
- 2 C02 (gas)
+ 3 H20 (liq)
+372.9
+1560
2 C02 (gas)
— 9 C
^ ^(graphite)
+ 2 02 (gas)
-188.2
-204.9
-787
-857
3 H20(liq)
= 3 H2 (gas)
+ 1.5 02 (gas)
-312/5
-1308
3 H2 (gas)
- 6 H (gas)
-343.4
-1437
2 C
^(graphite)
- 2 C (gas)
^2^6 (gas)
- 6 H (gas)
+ 2 C (gas)
-676.1 kcal
-2829 kJ
Fig. 1.11. Calculation of the heat of atomization of ethane at 25°C.
The D values may be easy or difficult to measure, and they can be estimated by var¬ ious techniques.90 When properly applied, “Pauling’s original electronegativity equa¬ tion accurately describes homolytic bond dissociation enthalpies of common covalent bonds, including highly polar ones, with an average deviation of (1.5 kcal mol1 [^6.3 kJ mol1] from literature values).”91 Whether measured or calculated, there is no question as to what D values mean. With E values the matter is not so simple. For methane, the total energy of conversion from CH4 to C + 4H (at 0 K) is 393 kcal mol-1 (1644 kJ mol1).92 Consequently, E for the C— H bond in methane is 98 kcal mol-1 (41 1 kJ mol-1) at 0 K. The more usual practice, though, is not to measure the heat of ato¬ mization (i.e., the energy necessary to convert a compound to its atoms) directly but to calculate it from the heat of combustion. Such a calculation is shown in Figure 1.11.
Heats of combustion are very accurately known for hydrocarbons.93 For methane the value at 25°C is 212.8 kcal mol-1 (890.4 kJ mol-1), which leads to a heat of atomi¬ zation of 398.0 kcal mol-1 (1665 kJ mol-1) or a value of E for the C— H bond at 25°C of 99.5 kcal mol-1 (416 kJ mol-1). This method is fine for molecules like methane in which all the bonds are equivalent, but for more complicated molecules assumptions must be made. Thus for ethane, the heat of atomization at 25° C is 676.1 kcal mol-1 or 2829 kJ mol-1 (Fig. 1.11), and we must decide how much of this energy is due to the C— C bond and how much to the six C— H bonds. Any assumption must be artificial, since there is no way of actually obtaining this information, and indeed the question has no real meaning. If we make the assumption that E for each of the C— H bonds is the same as E for the C— H bond in methane (99.5 kcal mol-1 or 416 kJ mol-1), then 6x 99.5 (or 416) = 597.0 (or 2498), leaving 79.1 kcal mol-1 (331 kJ mol-1) for the C-C bond. However, a similar calculation for propane gives a value of 80.3 (or 336) for the
90Cohen, N.; Benson, S.W. Chem. Rev 1993, 93, 2419; Korth, H.-G.; Sicking, W. J. Chem. Soc. Perkin Trans. 2 1997, 715.
91Matsunaga, N.; Rogers, D.W.; Zavitsas, A.A. J. Org. Chem, 2003, 68, 3158.
92For the four steps, D values are 101 to 102, 88, 124, and 80 kcal mol-1 (423-427, 368, 519, and 335 kJ mol-1), respectively, though the middle values are much less reliable than the other two: Knox, B.E.; Palmer, H.B. Chem. Rev. 1961, 61, 247; Brewer, R.G.; Kester, F.L. J. Chem. Phys. 1964, 40, 812; Linevsky, M.J. J. Chem. Phys. 1967, 47, 3485.
93For values of heats of combustion of large numbers of organic compounds: hydrocarbons and others, see Cox, J.D.; Pilcher, G., Thermochemistry of Organic and Organometallic Compounds, Academic Press, NY, 1970; Domalski, E.S. J. Phys. Chem. Ref. Data 1972, 1, 221-277 . For large numbers of heats-of- formation values (from which heats of combustion are easily calculated) see Stull, D.R.; Westrum, Jr., E.F.; Sinke, G.C. The Chemical Thermodynamics of Organic Compounds, Wiley, NY, 1969.
CHAPTER 1
BOND ENERGIES 29
C— C bond, and for isobutane, the value is 81.6 (or 341). A consideration of heats of atomization of isomers also illustrates the difficulty. E values for the C— C bonds in pentane, isopentane, and neopentane, calculated from heats of atomization in the same way, are (at 25°C) 81.1, 81.8, and 82.4 kcal mol-1 (339, 342, 345 kJ mol-1), respectively, even though all of them have twelve C— H bonds and four C— C bonds.
These differences have been attributed to various factors caused by the introduction of new structural features. Thus isopentane has a tertiary carbon whose C— H bond does not have exactly the same amount of s character as the C— H bond in pentane, which for that matter contains secondary carbons not possessed by methane. It is known that D values, which can be measured, are not the same for primary, secondary, and tertiary C— H bonds (see Table 5.3). There is also the steric factor. Hence, it is certainly not cor¬ rect to use the value of 99.5 kcal mol-1 (416 kJ mol-1) from methane as the E value for all C— H bonds. Several empirical equations have been devised that account for these factors; the total energy can be computed94 if the proper set of parameters (one for each structural feature) is inserted. Of course, these parameters are originally calculated from the known total energies of some molecules which contain the structural feature.
Table 1.7 gives E values for various bonds. The values given are averaged over a large series of compounds. The literature contains charts that take account of
TABLE 1.7. Bond Energy E Values at 25° C for Some Important Bond Types95a
Bond
kcal mol 1
kJ mol 1
Bond
kcal mol 1
kJ mol 1
0— H
110-111
460-464
c-s96
61
255
C-H
96-99
400-415
C-I
52
220
N-H
93
390
S-H
82
340
c=c
199-200
835
c=c
146-151
610-630
C-F
—
—
C-C
83-85
345-355
C-H
96-99
400-415
C-0
85-91
355-380
C=N
204
854
C-C
83-85
345-355
c=o
173-81
724-757
C-Cl
79
330
C-N97
69-75
290-315
C=N97
143
598
C-Br
66
275
O-O98
42.9
179.6 ±4.5
"The E values are arranged within each group in order of decreasing strength. The values are averaged over a large series of compounds.
94For a review, see Cox, J.D.; Pilcher, G. Thermochemistry of Organic and Organometallic Compounds , Academic Press, NY, 1970, pp. 531-597. See also, Gasteiger, J.; Jacob, P; Strauss, U. Tetrahedron 1979, 35, 139. 95These values, except where noted, are from Lovering, E.G.; Laidler, K.J. Can. J. Chem. 1960, 38, 2367 ; Levi, G.I.; Balandin, A.A. Bull. Acad. Sci. USSR, Div. Chem. Sci. 1960, 149.
96Grelbig, T.; Potter, B.; Seppelt, K. Chem. Ber. 1987, 120, 815.
97Bedford, A.F.; Edmondson, P.B.; Mortimer, C.T. J. Chem. Soc. 1962, 2927 .
98The average of the values obtained was DH°{ O—O). dos Santos, R.M.B.; Muralha, V.S.F.; Correia, C.F.; Simoes, J.A.M. J. Am. Chem. Soc. 2001, 123, 12670.
30 LOCALIZED CHEMICAL BONDING
hybridization (thus an sp 3 C— H bond does not have the same energy as an sp2 C— H bond).99 Bond dissociation energies, both calculated and experientially determined, are constantly being refined. Improved values are available for the 0—0 bond of peroxides,100 the C— H bond in alkyl amines,101 the N— H bond in aniline deriva¬ tives,102 the N— H bond in protonated amines,103 the O— H bond in phenols,104 the C— H bond in alkenes,105 amides and ketones,106 and in CH2X2 and CH3X derivatives (X = COOR, C=0, SR, N02, etc.),107 the 0-H and S-H bonds of alcohols and thiols,108 and the C— Si bond of aromatic silanes.109 Solvent plays a role in the E values. When phenols bearing electron-releasing groups are in aqueous media, calculations show that the bond dissociation energies of decrease due to hydrogen-bonding interactions with water molecules, while electron- withdrawing substituents on the phenol increase the bond dissociation energies.110
Certain generalizations can be derived from the data in Table 1.7.
1. There is a correlation of bond strengths with bond distances. A comparison of Tables 1.5 and 1.7 shows that, in general, shorter bonds are stronger bonds. Since we have already seen that increasing s character shortens bonds (p. 24), it follows that bond strengths increase with increasing s character. Calcula¬ tions show that ring strain has a significant effect on bond dissociation energy, particularly the C— H bond of hydrocarbons, because it forces the compound to adopt an undesirable hybridization.111
2. Bonds become weaker as we move down the Periodic Table. Compare C— O and C— S, or the carbon-halogen bonds C— F, C— Cl, C— Br, C— I. This is a consequence of the first generalization, since bond distances must increase as we go down the periodic table because the number of inner electrons increases. However, it is noted that “high-level ab initio molecular- orbital calculations confirm that the effect of alkyl substituents on R— X bond dissociation energies varies according to the nature of X (the stabilizing
"Cox, J.D.; Pilcher, G. Thermochemistry of Organic and Organometallic Compounds , Academic Press, NY, 1970, pp. 531-597; Cox, J.D. Tetrahedron 1962, 18, 1337.
100Bach, R.D.; Ayala, P.Y.; Schlegel, H.B. J. Am. Chem. Soc. 1996, 118, 12758.
101Wayner, D.D.M.; Clark, K.B.; Rauk, A.; Yu, D.; Armstrong, D.A. J. Am. Chem. Soc. 1997, 119, 8925. For the a C— H bond of tertiary amines, see Dombrowski, G.W.; Dinnocenzo, J.P.; Farid, S.; Goodman, J.L. Gould, I.R. J. Org. Chem. 1999, 64, 427.
102Bordwell, F.G.; Zhang, X.-M.; Cheng, J.-P. J. Org. Chem. 1993, 58, 6410. See also, Li, Z.; Cheng, J.-P. J. Org. Chem. 2003, 68, 7350.
103Liu, W.-Z.; Bordwell, F.G. J. Org. Chem. 1996, 61, 4778.
104Lucarini, M.; Pedrielli, R; Pedulli, G.F.; Cabiddu, S.; Fattuoni, C. J. Org. Chem. 1996, 61, 9259. For the O— H E of polymethylphenols, see de Heer, M.I.; Korth, H.-G.; Mulder, P. J. Org. Chem. 1999, 64, 6969. 105Zhang, X.-M. J. Org. Chem. 1998, 63, 1872.
106Bordwell, F.G.; Zhang, X.-M.; Filler, R. J. Org. Chem. 1993, 58, 6067.
107Brocks, J.J.; Beckhaus, H.-D.; Beckwith, A.L.J.; Ruchardt, C. J. Org. Chem. 1998, 63, 1935. 108Hadad, C.M.; Rablen, P.R.; Wiberg, K.B. J. Org. Chem. 1998, 63, 8668.
109Cheng, Y.-H.; Zhao, X.; Song, K.-S.; Liu, L.; Guo, Q.-X. J. Org. Chem. 2002, 67, 6638.
110Guerra, M.; Amorati, R.; Pedulli, G.F. J. Org. Chem. 2004, 69, 5460.
inFeng, Y.; Liu, L.; Wang, J.-T.; Zhao, S.-W.; Guo, Q.X. J. Org. Chem. 2004, 69, 3129; Song, K.-S.; Liu, L.; Guo, Q.X. Tetrahedron 2004, 60, 9909.
CHAPTER 1
BOND ENERGIES 31
influence of the ionic configurations to increase in the order Me < Et < i- Pr < t- Bu, accounting for the increase (rather than expected decrease) in the R— X bond dissociation energies with increasing alkylation in the R— OCH3, R— OH, and R— F molecules. This effect of X can be understood in terms of the increasing contribution of the ionic R+X“ configuration for electro¬ negative X substituents.”112
3. Double bonds are both shorter and stronger than the corresponding single bonds, but not twice as strong, because n overlap is less than a overlap. This means that a a bond is stronger than a n bond. The difference in energy between a single bond, say C— C, and the corresponding double bond is the amount of energy necessary to cause rotation around the double bond.113
112Coote, M.L.; Pross, A.; Radom, L. Org. Lett. 2003, 5, 4689.
113For a discussion of the different magnitdues of the bond energies of the two bonds of the double bond, see Miller, S.I. J. Chem. Educ. 1978, 55, 778.
CHAPTER 2
Delocalized Chemical Bonding
Although the bonding of many compounds can be adequately described by a single Lewis structure (p. 14), this is not sufficient for many other compounds. These com¬ pounds contain one or more bonding orbitals that are not restricted to two atoms, but that are spread out over three or more. Such bonding is said to be delocalized } In this chapter, we will see which types of compounds must be represented in this way.
The two chief general methods of approximately solving the wave equation, dis¬ cussed in Chapter 1, are also used for compounds containing delocalized bonds.2 In the valence-bond method, several possible Lewis structures (called canonical forms ) are drawn and the molecule is taken to be a weighted average of them. Each ^ in Eq. (1.3), Chapter 1,
^ = Cixj/j +Ci\|/! H -
represents one of these structures. This representation of a real structure as a weighted average of two or more canonical forms is called resonance. For benzene the canonical forms are 1 and 2. Double-headed arrows ( <-> ) are used to indicate resonance. When the wave equation is solved, it is found that the energy value obtained by considering that 1 and 2 participate equally is lower than that for 1 or 2 alone. If 3, 4, and 5 (called Dewar structures ) are also considered, the value
lrThe classic work on delocalized bonding is Wheland, G.W. Resonance in Organic Chemistry, Wiley, NY,
1955.
2There are other methods. For a discussion of the free-electron method, see Streitwieser Jr., A. Molecular Orbital Theory for Organic Chemists', Wiley, NY, 1961, pp. 27-29. For the nonpairing method, in which benzene is represented as having three electrons between adjacent carbons, see Hirst, D.M.; Linnett, J.W. J. Chem. Soc. 1962, 1035; Firestone, R.A. J. Org. Chem. 1969, 34, 2621.
March ’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Sixth Edition, by Michael B. Smith and Jerry March Copyright © 2007 John Wiley & Sons, Inc.
32
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 33
is lower still. According to this method, 1 and 2 each contribute 39% to the actual molecule and the others 7.3% each.3 The carbon-carbon bond order is 1.463 (not 1.5, which would be the case if only 1 and 2 contributed). In the valence-bond method, the bond order of a particular bond is the sum of the weights of those cano¬ nical forms in which the bonds is double plus 1 for the single bond that is present in all of them.4 Thus, according to this picture, each C— C bond is not halfway between a single and a double bond but somewhat less. The energy of the actual molecule is obviously less than that of any one Lewis structure, since otherwise it would have one of those structures. The difference in energy between the actual molecule and the Lewis structure of lowest energy is call the resonance energy. Of course, the Lewis structures are not real, and their energies can only be estimated.
Qualitatively, the resonance picture is often used to describe the structure of molecules, but quantitative valence-bond calculations become much more difficult as the structures become more complicated (e.g., naphthalene, and pyridine). Therefore, the molecular-orbital method is used much more often for the solution of wave equations.5 If we look at benzene by this method (qualitatively), we see that each carbon atom, being connected to three other atoms, uses sp2 orbitals to form a bonds, so that all 12 atoms are in one plane. Each carbon has a p orbital (containing one electron) remaining and each of these can overlap equally with the two adjacent p orbitals. This overlap of six orbitals (see Fig. 2.1) produces six new orbitals, three of which (shown) are bonding. These three (called n orbitals) all occupy approximately the same space.6 One of the three is of lower energy than the other two, which are degenerate. They each have the plane of the ring as a node and so are in two parts, one above and one below the plane. The two orbitals of higher energy (Fig. 2.1 b and c ) also have another node. The six electrons that occu¬ py this torus-shaped cloud are called the aromatic sextet. The carbon-carbon bond order for benzene, calculated by the molecular-orbital method, is 1.667. 7
For planar unsaturated and aromatic molecules, many molecular-orbital calcula¬ tions (MO calculations ) have been made by treating the a and n electrons sepa¬ rately. It is assumed that the a orbitals can be treated as localized bonds and the
3Pullman, A. Prog. Org. Chem. 1958, 4, 31, p. 33.
4For a more precise method of calculating valence-bond orders, see Clarkson, D.; Coulson, C.A.; Goodwin, T.H. Tetrahedron 1963, 19, 2153. See also Herndon, W.C.; Parkanyi, C. J. Chem. Educ. 1976, 53, 689.
5For a review of how MO theory explains localized and delocalized bonding, see Dewar, M.J.S. Mol. Struct. Energ., 1988, 5, 1.
6According to the explanation given here, the symmetrical hexagonal structure of benzene is caused by both the a bonds and the n orbitals. It has been contended, based on MO calculations, that this symmetry is caused by the a framework alone, and that the n system would favor three localized double bonds: Shaik, S.S.; Hiberty, PC.; Lefour, J.; Ohanessian, G. J. Am. Chem. Soc. 1987, 109, 363; Stanger, A.; Vollhardt, K.P.C. J. Org. Chem. 1988, 53, 4889. See also Cooper, D.L.; Wright, S.C.; Gerratt, J.; Raimondi, M. J. Chem. Soc. Perkin Trans. 2 1989, 255, 263; Jug, K.; Koster, A.M. J. Am. Chem. Soc. 1990, 112, 6772; Aihara, J. Bull. Chem. Soc. Jpn. 1990, 63, 1956.
7The molecular-orbital method of calculating bond order is more complicated than the valence-bond method. See Pullman, A. Prog. Org. Chem. 1958, 4, 31, p. 36; Clarkson, D.; Coulson, C.A.; Goodwin, T.H. Tetrahedron 1963, 19, 2153.
34 DELOCALIZED CHEMICAL BONDING
H H
Superposition of (a), (b), and (c). (d)
Fig. 2.1. The six p orbitals of benzene overlap to form three bonding orbitals, ( a ), (b), and (c). The three orbitals superimposed are shown in (d).
calculations involve only the n electrons. The first such calculations were made by Hiickel; such calculations are often called Hiickel molecular-orbital (HMO) calcu¬ lations. 8 Because electron-electron repulsions are either neglected or averaged out in the HMO method, another approach, the self-consistent field (SCF), or Hartree- Fock , method, was devised.9 Although these methods give many useful results for
8See Yates, K. Hiickel Molecular Orbital Theory, Academic Press, NY, 1978; Coulson, C.A.; O’Leary, B.; Mallion, R.B. Hiickel Theory for Organic Chemists, Academic Press, NY, 1978; Lowry, T.H.; Richardson, K.S. Mechanism and Theory in Organic Chemistry, 3rd ed., Harper and Row, NY, 1987, pp. 100-121.
9Roothaan, C.C.J. Rev. Mod. Phys. 1951, 23, 69; Pariser, R.; Parr, R.G. J. Chem. Phys. 1952, 21, 466, 767; Pople, J.A. Trans. Faraday Soc,. 1953, 49, 1375, J. Phys. Chem. 1975, 61, 6; Dewar, M J.S. The Molecular Orbital Theory of Organic Chemistry; McGraw-Hill, NY, 1969; Dewar, in Aromaticity, Chem. Soc.
Spec. Pub. no. 21, 1967, pp. 177-215.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 35
planar unsaturated and aromatic molecules, they are often unsuccessful for other molecules; it would obviously be better if all electrons, both a and n, could be included in the calculations. The development of modern computers has now made this possible.10 Many such calculations have been made11 using a number of methods, among them an extension of the Hiickel method (EHMO)12 and the application of the SCF method to all valence electrons.13
One type of MO calculation that includes all electrons is called ab initio.14 Despite the name (which means “from first principles”) this type does involve assumptions, though not very many. It requires a large amount of computer time, especially for molecules that contain more than about five or six atoms other than hydrogen. Treatments that use certain simplifying assumptions (but still include all electrons) are called semiempirical methods.15 One of the first of these was called CNDO (Complete Neglect of Differential Overlap),16 but as computers have become more powerful, this has been superseded by more modern methods, includ¬ ing MINDO/3 (Modified Intermediate Neglect of Differential Overlap),17 MNDO (Modified Neglect of Diatomic Overlap),17 and AMI (Austin Model 1), all of which were introduced by M.J. Dewar and co-workers.18 Semiempirical calculations are generally regarded as less accurate than ab initio methods,19 but are much faster and cheaper. Indeed, calculations for some very large molecules are possible only with the semiempirical methods.20
Molecular-orbital calculations, whether by ab initio or semiempirical methods, can be used to obtain structures (bond distances and angles), energies (e.g., heats of formation), dipole moments, ionization energies, and other properties of molecules,
10For discussions of the progress made in quantum chemistry calculations, see Ramsden, C.A. Chem. Ber. 1978, 14, 396; Hall, G.G. Chem. Soc. Rev. 1973, 2, 21.
nFor a review of molecular-orbital calculatons on saturated organic compounds, see Herndon, W.C. Prog. Phys. Org. Chem. 1972, 9, 99.
12Hoffmann, R. J. Chem. Phys. 1963, 39, 1397. See Yates, K. Hiickel Molecular Orbital Theory, Academic Press, NY, 1978, pp. 190-201.
13Dewar, M.J.S. The Molecular Orbital Theory of Chemistry, McGraw-Hill, NY, 1969, Jaffe, H.H. Acc. Chem. Res. 1969, 2, 136; Kutzelnigg, W.; Del Re, G.; Berthier, G. Fortschr. Chem. Forsch. 1971, 22, 1. 14Hehre, W.J.; Radom, L.; Schley er, P.v.R.; Pople, J.A. Ab Initio Molecular Orbital Theory, Wiley, NY, 1986, Clark, T. A Handbook of Computational Chemistry, Wiley, NY, 1985, pp. 233-317; Richards, W.G.; Cooper, D.L. Ab Initio Molecular Orbital Calculations for Chemists, 2nd ed., Oxford University Press: Oxford, 1983.
15For a review, see Thiel, W. Tetrahedron 1988, 44, 7393.
16Pople, J.A.; Santry, D.P.; Segal, G.A. J. Chem. Phys. 1965, 43, S129; Pople, J.A.; Segal, G.A. J. Chem. Phys. 1965, 43, SI 36; 1966, 44, 3289; Pople, J.A.; Beveridge, D.L. Approximate Molecular Orbital Theory, McGraw-Hill, NY, 1970.
17For a discussion of MNDO and MINDO/3, and a list of systems for which these methods have been used, with references, see Clark, T. A Handbook of Computational Chemistry, Wiley, NY, 1985, pp. 93- 232. For a review of MINDO/3, see Lewis, D.F.V. Chem. Rev. 1986, 86, 1111.
18First publications are, MINDO/3: Bingham, R.C.; Dewar, M.J.S. ; Lo, D.H. J. Am. Chem. Soc. 1975, 97, 1285; MNDO: Dewar, M.J.S.; Thiel, W. J. Am. Chem. Soc. 1977, 99, 4899; AMI: Dewar, M.J.S.; Zoebisch, E.G.; Healy, E.F.; Stewart, J.J.P. J. Am. Chem. Soc. 1985, 107, 3902.
19See, however, Dewar, M.J.S.; Storch, D.M. J. Am. Chem. Soc. 1985, 107, 3898.
20Clark, T. A Handbook of Computational Chemistry, Wiley, NY, 1985, p. 141.
36 DELOCALIZED CHEMICAL BONDING
ions, and radicals: not only of stable ones, but also of those so unstable that these properties cannot be obtained from experimental measurements.21 Many of these calculations have been performed on transition states (p. 302); this is the only way to get this information, since transition states are not, in general, directly obser¬ vable. Of course, it is not possible to check data obtained for unstable molecules and transition states against any experimental values, so that the reliability of the various MO methods for these cases is always a question. However, our confidence in them does increase when (7) different MO methods give similar results, and (2) a particular MO method works well for cases that can be checked against experimen¬ tal methods.22
Both the valence-bond and molecular-orbital methods show that there is deloca¬ lization in benzene. For example, each predicts that the six carbon-carbon bonds should have equal lengths, which is true. Since each method is useful for certain purposes, we will use one or the other as appropriate. Recent ab initio , SCF calcu¬ lations confirms that the delocalization effect acts to strongly stabilize symmetric benzene, consistent with the concepts of classical resonance theory.23
Bond Energies and Distances in Compounds Containing Delocalized Bonds
If we add the energies of all the bonds in benzene, taking the values from a source like Table 1.7, the value for the heat of atomization turns out to be less than that actually found in benzene (Fig. 2.2). The actual value is 1323 kcal mol-1 (5535 kJ mol-1). If we use E values for a C=C double bond obtained from cyclo¬ hexene (148.8 kcal mol-1; 622.6 kJ mol-1), a C— C single bond from cyclohexane (81.8 kcal mol-1, 342 kJ mol-1), and C-H bonds from methane (99.5 kcal mol-1, 416 kJ mol-1), we get a total of 1289 kcal mol-1 (5390 kJ mol-1) for structure 1 or 2. By this calculation the resonance energy is 34 kcal mol-1 (145 kJ mol-1). Of course, this is an arbitrary calculation since, in addition to the fact that we are calculating a heat of atomization for a nonexistent structure (1), we are forced to use E values that themselves do not have a firm basis in reality. The actual C— H bond energy for benzene has been measured to be 1 13.5 ± 0.5 kcal mol-1 at 300 K and estimated to be 1 12.0 =b 0.6 kcal mol-1 (469 kJ mol-1) at 0 K.24 The resonance energy can never be measured, only estimated, since we can measure the heat of atomization of the real molecule but can only make an intelligent guess at that of the Lewis structure of lowest energy.
21Another method of calculating such properies is molecular mechanics (p. $$$).
22Dias, J.R. Molecular Orbital Calculations Using Chemical Graph Theory, Spring -Verlag, Berlin, 1993. 23Glendening, E.D.; Faust, R.; Streitwieser, A.; Vollhardt, K.P.C.; Weinhold, F. J. Am. Chem.Soc. 1993, 115, 10952.
24Davico, G.E.; Bierbaum, V.M.; DePuy, C.H.; Ellison, G.B.; Squires, R.R. J. Am. Chem. Soc. 1995, 117, 2590. See also Barckholtz, C.; Barckholtz, T.A.; Hadad, C.M. J. Am. Chem. Soc. 1999, 121, 491; Pratt, D.A.; DiLabio, G.A.; Mulder, R; Ingold, K.U. Acc. Chem. Res. 2004, 37, 334.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 37
Energy of six carbon and six hydrogen atoms
0 1/
1
Resonance energy
Energy of structure 1 or 2
Energy of benzene
Fig. 2.2. Resonance energy in benzene.
Another method frequently used for estimation of resonance energy involves measurements of heats of hydrogenation.25 Thus, the heat of hydrogenation of cyclo¬ hexene is 28.6 kcal mol-1 (120 kJ mol-1), so we might expect a hypothetical 1 or 2 with three double bonds to have a heat of hydrogenation of about 85.8 kcal mol-1 (360 kJ mol-1). The real benzene has a heat of hydrogenation of 49.8 kcal mol-1 (208 kJ mol-1), which gives a resonance energy of 36 kcal mol-1 (152 kJ mol-1). By any calculation the real molecule is more stable than a hypothetical 1 or 2.
The energies of the six benzene orbitals can be calculated from HMO theory in terms of two quantities, a and (3. The parameter a is the amount of energy possessed by an isolated 2 p orbital before overlap, while (3 (called the resonance integral) is an energy unit expressing the degree of stabilization resulting from 71-orbital over¬ lap. A negative value of (3 corresponds to stabilization, and the energies of the six orbitals are (lowest to highest): a + 2(3, a + (3, a + (3, a — (3, oc — (3, and oc — 2(3.26 The total energy of the three occupied orbitals is 6a + 8(3, since there are two elec¬ trons in each orbital. The energy of an ordinary double bond is a + (3, so that struc¬ ture 1 or 2 has an energy of 6a + 6(3. The resonance energy of benzene is therefore 2(3. Unfortunately, there is no convenient way to calculate the value of (3 from mole¬ cular-orbital theory. It is often given for benzene as about 18 kcal mol-1 (76kJmol-1); this number being one-half of the resonance energy calculated from heats of combustion or hydrogenation. Using modern ab initio calculations, bond resonance energies for many aromatic hydrocarbons other than benzene have been reported 27
25For a review of heats of hydrogenation, with tables of values, see Jensen, J.L. Prog. Phys. Org. Chem. 1976 , 72, 189.
26For the method for calculating these and similar results given in this chapter, see Higasi, K.; Baba, H.; Rembaum, A. Quantum Organic Chemistry, Interscience, NY, 1965. For values of calculated orbital energies and bond orders for many conjugated molecules, see Coulson, C.A.; Streitwieser, Jr., A. Dictionary ofn Electron Calculations, W.H. Freeman, San Francisco, 1965.
27Aihara, J-i. J. Chem. Soc. Perkin Trans 2 1996, 2185.
38 DELOCALIZED CHEMICAL BONDING
Isodesmic and homodesmotic reactions are frequently used for the study of aromaticity from the energetic point of view.28 However, the energy of the reactions used experimentally or in calculations may reflects only the relative aromaticity of benzene and not its absolute aromaticity. A new homodesmo¬ tic reactions based on radical systems predict an absolute aromaticity of 29.13 kcal mol-1 (121.9 kJ mol-1) for benzene and an absolute antiaromaticity of 40.28 kcal mol-1 (168.5 kJ mol-1) for cyclobutadiene at the MP4(SDQ)/ 6-31 G-(d,p) level.29
We might expect that in compounds exhibiting delocalization the bond distances would lie between the values gives in Table 1.5. This is certainly the case for ben¬ zene, since the carbon-carbon bond distance is 1.40 A,30 which is between the 1.48 A for an sp2-sp 2 C—C single bond and the 1.32 A of the sp2-sp 2 C=C double bond.31
Kinds of Molecules That Have Delocalized Bonds
There are four main types of structure that exhibit delocalization:
1. Double (or Triple ) Bonds in Conjugation ,32 The double bonds in benzene are conjugated, of course, but the conjugation exists in acyclic molecules such as butadiene. In the molecular orbital picture (Fig. 2.3), the overlap of four orbitals gives two bonding orbitals that contain the four electrons and two vacant antibonding orbitals. It can be seen that each orbital has one more node than the one of next lower energy. The energies of the four orbitals are (lowest to highest): a+1.618p, a + 0.618p, a — 0.618P, and a — 1.618P; hence the total energy of the two occupied orbitals is 4a + 4.472p. Since the energy of two isolated double bonds is 4a + 4p, the resonance energy by this calculation is 0.472p.
In the resonance picture, these structures are considered to contribute:
ch2=ch-ch=ch2 ^ ch2-ch=ch-ch2 ^ ch2-ch=ch-ch2
6 7 8
28Hehre, W.J.; Ditchfield, R.; Radom, L.; Pople, J.A. J. Am. Chem.Soc. 1970, 92, 4796; Hehre, W.J.; Radom, L.; Pople, J.A. J. Am. Chem. Soc. 1971, 93, 289; George, P.; Trachtman, M.; Bock, C.W.; Brett, A.M. Theor. Chim. Acta, 1975, 38, 121; George, P.; Trachtman, M.; Bock, C.W.; Brett, A.M. J. Chem. Soc. Perkin Trans. 2 1976, 1222; George, P.; Trachtman, M.; Brett, A.M. Bock, C.W.; Tetrahedron 1976, 32, 317; George, P.; Trachtman, M.; Brett, A.M.; Bock, C.W. J. Chem. Soc. Perkin Trans. 2 1977, 1036. 29Suresh, C.H.; Koga, N. J. Org. Chem. 2002, 67, 1965.
30Bastiansen, O.; Fernholt, L.; Seip, H.M.; Kambara, H.; Kuchitsu, K. J. Mol. Struct. 1973, 18, 163; Tamagawa, K.; Iijima, T.; Kimura, M. J. Mol. Struct. 1976, 30, 243.
31The average C—C bond distance in aromatic rings is 1.38 A: Allen, F.H.; Kennard, O.; Watson, D.G.; Brammer, L.; Orpen, A.G.; Taylor, R. J. Chem. Soc. Perkin Trans. 2 1987, p. S8.
32For reviews of conjugation in open-chain hydrocarbons, see Simmons, H.E. Prog. Phys. Org. Chem. 1970, 7, 1; Popov, E.M.; Kogan, G.A. Russ. Chem. Rev. 1968, 37, 119.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 39
Bonding orbitals (71)
Fig. 2.3. The four rc-orbitals of butadiene, formed by overlap of four p orbitals.
In either picture, the bond order of the central bond should be > 1 and that of the other carbon-carbon bonds <2, although neither predicts that the three bonds have equal electron density. Molecular- orbital bond orders of 1.894 and 1.447 have been calculated.33
The existence of delocalization in butadiene and similar molecules has been questoned. The bond lengths in butadiene are 1.34 A for the double bonds and 1.48 A for the single bond.34 Since the typical single-bond distance of a bond that is not adjacent to an unsaturated group is 1.53 A (p. 26), it has been argued that the shorter single bond in butadiene provides evidence for resonance. However, this shortening can also be explained by hybridization changes (see p. 26); and other explanations have also been offered.35 Resonance energies for butadienes, calculated from heats of combustion or hydrogenation, are only about 4kcalmol 1 (17kJmol_1), and these values may not be entirely attributable to resonance. Thus, a calculation from heat of atomization data gives a resonance energy of 4.6kcalmol_1 (19kJmol_1) for cA-l,3-pentadiene, and — 0.2 kcal mol-1 (— 0.8 kJmol-1), for 1,4-pentadiene. These two compounds, each of which possesses two double bonds, two C— C single bonds, and eight C— H bonds, would seem to offer as similar a comparison as we could make of a conjugated with a nonconjugated compound, but they are nevertheless not strictly comparable. The former has three sp3 C— H and five sp 2 C— H bonds, while the latter has two and six, respectively. Also, the two single C— C bonds
33Coulson, C.A. Proc. R. Soc. London, Ser. A 1939, 169, 413.
34Marais, D.J.; Sheppard, N.; Stoicheff, B.P. Tetrahedron 1962, 17, 163.
35Bartell, L.S. Tetrahedron 1978, 34, 2891, J. Chem. Educ. 1968, 45, 754; Wilson, E.B. Tetrahedron 1962, 17, 191; Hughes, D.O. Tetrahedron 1968, 24, 6423; Politzer, P; Harris, D.O. Tetrahedron 1971, 27, 1567.
