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CONTENTS Preface ix 1 THE COVALENT BOND 1 1.1 MODELS OF CHEMICAL BONDING 1 Appendix 1 : Hybrid Orbitals 43 Appendix 2: Molecular Orbital Theory 50... The Electron Pair Bond-Lewis St

KATHLEEN SCHUELLER RICHARDSON

The Ohio State University

HARPER & ROW, PUBLISHERS

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To Nancy and Frank

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Copyright 0 1976 by Thomas H Lowry and Kathleen Schueller Richardson

All rights reserved Printed in the United States of America No part of this book may be used or reproduced in any manner whatsoever without written permission except in the case of brief quotations embodied in critical articles and reviews For information address Harper & Row, Publishers, Inc.,

10 East 53rd Street, New York, N Y 10022

Library of Congress Cataloging in Publication Data

Lowry, Thomas H

Mechanism and theory in organic chemistry

Includes bibliographical references and index

1 Chemistry Physical organic I Richardson,

Kathleen Schueller, joint author 11 Title

QD476.1A68 547' 1'3 75-43926

ISBN 0-06-044082-1

CONTENTS

Preface ix

1 THE COVALENT BOND 1

1.1 MODELS OF CHEMICAL BONDING 1

Appendix 1 : Hybrid Orbitals 43

Appendix 2: Molecular Orbital Theory 50

vi Contents

Rate Constant 1 13

Appendix 2: The Transition State Theory of Isotope Effects 120

3 ACIDS AND BASES 124

4 BIMOLECULAR SUBSTITUTION REACTIONS 170

4

177

GEN 318

7 ADDITION AND ELIMINATION REACTIONS 337

Contents vii

8 REACTIONS OF CARBONYL COMPOUNDS 402

8.3 ADDITION FOLLOWED BY ELIMINATION 424

8.4 ADDITION OF NITROGEN NUCLEOPHILES 432

11 THE THEORY OF PERICYCLIC REACTIONS 568

12 APPLICATIONS OF THE PERICYCLIC SELECTION

RULES 626

PROBLEMS 729

Index 731

PREFACE

This book is intended as a text for undergraduate and first-year graduate students who have completed a one-year course in organic chemistry Its aim

study in greater depth of individual organic reactions and of methods by which chemists obtain information about chemical processes

The primary focus of the book is on reaction mechanisms, not only because knowledge of mechanism is essential to understanding chemical processes but also because theories about reaction mechanisms can explain diverse chemical phenomena in terms of a relatively small number of general principles It is this latter capability of mechanistic theory which makes it important as an organizing device for the subject of organic chemistry as a whole

In treating mechanisms of the important classes of organic reactions, we have tried to emphasize the experimental evidence upon which mechanistic ideas are built and to point out areas of uncertainty and controversy where more work still needs to be done In this way we hope to avoid giving the impression that all organic mechanisms are well understood and completely agreed upon but instead to convey the idea that the field is a dynamic one, still very much alive and filled with surprises, excitement, and knotty problems

The organization of the book is traditional We have, however, b2en' selective in our choice of topics in order to be able to devote a significant portion

of the book to the pericyclic reaction theory and its applications and to include

a chapter on photochemistry

The pericyclic theory is certainly the most important development in mechanistic organic chemistry in the past ten years Because it is our belief that

x Preface

the ideas and method of thinking associated with the pericyclic theory will have

an increasing impact in both organic and inorganic chemistry in the future, we have given a more detailed discussion of its purely theoretical aspects than has heretofore been customary in books of this kind This discussion includes both the Woodward-Hoffmann approach and the Dewar-Zimmerman aromaticity approach and makes the connection between them Our treatment requires as background a more sophisticated understanding of covalent bonding than is ordinarily given in introductory courses; we have therefore included an exten- sive presentation of bonding theory I t begins at a basic level with a review of familiar concepts in Chapter 1 and introduces in Chapter 10 the terminology and ideas needed to understand the pericyclic theory and its ramifications The treatment is qualitative throughout Although quantitative molecular orbital calculations are not needed for our purposes, Appendix 2 to Chapter 1 sum- marizes the molecular orbital calculation methods in general use The Hiickel

M O method is covered in sufficient detail to allow the reader to apply it to simple systems

Another innovation in this text is the use of three-dimensional reaction coordinate diagrams, pioneered by Thornton, More O'Ferrall, and Jencks, in the discussions of nucleophilic substitutions, eliminations, and acid catalysis of carbonyl additions We hope that the examples may lead to more widespread use of these highly informative diagrams

A chapter on photochemistry provides a discussion of photophysical processes needed as background for this increasingly important area of chemistry and treats the main categories of light-induced reactions,

The text assumes elementary knowledge of the common organic spectro- scopic techniques Nevertheless, we have included a description of the recently developed method of chemically induced dynamic nuclear polarization (CIDNP), which has already proved to be of great importance in the study of radical reactions and which has not yet found its way into books covering spectroscopy of organic compounds

