Molecular Orbitals and Organic Chemical Reactions, Reference Edition E-Book

Ian Fleming

• Herausgeber: John Wiley & Sons
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch
• Veröffentlichungsjahr: 2011

Seitenzahl: 1127

Beschreibung

Winner of the PROSE Award for Chemistry & Physics 2010 Acknowledging the very best in professional and scholarly publishing, the annual PROSE Awards recognise publishers' and authors' commitment to pioneering works of research and for contributing to the conception, production, and design of landmark works in their fields. Judged by peer publishers, librarians, and medical professionals, Wiley are pleased to congratulate Professor Ian Fleming, winner of the PROSE Award in Chemistry and Physics for Molecular Orbitals and Organic Chemical Reactions. Molecular orbital theory is used by chemists to describe the arrangement of electrons in chemical structures. It is also a theory capable of giving some insight into the forces involved in the making and breaking of chemical bonds--the chemical reactions that are often the focus of an organic chemist's interest. Organic chemists with a serious interest in understanding and explaining their work usually express their ideas in molecular orbital terms, so much so that it is now an essential component of every organic chemist's skills to have some acquaintance with molecular orbital theory. Molecular Orbitals and Organic Chemical Reactions is both a simplified account of molecular orbital theory and a review of its applications in organic chemistry; it provides a basic introduction to the subject and a wealth of illustrative examples. In this book molecular orbital theory is presented in a much simplified, and entirely non-mathematical language, accessible to every organic chemist, whether student or research worker, whether mathematically competent or not. Topics covered include: * Molecular Orbital Theory * Molecular Orbitals and the Structures of Organic Molecules * Chemical Reactions -- How Far and How Fast * Ionic Reactions -- Reactivity * Ionic Reactions -- Stereochemistry * Pericyclic Reactions * Radical Reactions * Photochemical Reactions This expanded Reference Edition of Molecular Orbitals and Organic Chemical Reactions takes the content and the same non-mathematical approach of the Student Edition, and adds extensive extra subject coverage, detail and over 1500 references. The additional material adds a deeper understanding of the models used, and includes a broader range of applications and case studies. Providing a complete in-depth reference for a more advanced audience, this edition will find a place on the bookshelves of researchers and advanced students of organic, physical organic and computational chemistry. The student edition of Molecular Orbitals and Organic Chemical Reactions presents molecular orbital theory in a simplified form, and offers an invaluable first textbook on this important subject for students of organic, physical organic and computational chemistry. Further information can be viewed here. "These books are the result of years of work, which began as an attempt to write a second edition of my 1976 book Frontier Orbitals and Organic Chemical Reactions. I wanted to give a rather more thorough introduction to molecular orbitals, while maintaining my focus on the organic chemist who did not want a mathematical account, but still wanted to understand organic chemistry at a physical level. I'm delighted to win this prize, and hope a new generation of chemists will benefit from these books." --Professor Ian Fleming