40 DELOCALIZED CHEMICAL BONDING
of the 1,4-diene are both sp2-sp 3 bonds, while in the 1,3-diene, one is sp2-sp 3 and the other sp2-sp2. Therefore, it may be that some of the already small value of 4 kcalmol 1 (17kJmol_1) is not resonance energy but arises from differing energies of bonds of different hybridization.36
Although bond distances fail to show it and the resonance energy is low, the fact that butadiene is planar37 shows that there is some delocalization, even if not as much as previously thought. Similar delocalization is found in other conjugated systems (e.g., C=C— C=038 and C=C— C=N), in longer systems with three or more multiple bonds in conjugation, and where double or triple bonds are conjugated with aromatic rings. Diynes such as 1,3- butadiyne (9) are another example of conjugated molecules. Based on calculations, Rogers et al. reported that the conjugation stabilization of 1,3- butadiyne is zero.39 Later calculations concluded that consideration of hyperconjugative interactions provides a more refined measure of conjugative stabilization.40 When this measure is used, the conjugation energies of the isomerization and hydrogenation reactions considered agree with a conjuga¬ tive stabilization of 9.3 ( 0.5 kcalmol-1 for diynes and 8.2 (0.1 kcalmol-1 for dienes.
H— C=C— C=C— H
9
2. Double (or Triple ) Bonds in Conjugation with a p Orbital on an Adjacent Atom. Where a p orbital is on an atom adjacent to a double bond, there are three parallel p orbitals that overlap. As previously noted, it is a general rule that the overlap of n atomic orbitals creates n molecular orbitals, so overlap of a p orbital with an adjacent double bond gives rise to three new orbitals, as
36For negative views on delocalization in butadiene and similar molecules, see Dewar, Gleicher,
G J. J. Am. Chem. Soc. 1965, 87, 692; Brown, M.G. Trans. Faraday Soc. 1959, 55, 694; Somayajulu, G.R. J. Chem. Phys. 1959, 31, 919; Mikhailov, B.M. J. Gen. Chem. USSR 1966, 36, 379. For positive views, see Miyazaki, T.; Shigetani, T.; Shinoda, H. Bull. Chem. Soc. Jpn. 1971, 44, 1491; Berry, R.S. J. Chem. Phys. 1962, 30, 936; Kogan, G.A.; Popov, E.M. Bull. Acad. Sci. USSR Div. Chem. Sci. 1964, 1306; Altmann, J.A.; Reynolds, W.F. J. Mol. Struct., 1977, 36, 149. In general, the negative argument is that resonance involving excited structures, such as 7 and 8, is unimportant. See rule 6 on p. $$$. An excellent discussion of the controversy is found in Popov, E.M.; Kogan, G.A. Russ. Chem. Rev. 1968, 37, 119, pp. 119-124. 37Marais, D.J.; Sheppard, N.; Stoicheff, B.P. Tetrahedron 1962, 17, 163; Fisher, J.J.; Michl, J. J. Am. Chem. Soc. 1987, 109, 1056; Wiberg, K.B.; Rosenberg, R.E.; Rablen, PR. J. Am. Chem. Soc. 1991, 113, 2890.
38For a treatise on C=C— C=0 systems, see Patai, S.; Rappoport, Z. The Chemistry ofEnones, two parts; Wiley, NY, 1989.
39Rogers, D.W.; Matsunaga, N.; Zavitsas, A.A.; McLafferty, F.J.; Liebman, J.F. Org. Lett. 2003, 5, 2373; Rogers, D.W.; Matsunaga, N.; McLafferty, F.J.; Zavitsas, A. A.; Liebman, J.F. J. Org. Chem. 2004, 69, 7143.
40Jarowski, P.D.; Wodrich, M.D.; Wannere, C.S.; Schleyer, P.v.R.; Houk, K.N. J. Am. Chem. Soc. 2004, 126, 15036.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 41
Fig. 2.4. The three orbitals of an allylic carbon, formed by overlap of three p orbitals.
shown in Fig. 2.4. The middle orbital is a nonbonding orbital of zero bonding energy. The central carbon atom does not participate in the nonbonding orbital.
There are three cases: the original p orbital may have contained two, one, or no electrons. Since the original double bond contributes two electrons, the total number of electrons accommodated by the new orbitals is four, three, or two. A typical example of the first situation is vinyl chloride CH2=CH— Cl. Although the p orbital of the chlorine atom is filled, it still overlaps with the double bond (see 10). The four electrons occupy the two molecular orbitals of lowest energies. This is our first example of resonance involving overlap between unfilled orbitals and a filled orbital. Canonical forms for vinyl chloride are shown in 11.
Any system containing an atom that has an unshared pair and that is directly attached to a multiple-bond atom can show this type of delocalization.
42 DELOCALIZED CHEMICAL BONDING
Another example is the carbonate ion:
%'°e
The bonding in ally lie carbanions, for example, CH2=CH— CITf, is similar.
The other two cases, where the original p orbital contains only one or no electron, are generally found only in free radicals and cations, respectively. Ally lie free radicals have one electron in the nonbonding orbital. In allylic cations this orbital is vacant and only the bonding orbital is occupied. The orbital structures of the allylic carbanion, free radical, and cation differ from each other, therefore, only in that the nonbonding orbital is filled, half-filled, or empty. Since this is an orbital of zero bonding energy, it follows that the bonding n energies of the three species relative to electrons in the 2 p orbitals of free atoms are the same. The electrons in the nonbonding orbital do not contribute to the bonding energy, positively or negatively.41
By the resonance picture, the three species may be described as having double bonds in conjugation with, respectively, an unshared pair, an unpaired electron, and an empty orbital as in the allyl cation 12 (see Chapter 5).
X ,0* c
ii
o
©
ch2=ch-ch2
0
ch2-ch=ch2
ch2=ch-ch2
ch2-ch=ch2
© ©
CH2=CH-CH2 - - CH2-CH=CH2
12
3. n- Allyl and Other rj-Complexes. In the presence of transition metals, delocalized cations are stabilized by donating electrons to the metal.42 In a C— Metal bond, such as H3C— Fe, the carbon donates (shares) one electron with them metal, and is considered to be a one-electron donor. With a 7t-bond, such as that found in ethylene, both electrons can be donated to the metal to
41It has been contended that here too, as with the benzene ring (Ref. 6), the geometry is forced upon allylic systems by the a framework, and not the n system: Shaik, S.S.; Hiberty, P.C.; Ohanessian, G.; Lefour, J. Nouv. J. Chim., 1985, 9, 385. It has also been suggested, on the basis of ab initio calculations, that while the allyl cation has significant resonance stabilization, the allyl anion has little stabilization: Wiberg, K.B.; Breneman, C.M.; LePage, T.J. J. Am. Chem. Soc. 1990, 112, 61.
42Crabtree, R.H. The Organometallic Chemistry of the Transition Metals, Wiley-Interscience, NY, 2005', Hill, A.F. Organotransition Metal Chemistry, Wiley Interscience, Canberra, 2002.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 43
form a complex such as 14 by reaction of Wilkinson’s catalyst (13) with an alkene and hydrogen gas,43 and the Ti-bond is considered to be a two-electron donor. In these two cases, the electron donating ability of the group coordinated to the metal (the ligand) is indicated by terminology p1, r|2, p3, and so on, for a one-, two-, and three-electron donor, respectively.
Ph3P<
Ph3P'
» PPh3 Cl
\
/
/
c = c^, h2
- PPh3
13
H
Ph^>'<
C1^t
/C=CV
rCl
14
Wilkinson's catalyst
Ligands can therefore be categorized as p -ligands according to their electron donation to the metal. A hydrogen atom (as in 14) or a halogen ligand (as in 13) are p1 ligands and an amine (NR3), a phosphine (PR3, as in 13, 14, and 18), CO (as in 16 or 17), an ether (OR2) or a thioether (SR2) are p2 ligands. Hydrocarbon ligands include alkyl (as the methyl in 15) or aryl with a C— metal bond (p1), alkenes or carbenes (p2, see p. 116), 7i-allyl (p3), conjugated dienes such as 1,3-butadiene (p4), cyclopentadienyl (p5, as in 15 and see p 63), and arenes or benzene (p6).44 Note that in the formation of 14 from 13, the two electron donor alkene displaces a two-electron donor phosphine. Other typical complexes include chromium hexacarbonyl Cr(CO)6 (16), with six p2-CO ligands; p6-C6H6Cr(CO)3 (18), and tetrakis- triphenylphosphinopalladium (0), 17, with four p2-phosphine ligands.
In the context of this section, the electron-delocalized ligand 7i-allyl (12) is an p3 donor and it is well known that ally lie halides react with PdCl2 to form a &A-p3-complex 19 (see the 3D model 20) 45 Complexes, such as 19, react with nucleophiles to give the corresponding coupling product (10-60) 46 The
43Jardine, F.H., Osborn, J.A.; Wilkinson, G.; Young, G.F. Chem. Ind. ( London ) 1965, 560; Imperial Chem. Ind. Ltd., Neth. Appl. 6,602,062 [Chem. Abstr., 66: 10556y 196T\\ Bennett, M.A.; Longstaff, RA. Chem. Ind. 1965, 846.
44Davies, S.G. Organotransition Metal Chemistry, Pergamon, Oxford, 1982, p. 4.
45Trost, B.M.; Strege, P.E.; Weber, L.; Fullerton, T.J.; Dietsche, T.J. J. Am. Chem. Soc. 1978, 100, 3407. 46Trost, B.M.; Weber, L.; Strege, PE.; Fullerton, T.J.; Dietsche, T.J. J. Am. Chem. Soc., 1978 100, 3416.
44 DELOCALIZED CHEMICAL BONDING
reaction of allylic acetates or carbons and a catalytic amount of palladium (0) compounds also lead to an r|3-complex that can react with nucleophiles.47
4. Hyperconjugation . The type of delocalization called hyperconjugation , is discussed on p. 95.
We will find examples of delocalization that cannot be strictly classified as belonging to any of these types.
Cross Conjugation48
In a cross-conjugated compound, three groups are present, two of which are not conjugated with each other, although each is conjugated with the third. Some examples49 are benzophenone (21), triene 22 and divinyl ether 23. Using the
21 22
molecular-orbital method, we find that the overlap of six p orbitals in 22 gives six molecular orbitals, of which the three bonding orbitals are shown in Fig. 2.5, along with their energies. Note that two of the carbon atoms do not participate in the a + P orbital. The total energy of the three occupied orbitals is 6a + 6.900(3, so the resonance energy is 0.900(3. Molecular-orbital bond orders are 1.930 for the C-l,C-2 bond, 1.859 for the C-3,C-6 bond and 1.363 for the C-2,C-3 bond.49 Com- paring these values with those for butadiene (p. 39), we see that the C-l,C-2 bond contains more and the C-3,C-6 bond less double-bond character than the double bonds in butadiene. The resonance picture supports this conclusion, since each C-l,C-2 bond is double in three of the five canonical forms, while the C-3,C-6 bond is double in only one. In most cases, it is easier to treat cross-conjugated
47Melpolder, J.B.; Heck, R.F. J. Org. Chem. 1976, 41, 265; Trost, B.M.; Verhoeven, T.R. J. Am. Chem. Soc., 1976, 98, 630; 1978, 100, 3435; Takahashi, K.; Miyake, A.; Hata, G. Bull Chem. Soc. Jpn. 1970, 45, 230,1183; Trost, B.M.; Verhoeven, T.R. J. Org. Chem. 1976, 41, 3215; Trost, B.M.; Verhoeven, T.R. J. Am. Chem. Soc. 1980, 102, 4730.
48For a discussion, see Phelan, N.F.; Orchin, M. J. Chem. Educ. 1968, 45, 633.
49Compound 22 is the simplest of a family of cross-conjugated alkenes, called dendralenes. For a review of these compounds, see Hopf, H. Angew. Chem. Int. Ed. 1984, 23, 948.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 45
a+ 1.932/3
a + p
Fig. 2.5. The three bonding orbitals of 3-methylelene-l,4-pentadiene (22).
molecules by the molecular-orbital method than by the valence-bond method.
24
One consequence of this phenomenon is that the cross-conjugated C=C unit has a slightly longer bond length that the noncross conjugated bond. In 24, for example, the cross-conjugated bond is ~0.01 A longer.50 The conjugative effect of a C=C or C=C unit can be measured. An ethenyl substituent on a conjugated enone contri¬ butes 4.2 kcal mol-1 and an ethynyl substituent has a more variable effect but con¬ tributes ~2.3 kcal mol-1.51
The phenomenon of homoconjugation is related to cross-conjugation in that there are C=C units in close proximity, but not conjugated one to the other. Homo¬ conjugation arises when the termini of two orthogonal 7i-systems are held in close proximity by being linked by a spiro-tetrahedral carbon atom.52 Spiro [4.4] nonate- traene (25)53 is an example and it known that the HOMO (p. 1208) of 25 is raised relative to cyclopentadiene, whereas the LUMO is unaffected54 Another example
50Traetteberg, M.; Hopf, H. Acta Chem. Scand. B 1994, 48, 989.
51Traetteberg, M.; Liebman, J.F.; Hulce, M.; Bohn, A.A.; Rogers, D.W. J. Chem. Soc. Perkin Trans. 2 1997, 1925.
52Simons, H.E.; Fukunaga, R. J. Am. Chem. Soc. 1967, 89, 5208; Hoffmann, R.; Imamura, A.; Zeiss, G.D. J. Am. Chem. Soc. 1967, 89, 5215; Durr, H.; Gleiter, R. Angew. Chem. Int. Ed. 1978, 17, 559.
53For the synthesis of this molecule, see Semmelhack, M.F.; Foos, J.S.; Katz, S. J. Am. Chem. Soc. 1973, 95, 7325.
54Raman, J.V.; Nielsen, K.E.; Randall, L.H.; Burke, L.A.; Dmitrienko, G.I. Tetrahedron Lett. 1994, 35, 5973.
46 DELOCALIZED CHEMICAL BONDING
is 26, where there are bond length distortions caused by electronic interactions between the unsaturated bicyclic moiety and the cyclopropyl moiety.55 It is assumed that cyclopropyl homoconjugation is responsible for this effect.
The Rules of Resonance
We have seen that one way of expressing the actual structure of a molecule contain¬ ing delocalized bonds is to draw several possible structures and to assume that the actual molecule is a hybrid of them. These canonical forms have no existence except in our imaginations. The molecule does not rapidly shift between them. It is not the case that some molecules have one canonical form and some another. All the molecules of the substance have the same structure. That structure is always the same all the time and is a weighted average of all the canonical forms. In drawing canonical forms and deriving the true structures from them, we are guided by certain rules, among them the following:
1. All the canonical forms must be bona fide Lewis structures (see p. 14). For example, none of them may have a carbon with five bonds.
2. The positions of the nuclei must be the same in all the structures. This means that when we draw the various canonical forms, all we are doing is putting in the electrons in different ways. For this reason, shorthand ways of repre¬ senting resonance are easy to devise:
The resonance interaction of chlorine with the benzene ring can be repre¬ sented as shown in 27 or 28 and both of these representations have been used in the literature to save space. However, we will not use the curved-arrow method of 27 since arrows will be used in this book to express the actual movement of electrons in reactions. We will use representations like 28 or else write out the canonical forms. The convention used in dashed-line formulas like 28 is that bonds that are present in all canonical forms are drawn as solid lines while bonds that are not present in all forms are drawn as dashed lines. In most resonance, a bonds are not involved, and only the n or unshared electrons are put in, in different ways. This means that if we write one canonical form for a molecule, we can then write the others by merely moving n and unshared electrons.
55Haumann, T.; Benet-Buchholz, J.; Klarner, F.-G.; Boese, R. Liebigs Ann. Chem. 1997, 1429.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 47
3. All atoms taking part in the resonance, that is, covered by delocalized electrons, must lie in a plane or nearly so (see p. 48). This, of course, does not apply to atoms that have the same bonding in all the canonical forms. The reason for planarity is maximum overlap of the p orbitals.
4. All canonical forms must have the same number of unpaired electrons. Thus *CH2— CH=CH— CH2» is not a valid canonical form for butadiene.
5. The energy of the actual molecule is lower than that of any form, obviously. Therefore, delocalization is a stabilizing phenomenon.56
6. All canonical forms do not contribute equally to the true molecule. Each form contributes in proportion to its stability, the most stable form contributing most. Thus, for ethylene, the form +CH2— CH^ has such a high energy compared to CH2=CH2 that it essentially does not contribute at all. We have seen the argument that such structures do not contribute even in such cases as butadiene.36 Equivalent canonical forms, such as 1 and 2, contribute equally. The greater the number of significant structures that can be written and the more nearly equal they are, the greater the resonance energy, other things being equal.
It is not always easy to decide relative stabilities of imaginary structures; the che¬ mist is often guided by intuition.57 However, the following rules may be helpful:
a. Structures with more covalent bonds are ordinarily more stable than those with fewer (cf. 6 and 7).
b. Stability is decreased by an increase in charge separation. Structures with formal charges are less stable than uncharged structures. Structures with more than two formal charges usually contribute very little. An especially unfavor¬ able type of structure is one with two like charges on adjacent atoms.
c. Structures that carry a negative charge on a more electronegative atom are more stable than those in which the charge is on a less electronegative atom. Thus, 30 is more stable than 29. Similarly, positive charges are best carried on atoms of low electronegativity.
H2C-r'H H2C;r'H
II " ^ I
O o©
29 30
d. Structures with distorted bond angles or lengths are unstable, for example, the structure 31 for ethane.
56It has been argued that resonance is not a stabilizing phenomenon in all systems, especially in acyclic ions: Wiberg, K.B. Chemtracts: Org. Chem. 1989, 2, 85. See also, Siggel, M.R.; Streitwieser Jr., A.; Thomas, T.D. J. Am. Chem. Soc. 1988, 110, 8022; Thomas, T.D.; Carroll, T.X.; Siggel, M.R. J. Org. Chem. 1988, 53, 1812.
57A quantitative method for weighting canonical forms has been proposed by Gasteiger, J.; Sailer, H. Angew. Chem. Int. Ed. 1985, 24, 687.
H - H
H H
;c=<
H H
31
48 DELOCALIZED CHEMICAL BONDING
The Resonance Effect
Resonance always results in a different distribution of electron density than would be the case if there were no resonance. For example, if 32 were the actual structure of aniline, the two unshared electrons of the nitrogen would reside
nh2 ®nh2 @nh2 @nh2
entirely on that atom. The structure of 32 can be represented as a hybrid that includes contributions from the canonical forms shown, indicating that the electron density of the unshared pair does not reside entirely on the nitrogen, but is spread over the ring. This decrease in electron density at one position (and corresponding increase elsewhere) is called the resonance or mesomeric effect. We loosely say that the NH2 contributes or donates electrons to the ring by a resonance effect, although no actual contribution takes place. The “effect” is caused by the fact that the electrons are in a different place from that we would expect if there were no resonance. In ammonia, where resonance is absent, the unshared pair is located on the nitrogen atom. As with the field effect (p. 20), we think of a certain molecule (in this case ammonia) as a substrate and then see what happens to the electron density when we make a substitution. When one of the hydrogen atoms of the ammonia molecule is replaced by a benzene ring, the electrons are “withdrawn” by the resonance effect, just as when a methyl group replaces a hydrogen of benzene, electrons are “donated” by the the field effect of the methyl. The idea of donation or withdrawal merely arises from the comparison of a compound with a closely related one or a real com¬ pound with a canonical form.
Steric Inhibition of Resonance and the Influences of Strain
Rule 3 states that all the atoms covered by delocalized electrons must lie in a plane or nearly so. Many examples are known where resonance is reduced or prevented because the atoms are sterically forced out of planarity.
Bond lengths for the o- and p- nitro groups in picryl iodide are quite different.58 Distance a in 33 is 1.45 A, whereas b is 1.35 A. This phenomenon can be explained if the oxygens of the p-niixo group are in the plane of the ring and thus in resonance with it, so that b has partial double-bond character, while the oxygens of the 6>-nitro
58Wepster, B.M. Prog. Stereochem. 1958, 2, 99, p. 125. For another example of this type of steric inhibition of resonance, see Exner, O.; Folli, U.; Marcaccioli, S.; Vivarelli, P. J. Chem. Soc. Perkin Trans. 2 1983, 757.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 49
groups are forced out of the plane by the large iodine atom.
The Dewar-type structure for the central ring of the anthracene system in 34 is possible only because the 9,10 substituents prevent the system from being planar.59 34 is the actual structure of the molecule and is not in resonance with forms like 35, although in anthracene itself, Dewar structures and structures like 35 both contribute. This is a consequence of rule 2 (p. 46). In order for a 35-like structure to contribute to resonance in 34, the nuclei would have to be in the same positions in both forms.
35
36
Even the benzene ring can be forced out of planarity.60 In [5]paracyclophane (36), 61 the presence of a short bridge (this is the shortest para bridge known for a benzene ring) forces the benzene ring to become boat- shaped. The parent 36 has so far not proven stable enough for isolation, but a UV spectrum was obtained and showed that the benzene ring was still aromatic, despite the distorted ring.62 The 8,11-dichloro analog of 36 is a stable solid, and X-ray diffraction showed
59Applequist, D.E.; Searle, R. J. Am. Chem. Soc. 1964, 86, 1389.
60For a review of planarity in aromatic systems, see Ferguson, G.; Robertson, J.M. Adv. Phys. Org. Chem. 1963, 1, 203.
61For a monograph, see Keehn, P.M.; Rosenfeld, S.M. Cyclophanes, 2 vols., Academic Press, NY, 1983. For reviews, see Bickelhaupt, F. PureAppl. Chem. 1990, 62, 373; Vogtle, F.; Hohner, G. Top. Curr. Chem. 1978, 74, 1; Cram, D.J.; Cram, J.M. Acc. Chem. Res. 1971, 4, 204; Vogtle, F.; Neumann, P. reviews in Top. Curr. Chem. 1983, 113, 1; 1985, 115, 1.
62Jenneskens, L.W.; de Kanter, F.J.J.; Kraakman, P.A.; Turkenburg, F.A.M.; Koolhaas, W.E.; de Wolf, W.H.; Bickelhaupt, F.; Tobe, Y.; Kakiuchi, K.; Odaira, Y. J. Am. Chem. Soc. 1985, 107, 3716. See also Tobe, Y.; Kaneda, T.; Kakiuchi, K.; Odaira, Y. Chem. Lett. 1985, 1301; Kostermans, G.B.M.; de Wolf, W.E.; Bickelhaupt, F. Tetrahedron Lett. 1986, 27, 1095; van Zijl, P.C.M.; Jenneskens, F.W.; Bastiaan, E.W.; MacLean, C.; de Wolf, W.E.; Bickelhaupt, F. J. Am. Chem. Soc. 1986, 108, 1415; Rice, J.E.; Lee, T.J.; Remington, R.B.; Allen, W.D.; Clabo Jr., D.A.; Schaefer III, H.F. J. Am. Chem. Soc. 1987, 109, 2902.
50 DELOCALIZED CHEMICAL BONDING
that the benzene ring is boat-shaped, with one end of the boat bending ~27° out of the plane, and the other ~12°.63 This compound too is aromatic, as shown by UV and NMR spectra. [6]Paracyclophanes are also bent,64 but in [7]paracyclo- phanes the bridge is long enough so that the ring is only moderately distorted. Simi¬ larly, [ft,m]paracyclophanes (37), where n and m are both 3 or less (the smallest yet prepared is [2.2]paracyclophane), have bent (boat-shaped) benzene rings. All these compounds have properties that depart significantly from those of ordinary benzene compounds. Strained paracyclophanes exhibit both n- and a-strain, and the effect of the two types of strain on the geometry is approximately additive.65 In “belt” cyclophane 38, 66 the molecule has a pyramidal structure with C3 symmetry rather than the planar structure found in [18]-annulene. l,8-Dioxa[8](2,70-pyrenophane (39)67 is another severely distorted aromatic hydrocarbon, in which the bridge undergoes rapid pseudo-rotation (p. 212). A recent study showed that despite sub¬ stantial changes in the hybridization of carbon atoms involving changes in the a-electron structure of pyrenephane, such as 39, the aromaticity of the system decreases slightly and regularly upon increasing the bend angle 0 from 0 to 109. 2°. 68 Heterocyclic paracyclophane analogs have been prepared, such as the report of [2.^](2,5)pyridinophanes.69
63Jenneskens, L.W.; Klamer, J.C.; de Boer, H.J.R.; de Wolf, W.H.; Bickelhaupt, F.; Stam, C.H. Angew. Chem. Int. Ed. 1984, 23, 238.
64See, for example, Liebe, J.; Wolff, C.; Krieger, C.; Weiss, J.; Tochtermann, W. Chem. Ber. 1985, 118, 4144; Tobe, Y.; Ueda, K.; Kakiuchi, K.; Odaira, Y.; Kai, Y.; Kasai, N. Tetrahedron 1986, 42, 1851. 65Stanger, A.; Ben-Mergui, N.; Perl, S. Eur. J. Org. Chem. 2003, 2709.
66Meier, H.; Muller, K. Angew. Chem. Int. Ed., 1995, 34, 1437.
67Bodwell, G.J.; Bridson, J.N.; Houghton, T.J.; Kennedy, J.W.J.; Mannion, M.R. Angew. Chem. Int. Ed., 1996, 35, 1320.
68Bodwell, G.J.; Bridson, J.N.; Cyranski, M.K.; Kennedy, J.W.J.; Krygowski, T.M.; Mannion, M.R.; Miller, D.O. J. Org. Chem. 2003, 68, 2089; Bodwell, G.J.; Miller, D.O.; Vermeij, RJ. Org. Lett. 2001, 3, 2093
69Funaki, T.; Inokuma, S.; Ida, H.; Yonekura, T.; Nakamura, Y.; Nishimura, J. Tetrahedron Lett. 2004, 45, 2393.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 51
There are many examples of molecules in which benzene rings are forced out of planarity, including 7-circulene (40), 70 9, 8 -diphenyltetrabenz [a, c,h,j\ anthracene (41), 71 and 4272 (see also p. 230). These have been called tormented aromatic systems.73 The “record” for twisting an aromatic 71-electron system appears to be 9,10,ll,12,13,14,15,16-octaphenyldibenzo[a,c]naphthacene (43), 74 which has an end-to-end twist of 105°. This is >1.5 times as great as that observed in any previous polyaromatic hydrocarbon.
Perchlorotriphenylene has been reported in the literature and said to show severe molecular twisting, however, recent work suggests this molecule has not actually been isolated with perchlorofluorene-9-spirocyclohexa-2/,5/-diene being formed instead.75 The X-ray structure of the linear [3]phenylene (benzo[3,4]cyclobuta- [l,2-b]biphenylene, 44) has been obtained, and it shows a relatively large degree of bond alternation while the center distorts to a cyclic bis-allyl frame.76
70Yamamoto, K.; Harada, T.; Okamoto, Y.; Chikamatsu, H.; Nakazaki, M.; Kai, Y.; Nakao, T.; Tanaka, M.; Harada, S.; Kasai, N. J. Am. Chem. Soc. 1988, 110, 3578.
7 Pascal, Jr., R.A.; McMillan, W.D.; Van Engen, D.; Eason, R.G. J. Am. Chem. Soc. 1987, 109, 4660.
72Chance, J.M.; Kahr, B.; Buda, A.B.; Siegel, J.S. J. Am. Chem. Soc. 1989 , 111, 5940.
73Pascal, Jr., R.A. Pure Appl. Chem. 1993, 65, 105.
74Qiao, X.; Ho, D.M.; Pascal Jr., R.A. Angew. Chem. Int. Ed., 1997, 36, 1531.
75Campbell, M.S.; Humphries, R.E.; Munn, N.M. J. Org. Chem. 1992, 57, 641.
76Schleifenbaum, A.; Feeder, N.; Vollhardt, K.P.C. Tetrahedron Lett. 2001, 42, 7329.
52 DELOCALIZED CHEMICAL BONDING
It is also possible to fuse strained rings on benzene, which induces great strain on the benzene ring. In 45, the benzene ring is compressed by the saturated environ¬ ment of the tetrahydropyran units. In this case, the strain leads to distortion of the benzene ring in 45 into a boat conformation.77 Benzocyclopropene (46) and benzo- cyclobutene (47) are also molecules where the small annellated ring induces great strain on the benzene ring. In these cases, bonds of annellation and those adjacent to it are strained.
47
Strain-induced bond localization was introduced in 1930 by Mills and Nixon78 and is commonly referred to as the Mills-Nixon effect (see Chapter 11, p. 677). Ortho-fused aromatic compounds, such as 46, are known as cycloproparenes79 and are highly strained. Cyclopropabenzene (46) is a stable molecule with a strain energy of 68 kcal mol-1 (284.5 kJ mol-1).80 and the annellated bond is always the shortest, although in 47 the adjacent bond is the shortest.81 In cycloproparenes, there is the expectation of partial aromatic bond localization, with bond length alternation in the aromatic ring.82 When the bridging units are saturated, the ben¬ zene ring current is essentially unchanged, but annelation with one or more cyclo- butadieno units disrupts the benzene ring current.83 The chemistry of the cycloproparenes is dominated by the influence of the high strain energy. When fused to a benzene ring, the bicyclo[ 1.1.0] butane unit also leads to strain-induced localization of aromatic 7t-bonds.84
pn-dn Bonding: Ylids
We have mentioned (p. 10) that, in general, atoms of the second row of the Periodic table do not form stable double bonds of the type discussed in Chapter 1
77Hall, G.G J. Chem. Soc. Perkin Trans. 2 1993, 1491.
78Mills, W. H.; Nixon, I.G. J. Chem. Soc. 1930, 2510.
79Halton, B. Chem. Rev. 2003, 103, 1327; Halton, B. Chem. Rev. 1989, 89, 1 161, and reviews cited therein. 80Billups, W.E.; Chow, W.Y.; Leavell, K.H.; Lewis, E.S.; Margrave, J.L.; Sass, R.L.; Shieh, J.J.; Wemess, P.G.; Wood, J.L. J. Am. Chem. Soc. 1973, 95, 7878.; Apeloig, Y.; Arad, D. J. Am. Chem. Soc. 1986, 108, 3241. 81Boese, R.; Blaser, D.; Billups, W.E.; Haley, M.M.; Maulitz, A.H.; Mohler, D.L.; Vollhardt, K.P.C. Angew. Chem. Int. Ed., 1994, 33, 313.
82Halton, B. Pure Appl. Chem. 1990, 62, 541; Stanger, A. J. Am. Chem. Soc. 1998, 120, 12034; Maksic, Z.B.; Eckert-Maksic, M.; Pfeifer, K.-H. J. Mol. Struct. 1993, 300, 445; Mo, M.; Yanez, M.; Eckert- Maksic, M.; Maksic, Z.B. J. Org. Chem. 1995, 60, 1638; Eckert-Maksic, M.; Glasovac, Z.; Maksic, Z.B.; Zrinski, I. J. Mol. Struct. (THEOCHEM) 1996, 366, 173; Baldridge, K.K.; Siegel, J.S. J. Am. Chem. Soc. 1992, 114, 9583.
83Soncini, A.; Havenith, R.W.A.; Fowler, P.W.; Jenneskens, L.W.; Steiner, E. J. Org. Chem. 2002, 67, 4753 84Cohrs, C.; Reuchlein, H.; Musch, P.W.; Selinka, C.; Walfort, B.; Stalke, D.; Christl, M. Eur. J. Org. Chem. 2003, 901.
CHAPTER 2
DELOCALIZED CHEMICAL BONDING 53
(n bonds formed by overlap of parallel p orbitals). However, there is another type of double bond that is particularly common for the second-row atoms, sulfur and phos¬ phorus. For example, such a double bond is found in the compound H2S03,
H
ii
o
^0.©^0. H S
i
°0
H
as written on the left. Like an ordinary double bond, this double bond contains one s orbital, but the second orbital is not a n orbital formed by overlap of half-filled p orbitals; instead it is formed by overlap of a filled p orbital from the oxygen with an empty d orbital from the sulfur. It is called a pn-dn orbital*5 Note that we can represent this molecule by two canonical forms, but the bond is nevertheless loca¬ lized, despite the resonance. Some other examples of pn-dn bonding are Nitrogen
R
R— P=0 /
R
R
\© o
R— P— Cr /
R
Phosphine oxides
O
ii
R-S-R
ii
O
o
R— S+2-R
I
O0
Sulfones
H
H-P = 0 HO
H-P-O®
HO
Hypophorphoros acid
O
ii
/S.
R R
© R
Sulfoxides
analogs are known for some of these phosphorus compounds, but they are less stable because the resonance is lacking. For example, amine oxides, analogs of phosphine oxides, can only be written R3N+— O-. The pn-dn canonical form is impossible since nitrogen is limited to eight outer-shell electrons.
In all the examples given above, the atom that donates the electron pair is oxygen and, indeed, oxygen is the most common such atom. But in another important class of compounds, called ylids, this atom is carbon.86 There are three main types of ylids phosphorus, nitrogen, and sulfur ylids, although
85For a monograph, see Kwart, H.; King, K. d-Orbitals in the Chemistry of Silicon, Phosphorus, and Sulfur, Springer, NY, 1977.
86For a monograph, see Johnson, A.W. Ylid Chemistry ; Academic Press, NY, 1966. For reviews, see Morris, D.G., Surv. Prog. Chem. 1983, 10, 189; Hudson, R.F. Chem. Br., 1971, 7, 287; Lowe, PA. Chem. Ind. ( London ) 1970, 1070. For a review on the formation of ylids from the reaction of carbenes and carbenoids with heteroatom lone pairs, see Padwa, A.; Hornbuckle, S.F. Chem. Rev. 1991, 91, 263. 87Although the phosphorus ylid shown has three R groups on the phosphorus atom, other phosphorus ylids are known where other atoms, for example, oxygen, replace one or more of these R groups. When the three groups are all alkyl or aryl, the phosphorus ylid is also called a phosphorane.
88For a review of nitrogen ylids, see Musker, W.K. Fortschr. Chem. Forsch. 1970, 14, 295.
89For a monograph on sulfur ylids, see Trost, B.M.; Melvin Jr., L.S. Sulfur Ylids ; Academic Press, NY, 1975. For reviews, see Fava, A, in Bernardi, F.; Csizmadia, I.G.; Mangini, A. Organic Sulfur Chemistry, Elsevier, NY, 1985, pp. 299-354; Belkin, Yu.V.; Polezhaeva, N.A. Russ. Chem. Rev. 1981, 50, 481; Block, E. in Stirling, C.J.M. The Chemistry of the Sulphonium Group, part 2, Wiley, NY, 1981, pp. 680-702; Block, E. Reactions of Organosulfur Compounds', Academic Press, NY, 1978, pp. 91-127.