Problems of varying difficulty have been included at the ends of the chapters Some problems illustrate points discussed in the text, but others are meant to extend the text by leading the student to investigate reactions, or even whole categories of reactions, which we have had to omit because of limitations of space References to review articles and to original literature are given for all problems except those restricted to illustration of points that the text discusses in detail Problems that represent significant extensions of the text are included in the index

The book is extensively footnoted I t is neither possible nor desirable in

a book of this kind to present exhaustive reviews of the topics taken up, and we have made no effort to give complete references We have tried to include references to review articles and monographs wherever recent ones are avail- able, to provide key references to the original literature for the ideas discussed, and to give sources for all factual information presented The text also contains numerous cross references

The amount of material included is sufficient for a full-year course For

a one-semester course, after review of the first two chapters, material may be chosen to emphasize heterolytic reactions (Chapters 3-8), to cover a broader

Preface xi range including radicals and photochemistry (selections from Chapters 3-8 plus 9 and 13), or to focus primarily on pericyclic reactions (Chapters 10-12)

I n selecting material for a one-semester course, the following sections should be considered for possible omission : 3.5, 4.4, 4.5, 5.6, 6.3, 7.3, 7.5, 8.3, 9.5, 10.4, 11.6, 11.7

MTe would like to thank the following people for reviewing parts of the manuscript and for providing helpful comments: Professors D E Applequist,

C W Beck, J C Gilbert, R W Holder, W P Jencks, J R Keeffe, C Levin,

F B Mallory, D R McKelvey, N A Porter, P v R Schleyer, J Swenton, and T T Tidwell We are particularly grateful to Professor N A Porter, who reviewed and commented on the entire manuscript We owe special thanks to Professor Charles Levin for many enlightening discussions and to Carol Demp- ster for essential help and encouragement

Thomas H Lowry Kathleen Schueller Richardson

MECHANISM AND THEORY IN ORGANIC CHEMISTRY

Understanding and progress in natural science rest largely on models A little reflection will make it clear that much of chemical thinking is in terms of models, and that the models useful in chemistry are of many kinds Although we cannot see atoms, we have many excellent reasons for believing in them, and when we think about them we think in terms of models For some purposes a very simple model suffices Understanding stoichiometry, for example, requires only the idea

of atoms as small lumps of matter that combine with each other in definite pro- portions and that have definite weights The mechanism by which the atoms are held together in compounds is not of central importance for this purpose When thinking about stereochemistry, we are likely to use an actual physical model con- sisting of small balls of wood or plastic held together by springs or sticks Now the relative weights of atoms are immaterial, and we do not bother to reproduce them in the model; instead we try to have the holes drilled carefully so that the model will show the geometrical properties of the molecules Still other models

are entirely mathematical We think of chemical rate processes in terms of sets of differential equations, and the details of chemical bonding require still more ab- stract mathematical manipulations The point to understand is that there may be many ways of building a model for a given phenomenon, none of which is com-

plete but each of which serves its special purpose in helping us understand some

aspect of the physical reality

The Electron Pair Bond-Lewis Structures

The familiar Lewis structure is the simplest bonding model in common use in

organic chemistry It is based on the idea that, at the simplest level, the ionic

bonding force arises from the electrostatic attraction between ions of opposite

charge, and the covalent bonding force arises from sharing of electron pairs be-

tween atoms

The starting point for the Lewis structure is a notation for an atom and its

valence electrons The element symbol represents the core, that is, the nucleus and

V.3hmxkctrons are shown explicitly For elements in the third and later rows

:Br: :Se: :I:

Ions are obtained by adding or removing electrons The charge on an ion is

given by

charge = core charge - number of electrons shown exvlicidy

An ionic compound is indicated by writing the Lewis structures for the two ions

A covalent bond model is constructed by allowing atoms to share pairs of

electrons Ordinarily, a shared pair is designated by a line:

H-H

All valence electrons of all atoms in the structure must be shown explicitly Those

electrons not in shared covalent bonds are indicated as dots, for example:

If an ion contains two or more atoms covalently bonded to each other, the

total charge on the ion must equal the total core charge less the total number of

In order to write-correct Lewis structures, two more concepts are needed

First, consider the total number of electrons in the immediate neighborhood of

find it, all unshared electrons around the atom and all electrons in bonds leading

Models of Chemical Bonding 3

to the atom must be counted The valence-shell occupancy must not exceed 2 for hydrogen and must not exceed 8 for atoms of the first row of the periodic table For elements of the second and later rows, the valence-shell occupancy may exceed 8 The structures

are acceptable

The second idea is that of formal charge For purposes of determining formal charge, partition all the electrons into groups as follows: Assign - to each atom all of its unshared pair elec_trons and half of all electrons in bonds leading to -

i w d t h e n l l m h e r assigned tof,h_e.te@m by this p r ~ ~ W k ~ d e t # ~ ~ n

