Table Of ContentThis is Volume 46 of
ORGANIC CHEMISTRY
A series of monographs
Editor: HARRY H. WASSERMAN
A complete list of the books in this series appears at the end of the volume.
SAN DIEGO
THE PROTON:
APPLICATIONS TO
ORGANIC CHEMISTRY
Ross Stewart
Department of Chemistry
University of British Columbia
Vancouver, British Columbia
Canada
1985
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LIBRARY OF CONGRESS CATALOGING IN PUBLICATION DATA
Stewart, Ross, Date
The proton: applications to organic chemistry.
Includes index.
1. Chemistry, Physical organic. 2. Protons.
I. Title.
QD476.S74 1985 547.1'3 85-757
ISBN 0-12-670370-1 (alk. paper)
PRINTED IN THE UNITED STATES OF AMERICA
85 86 87 88 987654321
Preface
The proton occupies a central place in organic chemistry, which the present mono
graph attempts to describe. It is a quarter of a century since the appearance of
R. R Bell's important work, The Proton in Chemistry, a book that dealt with a
number of aspects of the subject drawn from the chemical sciences as a whole.
The present work differs from Bell's in being concerned specifically with organic
chemistry and, as befits this branch of the subject, being somewhat less rigorous
in its treatment.
An introductory chapter is followed by two chapters that treat in fairly exhaus
tive fashion the strengths of neutral organic acids and neutral organic bases. Sub
seq+uent #chapters deal with the mode of transfer of hydrogen in its three forms
(H , H, and H~), with alternative sites of protonation or deprotonation of or
ganic compounds, with the acid-base chemistry of unstable and metastable
species (including free radicals and excited states), and finally, with the activation
induced in organic molecules by proton addition or removal and the catalytic
effects that ensue.
The book is intended to be of use to practicing organic chemists of whatever
stripe, and I hope that for a number of years it will be the first book to be pulled
from the shelf when a question arises about the acid-base chemistry of some or
ganic compound. Should the answer not be found herein, there is a reasonable
likelihood that the reader will at least find a helpful reference to the chemical lit
erature. In this connection an attempt has been made to include as many recent
literature references as possible.
I am grateful to the following friends and colleagues who reviewed parts of
the manuscript: Drs. C. A. Bunton, R Y. Bruice, R. A. Cox, J.T. Edward,
A.J. Kresge, C. Sandorfy, R. Srinivasan, and K. Yates. Dr. J. Peter Guthrie,
who reviewed the entire manuscript, deserves special thanks for his very helpful
criticisms and suggestions. For anything dubious that might remain I take full
responsibility. I am also grateful to the Izaak Walton Killam Memorial Fund for
Advanced Studies for the award of a Senior Fellowship that enabled me to devote
an uninterrupted year to the completion of the work. Finally, thanks are due to
my wife, whose help in a much earlier work (ref. 18, Chapter 4) was attrib
uted in the preface to "a charming lady," an acknowledgment that was accurate
but considered to be open to misinterpretation. At any rate the same charming
lady helped prepare the manuscript and index of the present work.
vii
1
i
Introduction
The proton is a unique chemical species, being a bare nucleus. As a
consequence it has no independent existence in the condensed state and is
invariably found bound by a pair of electrons to another atom. It is proton
transfer that is at the heart of the concept of acidity, which is the cornerstone
on which much of modern chemical science has been built. Measurements of
the extent of proton transfer to and from such compounds as carboxylic acids,
phenols, sulfonic acids, and amines have provided a vast body of data from
which much of the theory of organic reactivity has emerged. Inductive effects,
linear free-energy relationships, and acid and base catalysis are all subjects
that owe their development to the availability of such data.
Acids and bases have been defined in a number of ways over the years, but
we shall follow herein the Br0nsted definition, in which an acid is a species
that gives up a proton and a base is a substance that accepts a proton (1-4).
Alternative definitions of acidity and basicity have been well described
elsewhere (5-7) and need not concern us, inasmuch as the present work's focus
is explicitly on the proton and the chemistry associated with the transfer of this
unique entity.