54 DELOCALIZED CHEMICAL BONDING
arsenic,90 selenium, and so on, ylids are also known. Ylids may be defined as com¬ pounds in which a positively charged atom from group 15 or 16 of the Periodic table is connected to a carbon atom carrying an unshared pair of electrons. Because of pn-dn bonding, two canonical forms can be written for phosphorus and sulfur, but there is only one for nitrogen ylids. Phosphorus ylids are much more stable than nitrogen ylids (see also p. 810). Sulfur ylids also have a low stability.
*3
1
II
n
"t
Rs© 0
— ► R — P — CR2
R
's=cr2
Rs© 0
— s-cr2
R © © r-n-cr2 /
R
R
/
R
/
R
R
Phosphorus ylids
Sulfur ylids
Nitrogen ylids
In almost all compounds that have pn-dn bonds, the central atom is connected to four atoms or three atoms and an unshared pair and the bonding is approximately tetrahedral. The pn-dn bond, therefore, does not greatly change the geometry of the molecule in contrast to the normal n bond, which changes an atom from tetrahedral to trigonal. Calculations show that nonstabilized phosphonium ylids have nonplanar ylidic carbon geometries whereas stabilized ylids have planar ylidic carbons.91
AROMATICITY92
In the nineteenth century, it was recognized that aromatic compounds93 differ greatly from unsaturated aliphatic compounds,94 but for many years chemists
90For reviews of arsenic ylids, see Lloyd, D.; Gosney, I.; Ormiston, R.A. Chem. Soc. Rev. 1987, 16, 45; Yaozeng, H.; Yanchang, S. Adv. Organomet. Chem. 1982, 20, 115.
91Bachrach, S.M. J. Org. Chem. 1992, 57, 4367.
92Krygowski, T.M.; Cyranski, M.K.; Czarnocki, Z.; Hafelinger, G.; Katritzky, A.R. Tetrahedron 2000, 56, 1783; Simkin, B.Ya.; Minkin, V.I.; Glukhovtsev, M.N., in Advances in Heterocyclic Chemistry, Vol. 56, Katritzky, A.R., Ed., Academic Press, San Diego, 1993, pp 303-428; Krygowski, T.M.; Cyranski, M.K. Chem. Rev. 2001, 101, 1385; Katritzky, A.R.; Jug, K.; Oniciu, D.C. Chem. Rev. 2001, 101, 1421; Katritzky, A.R.; Karelson, M.; Wells, A.P. J. Org. Chem. 1996, 61, 1619. See also Cyranski, M.K.; Krygowski, T.M.; Katritzky, A.R.; Schley er, P.v.R. J. Org. Chem. 2002, 67, 1333.
93For books on Aromaticity, see Lloyd, D. The Chemistry of Conjugated Cyclic Compounds, Wiley, NY, 1989; Non-Benzenoid Conjugated Carbocyclic Compounds, Elsevier, NY, 1984; Garratt, P.J. Aromaticity, Wiley, NY, 1986; Balaban, A.T.; Banciu, M.; Ciorba, V. Annulenes, Benzo-, Hetero-, Homo-Derivatives and their Valence Isomers, 3 vols., CRC Press, Boca Raton, FL 1987; Badger, G.M. Aromatic Character and Aromaticity, Cambridge University Press, Cambridge, 1969; Snyder, J.P. Nonbenzenoid Aromatics, 2 vols., Academic Press, NY, 1969-1971; Bergmann, E.D.; Pullman, B. Aromaticity, Pseudo-Aromaticity, and Anti- Aromaticity, Israel Academy of Sciences and Humanities, Jerusalem, 1971; Aromaticity; Chem. Soc. Spec. Pub. No. 21, 1967. For reviews, see Gorelik, M.V. Russ. Chem. Rev. 1990, 59, 116; Stevenson, G.R. Mol. Struct. Energ., 1986, 3, 57; Sondheimer, F. Chimia, 1974, 28, 163; Cresp, T.M.; Sargent, M.V. Essays Chem. 1972, 4, 91; Figeys, H.P. Top. Carbocyclic Chem. 1969, 1, 269; Garratt, P.J.; Sargent, M.V. papers in, Top. Curr. Chem. 1990, 153 and Pure Appl. Chem. 1980, 52, 1397.
94For an account of the early history of Aromaticity, see Snyder, J.P., in Snyder, J.P. Nonbenzenoid Aromatics, Vol. 1, Academic Press, NY, 1971, pp. 1-31. See also Balaban, A.T. Pure Appl. Chem. 1980, 52, 1409.
CHAPTER 2
AROMATICITY 55
were hard pressed to arrive at a mutually satisfactory definition of aromatic char¬ acter.95 Qualitatively, there has never been real disagreement. Definitions have taken the form that aromatic compounds are characterized by a special stability and that they undergo substitution reactions more easily than addition reactions. The difficulty arises because these definitions are vague and not easy to apply in borderline cases. Definitions of aromaticity must encompass molecules ranging form polycyclic conjugated hydrocarbons,96 to heterocyclic compounds97 of various ring sizes, to reactive intermediates. In 1925 Armit and Robinson,98 recog¬ nized that the aromatic properties of the benzene ring are related to the presence of a closed loop of electrons, the aromatic sextet (aromatic compounds are thus the arch examples of delocalized bonding), but it still was not easy to determine whether rings other than the benzene ring possessed such a loop. With the advent of magnetic techniques, most notably NMR, it is possible to determine experimen¬ tally whether or not a compound has a closed ring of electrons; aromaticity can now be defined as the ability to sustain an induced ring current. A compound with this ability is called diatropic. Although this definition also has its flaws,99 it is the one most commonly accepted today. There are several methods of determining whether a compound can sustain a ring current, but the most important one is based on NMR chemical shifts.100 In order to understand this, it is necessary to remember that, as a general rule, the value of the chemical shift of a proton in an NMR spectrum depends on the electron density of its bond; the greater the density of the electron cloud surrounding or partially surrounding a proton, the more upfield is its chemical shift (a lower value of 5). However, this rule has several exceptions; one is for pro¬ tons in the vicinity of an aromatic ring. When an external magnetic field is imposed upon an aromatic ring (as in an NMR instrument), the closed loop of aromatic elec¬ trons circulates in a diamagnetic ring current, which sends out a field of its own. As can be seen in Fig. 2.6, this induced field curves around and in the area of the proton is parallel to the external field, so the field “seen” by the aromatic protons is greater than it would have been in the absence of the diamagnetic ring current. The protons are moved downfield (to higher 5) compared to where they would be if electron
95For a review of the criteria used to define aromatic character, see Jones, A.J. Pure Appl. Chem. 1968, 18, 253. For methods of assigning Aromaticity, see Jug, K.; Koster, A.M. J. Phys. Org. Chem. 1991, 4, 163; Zhou, Z.; Parr, R.G. J. Am. Chem. Soc. 1989, 111, 7371; Katritzky, A.R.; Barczynski, P.; Musumarra, G.; Pisano, D.; Szafran, M. J. Am. Chem. Soc. 1989, 111, 7; Schaad, L.J.; Hess, Jr., B.A. J. Am. Chem. Soc. 1972, 94, 3068, J. Chem. Educ. 1974, 51, 640. See also, Bird, C.W. Tetrahedron 1985, 41, 1409; 1986, 42, 89; 1987, 43, 4725.
96Randic, M. Chem. Rev. 2003, 103, 3449.
97Balaban, A.T.; Oniciu, D.C.; Katritzky, A.R. Chem. Rev. 2004, 104, 2111.
98Armit, J.W.; Robinson; R. J. Chem. Soc. 1925, 127, 1604.
"Jones, A.J. Pure Appl. Chem. 1968, 18, 253, pp. 266-274; Mallion, R.B. Pure Appl. Chem. 1980, 52, 1541. Also see, Schleyer, P.v.R.; Jiao, H. Pure Appl. Chem. 1996, 68, 209.
100For a review of NMR and other magnetic properties with respect to aromaticity, see Haddon, R.C.; Haddon, V.R.; Jackman, L.M. Fortschr. Chem. Forsch. 1971, 16, 103. For an example of a magentic method other than NMR, see Dauben Jr., H.J.; Wilson, J.D.; Laity, J.L., in Snyder, J.P. Nonbenzenoid Aromatics, Vol. 2, Academic Press, NY, 1971, pp. 167-206.
56 DELOCALIZED CHEMICAL BONDING
density were the only factor. Thus ordinary alkene hydrogens are found at ~5-6 5, while the hydrogens of benzene rings are located at ~7-8 5. However, if there
(^^)ch2)10
48A
were protons located above or within the ring, they would be subjected to a decreased field and should appear at lower 5 values than normal CH2 groups (nor¬ mal 5 for CH2 is ~l-2). The nmr spectrum of [10]paracyclophane (48A) showed that this was indeed the case101 and that the CH2 peaks were shifted to lower 5 the closer they were to the middle of the chain. Examination of 48B shows that a por¬ tion of the methylene chain is positioned directly over the benzene ring, making it subject to the anisotropy shift mentioned above.
It follows that aromaticity can be determined from an NMR spectrum. If the pro¬ tons attached to the ring are shifted downfield from the normal alkene region, we can conclude that the molecule is diatropic, and hence aromatic. In addition, if the compound has protons above or within the ring (we shall see an example of the latter on p. 90), then if the compound is diatropic, these will be shifted upfield.
101Waugh, J.S.; Fessenden, R.W. J. Am. Chem. Soc. 1957, 79, 846. See also, Shapiro, B.L.; Gattuso, M.J.; Sullivan, G.R. Tetrahedron Lett. 1971, 223; Pascal, Jr., R.A.; Winans, C.G.; Van Engen, D. J. Am. Chem. Soc. 1989, 111, 3007.
CHAPTER 2
AROMATICITY 57
One drawback to this method is that it cannot be applied to compounds that have
no protons in either category, for example, the dianion of squaric acid (p. 92).
Unfortunately, 13C NMR is of no help here, since these spectra do not show ring 102 ^
currents.
Antiaromatic systems exhibit a paramagnetic ring current,103 which causes pro¬ tons on the outside of the ring to be shifted upfield while any inner protons are shifted downfield , in sharp contrast to a diamagnetic ring current, which causes shifts in the opposite directions. Compounds that sustain a paramagnetic ring cur¬ rent are called paratropic ; and are prevalent in four- and eight-electron systems. As with aromaticity, we expect that antiaromaticity will be at a maximum when the molecule is planar and when bond distances are equal. The diamagnetic and para¬ magnetic effects of the ring currents associated with aromatic and antiaromatic compounds (i.e., shielding and deshielding of nuclei) can be measured by a simple and efficient criterion known as nucleus independent chemical shift (NICS).104 The aromatic-antiaromatic ring currents reflect the extra 71-effects that the mole¬ cules experience. The unique near zero value of NICS at the cyclobutadiene ring center is due to cancelation by large and opposite anistropic components.105
There are at least four theoretical models for aromaticity, which have recently been compared and evaluated for predictive ability.106 The Hess-Schaad model 107 is good for predicting aromatic stability of benzenoid hydrocarbons, but does not predict reactivity. The Herndon model 108 is also good for predicting aromatic stability, but is unreliable for benzenoidicity and does not predict reactivity. The conjugated-circuit model 109 is very good for predicting aromatic stability, but not reactivity, and the hardness model 110 is best for predicting kinetic stability. Delocalization energy of 7i-electrons has also been used as an index for aromaticity in polycyclic aromatic hydrocarbons.111 The claims for linear relationships between aromaticity and energetics, geometries, and magnetic criteria were said to be invalid for any representative set of heteroaromatics in which the number of heteroatoms varies.112
It should be emphasized that the old and new definitions of aromaticity are not necessarily parallel. If a compound is diatropic and therefore aromatic under the
102For a review of 13C NMR spectra of aromatic compounds, see Gunther, H.; Schmickler, H. Pure Appl. Chem. 1975, 44, 807.
103Pople, J.A.; Untch, K.G. J. Am. Chem. Soc. 1966, 88, 4811; Longuet-Higgins, H.C. in Garratt, RJ. Aromaticity, Wiley, NY, 1986, pp. 109-111.
104Schleyer, P.v.R.; Maerker, C.; Dransfeld, A.; Jiao, H.; Hommes, N.J.R.v.E. J. Am. Chem. Soc. 1996, 118, 6317.
105Schleyer, P.v.R.; Manoharan, M.; Wang, Z.-X.; Kiran, B.; Jiao, H.; Puchta, R.; Hommes, N.J.R.v.E. Org. Lett. 2001, 3, 2465
106Plavic, D.; Babic, D.; Nikolic, S.; Trinajstic, N. Gazz. Chim. Ital., 1993, 123, 243.
107Hess, Jr., B.A.; Schaad, L.J. J. Am. Chem. Soc. 1971, 93, 305.
108Herndon, W.C. Isr. J. Chem. 1980, 20, 270.
109Randic,M. Chem.Phys.Lett. 1976, 38, 68.
110Zhou, Z.; Parr, R.G. J. Am. Chem. Soc. 1989, 111, 7371; Zhou, Z.; Navangul, H.V. J. Phys. Org. Chem. 1990, 3, 784.
niBehrens, S.; Koster, A.M.; Jug, K. J. Org. Chem. 1994, 59, 2546.
112Katritzky, A.R.; Karelson, M.; Sild, S.; Krygowski, T.M.; Jug, K. J. Org. Chem. 1998, 63, 5228.
58 DELOCALIZED CHEMICAL BONDING
new definition, it is more stable than the canonical form of lowest energy, but this does not mean that it will be stable to air, light, or common reagents, since this stability is determined not by the resonance energy, but by the difference in free energy between the molecule and the transition states for the reactions involved; and these differences may be quite small, even if the resonance energy is large. A unified theory has been developed that relates ring currents, resonance energies, and aromatic character.113 Note that aromaticity varies in magnitude relatively and sometimes absolutely with the molecular environment, which includes the polarity of the medium.114
The vast majority of aromatic compounds have a closed loop of six electrons in a ring (the aromatic sextet), and we consider these compounds first.115 Note that a “formula Periodic table” for the benzenoid poly aromatic hydrocarbons has been developed.116
Six-Membered Rings
Not only is the benzene ring aromatic, but so are many heterocyclic analogs in which one or more heteroatoms replace carbon in the ring.117 When nitrogen is the heteroatom, little difference is made in the sextet and the unshared pair of the nitrogen does not participate in the aromaticity. Therefore, derivatives such as A-oxides or pyridinium ions are still aromatic. However, for nitrogen hetero¬ cycles there are more significant canonical forms (e.g., 49) than for benzene. Where oxygen or sulfur is the heteroatom, it must be present in its ionic form (50) in order to possess the valence of 3 that participation in such a system demands. Thus, pyran (51) is not aromatic, but the pyrylium ion (49) is.118
49 50 51
113Haddon, R.C. J. Am. Chem. Soc. 1979, 101, 1722; Haddon, R.C.; Fukunaga, T. Tetrahedron Lett. 1980, 21, 1191.
1 14Katritzky, A.R.; Karelson, M.; Wells, A.P. J. Org. Chem. 1996, 61, 1619.
115Values of molecular-orbital energies for many aromatic systems, calculated by the HMO method, are given in Coulson, C.A.; Streitwieser, Jr., A. A Dictonary of n Electron Calculations, W.H. Freeman, San Francisco, 1965. Values calculated by a variation of the SCF method are given by Dewar, M.J.S.; Trinajstic, N. Collect. Czech. Chem. Commun. 1970, 35, 3136, 3484.
116Dias, J.R. Chem. Br. 1994, 384.
117For reviews of Aromaticity of heterocycles, see Katritzky, A.R.; Karelson, M.; Malhotra, N. Heterocycles 1991, 32, 127.
118For a review of pyrylium salts, see Balaban, A.T.; Schroth, W.; Fischer, G. Adv. Heterocycl. Chem. 1969,10,241.
CHAPTER 2
AROMATICITY 59
In systems of fused six-membered aromatic rings,119 the principal canonical forms are usually not all equivalent. Compound 52 has a central double bond and is thus different from the other two canonical forms of naphthalene, which are equivalent to each other.120 For naphthalene, these are the only forms that can be drawn
8 1
52
without consideration of Dewar forms or those with charge separation.121 If we assume that the three forms contribute equally, the 1,2 bond has more double¬ bond character than the 2,3 bond. Molecular-orbital calculations show bond orders of 1.724 and 1.603, respectively, (cf. benzene, 1.667). In agreement with these pre¬ dictions, the 1,2 and 2,3 bond distances are 1.36 and 1.415 A, respectively,122 and ozone preferentially attacks the 1,2 bond.123 This nonequivalency of bonds, called partial bond fixation, 124 is found in nearly all fused aromatic systems. In phenan- threne, where the 9,10 bond is a single bond in only one of five forms (53), bond fixation becomes extreme and this bond is readily attacked by many reagents:125 It has been observed that increased steric crowding leads to an increase in Dewar-benzene type structures.126
6
53
119For books on this subject, see Gutman, I.; Cyvin, SJ. Introduction to the Theory of Benzenoid Hydrocarbons, Springer, NY, 1989; Dias, J.R. Handbook of Polycyclic Hydrocarbons, Part A: Benzenoid Hydrocarbons, Elsevier, NY, 1987; Clar, E. Polycyclic Hydrocarbons, 2 vols., Academic Press, NY, 1964. For a “Periodic table” that systematizes fused aromatic hydrocarbons, see Dias, J.R. Acc. Chem. Res. 1985, 18, 241; Top. Curr. Chem. 1990, 253, 123; J. Phys. Org. Chem. 1990, 3, 765.
120As the size of a given fused ring system increases, it becomes more difficult to draw all the canonical forms. For discussions of methods for doing this, see Herndon, W.C. J. Chem. Educ. 1974, 51, 10; Cyvin, S.J.; Cyvin, B.N.; Brunvoll, J.; Chen, R. Monatsh. Chem. 1989, 120, 833; Fuji, Z.; Xiaofeng, G.; Rongsi, C. Top. Curr. Chem. 1990, 153, 181; Wenchen, H.; Wenjie, H. Top. Curr. Chem. 1990, 153, 195; Sheng, R. Top. Curr. Chem. 1990, 153, 211; Rongsi, C.; Cyvin, S.J.; Cyvin, B.N.; Brunvoll, J.; Klein, D.J. Top. Curr. Chem. 1990, 153, 227, and references cited in these papers. For a monograph, see Cyvin, S.J.; Gutman, I. Kekul6 Structures in Benzenoid Hydrocarbons; Springer, NY, 1988.
121For a modern valence bond description of naphthalene, see Sironi, M.; Cooper, D.L.; Gerratt, J.; Raimondi, M. J. Chem. Soc. Chem. Commun. 1989, 675.
122Cruickshank, D.W.J. Tetrahedron 1962, 17, 155.
123Kooyman, E.C. Reel. Trav. Chim. Pays-Bas, 1947, 66, 201.
124For a review, see Efros, L.S. Russ. Chem. Rev. 1960, 29, 66.
125See also Lai, Y. J. Am. Chem. Soc. 1985, 107, 6678.
126Zhang, J.; Ho, D.M.; Pascal Jr., R.A. J. Am. Chem. Soc. 2001, 123, 10919.
60 DELOCALIZED CHEMICAL BONDING
In general, there is a good correlation between bond distances in fused aromatic compounds and bond orders. Another experimental quantity that correlates well with the bond order of a given bond in an aromatic system is the NMR coupling constant for coupling between the hydrogens on the two carbons of the bond.127
The resonance energies of fused systems increase as the number of principal canonical forms increases, as predicted by rule 6 (p. 47). 128 Thus, for benzene, naphthalene, anthracene, and phenanthrene, for which we can draw, respectively, two, three, four, and five principal canonical forms, the resonance energies are, respectively, 36, 61, 84, and 92kcalmol-1 (152, 255, 351, and 385 kJmol-1), cal¬ culated from heat-of-combustion data.129 Note that when phenanthrene, which has a total resonance energy of 92kcalmol-1 (385 kJ mol-1), loses the 9,10 bond by attack of a reagent, such as ozone or bromine, two complete benzene rings remain, each with 36 kcal mol-1 (152 kJ mol-1) that would be lost if benzene was similarly attacked. The fact that anthracene undergoes many reactions across the 9,10 posi¬ tions can be explained in a similar manner. Resonance energies for fused systems can be estimated by counting canonical forms.130
8 9 1
5 10 4
Anthracene
Not all fused systems can be fully aromatic. Thus for phenalene (54) there is no way double bonds can be distributed so that each carbon has one single and one double bond.131 However, phenalene is acidic and reacts with potassium meth- oxide to give the corresponding anion (55), which is completely aromatic. So are the corresponding radical and cation, in which the resonance energies are the same (see p. 68). 132
H Br
H Br
etc.
55
127Jonathan, N.; Gordon, S.; Dailey, B.P. J. Chem. Phys. 1962, 36, 2443; Cooper, M.A.; Manatt, S.L. J. Am. Chem. Soc. 1969, 91, 6325.
128See Herndon, W.C.; Ellzey Jr., M.L. J. Am. Chem. Soc. 1974, 96, 6631.
129Wheland, G.W. Resonance in Organic Chemistry, Wiley, NY, 1955, p. 98.
130Swinborne-Sheldrake, R.; Herndon, W.C. Tetrahedron Lett. 1975, 755.
131For reviews of phenalenes, see Murata, I. Top. Nonbenzenoid Aromat. Chem. 1973, 1, 159; Reid, D.H. Q. Rev. Chem. Soc. 1965, 19, 274.
132Pettit, R. J. Am. Chem. Soc. 1960, 82, 1972.
CHAPTER 2
AROMATICITY 61
Molecules that contain fused rings, such as phenanthrene or anthracene, are generally referred to as linear or angular polyacenes. In a fused system, there are not six electrons for each ring.133 In naphthalene, if one ring is to have six, the other must have only four. One way to explain the greater reactivity of the ring system of naphthalene compared with benzene is to regard one of the naphthalene rings as aromatic and the other as a butadiene system.134 This effect can become extreme, as in the case of triphenylene.135 For this compound, there are eight canonical forms like 56, in which none of the three bonds marked a is a double bond and only one form (57) in which at least one of them is double. Thus the molecule behaves as if the 1 8 electrons were distributed so as to give each of the outer rings a sextet, while the middle ring is “empty.” Since none of the outer rings need share
any electrons with an adjacent ring, they are as stable as benzene; triphenylene, unlike most fused aromatic hydrocarbons, does not dissolve in concentrated sulfuric acid and has a low reactivity.136 This phenomenon, whereby some rings in fused sys¬ tems give up part of their aromaticity to adjacent rings, is called annellation and can be demonstrated by UV spectra119 as well as reactivities. In general, an increase of size of both linear and angular polyacenes is associated with a substantial edecrease in their aromaticity, with a greater decrease for the linear polyacenes.137
A six-membered ring with a circle is often used to indicate an aromatic system, and this will be used from time to time. Kekule structures, those having the C=C units rather than a circle, are used most often in this book. Note that one circle can be used for benzene, but it would be misleading to use two circles for naphthalene, for example, because that would imply 12 aromatic electrons, although naphthalene has only 10. 138
Five-, Seven-, and Eight-Membered Rings
Aromatic sextets can also be present in five- and seven-membered rings. If a five- membered ring has two double bonds, and the fifth atom possesses an unshared pair
133For discussions of how the electrons in fused aromatic systems interact to form An + 2 systems, see Glidewell, C.; Lloyd, D. Tetrahedron 1984, 40, 4455, J. Chem. Educ. 1986, 63, 306; Hosoya, H. Top. Curr. Chem. 1990, 153, 255.
134Meredith, C.C.; Wright, G.F. Can. J. Chem. 1960, 38, 1177.
135For a review of triphenylenes, see Buess, C.M.; Lawson, D.D. Chem. Rev. 1960, 60, 313.
136Clar, E.; Zander, M. J. Chem. Soc. 1958, 1861.
137Cyran ski, M.K.; St^pien, B.T.; Krygowski, T.M. Tetrahedron 2000, 56, 9663.
138See Belloli, R. J. Chem. Educ. 1983, 60, 190.
62 DELOCALIZED CHEMICAL BONDING
Fig. 2.7. Overlap of five p orbitals in molecules such as pyrrole, thiophene, and the cyclopentadienide ion
of electrons, the ring has five p orbitals that can overlap to create five new orbitals: three bonding and two antibonding (Fig. 2.7). There are six electrons for these orbitals: the four p orbitals of the double bonds each contribute one and the filled orbital contributes the other two. The six electrons occupy the bonding orbitals and
9 Q Q
H
Pyrrole Thiophene Furan
constitute an aromatic sextet. The heterocyclic compounds pyrrole, thiophene, and furan are the most important examples of this kind of aromaticity, although furan has a lower degree of aromaticity than the other two.139 Resonance energies for these three compounds are, respectively, 21, 29, and 16 kcal mol-1 (88, 121, and 67 kJ mol-1).140 The aromaticity can also be shown by canonical forms, for exam¬ ple, for pyrrole:
o
-
^ Oe —
m -
©
- a
N
1
N©
©N
N©
N©
1
H
A
H
H
H
H
139The order of aromaticity of these compounds is benzene > thiophene > pyrrole > furan, as calculated by an Aromaticity index based on bond distance measurements. This index has been calculated for five- and six-membered monocyclic and bicyclic heterocycles: Bird, C.W. Tetrahedron 1985, 41, 1409; 1986, 42, 89; 1987, 43, 4725.
140Wheland, G.W. Resonance in Organic Chemistry, Wiley, NY, 1955, p 99. See also, Calderbank, K.E.; Calvert, R.L.; Lukins, P.B.; Ritchie, G.L.D. Aust. J. Chem. 1981, 34, 1835.
CHAPTER 2
AROMATICITY 63
In contrast to pyridine, the unshared pair in canonical structure A in pyrrole is needed for the aromatic sextet. This is why pyrrole is a much weaker base than pyridine.
The fifth atom may be carbon if it has an unshared pair. Cyclopentadiene has unexpected acidic properties (pKa& 16) since on loss of a proton, the resulting car- banion is greatly stabilized by resonance although it is quite reactive. The cyclopen- tadienide ion is usually represented as in 58. Resonance in this ion is greater than in pyrrole, thiophene, and furan, since all five forms are equivalent. The resonance energy for 58 has been estimated to be 24-27 kcalmol-1 (100-113 kJ mol-1).141
That all five carbons are equivalent has been demonstrated by labeling the starting compound with 14C and finding all positions equally labeled when cyclopentadiene was regenerated142 As expected for an aromatic system, the cyclopentadienide ion is diatropic143 and aromatic substitutions on it have been successfully carried out.144 Average bond order has been proposed as a parameter to evaluate the aro¬ maticity of these rings, but there is poor correlation with non-aromatic and antiaro¬ matic systems.145 A model that relies on calculating relative aromaticity from appropriate molecular fragments has also been developed.146 Bird devised the aromatic index (/A, or aromaticity index),147 which is a statistical evaluation of the extent of ring bond order, and this has been used as a criterion of aromaticity. Another bond-order index was proposed by Pozharskii,148 which goes back to the work of Fringuelli and co-workers.149 Absolute hardness (see p. 377), cal¬ culated from molecular refractions for a range of aromatic and heteroaromatic com¬ pounds, shows good linear correlation with aromaticity.150 Indene and fluorene are also acidic (p Ka « 20 and 23, respectively), but less so than cyclopentadiene, since annellation causes the electrons to be less available to the five-membered ring. On the other hand, the acidity of l,2,3,4,5-pentakis(trifhioromethyl)cyclopen- tadiene (59) is greater than that of nitric acid,151 because of the electron-
141Bordwell, F.G.; Drucker, G.E.; Fried, H.E. J. Org. Chem. 1981, 46, 632.
142Tkachuk, R.; Lee, C.C. Can. J. Chem. 1959, 37, 1644.
143Bradamante, S.; Marchesini, A.; Pagani, G. Tetrahedron Lett. 1971, 4621.
144Webster, O.W. J. Org. Chem. 1967, 32, 39; Rybinskaya, M.I.; Korneva, L.M. Russ. Chem. Rev. 1971, 40, 247.
145Jursic, B.S. J. Heterocycl. Chem. 1997, 34, 1387.
146Hosmane, R.S.; Liebman, J.F. Tetrahedron Lett. 1992, 33, 2303.
147Bird, C.W. Tetrahedron 1985, 41, 1409; Tetrahedron 1992, 48, 335; Tetrahedron 1996, 52, 9945. 148Pozharskii, A.F. Khimiya Geterotsikl Soedin 1985, 867.
149Fringuelli, F. Marino, G.; Taticchi, A.; Grandolini, G. J. Chem. Soc. Perkin Trans. 2 1974, 332. 150Bird, C.W. Tetrahedron 1997, 53, 3319; Tetrahedron 1998, 54, 4641.
151Laganis, E.D.; Lemal, D.M. J. Am. Chem. Soc. 1980, 102, 6633.
64 DELOCALIZED CHEMICAL BONDING
withdrawing effects of the trifluoromethyl groups (see p. 381). Modifications of the Bird and Pozharskii systems have been introduced that are particularly useful for five-membered ring heterocycles.152 Recent work introduced a new local aromati¬ city measure, defined as the mean of Bader’s electron delocalization index (DI)153 of para-related carbon atoms in six-membered rings.154
Indene
Fluorene
As seen above, acidity of compounds can be used to study the aromatic character of the resulting conjugate base. In sharp contrast to cyclopentadiene (see p. 63) is cycloheptatriene (60), which has no unusual acidity. This would be hard to explain without the aromatic sextet theory, since, on the basis of resonance forms or a simple
60 61 62
consideration of orbital overlaps, 61 should be as stable as the cyclopentadienyl anion (58). While 61 has been prepared in solution,155 it is less stable than 58 and far less stable than 62, in which 60 has lost not a proton, but a hydride ion. The six double-bond electrons of 62 overlap with the empty orbital on the seventh carbon and there is a sextet of electrons covering seven carbon atoms. The cyclo- heptatrienyl cations (known as the tropylium ion , 62) is quite stable.156 Tropylium bromide (63), which could be completely covalent if the electrons of the bromine were sufficiently attracted to the ring, is actually an ionic compound:157 Many sub¬ stituted tropylium ions have been prepared to probe the aromaticity, structure, and reactivity of such systems.158 Just as with 58, the equivalence of the carbons
mKotelevskii, S.I.; Prezhdo, O.V. Tetahedron 2001, 57, 5715.
153See Bader, R.F.W. Atoms in Molecules: A Quantum Theory, Clarendon, Oxford, 1990; Bader, R.F.W. Acc. Chem. Res. 1985, 18, 9; Bader, R.F.W. Chem. Rev. 1991, 91, 893.
154Poater, J.; Fradera, X.; Duran, M.; Sola, M. Chem. Eur. J. 2003, 9, 400; 1113.
155Dauben Jr., H.J.; Rifi, M.R. J. Am. Chem. Soc. 1963, 85, 3041; also see Breslow, R.; Chang, H.W. J. Am Chem. Soc. 1965, 87, 2200.
156For reviews, see Pietra, F. Chem. Rev. 1973, 73, 293; Bertelli, D.J. Top. Nonbenzenoid Aromat. Chem. 1973, 1, 29; Kolomnikova, G.D.; Parnes, Z.N. Russ. Chem. Rev. 1967, 36, 735; Harmon, K.H., in Olah, G.A.; Schleyer, P.v.R. Carbonium Ions, Vol. 4, Wiley, NY, 1973, pp. 1579-1641.
157Doering, W. von E.; Knox, L.H. J. Am. Chem. Soc. 1954, 76, 3203.
158Pischel, U.; Abraham, W.; Schnabel, W.; Muller, U. Chem. Commun. 1997, 1383. See Komatsu, K.; Nishinaga, T.; Maekawa, N.; Kagayama, A.; Takeuchi, K. J. Org. Chem. 1994, 59, 7316 for a tropylium dication.
CHAPTER 2
AROMATICITY 65
in 62 has been demonstrated by isotopic labeling.159 The aromatic cycloheptatrie- nyl cations C7Me| and C7Ph| are known,160 although their coordination complexes with transition metals have been problematic, possibly because they assume a boat¬ like rather than a planar conformation161
Br
63
Another seven-membered ring that shows some aromatic character is tropone (64). This molecule would have an aromatic sextet if the two C=0 electrons stayed away from the ring and resided near the electronegative oxygen atom. In fact, tropones are stable compounds, and tropolones (65) are found in nature.162 How¬ ever, analyses of dipole moments, NMR spectra, and X-ray diffraction measure¬ ments show that tropones and tropolones display appreciable bond alternations.163
9
o
64
These molecules must be regarded as essentially non-aromatic, although with some aromatic character. Tropolones readily undergo aromatic substitution, emphasizing that the old and the new definitions of aromaticity are not always parallel. In sharp contrast to 64, cyclopentadienone (66) has been isolated only in an argon matrix <38 K.164 Above this temperature it dimerizes. Many earlier attempts to prepare it were unsuccessful.165 As in 64, the electronegative oxygen atom draws electron to itself, but in this case it leaves only four electrons and the molecule is
159Vorpin, M.E.; Kursanov, D.N.; Shemyakin, M.M.; Maimind, V.I.; Neiman, L.A. J. Gen. Chem. USSR 1959, 29, 3667.
160Takeuchi, K.; Yokomichi, Y.; Okamoto, K. Chem. Lett. 7977,1177; Battiste, M.A. J. Am. Chem. Soc. 1961,83,4101.
161Tamm, M.; DreBel, B.; Frohlich, R. J. Org. Chem. 2000, 65, 6795.
162For reviews of tropones and tropolones, see Pietra, F. Acc. Chem. Res. 1979, 12, 132; Nozoe, T. Pure Appl. Chem. 1971, 28, 239.
163Bertelli, D.J.; Andrews, Jr., T.G. J. Am. Chem. Soc. 1969, 91, 5280; Bertelli, D.J.; Andrews Jr., T.G.; Crews, P.O. J. Am. Chem. Soc. 1969, 91, 5286; Schaefer, J.P; Reed, F.L. J. Am. Chem. Soc. 1971, 93, 3902; Watkin, D.J.; Hamor, T.A. J. Chem. SOc. B 1971, 2167; Barrow, M.J.; Mills, O.S.; Filippini, G. J. Chem. Soc. Chem. Commun. 1973, 66.
164Maier, G.; Franz, F.H.; Fanz, K.; Reisenauer, H.P. Chem. Ber. 1985, 118, 3196.
165For a review of cyclopentadienone derivatives and of attempts to prepare the parent compound, see Ogliaruso, M.A.; Romanelli, M.G.; Becker, E.I. Chem. Rev. 1965, 65, 261.