-

f o -

-To illustrate formal charge, consider the hydroxide ion, OH- The electron ownership of H is 1, its core charge is + 1, and its formal charge is therefore zero The electron ownership of oxygen is 7, and the core charge is + 6 ; therefore the formal charge is - 1 All nonzero formal charges must be shown explicitly in the structure The reader should verify the formal charges shown in the following examples :

The algebraic sum of all formal charges in a structure is equal to the total charge Formal charge is primarily useful as a bookkeeping device for electrons, but

it also gives a rough guide to the charge distribution within a molecule

In writing Lewis structures, the following procedure is to be followed:

1 Count the total number of valence electrons contributed by the electri-

cally neutral atoms If the species being considered is an ion, add one electron to the total for each negative charge; subtract one for each positive charge

2 Write the core symbols for the atoms and fill in the number of electrons determined in Step 1 The electrons should be added so as to make the valence- shell occupancy of hydrogen 2 and the valence-shell occupancy of other atoms not less than 8 wherever possible

3 Valence-shell occupancy must not exceed 2 for hydrogen and 8 for a first-row atom; for a second-row atom it may be 10 or 12

4 Maximize the number of bonds, and minimize the number of unpaired erectrons, always taking care not to violate Rule 3

5 Find the formal charge on each atom

We shall illustrate the procedure with two examples

Left 0 Ownership 6 0 charge

Right 0 Ownership 7 - 1 charge

(More bonds to C would exceed its valence-shell limit.)

be

Models of Chemical Bonding 5

There is a class of structures, however, for which the properties are not those expected from the Lewis structure A familiar example is benzene, for which the heat of hydrogenation (Equation 1.1) is less exothermic by about 37 kcal mole -

than one would have expected from Lewis structure 1 on the basis of the measured

heat of hydrogenation of ethylene The thermochemical properties of various types of bonds are in most instances transferable with good accuracy from molecule

to molecule; a discrepancy of this magnitude therefore requires a fundamental modification of the bonding model

The difficulty with model 1 for benzene is that there is another Lewis

structure, 2, which is identical to 1 except for the placement of the double bonds

Whenever there are two-ake_r-n_ativeLxwis _s_tructuyg_s,_one - - alone - will - be inaccurate representation of t h ~ d e _ c u _ l a r g~uct_ur_e, A more accurate picture will be obtained by the s u p e r p ~ s ~ i ~ n ~ o f the two structures into a new-model, which - for - benzene is - - indicated - - by 3 The superposition of two o r more Lewis structures into a composite - picture is called resonance

This terminology is well established, but unfortunate, because the term resonance when applied to a pair of pictures tends to convey the idea of a chang- ing back and forth with time I t is therefore difficult to avoid the pitfall of think- ing of the benzene molecule as a structure with three conventional double bonds,

of the ethylene type, jumping rapidly back and forth from one location to another This idea is incorrect The electrons in the~molecule~m_ove in-a field of

three sides and make them different from the other three, as the two alternative

pictures 3 seem to imply that they do

The symmetry of the ring of nuclei (4) is called a sixfold symmetry because

rotating the picture by one-sixth of a circle will give the identical picture again

This sixfold symmetry must be reflected in the electron distribution A less mis-

leading picture would be 5, in which the circle in the middle of the ring implies a

distribution of the six double bond electrons of the same symmetry as the arrange-

ment of nuclei We shall nevertheless usually continue to use the notation 3, as it

has certain advantages for thinking about reactions

The most important features of structures for which resonance is needed

are, first, that the -f lower energy) than on3 would

expect from l o o k iof ~the- structures, and second, that the actual

distribution of ~ m I I S i ~ e e m ~ o I _ e c u l e is different - frqm whhat*r?_e would expect

on the basis of one of the structures Since the composite picture shows that cer-

tain electrons are free to move alarger area of the molecule than a single one

of the structures implies, resonanse is often referred to as delocalization We shall

have more to say about delocalization later in connection with molecular orbitals

While the benzene ring is the most familiar example of the necessity for

modifying the Lewis structure language by the addition of the resonance concept,

there are many others The carboy& acids, for example, are much - s t r o n ~ acids than the alcohols; t~s~ff~em.musbedudar& tgsreate~stabili& of

t h e k&b6k$iii i d 6 k v e r the alkoxide ion17) ; it is the p w k d

two equivalent Lewis structures for the c a r h a x ~ i e n - ~ b ; b t - a l 9 1 % u t 0 ~

&fference,

Another example is the allylic system The ally1 cation (8), anion (9), and

radical (lo), are all more stable than their saturated counterparts Again, there is

for each an alternativestructure :

Models of Chemical Bonding 7

I n all the examples we have considered so far, the alternative structures have been equivalent This will not always be the case, as the following examples illustrate :