I. Water and the Reference State
We live in what has been called an aquocentric environment, and it is only
natural that water should have become the standard medium for the
quantitative determination of acid and base strengths. Thus, infinite dilution
in water at 25°C is generally regarded as the standard state for acid-base
equilibria. Apart from its ubiquity water has a number of other advantages
over most alternative solvent systems. It has a high dielectric constant, thus
reducing the effect of ion pairing on acid-base equilibria, and it can be easily
obtained in a highly pure form. Above all, the teaching of chemical equilibria is
still based to a very great extent on acid-base reactions in water, and one
suspects that there are few practicing chemists who do not regard aqueous
pH 7 as being somehow the point of neutrality, even though their work may
normally involve reactions in organic media or in the solid state or in the gas
phase. Further reasons for aqueous systems having become the focus of
acid-base studies have been elegantly set forth by Arnett and Scorrano (8).
1
2 1. INTRODUCTION
On the other hand, there are some disadvantages to referring all questions
of acid and base strength to the aqueous state. The principal objections that
can be raised to such an approach are the following. First, there is extensive
solvation by water molecules of the components of an acid-base equilibrium,
particularly the ionic components, making the system less satisfactory for
examining fundamental properties than gas phase studies. Second, many
organic compounds whose acidity or basicity is of interest are not particularly
soluble in water. Third, the autoprotolysis constant of water, though fairly
small, is not nearly small enough to encompass the range of acidities that
would enable us to measure directly acid strengths as low as those of alkanes
and as high as those of sulfonic acids. Similarly, the base strengths of very
weak bases such as aldehydes or nitriles are well beyond the reach of direct
measurement in water. [The paucity of very strong neutral bases means that
water is able to accommodate measurements of virtually all of the stronger
neutral bases that do exist; the strongest reported neutral base, apart from
some cryptands that react very slowly, is the diaminonaphthalene 1, whose
pK is estimated to be near 16 (9), just beyond the range of measurements in
BH+
water.]
(C2H5)2N N(C2H5)2
How critical are these three points? They will be addressed in Section III,
but first it is important that the terms and symbols used to express acid and
base strengths be clarified.
II. Acid-Base Equilibria: Terms and Symbols
It has become common practice to use the symbol pK to express the
a
strengths of both acids and bases. In the latter case, of course, the quantity
referred to is actually the strength of the conjugate acid of the base. Thus, the
equilibrium5 constant for the dissociation of anilinium ion in water at 25°C is
2.5 x 10~ [Eq. (1-1)], the negative logarithm of this quantity, 4.60, being
+ + 5
C6H5NH3 H + C6H5NH2 K = 2.5 x 1(T (1-1)
sometimes referred to as the "pK of aniline." However, aniline is also an acid
a
in its own right, although a very weak one in water, and it is the pK for the
II. ACID-BASE EQUILIBRIA 3
equilibrium shown in Eq. (1-2) that would more aptly be so described (70).
+ 28
C6H5NH2 H + C6H5NH~ K « 1(T (1-2)
The strength in water of a base such as aniline can always be related to the
reaction shown in Eq. (1-3), with K and pX being the appropriate symbols
h + b 10
C6H5NH2 + H20 t=> C6H5NH3 + HCT K = 4.0 x HT (1-3)
to use in this case. In aqueous solution at 25°C the pK and pK for weak
a h
bases and their conjugate acids are related by Eq. (1-4). [For moderately
pK + pK = 14.0 (1-4)
a h
strong acids this relationship may not hold rigorously (7i).]
Why, in the case of bases, have K and pK been largely superseded by the
h b
terms K and pK, ambiguous though the latter terms sometimes can be? The
a a
advantage of writing all acid-base reactions in the direction of proton loss
and then using K and pK to describe the equilibria as written is that one has
a a
a unified scale running from the most powerful acids, whether they be neutral
species such as trifluoromethanesulfonic acid or protonated species such as the
conjugate acids of nitroalkanes, through to the weakest acids, the alkanes.
There are two ways of avoiding ambiguity when using the unified approach
whereby all acid-base reactions are written in the direction of proton loss.
