Organic Chemistry: Structure, Mechanism, Synthesis

1033 Pages·2018·96.16 MB·English

by  J. David Rawn, Robert J. Ouellette

O c :
rganic hemistry
s , m , s
tructure echanism ynthesis
O c :
rganic hemistry
s , m , s
tructure echanism ynthesis
r J. O
Obert uellette
Professor Emeritus, Department
of Chemistry, The Ohio State University
a
nd
J. d r
avid awn
Professor Emeritus,
Towson University
Academic Press is an imprint of Elsevier
125 London Wall, London EC2Y 5AS, United Kingdom
525 B Street, Suite 1800, San Diego, CA 92101-4495, United States
50 Hampshire Street, 5th Floor, Cambridge, MA 02139, United States
The Boulevard, Langford Lane, Kidlington, Oxford OX5 1GB, United Kingdom
© 2018, Elsevier Inc. All rights reserved.
No part of this publication may be reproduced or transmitted in any form or by any means,
electronic or mechanical, including photocopying, recording, or any information storage and
retrieval system, without permission in writing from the publisher. Details on how to seek
permission, further information about the Publisher’s permissions policies and our arrangements
with organizations such as the Copyright Clearance Center and the Copyright Licensing Agency,
can be found at our website: www.elsevier.com/permissions.
This book and the individual contributions contained in it are protected under copyright by the
Publisher (other than as may be noted herein).
Notices
Knowledge and best practice in this field are constantly changing. As new research and experience
broaden our understanding, changes in research methods, professional practices, or medical
treatment may become necessary.
Practitioners and researchers must always rely on their own experience and knowledge in evaluating
and using any information, methods, compounds, or experiments described herein. In using
such information or methods they should be mindful of their own safety and the safety of others,
including parties for whom they have a professional responsibility.
To the fullest extent of the law, neither the Publisher nor the authors, contributors, or editors,
assume any liability for any injury and/or damage to persons or property as a matter of products
liability, negligence or otherwise, or from any use or operation of any methods, products,
instructions, or ideas contained in the material herein.
Library of Congress Cataloging-in-Publication Data
A catalog record for this book is available from the Library of Congress
British Library Cataloguing-in-Publication Data
A catalogue record for this book is available from the British Library
ISBN 978-0-12-812838-1
For information on all Academic Press publications
visit our website at https://www.elsevier.com/books-and-journals
Publisher: John Fedor
Acquisition Editor: Emily M. McCloskey
Editorial Project Manager: Katerina Zaliva
Production Project Manager: Maria Bernard
Cover Designer: Christian Bilbow
Typeset by SPi Global, India
To our families

All life is repetition, but repetition of increasing concentration.
-
Egon Friedell, A Cultural History of the Modern Age
2
1 S B
tructure and onding
o c
in rganic ompoundS
S S
O
N N
O
O
Firefly Luciferin
Organic chemistry began to emerge as a science about 200 years ago. By the late eighteenth century,
substances had been divided into inorganic and organic compounds. In those days, early in the
history of organic chemistry, inorganic compounds were isolated from mineral sources, and organic
compounds were obtained only from plants or animals. Organic compounds were more difficult to
study in the laboratory, and decomposed more easily than inorganic compounds. The differences be-
tween inorganic and organic compounds were attributed to a doctrine called vitalism in which a “vital
force” was thought necessary for the synthesis of organic compounds. In 1828 Friedrich Wöhler
O
synthesized the organic compound urea from inorganic starting materials, showing that a compound
NH C NH synthesized by living cells could be synthesized in the laboratory without biological starting materials,
2 2
urea countering a fundamental tenet of vitalism.
The organic compounds we will discuss throughout this text contain carbon and a few other
elements, such as hydrogen, oxygen, and nitrogen. We will also examine compounds containing sul-
fur, phosphorus, and halogens. Many, more exotic, organic compounds are also known, and organic
compounds have been made that contain virtually every element in the periodic table.
The molecule shown above is firefly luciferin, a light-emitting compound that is responsible
for the characteristic emission of yellow light from many species of fireflies. Its molecular formula is
C HNOS. The structure of luciferin is an example of the amazing variety of structures of organic
11 8 2 3 2
compounds. They are everywhere in nature, including interstellar space. No known living organism
can exist without organic compounds, and synthetic organic compounds are an integral part of the
objects we use every day. Organic chemistry is everywhere.
