Proximity Effects. Kinetics, Mechanisms and Reactivity ...
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8-1977
Proximity Effects. Kinetics, Mechanisms and Reactivity Proximity Effects. Kinetics, Mechanisms and Reactivity
Correlations for the Acidic and Alkaline Hydrolysis of Ortho-Correlations for the Acidic and Alkaline Hydrolysis of Ortho-
Substituted-N-Methylbenzohydroxamic Acids Substituted-N-Methylbenzohydroxamic Acids
Irl E. Ward Western Michigan University
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Recommended Citation Recommended Citation Ward, Irl E., "Proximity Effects. Kinetics, Mechanisms and Reactivity Correlations for the Acidic and Alkaline Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids" (1977). Dissertations. 2789. https://scholarworks.wmich.edu/dissertations/2789
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PROXIMITY EFFECTS. KINETICS, MECHANISMS AND REACTIVITY CORRELATIONS FOR THE ACIDIC AND ALKALINE HYDROLYSIS OF
o r t h q - s u b s t i t u t e d -n -m e t h y l b e n z o h y d r o x a m i c ACIDS
by
Irl E. Ward
A Dissertation Submitted to the
Faculty of the Graduate College in partial fulfillment
of theDegree of Doctor of Philosophy
Western Michigan University Kalamazoo, Michigan
August 1977
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PROXIMITY EFFECTS. KINETICS, MECHANISMS AND REACTIVITY
CORRELATIONS FOR THE ACIDIC AND ALKALINE HYDROLYSIS OF
ORTHO-SUBSTITUTED-N-METHYLBENZOHYDROXAMIC ACIDS
Irl E. Ward, Ph.D.
Western Michigan University, 1977
The kinetics, mechanisms and correlations of observed rate data
by the Taft-Pavelich equation were studied for the acidic and alkaline
hydrolysis of ortho-substituted-N-methylbenzohydroxamic acids. The
results of the mechanism study are interpreted in terms of a bi-
molecular mechanism for acidic catalysis and as reaction of the hydrox-
amic acid conjugate base with water or hydroxide ion for basic catalysis
in the catalytic range investigated. Ionic strength effects and
specific ion effects are also reported for both catalytic systems.
Correlation of the log of the observed rate constants (i.e., log
with the Taft substituent parameters a * and Eg was made for various
ortho-substituents for both catalytic systems. Correlation of log
for acidic hydrolysis was shown to be very good (R = 0.989). The
F-test showed the correlation to be significant at the 1% level. For
the alkaline hydrolysis, correlation of log kQbg with a* and Eg values
was shown to be good (R = 0.928). The F-test showed this correlation
to be significant nearly within the 5% level. For both hydrolysis
systems, correlation with o* and Eg values together was always better
than with o* and Eg values alone. These results lend support to the
semi-empirical description of the "Ortho-Effect" proposed by Taft and
Pavelich as a first approximation to a quantitative approach, which
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is of the general form:
log k = o * p* + 5Eg + log kQ
The results support the descriptions of the steric effect of an ortho
substituent by Taft and by McCoy and Riecke. Taft states that Esvalues are good measures of the actual steric effect of an ortho
substituent, although for those substituents which exhibit direct
resonance interaction with the reaction site, there is a resonance
contribution to Eg . McCoy and Riecke separate Eg into independent
contributions from primary and secondary steric and resonance effects
only. The results of this work also support the qualitative conclusions
of McCoy and Riecke that the susceptibility of a reaction system to
steric effects by ortho-substituents varies with the structural
skeleton of the system.
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ACKNOWLEDGEMENTS
I would like to thank Dr. Don Berndt for his valuable suggestions
and assistance in the preparation of this work. I am also grateful to
the chemistry department and graduate college of W.M.U. for the Graduate
College Associateship appointment which allowed me to devote full time
to the project and for my earlier appointment to a graduate teaching
assistantship. My special thanks go to my wife, Sue, for her infinite
patience and great help in the writing of this dissertation.
Irl Eugene Ward
ii
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WARD, Irl E., Jr., 1949-PROXIMITY EFFECTS. KINETICS, MECHANISMS, AND REACTIVITY CORRELATIONS FOR THE ACIDIC AND ALKALINE HYDROLYSIS OF fiRTHfl-SUB- STITUTED-N.-METHYLBENZOHYDROXAMIC ACIDS.Western Michigan University,Ph.D., 1977 Chemistry, organic
Xerox University Microfilms, Ann Arbor, Michigan 48io6
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TABLE OF CONTENTS
CHAPTER PAGE
I INTRODUCTION............................................... 1
Substituent Constants and the Ortho-Effect.......... I
Mechanism Background.................................... 20
Purpose.............................. 32
II EXPERIMENTAL METHOD, APPARATUS AND SYNTHESES.......... 35
Preparation of Ortho-Substituted Benzoyl Chlorides.. 35
Preparation of Ortho-Substituted-N-Methylbenzohy-droxamic Acids........................................ 36
Preparation of Standard Hydrolysis Solvents......... 39
Preparation of Standard Ferric Chloride Solutions... 40
Preparation of Reaction Solutions and KineticsProcedure.............................................. 41
The Constant Temperature Oil Bath..................... 48
Determination of Rate Constants....................... 48
Reaction Product Analysis of Selected Hydroxamic 'Acids in Alkaline Solution.......................... 50
III RESULTS AND DISCUSSION.................................... 53
Acidic Hydrolysis Mechanism................ 53
Alkaline Hydrolysis Mechanism .................. 57
Proximity Effects for Acidic Hydrolysis............ 62
Proximity Effects for Alkaline Hydrolysis ..... 73
iy BIBLIOGRAPHY................................................ 80V VITA ........................... 83
iii
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LIST OF TABLES
TABLE PAGE
I Yields of Prepared Ortho-Substituted BenzoylChlorides............................................. 36
II Yields of Prepared Ortho-Substituted-N-Methyl-benzohydroxamic Acids............................... 38
III Elemental Analysis of Prepared Ortho-Substituted-N-Methylbenzohydroxamic Acids..................... 39
IV Rate Constants for Base Catalyzed Hydrolysis ofQ2-Chloro-N-Methylbenzohydroxamic Acid at 90.0 C 43
V Rate Constants for Acid Catalyzed Hydrolysis of2-Methyl-N-Methylbenzohydroxamic Acid........... 44
VI Rate Constants for Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C............................................. 45
VII Rate Constants for Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids in 7.31M NaOH at 90.0°C............................................. 46
VIII Rate Constants for Catalyzed and UncatalyzedHydrolysis of Ortho-Substituted-N-Methylbenzo- hydroxamic Acids at 90.0°C in the Presence of Salts.............................................. 47
IX Activation Parameters for Acidic Hydrolysis 55
X Observed and Calculated Rate Constants for AcidicHydrolysis of Ortho-Substituted-N-Methylbenzo- hydroxamic Acids in 0.764M HC1 at 90.0 C 63
XI Comparison of Observed Rate Constants with RateConstants Calculated from Equations (10) and (11) for the Acidic Hydrolysis of Ortho-Substituted-N- Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C............................................. 65
XII Comparison of Op and Op Values for Para-Substi-tuents of Benzene Derivatives in Aqueous Media 71
XIII Observed and Calculated Rate Constants for AlkalineHydrolysis of Ortho-Substituted-N-Methylbenzo- hydroxamic Acids in 7.31M NaOH at 90.0 C 74
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LIST OF FIGURES
FIGURE PAGE
1 Experimental Apparatus ...................... 49
2 Dependency of k0us (sec_1) on Catalytic AcidConcentration (S)................................. 54
3 Dependency of k , (sec"1) on Catalytic BaseConcentration0^ ) ................................. 58
4 Correlation of log with a * and Es Valuesfor Acidic Hydrolysis............................. 64
5 Correlation of log k bg with o* and Eg Valuesfor Alkaline Hydro?ysis........................... 75
v
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INTRODUCTION
Substituent Constants and the Ortho-Effect
Organic chemists have long been interested in discovering
quantitative methods of correlating structural changes with reactivity
and mechanisms. Many empirical relations between reactivities of
organic compounds have been shown to be linear functions involving
logarithms of rate or equilibrium constants. The linear relationship
between two similar reaction systems involving reactivity changes due
to identical changes in structure was initially manifested for side
chain reactions of benzene derivatives in the Hammett equation.^ This
equation, formulated by L.P. Hammett in the 1930’s, was originally
defined for the ionizations of meta- and para-substituted benzoic acids
in an attempt to empirically describe substituent effects on side chain
reactions of benzene derivatives. The empirical relation is given as:
log (k/kQ) = op (1)or
log (K/Kq) = ap (2)
where k or K represent the rate constant or equilibrium constant,
respectively, for the meta- or para-substituted benzene derivative,
and kQ or Kq represent the rate constant or equilibrium constants,
respectively, for the parent unsubstituted benzene derivative. The
polar constant, a, represents the polar effect of the substituent on
the reaction relative to hydrogen and is, by nature, independent of
the reaction type. The reaction constant, p, measures the suscepti
bility of the reaction to polar effects and is, by nature, a constant
1
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for all substituents and depends only on the reaction series, temperature
and solvent.2,3The Hammett equation has been extensively discussed and found to
be applicable to many meta- and para-substituted benzene systems which
are quite different from the defining systems. The equation failed,
however, for para-substituents which, either by electron withdrawal or
electron release, exhibit a direct resonance interaction with the3functional group undergoing reaction, or when substituent proximity, as
for ortho-substituents, introduced steric effects not accounted for by2,3Hammett's polar substituent constant, a.