Leseprobe

Contents

Preface

1 Molecular Orbital Theory

1.1 The Atomic Orbitals of a Hydrogen Atom

1.2 Molecules Made from Hydrogen Atoms

1.3 C—H and C—C Bonds

1.4 Conjugation—Hückel Theory

1.5 Aromaticity22

1.6 Strained s Bonds—Cyclopropanes and Cyclobutanes

1.7 Heteronuclear Bonds, C—M, C—X and C=O

1.8 The Tau Bond Model

1.9 Spectroscopic Methods

2 Molecular Orbitals and the Structures of Organic Molecules

2.1 The Effects of π Conjugation

2.2 Hyperconjugation—σ Conjugation

2.3 The Configurations and Conformations of Molecules

2.4 The Effect of Conjugation on Electron Distribution

2.5 Other Noncovalent Interactions185

3 Chemical Reactions—How Far and How Fast

3.1 Factors Affecting the Position of an Equilibrium

3.2 The Principle of Hard and Soft Acids and Bases (HSAB)

3.3 Transition Structures

3.4 The Perturbation Theory of Reactivity

3.5 The Salem-Klopman Equation

3.6 Hard and Soft Nucleophiles and Electrophiles

3.7 Other Factors Affecting Chemical Reactivity

4 Ionic Reactions—Reactivity

4.1 Single Electron Transfer (SET) in Ionic Reactions

4.2 Nucleophilicity

4.3 Ambident Nucleophiles

4.4 Electrophilicity

4.5 Ambident Electrophiles

4.6 Carbenes428

5 Ionic Reactions—Stereochemistry

5.1 The Stereochemistry of the Fundamental Organic Reactions

5.2 Diastereoselectivity509

6 Thermal Pericyclic Reactions

6.1 The Four Classes of Pericyclic Reactions

6.2 Evidence for the Concertedness of Bond Making and Breaking

6.3 Symmetry-allowed and Symmetry-forbidden Reactions

6.4 Explanations for the Woodward-Hoffmann Rules

6.5 Secondary Effects

7 Radical Reactions

7.1 Nucleophilic and Electrophilic Radicals

7.2 The Abstraction of Hydrogen and Halogen Atoms

7.3 The Addition of Radicals to p Bonds

7.4 Synthetic Applications of the Chemoselectivity of Radicals

7.5 Stereochemistry in some Radical Reactions

7.6 Ambident Radicals

7.7 Radical Coupling

8 Photochemical Reactions

8.1 Photochemical Reactions in General

8.2 Photochemical Ionic Reactions

8.3 Photochemical Pericyclic Reactions and Related Stepwise Reactions

8.4 Photochemically Induced Radical Reactions

8.5 Chemiluminescence

Reference

Index

This edition first published 2010

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Library of Congress Cataloging-in-Publication Data

Fleming, Ian, 1935–

Molecular orbitals and organic chemical reactions/Ian Fleming. — Reference ed. p. cm.

Includes bibliographical references and index.

ISBN 978-0-470-74658-5

1. Molecular orbitals. 2. Chemical bonds. 3. Physical organic chemistry. I. Title.

QD461.F533 2010

547'.2—dc22

2009041770

Preface

Molecular orbital theory is used by chemists to describe the arrangement of electrons in chemical structures. It provides a basis for explaining the ground-state shapes of molecules and their many other properties. As a theory of bonding it has largely replaced valence bond theory,1 but organic chemists still implicitly use valence bond theory whenever they draw resonance structures. Unfortunately, misuse of valence bond theory is not uncommon as this approach remains in the hands largely of the less sophisticated. Organic chemists with a serious interest in understanding and explaining their work usually express their ideas in molecular orbital terms, so much so that it is now an essential component of every organic chemist's skills to have some acquaintance with molecular orbital theory. The problem is to find a level to suit everyone. At one extreme, a few organic chemists with high levels of mathematical skill are happy to use molecular orbital theory, and its computationally more amenable offshoot density functional theory, much as theoreticians do. At the other extreme are the many organic chemists with lower mathematical inclinations, who nevertheless want to understand their reactions at some kind of physical level. It is for these people that I have written this book. In between there are more and more experimental organic chemists carrying out calculations to support their observations, and these people need to know some of the physical basis for what their calculations are doing.2

I have presented molecular orbital theory in a much simplified and entirely nonmathematical language. I have simplified the treatment in order to make it accessible to every organic chemist, whether student or research worker, whether mathematically competent or not. In order to reach such a wide audience, I have frequently used oversimplified arguments. I trust that every student who has the aptitude will look beyond this book for a better understanding than can be found here. Accordingly, I have provided over 1800 references to the theoretical treatments and experimental evidence, to make it possible for every reader to go further into the subject.