1.1 The physical and chemical properties of a molecule depend on the bonds that hold it together.
BRIEF REVIEW OF ATOMIC And these bonds depend on the electron configurations of its atoms. Therefore, we will review some
of the electronic features of atoms and the periodic properties of the elements before describing
STRUCTURE
bonding and its relation to structure in organic compounds.
Atomic Structure
Each atom has a central, small, dense nucleus that contains protons and neutrons that are embed-
ded in a sea of electrons. The atomic number, which equals the number of protons in the nucleus,
determines the identity of an atom. Because atoms have an equal number of protons and electrons
and are electrically neutral, the atomic number also equals the number of electrons in an atom.
The elements in the periodic table are arranged by atomic number. The elements are ar-
rayed in horizontal rows called periods and vertical columns called groups. In this text, we will
emphasize hydrogen in the first period and the elements carbon, nitrogen, and oxygen in the second
period. The electronic structure of an atom determines its chemical reactivity.
Atomic Orbitals
The electrons in an atom occupy atomic orbitals, which are designated by the letters s, p, d, and f.
Each orbital can contain a maximum of two electrons. An atomic orbital is a mathematical equation
Organic Chemistry. https://doi.org/10.1016/B978-0-12-812838-1.50001-3 Copyright © 2018 Elsevier Inc. All rights reserved. 1
that describes the energy of an electron. The square of the equation for the atomic orbital defines the
probability of finding an electron within a given region of space.
Orbitals are grouped in shells of increasing energy, designated by the integers 1, 2, 3, 4, . . . , n.
These integers are called principal quantum numbers. Each shell contains a unique number and
type of orbital. The first shell contains a single 1s orbital. The second shell contains one 2s orbital
and three 2p orbitals. Each orbital can contain no more than two electrons, and two electrons in any
orbital must have opposite spin. We need to consider only the orbitals of the first three shells for the
elements commonly found in organic compounds.
All s orbitals are spherically symmetrical (Figure 1.1a). The 2s orbital is larger than the 1s
orbital. A 2s orbital is farther from the nucleus, and it has a higher energy than a 1s orbital. The three
p orbitals in a shell are not spherically symmetrical. Electron density in each p orbital is concentrated
in two regions or lobes—one on each side of the nucleus. The two lobes together are the orbital. The
shapes of the p orbitals are shown in Figure 1.1b. The p orbitals are often designated as p, p, and
x y
p. They are mutually perpendicular to one another, and they are aligned along the x, y, and z axes.
z
Although the orientations of the p, p, and p orbitals differ, the electrons in each p orbital have equal
x y z
energies.
Orbitals of the same type within a shell constitute a group called a subshell. For example, an
s subshell has one orbital and can contain only two electrons. In contrast, a p subshell, which begins
in period two, contains three p orbitals and can contain a total of six electrons.
Electrons are distributed in subshells to give an electron configuration that has the lowest
energy. The order of increasing energy of subshells is 1s < 2s < 2p < 3s < 3p for elements of atomic
number less than 18. For any subshell, the lowest energy state is the arrangement that maximizes the
number of electrons having the same spin. This generalization is Hund’s Rule. This means that elec-
trons first occupy orbitals one at a time within subshells before pairing in a common orbital. Table
1.1 shows the atomic numbers and electron configurations for the first two periods in the periodic
table.
Figure 1.1 Shapes of 2s and 2p z 2pz
Orbitals
(a) An orbital is a boundary surface
enclosing a volume where electrons can
be located with 90% probability. An s x 2px
orbital has a spherical boundary surface.
(b) Boundary surfaces of the three
y 2p
mutually perpendicular 2p orbitals. y
Each orbital can hold a maximum of
two electrons. The + and – signs on the s orbital 2p orbitals
orbitals refer to the phase of the orbital, (a) (b)
not to the charge of the orbital.
Table 1.1
Electron Configurations of First and Second Period Elements
Element Atomic Number 1s 2s 2p 2p 2p Electron Configuration
x y z
H 1 1 1s1
He 2 2 1s2
Li 3 2 1 1s2 2s1
Be 4 2 2 1s2 2s2
B 5 2 2 1(↑) 1s2 2s2 2p1
C 6 2 2 1 (↑) 1 (↑) 1s2 2s2 2p2
N 7 2 2 1 (↑) 1 (↑) 1 (↑) 1s2 2s2 2p3
O 8 2 2 2(↑↓) 1 (↑) 1 (↑) 1s2 2s2 2p4
F 9 2 2 2 (↑↓) 2 (↑↓) 1 (↑) 1s2 2s2 2p5
Ne 10 2 2 2 (↑↓) 2 (↑↓) 2 (↑↓) 1s2 2s2 2p6
2
Valence Shell Electrons
Electrons in filled, lower energy shells of atoms have no role in determining the structure of molecules.