The failure of the Hammett equation for ortho-substituents was A 5studied by Kindler. ’ He first observed a relationship between the
rates of base catalyzed hydrolysis of ethyl meta- and para-substituted
cinnamates and the corresponding rates of ethyl meta- and para-benzoates.
The failure of ortho-substituents to obey the relation was attributed to
a steric "ortho effect" for the benzoate system.
Later, Ingold devised a general method for the separation of polariand resonance effects from steric effects in ester hydrolysis.
According to this method, the ratio of the rate constants of alkaline
to acidic hydrolysis is a function of the polarity of the substituent,
even though both reactions show steric effects.
Following this line, Taft proposed an equation to quantitatively
describe the polar substituent effect for a substituent R in the
hydrolysis of aliphatic esters of the form: G-COOR', and ortho
substituted benzoates of the form: ^0^-(jj-OCH^• The polar substituent
constant, a*, was given as:^’^
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3
a* = [log (k/k0)B - log (k/ko)A J / 2 . 4 8 (3)
k represents the rate constant for the hydrolysis of the substituted
ester, G-COOR* o r COOCHg, while kQ is the rate constant for the
hydrolysis of the parent ester, CHo-COOR' or(oV COOCH-. The subscriptsCH3B and A refer to the base or acid catalyzed hydrolysis. The factor
2.48 is a constant introduced so that the values of a * will be put on
approximately the same scale as the Hammett a values. Such a constant
was obtained from ratios of o * values for selected substituents with
the purely inductive o' values for the same substituents derived by
Roberts and Moreland^ from the ionization constants, analogous to those
in equation (2), of saturated 4-substituted-l-carboxy~C2*2.2J bicyclo-
octanes which were geometrically similar to and had similar reactivities
as meta- and para-substituted benzoic acids. Values for a 1 were, there
fore, on the same scale as Hammett a values and represented the purely
inductive effects for a substituent which obeys the Hammett equation.
The terms on the right side of equation (3) have the following signifi
cance: log (k/kQ)B represents the sum of polar, resonance and steric
effects of G; log (k/kQ)A represents the sum of steric and resonance
effects of G, the difference giving the purely polar effect of G (see
assumption 2 below).
The validity of equation (3) as a measure of the polar effect of2,4a substituent is based on three assumptions:
1. The relative free energy of activation can be treated as a sum of independent contributions from polar, resonance and steric effects.
2. In corresponding acid and base catalyzed hydrolyses, the steric and resonance effects are cancelled in the difference:
log (k/kQ)B - log (k/kQ)A
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3. The polar effects of substituents are much greater in base than in acid catalyzed ester hydrolyses.
If assumption (1) were invalid, Taft’s entire argument would be
voided, for it is on this principle that separation analysis is feasible.
Support for this assumption is provided by the usefulness of results
obtained in applying a * values to reactivities. Since Taft’s definition
of a * arises from acidic and basic ester hydrolyses which are known to
occur through tetrahedral intermediates, justification of assumption (2)
is obtained from the similarity of these intermediates. It was proposed
that since these tetrahedral intermediates are saturated and differ in
size by only two protons, that steric and resonance effects on the rate
constants must be similar. Support for assumption (3) is based upon
hydrolysis rate data obtained for meta- and para-substituted benzoates.
Recorded values of p (meta and para), which lie between -0.2 and +0.5
for the acid catalyzed hydrolysis, are, in most cases, nearly z e r o J
This is in contrast to recorded p (meta and para) values for the base
catalyzed hydrolysis which commonly lie within the range of +2.2 to 2.8.
Hammett studies have shown that, for systems which obey equations
(1) or (2), p ̂ p and that differences in reactivity between * ’meta ^parameta- and para-substituted systems arise from differences in aHanunet;t;
(meta) and pt(, (para) values. Taft proposed that since p (meta)£J
p (para) for benzene deriviatives, it is reasonable that p (ortho)2£2
p (para). Therefore, this premise and assumption (3) (i.e., o)
allowed Taft to define the purely steric effect of a substituent, Eg , as:
log (k/k ) = 6E (4)o A s
6 represents the reaction’s susceptibility to steric effects and is
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3. The polar effects of substituents are much greater in base than in acid catalyzed ester hydrolyses.
If assumption (1) were invalid, Taft's entire argument would be
voided, for it is on this principle that separation analysis is feasible.
Support for this assumption is provided by the usefulness of results
obtained in applying a * values to reactivities. Since Taft's definition
of a * arises from acidic and basic ester hydrolyses which are known to
occur through tetrahedral intermediates, justification of assumption (2)
is obtained from the similarity of these intermediates. It was proposed
saturated and differ in
Hammett studies have shown that, for systems which obey equations
meta- and para-substituted systems arise from differences in aHammett
(meta) and (para) values. Taft proposed that since p (meta)£S
p (para) for benzene deriviatives, it is reasonable that p (ortho)^
p (para). Therefore, this premise and assumption (3) (i.e., p o)
allowed Taft to define the purely steric effect of a substituent, Eg , as:
6 represents the reaction's susceptibility to steric effects and is
constants must
hydrolysis rat<
Recorded value.-
for the acid can
This is in contrast"
catalyzed hydrolysis whici^^ u.y lie within the range of +2.2 to 2.8.
Wnd para) va'lues for the base
Resonance effects on the rate
V i e between -0.2 and +0.5
rrost cases, nearly zero.^
Option (3) is based upon
B.ra-substituted benzoates.
(1) or (2), p}meta^ ppara and that differences in reactivity between
log (k/kQ)A = 6Eg (4)
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defined as 5 = 1.00 for the acidic hydrolysis of^O^-COOCH^ and G-COOCH^.
The polar effect of -G, defined by p* a * for the ortho-substituted
benzoate system and by p* a * for the substituted acetate hydrolysis, are
necessarily zero for the acid catalyzed reactions (i.e., p* represents
the susceptibility of the ortho-substituted benzoate system to polar
effects and is, according to Taft's above premise and the support for
assumption (3), nearly zero, a * is as defined in equation (3)).
Equation (4) implies, for reasons analogous to those for a * values
discussed below, that Eg values are defined on the basis of two reaction
types; aliphatic systems (i.e., Eg and ortho-aromatic systems (i.e.,
E g°) in which for both systems, as with a * values, Eg (-CH^) = 0.
Although defined to be a measure of the steric effect of a substituent,
for both aliphatic and aromatic systems in which -G contains a n-system
conjugated with the ester function, there is a resonance contributionA
to E . Equation (4) has been found to correlate data for several 4 8 9reaction systems. * ’ This provides experimental evidence for the
validity of the above premise (i.e., the equality of p values) as applied4 8 9to ortho-substituted aromatic systems as well as aliphatic systems.
The equation for the polar substituent constant, o*, was defined
for two reaction systems, since each system type results in a different
value of a* for the same substituent. The difference between a *
(aliphatic) and a* (ortho) stems from differences in the defining systems.
These differences arise from a difference in the geometric and electronic
environments between the substituent on the aliphatic ester, of the form
G-C^-COOR', and the ortho-substituent on the corresponding benzoate, of
the form (o^-COOCHg, which is manifested in a difference in resonance
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6
interactions, field effects and conformational possibilities between
the two systems.