Molecular orbital theory is not only a theory of bonding, it is also a theory capable of giving some insight into the forces involved in the making and breaking of chemical bonds—the chemical reactions that are often the focus of an organic chemist's interest. Calculations on transition structures can be carried out with a bewildering array of techniques requiring more or less skill, more or fewer assumptions, and greater or smaller contributions from empirical input, but many of these fail to provide the organic chemist with insight. He or she wants to know what the physical forces are that give the various kinds of selectivity that are so precious in learning how to control organic reactions. The most accessible theory to give this kind of insight is frontier orbital theory, which is based on the perturbation treatment of molecular orbital theory, introduced by Coulson and Longuet-Higgins,3 and developed and named as frontier orbital theory by Fukui.4 Earlier theories of reactivity concentrated on the product-like character of transition structures—the concept of localisation energy in aromatic electrophilic substitution is a well-known example. The perturbation theory concentrates instead on the other side of the reaction coordinate. It looks at how the interaction of the molecular orbitals of the starting materials influences the transition structure. Both influences are obviously important, and it is therefore helpful to know about both if we want a better understanding of what factors affect a transition structure, and hence affect chemical reactivity.

Frontier orbital theory is now widely used, with more or less appropriateness, especially by organic chemists, not least because of the success of the predecessor to this book, Frontier Orbitals and Organic Chemical Reactions, which survived for more than thirty years as an introduction to the subject for a high proportion of the organic chemists trained in this period. However, there is a problem—computations show that the frontier orbitals do not make a significantly larger contribution than the sum of all the orbitals. One theoretician put it to me as: 'It has no right to work as well as it does.' The difficulty is that it works as an explanation in many situations where nothing else is immediately compelling. In writing this book, I have therefore emphasised more the molecular orbital basis for understanding organic chemistry, about which there is less disquiet. Thus I have completely rewritten the earlier book, enlarging especially the chapters on molecular orbital theory itself. I have added a chapter on the effect of orbital interactions on the structures of organic molecules, a section on the theoretical basis for the principle of hard and soft acids and bases, and a chapter on the stereochemistry of the fundamental organic reactions. I have introduced correlation diagrams into the discussion of pericyclic chemistry, and a great deal more in that, the largest chapter. I have also added a number of topics, both omissions from the earlier book and new work that has taken place in the intervening years. I have used more words of caution in discussing frontier orbital theory itself, making it less polemical in furthering that subject, and hoping that it might lead people to be more cautious themselves before applying the ideas uncritically in their own work.

For all their faults and limitations, frontier orbital theory and the principle of hard and soft acids and bases remain the most accessible approaches to understanding many aspects of reactivity. Since they fill a gap between the chemist's experimental results and a state of the art theoretical description of his or her observations, they will continue to be used, until something better comes along.

In this book, there is much detailed and not always convincing material, making it less suitable as a textbook for a lecture course; in consequence I have also written a second and shorter book on molecular orbital theory designed specifically for students of organic chemistry, Molecular Orbitals and Organic Chemistry—The Student Edition,5 which serves in a sense as a long awaited second edition to my earlier book. The shorter book uses a selection of the same material as in this volume, with appropriately revised text, but dispenses with most of the references, which can all be found here. The shorter book also has problem sets at the ends of the chapters, whereas this book has the answers to most of them in appropriate places in the text. I hope that everyone can use whichever volume suits them, and that even theoreticians might find unresolved problems in one or another of them.

As in the earlier book, I begin by presenting some experimental observations that chemists have wanted to explain. None of the questions raised by these observations has a simple answer without reference to the orbitals involved.

(i) Why does methyl tetrahydropyranyl ether largely adopt the conformation P.1, with the methoxy group axial, whereas methoxycyclohexane adopts largely the conformation P.2 with the methoxy group equatorial?

(ii) Reduction of butadiene P.3 with sodium in liquid ammonia gives more cis-2-butene P.4 than trans-2-butene P.5, even though the trans isomer is the more stable product.

(iii) Why is the inversion of configuration at nitrogen made slower if the nitrogen is in a small ring, and slower still if it has an electronegative substituent attached to it, so that, with the benefit of both features, an N-chloroaziridine can be separated into a pair of diastereoisomers P.6 and P.7?