One is always to refer to the acid or conjugate acid by its appropriate name or
formula. Th+us, it would always be "the pK of anilinium ion" or "the pK of
C H NH " that is given as 4.60 and never "the pK of aniline" that is so
6 5 3
designated. There is little difficulty in following this stricture with regard to
chemical formulas; problems arise, however, with regard to names for the
conjugate acids of ketones, esters, nitro compounds, and others. One can
always attach the prefix "conjugate acid of" to the name of the neutral
compound, but this practice can weary both the writer and the reader if
followed rigorously.
The second way to avoid ambiguity, and the one that will be followed
herein, is to use the terms K and pK for the acid dissociation of the species
HA HA
in question and the terms K and pK for its conjugate acid. Thus, we
B+H BH+
shall write "the pK of aniline is 28" and the "pK of aniline is 4.60." [The
HA BH+
term "pX (or pK) of anilinium ion" causes no difficulty, but the need for
HA a
such a term will only rarely be encountered herein.] Throughout this work the
terms pK and pK will usually be used, since this practice allows both
HA BH+
weakly acidic and weakly basic organic compounds to be referred to by their
usual names and formulas.
There remains the problem of molecules containing more than one acidic or
basic site, for example, polycarboxylic acids or polyamino compounds. The
4
1. INTRODUCTION
traditional terms pl^, pK , and so on seldom give rise to ambiguity in the case
2
of polyprotic acids since it is usually perfectly clear which processes are being
referred to; doubts are more likely to arise with polybasic compounds, and in
these cases reference to the appropriate chemical equation can be made. When
it is necessary to refer to the basicity of a polycarboxylic acid, the term pK
BH+
(plus the name or formula of the neutral compound) is available, and when the
acidity of a polyamino compound is in question, the term pK (plus name or
HA
formula) can be used, just as for monofunctional compounds.
In the case of unsymmetrical polyprotic molecules, microscopic equilibrium
processes become relevant. In these cases the microscopic dissociation in
question should be unambiguously identified by, for example, a chemical
equation; the unvarnished symbol K (or pK) is really all that is needed if an
equation is used, since the reaction to which the equilibrium constant refers
has been specified.
In summary, pKH Arefers to the acid strength of the species whose name or formula is given
(usually, but not necessarily, a neutral molecule); pKB+H refers to the acid strength of the
conjugate acid of the species whose name or formula is given (again, usually a neutral
molecule). For proton dissociation from polyprotic species, pK1, pK2, and so on will be
used without explanation when there is no possibility of ambiguity; when this possibility
exists or when microscopic dissociations are being considered, the reaction in question will
be suitably identified (e.g., by means of an equation).
With regard to amphiprotic molecules, such as glycine and other simple
amino acids, there are four microscopic acid dissociation processes that must
be considered; they and their associated pK values are shown in Eqs. (1-5) to
(1-8) (12).
+
H3NCH2C02H t=z H+ + H3NCH2CXV = 2.35 (1-5)
H3NCH2C02H H++ H2NCH2C02H pK * 7.6 (1-6)
H3NCH2C02" H+ + H2NCH2C02" pK = 9.78 (1-7)
H2NCH2C02H t=i H + H2NCH2C02" pK * 4.5 (1-8)
The quantities pK and pX , usually used for the experimentally deter
x 2
mined dissociation constants of (protonated) glycine, are identical, within
experimental error, to the pK values shown with Eqs. (1-5) and (1-7). The pX
x
value is actually a composite constant for reactions (1-5) and (1-6), with
reaction (1-5) being the overwhelmingly dominant quantity, as can be seen
from the size of the respective dissociation constants. [The estimate of 7.6 for
the pK of reaction (1-6) comes from the value of pK + of glycine ethyl ester.]
BH
The pK value of glycine includes the reactions shown in Eqs. (1-7) and (1-8),
2
III. LIMITATIONS ON WATER 5
but because almost all glycine of zero net charge is in the dipolar form, the pK
for the dissociation of the ammonio group [Eq. (1-7)] is virtually identical
with the experimental pK . A complete analysis of the relationships between
2
microscopic and macroscopic dissociation constants is given in Chapter 5.