Whenever t h e r e 3 n o n e q u i v a Q to contribute^^

stable Ilowest-energy ) molecule - w ~ u c L m _ 1 _ e c u l e _a.ctualk foe xi st,^ .conzi- butes_the-m.~st_to_ the composite,~ an6.others successively -less as they represent

h i & e r e e c u l e s

I t is because the lowest-energy structures are most important that we speci- fied in the rules for writing Lewis structures that the number of bonds should be maximum and the valence-shell occupancy not less than 8 whenever possible Structures that violate these stipulations, such as 11 and 12, represent high-energy forms and hence do not contribute significantly to the structural pictures, which

are quite adequately represented by 13 and 14:

The followinp; rules are useful in using resonance notatinn:

1 All nuclei must be in the same location in every structure Structures with

nuclei in different locations, for example 15 and 16, are chemically distinct sub-

stances, and interconversions between them are actual chemical changes, always designated by +

2 sStr_uc!ur~_w_ithfewer_ bon_d_s_so_ro_rw&hgeatte~I:~eparati.o-~~.of for~al-&a_rge are

less stable than those with more bonds or less charge sparation Thus .- - - 11 and

12 are higher-energy, respectively, than 13 and 14

3 W h e r e t w ~ ~ c u - e s with f~maL&@avt= t W enumber of bmds

and appra-xhubely t h e s m l u h r g ~ - s _ e ~ a r a t i o n ? the structure with c h a g ~ ~ & e

more electronegative atom will usually besomewhat lower in energy, but the_

difference will ordinarilv be small enough that both structures _- must

be-in the - composit~pi~ttllre Thus in 17a t, 17b, 17a should be more stable, but the

any structure must lie in t h e x m e plane For example, the structure 18b cannot

contribute, because the bridged ring prevents carbons 6 and 7 from lying in the

same plane as carbon 3 and the hydrogen on carbon 2 The i m p ~ s i b i l i t y d s t m -

a i t h d o ~ _ b _ l e h n n ~ ~ a l L h r l d g ~ i s ~ B r & s ~

&.l Double bonds can occur at a bridgehead if the rings are sufficiently large

Molecular Geometry

Lewis structures provide a simple method of estimating molecular shapes The

geometry about any atom covalently bonded to two or more other atoms is found

by counting the number of electron groups around the atom Each unshaared pair

counts as one group, and each bond, w h e t h e ~ s ~ n ~ l ~ r multiple4 counts a$ one

group The number of electron a r o u p a m n d a ~ ~ q uto the a l

sum of the number of electron pairs on the atomand-the number of other atoms

bonded to it The Peometry islinear if the number of electron goups is two, tri-

gonal if the number is thre~md-r

The rule is based on the electron-pair repulsion model, which postulates that-

m eelectron pairs repel each other, thev will try to stay as far apart as possible

I n trigonal and tetrahedral geometries, the shape will be exactly trigonal (120"

bond angles), or exactly tetrahedral (109.5" bond angles) if the electron groups

are all equivalent, as for example in BH, or CH, + (trigonal), or in CH, or NH, +

(tetrahedral) -

( a ) F S Fawcett, Chem Rev., 47,219 (1950); (b) J R Wiseman and W A, Pletcher, J Amer Chem

Sac., 92, 956 (1970); (c) C B Quinn and J R Wiseman, J Amer Chem Sac., 95, 6120 (1973);

(d) C B Quinn, J R Wiseman, and J C Calabrese, J A m r Chem Soc., 95, 6121 (1973)

Molecular Orbitals 9

If the groups are not all equivalent, the angles will deviate from the ideal values Thus in NH, (four electron <groups, three in N-H bonds, one an unshared pair), the unshared pair, being attracted only by the nitrogen nucleus, will be closer to the nitrogen on the average than will the bonding pairs, which are also attracted by a hydrogen nucleus Therefore the repulsion between the unshared pair and a bonding pair is greater than between two bonding pairs, and the bonding pairs will be pushed closer to each other The H-N-H angle should therefore be less than 109.5" I t is found experimentally to be 107" Similarly, in H,O (four electron groups, two unshared pairs, and two 0-H bonds), the angle

is 104.5"

Ambiguity may arise when more than one structure contributes Then un- shared pairs in one structure may become multiple bonds in another, so that the number of electron groups around a given atom is not the same in both structures

An example is methyl azide (19) The central nitrogen is clearly linear (two electron groups), but the nitrogen bonded to CH, has three electron groups in

19a and four in 19b I n such a situation, the number of electron groups is deter- mined from the structure with the larger number of honds Thus the nitrogen in

question in 19 is trigonal, not tetrahedral

Conventions for Structural Formulas

This book contains large numbers of Lewis structural formulas Frequently we shall not write out the full Lewis structure; unshared pairs of electrons not shown explicitly are implied When there are two or more contributing structures, we shall show them all only if that is essential to the point being illustrated; again, it will be assumed that the reader will understand that the missing structures are implied