(iv) Why do enolate ions P.8 react more rapidly with protons on oxygen, but with primary alkyl halides on carbon?

(v) Hydroperoxide ion P.9 is much less basic than hydroxide ion P.10. Why, then, is it so much more nucleophilic?

(vi) Why does butadiene P.11 react with maleic anhydride P.12, but ethylene P.13 does not?

(vii) Why do Diels-Alder reactions of butadiene P.11 go so much faster when there is an electron-withdrawing group on the dienophile, as with maleic anhydride P.12, than they do with ethylene P.13?

(viii) Why does diazomethane P.15 add to methyl acrylate P.16 to give the isomer P.17 in which the nitrogen end of the dipole is bonded to the carbon atom bearing the methoxycarbonyl group, and not the other way round P.14?

(ix) When methyl fumarate P.18 and vinyl acetate P.19 are copolymerised with a radical initiator, why does the polymer P.20 consist largely of alternating units?

(x) Why does the Paterno-Biichi reaction between acetone and acrylonitrile give only the isomer P.21 in which the two 'electrophilic' carbon atoms become bonded?

In the following chapters, each of these questions, and many others, receives a simple answer. Other books commend themselves to anyone able and willing to go further up the mathematical slopes towards a more acceptable level of explanation—a few introductory texts take the next step up,6,7 and several others8–11 take the story further.

I have been greatly helped by a number of chemists: first and foremost Professor Christopher Longuet-Higgins, whose inspiring lectures persuaded me to take the subject seriously at a time when most organic chemists who, like me, had little mathematics, had abandoned any hope of making sense of the subject; secondly, and more particularly those who gave me advice for the earlier book, and who therefore made their mark on this, namely Dr W. Carruthers, Professor R. F. Hudson, Professor A. R. Katritzky and Professor A. J. Stone. In addition, for this book, I am indebted to Dr Jonathan Goodman for help with computer programs, to Professor Wes Borden for some helpful discussions and collaboration on one topic, and to Professor A. D. Buckingham for several important corrections. More than usually, I must absolve all of them for any errors left in the book.

1

Molecular Orbital Theory

1.1 The Atomic Orbitals of a Hydrogen Atom

To understand the nature of the simplest chemical bond, that between two hydrogen atoms, we look at the effect on the electron distribution when two atoms are held within bonding distance, but first we need a picture of the hydrogen atoms themselves. Since a hydrogen atom consists of a proton and a single electron, we only need a description of the spatial distribution of that electron. This is usually expressed as a wave functionϕ, where ϕ2dτ is the probability of finding the electron in the volume dτ, and the integral of ϕ2dτ over the whole of space is 1. The wave function is the underlying mathematical description, and it may be positive or negative; it can even be complex with a real and an imaginary part, but this will not be needed in any of the discussion in this book. Only when squared does it correspond to anything with physical reality—the probability of finding an electron in any given space. Quantum theory12 gives us a number of permitted wave equations, but the only one that matters here is the lowest in energy, in which the distribution of the electron is described as being in a 1s orbital. This is spherically symmetrical about the nucleus, with a maximum at the centre, and falling off rapidly, so that the probability of finding the electron within a sphere of radius 1.4 Å is 90% and within 2 Å better than 99%. This orbital is calculated to be 13.60 eV lower in energy than a completely separated electron and proton.

We need pictures to illustrate the electron distribution, and the most common is simply to draw a circle, Fig. 1.1a, which can be thought of as a section through a spherical contour, within which the electron would be found, say, 90% of the time. This picture will suffice for most of what we need in this book, but it might be worth looking at some others, because the circle alone disguises some features that are worth appreciating. Thus a section showing more contours, Fig. 1.1b, has more detail. Another picture, even less amenable to a quick drawing, is to plot the electron distribution as a section through a cloud, Fig. 1.1c, where one imagines blinking one’s eyes a very large number of times, and plotting the points at which the electron was at each blink. This picture contributes to the language often used, in which the electron population in a given volume of space is referred to as the electron density.