III. Limitations on Water as a Medium
for Acid-Base Measurements
It was pointed out in Section I that there are three disadvantages to using
water as the standard state for acidity measurements. First, rather strong
solvation effects will generally be present in aqueous solution; second, many
compounds of interest have only slight solubility in water; and third, the size
of the autoprotolysis constant of water sets a limit on the range of pX's that
can be measured directly in this medium.
With respect to the first point, it does indeed seem reasonable to regard the
fundamental properties of a compound as those of the molecule in isolation,
that is, in the infinitely dilute gas phase, since molecular interactions in the
liquid and solid states modify the intrinsic properties that molecules in
isolation possess. Whether such a state is the most relevant starting point for
discussing chemical reactivity is another matter altogether. Most chemistry
is conducted in the condensed phase, where molecules are affected by the
presence of neighboring matter, even when the molecules comprising such
matter are completely unreactive and nonpolar.
Thus, a condensed medium, by virtue of its polarizability, will provide a
substantial degree of interaction with solute molecules, even though the
medium might have an extremely low dielectric constant. The difference in this
respect between a medium such as liquid hexane, with a dielectric constant of
1.9, and a vacuum, with its dielectric constant of 1.0, is very much greater than
a simple comparison of the two numbers would suggest.
Partly for these reasons the condensed state (and in particular the aqueous
state) will be used in this work as the reference state for questions of acid and
base strength. That is not to say that the important work done since the mid-
1970s on proton transfer in the gas phase by Kebarle, Mclver, Taft, Boehme,
DePuy, McDonald, and others (13-18) is irrelevant to our concerns; far from
it. It tells us a great deal about the fundamental properties of acids and bases,
and reference to such work will be made where it seems appropriate to do so.
With regard to the limited solubility in water of many organic compounds,
it turns out that this factor is not of particular consequence as far as extremely
weak acids and bases are concerned. Most determinations of the strengths of
such compounds in water require the use of mixed solvents in any case, which
generally dissolve organic compounds to a greater extent than does water.
6 1. INTRODUCTION
This solubilizing effect is observed not only with aqueous dimethyl sulfoxide
(DMSO) and like systems that are often used in pK measurements of very
HA
weak acids, but also with aqueous sulfuric acid and like systems that are often
used in pX measurements of very weak bases. For example, nitrobenzene
B+H
is ~ 10 times as soluble in 80 wt % sulfuric acid as it is in water (79), with the
increase not being caused by protonation of the substrate, which is in fact
negligible at this acidity. In those other cases where substrate solubility is
insufficient, water can be replaced by ethanol (20-24), dioxane (25), or other
solvent (26) and corrections made in order to refer the pK values so
BH+
obtained back to the standard state of water (26).
There remains the matter of the magnitude of the autoprotolysis constant
of water. Dimethyl sulfoxide has a much smaller degree of autoprotolysis than
does water [Eqs. (1-9) and (1-10)] (27-29), and it has been pointed out by
Bordwell (30) and others that DMSO as a reference state allows direct
measurements to be made on nitrogen and carbon acids whose degree of
ionization is far too small to be directly measured in even the most alkaline
aqueous solutions. Likewise, DMSO solutions containing strong acids are
highly acidic and can protonate weak organic bases.
+
2H20 ^=±+ H30 +HO~ 14 (1-9)
*.„« , = [H30 ][HO-] = l<r +
O OH O"
II II I
2CH3SCH3 H3C—S—CH3 + H3C—S=CH2 (1-10)
O"
I
H3C—S=CH2
Despite this advantage the DMSO standard state and its attendant pK scale
are unlikely to supplant the long-established aqueous scale, and the most that
might be expected would be an indefinite period of confusion if attempts were
made to supplant aqueous pK values with those based on DMSO medium,
even though in the case of very weak acids the latter would probably be the
more firmly established of the two. [Particular care, however, is required in
making measurements in DMSO because of the great effect that small
amounts of water and other likely impurities have on the acidity or basicity of
the medium (37).]
A distinctive feature of pK measurements in aqueous solution is the
HA
absence of ion-pairing effects, except at very high solute concentrations.
Although solvation is a powerful influence, it is at least comparatively uniform
from compound to compound. The situation is much more difficult in most
other solvents, where ion pairing can be a serious problem; even in DMSO, ion
pairing is significant in all but the most dilute solutions.