The difference between o * , as defined by Taft, and a, as defined
by Hammett, stems from their different origins. These differences are
three-fold:
1. While Hammett a values are defined relative to a hydrogen standard, o * values are defined relative to a -CHg standard for benzene derivatives.
2. While Hammett a values were defined from the uni- molecular ionization of benzoic acids, o * values were defined from a bimolecular hydrolysis reaction, the mechanism of which involved a tetrahedral intermediate.
3. While a* values are defined for substituents in close proximity to the reaction center, a values are calculated for the substituents when no such steric effects are possible.
Taft concluded that, except for groups that are unsaturated and
in conjugation with the ester function, or for groups which give rise
to changes in attractive interactions from the reactant state to the
transition state (i.e., changes in hydrogen bonding, field effects,
etc.):
log (k/k0)A i Es (5)
is a near quantitative description of the total steric effect of a
substituent relative to -CH^ for both G-COOR' and ^O^-COOCH^ acidic
hydrolysis. Studies showed for these types of substituents that
equation (4) was applicable to many aliphatic and some ortho-substituted4
aromatic systems other than those used in defining Eg values. Taft
also found that for aliphatic systems, values of Eg for both symmetrical
and unsymmetrical substituents paralleled their van der Waals radii as
determined by Pauling. For ortho-substituted aromatic systems, however,
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7
only those "symmetrical-top" substituents (i.e., CH-j, t-Bu, CX^, X)
were found to have Eg values which roughly paralleled their van der
Waals radii. For the unsymmetrical substituents, (i.e., -CI^X,
-CH2R, etc.), Eg values did not parallel their van der Waals radii,
but did conform to the qualitative idea that steric effects, as
described by E values, do increase with the overall bulk of the J s
substituent.
M. Charton investigated the validity of Taft's conclusion1^ ’11
that, except for substituents that are unsaturated and in conjugation
with the reaction function, or for groups which give rise to changes
in attractive interactions from the reactant state to the transition
state, that equation (5) is a quantitative measure of the steric effect
of a substituent.
Charton studied the acid catalyzed hydrolysis of aliphatic esters
of the form G-CH2-COOR. He found a quantitative linear relation
between Taft's E values and van der Waals radii, as calculated from s12works of Bondi, which was given as:
E = ip /— + h (6)s,x r ' v,x
i|i represents the susceptibility of the reaction system to steric effects
and is analogous to Taft's 6 constant in equation (4), x represents12the substituent's van der Waals radii, h represents a discrepancy
constant of undescribed composition which Charton employed when the
use of hydrogen as the standard substituent, rather than methyl,
decreased the validity of the correlation. Charton concluded that for
acidic hydrolysis of aliphatic esters and related aliphatic reaction
systems,1"* Eg , as defined by Taft in equation (5), is a true measure
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of a substituents' steric effect. However, Taft also used equation (5)
to define E as a steric substituent constant from ortho-substituted sbenzoate hydrolysis as well. Charton disagreed with this definition
and represented the Eg (ortho) values in terms of a combination of11,14contributions from inductive, resonance and steric effects given as:
Es <ortho) E Es,x 5 “°I,X + S“r ,X + * ^ , x + h <7)
« and g represent reaction susceptibility constants to inductive and
resonance effects, respectively, ip and h are similar to those described
in equation (6). ^ is the inductive effect of the substituent for
which values were compiled from the new set of a^ values determined by
Charton from the ionization constants of substituted acetic acids.^
O - (Charton) values are analogous to o' (Roberts and Moreland).^ a„1 k
represents the resonance effect of the substituent for which values were
obtained from the equation:
°R " °P - °I (8)
Op is the para-substituent constant for which values were taken from
the compilation of McDaniel and Brown.^ values are obtained from
Charton's compilation.
Charton noted, however, that since values of "h", resulting from
hi's correlations with a large number of reaction series, varied with
the reaction system, there could be no single value of E° ^ for a
s u b s t i t u e n t . ^ He also noted, for ortho-substituted benzoic acid
ionizations, benzoate hydrolyses, and other similar reactions, that
i p & O which, he concluded, suggests that E° ^ is primarily an electrical
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9
Charton also studied the validity of Taft's contention that, for
both the aliphatic and ortho-aromatic reaction systems to which a * has
been applied, a * (aliphatic) and a * (ortho-aromatic) is a measure of
the inductive effect of a substituent. Charton noted that in his
derivation of o * values, Taft makes the same assumption he did in his
derivation of E° x values; namely, that P|aft (ortho)^ pHammett (Para) *
Charton disagreed with this assumption and reported that, in general,
°Taft i.15.1
effects operate equally from the para- or ortho- positions. He concluded
that a * is not a purely inductive effect, and he represented the ortho-
polar substituent effect, derived from different reaction systems than
those used for equation (7), in general a s : ^
(9)
^ was derived using hydrogen as the standard rather than ortho-CH^
for ortho-substituted aromatic systems in which steric effects were
minimized. A and 6' represent constants defining the relative importance
of inductive and resonance contributions, respectively, to CT0 x > h
represents a constant analogous to that in equation (7). Charton
calculated values for 5 '/A using the method of multiple linear regression
analysis for 19 reaction series involving mostly ortho-substituted acid
ionizations.1^ He found values of this ratio ranged from 0.2 to 1.4
one reaction system to another. Charton, therefore, concluded that, as
with values for E° , there was no single value of a for a substituent s,x o,xapplicable to a variety of reaction systems.
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10
On the basis of work done by Kindler, Taft, Charton and others,
the total effect of a substituent in the ortho- position on the rate
constant or equilibrium constant of a reaction, referred to in general
as the "Ortho-Effect", has been described in terms of polar, steric,
and a combination of polar and steric parameters. As discussed above,
the separation of polar, steric and resonance effects of an ortho
substituent has led various authors to different descriptions of
substituent constants, and of the "Ortho-Effect".
Taft found for some reaction systems (i.e., acid hydrolyses of4
ortho-substituted benzamides, ortho-substituted benzoates, etc.) that
the "Ortho-Effect" could be suitably described in terms of a single
steric effect given as:
log k = 6E + log k (10)x s o
For some other systems (i.e., ionizations of ortho-substituted4
anilinium ions, ortho-substituted benzoic acids, etc.) Taft found that
the equation:
log kx = a*p* + log kQ (11)
suitably described the "Ortho-Effect". However, for most ortho
substituted reaction systems, Taft described the "Ortho-Effect" as a
combination of polar and steric effects as:
log kx = a*p* + 6Eg + log kQ (12)
This equation has found wide applicability although it has failed for
some ortho-substituted systems.
The description of the "Ortho-Effect" is somewhat dependent upon
whether or not the author assumes p * # p TT .. (para) for the describedo Hammett
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11
system. As did Taft, Kindler and Newman found for many systems that
p* does approximate pHamn|ett (para) and that 0o « 0Hainmett (para) (i.e.,
oq represents the ortho-substituent constant as defined by Taft in
equation (3) using hydrogen rather than ortho-methyl- as the standard).
They used these relations to describe the "Ortho-Effect" as:
“O.K. <13)nstant
substituted reaction respectively and
in the effect of an ortho- vs. para-substituent.
Charton, as discussed, did not agree with Taft's assumption that
p * « p „ (para) and, using his own description for a from equationo Hammett o,x(9), described the "Ortho-Effect", for reaction systems which are
insulated from steric effects by separation of the reaction site and the
ortho-substituent using a -CH2 or -0CH2 group between the ring and
functional group, or ring and substituent, a s : ^ ’^ ’^
(14)
Qx represents the quantity measured as log kx or log Ka< a ' and 3'
represent the importance of the inductive and resonance contributions,
aT v and o_ v , to Q , respectively (i.e., «' = p*X and 3' = P*5')*1, A K, A X O OFor some aliphatic, but mostly for aromatic systems in which proximity
effects are expected to be large, Charton described the general "Ortho-
Effect" as a combination of a and E° values from equations (7) and o,x s,x(9) and represented it as:.14,20,21
Qx ' * ' 0I , X + S '0R,X + 'l''7,x + h (l5)
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12
«' and 3 ’ have meanings analogous to those in equation (14). h also is
analogous to that in equation (6). From his determination that \p 0
in equation (7), Charton concluded that the general "Ortho-Effect" is
primarily electrical in nature and can be finally represented as:
Qx = ‘,0I,X + b'°r,x + h ’ (16)where h' includes any steric effects present which, since tj; 0, are
either constant or negligible. Charton claimed that this equation is
applicable to mostly benzoic acid ionizations, acid and base catalyzed
aromatic ester hydrolysis, ionizations of other aromatic acids, phenols,
anilinium ions, etc.