Fig. 1.1 The 1s atomic orbital of a hydrogen atom

Fig. 1.2 Radial probability plots for the 1s orbital of a hydrogen atom

Taking advantage of the spherical symmetry, we can also plot the fraction of the electron population outside a radius r against r, as in Fig. 1.2a, showing the rapid fall off of electron population with distance. The van der Waals radius at 1.2 Å has no theoretical significance—it is an empirical measurement from solidstate structures, being one-half of the distance apart of the hydrogen atom in a C—H bond and the hydrogen atom in the C—H bond of an adjacent molecule.13 It does not even have a fixed value, but is an average of several measurements. Yet another way to appreciate the electron distribution is to look at the radial density, where we plot the probability of finding the electron between one sphere of radius r and another of radius r + dr. This has a revealing form, Fig. 1.2b, with a maximum 0.529 Å from the nucleus, showing that, in spite of the wave function being at a maximum at the nucleus, the chance of finding an electron precisely there is very small. The distance 0.529 Å proves to be the same as the radius calculated for the orbit of an electron in the early but untenable planetary model of a hydrogen atom. It is called the Bohr radius a0, and is often used as a unit of length in molecular orbital calculations.

1.2 Molecules Made from Hydrogen Atoms

1.2.1 The H2 Molecule

To understand the bonding in a hydrogen molecule, we have to see what happens when two hydrogen atoms are close enough for their atomic orbitals to interact. We now have two protons and two nuclei, and even with this small a molecule we cannot expect theory to give us complete solutions. We need a description of the electron distribution over the whole molecule—a molecular orbital. The way the problem is handled is to accept that a first approximation has the two atoms remaining more or less unchanged, so that the description of the molecule will resemble the sum of the two isolated atoms. Thus we combine the two atomic orbitals in a linear combination expressed in Equation 1.1, where the function which describes the new electron distribution, the molecular orbital, is called σ and ϕ1 and ϕ2 are the atomic 1s wave functions on atoms 1 and 2.

1.1

The coefficients, c1 and c2, are a measure of the contribution which the atomic orbital is making to the molecular orbital. They are of course equal in magnitude in this case, since the two atoms are the same, but they may be positive or negative. To obtain the electron distribution, we square the function in Equation 1.1, which is written in two ways in Equation 1.2.

1.2

Taking the expanded version, we can see that the molecular orbital σ2 differs from the superposition of the two atomic orbitals (c1ϕ1)2+(c2ϕ2)2 by the term 2c1ϕ1c2ϕ2. Thus we have two solutions (Fig. 1.3). In the first, both c1 and c2 are positive, with orbitals of the same sign placed next to each other; the electron population between the two atoms is increased (shaded area), and hence the negative charge which these electrons carry attracts the two positively charged nuclei. This results in a lowering in energy and is illustrated in Fig. 1.3, where the horizontal line next to the drawing of this orbital is placed low on the diagram. In the second way in which the orbitals can combine, c1 and c2 are of opposite sign, and, if there were any electrons in this orbital, there would be a low electron population in the space between the nuclei, since the function is changing sign. We represent the sign change by shading one of the orbitals, and we call the plane which divides the function at the sign change a node. If there were any electrons in this orbital, the reduced electron population between the nuclei would lead to repulsion between them; thus, if we wanted to have electrons in this orbital and still keep the nuclei reasonably close, energy would have to be put into the system. In summary, by making a bond between two hydrogen atoms, we create two new orbitals, and *, which we call the molecular orbitals; the former is and the latter (an asterisk generally signifies an antibonding orbital). In the ground state of the molecule, the two electrons will be in the orbital labelled . There is, therefore, when we make a bond, a lowering of energy equal to twice the value of in ( the value, because there are two electrons in the bonding orbital).