66 DELOCALIZED CHEMICAL BONDING
unstable. Some derivatives of 66 have been prepared.130
66 Ferrocene
Another type of five-membered aromatic compound is the metallocenes (also called sandwich compounds ), in which two cyclopentadienide rings form a sand¬ wich around a metallic ion. The best known of these is ferrocene, where the r|5- coordination of the two cyclopentadienyl rings to iron is apparent in the 3D model 67. Other sandwich compounds have been prepared with Co, Ni, Cr, Ti, V, and many other metals.166 As a reminder (see p. 43), the r\ terminology refers to 71-donation of electrons to the metal (r|3 for 7t-allyl systems, r|6 for coordination to a benzene ring, etc.), and r|5 refers to donation of five 7i-electrons to the iron. Ferrocene is quite stable, subliming >100°C and unchanged at 400°C. The two rings rotate freely.167 Many aromatic substitutions have been carried out on metallocenes.168 Metallocenes containing two metal atoms and three cyclo¬ pentadienyl rings have also been prepared and are known as triple-decker sand¬ wiches.169 Even tetradecker, pentadecker, and hexadecker sandwiches have been reported.170
The bonding in ferrocene may be looked upon in simplified molecular- orbital terms as follows.171 Each of the cyclopentadienide rings has five molecular orbitals: three filled bonding and two empty antibonding orbitals (p. 62). The outer
166For a monograph on metallocenes, see Rosenblum, M. Chemistry of the Iron Group Metallocenes , Wiley, NY, 1965. For reviews, see Lukehart, C.M. Fundamental Transition Metal Organometallic Chemistry, Brooks/Cole, Monterey, CA, 1985, pp. 85-118; Lemenovskii, D.A.; Fedin, V.P. Russ. Chem. Rev. 1986, 55, 127; Sikora, D J.; Macomber, D.W.; Rausch, M.D. Adv. Organomet. Chem. 1986, 25, 317; Pauson, PL. Pure Appl. Chem. 1977, 49, 839; Nesmeyanov, A.N.; Kochetkova, N.S. Russ. Chem. Rev. 1974, 43, 710; Shul’pin, G.B.; Rybinskaya, M.I. Russ. Chem. Rev. 1974, 43, 716; Perevalova, E.G.; Nikitina, T.V. Organomet. React., 1972, 4, 163; Bublitz, D.E.; Rinehart Jr., K.L. Org. React., 1969, 17, 1; Leonova, E.V.; Kochetkova, N.S. Russ. Chem. Rev. 1973, 42, 278; Rausch, M.D. Pure Appl. Chem. 1972, 30, 523. For a bibliography of reviews on metallocenes, see Bruce, M.I. Adv. Organomet. Chem. 1972, 10, 273, pp. 322-325.
167For a discussion of the molecular structure, see Haaland, A. Acc. Chem. Res. 1979, 12, 415.
168For a review on aromatic substitution on ferrocenes, see Plesske, K. Angew. Chem. Int. Ed. 1962, 1, 312, 394.
169For a review, see Werner, H. Angew. Chem. Int. Ed. 1977, 16,1.
170See, for example, Siebert, W. Angew. Chem. Int. Ed. 1985, 24, 943.
171Rosenblum, M. Chemistry of the Iron Group Metallocnes, Wiley, NY, 1965, pp. 13-28; Coates, G.E.; Green, M.L.H.; Wade, K. Organometallic Compounds, 3rd ed., Vol. 2, Methuene, London, 1968, pp. 97- 104; Grebenik, P.; Grinter, R.; Perutz, R.N. Chem. Soc. Rev. 1988, 17, 453; 460.
CHAPTER 2
AROMATICITY 67
shell of the Fe atom possesses nine atomic orbitals, that is, one 4s, three 4 p, and five 3 d orbitals. The six filled orbitals of the two cyclopentadienide rings overlap with the s, three p, and two of the d orbitals of the Fe to form twelve new orbitals, six of which are bonding. These six orbitals make up two ring-to-metal triple bonds. In addition, further bonding results from the overlap of the empty antibonding orbitals of the rings with additional filled d orbitals of the iron. All told, there are 18 elec¬ trons (10 of which may be considered to come from the rings and 8 from iron in the zero oxidation state) in nine orbitals; six of these are strongly bonding and three weakly bonding or nonbonding.
The tropylium ion has an aromatic sextet spread over seven carbon atoms. An analogous ion, with the sextet spread over eight carbon atoms, is 1,3,5,7-tetra- methylcyclooctatetraene dictation (68). This ion, which is stable in solution at — 50°C, is diatropic and approximately planar. The dication 68 is not stable above about — 30°C.172
68
Other Systems Containing Aromatic Sextets
Simple resonance theory predicts that pentalene (69), azulene (70), and heptalene (71) should be aromatic, although no nonionic canonical form can have a double bond at the ring junction. Molecular-orbital calculations show that azulene should be stable but not the other two, and this is borne out by experiment. Heptalene has been prepared,173 but reacts readily with oxygen, acids, and bromine, is easily hydrogenated, and polymerizes on standing. Analysis of its NMR spectrum shows
69 70 71
172This and related ions were prepared by Olah, G.A.; Staral, J.S.; Liang, G.; Paquette, L.A.; Melega, W.P.; Carmody, MJ. J. Am. Chem. Soc. 1977 , 99, 3349. See also Radom, L.; Schaefer III, H.F. J. Am. Chem. Soc. 1977, 99, 7522; Olah, G.A.; Liang, G. J. Am. Chem. Soc. 1976, 98, 3033; Willner, I.; Rabinovitz, M. Nouv. J. Chim., 1982, 6, 129.
173Dauben, Jr., H.J.; Bertelli, D.J. J. Am. Chem. Soc. 1961, 83, 4659; Vogel, E.; Konigshofen, H.; Wassen, J.; Mullen, K.; Oth, J.F.M. Angew. Chem. Int. Ed. 1974, 13, 732; Paquette, L.A.; Browne, A.R.; Chamot, E. Angew. Chem. Int. Ed. 1979, 18, 546. For a review of heptalenes, see Paquette, L.A. Isr. J. Chem. 1980, 20, 233.
68 DELOCALIZED CHEMICAL BONDING
that it is not planar.174 The 3,8-dibromo and 3,8-dicarbomethoxy derivatives of 71 are stable in air at room temperature but are not diatropic.175 A number of methylated heptalenes and dimethyl 1 ,2-heptalenedicarboxylates have also been prepared and are stable nonaromatic compounds.176 Pentalene has not been prepared,177 but the hexaphenyl178 and 1,3,5-tri-te/t-butyl derivatives179 are known. The former is air sensitive in solution. The latter is stable, but X-ray diffrac¬ tion and photoelectron spectral data show bond alternation.180 Pentalene and its methyl and dimethyl derivatives have been formed in solution, but they dimerize before they can be isolated.181 Many other attempts to prepare these two systems have failed.
72 73 74 75 76
In sharp contrast to 69 and 71, azulene, a blue solid, is quite stable and many of its derivatives are known.182 Azulene readily undergoes aromatic substitution. Azulene may be regarded as a combination of 58 and 62 and, indeed, possesses a dipole moment of 0.8 D (see 72). 183 Interestingly, if two electrons are added to pentalene, a stable dianion (73) results.184 It can be concluded that an aromatic system of electrons will be spread over two rings only if 10 electrons (not 8 or 12) are available for aromaticity. [n,m\- Fluvalenes (n ^ m, where fulvalene is 74) as well as azulene are known to shift their ^-electrons due to the influence of dipolar aromatic resonance structures.185 However, calculations showed that
174Bertelli, D.J., in Bergmann, E.D.; Pullman, B. Aromaticity, Pseudo- Aromaticity, and Anti-Aromaticity, Israel Academy of Sciences and Humanities, Jerusalem, 1971, p. 326. See also Stegemann, J.; Lindner, HJ. Tetrahedron Lett. 1977, 2515.
175Vogel, E.; Ippen, J. Angew. Chem. Int. Ed. 1974, 13, 734; Vogel, E.; Hogrefe, F. Angew. Chem. Int. Ed. 1974, 13, 735.
176Hafner, K.; Knaup, G.L.; Lindner, H.J. Bull. Soc. Chem. Jpn. 1988, 61, 155.
177Metal complexes of pentalene have been prepared: Knox, S.A.R.; Stone, F.G.A. Acc. Chem. Res. 1974, 7, 321.
178LeGoff, E. J. Am. Chem. Soc. 1962, 84, 3975. See also Hafner, K.; Bangert, K.F.; Orfanos, V. Angew. Chem. Int. Ed. 1967, 6, 451; Hartke, K.; Matusch, R. Angew. Chem. Int. Ed. 1972, 11, 50.
179Hafner, K.; Suss, H.U. Angew. Chem. Int. Ed. 1973, 12, 575. See also Hafner, K.; Suda, M. Angew. Chem. Int. Ed. 1976, 15, 314.
180Kitschke, B.; Lindner, H.J. Tetrahedron Lett. 1977, 2511; Bischof, P.; Gleiter, R.; Hafner, K.; Knauer, K.H.; Spanget-Larsen, J.; Suss, H.U. Chem. Ber. 1978, 111, 932.
181Bloch, R.; Marty, R.A.; de Mayo, P. J. Am. Chem. Soc. 1971, 93, 3071; Bull. Soc. Chim. Fr., 1972, 2031; Hafner, K.; Donges, R.; Goedecke, E.; Kaiser, R. Angew. Chem. Int. Ed. 1973, 12, 337.
182For a review on azulene, see Mochalin, V.B.; Porshnev, Yu.N. Russ. Chem. Rev. 1977, 46, 530. 183Tobler, H.J.; Bauder, A.; Giinthard, H.H. J. Mol. Spectrosc., 1965, 18, 239.
184Katz, T.J.; Rosenberger, M.; O’Hara, R.K. J. Am. Chem. Soc. 1964, 86, 249. See also, Willner, I.; Becker, J.Y.; Rabinovitz, M. J. Am. Chem. Soc. 1979, 101, 395.
185Mollerstedt, H.; Piqueras, M.C.; Crespo, R.; Ottosson, H. J. Am. Chem. Soc. 2004, 126, 13938.
CHAPTER 2
AROMATICITY 69
dipolar resonance structures contribute only 5% to the electronic structure of hep- tafulvalene (75), although 22-31% to calicene (76). 186 Based on Baird’s theory,187 these molecules are influenced by aromaticity in both the ground and excited states, therefore acting as aromatic “chameleons.” This premise was confirmed in work by Ottosson and co-workers.185 Aromaticity indexes for various substituted fulva- lene compounds has been reported.188
Alternant and Nonalternant Hydrocarbons189
Aromatic hydrocarbons can be divided into alternant and nonaltemant hydrocar¬ bons. In alternant hydrocarbons, the conjugated carbon atoms can be divided into two sets such that no two atoms of the same set are directly linked. For conveni¬ ence, one set may be starred. Naphthalene is an alternant and azulene a nonalternant hydrocarbon:
*
In alternant hydrocarbons, the bonding and antibonding orbitals occur in pairs; that is, for every bonding orbital with an energy — E there is an antibonding one with energy +E (Fig. 2.8190). Even-alternant hydrocarbons are those with an even number of conjugated atoms, that is, an equal number of starred and unstarred atoms. For these hydrocarbons, all the bonding orbitals are filled and the n electrons are uniformly spread over the unsaturated atoms.
As with the allylic system, odd-alternant hydrocarbons (which must be carbo- cations, carbanions, or radicals) in addition to equal and opposite bonding and anti¬ bonding orbitals also have a nonbonding orbital of zero energy. When an odd number of orbitals overlap, an odd number is created. Since orbitals of alternant hydrocarbons occur in —E and +E pairs, one orbital can have no partner and must therefore have zero bonding energy. For example, in the benzylic system the cation has an unoccupied nonbonding orbital, the free radical has one electron there and the carbanion two (Fig. 2.9). As with the allylic system, all three species have the same bonding energy. The charge distribution (or unpaired-electron distribution)
186Scott, A.P.; Agranat, A.; Biedermann, P.U.; Riggs, N.V.; Radom, L. J. Org. Chem. 1997, 62, 2026. 187Baird, N.C. J. Am. Chem. Soc. 1972, 94, 4941.
188Stepien, B.T.; Krygowski, T.M.; Cyranski, M.K. J. Org. Chem. 2002, 67, 5987.
189For discussions, see Jones, R.A.Y. Physical and Mechanistic Organic Chemistry, 2nd ed.; Cambridge University Press, Cambridge, 1984, pp. 122-129; Dewar, M.J.S. Prog. Org. Chem. 1953, 2, 1.
190Taken from Dewar, M.J.S Prog. Org. Chem. 1953, 2, 1, p. 8.
70 DELOCALIZED CHEMICAL BONDING
_ J_ J
X X X J! XX X _tl
X v X X 1
Even a.h. Y
odd a.h
Fig. 2.8. Energy levels in odd- and even-alternant hydrocarbons.190 The arrows represent electrons. The orbitals are shown as having different energies, but some may be degenerate.
over the entire molecule is also the same for the three species and can be calculated by a relatively simple process.189
For nonalternant hydrocarbons the energies of the bonding and antibonding orbi¬ tals are not equal and opposite and charge distributions are not the same in cations, anions, and radicals. Calculations are much more difficult but have been carried
Energy a - 2.101(3
a - 1.259(3
t ti
Fig. 2.9. Energy levels for the benzyl cation, free radical, and carbanion. Since a is the energy of a /xorbital (p. 36), the nonbonding orbital has no bonding energy.
CHAPTER 2
AROMATICITY 71
out.191 Theoretical approaches to calculate topological polarization and reactivity of these hydrocarbons have been reported.192
Aromatic Systems with Electron Numbers Other Than Six
Ever since the special stability of benzene was recognized, chemists have been thinking about homologous molecules and wondering whether this stability is also associated with rings that are similar but of different sizes, such as cyclobutadiene (77), cyclooctatetraene (78), cyclodecapentaene (79) 193, and so on. The general
Q O
77 78
name annulene is given to these compounds, benzene being [6]annulene, and 77-79 being called, respectively, [4], [8], and [10] annulene. By a naive consideration of resonance forms, these annulenes and higher ones should be as aromatic as ben¬ zene. Yet they proved remarkably elusive. The ubiquitous benzene ring is found in thousands of natural products, in coal and petroleum, and is formed by strong treatment of many noncyclic compounds. None of the other annulene ring systems has ever been found in nature and, except for cyclooctatetraene, their synthesis is
not simple. Obviously, there is something special about the number six in a cyclic
system of electrons.
t
t
t t
tt h
t*
JL
— jj_—
Jj_
JL
Duet
Quartet
Sextet
Octet
(aromatic)
(diradical)
(aromatic)
(diradical)
Huckel’s rule , based on molecular-orbital calculations,194 predicts that elec¬ tron rings will constitute an aromatic system only if the number of electrons in the ring is of the form 4 n + 2, where n is zero or any position integer. Systems that contain 4 n electrons are predicted to be nonaromatic. The rule predicts that
191Peters, D. J. Chem. Soc. 1958, 1023, 1028, 1039; Brown, R.D.; Burden, F.R.; Williams, G.R. Aust. J. Chem. 1968, 21, 1939. For reviews, see Zahradnik, R., in Snyder, J.R Nonbenzenoid Aromatics vol. 2, Academic Press, NY, 1971, pp. 1-80; Zahradnik, R. Angew. Chem. Int. Ed. 1965, 4, 1039.
192Langler, R.F. Aust. J. Chem. 2000, 53, 471; Fredereiksen, M.U.; Langler, R.F.; Staples, M.A.; Verma, S.D. Aust. J. Chem. 2000, 53, 481.
193The cyclodecapentaene shown here is the cis-trans-cis-cis-trans form. For other stereoisomers, see p. 79.
194For reviews of molecular-orbital calculations of nonbenzenoid cyclic conjugated hydrocarbons, see Nakajima, T. Pure Appl. Chem. 1971, 28, 219; Fortschr. Chem. Forsch. 1972, 32, 1.
72 DELOCALIZED CHEMICAL BONDING
rings of 2, 6, 10, 14, and so on, electrons will be aromatic, while rings of 4, 8, 12, and so on, will not be. This is actually a consequence of Hund’s rule. The first pair of electrons in an annulene goes into the n orbital of lowest energy. After that the bonding orbitals are degenerate and occur in pairs of equal energy. When there is a total of four electrons, Hund’s rule predicts that two will be in the lowest orbital but the other two will be unpaired, so that the system will exist as a diradical rather than as two pairs. The degeneracy can be removed if the molecule is distorted from max¬ imum molecular symmetry to a structure of lesser symmetry. For example, if 77 assumes a rectangular rather than a square shape, one of the previously degenerate orbitals has a lower energy than the other and will be occupied by two electrons. In this case, of course, the double bonds are essentially separate and the molecule is still not aromatic. Distortions of symmetry can also occur when one or more car¬ bons are replaced by heteroatoms or in other ways.195
In the following sections systems with various numbers of electrons are dis¬ cussed. When we look for aromaticity we look for (7) the presence of a diamagnetic ring current; (2) equal or approximately equal bond distances, except when the symmetry of the system is disturbed by a heteroatom or in some other way; (2) pla¬ narity; (4) chemical stability; (5) the ability to undergo aromatic substitution.
Systems of Two Electrons196
Obviously, there can be no ring of two carbon atoms though a double bond may be regarded as a degenerate case. However, in analogy to the tropylium ion, a three- membered ring with a double bond and a positive charge on the third atom (the cyclopropenyl cation ) is a An + 2 system and hence is expected to show aromaticity. The unsubstituted 80 has been prepared,197 as well as several derivatives, e.g.,
v — @v — V0
80
the trichloro, diphenyl, and dipropyl derivatives, and these are stable despite the angles of only 60°. In fact, the tripropylcyclopropenyl,198 tricyclopropylcyclo- propenyl,199 chlorodipropylcyclopropenyl,200 and chloro-bisdialkylaminocyclo- propenyl201 cations are among the most stable carbocations known, being stable
195For a discussion, see Hoffmann, R. Chem. Commun. 1969, 240.
196For reviews, see Billups, W.E.; Moorehead, A.W., in Rappoport The Chemistry of the Cyclopropyl Group, pt. 2, Wiley, NY, 1987, pp. 1533-1574; Potts, K.T.; Baum, J.S. Chem. Rev. 1974, 74, 189; Yoshida, Z. Top. Curr. Chem. 1973, 40, 47; D’yakonov, I. A.; Kostikov, R.R. Russ. Chem. Rev. 1967, 36, 557; Closs, G.L. Adv. Alicyclic Chem. 1966, 1, 53, pp. 102-126; Krebs, A.W. Angew. Chem. Int. Ed. 1965, 4, 10. 197Farnum, D.G.; Mehta, G.; Silberman, R.G. J. Am. Chem. Soc. 1967, 89, 5048; Breslow, R.; Groves, J.T. J. Am. Chem. Soc. 1970, 92, 984.
198Breslow, R.; Hover, H.; Chang, H.W. J. Am. Chem. Soc. 1962, 84, 3168.
199Komatsu, K.; Tomioka, K.; Okamoto, K. Tetrahedron Lett. 1980, 21, 947 ; Moss, R.A.; Shen, S.; Krogh- Jespersen, K.; Potenza, J.A.; Schugar, H.J.; Munjal, R.C. J. Am. Chem. Soc. 1986, 108, 134.
200Ito, S.; Morita, N.; Asao, T. Tetrahedron Lett. 1992, 33, 3773.
201Taylor, M.J.; Surman, P.W.J.; Clark, G.R. J. Chem. Soc. Chem. Commun. 1994, 2517 .
CHAPTER 2
AROMATICITY 73
even in water solution. The tri-fp/Y-butylcyclopropenyl cation is also very stable.202 In addition, cyclopropenone and several of its derivatives are stable
Y
O
Cyclopropenone
R ++ R
O
R R
81
compounds,203 in accord with the corresponding stability of the tropones.204 The ring system 80 is nonaltemant and the corresponding radical and anion (which do not have an aromatic duet) have electrons in antibonding orbitals, so that their energies are much higher. As with 58 and 62, the equivalence of the three carbon atoms in the triphenylcyclopropenyl cation has been demonstrated by 14C labeling experiments.205 The interesting dications 81 (R = Me or Ph) have been prepared,206 and they too should represent aromatic systems of two electrons.207
Systems of Four Electrons: Antiaromaticity
The most obvious compound in which to look for a closed loop of four electrons is cyclo¬ butadiene (77).208 Hiickel’s rule predicts no aromatic character here, since 4 is not a number of the form 4 n + 2. There is a long history of attempts to prepare this com¬ pound and its simple derivatives, and those experiments fully bear out Huckel’s pre¬ diction. Cyclobutadienes display none of the characteristics that would lead us to call them aromatic, and there is evidence that a closed loop of four electrons is actually antiaromatic .209 If such compounds simply lacked aromaticity, we would expect
202Ciabattoni, J.; Nathan III, E.C. J. Am. Chem. Soc. 1968, 90, 4495.
203See, for example, Kursanov, D.N.; Vol’pin, M.E.; Koreshkov, Yu.D. J. Gen. Chem. USSR 1960, 30, 2855; Breslow, R.; Oda, M. J. Am. Chem. Soc. 1972, 94, 4787; Yoshida, Z.; Konishi, H.; Tawara, Y.; Ogoshi, H. J. Am. Chem. Soc. 1973, 95, 3043; Ciabattoni, J.; Nathan III, E.C. J. Am. Chem. Soc. 1968, 90, 4495.
204For a reveiw of cyclopropenones, see Eicher, T.; Weber, J.L. Top. Curr. Chem. Soc. 1975, 57, 1. For discussions of cyclopropenone structure, see Shafer, W.; Schweig, A.; Maier, G.; Sayrac, T.; Crandall, J.K. Tetrahedron Lett. 1974, 1213; Tobey, S.W., in Bergmann, E.D.; Pullman, B. Aromaticity, Pseudo- Aromaticity, and Anti-Aromaticity, Israel Academy of Sciences and Humanities, Jerusalem, 1971, pp. 351-362; Greenberg, A.; Tomkins, R.P.T.; Dobrovolny, M.; Liebman, J.F. J. Am. Chem. Soc. 1983, 105, 6855.
205D’yakonov, I.A.; Kostikov, R.R.; Molchanov, A.P J. Org. Chem. USSR 1969, 5, 171; 1970, 6, 304. 206Freedman, H.H.; Young, A.E. J. Am. Chem. Soc. 1964, 86, 734; Olah, G.A.; Staral, J.S. J. Am. Chem. Soc. 1976, 98, 6290. See also Lambert, J.B.; Holcomb, A.G. J. Am. Chem. Soc. 1971, 93, 2994; Seitz, G.; Schmiedel, R.; Mann, K. Synthesis, 1974, 578.
207See Pittman Jr., C.U.; Kress, A.; Kispert, L.D. J. Org. Chem. 1974, 39, 378. See, however, Krogh- Jespersen, K.; Schleyer, P.v.R.; Pople, J.A.; Cremer, D. J. Am. Chem. Soc. 1978, 100, 4301.
208For a monograph, see Cava, M.P.; Mitchell, M.J. Cyclobutadiene and Related Compounds', Academic Press, NY, 1967. For reviews, see Maier, G. Angew. Chem. Int. Ed. 1988, 27, 309; 1974, 13, 425-438; Bally, T.; Masamune, S. Tetrahedron 1980, 36, 343; Vollhardt, K.P.C. Top. Curr. Chem. 1975, 59, 113. 209For reviews of antiaromaticity, see Glukhovtsev, M.N.; Simkin, B.Ya.; Minkin, V.I. Russ. Chem. Rev. 1985, 54, 54; Breslow, R. Pure Appl. Chem. 1971, 28, 111; Acc. Chem. Res. 1973, 6, 393.
74 DELOCALIZED CHEMICAL BONDING
them to be about as stable as similar nonaromatic compounds, but both theory and experiment show that they are much less stable?10 An antiaromatic compound may be defined as a compound that is destabilized by a closed loop of electrons.
After years of attempts to prepare cyclobutadiene, the goal was finally reached by Pettit and co-workers.211 It is now clear that 77 and its simple derivatives are extremely unstable compounds with very short lifetimes (they dimerize by a Diels-Alder reaction; see 15-60) unless they are stabilized in some fashion, either at ordinary temperatures embedded in the cavity of a hemicarcerand212 (see the structure of a carcerand on p. 128), or in matrices at very low temperatures (gen¬ erally under 35 K). In either of these cases, the cyclobutadiene molecules are forced to remain apart from each other, and other molecules cannot get in. The structures of 77 and some of its derivatives have been studied a number of times using the low-temperature matrix technique.213 The ground- state structure of 77 is a rectan¬ gular diene (not a diradical) as shown by the ir spectra of 77 and deuterated 77 trapped in matrices,214 as well as by a photoelectron spectrum.215 Molecular- orbital calculations agree.216 The same conclusion was also reached in an elegant experi¬ ment in which 1,2-dideuterocyclobutadiene was generated. If 77 is a rectangular diene, the dideutero compound should exist as two isomers:
D
y
and
\
D
D D
The compound was generated (as an intermediate that was not isolated) and two isomers were indeed found.217 The cyclobutadiene molecule is not static, even in the matrices. There are two forms (77a and 77b), which rapidly interconvert.218
210For a discussion, see Bauld, N.L.; Welsher, T.L.; Cessac, J.; Holloway, R.L. J. Am. Chem. Soc. 1978, 100, 6920.
211Watts, L.; Fitzpatrick, J.D.; Pettit, R. J. Am. Chem. Soc. 1965, 87, 3253, 1966, 88, 623. See also, Cookson, R.C.; Jones, D.W. J. Chem. Soc. 1965, 1881.
212Cram, D.J.; Tanner, M.E.; Thomas, R. Angew. Chem. Int. Ed. 1991, 30, 1024.
213See, for example, Lin, C.Y.; Krantz, A. J. Chem. Soc. Chem. Commun. 1972, 1111; Chapman, O.L.; McIntosh, C.L.; Pacansky, J. J. Am. Chem. Soc. 1973, 95, 614; Maier, G.; Mende, U. Tetrahedron Lett. 1969, 3155. For a review, see Sheridan, R.S. Org. Photochem. 1987, 8, 159; pp. 167-181. 214Masamune, S.; Souto-Bachiller, F.A.; Machiguchi, T.; Bertie, J.E. J. Am. Chem. Soc. 1978, 100, 4889. 215Kreile, J.; Munzel, N.; Schweig, A.; Specht, H. Chem. Phys. Lett. 1986, 124, 140.
216See, for example, Borden, W.T.; Davidson, E.R.; Hart, P. J. Am. Chem. Soc. 1978, 100, 388; Kollmar, H.; Staemmler, V. J. Am. Chem. Soc. 1978, 100, 4304; Jafri, J.A.; Newton, M.D. J. Am. Chem. Soc. 1978, 100, 5012; Ermer, O.; Heilbronner, E. Angew. Chem. Int. Ed. 1983, 22, 402; Voter, A.F.; Goddard III, W.A. J. Am. Chem. Soc. 1986, 108, 2830.
217Whitman, D.W.; Carpenter, B.K. J. Am. Chem. Soc. 1980, 102, 4272. See also Whitman, D.W.; Carpenter, B.K. J. Am. Chem. Soc. 1982, 104, 6473.
218Carpenter, B.K. J. Am. Chem. Soc. 1983, 105, 1700; Huang, M.; Wolfsberg, M. J. Am. Chem. Soc. 1984, 106, 4039; Dewar, M.J.S.; Merz, Jr., K.M.; Stewart, J.J.P. J. Am. Chem. Soc. 1984, 106, 4040; Orendt, A.M.; Arnold, B.R.; Radziszewski, J.G.; Facelli, J.C.; Malsch, K.D.; Strub, H.; Grant, D.M.; Michl, J. J. Am. Chem. Soc. 1988, 110, 2648. See, however, Arnold, B.R.; Radziszewski, J.G.; Campion, A.; Perry, S.S.; Michl, J. J. Am. Chem. Soc. 1991, 113, 692.
CHAPTER 2
AROMATICITY 75
Note that there is experimental evidence that the aromatic and antiaromatic characters of neutral and dianionic systems are measurably increased via deuteration.219
1 _ 2 '“Buv /rBu
=n
1 _ 2
jni
4IJ - "3
77a
t- Bu
H
77b
82
There are some simple cyclobutadienes that are stable at room temperature for varying periods of time. These either have bulky substituents or carry certain other stabilizing substituents such as seen in tri-terf-butylcyclobutadiene (83). 220 Such compounds are relatively stable because dimerization is sterically hindered. Exam¬ ination of the NMR spectrum of 83 showed that the ring proton (5 = 5.38) was shifted upfield , compared with the position expected for a nonaromatic proton, for example, cyclopentadiene. As we will see (pp. 89-90), this indicates that the compound is antiaromatic.
Et2N
COOEt
Et2N
O
i
X.C'OEt
etc.
EtOOC NEt2
83
EtOOC NEt2 ©
The other type of stable cyclobutadiene has two electron-donating and two electron- withdrawing groups, and is stable in the absence of water. An example is 58. The stability of these compounds is generally attributed to the reso¬ nance shown, a type of resonance stabilization called the push-pull or captodative effect ,223 although it has been concluded from a photoelectron spectroscopy study that second-order bond fixation is more important.224 An X-ray crystallographic study of 83 has shown225 the ring to be a distorted square with bond lengths of 1.46 A and angles of 87° and 93°.
219For experiments with [16]-annulene (see p 82), see Stevenson, C.D.; Kurth, T.L. J. Am. Chem. Soc. 1999, 121, 1623
220Masamune, S.; Nakamura, N.; Suda, M.; Ona, H. J. Am. Chem. Soc. 1973, 95, 8481; Maier, G.; Alzerreca, A. Angew. Chem. Int. Ed. 1973, 12, 1015. For a discussion, see Masamune, S. Pure Appl. Chem. 1975, 44, 861.
221The presence of electron-donating and -withdrawing groups on the same ring stabilizes 4 n systems and destabilizes An + 2 systems. For a review of this concept, see Gompper, R.; Wagner, H. Angew. Chem. Int. Ed. 1988, 27, 1437.
222Neuensch wander, M.; Niederhauser, A. Chimia, 1968, 22, 491, Helv. Chim. Acta, 1970, 53, 519; Gompper, R.; Kroner, J.; Seybold, G.; Wagner, H. Tetrahedron 1976, 32, 629.
223Manatt, S.L.; Roberts, J.D. J. Org. Chem. 1959, 24, 1336; Breslow, R.; Kivelevich, D.; Mitchell, M.J.; Fabian, W.; Wendel, K. J. Am. Chem. Soc. 1965, 87, 5132; Hess Jr., B.A.; Schaad, L.J. J. Org. Chem. 1976, 41, 3058.
224Gompper, R.; Holsboer, F.; Schmidt, W.; Seybold, G. J. Am. Chem. Soc. 1973, 95, 8479.
225Lindner, H.J.; von Ross, B. Chem. Ber. 1974, 107, 598.
76 DELOCALIZED CHEMICAL BONDING
It is clear that simple cyclobutadienes, which could easily adopt a square planar shape if that would result in aromatic stabilization, do not in fact do so and are not aromatic. The high reactivity of these compounds is not caused merely by steric strain, since the strain should be no greater than that of simple cyclopropenes, which are known compounds. It is probably caused by antiaromaticity.226
R
Fe(CO)3 =
84
The cyclobutadiene system can be stabilized as a r|4-complex with metals,227 as with the iron complex 84 (see Chapter 3), but in these cases electron density is withdrawn from the ring by the metal and there is no aromatic quartet. In fact, these cyclobutadiene-metal complexes can be looked upon as systems containing an aromatic duet. The ring is square planar,228 the compounds undergo aromatic sub¬ stitution,229 and nmr spectra of monosubstituted derivatives show that the C-2 and C-4 protons are equivalent.229
V — ve
©
85
86
Other systems that have been studied as possible aromatic or antiaromatic four- electron systems include the cyclopropenyl anion (86), the cyclopentadienyl cation (8?) 230 with respect to 86, HMO theory predicts that an unconjugated 85 (i.e., a single canonical form) is more stable than a conjugated 86,231 so that 85 would actually lose stability by forming a closed loop of four electrons. The HMO theory
226For evidence, see Breslow, R.; Murayama, D.R.; Murahashi, S.; Grubbs, R. J. Am. Chem. Soc. 1973, 95, 6688; Herr, M.L. Tetrahedron 1976, 32, 2835.
227For reviews, see Efraty, A. Chem. Rev. 1977, 77, 691; Pettit, R. Pure Appl. Chem. 1968, 17, 253; Maitlis, P.M. Adv. Organomet. Chem. 1966, 4, 95; Maitlis, P.M.; Eberius, K.W., in Snyder, J.P. Nonbenzenoid Aromatics, vol. 2, Academic Press, NY, 1971, pp. 359-409.
228Dodge, R.P.; Schomaker, V. Acta Crystallogr. 1965, 18, 614; Nature ( London ) 1960, 186, 798; Dunitz, J.D.; Mez, H.C.; Mills, O.S.; Shearer, H.M.M. Helv. Chim. Acta, 1962, 45, 647; Yannoni, C.S.; Ceasar, G.P.; Dailey, B.P. J. Am. Chem. Soc. 1967, 89, 2833.
229Fitzpatrick, J.D.; Watts, L.; Emerson, G.F.; Pettit, R. J. Am. Chem. Soc. 1965, 87, 3255. For a discussion, see Pettit, R. J. Organomet. Chem. 1975, 100, 205.
230For a review of cyclopentadienyl cations, see Breslow, R. Top. Nonbenzenoid Aromat. Chem. 1973, 1, 81.
231Clark, D.T. Chem. Commun. 1969, 637 ; Glukhovtsev, M.N.; Simkin, B.Ya.; Minkin, V.I. Russ. Chem. Rev. 1985, 54, 54; Breslow, R. Pure Appl. Chem. 1971, 28, 111; Acc. Chem. Res. 1973, 6, 393.