Lewis structures serve admirably for many aspects of mechanistic organic chemistry Frequently, however, we need a more accurate bonding model

Models Based on the Quantum Theory

The description of chemical bonding must ultimately be based on a n understand- ing of the motions of electrons I n order to improve our model, we need to appeal

to the quantum theory, which summarizes the current understanding of the be- havior of particles of atomic and subatomic size

The quantum theory provides the mathematical framework for describing the motions of electrons in molecules When several electrons are present, all interacting strongly with each other through their mutual electrostatic repulsion, the complexity is so great that exact solutions cannot be found Therefore approximate methods must be used even for simple molecules These methods

take various forms, ranging from complex ab initio calculations, which begin from first principles and have no parameters adjusted to fit experimental data, to highly approximate methods such as the Hiickel theory, which is discussed further

in Appendix 2 The more sophisticated of these methods now can give results

of quite good accuracy for small molecules, but they require extensive use

of computing e q ~ i p m e n t ~ Such methods are hardly suited to day-to-day qualita- tive chemical thinking Furthermore, the most generally applicable and therefore most powerful methods are frequently simple and qualitativẹ

Our ambitions in looking at bonding from the point of view of the quantum theory are therefore modest We want to make simple qualitative arguments that will provide a practical bonding model

Atomic Orbitals

The quantum theory specifies the mathematical machinery required to obtain a complete description of the hydrogen atom There are a large number of func- tions that are solutions to the appropriate equation; they are functions of the x,

y, and z coordinates of a coordinate system centered at the n u c l e u ~ ~ Each of these functions describes a possible condition, or state, of the electron in the atom, and each has associated with it an energy, which is the total energy (kinetic plus potential) of the electron when it is in the state described by the function in question

The functions we are talking about are the familiar Is, 2s, 2P, 3s,

atomic orbitals, which are illustrated in textbooks by diagrams like those in Figure 1.1 Each orbital function (or wave function) is a solution to the quantum mechanical equation for the hydrogen atom called the Schrodinger equation The functions are ordinarily designated by a symbol such as g,, X, $, and so on

We shall call atomic orbitals g, or X, and designate by a subscript the orbital meant, as for example g,,,, g,,,, and so on Later, we may abbreviate the notation

by simply using the symbols Is, 2s, , to indicate the corresponding orbital functions Each function has a certain numerical value at every point in space; the value at any point can be calculated once the orbital function is known We shall never need to know these values, and shall therefore not give the formulas; they can be found in other sourcệ^ The important things for our purposes a s

fiist, that t k e m e sare positive in certain regions ocspace and neg? tive in other regions, and second, that the value of each function approaches zero

a A number of texts cover methods for obtaining complete orbital descriptions of molecules Ex-

amples, in approximate order of increasing coverage, are (a) Ạ Liberles, Introduction to Molecular- Orbital Theory, Holt, Rinehart, and Winston, New York, 1966; (b) J D Roberts, Notes on Mokcular Orbital Theory, W Ạ Benjamin, Menlo Park, Calif., 1962; (c) K B Wiberg, Physiral Organic Chemistry, Wiley, New York, 1964; (d) Ạ Streitwieser, Jr., Molecular Orbital Theory for Organic Chemists, Wiley, New York, 1961; (e) M J S Dewar, The Molecular Orbital Theory of Organic Chemistry, McGraw-Hill, New York, 1969; (f) P ÓD Offenhartz, Atomic and Mokcular Orbital Theory, McGraw-Hill, New York, 1970; (g) S P McGlynn, L G Vanquickenborne, M Kinoshita,

and D G Carroll, Introdudion to Applied Quantum Chemistry, Holt, Rinehart, and Winston, New York,

Molecular Orbitals 11

Figure 1.1 Hydrogen atomic orbital functions (a) Is; (b) 2p; (c) 3d The edges drawn are

artificial, because orbitals have no edges but merely decrease in magnitude as distance from the nucleus increases The important features of the orbitals are the nodal planes indicated, and the algebraic signs of the orbital functions, posi- tive in the shaded regions and negative in the unshaded regions

grams used throughout this book, positive regions are shaded and negative regions are unshaded

Imagine walking around inside a n orbital, and suppose that there is some

region, you must pass through some point where the value is zero T-ctions

of all ad.jacent p o i n ~ ~ h i c k a L n c t ~ - z e r a z e r a ~ ~ e e c ~ ~ d ~ o ~ e s ; they are surfaces

in three-dimensional space, and most of the important ones for our purposes are

can also be spherical, and of other shapes, but these are of less concern to us.)