They are called core electrons. They do not participate in chemical reactions because they are held
too tightly to the nucleus. Only the higher energy electrons, which are located in the outermost
shell, called the valence shell, participate in chemical bonding. These are the valence electrons.
For example, the single electron of the hydrogen atom is a valence electron. The number of valence
electrons for the common atoms contained in organic molecules is given by their group number
in the periodic table. Thus, carbon, nitrogen, and oxygen atoms have four, five, and six valence
electrons, respectively. With this information we can understand how these elements combine to
form organic compounds.
1.2 The elements in the periodic table are arranged by atomic number. The elements are arranged in
ATOMIC PROPERTIES horizontal rows called periods and vertical columns called groups. The physical and chemical
properties of an element can be estimated from its position in the periodic table. Two properties that
help us explain the properties of organic compounds are the atomic radius and electronegativity.
Atomic Radius
The overall shape of an atom is spherical, and its volume depends both on the number of electrons
and on the energies of the orbitals the electrons occupy. The sizes of some atoms, expressed as the
atomic radius, in picometers (pm,10-12 m), are given in Figure 1.2 in a greatly abbreviated peri-
odic table that shows the atoms we will most commonly encounter in our discussion of organic
compounds. Atomic radii increase from top to bottom in a group of the periodic table because the
electrons in each new shell are located at greater distances from the nucleus. Thus, the atomic radius
of sulfur is greater than that of oxygen, and the radii of the halogens increase in the order F < Cl <
Br < I.
The atomic radius decreases from left to right across a period. Although electrons are lo-
cated in the same energy level within the s and p orbitals of the elements, the nuclear charge increases
from left to right within a period. These electrons are not shielded very well from the nuclear charge,
and the atomic radius decreases. The radii of the common elements in organic compounds decrease
in the order C > N > O.
Figure 1.2
H
Atomic Radii in Picometers,
37
pm (10-12 m)
Li Be B C N O F
152 111 88 77 70 66 64
Na Mg Al Si P S Cl
186 160 143 117 110 104 99
Br
114
I
133
Electronegativity
Electronegativity is an index of the tendency of an atom to attract electrons. It is proportional to
the difference between an atom’s ionization potential and its electron affinity. Linus Pauling placed
electronegativity values on a scale of slightly less than 1.0 for alkali metals to a maximum of 4.0 for
fluorine (Figure 1.3). The alkali metals and alkaline earth metals tend to lose an electron to gain an
inert gas configuration. Thus, groups I and II contain the least electronegative atoms. In fact they
are electropositive. On the other end of the scale, halogens, in group VII, tend to gain an electron
to give an inert gas configuration. Thus, we find that electronegativity increases from left to right
across the periodic table. Electronegativity values increase in period 2 in the order C < N < O < F.
Electronegativity values decrease from top to bottom within a group of elements. We will often use
these periodic trends to interpret the chemical and physical properties of organic compounds.
3
Figure 1.3 H
Electronegativity
2.1
Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Na Mg Al Si P S Cl
0.9 1.2 1.5 1.8 2.1 2.5 3.0
Br
2.8
I
2.5
Problem 1.1
A few proteins contain selenocysteine, which contains a selenium atom in place of the sulfur atom
of the amino acid cysteine. Selenium is in the fourth period, just below sulfur. Is sulfur or selenium
more electronegative?
1.3 In 1916, long before the advent of modern theory, the American chemist G. N. Lewis proposed that
IONIC AND COVALENT elements react to obtain the electron confi gurations of the inert gases. Th is hypothesis is summarized
in the Lewis octet rule for second period atoms: Atoms in the second period tend to form bonds by
BONDS
“sharing” electrons so that each atom contains eight electrons in its valence shell. When an atom is sur-
rounded by the electon confi guration of an inert case it is in a lower energy state.
Ionic Bonds
Ionic bonds are formed between two or more atoms by the transfer of one or more electrons between
atoms. Electron transfer produces negative ions called anions and positive ions called cations. Ionic
substances exist as crystalline solids. When the solid dissolves, the ions dissociate and can diff use
freely in solution.