CHAPTER 2
AROMATICITY 77
is supported by experiment. Among other evidence,
Ph
Ph
/
K ><
H
R
Ph
Ph
88
89
it has been shown that 88 (R = COPh) loses its proton in hydrogen-exchange reac¬ tions ^6000 times more slowly than 89 (R = COPh).232 Where R = CN, the ratio is ~10,000.233 This indicates that 88 are much more reluctant to form carbanions (which would have to be cyclopropenyl carbanions) than 89, which form ordinary carbanions. Thus the carbanions of 88 are less stable than corresponding ordinary carbanions. Although derivatives of cyclopropenyl anion have been prepared as fleeting intermediates (as in the exchange reactions mentioned above), all attempts to prepare the ion or any of its derivatives as relatively stable species have so far met with failure.234
In the case of 87, the ion has been prepared and has been shown to be a diradical in the ground state,235 as predicted by the discussion on p. 73. 236 Evidence that 87 is not only nonaromatic, but also antiaromatic comes from studies on 90 and 92. 237 When 90 is treated with silver perchlorate in propionic acid, the molecule is rapidly solvolyzed (a reaction in which the intermediate 91 is formed; see Chapter 5). Under the same conditions, 92 undergoes no solvolysis at all; that is, 87 does not form. If 87 were merely nonaromatic, it should be about as stable as 91 (which of course has no resonance stabilization at all). The fact that it is so much more reluc¬ tant to form indicates that 87 is much less stable than 91. It is noted that under cer¬ tain conditions, 91 can be generated solvolytically.238
90
91
92
87
232B reslow, R.; Brown, J.; Gajewski, J.J. J. Am. Chem. Soc. 1967, 89, 4383.
233Breslow, R.; Douek, M. J. Am. Chem. Soc. 1968, 90, 2698.
234See, for example, Breslow, R.; Cortes, D.A.; Juan, B.; Mitchell, R.D. Tetrahedron Lett. 1982, 23, 795. A triphenylcyclopropyl anion has been prepared in the gas phase, with a lifetime of 1-2 s: Bartmess, J.E.; Kester, J.; Borden, W.T.; Koser, H.G. Tetrahedron Lett. 1986, 27, 5931.
235Saunders, M.; Berger, R.; Jaffe, A.; McBride, J.M.; O’Neill, J.; Breslow, R.; Hoffman Jr., J.M.; Perchonock, C.; Wasserman, E.; Hutton, R.S.; Kuck, V.J. J. Am. Chem. Soc. 1973, 95, 3017. 236Derivatives of 87 show similar behavior. Volz, H. Tetrahedron Lett. 1964, 1899; Breslow, R.; Chang, H.W.; Hill, R.; Wasserman, E. J. Am. Chem. Soc. 1967, 89, 1112; Gompper, R.; Glockner, H. Angew. Chem. Int. Ed. 1984, 23, 53.
237Breslow, R.; Mazur, S. J. Am. Chem. Soc. 1973, 95, 584. For further evidence, see Lossing, F.P.; Treager, J.C. J. Am. Chem. Soc. 1975, 97, 1579. See also, Breslow, R.; Canary, J.W. J. Am. Chem. Soc. 1991, 113, 3950.
238Allen, A.D.; Sumonja, M.; Tidwell, T.T. J. Am. Chem. Soc. 1997, 119, 2371.
78 DELOCALIZED CHEMICAL BONDING
It is strong evidence for Hiickel’s rule that 86 and 87 are not aromatic while the cyclopropenyl cation (80) and the cyclopentadienyl anion (58) are, since simple resonance theory predicts no difference between 86 and 80 or 87 and 58 (the same number of equivalent canonical forms can be drawn for 86 as for 80 and for 87 as for 58).
H
H H
78a
Systems of Eight Electrons
Cyclooctatetraene239 ([8]annulene, 78a) is not planar, but tub-shaped.240 There¬ fore we would expect that it is neither aromatic nor antiaromatic, since both these conditions require overlap of parallel p orbitals. The reason for the lack of planarity is that a regular octagon has angles of 135°, while sp2 angles are most stable at 120°. To avoid the strain, the molecule assumes a nonplanar shape, in which orbital overlap is greatly diminished.241 Single- and double-bond distances in 78 are, respectively, 1.46 and 1.33 A, which is expected for a compound made up of four individual double bonds.240 The reactivity is also what would be expected for a linear polyene. Reactive intermediates can be formed in solution. Dehydroha- logenation of bromocyclooctatetraene at — 100°C has been reported, for example, and trapping by immediate electron transfer gave a stable solution of the [8]annu- lyne anion radical.242
The cyclooctadiendiynes 93 and 94 are planar conjugated eight-electron systems (the four extra triple-bond electrons do not participate), which nmr evidence show to be antiaromatic.243 There is evidence that part of the reason for the lack of planarity in 78 itself is that a planar molecular would have to be antiaromatic.244 The cycloheptatrienyl anion (61) also has eight electrons, but does not behave like an aromatic system.151 The bond lengths for a series of molecules containing the cycloheptatrienide anion have recently been published.245 The NMR spectrum
239For a monograph, see Fray, G.I.; Saxton, R.G. The Chemistry of Cyclooctatetraene and its Derivatives', Cambridge University Press: Cambridge, 1978. For a review, see Paquette, L.A. Tetrahedron 1975, 31, 2855. For reviews of heterocyclic 871 systems, see Kaim, W. Rev. Chem. Intermed. 1987, 8, 247; Schmidt, R.R. Angew. Chem. Int. Ed. 1975, 14, 581.
240Bastiansen, O.; Hedberg, K.; Hedberg, L. J. Chem. Phys. 1957, 27, 1311.
241The compound perfluorotetracyclobutacyclooctatetraene has been found to have a planar cyclooctate¬ traene ring, although the corresponding tetracyclopenta analog is nonplanar: Einstein, F.W.B.; Willis, A.C.; Cullen, W.R.; Soulen, R.L. J. Chem. Soc. Chem. Commun. 1981, 526. See also, Paquette, L.A.; Wang, T.; Cottrell, C.E. J. Am. Chem. Soc. 1987, 109, 3730.
242Peters, S.J.; Turk, M.R.; Kiesewetter, M.K.; Stevenson, C.D. J. Am. Chem. Soc. 2003, 125, 11264. 243For a review, see Huang, N.Z.; Sondheimer, F. Acc. Chem. Res. 1982, 15, 96. See also, Durr, H.; Klauck, G.; Peters, K.; von Schnering, H.G. Angew. Chem. Int. Ed. 1983, 22, 332; Chan, T.; Mak, T.C.W.; Poon, C.; Wong, H.N.C.; Jia, J.H.; Wang, L.L. Tetrahedron 1986, 42, 655.
244Figeys, H.P; Dralants, A. Tetrahedron Lett. 1971, 3901; Buchanan, G.W. Tetrahedron Lett. 1972, 665. 245Dietz, F.; Rabinowitz, M.; Tadjer, A.; Tyutyulkov, N. J. Chem. Soc. Perkin Trans. 2 1995, 735.
CHAPTER 2
AROMATICITY 79
of the benzocycloheptatrienyl anion (95) shows that, like 82, 93, and 94, this com¬ pound is antiaromatic.246 A new antiaromatic compound 1 ,4-biphenylene quinone (96) was prepared, but it rapidly dimerizes due to instability.247
o
Systems of Ten Electrons248
There are three geometrically possible isomers of [10]annulene: the all -cis (97), the mono-trans (98), and the cis-trans-cis-cis-trans (79). If Hiickel’s rule applies, they should be planar. But it is far from obvious that the molecules would adopt a planar
79 97 98
shape, since they must overcome considerable strain to do so. For a regular decagon (97) the angles would have to be 144°, considerably larger than the 120° required for sp2 angles. Some of this strain would also be present in 98, but this kind of strain is eliminated in 79 since all the angles are 120°. However, it was pointed out by Mislow249 that the hydrogens in the 1 and 6 positions should interfere with each other and force the molecule out of planarity.
99 100 101 102
Compounds 97 and 98 have been prepared250 as crystalline solids at — 80° C. The NMR spectra show that all the hydrogens lie in the alkene region and it was concluded that neither compound is aromatic. Calculations on 98 suggest that
246Staley, S.W.; Orvedal, A.W. J. Am. Chem. Soc. 1973, 95, 3382.
247Kili?, H.; Balci, M. J. Org. Chem. 1997, 62, 3434.
248For reviews, see Kemp-Jones, A.V.; Masamune, S. Top. Nonbenzenoid Aromat. Chem. 1973, 1, 121; Masamune, S.; Darby, N. Acc. Chem. Res. 1972, 5, 272; Burkoth, T.L.; van Tamelen, E.E., in Snyder, J.P. Nonbenzenoid Aromaticity, Vol. 1, Academic Press, NY, 1969, pp. 63-116; Vogel, E., in Garratt, P.J. Aromaticity, Wiley, NY, 1986, pp. 113-147.
249Mislow, K. J. Chem. Phys. 1952, 20, 1489.
250Masamune,S.; Hojo, K.; Bigam, G.; Rabenstein, D.L. J. Am. Chem. Soc. 1971, 93, 4966. [10]Annulenes had previously been prepared, but it was not known which ones: van Tamelen, E.E.; Greeley, R.H. Chem. Commun. 1971, 601; van Tamelen, E.E.; Burkoth, T.L.; Greeley, R.H. J. Am. Chem. Soc. 1971, 93, 6120.
80 DELOCALIZED CHEMICAL BONDING
it may indeed be aromatic, although the other isomers are not.251 It is known that the Hartree-Fock (HF) method incorrectly favors bond-length-alternating structures for [10]annulene, and aromatic structures are incorrectly favored by density func¬ tional theory. Improved calculations predict that the twist conformation is lowest in energy, and the naphthalene-like and heart-shaped conformations lie higher than the twist by 1.40 and 4.24 kcal mol-1, respectively.252 From 13C and proton (H1) nmr spectra it has been deduced that neither is planar. However, that the angle strain is not insurmountable has been demonstrated by the preparation of several compounds that have large angles, but that are definitely planar 10-electron aromatic systems. Among these are the dianion 99, the anions 100 and 101, and the azonine 102 253 Compound 99254 has angles of ^135°, while 100255 and 101256 have angles of ^140°, which are not very far from 144°. The inner proton in 101257 (which is the mono-trans isomer of the all-cis 100) is found far upfield in the NMR (—3.5 5). For 97 and 98, the cost in strain energy to achieve planarity apparently out¬ weighs the extra stability that would come from an aromatic ring. To emphasize the delicate balance between these factors, we may mention that the oxygen analog of 102 (X = O, oxonin) and the N- carb ethoxy derivative of 102 (X = CH) are nonaromatic and nonplanar, while 102 (X = N) is aromatic and planar.258 Other azaannulenes are known, including Vogel’s 2,7-methanoazaannulene,259 as well
251Sulzbach, H.M.; Schleyer, P.v.R.; Jiao, H.; Xie, Y.; Schaefer III, H.F. J. Am. Chem. Soc. 1995, 117, 1369. Also see, Sulzbach, H.M.; Schaefer III, H.F.; Klopper, W.; Luthi, H.P. J. Am. Chem. Soc. 1996, 118, 3519 for a discussion of Aromaticity calculations for [10]annulene.
252King, R.A.; Crawford, T.D.; Stanton, J.F.; Schaefer, III, H.F. J. Am. Chem. Soc. 1999, 121, 10788. 253For reviews of 102 (X = N) and other nine-membered rings containing four double bonds and a hetero atom (heteronins), see Anastassiou, A.G. Acc. Chem. Res. 1972, 5, 281, Top. Nonbenzenoid Aromat. Chem. 1973, 1, 1, Pure Appl. Chem. 1975, 44, 691. For a review of heteroannulenes in general, see Anastassiou; Kasmai, H.S. Adv. Heterocycl. Chem. 1978, 23, 55.
254Katz, T.J. J. Am. Chem. Soc. 1960, 82, 3784, 3785; Goldstein, M.J.; Wenzel, T.T. J. Chem. Soc. Chem. Commun. 1984, 1654; Garkusha, O.G.; Garbuzova, I. A.; Lokshin, B.V.; Todres, Z.V. J. Organomet. Chem. 1989, 371, 279. See also, Noordik, J.H.; van den Hark, T.E.M.; Mooij, J.J.; Klaassen, A.A.K. Acta Crystallogr. Sect. B. 1974, 30, 833; Goldberg, S.Z.; Raymond, K.N.; Harmon, C.A.; Templeton, D.H. J. Am. Chem. Soc. 1974, 96, 1348; Evans, W.J.; Wink, D.J.; Wayda, A.L.; Little, D.A. J. Org. Chem. 1981, 46, 3925; Heinz, W.; Langensee, P; Mullen, K. J. Chem. Soc. Chem. Commun. 1986, 947 .
255Katz, T.J.; Garratt, PJ. J. Am. Chem. Soc. 1964, 86, 5194; LaLancette, E.A.; Benson, R.E. J. Am. Chem. Soc. 1965, 87, 1941; Simmons, H.E.; Chesnut, D.B.; LaLancette, E.A. J. Am. Chem. Soc. 1965, 87, 982; Paquette, L.A.; Ley, S.V.; Meisinger, R.H.; Russell, R.K.; Oku, M. J. Am. Chem. Soc. 1974, 96, 5806; Radlick, P; Rosen, W. J. Am. Chem. Soc. 1966, 88, 3461.
256Anastassiou, A.G.; Gebrian, J.H. Tetrahedron Lett. 1970, 825.
257Boche, G.; Weber, H.; Martens, D.; Bieberbach, A. Chem. Ber. 1978, 111, 2480. See also, Anastassiou, A.G.; Reichmanis, E. Angew. Chem. Int. Ed. 1974, 13, 728; Boche, G.; Bieberbach, A. Tetrahedron Lett. 1976, 1021.
258Anastassiou, A.G.; Gebrian, J.H. J. Am. Chem. Soc. 1969, 91, 4011; Chiang, C.C.; Paul, I.C.; Anastassiou, A.G.; Eachus, S.W. J. Am. Chem. Soc. 1974, 96, 1636.
259Vogel, E.; Roth, H.D. Angew. Chem. Int. Ed. 1964, 3, 228; Vogel, E.; Biskup, M.; Pretzer, W.; Boll, W.A. Angew. Chem. Int. Ed. 1964, 3, 642.; Vogel, E.; Meckel, M.; Grimme, W. Angew. Chem. Int. Ed. 1964, 3, 643; Vogel, E.; Pretzer, W.; Boll, W.A. Tetrahedron Lett. 1965, 3613; Sondheimer, F.; Shani, A. J. Am. Chem. Soc. 1964, 86, 3168; Shani, A.; Sondheimer, F. J. Am. Chem. Soc. 1967, 89, 6310; Bailey, N.A.; Mason, R. J. Chem. Soc. Chem. Commun. 1967, 1039.
CHAPTER 2
AROMATICITY 81
as 3,8-methanoaza[10]annulene,260 and their alkoxy derivatives.261 Calculations for aza[10]annulene concluded that the best olefinic twist isomer is 2.1kcalmol_1 (8.8 kJ mol-1) more stable than the aromatic form,262 and is probably the more stable form.
ch2 o nh
103 104 105
So far, 79 has not been prepared despite many attempts. However, there are var¬ ious ways of avoiding the interference between the two inner protons. The approach that has been most successful involves bridging the 1 and 6 positions.263 Thus, 1,6- methano[10]annulene (103)264 and its oxygen and nitrogen analogs, 104265 and 105, 266 have been prepared and are stable compounds that undergo aromatic sub¬ stitution and are diatropic.267 For example, the perimeter protons of 103 are found at 6. 9-7. 3 5, while the bridge protons are at —0.5 5. The crystal structure of 103 shows that the perimeter is nonplanar, but the bond distances are in the range 1.37- 1.42 A.268 It has therefore been amply demonstrated that a closed loop of 10 elec¬ trons is an aromatic system, although some molecules that could conceivably have such a system are too distorted from planarity to be aromatic. A small distortion from planarity (as in 103) does not prevent aromaticity, at least in part because the s orbitals so distort themselves as to maximize the favorable (parallel) overlap
260Schafer-Ridder, M.; Wagner, A.; Schwamborn, M.; Schreiner, H.; Devrout, E.; Vogel, E. Angew. Chem. Int. Ed. 1978, 17, 853.; Destro, R.; Simonetta, M.; Vogel, E. J. Am. Chem. Soc. 1981, 103, 2863. 261Vogel, E. Presented at the 3rd International Symposium on Novel Aromatic Compounds (ISNA 3), San Francisco, Aug 1977; Golz, H.-J.; Muchowski, J.M.; Maddox, M.L. Angew. Chem. Int. Ed. 1978, 17, 855; Schleyer, P.v.R.; Jiao, H.; Sulzbach, H.M.; Schaefer III H.F. J. Am. Chem. Soc. 1996, 118, 2093. 262Bettinger, H.F.; Sulzbach, H.M.; Schleyer, P.v.R.; Schaefer III, H.F. J. Org. Chem. 1999, 64, 3278. 263For reviews of bridged [10]-, [14]-, and [18]annulenes, see Vogel, E. Pure Appl. Chem. 1982, 54, 1015; Isr. J. Chem. 1980, 20, 215; Chimia, 1968, 22, 21; Vogel, E.; Gunther, H. Angew. Chem. Int. Ed. 1967, 6, 385.
264Vogel, E.; Roth, H.D. Angew. Chem. Int. Ed. 1964, 3, 228; Vogel, E.; Boll, W.A. Angew. Chem. Int. Ed. 1964, 3, 642; Vogel, E.; Boll, W.A.; Biskup, M. Tetrahedron Lett. 1966, 1569.
265Vogel, E.; Biskup, M.; Pretzer, W.; Boll, W.A. Angew. Chem. Int. Ed. 1964, 3, 642; Shani, A.; Sondheimer, F. J. Am. Chem. Soc. 1967, 89, 6310; Bailey, N.A.; Mason, R. Chem. Commun. 1967, 1039.
266Vogel, E.; Pretzer, W.; Boll, W.A. Tetrahedron Lett. 1965, 3613. See also, Vogel, E.; Biskup, M.; Pretzer, W.; Boll, W.A. Angew. Chem. Int. Ed. 1964, 3, 642.
267For another type of bridged diatropic [10]annulene, see Lidert, Z.; Rees, C.W. J. Chem. Soc. Chem. Commun. 1982, 499; Gilchrist, T.L.; Rees, C.W.; Tuddenham, D. J. Chem. Soc. Perkin Trans. 1 1983, 83; McCague, R.; Moody, C.J.; Rees, C.W. J. Chem. Soc. Perkin Trans. 1 1984, 165, 175; Gibbard, H.C.; Moody, C.J.; Rees, C.W. J. Chem. Soc. Perkin Trans. 1 1985, 731, 735.
268Bianchi, R.; Pilati, T.; Simonetta, M. Acta Crystallogr., Sect. B 1980, 36, 3146. See also Dobler, M.; Dunitz, J.D. Helv. Chim Acta, 1965, 48, 1429.
82 DELOCALIZED CHEMICAL BONDING
of p orbitals to form the aromatic 10-electron loop.269
106
107
In 106, where 103 is fused to two benzene rings in such a way that no canonical form can be written in which both benzene rings have six electrons, the aromaticity is reduced by annellation, as shown by the fact that the molecule rapidly converts to the more stable 107, in which both benzene rings can be fully aromatic270 (this is similar to the cycloheptatriene-norcaradiene conversions discussed on p. 1664).
n
C13C
C13C
108
Molecules can sustain significant distortion from planarity and retain their aro¬ matic character. l,3-Bis(trichloroacetyl)homoazulene (108) qualifies as aromatic using the geometric criterion that there is only a small average deviation from the C— C bond length in the [10]annulene perimeter.271 X-ray crystal structure shows that the 1,5-bridge distorts the [10]-annulene 71-system away from planarity (see the 3D model) with torsion angles as large as 42.2° at the bridgehead position, but 108 does not lose its aromaticity.
Systems of More than Ten Electrons: An + 2 Electrons272
Extrapolating from the discussion of [10]annulene, we expect larger 4n + 2 systems to be aromatic if they are planar. Mislow249 predicted that [14]annulene (109)
269For a discussion, see Haddon, R.C. Acc. Chem. Res. 1988, 21, 243.
270Hill, R.K.; Giberson, C.B.; Silverton, J.V. J. Am. Chem. Soc. 1988, 110, 497. See also, McCague, R.; Moody, C.J.; Rees, C.W.; Williams, D.J. J. Chem. Soc. Perkin Trans. 1 1984, 909.
271Scott, L.T.; Sumpter, C.A.; Gantzel, P.K.; Maverick, E.; Trueblood, K.N. Tetrahedron 2001 , 57, 3795. 272For reviews of annulenes, with particular attention to their nmr spectra, see Sondheimer, F. Acc. Chem. Res. 1972, 5, 81-91, Pure Appl. Chem. 1971, 28, 331, Proc. R. Soc. London. Ser. A, 1967, 297, 173; Sondheimer, F.; Calder, I.C.; Elix, J.A.; Gaoni, Y; Garratt, P.J.; Grohmann, K.; di Maio, G.; Mayer, J.; Sargent, M.V.; Wolovsky, R. in Garratt, PG. Aromaticity, Wiley, NY, 1986, pp. 75-107; Haddon, R.C.; Haddon, V.R.; Jackman, F.M. Fortschr. Chem. Forsch. 1971, 16, 103. For a review of annulenoannulenes (two annulene rings fused together), see Nakagawa, M. Angew. Chem. Int. Ed. 1979, 18, 202. For a review of reduction and oxidation of annulenes; that is, formation of radical ions, dianions, and dications, see Mullen, K. Chem. Rev. 1984, 84, 603. For a review of annulene anions, see Rabinovitz, M. Top. Curr. Chem. 1988, 146, 99. Also see Cyvin, S.J.; Brunvoll, J.; Chen, R.S.; Cyvin, B.N.; Zhang, FJ. Theory of Coronoid Hydrocarbons II, Springer- Verlag, Berlin, 1994.
CHAPTER 2
AROMATICITY 83
would possess the same type of interference as 79, although in lesser degree. This is
borne out by experiment. Compound 109 is aromatic (it is diatropic; inner protons at 0.00 5, outer protons at 7.6 5), 273 but is completely destroyed by light and air in 1 day. X-ray analysis shows that although there are no alternating single and double bonds, the molecule is not planar.274 A number of stable bridged [14]annulenes have been prepared,275 for example, frafts-15,16-dimethyldi- hydropyrene (110), yyft-l,6:8,13-diimino[14]annulene (111), and syn- and aftf/-l,6:8,13-bis(methano[14]annulene) (112 and 113).278 The dihydropyrene 110
H H
(and its diethyl and dipropyl homologs) is undoubtedly aromatic: the n peri¬ meter is approximately planar;279 the bond distances are all 1.39-1.40 A; and the
273Gaoni, Y.; Melera, A.; Sondheimer, F.; Wolovsky, R. Proc. Chem. Soc. 1964, 397.
274Bregman, J. Nature ( London ) 1962, 194, 679; Chiang, C.C.; Paul, I.C. J. Am. Chem. Soc. 1972, 94, 4741. Another 14-electron system is the dianion of [12]annulene, which is also apparently aromatic though not planar: Oth, J.F.M.; Schroder, G. J. Chem. Soc. B, 1971, 904. See also Garratt, P.J.; Rowland, N.E.; Sondheimer, F. Tetrahedron 1971, 27, 3157; Oth, J.F.M.; Mullen, K.; Konigshofen, H.; Mann, M.; Sakata, Y.; Vogel, E. Angew. Chem. Int. Ed. 1974, 13, 284. For some other 14-electron aromatic systems, see Anastassiou, A.G.; Elliott, R.L.; Reichmanis, E. J. Am. Chem. Soc. 1974, 96, 7823; Wife, R.L.; Sondheimer, F. J. Am. Chem. Soc. 1975, 97, 640; Ogawa, H.; Kubo, M.; Saikachi, H. Tetrahedron Lett. 1971, 4859; Oth, J.F.M.; Mullen, K.; Konigshofen, H.; Wassen, J.; Vogel, E. Helv. Chim. Acta, 1974, 57, 2387; Willner, I.; Gutman, A.L.; Rabinovitz, M. J. Am. Chem. Soc. 1977, 99, 4167; Rottele, H.; Schroder, G. Chem. Ber. 1982, 115, 248.
275For a review, see Vogel, E. Pure Appl. Chem. 1971, 28, 355.
276Boekelheide, V.; Phillips, J.B. J. Am. Chem. Soc. 1967, 89, 1695; Boekelheide, V.; Miyasaka, T. J. Am. Chem. Soc. 1967, 89, 1709. For reviews of dihydropyrenes, see Mitchell, R.H. Adv. Theor. Interesting Mol. 1989, 1, 135; Boekelheide, V. Top. Nonbenzoid Arom. Chem. 1973, 7, 47; Pure Appl. Chem. 1975, 44, 807.
277Vogel, E.; Kuebart, F.; Marco, J.A.; Andree, R.; Gunther, H.; Ay din, R. J. Am. Chem. Soc. 1983, 105, 6982; Destro, R.; Pilati, T.; Simonetta, M.; Vogel, E. J. Am. Chem. Soc. 1985, 107, 3185, 3192. For the di -O- analog of 102, see Vogel, A.; Biskup, M.; Vogel, E.; Gunther, H. Angew. Chem. Int. Ed. 1966, 5, 734. 278Vogel, E.; Sombroek, J.; Wagemann, W. Angew. Chem. Int. Ed. 1975, 14, 564.
279Hanson, A.W. Acta Crystallogr. 1965, 18, 599, 1967, 23, 476.
84 DELOCALIZED CHEMICAL BONDING
molecule undergoes aromatic substitution276 and is diatropic.280 The outer protons are found at 8.14-8.67 5, while the CH3 protons are at —4.25 5. Other nonplanar aromatic dihydropyrenes are known.281 Annulenes 111 and 112 are also diatro¬ pic,282 although X-ray crystallography indicates that the n periphery in at least 111 is not quite planar.283 However, 113, in which the geometry of the molecule greatly reduces the overlap of the p orbitals at the bridgehead positions with adja¬ cent p orbitals, is definitely not aromatic,284 as shown by NMR spectra278 and X-ray crystallography, from which bond distances of 1.33-1.36 A for the double bonds and 1.44-1.49 A for the single bonds have been obtained.285 In contrast, all the bond distances in 111 are ~ 1.38-1. 40 A.283
Another way of eliminating the hydrogen interferences of [14]annulene is to intro¬ duce one or more triple bonds into the system, as in dehydro [1 4] annulene (114).286 All five known dehydro [14] annulenes are diatropic, and 87 can be nitrated or sulfonated.287 The extra electrons of the triple bond do not form part of the aromatic system, but simply
exist as a localized bond. There has been a debate concerning the extent of deloca¬ lization in dehydrobenzoannulenes,288 but there is evidence for a weak, but discern¬ ible ring current.289 3,4,7,8,9,10,13,14-Octahydro[14]annulene (116) has been
280A number of annellated derivatives of 110 are less diatropic, as would be expected from the discussion on p. $$$: Mitchell, R.H.; Williams, R.V.; Mahadevan, R.; Lai, Y.H.; Dingle, T.W. J. Am. Chem. Soc. 1982, 104, 2571 and other papers in this series.
281Bodwell, G.J.; Bridson, J.N.; Chen, S.-L.; Poirier, R.A. J. Am. Chem. Soc. 2001, 123, 4704; Bodwell, G.J.; Fleming, J.J.; Miller, D.O. Tetrahedron 2001, 57, 3511.
282 As are several other similarly bridged [14] annulenes; see, for example, Flitsch, W.; Peeters, H. Chem. Ber. 1973, 106, 1731; Huber, W.; Lex, J.; Meul, T.; Mullen, K. Angew. Chem. Int. Ed. 1981, 20, 391; Vogel, E.; Nitsche, R.; Krieg, H. Angew. Chem. Int. Ed. 1981, 20, 811; Mitchell, R.H.; Anker, W. Tetrahedron Lett. 1981, 22, 5139; Vogel, E.; Wieland, H.; Schmalstieg, L.; Lex, J. Angew. Chem. Int. Ed. 1984, 23, 111 ; Neumann, G.; Mullen, K. J. Am. Chem. Soc. 1986, 108, 4105.
283Ganis, R; Dunitz, J.D. Helv. Chim. Acta, 1967, 50, 2369.
284For another such pair of molecules, see Vogel, E.; Nitsche, R.; Krieg, H. Angew. Chem. Int. Ed. 1981, 20, 811. See also, Vogel, E.; Schieb, T.; Schulz, W.H.; Schmidt, K.; Schmickler, H.; Lex, J. Angew. Chem. Int. Ed. 1986, 25, 723.
285Gramaccioli, C.M.; Mimun, A.; Mugnoli, A.; Simonetta, M. Chem. Commun. 1971, 796. See also, Destro, R.; Simonetta, M. Tetrahedron 1982, 38, 1443.
286For a review of dehydroannulenes, see, Nakagawa, M. Top. Nonbenzenoid Aromat. Chem. 1973, 1, 191. 287Gaoni, Y.; Sondheimer, F. J. Am. Chem. Soc. 1964, 86, 521.
288Balaban, A.T.; Banciu, M.; Ciorba, V. Annulenes, Benzo-, Hetero-, Homo- Derivatives and their Valence Isomers, Vols. 1-3, CRC Press, Boca Raton, FL, 1987; Garratt, PJ. Aromaticity, Wiley, NY, 1986; Minkin, V.I.; Glukhovtsev, M.N.; Simkin, B.Ya. Aromaticity and Antiaromaticity, Wiley, NY, 1994. 289Kimball, D.B.; Wan, W.B.; Haley, M.M. Tetrahdron Lett. 1998, 39, 6795; Bell, M.L.; Chiechi, R.C.; Johnson, C.A.; Kimball, D.B.; Matzger, A.J.; Wan, W.B.; Weakley, T.J.R.; Haley, M.M. Tetahedron 2001, 57, 3507; Wan, W.B.; Chiechi, R.C.; Weakley, T.J.R.; Haley, M.M. Eur. J. Org. Chem. 2001, 3485.
CHAPTER 2
AROMATICITY 85
prepared, for example, and the evidence supported its aromaticity.290 This study suggested that increasing benzoannelation of the parent, 116, led to a step-down in aromaticity, a result of competing ring currents in the annulenic system.
[18]Annulene (115) is diatropic:291 the 12 outer protons are found at = 9 and the 6 inner protons at ^5 = — 3. X-ray crystallography292 shows that it is nearly planar, so that interference of the inner hydrogens is not important in annulenes this large. Compound 115 is reasonably stable, being distillable at reduced pres¬ sures, and undergoes aromatic substitutions.293 The C— C bond distances are not equal, but they do not alternate. There are 12 inner bonds of ^1.38 A and 6 outer bonds of ~1. 42 A.292 Compound 115 has been estimated to have a resonance energy of ^37 kcal mol-1 (155 kJ mol-1), similar to that of benzene.294
The known bridged [18] annulenes are also diatropic295 as are most of the known
"7
dehydro [ 18] annulenes. The dianions of open and bridged [16] annulenes are
also 18-electron aromatic systems,298 and there are dibenzo[ 18] annulenes.299
[22]Annulene300 and dehydro [22] annulene301 are also diatropic. A dehydro- benzo[22]annulene has been prepared that has eight C=C units, is planar and possesses a weak induced ring current.302 In the latter compound there are 13 outer protons at 6.25-8.45 5 and 7 inner protons at 0.70-3.45 5. Some aromatic bridged
290Bodyston, A.J.; Haley, M.M. Org. Lett. 2001, 3, 3599; Boydston, A.J.; Haley, M.M.; Williams, R.V.; Armantrout, J.R. J. Org. Chem. 2002, 67, 8812.
291Jackman, L.M.; Sondheimer, F.; Amiel, Y.; Ben-Efraim, D.A.; Gaoni, Y.; Wolovsky, R.; Bothner-By, A. A. J. Am. Chem. Soc. 1962, 84, 4307; Gilles, J.; Oth, Sondheimer, F.; Woo, E.R J. Chem. Soc. B,
1971, 2111. For a thorough discussion, see Baumann, H.; Oth, J.F.M. Helv. Chim. Acta, 1982, 65, 1885. 292Bregman, J.; Hirshfeld, F.L.; Rabinovich, D.; Schmidt, G.M.J. Acta Crystallogr., 1965, 19, 227; Hirshfeld, F.L.; Rabinovich, D. Acta Crystallogr., 1965, 19, 235.
293Sondheimer, F. Tetrahedron 1970, 26, 3933.
294Oth, J.F.M. ; Biinzli, J.; de Julien de Zelicourt, Y. Helv. Chim. Acta, 1974, 57, 2276.
295For some examples, see DuVernet, R.B.; Wennerstrom, O.; Lawson, J.; Otsubo, T.; Boekelheide, V. J. Am. Chem. Soc. 1978, 100, 2457; Ogawa, H.; Sadakari, N.; Imoto, T.; Miyamoto, I.; Kato, H.; Taniguchi, Y. Angew. Chem. Int. Ed. 1983, 22, 417; Vogel, E.; Sicken, M.; Rohrig, R; Schmickler, H.; Lex, J.; Ermer, O. Angew. Chem. Int. Ed. 1988, 27, 411.
296Okamura, W.H.; Sondheimer, F. J. Am. Chem. Soc. 1967, 89, 5991; Ojima, J.; Ejiri, E.; Kato, T.; Nakamura, M.; Kuroda, S.; Hirooka, S.; Shibutani, M. J. Chem. Soc. Perkin Trans. 1 1987, 831; Sondheimer, F. Acc. Chem. Res. 1972, 5, 81. For two that are not, see Endo, K.; Sakata, Y.; Misumi, S. Bull. Chem. Soc. Jpn. 1971, 44, 2465.
297For a review of this type of polycyclic ion, see Rabinovitz, M.; Willner, I.; Minsky, A. Acc. Chem. Res. 1983, 16, 298.
298Mitchell, R.H.; Boekelheide, V. Chem. Commun. 1970, 1557; Oth, J.F.M.; Baumann, H.; Gilles, J.; Schroder, G. J. Am. Chem. Soc. 1972, 94, 3948. See also Brown, J.M.; Sondheimer, F. Angew. Chem. Int. Ed. 1974, 13, 337; Cresp, T.M.; Sargent, M.V. J. Chem. Soc. Chem. Commun. 1974, 101; Schroder, G.; Plinke, G.; Smith, D.M.; Oth, J.F.M. Angew. Chem. Int. Ed. 1973, 12, 325; Rabinovitz, M.; Minsky, A. Pure Appl. Chem. 1982, 54, 1005.
299Michels, H.P.; Nieger, M.; Vogtle, F. Chem. Ber. 1994, 127, 1167.
300McQuilkin, R.M.; Metcalf, B.W.; Sondheimer, F. Chem. Commun. 1971, 338.
301McQuilkin, R.M.; Sondheimer, F. J. Am. Chem. Soc. 1970, 92, 6341; Iyoda, M.; Nakagawa, M. J. Chem. Soc. Chem. Commun. 1972, 1003. See also, Akiyama, S.; Nomoto, T.; Iyoda, M.; Nakagawa, M. Bull. Chem. Soc. Jpn. 1976, 49, 2579.