It is significant that Charton's equation is claimed to apply to
the acid hydrolysis of ortho-substituted benzoates, for it is this
reaction which Taft used to define his steric substituent constant given
in equation (5) as:
log (k/kQ)A = 6Eg
where 6 = 1.00 by definition for this system. This definition was based
on the assumptions that: (1) polar effects for such a reaction are very
small; and (2) p *c ? p „ (para). E is, according to Taft, aKo ~ KHammett vr sdescription of the purely steric effect of a substituent relative to the
methyl standard for which Eg (ortho-CH^) = 0. Charton argues, though,
that since his studies have shown that Taft's above assumption (2) is
incorrect in general, this invalidates Taft's definition of Eg as a
purely steric effect. Therefore, that the "Ortho-Effect" description
in equation (16) is applicable to acid catalyzed ortho-substituted
benzoate hydrolysis is not, claims Charton, inconsistant with Taft.
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13
Further, Charton concludes that since values of S'/X vary with the
reaction system, values of «' and 3 ’ in equation (16) are character
istic of the system studied.
L. McCoy and E. Riecke have also studied ionizations of ortho- 22substituted benzoic and related acids and interpret the results in
terms of an "Ortho-Effect" which contradicts Charton's general
"Electrical Ortho-Effect" as described in equation (16). McCoy and
Riecke point out that in Charton's expanded Hammett-like equation:
“ '<’l,X + S'°R,X+ * ^ , x + h (17)
which, when applied to ortho-substituted benzene systems, takes the
form of equation (16), the standard reaction system employed hydrogen
rather than ortho-methyl. As a result, in all of his correlations,
Charton excluded the parent compound. His exclusion was based on the
premise that the unsubstituted compound "did not represent a typical
member of the ortho-substituted se t " . ^ In nearly one-half of his
correlation sets, Charton found o Q R = 0 was not a satisfactory standard,
and the introduction of the constant h, as a discrepancy constant
described earlier, varied from one reaction system to another. Charton
considered this as evidence for the non-existence of a single value of
o q x for any substituent. McCoy and Riecke, however, dispute Charton's
logic on the basis of Charton's inability to define the composition of
"h". They believe it is inconsistent for Charton to separate polar
effects into inductive and resonance contributions, as in equation (9),
and then to imply that steric effects are either constant or negligible
and attempt to represent them by a single parameter, h', as in
equation (16).
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14
McCoy and Riecke also extensively criticized Charton's represen
tation of E° x in equation (7) as primarily an electrical effect. Their
criticism is based on two points. First, Charton, as in his equation
for a , does not define the composition of h. Second, he also attempts0.xto represent the entire steric effect in terms of one parameter,
which is incorporated in h' in equation (16). McCoy and Riecke state
that the steric factor of an ortho-substituent, which Charton attempted
to describe, represents "space-filling" interactions. This implies that
an ortho-substituent will have no steric effect until it is large enough
in volume. They interpret the steric effect of a substituent in terms
of contributions from two factors:
1. Primary Effect: This effect represents direct spatial interactions between the ortho-substituent and the approaching solvent or nucleophile.
2. Secondary Effect: This effect represents steric hindrance by the ortho-substituent to resonance in the reactant state due to bond twisting and bending of conjugated unsaturated groups.
McCoy and Riecke state that such effects are applicable for unimolecular22acid ionizations and bimolecular substitutions.
Steric effects in these systems are represented by a graph,
according to McCoy and Riecke, relating the total steric effect of22ortho-G with increasing substituent volume:
•rla)4J
cdoH
Increasing Size of Substituent
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The graph has been interpreted by McCoy and Riecke in the following
manner: Point A represents the size, or volume, of the substituent
(i.e., -G) which is just large enough to enter the space of the solvent
shell at the reaction site, and over some size increase in ortho-G,
from A to B, it increasingly excludes some solvent or attacking reagent
molecules (i.e., the primary effect). This increase in solvent or
attacking reagent exclusion as the ortho-substituent increases in size
will depend on the shapes of the solvent or attacking reagent molecules,
the substituent, and the mode of attack at the reaction site. "Although
preventing the solvent (or attacking reagent) molecules from occupying
this volume, substituents at minimal size B will not themselves occupy 22the total volume excluded. So, from B to C, there is little increase
in the total steric effect since ortho-G simply occupies more of the
space already excluded, "or because steric hindrance to solvation,
region A to B, and steric inhibition of resonance, region C to D,
overlap and operate in opposite directions" (see below). At size C,
the ortho-substituent is in direct contact with the reaction site.
Between C and D, it is expected that either ortho-G, the reaction site,
or both would be increasingly bent, twisted or distorted, due to direct
electron cloud interactions, in a fashion which minimizes the volume-
filling interactions. Such distortion of the carbonyl reaction site
has the effect of lessening the extent of coplanar conjugation between
the ring and the carbonyl group in the reactant state (i.e., the
secondary steric effect). At size D, maximum bond twisting of 90° has
been reached at which point resonance of the carbonyl reaction site
with the ring is lost. Between D and E, the degree of interaction
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16
again would change, "but probably to a lesser rate of change with 22increasing size". Interactions within this region may involve a
conformational factor, "but the limits and degree of such interactions
of direct contact may not coincide with those for the steric inhibition
to resonance",^ (i.e., the secondary effect), proposed in the region
between C and D.
This description of the steric effect by McCoy and Riecke as a
combination of two factors depending on substituent size is in direct
contrast to Charton's depiction of the steric effect in equations (7),
(9) and (15). McCoy and Riecke conclude, therefore, that it is
inadequate to describe, as Charton does, both the primary and secondary
‘steric interactions of an ortho-substituent in terms of a single steric
parameter Further, a single reaction parameter, ip, to describe
the susceptibility or importance of these two steric interactions is
also inadequate.
McCoy and Riecke substantiate the correctness of their interpre
tations by pointing out the consistency of their results with observations
by both Taft and Charton that Eg ^ values for substituents in aliphatic
systems are directly proportional to their van der Waals radii.
They point out that for aliphatic systems where there is no resonance in
the reactant state, there is no "conformational factor" and no steric
hindrance to resonance by the substituent. The steric effect of the
substituent will, therefore, be primary and dependent only upon the
substituent size as related to van der Waals radii as observation has
indicated.
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17
McCoy and Riecke note that even though Charton's descriptions of
the steric effect of an ortho-substituent in equations (7) and (15) are
inadequate, Charton's data provides two pieces of valuable information:
1. Charton's constant "h" is probably the sum of two oppositely acting steric effects.
2. Charton's isolation of his "h" values, which aretypically of the same magnitude as other contributionsto Q , a or E , leads to a misleading correlation x o,x s,xbetween E ^ , oIjX and oR>x.
Justification for these statements comes from McCoy and Riecke's studies
of the ionizations of ortho-substituted benzoic acids in 50/50 (w/w)22methanol: water, and application of Charton s data to their graphical
interpretation of an ortho-substituent's steric effect. They conclude
from their data on the ionization of benzoic acids that steric hindrance
to solvation by an ortho-substituent (i.e., primary effect) decreases
acidity. However, steric hindrance to resonance of the carbonyl group
with the ring caused by bond angle twisting (i.e., secondary effect)
increases acidity. This results from the loss of a greater amount of
resonance stabilization in the reactant state than in the product state.
Assuming the same effects are present in esters, they claim, the net
steric effect for a certain class of substituents resulting from the
opposite effects of hindrance to nucleophilic attack and resonance would
be a constant (i.e., Charton's "h") or would vary slowly with substituent
size. Such a constant value for the steric effect of substituents could
well lead, as McCoy and Riecke's second statement concluded, to a mis
leading correlation between E° x and polar parameters as in Charton's
equation (7), even though "h" is typically of the same or greater
magnitude than the polar contributions.