Sodium chloride is an example of an ionic solid. A sodium atom, which has 11 protons and
11 electrons, has a single valence electron in its 3s subshell. A chlorine atom, which has 17 protons
and 17 electrons, has seven valence electrons in its third shell, represented as 3s23p5. In forming an
ionic bond, the sodium atom, which is electropositive, loses its valence electron to chlorine. Th e
resulting sodium ion has the same electron confi guration as neon (1s2 2s22p6 ). It has a +1 charge,
because there are 11 protons in the nucleus, but only 10 electrons around the nucleus of the ion. Th e
chlorine atom, which has a high electronegativity, gains an electron and is converted into a chloride
ion that has the same electron confi guration as argon (1s2 2s22p6 3s23p6). Th e chloride ion has a –1
charge because there are 17 protons in the nucleus, but there are 18 electrons around the nucleus of
the ion.
In the crystal structure, each sodium ion is surrounded by six chloride ions and each chlo-
ride ion is surrounded by six sodium ions. Each ion has a complete electron shell that corresponds to
the nearest inert gas; neon for a sodium ion, argon for a chloride ion (Figure 1.4).
Figure 1.4 Sodium Chloride
Crystal
In the ionic solid, sodium chloride,
each sodium ion is surrounded by 6
chloride ions and each chloride ion is
surrounded by 6 sodium ions.
4
Covalent Bonds
Covalent bonds are much more common in organic chemistry than ionic bonds. A covalent bond
consists of the simultaneous attraction of two nuclei for one or more pairs of electrons. The electrons lo-
cated between the two nuclei are bonding electrons. Covalent bonds occur between identical atoms
or between different atoms whose difference in electronegativity is insufficient to allow transfer of
electrons to form ions.
Let’s consider the covalent bond in the hydrogen molecule. A hydrogen molecule forms
from two hydrogen atoms, each with one electron in a 1s orbital. The two hydrogen atoms are at-
tracted to the same pair of electrons in the covalent bond. The bond is represented either as a pair
of “dots” or as a solid line. Each hydrogen atom acquires a helium-like electron configuration. The
atomic nuclei are separated by an equilibrium distance sensibly called the bond length.
H + H H H or H H
Energy is released when the electrons associated with the two hydrogen atoms form a covalent bond.
Bond formation releases heat; therefore, it is exothermic. The heat released when one molecule of a
compound forms at 298K is the standard enthalpy change (DHo) for the process. DHo for forming
a mole of hydrogen from two hydrogen atoms is –435 kJ mole–1. Because energy is released in the
reaction, the hydrogen molecule is more stable than two isolated hydrogen atoms. The reverse pro-
cess, breaking the H—H bond, requires 435 kJ mole–1, a quantity called the bond strength.
Another example of a covalent bond is the Cl—Cl bond in a chlorine molecule. Two chlorine
atoms are attracted to the same pair of electrons. Each chlorine atom has seven valence electrons in
the third energy level and requires one more electron to form an electron core with an argon electron
configuration. Each chlorine atom contributes one electron to the bonded pair shared by the two
atoms. The remaining six valence electrons of each chlorine atom are not involved in bonding. They
are variously called nonbonding electrons, lone pair electrons, or unshared electron pairs.
Cl Cl nonbonding electrons
As we noted earlier, a covalent bond is drawn as a dash in a Lewis structure. In a Lewis structure,
nonbonding electron pairs are shown as “dots.” The Lewis structures of four simple organic com-
pounds: methane, aminomethane, methanol, and chloromethane are shown below with both bond-
ing and nonbonding electrons.
H H H H
H C H H C N H H C O H H C Cl
H H H H H
methane aminomethane methanol chloromethane
The hydrogen atom and the halogen atoms form only one covalent bond to other atoms in stable
neutral compounds. The number of covalent bonds an atom can form is called the valence of the
atom. The valence of a given atom is the same in most stable, neutral organic compounds. Table 1.2
lists the valences of some common elements contained in organic compounds.
Structural Formulas
A molecular formula tells us the composition of a molecule. A structural formula shows the ar-
rangement of atoms and bonds in a molecule. The structural formulas for methane, aminomethane,
methanol, and chloromethane show all of the bonds connecting the constituent atoms. Structural
formulas are often drawn in abbreviated or condensed versions to save time and space. Condensed
structural formulas show only specific bonds; other bonds are implied, but not shown. The degree
of condensation depends on which bonds are shown and which are only implied. For example,
because hydrogen forms only a single bond to carbon, the C—H bond need not be shown in the
condensed structure. Similarly, the two nitrogen–hydrogen bonds in aminomethane and the oxy-
5