302Wan, W.B.; Kimball, D.B.; Haley, M.M. Tetrahedron Lett. 1998, 39, 6795.
86 DELOCALIZED CHEMICAL BONDING
[22]annulenes are also known.303 [26]Annulene has not yet been prepared, but several dehydro [26] annulenes are aromatic.304 Furthermore, the dianion of 1,3,7,9,13,- 15,19,21-octadehydro[24]annulene is another 26-electron system that is aromatic.305 Ojima and co-workers have prepared bridged dehydro derivatives of [26], [30], and [34] annulenes.306 All of these are diatropic. The same workers prepared a bridged tet- radehydro[38]annulene,306 which showed no ring current. On the other hand, the dianion of the cyclophane, 117, also has 38 perimeter electrons, and this species is diatropic.307
There is now no doubt that An + 2 systems are aromatic if they can be planar, although 97 and 113 among others, demonstrate that not all such systems are in fact planar enough for aromaticity. The cases of 109 and 111 prove that absolute planarity is not required for aromaticity, but that aromaticity decreases with decreasing planarity.
H H H H
303For example see Broadhurst, M J.; Grigg, R.; Johnson, A.W. J. Chem. Soc. Perkin Trans. 1 1972, 2111; Ojima, J.; Ejiri, E.; Kato, T.; Nakamura, M.; Kuroda, S.; Hirooka, S.; Shibutani, M. J. Chem. Soc. Perkin Trans. 1 1987, 831; Yamamoto, K.; Kuroda, S.; Shibutani, M.; Yoneyama, Y.; Ojima, J.; Fujita, S.; Ejiri, E.; Yanagihara, K. J. Chem. Soc. Perkin Trans. 1 1988, 395.
304Metcalf, B.W.; Sondheimer, F. J. Am. Chem. Soc. 1971, 93, 5271; Iyoda, M.; Nakagawa, M. Tetrahedron Lett. 1972, 4253; Ojima, J.; Fujita, S.; Matsumoto, M.; Ejiri, E.; Kato, T.; Kuroda, S.; Nozawa, Y.; Hirooka, S.; Yoneyama, Y.; Tatemitsu, H. J. Chem. Soc. Perkin Trans. 1 1988, 385. 305McQuilkin, R.M.; Garratt, P.J.; Sondheimer, F. J. Am. Chem. Soc. 1970, 92, 6682. See also, Huber, W.; Mullen, K.; Wennerstrom, O. Angew. Chem. Int. Ed. 1980, 19, 624.
306Ojima, J.; Fujita, S.; Matsumoto, M.; Ejiri, E.; Kato, T.; Kuroda, S.; Nozawa, Y.; Hirooka, S.; Yoneyama, Y.; Tatemitsu, H. J. Chem. Soc., Perkin Trans. 1 1988, 385.
307Mullen, K.; Unterberg, H.; Huber, W.; Wennerstrom, O.; Norinder, U.; Tanner, D.; Thulin, B. J. Am. Chem. Soc. 1984, 106, 7514.
CHAPTER 2
AROMATICITY 87
The proton NMR C H NMR) spectrum of 118 (called kekulene) showed that in a case where electrons can form either aromatic sextets or larger systems, the sextets are preferred.308 There was initial speculation that kekulene might be superaro- matic , that is, it would show enhanced aromatic stabilization. Recent calculations suggest that there is no enhanced stabilization.309 The 48 n electrons of 118 might, in theory, prefer structure 118a, where each ring is a fused benzene ring, or 118b, which has a [30]annulene on the outside and an [18]annulene on the inside. The XH NMR spectrum of this compound shows three peaks at 5 = 7.94, 8.37, and 10.45 in a ratio of 2:1:1. It is seen from the structure that 118 contains three groups of protons. The peak at 7.94 5 is attributed to the 12 ortho protons and the peak at 8.37 5 to the six external para protons. The remaining peak comes from the six inner protons. If the molecule preferred 118b, we would expect to find this peak upfield, probably with a negative 5, as in the case of 115. The fact that this peak is far downfield indicates that the electrons prefer to be in benzenoid rings. Note that in the case of the dianion of 117, we have the opposite situation. In this ion, the 38-electron system is preferred even though 24 of these must come from the six benzene rings, which therefore cannot have aromatic sextets.
120
121
Phenacenes are a family of “graphite ribbons,” where benzene rings are fused together in an alternating pattern. Phenanthrene is the simplest member of this family and other members include the 22-electron system picene (119); the 26- electron system fulminene (120); and the larger member of this family, the 30 elec¬ tron [7]-phenancene, with seven rings (121). 310 In the series benzene to heptacene, reactivity increases although acene resonance energies per n electron are nearly constant. The inner rings of the “acenes” are more reactive, and calculations shown that those rings are more aromatic than the outer rings, and even more aromatic than benzene itself.311
308Staab, H.A.; Diederich, F. Chem. Ber. 1983, 116, 3487; Staab, H.A.; Diederich, F.; Krieger, C.; Schweitzer, D. Chem. Ber. 1983, 116, 3504. For a similar molecule with 10 instead of 12 rings, see Funhoff, D.J.H.; Staab, H.A. Angew. Chem. Int. Ed. 1986, 25, 742.
309Jiao, H.; Schleyer, P.v.R. Angew. Chem. Int. Ed., 1996, 35, 2383.
310Mallory, F.B.; Butler, K.E.; Evans, A.C.; Mallory, C.W. Tetrahedron Lett. 1996, 37, 7173. 311Schleyer, P.v.R.; Manoharan, M.; Jiao, H.; Stahl, F. Org. Lett. 2001, 3, 3643.
88 DELOCALIZED CHEMICAL BONDING
A super ring molecule is formed by rolling a polyacene molecule into one ring with one edge benzene ring folding into the other. These are called cyclopolyacenes or cyclacenes.312 Although the zigzag cyclohexacenes (122) are highly aromatic (this example is a 22-electron system), the linear cyclohexacenes (e.g., the 24 elec¬ tron 123) are much less aromatic.313
Systems of More Than Ten Electrons: 4 n Electrons224
As we have seen (p. 74), these systems are expected to be not only nonaromatic, but actually antiaromatic.
The [12]annulene 124 has been prepared.314 In solution, 124 undergoes rapid con¬ formational mobility (as do many other annulenes),315 and above — 150°C in this par- tiuclar case, all protons are magnetically equivalent. However, at —170° C the mobility is greatly slowed and the three inner protons are found at 5 while the nine outer protons are at ^6 5. Interaction of the “internal” hydrogens in annulene 124 leads to nonplanarity. Above — 50°C, 124 is unstable and rearranges to 125. Several bridged
o
312Ashton, P.R.; Issacs, N.S.; Kohnke, F.H.; Slawin, A.M.Z.; Spencer, C.M.; Stoddart, J.F.; Williams, D.J. Angew. Chem. Int. Ed. 1988, 27, 966; Ashton, P.R.; Brown, G.R.; Issacs, N.S.; Giuffrida, D.; Kohnke, F.H.; Mathias, J.P.; Slawin, A.M.Z.; Smith, D.R.; Stoddart, J.F.; Williams, D.J. J. Am. Chem. Soc. 1992, 114, 6330; Ashton, P.R.; Girreser, U.; Giuffrida, D.; Kohnke, F.H.; Mathias, J.P.; Raymo, F.M.; Slawin, A.M.Z.; Stoddart, J.F.; Williams, D.J. J. Am. Chem. Soc. 1993, 115, 5422.
313Aihara, J-i. J. Chem. Soc. Perkin Trans. 2 1994, 971.
314Oth, J.F.M.; Rottele, H.; Schroder, G. Tetrahedron Lett. 1970, 61; Oth, J.F.M.; Gilles, J.; Schroder, G. Tetrahedron Lett. 1970, 67.
315For a review of conformational mobility in annulenes, see Oth, J.F.M. Pure Appl. Chem. 1971, 25, 573.
CHAPTER 2
AROMATICITY 89
and dehydro [1 2] annulenes are known, for example, 5-bromo- 1,9-didehydro [12] annu- lene (126), 316 cycl[3.3.3]azine (127), 317 s-indacene (128), 318 and l,7-methano[12]annu- lene (129).319 s-Indacene is a planar, conjugated system perturbed by two cross-links, and studies showed that the low-energy structure has localized double bonds. In these compounds, both hydrogen interference and conformational mobility are prevented. In 127-129, the bridge prevents conformational changes, while in 126 the bromine atom is too large to be found inside the ring. The NMR spectra show that all four compounds are paratropic, the inner proton of 126 being found at 16.4 5. The dication
QA 1
of 112 and the dianion of 103 are also 12-electron paratropic species. An inter¬ esting 12-electron [13]-annulenone has recently been reported. 5,1 0-Dimethyl [1 3] an- nulenone (130) is the first monocyclic annulene larger than tropane,322 and a linearly fused benzodehydro[ 12] annulene system has been reported.323
The results for [16] annulene are similar. The compound was synthesized in two different ways,324 both of which gave 131, which in solution is in equili¬ brium with 132. Above — 50°C there is conformational mobility, resulting in the magnetic equivalence of all protons, but at — 130°C the compound is clearly paratropic: there are 4 protons at 10.56 5 and 12 at 5.35 5. In the solid state, where the compound exists entirely as 131, X-ray crystallography325 shows that the molecules are nonplanar with almost complete bond alternation: the single bonds are 1.44-1.47 A and the double bonds 1.31-1.35 A. A number of dehydro and bridged [16] annulenes are also paratropic,326 as are [20] annulene327 and
316Untch, K.G.; Wysocki, D.C. J. Am. Chem. Soc. 1967, 89, 6386.
317Farquhar, D.; Leaver, D. Chem. Commun. 1969, 24. For a review, see Matsuda, Y.; Gotou, H. Heterocycles 1987, 26, 2151.
318Hertwig, R.H.; Holthausen, M.C.; Koch, W.; Maksic, Z.B. Angew. Chem. Int. Ed. 1994, 33, 1192. 319Vogel, E.; Konigshofen, H.; Mullen, K.; Oth, J.F.M. Angew. Chem. Int. Ed. 1974, 13, 281. See also, Mugnoli, A.; Simonetta, M. J. Chem. Soc. Perkin Trans. 2 1976, 822; Scott, L.T.; Kirms, M.A.; Gunther,
H. ; von Puttkamer, H. J. Am. Chem. Soc. 1983, 105, 1372; Destro, R.; Ortoleva, E.; Simonetta, M.; Todeschini, R. J. Chem. Soc. Perkin Trans. 2 1983, 1227.
320Mullen, K.; Meul, T.; Schade, P.; Schmickler, H.; Vogel, E. J. Am. Chem. Soc. 1987 , 109, 4992. This paper also reports a number of other bridged paratropic 12-, 16-, and 20-electron dianions and dications. See also Hafner, K.; Thiele, G.F. Tetrahedron Lett. 1984, 25, 1445.
321Schmalz, D.; Gunther, H. Angew. Chem. Int. Ed. 1988, 27, 1692.
322Higuchi, H.; Hiraiwa, N.; Kondo, S.; Ojima, J.; Yamamoto, G. Tetrahedron Lett. 1996, 37, 2601. 323Gallagher, M.E.; Anthony, J.E. Tetrahedron Lett. 2001, 42, 7533.
324Schroder, G.; Oth, J.F.M. Tetrahedron Lett. 1966, 4083; Oth, J.F.M. ; Gilles, J. Tetrahedron Lett. 1968, 6259; Calder, I.C.; Gaoni, Y.; Sondheimer, F. J. Am. Chem. Soc. 1968, 90, 4946. For monosubstituted [16] annulenes, see Schroder, G.; Kirsch, G.; Oth, J.F.M. Chem. Ber. 1974, 107, 460.
325Johnson, S.M.; Paul, I.C.; King, G.S.D. J. Chem. Soc. B 1970, 643.
326For example, see Calder, I.C.; Garratt, P.J.; Sondheimer, F. J. Am. Chem. Soc. 1968, 90, 4954; Murata,
I. ; Okazaki, M.; Nakazawa, T. Angew. Chem. Int. Ed. 1971, 10, 576; Ogawa, H.; Kubo, M.; Tabushi, I. Tetrahedron Lett. 1973, 361; Nakatsuji, S.; Morigaki, M.; Akiyama, S.; Nakagawa, M. Tetrahedron Lett. 1975, 1233; Elix, J.A. Aust. J. Chem. 1969, 22, 1951; Vogel, E.; Kurshner, U.; Schmickler, H.; Lex, J.; Wennerstrom, O.; Tanner, D.; Norinder, U.; Kruger, C. Tetrahedron Lett. 1985, 26, 3087.
327Metcalf, B.W.; Sondheimer, F. J. Am. Chem. Soc. 1971, 93, 6675. See also Oth, J.F.M.; Woo, E.P.; Sondheimer, F. J. Am. Chem. Soc. 1973, 95, 7337; Nakatsuji, S.; Nakagawa, M. Tetrahedron Lett. 1975, 3927; Wilcox, Jr., C.F.; Farley, E.N. J. Am. Chem. Soc. 1984, 106, 7195.
90 DELOCALIZED CHEMICAL BONDING
[24]annulene.328 However, a bridged tetradehydro[32]annulene was atropic.306
Both pyracyclene (133)329 (which because of strain is stable only in solution) and dipleiadiene (134)330 are paratropic, as shown by NMR spectra. These mole¬ cules might have been expected to behave like naphthalenes with outer bridges, but the outer n frameworks (12 and 16 electrons, respectively) constitute antiaromatic systems with an extra central double bond. With respect to 133, the An + 2 rule pre¬ dicts pyracylene to be “aromatic” if it is regarded as a 10-Ti-electron naphthalene unit connected to two 2-7i-electron etheno systems, but “antiaromatic” if it is viewed as a 12-Ti-electron cyclododecahexaene periphery perturbed by an internal cross-linked etheno unit.331 Recent studies have concluded on energetic grounds that 133 is a “borderline” case, in terms of aromaticity-antiaromaticity charac¬ ter.329 Dipleiadiene appears to be antiaromatic.330
The fact that many An systems are paratropic, even though they may be nonpla- nar and have unequal bond distances, indicates that if planarity were enforced, the ring currents might be even greater. That this is true is dramatically illustrated by the NMR spectrum of the dianion of 110332 (and its diethyl and dipropyl homo¬ logs).333 We may recall that in 110, the outer protons were found at 8.14-8.67 5 with the methyl protons at —4.25 5. For the dianion, however, which is forced to have approximately the same planar geometry, but now has 16 electrons, the outer protons are shifted to about —35 while the methyl protons are found at ~21 5, a shift of ~25 5! We have already seen where the converse shift was made, when [16]annulenes that were antiaromatic were converted to 18-electron dianions that were aromatic.254 In these cases, the changes in nmr chemical shifts were almost
328Calder, I.C.; Sondheimer, F. Chem. Commun. 1966, 904. See also, Stockel, K.; Sondheimer, F. J. Chem. Soc. Perkin Trans. 1 1972, 355; Nakatsuji, S.; Akiyama, S.; Nakagawa, M. Tetrahedron Lett. 1976, 2623; Yamamoto, K.; Kuroda, S.; Shibutani, M.; Yoneyama, Y.; Ojima, J.; Fujita, S.; Ejiri, E.; Yanagihara, K. J. Chem. Soc., Perkin Trans. 1 1988, 395.
329Trost, B.M.; Herdle, W.B. J. Am. Chem. Soc. 1976, 98, 4080.
330Vogel, E.; Neumann, B.; Klug, W.; Schmickler, H.; Lex, J. Angew. Chem. Int. Ed. 1985, 24, 1046. 331Diogo, H.P.; Kiyobayashi, T.; Minas da Piedade, M.E.; Burlak, N.; Rogers, D.W.; McMasters, D.; Persy, G.; Wirz, J.; Liebman, J.F. J. Am. Chem. Soc. 2002, 124, 2065.
332For a review of polycyclic dianions, see Rabinovitz, M.; Cohen, Y. Tetrahedron 1988, 44, 6957 . 333Mitchell, R.H.; Klopfenstein, C.E.; Boekelheide, V. J. Am. Chem. Soc. 1969, 91, 4931. For another example, see Deger, H.M.; Mullen, K.; Vogel, E. Angew. Chem. Int. Ed. 1978, 17, 957 .
CHAPTER 2
AROMATICITY 91
as dramatic. Heat-of-combustion measures also show that [16]annulene is much less stable than its dianion.334
We can therefore conclude that A n systems will be at a maximum where a mole¬ cule is constrained to be planar (as in 86 or the dianion of 110) but, where possible, the molecule will distort itself from planarity and avoid equal bond distances in order to reduce. In some cases, such as cyclooctatraene, the distortion and bond alternation are great enough to be completely avoided. In other cases, for example, 124 or 131, it is apparently not possible for the molecules to avoid at least some /7-orbital overlap. Such molecules show evidence of paramagnetic ring currents, although the degree of is not as great as in molecules such as 86 or the dianion of 110.
The concept of “Mobius aromaticity” was conceived by Helbronner in 196 4335 when he suggested that large cyclic [4^]annulenes might be stabilized if the Ti-orbitals were twisted gradually around a Mobius strip. This concept is illustrated by the diagrams labeled Hiickel, which is a destabilized [An] system, in contrast to the Mobius model, which is a stabilized [An] system.336 Zimmerman generalized this idea and applied the “Hiickel-Mobius concept” to the analysis of ground-state systems, such as barrelene (135). 337 In 1998, a computational reinterpretation of existing experimental evidence for (CH)^ as a Mobius
334Stevenson, G.R.; Forch, B.E. J. Am. Chem. Soc. 1980, 102, 5985.
335Heilbronner, E. Tetrahedron Lett. 1964, 1923.
336Kawase, T; Oda, M. Angew. Chem. Int. Ed., 2004, 43, 4396.
337Zimmerman, H.E. J. Am. Chem. Soc. 1966, 88, 1564.; Zimmerman, H.E. Acc. Chem. Res. 1972, 4, 272.
92 DELOCALIZED CHEMICAL BONDING
aromatic cyclic annulene with An 7i-electrons was reported.338 A recent computa¬ tional study predicted several Mobius local minima for [12]-, [16]-, and [20]annu- lenes.339 A twisted [16] annulene has been prepared and calculations suggested it should show Mobius aromaticity.340 High-performance liquid chromatography (HPLC) separation of isomers gave 136, which the authors concluded is Mobius aromatic.
Other Aromatic Compounds
We will briefly mention three other types of aromatic compounds.
1. Mesoionic Compounds. 3 41 These compounds cannot be satisfactorily represen¬ ted by Lewis structures not involving charge separation. Most of them contain five-membered rings. The most common are the sydnones , stable aromatic compounds that undergo aromatic substitution when R' is hydrogen.
R'
/
R — N
\
N
O©
R'
\
®/CZ
-N
\\
N'
R'
\
C^.
/ c
R — N © I
\
N
O
o°
Sydnone
2. The Dianion of Squaric Acid.342 The stability of this system is illustrated by the fact that the p^i of squaric acid343 is ~1.5 and the p K2 is ~3.5,344 which means that even the second proton is given up much more readily than the proton of acetic acid, for example.345 The analogous three-,346
338Mauksch, M.; Gogonea, V.; Jiao, H.; Schleyer, Rv.R. Angew. Chem. Int. Ed., 1998, 37, 2395. 339Castro, C.; Isborn, C.M.; Karney, W.L.; Mauksch, M.; Schleyer, Rv.R. Org. Lett. 2002, 4, 3431. 340Ajami, D.; Oeckler, O.; Simon, A.; Herges, R. Nature ( London ) 2003, 426, 819.
341For reviews, see Newton, C.G.; Ramsden, C.A. Tetrahedron 1982, 38, 2965; Ollis, W.D.; Ramsden, C.A. Adv. Heterocycl. Chem. 1976, 19, 1; Ramsden, C.A. Tetrahedron 1977, 33, 3203; Yashunskii, V.G.; Kholodov, L.E. Russ. Chem. Rev. 1980, 49, 28; Ohta, M.; Kato, H., in Snyder, J.P. Nonbenzenoid Aromaticity, Vol. 1, Academic Press, NY, 1969, pp. 117-248.
342West, R.; Powell, D.L. J. Am. Chem. Soc. 1963, 85, 2577; Ito, M.; West, R. J. Am. Chem. Soc. 1963, 85, 2580.
343For a review of squaric acid and other nonbenzenoid quinones, see Wong, H.N.C.; Chan, T.; Luh, T., in Patai, S.; Rappoport, Z. The Chemistry of the Quinonoid Compounds, Vol. 2, pt. 2, Wiley, NY, 1988, pp. 1501-1563.
344Ireland, D.T.; Walton, H.F. J. Phys. Chem. 1967, 71, 75 1 ; MacDonald, D.J. J. Org. Chem. 1968, 33, 4559. 345There has been a controversy as to whether this dianion is in fact aromatic. See Aihara, J. J. Am. Chem. Soc. 1981, 103, 1633.
346Eggerding, D.; West, R. J. Am. Chem. Soc. 1976, 98, 3641; Pericas, M.A.; Serratosa, F. Tetrahedron Lett. 1977, 4437; Semmingsen, D.; Groth, P. J. Am. Chem. Soc. 1987 , 109, 7238.
CHAPTER 2
AROMATICITY 93
five-, and six-membered ring compounds are also known.347
J
OH
OH
O
O
°w°
etc.
O o
oO
Oo
3. Homoaromatic Compounds. When cyclooctatetraene is dissolved in con¬ centrated H2S04, a proton adds to one of the double bonds to form the homotropylium ion 137. 348 In this species, an aromatic sextet is spread over seven carbons, as in the tropylium ion. The eighth carbon is an sp 3 carbon and so cannot take part in the aromaticity. The NMR spectra show the presence of a diatropic ring current: is found at 5 = —0.3; at
5.1 5; Hi and H7 at 6.4 5; H2-H6 at 8.5 5. This ion is an example of a homoaromatic compound, which may be defined as a compound that contains one or more349 ^-hybridized carbon atoms in an otherwise conjugated cycle.350
In order for the orbitals to overlap most effectively so as to close a loop, the sp 3 atoms are forced to lie almost vertically above the plane of the
347For a monograph, see West, R. Oxocarbons\ Academic Press, NY, 1980. For reviews, see Serratosa, F. Acc. Chem. Res. 1983, 16, 170; Schmidt, A.H. Synthesis 1980, 961; West, R. Isr. J. Chem. 1980, 20, 300; West, R.; Niu, J., in Snyder, J.P. Nonbenzenoid Aromaticity, Vol. 1, Academic Press, NY, 1969, pp. 311- 345, and in Zabicky, J. The Chemistry of the Carbonyl Group, Vol. 2, Wiley, NY, 1970, pp. 241-275; Maahs, G.; Hegenberg, P. Angew. Chem. Int. Ed. 1966, 5, 888.
348Rosenberg, J.L.; Mahler, J.E.; Pettit, R. J. Am. Chem. Soc. 1962, 84, 2842; Keller, C.E.; Pettit, R. J. Am. Chem. Soc. 1966, 88, 604, 606; Winstein, S.; Kreiter, C.G.; Brauman, J.I. J. Am. Chem. Soc. 1966, 88, 2047; Haddon, R.C. J. Am. Chem. Soc. 1988, 110, 1108. See also, Childs, R.F.; Mulholland, D.L.; Varadarajan, A.; Yeroushalmi, S. J. Org. Chem. 1983, 48, 1431. See also, Alkorta, I.; Elguero, J.; Eckert- Maksic, M.; Maksic, Z.B. Tetrahedron 2004, 60, 2259.
349If a compound contains two such atoms it is bishomoaromatic; if three, trishomoaromatic, and so on. For examples see Paquette, L.A. Angew. Chem. Int. Ed. 1978, 17, 106.
350For reviews, see Childs, R.F. Acc. Chem. Res. 1984, 17, 347; Paquette, L.A. Angew. Chem. Int. Ed. 1978, 17, 106; Winstein, S. Q. Rev. Chem. Soc. 1969, 23, 141; Garratt, P.J. Aromaticity, Wiley, NY, 1986, pp. 5-45; and in Olah, G.A.; Schleyer, P.v.R. Carbonium Ions, Wiley, NY, Vol. 3, 1972, the reviews by Story, P.R.; Clark, Jr., B.C. 1007-1098, pp. 1073-1093; Winstein, S. 965-1005. (The latter is a reprint of the Q. Rev. Chem. Soc. review mentioned above.)
94 DELOCALIZED CHEMICAL BONDING
aromatic atoms.351 In 137, is directly above the aromatic sextet, and so is shifted far upheld in the nmr. All homoaromatic compounds so far discovered are ions, and it is questionable352 as to whether homoaromatic character can exist in uncharged systems.353 Homoaromatic ions of 2 and 10 electrons are also known.
New conceptual applications to 3D homoaromatic systems with cubane, dodecahedrane, and adamantane frameworks has been presented.354 This concept includes families of spherical homoaromatics with both 2 and 8 mobile electrons. Each set has complete spherical homo aromaticity, that is, all the sp 2 carbon atoms in a highly symmetrical frameworks are separated by one or two sp3 -hybridized atoms.
4. Fullerenes. Fullerenes are a family of aromatic hydrocarbons based on the parent buckminsterfullerene (138; C60)355 that have a variety of very interesting properties.356 Molecular-orbital calculations showed that “fullerene aromaticity lies within 2kcalmol-1 (8.4 kJ mol-1) per carbon of a hypothetical ball of rolled up graphite.357 Another class of polynuclear aromatic hydrocarbons are the buckybowls , which are essen¬ tially fragments of 138. Corannulene (139)358 (also called 5-circulene), for example, is the simplest curved-surface hydrocarbon possessing a carbon framework that is identified with the buckminsterfullerene
35 Calculations show that only ~60% of the chemical shift difference between Hfl and is the result of the aromatic ring current, and that even Ha is shielded; it would appear at 5 ~ 5.5 without the ring current: Childs, R.F.; McGlinchey, M.J.; Varadarajan, A. J. Am. Chem. Soc. 1984, 106, 5974.
352Houk, K.N.; Gandour, R.W.; Strozier, R.W.; Rondan, N.G.; Paquette, L.A. J. Am. Chem. Soc. 1979, 101, 6797; Paquette, L.A.; Snow, R.A.; Muthard, J.L.; Cynkowski, T. J. Am. Chem. Soc. 1979, 101, 6991. See however, Liebman, J.F.; Paquette, L.A.; Peterson, J.R.; Rogers, D.W. J. Am. Chem. Soc. 1986, 108, 8267.
353Examples of uncharged homoantiaromatic compounds have been claimed: Wilcox, Jr., C.F.; Blain, D.A.; Clardy, J.; Van Duyne, G.; Gleiter, R.; Eckert-Maksic, M. J. Am. Chem. Soc. 1986, 108, 7693; Scott, L.T.; Cooney, M.J.; Rogers, D.W.; Dejroongruang, K. J. Am. Chem. Soc. 1988, 110, 7244.
354Chen, Z.; Haijun Jiao, H.; Andreas Hirsch, A.; Schleyer, P.v.R. Angew. Chem. Int. Ed., 2002, 41, 4309
355Billups, W.E.; Ciufolini, M.A. Buckminsterfullerene s, VCH, NY, 1993’, Taylor, R. The Chemistry of Fullerenes, World Scientific, River Edge, NJ, Singapore, 1995’, Aldersey-Williams, H. The Most Beautiful Molecule: The Discovery of the Buckyball, Wiley, NY, 1995’, Baggott, J.E. Perfect Symmetry: the Accidental Discovery of Buckminsterfullerene, Oxford University Press, Oxford, NY, 1994. Also see Kroto, H.W.; Heath, J.R.; O’Brien, S.C.; Curl, R.F.; Smalley, R.E. Nature ( London ) 1985, 318, 162.
356Smalley, R.E. Acc. Chem. Res. 1992, 25, 98; Diederich, F.; Whetten, R.L. Acc. Chem. Res. 1992, 25, 119; Hawkins, J.M. Acc. Chem. Res. 1992, 25, 150; Wudl, F. Acc. Chem. Res. 1992, 25, 157; McElvany, S.W.; Ross, M.M.; Callahan, J.H. Acc. Chem. Res. 1992, 25, 162; Johnson, R.D.; Bethune, D.S.; Yannoni, C.S. Acc. Chem. Res. 1992, 25, 169.
357Warner, P.M. Tetrahedron Lett. 1994, 35, 7173.
358Barth, W.E.; Lawton, R.G. J. Am. Chem. Soc. 1971, 93, 1730; Scott, L.T.; Hashemi, M.M.; Meyer, D.T.; Warren, H.B. J. Am. Chem. Soc. 1991, 113, 7082.
CHAPTER 2
HYPERCONJUGATION 95
surface. It has been synthesized by Scott,352 and several other groups.359 Corannulene is a flexible molecule, with a bowl-to-bowl inversion barrier of ~10— 1 1 kcal mol-1 (41.8-46.0 kJ mol-1).360 Benzocorannulenes are known,361 and other bowl-shaped hydrocarbons include acenaphtho[3,2,l,8- ijklm] diindeno [4 , 3 ,2 , 1 -cdef- 1 ' ,2' , 3' A'pqra] tripheny lene . 362 The inversion barrier to buckybowl inversion has been lowered by such benzannelation of the rim.363 Other semibuckminsterfullerenes include C2v-C3oH12 and C3- C30Hi2.358 Larger fullerenes include C60,C80, C84, and fullerenes are known that contain an endohedral metal, such as scandium or even Sc3N.364 Syn¬ thetic methods often generate mixtures of fullerenes that must be separated, as in the report of new methods for separating C84-fullerenes.365 A homo- fullerene has been prepared.366
HYPERCONJUGATION
All of the delocalization discussed so far involves n electrons. Another type, called hyperconjugation , involves a electrons.367 When a carbon attached
359Borchardt, A.; Fuchicello, A.; Kilway, K.V.; Baldridge, K.K.; Siegel, J.S. J. Am. Chem. Soc. 1992, 114, 1921; Liu, C.Z.; Rabideau, P.W. Tetrahedron Lett. 1996, 37, 3437.
360Biedermann, P.U.; Pogodin, S.; Agranat, I. J. Org. Chem. 1999, 64, 3655; Rabideau, P.W.; Sygula, A. Acc. Chem. Res. 1996, 29, 235; Mehta, G.; Panda, G. Chem. Comm., 1997, 2081; Rabideau, P.W.; Abdourazak, A.H.; Folsom, H.E.; Marcinow, Z.; Sygula, A.; Sygula, R. J. Am. Chem. Soc. 1994, 116, 7891; Hagan, S.; Bratcher, M.S.; Erickson, M.S.; Zimmermann, G.; Scott, L.T. Angew. Chem. Int. Ed., 1997, 36, 406. See also, Dinadayalane, T.C.; Sastry, G.N. Tetrahedron 2003, 59, 8347.
361Dinadayalane, T.C.; Sastry, G.N. J. Org. Chem. 2002, 67, 4605.
362Marcinow, Z.; Grove, D.I.; Rabideau, P.W. J. Org. Chem. 2002, 67, 3537.
363Marcinow, Z.; Sygula, A.; Ellern, D.A.; Rabideau, P.W. Org. Lett. 2001, 3, 3527 .
364Stevenson, S.; Rice, G.; Glass, T.; Harich, K.; Cromer, F.; Jordan, M.R.; Craft, J.; Hadju, E.; Bible, R.; Olmstead, M.M.; Maitra, K.; Fisher, A.J.; Balch, A.L.; Dorn, H.C. Nature ( London ) 1999, 401, 55. 365Wang, G. W.; Saunders, M.; Khong, A.; Cross, RJ. J. Am. Chem. Soc. 2000, 122, 3216.
366Kiely, A.F.; Haddon, R.C.; Meier, M.S.; Selegue, J.P; Brock, C.P.; Patrick, B.O.; Wang, G.-W.; Chen, Y. J. Am. Chem. Soc. 1999, 121, 7971.
367For monographs, see Baker, J.W. Hyperconjugation, Oxford University Press, Oxford, 1952; Dewar, M.J.S. Hyperconjugation, Ronald Press, NY, 1962. For a review, see de la Mare, P.B.D. Pure Appl. Chem. 1984, 56, 1755.
96 DELOCALIZED CHEMICAL BONDING
forms there is no bond at all between the carbon and hydrogen. The effect of 140 on the actual molecule is that the electrons in the C— H bond are closer to the carbon than they would be if 140 did not contribute at all.
R
i
R^
/
R
\
H
-R
© , H R
R
*CVR
R
140
Hyperconjugation in the above case may be regarded as an overlap of the a orbital of the C— H bond and the n orbital of the C— C bond, analogous to the 71-71 orbital overlap previously considered. As might be expected, those who reject the idea of resonance in butadiene (p. 39) believe it even less likely when it involves no-bond structures.
The concept of hyperconjugation arose from the discovery of apparently anomalous electron-release patterns for alkyl groups. By the field effect alone, the order of electron release for simple alkyl groups connected to an unsaturated system is tert- butyl > isopropyl > ethyl > methyl, and this order is observed in many phenomena. Thus, the dipole moments in the gas phase of PhCH3, PhC2H5, PhCH(CH3)2, and PhC(CH3)3 are, respectively, 0.37, 0.58, 0.65, and 0.70 D.368
However, Baker and Nathan369 observed that the rates of reaction with pyridine of /?<2ra-substituted benzyl bromides (see reaction 10-31) were opposite that expected from electron release by the field effect. That is, the methyl-substituted compound reacted fastest and the te/t-butyl- substituted compounded reacted slowest.
CH2Br + C5H5N
©
CH2NC5H5 Br-
This came to be called the Baker-Nathan effect and has since been found in many processes. Baker and Nathan explained it by considering that hyperconjuga- tive forms contribute to the actual structure of toluene:
For the other alkyl groups, hyperconjugation is diminished because the number of C— H bonds is diminished and in tert- butyl there are none; hence, with
368Baker, J.W.; Groves, L.G. J. Chem. Soc. 1939, 1144. 369Baker, J.W.; Nathan, W.S. J. Chem. Soc. 1935, 1840, 1844.
CHAPTER 2
HYPERCONJUGATION 97
respect to this effect, methyl is the strongest electron donor and tert- butyl is the weakest.