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18
McCoy and Riecke made two final points. First, when Charton used
substituents which could exhibit resonance with the reaction site
(i.e., -OR, -CN, -CC^, etc.), the equation which best defined E° x 20w a s :
(18)
For these cases, Charton’s data showed that in most of the correlations
the magnitude of h > BoR *20 This suggests that not only are both
primary and secondary steric effects present, but also an independent
resonance contribution from those substituents which exhibit a direct
resonance interaction with the reaction site. This is consistent with
certain substituents, as defined in equation (4). Second, they claim
that Charton typically uses substituents in the B to C size range where
the primary and secondary steric effects nearly cancel each other
yielding a constant steric value which is probably Charton's "h" value.
Support for this interpretation again comes from Charton's own data.
In his correlations, Charton invariably excludes bulky substituents,
as -0 (0 ^) 3 , since these substituents yield poor correlations and
increase the value of h. McCoy and Riecke state that such substituents
are not within the B to C size range and thus do not yield a cancelling
of the primary and secondary steric effects. Since Charton's equation
for E° (i.e., equation (7) when f y & O ) does not contain steric parameters,
this results in a poor correlation with his polar parameters. Further,
since h' in Charton's equation (16) contains constant steric effects,
as x > inclusion of substituents as - C^Hg)^ should increase its
value. This interpretation is consistent with Charton's observations.
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19
McCoy and Riecke conclude, therefore, that the ortho-substituent
steric constant is best represented by a sum of steric and resonance
contributions. This may be represented as:
E° = «« v + t y r + + h (19)s,x R,X v,x x
rp and 6 *1 represent reaction susceptibility to the primary and secondary
steric effects respectively (i.e., x and a O and are relatively
constant for benzene systems, h represents an intercept constant incor
porating other steric contributions. Since equation (19) was derived
on the basis of a qualitative graphical interpretation of a substituent’s
steric effect, it is a qualitative equation in which no values for the
parameters have been calculated for any specified reaction system.
However, it is qualitatively applicable, in general, to ortho-substituted
benzene derivative reactions of both unimolecular and tetrahedral type
mechanisms.
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20
Mechanism Background
The hydrolysis mechanisms of carboxylic acid derivatives such as
esters, imidates, amides, hydroxamic acids, etc. have been reported to23-29vary with structure, solvent and acid or base concentration. Of
these, the most widely studied have been the esters. In moderate
basicity or acidity, aliphatic and aromatic esters have been shown to24hydrolyze via a mechanism involving a tetrahedral intermediate
(i.e., B ^ 2 or A ^ 2 type mechanism respectively). For a compound of
general form R-C-G, where G = -OR, -NROH, etc., the general
tetrahedral-substitution mechanism in moderate basicity and acidity18 23-25 30-32has been supported by 0 -exchange studies, ’ product
23 24 27 28 27—29 33analyses, 5 * ’ and kinetic data * and is generally repre
sented by Schemes I and II:
0II + KR-C-G + H30 — —
r ° h hi
R-£-G
Jfkl
+ H 2°
OH1R-C-G
+ 0H>11o
jrOH
R-C-&H r -c-£hOH
k (H 0) 'jj --- R-C-OH.
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The acid and base catalyzed hydrolysis of amides was generally
believed to follow this type of mechanism. Product analysis, O1^-
exchange studies and kinetic data supported such a mechanism for the
base hydrolysis, but the total lack of O^-exchange in the recovered
unreacted amide for the acid hydrolysis shed doubt upon a tetrahedral 28 34 36type mechanism for this system. ’ ’ There was also some confusion
as to the position of protonation for the acidic hydrolysis of23 35 36 18amides. ’ ’ The lack of observed 0 -exchange lent support to a
proposed one-step concerted S^2 type mechanism in which water directly
displaces an amine molecule from the _N-protonated a m i d e . H o w e v e r ,
it was found in these studies that the derived rate law was consistent
with the tetrahedral intermediate type mechanism illustrated in scheme
I at moderate acidity.
Such an apparent discrepancy in observed data prompted the study ?Rof benzimidates (i.e., R'-C = N R ’’) as suitable structural models for
benzamide acidic hydrolysis. ^ Benzimidates and N-methylated
benzimidates were known to hydrolyze via tetrahedral mechanisms in
both acidic and basic solvents analogous to schemes I and 11.^'"*®
In acidic media, benzimidates hydrolyze via attack of water on the
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22
protonated imidate to yield an ester and amine product,^ which is
consistent with the tetrahedral type mechanism for the acid hydrolysis
of esters. It was thought that since benzimidates have been shown to
be such good models for ester hydrolysis, in both acidic and basic
media, they would also prove to be good models for amide hydrolysis as
well.
Studies with N-methylated benzimidates showed monotonic changes
in the pseudo-first order rate constant (i.e., k^) and the enthalpy of
activation (i.e., AH*) with successive N-methylations.^ Therefore if
imidates and amides react via similar tetrahedral intermediates, similar
changes in kr and AH* for acidic benzamide hydrolysis should occur with
successive N-methylations. However, for the benzamides studied, Smith
and Yates found that these changes were not monotonic.^ Further, they
found that the order of reactivity with successive N-methylations for
benzamides was primary > tertiary > secondary which contrasted that for
the benzimidates, which were primary > secondary > tertiary.^ Smith
and Yates concluded t h a t : ^ ’^
1. There is a change in the hydrolysis mechanism from benzamide to N-methyl- and N, N-dimethylbenzamide.
or
2. All three benzamides do not conform to an oxygen protonated, tetrahedral type mechanism (i.e., Aq2, which is a subset of the AAC;2 type).
Changes in reaction mechanism with structure have been observed
in the hydrolysis of carboxylate esters, imidate esters, other than
the above, aminolysis of carboxylate esters and in the alcoholysis of 24carboxylic acids. It was not inconsistent, therefore, that successive
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23
N-methylations could cause a mechanistic change for the acid catalyzed
amide hydrolysis. However, use of benzimidates as structural models
for these reactions is inadequate since, in addition to the above
differences, extensive O^-exchange was observed for both acid and
base catalyzed imidate hydrolysis.32In 1975, R. McClelland found a small but detectable and repro
ducible amount of O^-exchange in the acid hydrolysis of benzamide in
oxygen labeled water. This result has been interpreted as support for Tthe Aq2 type mechanism by McClelland. Such a mechanism is illustrated
in scheme III:
Scheme III
OHI
r -£-n h
OHE t - c - f o i ,
OHIR-C-NH0
+ o 1bh 2
+ ?H 2 • R-C-NH„’ i - H 2
k ‘
McClelland argued that this mechanism implies, depending on the size of
He noted that, although previously suggested,28’8^ ’88 if k, /kh elarge enough, (i.e., Bender and Ginger8^ have placed a limiting value
of kh/kg = 37A as the largest ratio from which 018-exchange can be
observed for amide hydrolysis) decomposition of the tetrahedral inter
mediate may occur without any significant 0̂ "8-exchange. The fact that
a small amount of 0^8-exchange was observed simply implies that the
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24
ratio was near its limiting value, according to McClelland. McClelland's
findings supported Smith and Yates' initial conclusion regarding the
mechanisms behind the acid hydrolysis of benzamide, N-methyl- and
N, N-dimethylbenzamide. Therefore, it seems likely that although Tbenzamide hydrolyzes via the Aq2 mechanism at moderate acidity, there is
a change in the mechanism for the N-methylated benzamides. Such a
conclusion was in contrast to that of the benzimidates which are
structural analogs of the benzamides. Successive N-methylations of the28benzimidates showed no evidence for a change in mechanism.
38 39Changes in mechanism also occur with changes in pH. McClelland
and Smith and Y a t e s ^ studied the changes in the acidic hydrolysis
mechanism of substituted benzimidates with pH. At moderate acidity
(i.e., 7 > pH > 1), the normal hydrolysis mechanism involving a tetra
hedral intermediate produced the expected ester and amine products28 40 18exclusively. * As expected with such a mechanism, extensive 0 -
exchange was observed. At increased acidity (i.e., pH < 1), the break-28down pathway of the tetrahedral intermediate was reportedly changed.
Products obtained from the oxygen-labeled imidate hydrolysis included a
significant concentration of starting imidate, unlabeled carboxylic28acid, labeled alcohol and unlabeled amide. Smith and Yates interpreted
these results in terms of a competition between proton addition to the
- O ^ R and -NHR groups in the tetrahedral intermediate at higher acidity.