The Physical Significance of Atomic Orbital Functions

has no particular physical significance for the behavior of an electron that finds itself in the state defincd by the orbital (We shall scc shortly that the significance

of the signs comes from the way in which orbitals can be combined with each other.) The quantity that has physical meaning is the value at each point of the

<arying density showing the relative probabilityoffinding the electron in various regions or, more succinctly, the electron dzrtrzbuLzon or electron denrzty, are ~ t u a l l y

pi-res_of w2 not of o, itself T h e g e n e m l s h p e a f ~ will be s i m i l a a ~ h e s h a p e of2 T h e orbitals and their squares have no edges, even though definite outlines are usually drawn in diagrams; the values merely approach closer and closcr to zero as one goes farther and farther from the nucleus

Extension to Other Atoms

The hydrogen atomic orbitals would not do us a great deal of good if orbitals of other atoms were radically different, since in that case different pictures would

be required for each atom But the feature of the hydrogen atom problem that determines the most important characteristics of the hydrogen atom orbitals is the spherical symmetry Since all the atoms are spherically symmetric, the atomic orbitals of all atoms are similar, the main difference being in their radial depen- dence, that is, in how rapidly they approach zero as one moves away from the nucleus Because the radial dependence is of minimal importance in qualitative

Figure 1.2 Electron density, v2, for 1s and 2 p atomic orbitals T h e density of shading is

roughly proportional to v2

Molecular Orbitals 13

applications, one may simply use orbitals of the shapes found for hydrogen to describe behavior of electrons in all the atoms

Ground and Excited States

this statement in a more abbreviated form by saying that the electron is in one of

statement henceforth

T h e orbital that has associated with it the lowest energy is y,,; if the electron

electronic ground state If we were to give the electron more energy, say enough to

put it in the 932px orbital, the atom would be in a n electronic excited slate I n general,

sible energy orbitals (remembering always that the Pauli exclusion principle prevents more than two electrons from occupying the same orbital) is the elec- tronic ground state Any higher-energy state is an electronic excited state

An Orbital Model for the Covalent Bond

Suppose that we bring together two ground-state hydrogen atoms Initially, the

atoms a r e very close, say within 1 A ( = l o - * cm) of each other, each electron will feel strongly the attractive force of the other nucleus as well as of its own Clearly, then, the spherical p,, orbitals will no longer be appropriate to the description of the electron motions We need to find new orbital functions appro- priate to the new situation, but we would prefer to do so in the simplest way possible, since going back to first principles and calculating the correct new orbi-

We therefore make a guess that a possible description for a new orbital

of p,,, and adding the two numbers tog-ether This process will give us a new

also be positive everywhere Figure 1.3 illustrates the procedure Mathematically,

orbital is any orbital function that extends over more than one atom.Since_the-

Figure 1.3 T h e linear combination of 1s orbital functions on hydrogen atoms A and B to

yield a new orbital function, I,!JMO = q~~~~ + v l P ~

inthi.s_mol.ec&ar orbital, because a t - a g p o i n t - + equidistant from-the - two nuclei

- -

zero

The procedures of the quantum theory require that the negative combina-

with #.&dl-bc higher - athan that of FICA and ?La

Energies of Molecular Orbitals

We can summarize the process of constructing our bonding model in an energy-

formation of new orbitals by combining others O n either side we place the starting orbitals, and a t the center the orbitals resulting from the combination

Molecular Orbitals 15

Figure 1.4 T h e linear c o m b i n a t i o n of Is o r b i t a l functions o n h y d r o g e n a t o m s A a n d B t o

yield o r b i t a l function $Go = - tpl,~

91,,,, will come together to give Hz with a pair of electrons in $, and will in the

4 AE to the ground-state molecule and placing either one or both electrons in

*Go

destabilizing Ther-efore w e - ~ a l l _ + ~ -a bonding- o h i ~ a ! and $$a t;,n-g&bonding-

orbital - I n antibonding orbitqls there is - always - a node between the nuclei, so that

- -

T h e energy change o n formation of the molecule, known from experiment to be 104 kcal mole-', will not actually be equal to 2 AE, because the quantum mechanical procedures count the mutual

repulsion of the electrons twice and neglect the mutual repulsion of the nuclei T h e two corrections

to 2 AE are opposite in sign and roughly cancel, but they are both large numbers (on the order of

4 0 0 4 5 0 kcal mole-' for Hz), a n d their difference (about 35 kcal mole-') is significant T h e actual energy lowering is less than 2 AE by this amount; in other words, for hydrogen the actual experi-

mental dissociation energy is 104 kcal mole-', but 2 AE calculated from theory is about 139 kcal

mole-' and AE is about 69 kcal mole-' See C A Coulson, Valence, 2nd ed., Oxford University

Press, London, 1963, p 90

Figure 1.5 T h e energy-level d i a g r a m for the interaction o f v l S A with v,,, On either side a r e

the atomic orbitals before interaction; a t the center a r e t h e two molecular orbitals Orbital occupancies a r e indicated for the two separate hydrogen atoms

a n d for the molecule

electrons - are - excluded from that r e ~ i o n , whereas bonding orbitals have no such