However, the Baker-Nathan effect has now been shown not to be caused by hyperconjugation, but by differential solvation.370 This was demonstrated by the finding that in certain instances where the Baker-Nathan effect was found to apply in solution, the order was completely reversed in the gas phase.371 Since the mole¬ cular structures are unchanged in going from the gas phase into solution, it is evi¬ dent that the Baker-Nathan order in these cases is not caused by a structural feature (hyperconjugation), but by the solvent. That is, each alkyl group is solvated to a different extent.372
H H
H H
1
©
-U-
1
-o-
1
- ► H— C~C
1 1
H H
©H H
H H
H H
H— C~C* -
- ► H— C = C
H H
* H H
There is a large body of evidence against hyperconjugation in the ground states of neutral molecules.373 A recent study of the one-bond coupling constants for the aromatic system 141, however, appears to provide the first structural evi¬ dence for hyperconjugation in a neutral ground state.374 In hyperconjugation
x
M = C, Si, Ge, Sn X = N02, CN, H, Me, OMe
©
MMe3
in the ground state of neutral molecules, which Muller and Mulliken call sacrificial hyperconjugation,375 the canonical forms involve not only no-bond resonance, but also a charge separation not possessed by the main form (see 141). For carbocations and free radicals376 and for excited states of molecules,377 there is evidence that hyperconjugation is important. In free radicals and carbocations, the canonical
370This idea was first suggested by Schubert, W.M.; Sweeney, W.A. J. Org. Chem. 1956, 21, 119. 371Hehre,W.J.; Mclver, Jr., R.T.; Pople, J.A.; Schleyer, P.v.R. J. Am. Chem. Soc. 1974, 96, 7162; Arnett, E.M.; Abboud, J.M. J. Am. Chem. Soc. 1975, 97, 3865; Glyde, E.; Taylor, R. J. Chem. Soc. Perkin Trans. 2 1977, 678. See also, Taylor, R. J. Chem. Res. (S), 1985, 318.
372For an opposing view, see Cooney, B.T.; Happer, D.A.R. Aust. J. Chem. 1987, 40, 1537.
373For some evidence in favor, see Laube, T.; Ha, T. J. Am. Chem. Soc. 1988, 110, 5511.
374Lambert, J.B.; Singer, R.A. J. Am. Chem. Soc. 1992, 114, 10246.
375Muller, N.; Mulliken, R.S. J. Am. Chem. Soc. 1958, 80, 3489.
376Symons, M.C.R. Tetrahedron 1962, 18, 333.
377Rao, C.N.R.; Goldman, G.K.; Balasubramanian, A. Can. J. Chem. 1960, 38, 2508.
98 DELOCALIZED CHEMICAL BONDING
forms display no more charge separation than the main form. Muller and Mulliken call this isovalent hyperconjugation : Even here the main form contributes more to the hybrid than the others.
TAUTOMERISM378
There remains one topic to be discussed in our survey of chemical bonding in organic compounds. For most compounds, all the molecules have the same struc¬ ture, whether or not this structure can be satisfactorily represented by a Lewis formula. But for many other compounds there is a mixture of two or more struc¬ turally distinct compounds that are in rapid equilibrium. When this phenomenon, called tautomerism ,379 exists, there is a rapid shift back and forth among the molecules. In most cases, it is a proton that shifts from one atom of a molecule to another.
Keto-Enol Tautomerism380
A very common form of tautomerism is that between a carbonyl compound containing an a hydrogen and its enol form:381 Such equilibria are pH dependent, as in the case of 2-acetylcyclohexanone.382
R'
H i R C
.R2
ii
o
Keto form
i
Ov
H
Enol form
In simple cases (R2 = H, alkyl, OR, etc.) the equilibrium lies well to the left (Table 2.1). The reason can be seen by examining the bond energies in Table 1.7.
378Baker, J.W. Tautomerism ; D. Van Nostrand Company, Inc., New York, 1934; Minkin, V.I.; Olekhnovich, L.R; Zhdanov, Y.A. Molecular Design of Tautomeric Compounds , D. Reidel Publishing Co.: Dordrecht, Holland, 1988.
379For reviews, see Toullec, J. Adv. Phys. Org. Chem. 1982, 18, 1; Kolsov, A.I.; Kheifets, G.M. Russ. Chem. Rev. 1971, 40, 773; 1972, 41, 452-467; Forsen, S.; Nilsson, M., in Zabicky, J. The Chemistry of the Carbonyl Group, Vol. 2, Wiley, NY, 1970, pp. 157-240.
380The mechanism for conversion of one tautomer to another is discussed in Chapter 12 (reaction 12-3).
381Capponi, M.; Gut, I.G.; Hellrung, B.; Persy, G.; Wirz, J. Can. J. Chem. 1999, 77, 605. For a treatise, see Rappoport, Z. The Chemistry of Enols, Wiley, NY, 1990.
382Iglesias, E. J. Org. Chem, 2003, 68, 2680.
CHAPTER 2
TAUTOMERISM 99
TABLE 2.1. The Enol Content of Some Carbonyl Compounds
Compound
Enol Content, %
References
Acetone
6 x 1(T7
383
PhCOCH3
1.1 x 1(T6
384
Cyclopentanone
1 x 1(T6
385
CH3CHO
6 x 10“5
386
Cyclohexanone
4 x 10“5
385
Butanal
5.5 x 1CT4
387
(CH3)2CHCHO
1.4 x 1(T2
388,387
Ph2CHCHO
9.1
389
CH3COOEt
No enol found"
385
CH3COCH2COOEt
8.4
390
CH3COCH2COCH3
80
322
PhCOCH2COCH3
89.2
385
EtOOCCH2COOEt
7.7 x 10“3
385
N=C-CH2COOEt
2.5 x 10-1
385
Indane-l-one
3.3 x 10“8
391
Malonamide
No enol found
392
aLess than 1 part in 10 million.
The keto form differs from the enol form in possessing a C— H, a C— C, and a C=0 bond, where the enol has a C=C, a C— O, and an O— H bond. The approximate sum of the first three is 359kcalmol 1 (1500 kJ mol-1) and of the second three is 347kcalmol-1 (1452 kJ mol-1). The keto form is therefore thermodynamically more stable by ^12kcalmol-1 (48kJmol-1) and enol forms cannot normally be isolated.393 In certain cases, however, a larger amount of the enol form is present,
383Tapuhi, E.; Jencks, W.P. /. Am. Chem. Soc. 1982, 104, 5758; Chiang, Y.; Kresge, A.J.; Tang, Y.S.; Wirz, J. J. Am. Chem. Soc. 1984, 106, 460. See also, Hine, J.; Arata, K. Bull. Chem. Soc. Jpn. 1976, 49, 3089; Guthrie, J.P. Can. J. Chem. 1979, 57, 797, 1177; Dubois, J.E.; El-Alaoui, M.; Toullec, J. J. Am. Chem. Soc. 1981, 103, 5393; Toullec, J. Tetrahedron Lett. 1984, 25, 4401; Chiang, Y.; Kresge, A.J.; Schepp, N.P. J. Am. Chem. Soc. 1989, 111, 3977.
384Keeffe, J.R.; Kresge, A.R.; Toullec, J. Can. J. Chem. 1986, 64, 1224.
385Gero, A. J. Org. Chem. 1954, 19, 469, 1960; Keeffe, J.R., Kresge, A.J.; Schepp, N.P. J. Am. Chem. Soc. 1990, 112, 4862; Iglesias, E. J. Chem. Soc. Perkin Trans. 2 1997, 431. See these papers for values for other simple compounds.
386Chiang, Y.; Hojatti, M.; Keeffe, J.R.; Kresge, A.J.; Schepp, N.P.; Wirz, J. J. Am. Chem. Soc. 1987, 109, 4000. 387Bohne, C.; MacDonald, I.D.; Dunford, H.B. J. Am. Chem. Soc. 1986, 108, 7867.
388Chiang, Y.; Kresge, A.J.; Walsh, PA. J. Am. Chem. Soc. 1986, 108, 6314.
389Chiang, Y.; Kresge, A.J.; Krogh, E.T. J. Am. Chem. Soc. 1988, 110, 2600.
390Moriyasu, M.; Kato, A.; Hashimoto, Y. J. Chem. Soc. Perkin Trans. 2 1986, 515. For enolization of |3- ketoamides, see Hynes, M.J.; Clarke, E.M. J. Chem. Soc. Perkin Trans. 2 1994, 901.
391Jefferson, E.A.; Keeffe, J.R.; Kresge, A.J. J. Chem. Soc. Perkin Trans. 2 1995, 2041.
392Williams, D.L.H.; Xia, L. J. Chem. Soc. Chem. Commun. 1992, 985.
393For reviews on the generation of unstable enols, see Kresge, A.J. Pure Appl. Chem. 1991, 63, 213; Capon, B., in Rappoport, Z. The Chemistry of Enols, Wiley, NY, 1990, pp. 307-322.
1 00 DELOCALIZED CHEMICAL BONDING
and it can even be the predominant form.394 There are three main types of the more stable enols:395
1. Molecules in which the enolic double bond is in conjugation with another double bond. Some of these are shown in Table 2.1. As the table shows, carboxylic esters have a much smaller enolic content than ketones. In molecules like acetoacetic ester (142), the enol is also stabilized by internal hydrogen bonding, which is unavailable to the keto form:
H
h3cx ,OEt
c vc
II I
,0
H
142
2. Molecules that contain two or three bulky aryl groups.396 An example is 2,2- dimesitylethenol (143). In this case the keto content at equilibrium is only 5%. 397 In cases such as this, steric hindrance (p. 230) destabilizes the keto form. In 143, the two aryl groups are ~120° apart, but in 144 they must move closer together (~ 109.5°). Such compounds are often called Fuson-type enols.39S There is one example of an amide with a bulky aryl group [N- methyl bis(2,4,6-triisopropylphenyl)acetamide] that has a measurable enol content, in sharp contrast to most amides.399
Ar =
143 144
394For reviews of stable enols, see Kresge, A.J. Acc. Chem. Res. 1990, 23, 43; Hart, H.; Rappoport, Z.; Biali, S.E., in Rappoport, Z. The Chemistry of Enols, Wiley, NY, 1990, pp. 481-589; Hart, H. Chem. Rev, 1979, 79, 515; Hart, H.; Sasaoka, M. J. Chem. Educ. 1980, 57, 685.
395For some examples of other types, see Pratt, D.V.; Hopkins, P.B. J. Am. Chem. Soc. 1987, 109, 5553; Nadler, E.B.; Rappoport, Z.; Arad, D.; Apeloig, Y. J. Am. Chem. Soc. 1987, 109, 7873.
396For a review, see Rappoport, Z.; Biali, S.E. Acc. Chem. Res. 1988, 21, 442. For a discussion of their structures, see Kaftory, M.; Nugiel, D.A.; Biali, D.A.; Rappoport, Z. J. Am. Chem. Soc. 1989, 111, 8181. 397Biali, S.E.; Rappoport, Z. J. Am. Chem. Soc. 1985, 107, 1007. See also, Kaftory, M.; Biali, S.E.; Rappoport, Z. J. Am. Chem. Soc. 1985, 107, 1701; Nugiel, D.A.; Nadler, E.B.; Rappoport, Z. J. Am. Chem. Soc. 1987, 109, 2112; O’Neill, P; Hegarty, A.F. J. Chem. Soc. Chem. Commun. 1987, 744; Becker, H.; Andersson, K. Tetrahedron Lett. 1987, 28, 1323.
398First synthesized by Fuson, R.C.; see, for example, Fuson, R.C.; South wick, P.L.; Rowland, S.P J. Am. Chem. Soc. 1944, 66, 1109.
399Frey, J.; Rappoport, Z. J. Am. Chem. Soc. 1996, 118, 3994.
CHAPTER 2
TAUTOMERISM 101
3. Highly fluorinated enols, such as 145. 400
OH
X
200°C
o
X
3 h
^ cf3
F2CHC '
145
146
In this case, the enol form is not more stable than the keto form (146). The enol form is less stable, and converts to the keto form upon prolonged heating). It can, however, be kept at room temperature for long periods of time because the tautomerization reaction (12-3) is very slow, owing to the electron- withdrawing power of the fluorines.
Frequently, when the enol content is high, both forms can be isolated. The pure keto form of acetoacetic ester melts at — 39°C, while the enol is a liquid even at — 78°C. Each can be kept at room temperature for days if catalysts, such as acids or bases, are rigorously excluded.401 Even the simplest enol, vinyl alcohol CH2=CHOH, has been prepared in the gas phase at room temperature, where it has a half-life of ~30 min.402 The enol Me2C=CCHOH is indefinitely stable in the solid state at — 78°C and has a half-life of ~24 h in the liquid state at 25°C.403 When both forms cannot be isolated, the extent of enolization is often measured by NMR.404
H
R'
.R
R
i
(X
R
H
R
i
M
O
147
Rx
148
400For a review, see Bekker, R.A.; Knunyants, I.L. Sov. Sci. Rev. Sect. B 1984, 5, 145.
401For an example of particularly stable enol and keto forms, which could be kept in the solid state for more than a year without significant interconversion, see Schulenberg, J.W. J. Am. Chem. Soc. 1968, 90, 7008.
402Saito, S. Chem. Phys. Lett. 1976, 42, 399. See also, Capon, B.; Rycroft, D.S.; Watson, T.W.; Zucco, C. J. Am. Chem. Soc. 1981, 103, 1761; Holmes, J.L.; Lossing, F.R J. Am. Chem. Soc. 1982, 104, 2648; McGarrity, J.F.; Cretton, A.; Pinkerton, A. A.; Schwarzenbach, D.; Flack, H.D. Angew. Chem. Int. Ed. 1983, 22, 405; Rodler, M.; Blom, C.E.; Bauder, A. J. Am. Chem. Soc. 1984, 106, 4029; Capon, B.; Guo, B.; Kwok, F.C.; Siddhanta, A.K.; Zucco, C. Acc. Chem. Res. 1988, 21, 135.
403Chin, C.S.; Lee, S.Y.; Park, J.; Kim, S. J. Am. Chem. Soc. 1988, 110, 8244.
404Cravero, R.M.; Gonzalez-Sierra, M.; Olivieri, A.C. J. Chem. Soc. Perkin Trans. 2 1993, 1067.
1 02 DELOCALIZED CHEMICAL BONDING
The extent of enolization405 is greatly affected by solvent,406 concentration, and temperature. Lactone enols, for example, have been shown to be stable in the gas phase, but unstable in solution 407 Thus, acetoacetic ester has an enol content of 0.4% in water and 19.8% in toluene.408 In this case, water reduces the enol concen¬ tration by hydrogen bonding with the carbonyl, making this group less available for internal hydrogen bonding. As an example of the effect of temperature, the enol content of pentan-2,4-dione, CH3COCH2COCH3, was found to be 95, 68, and 44%, respectively, at 22, 180, and 275°C.409 When a strong base is present, both the enol and the keto form can lose a proton. The resulting anion (the enolate ion ) is the same in both cases. Since 147 and 148 differ only in placement of elec¬ trons, they are not tautomers, but canonical forms. The true structure of the enolate ion is a hybrid of 147 and 148 although 148 contributes more, since in this form the negative charge is on the more electronegative atom.
Other Proton-Shift Tautomerism
In all such cases, the anion resulting from removal of a proton from either tautomer is the same because of resonance. Some examples are:410
1. Phenol-Keto Tautomerism ,411
Phenol Cyclohexadienone
For most simple phenols, this equilibrium lies well to the side of the phenol, since only on that side is there aromaticity. For phenol itself, there is no evidence for the existence of the keto form.412 However, the keto form
405For a review of keto-enol equilibrium constants, see Toullec, J. in Rappoport, Z. The Chemistry of Enols , Wiley, NY, 1990, pp. 323-398.
406For an extensive study, see Mills, S.G.; Beak, R J. Org. Chem. 1985, 50, 1216. For keto-enol tautomerism in aqueous alcohol solutions, see Blokzijl, W.; Engberts, J.B.F.N.; Blandamer, M.J. J. Chem. Soc. Perkin Trans. 2 1994, 455; For theoretical calculations of keto-enol tautomerism in aqueous solutions, see Karelson, M.; Maran, U.; Katritzky, A.R. Tetrahedron 1996, 52, 11325.
407Turec ek, F.; Vivekananda, S.; Sadflek, M.; Polas ek, M. J. Am. Chem. Soc,. 2002, 124, 13282. 408Meyer, K.H. Leibigs Ann. Chem. 1911, 380, 212. See also, Moriyasu, M.; Kato, A.; Hashimoto, Y. J. Chem. Soc. Perkin Trans. 2 1986, 515.
409Hush, N.S.; Livett, M.K.; Peel, J.B.; Willett, G.D. Aust. J. Chem. 1987, 40, 599.
410For a review of the use of X-ray crystallography to determine tautomeric forms, see Furmanova, N.G. Russ. Chem. Rev. 1981, 50, 775.
41 Tor reviews, see Ershov, V.V.; Nikiforov, G.A. Russ. Chem. Rev. 1966, 35, 817; Forsen, S.; Nilsson, M., in Zabicky, J. The Chemistry of the Carbonyl Group, Vol. 2, Wiley, NY, 1970, pp. 168-198.
412Keto forms of phenol and some simple derivatives have been generated as intermediates with very short lives, but long enough for spectra to be taken at 77 K. Lasne, M.; Ripoll, J.; Denis, J. Tetrahedron Lett. 1980, 21, 463. See also, Capponi, M.; Gut, I.; Wirz, J. Angew. Chem. Int. Ed. 1986, 25, 344.
CHAPTER 2
TAUTOMERISM 103
becomes important and may predominate: (7) where certain groups, such as a second OH group or an N=0 group, are present;413 (2) in systems of fused aromatic rings;414 (3) in heterocyclic systems. In many heterocyclic com¬ pounds in the liquid phase or in solution, the keto form is more stable,415 although in the vapor phase the positions of many of these equilibria are reversed.416 For example, in the equilibrium between 4-pyridone (149) and 4- hydroxypyridine (150), 149 is the only form detectable in ethanolic solution, while 150 predominates in the vapor phase.416 In other heterocycles, the hydroxy-form predominates. 2-Hydroxypyridone (151) and pyridone-2-thiol (153)417 are in equilibrium with their tautomers, 2-pyridone 152 and pyridine- 2-thione 154, respectively. In both cases, the most stable form is the hydroxy tautomer, 151 and 153 418
O OH
153 154
2. Nitroso-Oxime Tautomerism.
O-H _ ► O
H2C=n ^ - H3C — N
The equiblirum shown for formaldhyde oxime and nitrosomethane illus¬ trates this process.419 In molecules where the products are stable, the equilibrium lies far to the right, and as a rule nitroso compounds are stable only when there is not a hydrogen.
413Ershov, V.V.; Nikiforov, G.A. Russ. Chem. Rev. 1966, 35, 817. See also, Highet, R.J.; Chou, F.E. J. Am. Chem. Soc. 1977, 99, 3538.
414See, for example, Majerski, Z.; Trinajstic, N. Bull. Chem. Soc. Jpn. 1970, 43, 2648.
415For a monograph on tautomerism in heterocyclic compounds, see Elguero, J.; Marzin, C.; Katritzky, A.R.; Linda, R The Tautomerism of Heterocycles, Academic Press, NY, 1976. For reviews, see Katritzky, A.R.; Karelson, M.; Harris, P.A. Heterocycles 1991, 32, 329; Beak, P Acc. Chem. Res. 1977, 10, 186; Katritzky, A.R. Chimia, 1970, 24, 134.
416Beak, P; Fry, Jr., F.S.; Lee, J.; Steele, F. J. Am. Chem. Soc. 1976, 98, 171.
417Moran, D.; Sukcharoenphon, K.; Puchta, R.; Schaefer III, H.F.; Schleyer, P.v.R.; Hoff, C.D. J. Org. Chem. 2002, 67, 9061.
418Parchment, O.G.; Burton, N.A.; Hillier, I.H.; Vincent, M.A. J. Chem. Soc. Perkin Trans. 2 1993, 861. 419Long, J.A.; Harris, N.J.; Lammertsma, K. J. Org. Chem. 2001, 66, 6762.
1 04 DELOCALIZED CHEMICAL BONDING
3. Aliphatic Nitro Compounds Are in Equilibrium with Aci Forms.
0
© /
r2ch — N^
,o
o
r2c =
,OH
O ©
Nitro form
Aci form
The nitro form is much more stable than the aci form in sharp contrast to the parallel case of nitroso-oxime tautomerism, undoubtedly because the nitro form has resonance not found in the nitroso case. Aci forms of nitro compounds are also called nitronic acids and azinic acids.
4. Imine-Enamine Tautomerism.420
R2CH — CR=NR
R2C=CR — NHR
Imine
Enamine
Enamines are normally stable only when there is no hydrogen on the nitrogen (R2C=CR— NR2). Otherwise, the imine form predominates.421 The energy of various imine-enamine tautomers has been calculated.422 In the case of 6- aminofulvene-l-aldimines, tautomerism was observed in the solid state, as well as in solution.423
5. Ring-Chain Tautomerism. Ring-chain tautomerism424 occurs in sugars (alde¬ hyde vs. the pyranose or furanose structures), and in y-oxocarboxylic acids 425 In benzamide carboxaldehyde, 156, whose ring-chain tautomer is 155, the equilibrium favors the cyclic form (156) 426 Similarly, benzoic acid 2-carboxy aldehyde (157) exists largely as the cyclic form (158) 427 In these latter cases, and in many others, this tautomerism influences chemical reactivity. Conversion of 157 to an ester, for example, is difficult since most standard methods lead to the OR derivative of 158 rather than the ester of 157. Ring-chain tautomerism also occurs in spriooxathianes,428 and in
420For reviews, see Shainyan, B.A.; Mirskova, A.N. Russ. Chem. Rev. 1979, 48, 107; Mamaev, V.R;
Lapachev, V.V. Sov. Sci. Rev. Sect. B. 1985, 7, 1. The second review also includes other closely related
types of tautomerization.
421For examples of the isolation of primary and secondary enamines, see Shin, C.; Masaki, M.; Ohta, M.
Bull. Chem. Soc. Jpn. 1971, 44, 1657; de Jeso, B.; Pommier, J. J. Chem. Soc. Chem. Commun. 1977, 565.
422Lammertsma, K.; Prasad, B.V. J. Am. Chem. Soc. 1994, 116, 642.
423Sanz, D.; Perez-Torralba, M.; Alarcon, S.H.; Claramunt, R.M.; Foces-Foces, C.; Elguero, J. J. Org.
Chem. 2002, 67, 1462.
424For a monograph, see Valters, R.E.; Flitsch, W. Ring-Chain Tautomerism, Plenum, NY, 1985. For
reviews, see Valters, R.E. Russ. Chem. Rev. 1973, 42, 464; 1974, 43, 665; Escale, R.; Verducci, J. Bull.
Soc. Chim. Fr., 1974, 1203.
425Fabian, W.M.F.; Bowden, K. Eur. J. Org. Chem. 2001, 303.
426Bowden, K.; Hiscocks, S.P.; Perjessy, A. J. Chem. Soc. Perkin Trans. 2 1998, 291.
427Ring chain tautomer of benzoic acid 2-carboxaldehdye.
428Terec, A.; Grosu, I.; Muntean, L.; Toupet, L.; Pie, G.; Socaci, C.; Mager, S. Tetrahedron 2001, 57,
8751; Muntean, L.; Grosu, I.; Mager, S.; Pie, G.; Balog, M. Tetrahedron Lett. 2000, 41, 1967 .
CHAPTER 2
TAUTOMERISM 105
decahydroquinazolines, such as 159 and 160, 429 as well as other 1,3-hetero- cycles.430
159 160
There are many other highly specialized cases of proton-shift tautomerism, including an internal Michael reaction (see 15-24) in which 2-(2,2-dicyano-l- methylethenyl)benzoic acid (161) exists largely in the open chain form rather an its tautomer (162) in the solid state, but in solution there is an increasing amount of 162 as the solvent becomes more polar 431
Valence Tautomerism
This type of tautomerism is discussed on p. 105.
429Lazar, L.; Goblyos, A.; Martinek, T.A.; Fulop, F. J. Org. Chem. 2002, 67, 4734.
430Lazar, L.; Fulop, F. Eur. J. Org. Chem. 2003, 3025.
431Kolsaker, P.; Arukwe, J.; Barcoczy, J.; Wiberg, A.; Fagerli, A.K. Acta Chem. Scand. B 1998, 52, 490.
CHAPTER 3
Bonding Weaker than Covalent
In the first two chapters, we discussed the structure of molecules each of which is an aggregate of atoms in a distinct three-dimensional (3D) arrangement held together by bonds with energies on the order of 50-100 kcal mol-1 (200-400 kJ mol-1). There are also very weak attractive forces between molecules, on the order of a few tenths of a kilocalorie per mole. These forces, called van der Waals forces, are caused by electrostatic attractions, such as those between dipole and dipole, induced dipole, and induced dipole, and are responsible for liquefaction of gases at sufficiently low temperatures. The bonding discussed in this chapter has energies of the order of 2-10 kcal mol-1 (9-40 kJ mol-1), intermediate between the two extremes, and produces clusters of molecules. We will also discuss compounds in which portions of molecules are held together without any attractive forces at all.
HYDROGEN BONDING
A hydrogen bond is a bond between a functional group A— H and an atom or group of atoms B in the same or a different molecule.1 With exceptions to be noted later, hydrogen bonds are assumed to form only when A is oxygen, nitrogen, or fluorine and when B is oxygen, nitrogen, or fluorine.2 The oxygen may be singly or doubly
1 For a treatise, see Schuster, P.; Zundel, G.; Sandorfy, C. The Hydrogen Bond, 3 vols., North-Holland Publishing Co.: Amsterdam, The Netherlands, 1976. For a monograph, see Joesten, M.D.; Schaad, L.J. Hydrogen Bonding ; Marcel Dekker, NY, 1974. For reviews, see Meot-Ner, M. Mol. Struct. Energ. 1987, 4, 71; Deakyne, C.A. Mol. Struct. Energ. 1987, 4, 105; Joesten, M.D. J. Chem. Educ. 1982, 59, 362; Gur’yanova, E.N.; Gol’dshtein, I.P.; Perepelkova, T.I. Russ. Chem. Rev. 1976, 45, 792; Pimentel, G.C.; McClellan, A.L. Annu. Rev. Phys. Chem. 1971, 22, 347; Kollman, PA.; Allen, L.C. Chem. Rev. 1972, 72, 283; Huggins, M.L. Angew. Chem. Int. Ed. 1971, 10, 147; Rochester, C.H., in Patai, S. The Chemistry of the Hydroxyl Group, pt. 1; Wiley, NY, 1971, pp. 327-392, 328-369. See also Hamilton, W.C.; Ibers, J.A. Hydrogen Bonding in Solids, W.A. Benjamin, NY, 1968. Also see, Chen, J.; McAllister, M.A.; Lee, J.K.; Houk, K.N. J. Org. Chem. 1998, 63, 4611 for a discussion of short, strong hydrogen bonds.
2The ability of functional groups to act as hydrogen bond acids and bases can be obtained from either equilibrium constants for 1:1 hydrogen bonding or overall hydrogen bond constants. See Abraham, M.H.; Platts, J.A. J. Org. Chem. 2001, 66, 3484.
March ’s Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, Sixth Edition, by Michael B. Smith and Jerry March Copyright © 2007 John Wiley & Sons, Inc.
106
CHAPTER 3
HYDROGEN BONDING 107
bonded and the nitrogen singly, doubly, or triply bonded. The bonds are usually represented by dotted or dashed lines, as shown in the following examples:
MeN MeN
O H — F or 0<
Me Me
O Me
H ^O-H
H— N - H
H
Me. Me P_H -
'H-F Y Me-T / Me
°-HV€H3 °--H-°
F\
F— ^ — F H
F
MeN
O'
i
H
,Me
O
Hydrogen bonds can exist in the solid3 and liquid phases and in solution.4 Many organic reactions that will be discussed in later chapters can be done in aqueous media,5 and their efficacy is due, in part, to the hydrogen bonding nature of aqueous media.6 Even in the gas phase, compounds that form particularly strong hydrogen bonds may remain associated.7 Acetic acid, for example, exists in the gas phase as a dimer, as shown above, except at very low pressures.8 In solution and in the liquid phase, hydrogen bonds rapidly form and break. The mean lifetime of the NH3m*H20 bond is 2 x 10“ 12 s.9 Except for a few very strong hydrogen bonds,10 such as the FH"*F“bond (which has an energy of ~ 50kcal mol-1 or 210 kJ mol-1), the strongest hydrogen bonds are the FH"*F bond and the bonds connecting one carboxylic acid with another. The energies of these bonds are in the range of 6-8 kcal mol-1 or 25- 30 kJ mol-1 (for carboxylic acids, this refers to the energy of each bond). In general, short contact hydrogen bonds between fluorine and HO or NH are rare.11 Other OH—O and NH—N bonds12 have energies of 3-6 kcal mol-1 (12-25 kJ mol-1).
3Steiner, T. Angew. Chem. Int. Ed. 2002, 41, 48. See also Damodharan, L.; Pattabhi, V. Tetrahedron Lett. 2004, 45, 9427.
4See Nakahara, M.; Wakai, C. Chem. Lett. 1992, 809 for a discussion of monomeric and cluster states of water molecules in organic solvents due to hydrogen bonding.
5Li, C.-J.; Chen, T.-H. Organic Reactions in Aqueous Media, Wiley, NY, 1997.
6Li, C.-J. Chem. Rev. 1993, 93, 2023.
vFor a review of energies of hydrogen bonds in the gas phase, see Curtiss, L.A.; Blander, M. Chem. Rev. 1988, 88, 827.
8For a review of hydrogen bonding in carboxylic acids and acid derivatives, see Hadzi, D.; Detoni, S., in Patai, S. The Chemistry of Acid Derivatives, pt. 1, Wiley, NY, 1979, pp. 213-266.
9Emerson, M.T.; Grunwald, E.; Kaplan, M.L.; Kromhout, R.A. J. Am. Chem. Soc. 1960, 82, 6307.
10For a review of very strong hydrogen bonding, see Emsley, J. Chem. Soc. Rev. 1980, 9, 91.
“Howard, J.A.K.; Hoy, V.J.; O’Hagan, D.; Smith, G.T. Tetrahedron 1996, 52, 12613.
12For an ab initio study of diamine hydrogen bonds see Sorensen, J.B.; Lewin, A.H.; Bowen, J.P. J. Org. Chem. 2001, 66, 4105.
1 08 BONDING WEAKER THAN COVALENT
The intramolecular 0-H#~N hydrogen bond in hydroxy amines is also rather strong.13
To a first approximation, the strength of hydrogen bonds increases with increas¬ ing acidity of A— H and basicity of B, but the parallel is far from exact.14 A quan¬ titative measure of the strengths of hydrogen bonds has been established, involving the use of an a scale to represent hydrogen-bond donor acidities and a (3 scale for hydrogen-bond acceptor basicities.15 The use of the (3 scale, along with another parameter, allows hydrogen-bond basicities to be related to proton-transfer basi¬ cities (p K values).16 A database has been developed to locate all possible occur¬ rences of bimolecular cyclic hydrogen-bond motifs in the Cambridge Structural Database,17 and donor-acceptor as well as polarity parameters have been calculated for hydrogen-bonding solvents.18
When two compounds whose molecules form hydrogen bonds with each other are both dissolved in water, the hydrogen bond between the two molecules is usual¬ ly greatly weakened or completely removed,19 because the molecules generally form hydrogen bonds with the water molecules rather than with each other, espe¬ cially since the water molecules are present in such great numbers. In amides, the oxygen atom is the preferred site of protonation or complexation with water.20 In the case of dicarboxylic acids, arguments have been presented that there is little or no evidence for strong hydrogen bonding in aqueous solution,21 although recent studies concluded that strong, intramolecular hydrogen bonding can exist in aqu¬ eous acetone solutions (0.31 mole-fraction water) of hydrogen maleate and hydro¬ gen c A-cyclohexane- 1 ,2-dicarboxylate.22
Many studies have been made of the geometry of hydrogen bonds,23 and the evi¬ dence shows that in most (though not all) cases, the hydrogen is on or near the
13Grech, E.; Nowicka-Scheibe, J.; Olejnik, Z.; Lis, T.; Paweka, Z.; Malarski, Z.; Sobczyk, L. J. Chem. Soc., Perkin Trans. 2 1996 , 343. See Steiner, T. J. Chem. Soc., Perkin Trans. 2 1995, 1315 for a discussion of hydrogen bonding in the crystal structure of a-amino acids.
14For reviews of the relationship between hydrogen-bond strength and acid-base properties, see Pogorelyi, V.K.; Vishnyakova, T.B. Russ. Chem. Rev. 1984, 53, 1 154; Epshtein, L.M. Russ. Chem. Rev. 1979, 48, 854. 15For reviews, see Abraham, M.H.; Doherty, R.M.; Kamlet, M.J.; Taft, R.W. Chem. Br. 1986, 551; Kamlet, M.J.; Abboud, J.M.; Taft, R.W. Prog. Phys. Org. Chem. 1981, 13, 485. For a comprehensive table and a and P values, see Kamlet, M.J.; Abboud, J.M.; Abraham, M.H.; Taft, R.W. J. Org. Chem. 1983, 48, 2877. For a criticism of the P scale, see Laurence, C.; Nicolet, R; Helbert, M. J. Chem. Soc., Perkin Trans. 2 1986, 1081. See also Nicolet, P.; Laurence, C.; Lu£on, M. J. Chem. Soc., Perkin Trans. 2 1987, 483; Abboud, J.M.; Roussel, C.; Gentric, E.; Sraidi, K.; Lauransan, J.; Guiheneuf, G.; Kamlet, M.J.; Taft, R.W. J. Org. Chem. 1988, 53, 1545; Abraham, M.H.; Grellier, P.L.; Prior, D.V.; Morris, J.J.; Taylor, PJ. J. Chem. Soc., Perkin Trans. 2 1990, 521.
16Kamlet, M.J.; Gal, J.; Maria, R; Taft, R.W. J. Chem. Soc., Perkin Trans. 2 1985, 1583.
17 Allen, F.H.; Raithby, P.R.; Shields, G.P.; Taylor, R. Chem. Commun. 1998, 1043.
18Joerg, S.; Drago, R.S.; Adams, J. J. Chem. Soc., Perkin Trans. 2 1997, 2431.
19Stahl, N.; Jencks, W.P. J. Am. Chem. Soc. 1986, 108, 4196.
20Scheiner, S.; Wang, L. J. Am. Chem. Soc. 1993, 115, 1958.
21Perrin, C.L. Annu. Rev. Phys. Org. Chem. 1997, 48, 511.
22Lin, J.; Frey, P.A. J. Am. Chem. Soc. 2000, 122, 11258.
23For reviews, see Etter, M.C. Acc. Chem. Res. 1990, 23, 120; Taylor, R.; Kennard, O. Acc. Chem. Res. 1984, 17, 320.