At the higher acidity, proton addition becomes less selective, and
addition to the - O ^ R group, forming the less stable Ҥ18r cation in the
tetrahedral intermediate, becomes more pronounced yielding the above
mentioned products following intermediate decomposition. The competition
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25
9ieR ?18r ,, ^ 918rk, (H_o) |R'-C=NR" + H o0t k ^ R ’-C=ShR m + H20 - - -- - R ’-C-NHR"
J FI5 OHI High
Low Acidity_______________1 Acidity
+ 0 18RI , 1-C-^H2R" + H20 r '-c-n h r "OH OH
0 0II IIR ’-C + NH R " R ’-C-NHRNn 1br i nU R + R018H
38 39At very high acidity (i.e., pH < -1, -^65% H 2SO^) McClelland found ’
that the hydrolysis of oxygen-labeled ortho-substituted benzimidate
yielded labeled amide and unlabeled alcohol. These results were inter
preted in terms of a competition between the accepted tetrahedral
acidity) and an SN2 type mechanism not involving a tetrahedral inter
mediate (i.e., an A ^ 2 mechanism in which AL = alkyl oxygen cleavage).
McClelland concluded that as acidity becomes very high, "two competing
pathways for imidate hydrolysis ... appear to be possible. Both
involve attack of water on the protonated imidate but differ in the
competition between these mechanisms, due to changes in pH, is
illustrated in scheme V . ^
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26
0 18R o 18rR-C=NR" + H_0+ ---- ^ R-C=5h R" + H.O3
& High Acidity Super High j [•' (H20)jl Acidity
0 18R' O 181 + nR-C-NHR" + H„0 R-C-NHRIOH
cf. scheme III
R-C-NHR1 + R't)18H
Changes in mechanism with changing pH have also been reported for
carboxylate e s t e r s , a m i d e s ^2 *2^ ’^2 and hydroxamic acids.2^ In
some cases, the hydrolysis mechanism has been discussed in terms of
Hammett or Taft reaction and substituent parameters (i.e., p, 6 , a or . „ . 2,26,29,30,43,44
Es>-For example, Buglass and coworkers have studied the hydrolysis of
para-substituted benzohydroxamic acids over a relatively large acidity
range (i.e., [HCIO^/ = 1.0 to 5.0 M ).29 They have interpreted the
mechanism in terms of a previously accepted two-step bimolecular
and benzimidates in moderate acidity. This specific mechanism is
illustrated in scheme VI:
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27
Scheme VI
OHI + KR-C-NHOH + H30 ' ^IZ± R-C-NHOH
+OH,2
wR-C-NH„OH I 2
OH I H
\'OH 3 IOH
They have calculated rate constants for the second step in the
mechanism, (i.e., nucleophilic attack by water to form the tetrahedral
intermediate), and reported a positive Hammett p value for the
correlation of those rate constants. Such a value is consistent with
the above bimolecular mechanism in which electron withdrawing groups
enhance nucleophilic attack. Further examination of their data
indicates a fair correlation between their observed overall rate
constants and Hammett a values with the overall p value being negative.
This result, which is consistent with the aliphatic hydroxamic acid
series,^ is also consistent with our previous work on the ortho-45substituted benzohydroxamic acid series in which a bimolecular
mechanism analogous to scheme VI was supported. From these results
and the studies of Buglass and coworkers, the first step in the
mechanism illustrated in scheme VI appears to be susceptible to polar
effects to a greater extent than is the second step [ i . e . , p (overall)
represents the sum of p's for the two steps in the mechanismJ.
Ahmad, Socha and V e c e r a ^ have studied the alkaline hydrolysis
of benzohydroxamic acid over a wide hydroxide ion concentration range
(i.e., 0.12 M to 2.18 M ) . The authors concluded that, depending upon
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28
pH, the hydrolysis was a function of contributions from competing
mechanisms in which either hydroxide ion or water is the attacking
nucleophile. Their two competing mechanisms are illustrated in schemes
VII and VIII.
Scheme VII
0IIPh-C-NHO + H 20
?" ?'------- Ph-C-NHOH ■ — Ph-C-NH.OHI . . -- 1 2
0II
Ph-C-NHOH + OH
Scheme VIII
(A) At lower pH (i.e., 7 to 9):
II kiPh-C-NHOH + H„0------0 0
Ph-C-NHOH ^ Ph-C-fe0OHJ ^---- 1+0H2 oh
H o0 + Ph-COO + NH OH J l(B) At high pH (i.e., 12 to 14):
Ph-C-NHOH + OH ■ Ph-C-NHO + H 90 -0
- Ph-C-NH0_ +0H„I I0
|Ph-C-SH„OH I 2
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29
It should be noted that other tautomeric structures for the charged
intermediates presented in schemes VII and VIII may exist, and these
presented here simply represent a recitation of the authors' proposed
mechanisms.^
These two schemes yield observed rate constants (i.e.,
which, at high pH, are independent of hydroxide ion concentration.
This is consistent, they claimed, with their observed data which was
used to construct a profile on the pH dependence of kQbs over a range
of ^ 4 to 14 for benzohydroxamic acid and for meta-nitrobenzohydroxamic
acid, the reported rate constants in the acid range were apparently
obtained from other sources. The profile shows that at pH = 4 t o — ?,
the hydrolysis rates of both hydroxamic acids are pH independent. At
pH > 7, kQbg increases with pH until at pH Xs 11, the rate constants
for both hydroxamic acids become pH independent. Ahmad and coworkers
interpreted this data in terms of a mechanistic competion between
schemes VII and VIII. At low pH, most of the hydroxamic acid is in
the undissociated form. Water is the predominant attacking nucleophile
within the pH range implying mechanism (A) in scheme VIII as the sole
contributor to the rate law. This conclusion, claim the authors, is
consistent with their observation of pH independence of kQbg within
this range. As pH rises, more hydroxamic acid is converted to its'
conjugate base form and contributions to kQbg from the mechanism of
scheme VII and, even more, from mechanism (B)-scheme VIII increase.
Finally, at high pH (i.e., pH > 11), all the hydroxamic acid is in its
conjugate base form and any further addition of hydroxide ion has no
effect on the conjugate base concentration. At this pH, mechanism
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30
(B)-scheme VIII predominates implying independence of kQbs on hydroxide
ion concentration. This conclusion, according to Ahmad and coworkers,
is again supported by their observations illustrated in the profile.2^
The authors note that the derived rate laws for scheme VII and
mechanism (B)-scheme VIII are kinetically indistinguishable, both
resulting in independence from hydroxide ion concentration. However,
because of the great excess of water present at all base concentrations
considered, the authors concluded that the general rate law describing
the hydrolysis over the entire pH range studied (i.e.,^-4 to 14) could
be represented by a combination of the mechanisms in scheme VIII only.
Their rate law is given as equation (20), in which this particular
form assumes the slow step as k2>
kobs = [ } l + (k2K/|HHJ ) J / l + (K/[H+J)
or (20)
kobs = (kl + k2 (K/Ka))C0H ■2 / 1 + (K/Kio) (0H~J
A Hammett correlation with meta- and para-substituted benzo
hydroxamic acids was made. The value of PQ^S = +0.118 supports the
above rate law according to Ahmad and coworkers. They argued that
since PQ|JS is an overall value (i.e., PQ|JS = PK + P^ )> the value of
p for the nucleophilic attack step in the hydrolysis mechanism
illustrated by scheme VIII-B can be easily calculated from pKa values
reported for the hydroxamic acids studied. They obtained a value of
p = + 1.0 for that step which is consistent with the fact that electron
withdrawing groups enhance nucleophilic attack by water.
This rate law and proposed mechanism for the alkaline hydrolysis
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of benzohydroxamic acid at pH > 13 is in contrast to earlier work
mechanisms contributing to at hydroxide ion concentrations
n . k2R-C-NHOH + OH(A) (D)
K- it." NC
(E)k 2D + H O — Products k3E + H„0 -----> Products
- k4D + OH -----> Products- k5E + OH — Products
27Their derived rate law from this scheme is -given as equation (21):
k2K2 + k3K3 + [k4K2 + k5KJ [0H"Jl/fOH"J + Kkobs - ^ (21)
studied, K 2 and K3 >> and equation (21) reduces to:
' [0H"J (22)
utions to kQks by the
via the "one hydroxide ion pathway", and the "two hydroxide ion pathway"
respectively. The authors calculated values for k' and k" from a plot
of k , versus [OH J and found k' = 0.041 hrs. 1 and k " = 0.012 hrs.-1 .