Interaction of Orbitals

q ? l s ~ and v,,, as interacting with each other to produce the two new orbitals

+,

O C C U ~ S ~ - $ ~ ~ moves-do-wn by interaction energy AE and I/&, moves u p by inter-

-

a c t i o ~ e n e r g y AE.7

A somewhat more careful treatment shows that +hco will actually have moved up above the a,,,

level by somewhat more than t,h,o moved down This fact will be important in certain applications later, b u t necd not concern us now

Molecular Orbitals 17

Figure 1.6 The three-dimensional shapes of #so and : #: Each has infinite-fold rotational

symmetry, because one can rotate each picture around the internuclear axis in

an infinite number of steps and have at every step an identical picture

As we have noted above, A E can be calculated, but for our purposes we need only to know what quantities affect its magnitude x h e h a c t i o n energy is reater the more s t r o n ~ l y the two interacting orbitals overlap~c.ye_r_Ia~ is large

when both orbitals have yahe.s in the same region of space T&e,v_-ay-of two o r b i t ~ ~ ~ ~ - ~ w ~ 2 ~ i f - O b t a W : d - ~ - ~ p Y ; n g &e-valEs of thetwo functions a t each ~ o i n t pr oducts over all poi-nns, in-ot_herrw&s

The second factor - affecting- e of A E is whether or not t h e tky- interactinp orbitaTsare nf t h p c t i o n - i s-maxi- mum .- when the energies of t h e intPr;lrtiffP.- and_ becor@$ smaller the farthchapart.i.nincrgy they are We shall return to consider the over- lap and the energy differences between the initial orbitals in more detail in Chapter 10

The H, model has illustrated an important point about orbital interactions which must be remembered : - W h e n e v ~ r h i s orbital d&&ionr interact to form new orbital junctions, the number new f i r o btained i s equal-to the number bmzs

mans used

aBondsand~Bonds

In Figure 1.6 are shown the three-dimensiona1 shapes of the electron distributions

#iO and $5; corresponding to the H, molecular orbitals Suppose that we were

Figure 1.7 Combination of two p orbitals to give o molecular orbitals (a) Bonding com-

bination (b) Antibonding combination

to rotate one of these pictures around an axis coinciding with the line joining the nuclei We can rotate around this axis by any angle at all, and we shall get an identical picture If you were to close your eyes while the rotation was done and then to open them, you would have no way of telling that any change had been made T o state this idea another way, we can say that we could divide one full rotation around the axis into an infinite number ofsteps, and have after each step

a n indentical picture

Molecular Orbitals 19

Figure 1.8 Combination of p orbitals to give rr molecular orbitals ( a ) Bonding combina-

tion ( b ) Antibonding combination

- _I

is called a - o -orbztal Both , (G, and (GCo of o u r hydrogen molecule model are a

orbitals

y ! ~ & ~ ~ ~ ~ and I,!I&~,~~ Now the symmetry is different: O n e fhe-ul

m a e l e m e n t is called a 6, axis An mhitaly&h this kind o f z ~ m e t r y is called2

-rr orbital Atomic orbitals of the s type can form only o molecular orbitals; atomic

Figure 1.9 The symmetry of the electron distributions I ) $ ~ ~ ~ ~ and I)~~,~, Rotational

Suppose that we wish to construct an LCAO bonding model for methane We set

up the problem by defining an x, y, z coordinate system and placing the carbon at the origin The molecule is tetrahedral, as determined from the electron-pair repulsion model The orientation of the molecule is arbitrary; we choose to arrange it as shown in 21, with the hydrogen atoms in the + x, +y, + z quadrant, the - x, -y, + z quadrant, the +x, -y, - z quadrant, and the - x, +y, - z

quadrant

Hybrid Orbitals 21

Figure 1.10 The valence atomic orbitals of the carbon and four hydrogens in methane

Zp,, and 2p, (Figure 1.10)

The Need for Hybrid Orbitals

We could simply proceed to inspect these orbitals to see which overlap with each other, and then begin to make molecular orbitals in the way described in the

previous section Unfortunately, the situation is now quite complicated The y,,, orbital of hydrogen number 1 interacts with all four of the carbon valence orbitals

T h e quantum theory gives procedures for dealing with this situation; for calculations done with the aid of a computer, there is no disadvantage in using

cumbersome; we are looking for a simplcr model that will allow us to see quickly and clearly what the final outcome of this complex set of interactions will be

Constructing Hybrids

The strategy we adopt is to look first a t the atomic orbitals of the central atom, and to decide on the basis of the geometry which orbitals are going to interact

in uthis instance as p sp3 because it is form-ed-from a n s and three p orbitals

T h e process of forming hybrids is not the s a m e a s the o_r_bdal intcraction process~h_a_t occ-urs on bringing - - two - atoms together There is no molecular orbital formation involved, because_weeare still talking about only one atom, and there is no-energy lowering T h e energy of a hybridorbital is between the energies of the orbitals - from-trmadk J r a t h < i h a n ~ ~ f i ~ ~ K i h e r - or lower