CHAPTER 3
HYDROGEN BONDING 109
straight line formed by A and B.24 This is true both in the solid state (where X-ray crystallography and neutron diffraction have been used to determine structures),25 and in solution.26 It is significant that the vast majority of intramolecular hydrogen bonding occurs where six-membered rings (counting the hydrogen as one of the six) can be formed, in which linearity of the hydrogen bond is geometrically favorable, while five-membered rings, where linearity is usually not favored (though it is known), are much rarer. A novel nine-membered intramolecular hydrogen bond has been reported.27
In certain cases, X-ray crystallography has shown that a single H-A can form simultaneous hydrogen bonds with two B atoms ( bifurcated or three-center hydro¬ gen bonds). An example is an adduct (1) formed from pentane-2, 4-dione (in its enol form; see p. 98) and diethylamine, in which the O-H hydrogen simultaneously bonds28 to an O and an N (the N-H hydrogen forms a hydrogen bond with the O of another pentane-2, 4-dione molecule).29 On the other hand, in the adduct (2) formed from 1,8-biphenylenediol and hexamethylphosphoramide (HMPA), the B atom (in this case oxygen) forms simultaneous hydrogen bonds with two A— H hydrogens.30 Another such case is found in methyl hydrazine carboxylate 3.31 Except for the special case of FH*"F“ bonds (see p. 107), the hydrogen is not equi¬ distant between A and B. For example, in ice the O-H distance is 0.97 A, while the H***0 distance is 1.79 A.32 A theoretical study of the vinyl alcohol-vinyl alcoholate system concluded the hydrogen bonding is strong, but asymmetric.33 The hydrogen bond in the enol of malonaldehyde, in organic solvents, is asymmetric with the hydrogen atom closer to the basic oxygen atom.34 There is recent evidence, how¬ ever, that symmetrical hydrogen bonds to carboxylates should be regarded as two- center rather than three-center hydrogen bonds, since the criteria traditionally used to infer three-center hydrogen bonding are inadequate for carboxylates.35 There is
24See Stewart, R. The Proton: Applications to Organic Chemistry ; Academic Press, NY, 1985, pp. 148-153. 25A statisical analysis of X-ray crystallographic data has shown that most hydrogen bonds in crystals are nonlinear by ~10-15°: Kroon, J.; Kanters, J.A.; van Duijneveldt-van de Rijdt, J.G.C.M.; van Duijneveldt, F.B.; Vliegenthart, J.A. J. Mol. Struct. 1975, 24, 109. See also, Ceccarelli, C.; Jeffrey, G.A.; Taylor, R. J. Mol. Struct. 1981, 70, 255; Taylor, R.; Kennard, O.; Versichel, W. J. Am. Chem. Soc. 1983, 105, 5761; 1984, 106, 244.
26For reviews of a different aspect of hydrogen-bond geometry: the angle between A###H###B and the rest of the molecule, see Legon, A.C.; Millen, DJ. Chem. Soc. Rev. 1987, 16, 467, Acc. Chem. Res. 1987, 20, 39. 27Yoshimi, Y.; Maeda, H.; Sugimoto, A.; Mizuno, K. Tetrahedron Lett. 2001, 42, 2341.
28Emsley, J.; Freeman, N.J.; Parker, R.J.; Dawes, H.M.; Hursthouse, M.B. J. Chem. Soc., Perkin Trans. 1 1986, 471.
29For some other three-center hydrogen bonds, see Taylor, R.; Kennard, O.; Versichel, W. J. Am. Chem. Soc. 1984, 106, 244; Jeffrey, G.A.; Mitra, J. J. Am. Chem. Soc. 1984, 106, 5546; Staab, H.A.; Elbl, K.; Krieger, C. Tetrahedron Lett. 1986, 27, 5719.
30Hine, J.; Hahn, S.; Miles, D.E. J. Org. Chem. 1986, 51, 577.
31Caminati, W.; Fantoni, A.C.; Schafer, L.; Siam, K.; Van Alsenoy, C. J. Am. Chem. Soc. 1986, 108, 4364. 32Pimentel, G.C.; McClellan, A.L. The Hydrogen Bond ; W.H. Freeman: San Francisco, 1960, p. 260. 33Chandra, A.K.; Zeegers-Huyskens, T., J. Org. Chem. 2003, 68, 3618.
34Perrin, C.L.; Kim, Y.-J. J. Am. Chem. Soc. 1998, 120, 12641.
35Gorbitz, C.H.; Etter, M.C. J. Chem. Soc., Perkin Trans. 2 1992, 131.
110 BONDING WEAKER THAN COVALENT
also an example of cooperative hydrogen bonding (0-H*"C=C-H"*Ph) in crystal¬ line 2-ethynyl-6,8-diphenyl-7/7-benzocyclohepten-7-ol (4). 36
NMe?
12 3 4
Hydrogen bonding has been detected in many ways, including measurements of dipole moments, solubility behavior, freezing-point lowering, and heats of mixing, but one important way is by the effect of the hydrogen bond on IR.37 The IR fre¬ quencies of groups, such as O-H or C=0, are shifted when the group is hydrogen bonded. Hydrogen bonding always moves the peak toward lower frequencies, for both the A— H and the B groups, though the shift is greater for the former. For example, a free OH group of an alcohol or phenol absorbs at ^3590-3650 cm-1, while a hydrogen-bonded OH group is found ~50-100 cm-1 lower.38 In many cases, in dilute solution, there is partial hydrogen bonding, that is, some OH groups are free and some are hydrogen bonded. In such cases, two peaks appear. Infrared spec¬ troscopy can also distinguish between inter- and intramolecular hydrogen bonding, since intermolecular peaks are intensified by an increase in concentration while intramolecular peaks are unaffected. Other types of spectra that have been used for the detection of hydrogen bonding include Raman, electronic,39 and NMR.40 Since hydrogen bonding involves a rapid movement of protons from one atom to another, nmr records an average value. Hydrogen bonding can be detected because it usually produces a chemical shift to a lower field. For example, carboxylic acid-carboxylate systems arising from either mono- or diacids generally exhibit a downfield resonance (16-22 ppm), which indicates “strong” hydrogen bonding
36Steiner, T.; Tamm, M.; Lutz, B.; van der Maas, J. Chem. Commun. 1996, 1127.
37For reviews of the use of ir spectra to detect hydrogen bonding, see Symons, M.C.R. Chem. Soc. Rev. 1983, 12, 1; Egorochkin, A.N.; Skobeleva, S.E. Russ. Chem. Rev. 1979, 48, 1198; Tichy, M. Adv. Org. Chem. 1965, 5, 115; Ratajczak, H.; Orville-Thomas, WJ. J. Mol. Struct. 1968, 1, 449. For a review of studies by ir of the shapes of intramolecular hydrogen-bonded compounds, see Aaron, H.S. Top. Stereochem. 1979, 11, 1. For a review of the use of rotational spectra to study hydrogen bonding, see Legon, A.C. Chem. Soc. Rev. 1990, 19, 197.
38Tichy, M. Adv. Org. Chem. 1965, 5, 1 15 contains a lengthy table of free and intramolecularly hydrogen¬ bonding peaks.
39For a discussion of the effect of hydrogen bonding on electronic spectra, see Lees, W.A.; Burawoy, A. Tetrahedron 1963, 19, 419.
40For a review of the use of nmr to detect hydrogen bonding, see Davis, Jr., J.C.; Deb, K.K. Adv. Magn. Reson. 1970, 4, 201. Also see, Kumar, G.A.; McAllister, M.A. J. Org. Chem. 1998, 63, 6968, which shows the relationship between XH NMR chemical shift and hydrogen bond strength.
CHAPTER 3
HYDROGEN BONDING 111
in anhydrous, aprotic solvents.41 Hydrogen bonding changes with temperature and concentration, and comparison of spectra taken under different conditions also serves to detect and measure it. As with IR spectra, intramolecular hydrogen bond¬ ing can be distinguished from intermolecular by its constancy when the concentra¬ tion is varied. The spin-spin coupling constant across a hydrogen bond, obtained by NMR studies, has been shown to provide a “fingerprint” for hydrogen-bond type.42
Hydrogen bonds are important because of the effects they have on the properties of compounds, among them:
1. Intermolecular hydrogen bonding raises boiling points and frequently melting points.
2. If hydrogen bonding is possible between solute and solvent, this greatly increases solubility and often results in large or even infinite solubility where none would otherwise be expected.
3. Hydrogen bonding causes lack of ideality in gas and solution laws.
4. As previously mentioned, hydrogen bonding changes spectral absorption positions.
5. Hydrogen bonding, especially the intramolecular variety, changes many chemical properties. For example, it is responsible for the large amount of enol present in certain tautomeric equilibria (see p. 98). Also, by influencing the conformation of molecules (see Chapter 4), it often plays a significant role in determining reaction rates.43 Hydrogen bonding is also important in maintaining the 3D structures of protein and nucleic acid molecules.
Besides oxygen, nitrogen, and fluorine, there is evidence that weaker hydrogen bonding exists in other systems 44 Although many searches have been made for hydrogen bonding where A is carbon,45 only three types of C-H bonds have been found that are acidic enough to form weak hydrogen bonds 46 These are found in terminal alkynes, RC=CH,47 chloroform and some other halogenated alkanes, and HCN. Sterically unhindered C-H groups (CHC13, CH2C12, RC=CH) form short contact hydrogen bonds with carbonyl acceptors, where there is a significant preference for coordination with the conventional carbonyl lone-pair direction.48
41Bruck, A.; McCoy, L.L.; Kilway, K.V. Org. Lett. 2000, 2, 2007.
42Del Bene, J.E.; Perera, S.A.; Bartlett, R.J. J. Am. Chem. Soc. 2000, 122, 3560.
43For reviews of the effect of hydrogen bonding on reactivity, see Hibbert, F.; Emsley, J. Adv. Phys. Org. Chem. 1990, 26, 255; Sadekov, I.D.; Minkin, V.I.; Lutskii, A.E. Russ. Chem. Rev. 1970, 39, 179.
44For a review, see Pogorelyi, V.K. Russ. Chem. Rev. 1977, 46, 316.
45For a monograph on this subject, see Green, R.D. Hydrogen Bonding by C-H Groups; Wiley, NY, 1974. See also Taylor, R.; Kennard, O. J. Am. Chem. Soc. 1982, 104, 5063; Harlow, R.L.; Li, C.; Sammes, M.P. J. Chem. Soc., Perkin Trans. 1 1984, 547; Nakai, Y.; Inoue, K.; Yamamoto, G.; O ki, M. Bull. Chem. Soc. Jpn. 1989, 62, 2923; Seiler, P; Dunitz, J.D. Helv. Chim. Acta 1989, 72, 1125.
46For a theoretical study of weak hydrogen-bonds, see Calhorda, M.J. Chem. Commun. 2000, 801. 47For a review, see Hopkinson, A.C., in Patai, S. The Chemistry of the Carbon-Carbon Triple Bond, pt. 1, Wiley, NY, 1978, pp. 75-136. See also DeLaat, A.M.; Ault, B.S. J. Am. Chem. Soc. 1987, 109, 4232. 48Streiner, T.; Kanters, J.A.; Kroon, J. Chem. Commun. 1996, 1277.
112 BONDING WEAKER THAN COVALENT
Weak hydrogen bonds are formed by compounds containing S-H bonds.49 There has been much speculation regarding other possibilities for B. There is evidence that Cl can form weak hydrogen bonds,50 but Br and I form very weak bonds if at all.51 How¬ ever, the ions Cl-, Br-, and I- form hydrogen bonds that are much stronger than those of the covalently bonded atoms.52 As we have already seen, the FH--F bond is especially strong. In this case, the hydrogen is equidistant from the fluor¬ ines.53 Similarly, a sulfur atom49 can be the B component in weak hydrogen bonds,54 but the -SH ion forms much stronger bonds.55 There are theoretical studies of weak hydrogen bonding.56 Hydrogen bonding has been directly observed (by NMR and IR) between a negatively charged carbon (see Carbanions, Chapter 5) and an OH group in the same molecule.57 Another type of molecule in which carbon is the B component are isocyanides, R-+N=C- which form rather strong hydrogen bonds.58 There is evidence that double and triple bonds, aro¬ matic rings,59 and even cyclopropane rings60 may be the B component of hydro¬ gen bonds, but these bonds are very weak. An interesting case is that of the m-bicyclo[4.4.4]-l-tetradecyl cation 5 (see in-out isomerism, p. 189). The NMR and IR spectra show that the actual structure of this ion is 6, in which both the A and the B component of the hydrogen bond is a carbon 61 These are sometimes
49For reviews of hydrogen bonding in sulfur- containing compounds, see Zuika, I.V.; Bankovskii, Yu.A. Russ. Chem. Rev. 1973 , 42, 22; Crampton, M.R., in Patai, S. The Chemistry of the Thiol Group, pt. 1; Wiley, NY, 1974, pp. 379-396; Pogorelyi, V.K. Russ. Chem. Rev. 1977, 46, 316.
50For a review of hydrogen bonding to halogens, see Smith, J.W., in Patai, S. The Chemistry of the Carbon-Halogen Bond, pt. 1; Wiley, NY, 1973, pp. 265-300. See also, Bastiansen, O.; Fernholt, L.; Hedberg, K.; Seip, R. J. Am. Chem. Soc. 1985, 107, 7836.
5 ‘West, R.; Powell, D.L.; Whatley, L.S.; Lee, M.K.T.; Schleyer, P.v.R. J. Am. Chem. Soc. 1962, 84, 3221; Fujimoto, E.; Takeoka, Y.; Kozima, K. Bull. Chem. Soc. Jpn. 1970, 43, 991; Azrak, R.G.; Wilson, E.B. J. Chem. Phys. 1970, 52, 5299.
52Allerhand, A.; Schleyer, P.v.R. J. Am. Chem. Soc. 1963, 85, 1233; McDaniel, D.H.; Vallee, R.E. Inorg. Chem. 1963, 2, 996; Fujiwara, F.Y.; Martin, J.S. J. Am. Chem. Soc. 1974, 96, 7625; French, M.A.; Ikuta, S.; Kebarle, P. Can. J. Chem. 1982, 60, 1907.
53 A few exceptions have been found, where the presence of an unsymmetrical cation causes the hydrogen to be closer to one fluorine than to the other: Williams, J.M.; Schneemeyer, L.F. J. Am. Chem. Soc. 1973, 95, 5780.
54Vogel, G.C.; Drago, R.S. J. Am. Chem. Soc. 1970, 92, 5347; Mukherjee, S.; Palit, S.R.; De, S.K. J. Phys. Chem. 1970, 74, 1389; Schaefer, T.; McKinnon, D.M.; Sebastian, R.; Peeling, J.; Penner, G.H.; Veregin, R.P. Can. J. Chem. 1987, 65, 908; Marstokk, K.; Mpllendal, H.; Uggerrud, E. Acta Chem. Scand. 1989, 43, 26.
55McDaniel, D.H.; Evans, W.G. Inorg. Chem. 1966, 5, 2180; Sabin, J.R. J. Chem. Phys. 1971, 54, 4675. 56Calhorda, M.J. Chem. Commun. 2000, 801.
57Ahlberg, P; Davidsson, O.; Johnsson, B.; McEwen, I.; Ronnqvist, M. Bull. Soc. Chim. Fr. 1988, 111 . 58Ferstandig, L.L. J. Am. Chem. Soc. 1962, 84, 3553; Allerhand, A.; Schleyer, P.v.R. J. Am. Chem. Soc. 1963, 85, 866.
59For example, see Bakke, J.M.; Chadwick, D.J. Acta Chem. Scand. Ser. B 1988, 42, 223: Atwood, J.L.; Hamada, F.; Robinson, K.D.; Orr, G.W.; Vincent, R.L. Nature ( London ) 1991, 349, 683.
60Joris, L.; Schleyer, P.v.R.; Gleiter, R. J. Am. Chem. Soc. 1968, 90, 327; Yoshida, Z.; Ishibe, N.; Kusumoto, H. J. Am. Chem. Soc. 1969, 91, 2279.
61McMurry, J.E.; Lectka, T.; Hodge, C.N. J. Am. Chem. Soc. 1989, 111, 8867. See also, Sorensen, T.S.; Whitworth, S.M. J. Am. Chem. Soc. 1990, 112, 8135.
CHAPTER 3
HYDROGEN BONDING 113
called 3-center-2-electron C-H-C bonds.62 A technique called generalized popula¬ tion analysis has been developed to study this type of multicenter bonding.63
A weak (~1.5 kcalmol-1) and rare C-H*"0=C hydrogen bond has been reported in a class of compounds known as a [6]semirubin (a dipyrrinone).64 There is also evidence for a C-H#mN/CH##*OH bond in the crystal structures of a,P-unsaturated ketones carrying a terminal pyridine subunit,65 and for R3N+-C- H*"0=C hydrogen bonding.66
Deuterium also forms hydrogen bonds; in some systems these seem to be stron¬ ger than the corresponding hydrogen bonds; in others, weaker.67
OH
Pc.;
Ts
Weak hydrogen bonds can be formed between an appropriate hydrogen and a n bond, both with alkenes and with aromatic compounds. For example, IR data in dilute dichloromethane suggests that the predominant conformation for bis (amide) 7 contains an hydrogen bond involving the C=C unit.68 The
strength of an intramolecular 71-facial hydrogen bond between an NH group and an aromatic ring in chloroform has been estimated to have a lower limit of —4.5 zb 0.5 kcal mol-1(— 18.8 kJ mol-1).69 A neutron diffraction study of crystal¬ line 2-ethynyladamantan-2-ol (8) shows the presence of an unusual 0-H*"7i
62McMurry, J.E.; Lectka, T. Acc. Chem. Res. 1992, 25, 47.
63Ponec, R.; Yuzhakov, G.; Tantillo, DJ. J. Org. Chem. 2004, 69, 2992.
64Huggins, M.T.; Lightner, D.A. J. Org. Chem. 2001, 66, 8402.
65Mazik, M.; B laser, D.; Boese, R. Tetrahedron 2001, 57, 5791.
66Cannizzaro, C.E.; Houk, K.N. J. Am. Chem. Soc. 2002, 124, 7163.
67Dahlgren Jr., G.; Long, F.A. J. Am. Chem. Soc. 1960, 82, 1303; Creswell, C.J.; Allred, A.L. J. Am. Chem. Soc. 1962, 84, 3966; Singh, S.; Rao, C.N.R. Can. J. Chem. 1966, 44, 2611; Cummings, D.L.; Wood, J.L. J. Mol. Struct. 1974, 23, 103.
68Gallo, E.A.; Gelman, S.H. Tetrahedron Lett. 1992, 33, 7485.
69 Adams, H.; Harris, K.D.M.; Hembury, G.A.; Hunter, C.A.; Livingstone, D.; McCabe, J.F. Chem. Commun. 1996, 2531. See Steiner, T.; Starikov, E.B.; Tamm, M. J. Chem. Soc., Perkin Trans. 2 1996, 67 for a related example with 5-ethynyl-5//-dibenzo[a,<i]cyclohepten-5-ol.
114 BONDING WEAKER THAN COVALENT
hydrogen bond, which is short and linear, as well as the more common 0-H**»0 and C-H*"0 hydrogen bonds.70
Ti-Ti INTERACTIONS
The 7i-7i interactions are fundamental to many supramolecular organization and recognition processes.71 There are many theoretical and experimental studies that clearly show the importance of n-n interactions.72 Perhaps the simplest prototype of aromatic n-n interactions is the benzene dimer.73 Within dimeric aryl systems such as this, possible n-n interactions are the sandwich and T-shaped interactions shown. It has been shown that all substituted sandwich dimers bind more strongly than benzene dimer, whereas the T-shaped configurations bind more or less favorably depending on the substituent.74 Electrostatic, dispersion, induction, and exchange- repulsion contributions are all significant to the overall binding energies.74
x
T-shaped (1)
T-shaped (2)
The 7i-electrons of aromatic rings can interact with charged species, yielding strong cation-71 interactions dominated by electrostatic and polarization effects.75 Interactions with CH units is also possible. For CH-ti interactions in both alkyl- and aryl-based model systems, dispersion effects dominate the interaction, but the electrostatics term is also relevant for aryl CH-ti interactions.76
70Allen, F.H.; Howard, J.A.K.; Hoy, V.J.; Desiraju, G.R.; Reddy, D.S.; Wilson, C.C. J. Am. Chem. Soc. 1996, 118, 4081.
71Meyer, E.A.; Castellano, R.K.; Diederich, F. Angew. Chem. Int. Ed. 2003, 42, 1210.
72T suzuki, T.; Uchimaru, T.; Tanabe, K. J. Mol. Struct. ( THEOCHEM) 1994, 307, 107; Hobza, P.; Selzle, H.L.; Schlag, E.W. J. Phys. Chem. 1996, 100, 18790; Tsuzuki, S.; Liithi, H.P. J. Chem. Phys. 2001, 114, 3949; Steed, J.M.; Dixon, T.A.; Klemperer, W. J. Chem. Phys. 1979, 70, 4940.; Arunan, E.; Gutowsky, H.S. J. Chem. Phys. 1993, 98, 4294; Law, K.S.; Schauer, M.; Bernstein, E.R. J. Chem. Phys. 1984, 81, 4871; Felker, P.M.; Maxton, P.M.; Schaeffer, M.W. Chem. Rev. 1994, 94, 1787; Venture, V.A.; Felker, PM. J. Chem. Phys. 1993, 99, 748; Tsuzuki, S.; Honda, K.; Uchimaru, T.; Mikami, M.; Tanabe, K. J. Am. Chem. Soc. 2002, 124, 104; Hobza, P; Jurecka, P. J. Am. Chem. Soc. 2003, 125, 15608.
73Sinnokrot, M.O.; Valeev, E.F.; Sherrill, C.D. J. Am. Chem. Soc. 2002, 124, 10887.
74Sinnokrot, M.O.; Sherrill, C.D. J. Am. Chem. Soc. 2004, 126, 7690
75Lindeman, S.V.; Kosynkin, D.; Kochi, J.K. J. Am. Chem. Soc. 1998, 120, 13268; Ma, J.C.; Dougherty, D.A. Chem. Rev. 1997, 97, 1303; Dougherty, D.A. Science 1996, 271, 163; Cubero, E.; Luque, F.J.; Orozco, M. Proc. Natl. Acad. Sci. U.S.A. 1998, 95, 5976.
76Ribas, J.; Cubero, E.; Luque, F. J.; Orozco, M. J. Org. Chem. 2002, 67, 7057.
CHAPTER 3
ADDITION COMPOUNDS 115
Detection of n-n interactions has largely relied on NMR-based techniques, such as chemical shifts variations,77 and Nuclear Overhauser Effect Spectroscopy (NOESY) or Rotating-Frame NOE Spectroscopy (ROESY).78 Diffusion-ordered NMR spectroscopy (DOSY) has also been used to detect n-n stacked complexes.79
ADDITION COMPOUNDS
When the reaction of two compounds results in a product that contains all the mass of the two compounds, the product is called an addition compound. There are sev¬ eral kinds. In the rest of this chapter, we will discuss addition compounds in which the molecules of the starting materials remain more or less intact and weak bonds hold two or more molecules together. We can divide them into four broad classes: electron donor-acceptor complexes, complexes formed by crown ethers and similar compounds, inclusion compounds, and catenanes.
Electron Donor-Acceptor (EDA) Complexes80
In EDA complexes ,81 there is always a donor and an acceptor molecule. The donor may donate an unshared pair (an n donor) or a pair of electrons in a n orbital of a double bond or aromatic system (a n donor). One test for the presence of an EDA complex is the electronic spectrum. These complexes generally exhibit a spectrum (called a charge-transfer spectrum) that is not the same as the sum of the spectra of the two individual molecules.82 Because the first excited state of the complex is relatively close in energy to the ground state, there is usually a peak in the visible or near-uv region and EDA complexes are often colored. Many EDA complexes are unstable and exist only in solutions in equilibrium with their components, but others are stable solids. In most EDA complexes the donor and acceptor molecules are present in an integral ratio, most often 1:1, but complexes with nonintegral ratios are also known. There are several types of acceptor molecules; we will discuss com¬ plexes formed by two of them.
77Petersen, S.B.; Led, J.J.; Johnston, E.R.; Grant, D.M. J. Am. Chem. Soc. 1982, 104, 5007.
78Wakita, M.; Kuroda, Y.; Fujiwara, Y.; Nakagawa, T. Chem. Phys. Lipids 1992, 62, 45.
79Viel, S.; Mannina, L.; Segre, A. Tetrahedron Lett. 2002, 43, 2515. See also, Ribas, J.; Cubero, E.; Luque, F.J.; Orozco, M. J. Org. Chem. 2002, 67, 7057.
80For monographs, see Foster, R. Organic Charge-Transfer Complexes, Academic Press, NY, 1969, Mulliken, R.S.; Person, W.B. Molecular Complexes, Wiley, NY, 1969\ Rose, J. Molecular Complexes, Pergamon, Elmsford, NY, 1967. For reviews, see Poleshchuk, O.Kh.; Maksyutin, Yu.K. Russ. Chem. Rev. 1976, 45, 1077; Banthorpe, D.V. Chem. Rev. 1970, 70, 295; Kosower, E.M. Prog. Phys. Org. Chem. 1965, 3, 81; Foster, R. Chem. Br. 1976, 12, 18.
81These have often been called charge-transfer complexes, but this term implies that the bonding involves charge transfer, which is not always the case, so that the more neutral name EDA complex is preferable. See Mulliken, R.S.; Person, W.B. J. Am. Chem. Soc. 1969, 91, 3409.
82For examples of EDA complexes that do not show charge-transfer spectra, see Bentley, M.D.; Dewar, M.J.S. Tetrahedron Lett. 1967, 5043.
116 BONDING WEAKER THAN COVALENT
1. Complexes in Which the Acceptor Is A Metal Ion and the Donor an Alkene or an Aromatic Ring (n donors do not give EDA complexes with metal ions but form covalent bonds instead).83 Many metal ions form complexes, that are often stable solids, with alkenes, dienes (usually conjugated, but not always), alkynes, and aromatic rings. The donor (or ligand) molecules in these complexes are classified by the prefix hapto 84 and/or the descriptor r\n (the Greek letter eta), where n indicates how many atoms the ligand uses to bond with the metal.85 The generally accepted picture of the bonding in these complexes,86 first proposed by Dewar,87 can be
illustrated by the ethylene complex with silver, 9, in which the alkene unit forms an r|2-complex with the silver ion (the alkene functions as a 2-electron donating ligand to the metal). There is evidence of 7i-complexation of Na+ by C=C.88
83For monographs, see Collman, J.P.; Hegedus, L.S.; Norton, J.R.; Finke, R.G. Principles and Applications of Organotransition Metal Chemistry, 2nd ed, University Science Books, Mill Valley, CA, 1987; Alper, H. Transition Metal Organometallics in Organic Synthesis, 2 vols., Academic Press, NY, 1976, 1978; King, R.B. Transition-Metal Organic Chemistry, Academic Press, NY, 1969; Green, M.L.H. Organometallic Compounds, Vol. 2, Methuen, London, 1968; For general reviews, see Churchill, M.R.; Mason, R. Adv. Organomet. Chem. 1967, 5, 93; Cais, M., in Patai, S. The Chemistry of Alkenes, Vol. 1, Wiley, NY, 1964, pp. 335-385. Among the many reviews limited to certain classes of complexes are transition metals-dienes, Nakamura, A. J. Organomet. Chem. 1990, 400, 35; metals-cycloalkynes and arynes, Bennett, M.A.; Schwemlein, H.P. Angew. Chem. Int. Ed. 1989, 28, 1296; metals-pentadienyl ions, Powell, P. Adv. Organomet. Chem. 1986, 26, 125; complexes of main-group metals, Jutzi, P. Adv. Organomet. Chem. 1986, 26, 217; intramolecular complexes, Omae, I. Angew. Chem. Int. Ed. 1982, 21, 889; transition metals-olefins and acetylenes, Pettit, L.D.; Barnes, D.S. Fortschr. Chem. Forsch. 1972, 28, 85; Quinn, H.W.; Tsai, J.H. Adv. Inorg. Chem. Radiochem. 1969, 12, 217; Pt- and Pd-olefins and acetylenes, Hartley, F.R. Chem. Rev. 1969, 69, 799; silver ions-olefins and aromatics, Beverwijk, C.D.M.; van der Kerk, G.J.M.; Leusink, J.; Noltes, J.G. Organomet. Chem. Rev. Sect. A 1970, 5, 215; metals- substituted olefins, Jones, R. Chem. Rev. 1968, 68, 785; transition metals-allylic compounds, Clarke, H.L. J. Organomet. Chem. 1974, 80, 155; transition metals-arenes, Silverthorn, W.E. Adv. Organomet. Chem. 1976, 14, 47; metals-organosilicon compounds, Haiduc, I.; Popa, V. Adv. Organomet. Chem. 1977, 15, 113; metals-carbocations, Pettit, L.D.; Haynes, L.W., in Olah, G.A.; Schleyer, P.v.R. Carbonium Ions, Vol. 5, Wiley, NY, 1976, pp. 2263-2302; metals-seven-and eight-membered rings, Bennett, M.A. Adv. Organomet. Chem. 1966, 4, 353. For a list of review articles on this subject, see Bruce, M.I. Adv. Organomet. Chem. 1972, 10, 273, pp. 317-321.
84For a discussion of how this system originated, see Cotton, F.A. J. Organomet. Chem. 1975, 100, 29. 85 Another prefix used for complexes is p (mu), which indicates that the ligand bridges two metal atoms. 86For reviews, see Pearson, A.J. Metallo- organic Chemistry, Wiley, NY, 1985; Ittel, S.D.; Ibers, J.A. Adv. Organomet. Chem. 1976, 14, 33; Hartley, F.R. Chem. Rev. 1973, 73, 163; Angew. Chem. Int. Ed. 1972, 11, 596.
87Dewar, M.J.S. Bull. Soc. Chim. Fr. 1951, 18, C79.
88Hu, J.; Gokel, G.W.; Barbour, LJ. Chem. Commun. 2001, 1858.
CHAPTER 3
ADDITION COMPOUNDS 117
In the case of the silver complex, the bond is not from one atom of the C=C unit to the silver ion, but from the n center such that two electrons are transferred from the alkene to the metal ion.89 Ethene has two Ti-electrons and is a dihapto or p2 ligand, as are other simple alkenes. Similarly, benzene has six 71-electrons and is a hexahapto or r|6 ligand. Ferrocene (10) has two cyclopentadienyl ligands (each is a five-electron donor or an r|5 ligand), and ferrocene is properly called bis(r|5-cyclopentadienyl)iron(II). This system can be extended to compounds in which only a single a bond connects the organic group to the metal, for example, C6H5-Li (a mono- hapto or p1 ligand), and to complexes in which the organic group is an ion, for example, Ti-allyl complexes, such as 11, in which the allyl ligand is trihapto or p3. Note that in a compound such as allyllithium, where a a bond connects the carbon to the metal, the allyl group is referred to as monohapto or p1.
CH2 — CH — CH2 — Li Allyllithium
As mentioned, benzene is an p6 ligand that forms complexes with silver and other metals.90 When the metal involved has a coordination number >1, more than one donor molecule (ligand) participates. The CO group is a common ligand (a two-electron donating or p2 ligand), and in metal complexes the CO group is classified as a metal carbonyl. Benzenechro- mium tricarbonyl (12) is a stable compound91 that illustrates both benzene and carbonyl ligands. Three arrows are shown to represent the six-electron donation (an p6 ligand), but the accompanying model gives a clearer picture of the bonding. Cyclooctatetraene is an eight-electron donating or p8 ligand that also forms complexes with metals. Metallocenes (see 10) may be considered a special case of this type of complex, although the bonding in
89For a discussion of how the nature of the metal ion affects the stability of the complex, see p. $$$. 90For a monograph, see Zeiss, H.; Wheatley, P.J.; Winkler, H.J.S. BenzenoidMetal Complexes ; Ronald Press, NY, 1966.
91Nicholls, B.; Whiting, M.C. J. Chem. Soc. 1959, 551. For reviews of arene-transition-metal complexes, see Uemura, M.Adv. Met.-Org. Chem. 1991, 2, 195; Silverthorn, W.E. Adv. Organomet. Chem. 1975, 13, 47.
118 BONDING WEAKER THAN COVALENT
metallocenes is much stronger.
13
In a number of cases, alkenes that are too unstable for isolation have been isolated in the form of metal complexes. As example is norbornadienone, which was isolated in the form of its iron-tricarbonyl complex (13), 92 where the norbornadiene unit is an r|4 ligand, and each of the carbonyl units are r|2 ligands. The free dienone spontaneously decomposes to carbon monoxide and benzene (see reaction 17-28).
2. Complexes in Which the Acceptor Is an Organic Molecule. Picric acid, 1,3, 5 -trinitrobenzene, and similar polynitro compounds are the most important of these.93 Picric acid forms addition compounds with many
OH
aromatic hydrocarbons, aromatic amines, aliphatic amines, alkenes, and other compounds. These addition compounds are usually solids with definite melt¬ ing points and are often used as derivatives of the compounds in question. They are called picrates, though they are not salts of picric acid, but addition compounds. Unfortunately, salts of picric acid are also called picrates. Simi¬ lar complexes are formed between phenols and quinones (quinhydrones).94
O
Fe(CO)3
92Landesberg, J.M.; Sieczkowski, J. J. Am. Chem. Soc. 1971, 93, 912.
93For a review, see Parini, V.P. Russ. Chem. Rev. 1962, 31, 408; for a review of complexes in which the acceptor is an organic cation, see Kampar, V.E. Russ. Chem. Rev. 1982, 51, 107; also see Ref. 80.
94For a review of quinone complexes, see Foster, R.; Foreman, M.I., in Patai, S. The Chemistry of the Quinonoid Compounds, pt. 1, Wiley, NY, 1974, pp. 257-333.
CHAPTER 3
ADDITION COMPOUNDS 119
Alkenes that contain electron-withdrawing substituents also act as acceptor molecules, as do carbon tetrahalides95 and certain anhydrides.96 A particu¬ larly strong alkene acceptor is tetracyanoethylene.97
The bonding in these cases is more difficult to explain than in the previous case, and indeed no really satisfactory explanation is available.98 The difficulty is that although the donor has a pair of electrons to contribute (both n and n donors are found here), the acceptor does not have a vacant orbital. Simple attraction of the dipole-induced dipole type accounts for some of the bonding,99 but is too weak to explain the bonding in all cases;100 for example, nitromethane, with about the same dipole moment as nitrobenzene, forms much weaker complexes. Some other type of bonding clearly must also be present in many EDA complexes. The exact nature of this bonding, called charge -tra