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32
These values Indicate that when hydroxide ion concentration is less
than 0.1 M the contribution to kQbs from the hydroxide ion dependent
k" term becomes negligible and the rate law takes a form which is
independent of hydroxide ion concentration. This is consistent with
the observations and conclusions of Ahmad and coworkers.^ However,
at pH > 13, the contribution from the k" term to kQbg increases and
at CoH U = 0.12 to 2.2 M, kQbs shows a linear dependence on hydroxide
ion concentration corresponding to a significant contribution by
k' (pH J. This conclusion is in contrast to that of Ahmad and co-
workers^ who claimed hydroxide ion independence of kQbg above27p H £12. The results of Berndt and Fuller's work indicates that
Ahmad and coworkers^ have neglected the contribution of the two-
hydroxide ion pathway in their extrapolation beyond pH = 13, and that
Ahmad and coworkers'^ derived rate law, implying no mechanism change
beyond pH = 13, must also be invalid at higher hydroxide ion concen
trations.
Purpose
The system chosen for study is the acidic and basic hydrolysis
of ortho-substituted-N-methylbenzohydroxamic acids. This system is a
structural analog of the amides and imidates discussed above, and is26,27,29more hindered than the previously studied benzohydroxamic and
45ortho-substituted benzohydroxamic acid series. Increased hindrance
over these two previously studied systems is provided by N-methylation.
Since both of these systems have been shown to hydrolyze, in bothTacidic and basic media, via a tetrahedral type mechanism (i.e., A q 2)
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33
it is of interest to determine if, as with the N-methylatedOQ OO OA
amides * * ’ discussed above, the hydrolysis mechanism changes
with increased hindrance, or if, as with the N-methylated benzimi- 28 36 40dates ’ ’ the mechanism remains unchanged.
Secondly, it has been shown that the hydrolysis mechanism for
benzohydroxamic acid changes with p H . ^ However, the general rate
law of Ahmad and coworkers^ is in contrast to that of earlier work27by Berndt and Fuller at higher basicity. A study of the hydroxide
ion concentration dependence of kQ^g at high basicity (i.e., [pH J ='VJ.O to '*'7.5 M) will provide valuable support or contradiction to the
general rate law proposed by Berndt and Fuller. Further, determination
of the rate law at such basicity levels will provide evidence for a
further mechanism change beyond that proposed by Berndt and Fuller.
Thirdly, in an attempt to study the "Ortho-Effect" and the validity
of ortho-substituent and reaction parameters as defined previously,
application of the two-parameter Taft-Pavelich equation (i.e.,
equation (12)) to this "hindered" system will be attempted. Successful
application would provide four useful conclusions:
1. The same mechanism must be operating for all compounds in the series.
2. Calculated values for p* and 6 will help support or refute a bimolecular mechanism similar to that found in our earlier study as the hydrolysis mechanism in acidic and basic media.
3. A comparison of p* and 6 values for the acidic hydrolysis of the ortho-substituted-N-methylbenzohy- ^ droxamic acid series with those from our previous work will provide useful information in the qualitative determination of McCoy and Riecke's interpretation of the "Steric Ortho Effect"^ over that presented by Charton.1
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4. Successful application of Taft's a * and E° values to a more hindered system than that previously studied provides further support for Taft's separation of polar, steric and resonance effects assumption discussed above for equation (3).
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EXPERIMENTAL METHOD, APPARATUS AND SYNTHESES
Preparation of Ortho-Substituted Benzoyl Chlorides
The preparation of the acid chlorides for use in the synthesis of
ortho-substituted-N-methylbenzohydroxamic acids followed a general46procedure, except for the preparation of ortho-nitrobenzoyl chloride
which will be described separately. Such a general procedure is47illustrated by the following synthesis of ortho-methylbenzoyl chloride
and is, therefore, also applicable to the preparation of the ortho-48 48 49chloro, -bromo, and -methoxybenzoyl chlorides.
ortho-Methylbenzoic acid (0.2 moles) was refluxed with thionyl
chloride (80 grams, 0.67 moles) for six hours, which was one-half hour
after the mixture turned to a clear solution. The resulting solution
was distilled at room temperature and reduced pressure (10-15 mm Hg) to
remove unreacted thionyl chloride. The remaining residue was distilled
under reduced pressure (3-4 mm Hg, 66°-68°C) to yield the acid 47 52chloride, * which was used without further purification.
To ortho-nitrobenzoic acid (0.2 moles) cooled in an ice-water bath,
thionyl chloride (50 grams) was slowly added. The mixture was then
brought to room temperature and left standing approximately ten minutes.
It was then gently refluxed for nearly one hour at which time most solid
had dissolved forming a dark yellow mixture. The solution was then
suction filtered while still hot to remove impurities. On cooling to
room temperature, a precipitate formed. Unreacted thionyl chloride was
then evaporated by purging with nitrogen. The remaining mother liquor
35
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36
was separated by suction filtration and again purged with nitrogen to
remove final traces of thionyl chloride to yield 13.0 grams of the
yellow-orange acid c h l o r i d e . T h e acid chloride was used without
further purification.
The following table illustrates the total yields obtained for
each acid chloride prepared.
Table I
Yields of Prepared Ortho-Substituted Benzoyl Chloridesn 1̂
______ Substituent Yield,_Percent_______
-CH3 71.0
-och3 75.5
-Cl 73.5
-Br 96.4
-I 39.0
-N°2 70.0
aSee refs. 46-49. ^Yields are based on 0.2 moles of the corresponding acid.
Preparation of Ortho-Substituted-N-Methylbenzohydroxamic Acids
The preparation of the hydroxamic acids followed a general
procedure adapted from the method of Ulrich and Sayigh"^ except for
the preparation of ortho-nitro-N-methylbenzohydroxamic acid which will
be described separately. Such a general procedure is illustrated by
the following preparation of ortho-bromo-N-methylbenzohydroxamic acid
and is, therefore, also applicable to ortho-methyl, -chloro, -iodo and
methoxy-N-methylbenzohydroxamic acids.
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37
N-Methylhydroxylamine hydrochloride (0.05 moles) and sodium
carbonate monohydrate (0.05 moles) were mixed in methanol (40 ml).
The mixture was stirred to maintain the pH £ 7. Ortho-bromobenzoyl
chloride (0.05 moles) was added dropwise to the constantly stirred,
ice-water-bath-cooled mixture over a period of thirty to forty minutes.
The pH was frequently tested, and sodium carbonate was added when
needed to maintain neutral or basic conditions. The milky white
mixture was suction filtered, and the residue was washed with methanol
(10-20 ml). The washings and filtrate were combined and the solvent
evaporated using a rotary evaporator. The resulting oil, which
solidified on cooling, yielded a precipitate which gave a positive
ferric chloride test (for explanation of the ferric chloride test, see
"Preparation of Reaction Solutions and Kinetics Procedure", below).
The precipitate was extracted with hot benzene by gently refluxing it
for five to ten minutes. The benzene layer was then separated and
cooled ('~'10°C) to yield off-white crystals (2.9 grams) which gave an
intensely positive ferric chloride test. The crystals were then twice
recrystallized from benzene and once from carbon tetrachloride to
finally yield the hydroxamic acid (1.45 grams, see Tables II and III).
ortho-Nitrobenzoyl chloride (0.1 mole) was added dropwise to the
ice-water-bath-cooled N-methylhydroxylamine solution described above
over a thirty to forty minute period. pH tests were frequently made
on the constantly stirred mixture, and sodium carbonate was added when
required to maintain pH - 7. The mixture was then suction filtered and
the residue was washed with methanol (''-'50 ml) . The washings and clear
orange filtrate were combined, cooled overnight and filtered to yield
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38
bright yellow crystals (7.5 grams). The mother liquor was evaporated
to about 25 ml using a rotary evaporator and cooled in a refrigerator
for two hours. The resulting off-yellow crystals (3.45 grams) were
separated, and both sets of crystals gave a positive ferric chloride
test although reaction was not immediate. Recrystallization was
attempted with benzene, toluene and chloroform, but only successfully
accomplished with ethyl acetate. Two recrystallizations of the combined
solid yielded the hydroxamic acid as light-yellow needles (6.5 grams,
Tables II and III) which gave an intensely positive ferric chloride
test although reaction was not immediate.