T h e reader should convince himselfthat3lEGlfowing four ways of adding

identical in shape to ,ySP31, shown in Figure 1.11, but each oriented toward a dif- ferent one of the [our hydrogen atoms:

T h e actual correct mathematical forms are not exactly as indicated in Equations 1.6 T h e sign of each term, which is the important quantity for our present purpose, is correctly represented there, but each orbital function must be

desired geometry is given in Appendix 1 to this chapter

Figure 1.11 T h e formation of a n sp3 hybrid by adding together the four valence atomic

orbitals Orbital shapes and locations of the nodes are approximate in these diagrams For a more accurate description, see Wiberg, Physical Organic

Chemistry, pp 29-33

Hybrid Orbitals 23

Figure 1.12 Formation of a bonding molecular orbital from F,,, and ~~~3~

Figure 1.13 T h e energy relationships in MO formation from ~~~3~ and p , , ~

T h e advantage we gain by making hybrid orbitals is that we now have four new atomic orbitals on carbon, each one oriented directly toward one of the hydrogen atoms Each hybrid will have a large overlap and therefore a large

orbitals, is now replaced by four separate but simple problems

MOs from Hybrid Orbitals

and x S p s l ; there will also be a n antibonding combination which has a node be-

the same and have different energies T h e energy difference in this instance is not large, and makes no fundamental change in our model We shall return to this

Figure 1.14 Formation of s,h2 hybrids from an s and two p orbitals

Hybrid Orbitals 25

Figure 1.15 Formation of sp hybrids from an s and one p orbital

atomic orbitals

Other Types of Hybridization

T h e relative _ contri4-Lutir>ns.of~andp _ _ orbitals - - to - the hpbrids_isidiffer-g-nt f o _ ~

electrons - - Y which - - have a node at_the-nu_cl~s,_s_electrons are held more tightly Therefore a n atom is effectively more electronegative in bonds that use a larger

tions change with changing geometry

a And T Bonding Ethylene

Figure 1.16 T h e basis orbitals for the a MO's of ethylenc

1.16 shows these hybrids, together with the hydrogen 1s orbitals T h e orbitals are

Figure 1.17, which also shows how the 12 valence electrons are assigned to the molecular orbitals in the electronic ground state

levels will be a t different energies rather than all the same as shown in the figure

Figure 1.17 Energy-level diagram for the bonding model of ethylene T h e ow levels are not

actually all a t the same energy, but are lower than T C C

Delocalized rr Bonding 27 But the important point for most purposes is that the highest-energy bonding M O

1.4 DELOCALIZED 7C BONDING

23

orbitals

Formation of rr Systems

one other orbital One approach is to go to the quantum theory rules and work

through the prescribed procedures to find how the three orbitals will combine

The method, a t a n approximate level, is the Hiickel theory I t is described in detail

T h e next-higher e n e r ~ - - M Q i s

- - -

I t looks like 26:

a,!~,,, has a node cutting across it and passing through the central carbon; basis

is the same as that of the basis orbitals themselves: so that electrons in it do not

# M O ~ - _ = PI - - _ P 2 - + @3 -

a n d looks like 27:

I t has a node between each bonded pair of carbons and is antibonding Figure

concentrated a t the ends of the c h a h ; the molecular orbital pictures for these

which show the charge or unpaired electron to be concentrated a t the ends

pentadienyl I n each case the lowest-energy orbital has no veitical nodes, and each higher-energy orbital has one more vertical node than the orbital below it had, with the highest-energy orbital always having a node between every adja- cent pair of atoms Chains with a n odd number of atoms have a nonbonding

than encompassing only two, as have the MO's we considered earlier Orbitals

1.5 AROMATICITY

T h e concept of aromaticity has been extremely fruitful for both theoretical and experimental organic chemists Aromatic compounds are cyclic unsaturated molecules characterize8 by certain magnetic effects and by substantially lower chemical reactivity and greater thermodynamic stability than would be expec~ed from localized bond models

Figure 1.18 T h e .rr MO's of the ally1 system T h e basis orbitals from which the rr MO's are

constructed are shown at the top of the figure, and below are the molecular orbitals in an energy-level diagram

Resonance and Aromaticity

The familiar properties of benzene illustrate the characteristics of aromatic com- pounds Benzene is much less reactive toward electrophiles, such as molecular halogens, than are simple olefins; and the heat evolved on hydrogenation is less

type double bonds Furthermore, the nuclear magnetic resonance spectrum of benzene and its derivatives shows the protons bonded to the ring to be experienc- ing a stronger effective magnetic field than do protons attached to simple olefins

As we have seen, these properties are accounted for in the resonance picture

by modifying the model through inclusion of a second structure with double bonds in the alternative locations