Each of the new hydroxamic acids prepared were analyzed by infrared
and nuclear magnetic resonance spectroscopy. The elemental analysis
for percent carbon, nitrogen and hydrogen was performed by Galbraith
Labs., Knoxville, Tennessee. The following tables list yields, observed
melting points and results of the elemental analysis for each new
hydroxamic acid prepared.
Table II
Yields of Prepared Ortho-Substituted-N-Methylbenzohydroxamic Acids3
Substituent___________ Yield, Percent_________ Observed M.P., °C
“ CH3 17.6 117.0 - 118.0
- o c h 3 15.8 138.5 - 139.2
-Cl 22.0 118.0 - 119.0
-Br 13.1 135.0 - 135.8
-I 9.1 145.1 - 145.8
-N°2b 33.1 170.8 - 171.6 (d)
aAll prepared hydroxamic acids are new compounds. ^Compoundobserved to decompose on melting.
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39
Table III
Elemental Analyses of Prepared Ortho-Substltuted-N-Methylbenzohydroxamic Acids
Substituent_______Analysis3________ %_C_________ % N__________ % H
-CH_ Theoretical 65.45 8.48 6.673 Observed 65.14 8.47 6.48
-0CH„ Theoretical 59.67 7.73 6.083 Observed 59.52 7.64 6.14
-Cl Theoretical 51.76 7.55 4.31Observed 51.61 7.59 4.38
-Br Theoretical 41.76 6.09 3.50Observed 41.97 5.89 3.36
-I Theoretical 34.68 5.05 2.91Observed 34.87 4.93 3.07
-N0„ Theoretical 48.97 14.28 4.12I Observed 48.87 14.26 4.04
aObserved analysis performed by Galbraith Labs., Knoxville, Tennessee.
Preparation of Standard Hydrolysis Solvents
For both acid and base catalyzed systems, water was the chosen
solvent. The water used in preparation of the standard solutions was
doubly distilled and all prepared hydroxamic acids were found to be
soluble at levels greater than those employed in the kinetic runs
(0.01M) and at temperatures lower than those employed (90.0 - 0.1°C).
Standard acid solvents were prepared from hydrochloric acid
(~37%). Each acid solvent was standardized by titration.
The standard sodium hydroxide solutions were prepared from a
saturated sodium hydroxide solution in order to minimize levels of
absorbed carbon dioxide from the air. The saturated solution was
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40
prepared from excess sodium hydroxide pellets dissolved in redistilled
water purged with nitrogen. The solution was then filtered through a
sintered glass funnel, decanted into a nalgene container and again
purged with nitrogen to insure an inert atmosphere. Each standard base
solution was prepared by careful dilution of the saturated solution
with nitrogen purged, heated, redistilled water. Each prepared solution
was then subjected to further nitrogen bubbling to insure that there was
an inert atmosphere and then stored in airtight nalgene bottles.
Standardization was accomplished by titration with a standard hydro
chloric acid solution.
For both acidic and basic solvent systems, ionic strength was kept
constant for the various acid and base concentrations using monovalent
salts; sodium chloride for the base system and potassium chloride for
the acid system. The salts were of A.C.S. grade and dried in a
desiccator oven overnight before weighing and addition to the prepared
acidic or basic solvent.
Preparation of Standard Ferric Chloride Solutions
The ferric chloride solution used in the acid catalyzed hydrolysis
was prepared by dissolution of ferric chloride hexahydrate (2 grams)
in 200 ml of distilled water with 5-7 drops of hydrochloric acid (37%)
added. The solution was suction filtered, and 10 ml aliquots were
pipetted into 25 ml volumetric flasks. One flask was diluted to 25 ml
with distilled water and used as the blank solution. The others were
used as sample flasks.
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41
The ferric chloride solution used in the base catalyzed hydrolysis
was prepared as described above with some modification. To maintain
acidity of each 10 ml aliquot of standard ferric chloride solution
after addition of the 1.0 ml reaction solution aliquot, (for explanation
see "Preparation of Reaction Solutions and Kinetics Procedure", below),
the initial ferric chloride solution was prepared with a 2.5 fold excess
of acid over the amount of base present in the added 1 ml reaction
solution aliquot. Such an excess was necessary to prevent formation of
ferric hydroxide and insure proper complexation of the ferric ion with
the unreacted hydroxamic acid (for explanation see "Preparation of
Reaction Solutions and Kinetics Procedure", below).
Preparation of Reaction Solutions and Kinetics Procedure
The reaction solutions used for rate measurements for the acid and
base catalyzed hydrolyses at all concentrations of acid and base
employed were prepared in the following manner: a 0.01M solution of
the chosen hydroxamic acid was made by dissolving the appropriate weight
of the acid in a 20 ml nalgene cell into which 15 ml of the appropriate
standard acid or base solvent was pipetted. Complete solution was
attained by steam heating the airtight stopped cell. The solution was
then placed in a constant temperature oil bath held at 90.0 - 0.1°C.
The reaction solution was typically given five to 10 minutes to come
to temperature equilibrium. At this point, a 1.00 ml aliquot was
pipetted into one of the previously prepared sample flasks containing
the standard ferric chloride solution. A different pipet was used for
the acid and base systems to prevent any chance of contamination. The
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42
sample was then diluted to 25 ml with distilled water, and the absorbance
of the solution relative to the previously prepared blank was determined
with a Beckman D.U. spectrophotometer using two 1-cm Beckman quartz
cells at a predetermined wavelength. For the acid catalyzed reactions45at all acid solvent concentrations used, the wavelength was 520 nm for
all hydroxamic acid runs except for ortho-nitro-N-methylbenzohydroxamic
acid for which 500 nm was used. For the base catalyzed reactions at all
base concentrations used, the wavelength was also 520 nm except for
ortho-nitro-, ortho-bromo- and ortho-iodo-N-methylbenzohydroxamic acids 45for which 500 nm was used.
The relative quantity of remaining unreacted hydroxamic acid is
determined from the absorbance of the ferric ion-hydroxamic acid complex
which forms in the sample flask. As the hydroxamic acid concentration
decreases during hydrolysis, so does the concentration of the complex
and thus the absorbance. The complex which forms is the characteristic
purple magenta complex between ferric ion and the hydroxamic acid
functional group.^ At the complex concentration range under study,43Beer's law has been shown to apply.
The spectrophotometer cells were calibrated by filling both with
distilled water and measuring the absorbance of one relative to the
other at the employed wavelengths. The cells were found to be identical
within 0.002 absorbance units.
Each pipetted 1 ml sample from the reaction solution was taken at
a specified time interval depending on the reaction rate. The following
tables illustrate the number of runs taken for each hydroxamic acid in
both the acid and base catalyzed systems and the observed rate constants
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43
at the acid and base concentrations employed. Also illustrated are
the observed rate constants for reactions of specified hydroxamic
acids at various temperatures, with various salts and salt concentra
tions. The high ionic strength maintained in the base catalyzed system
prompted the expectation of specific salt effects on the hydrolysis
rate constants. Such effects, manifested as specific ion effects, will
be discussed in the next section.
Table IV
Rate Constants for Base Catalyzed Hydrolysis of 2-Chloro-N-Methylbenzohydroxamic Acid at 90.0°C
NaOH, M Trial 106k . aobs % Mean Deviation
7.31 1234
Average
2.532.56 2.61+2.59-(0.006)2.57 1.07
6.58 12
Average
2.23.2.32-(0.003)2.27 1.98
5.47 1234
Average
1.811.921.90-(0.008)1.93 1.89 2.12
4.40 12345
Average
1.52-(0.007)1.421.471.541.391.44 3.89
3.23 12
Average
0.890.94-(0.001) 0.92 2.72
Pseudo first order rate constant in sec“l. Ionic strength maintained at 7.31M with IjlaCl. An average of 6 points was used for slope calculations. - represents typical values for one standard deviation as determined by least squares.
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Table V
Rate Constants for Acid Catalyzed Hydrolysis of 2-Methyl-N-Methylbenzohydroxamic Acid
HC1, M Trial Temp, °Ca 1()5kobsb % MeanDeviation
0.149 1 1.452
Average 90.01.51-(0.009) 1.48 2.03
0.225 12
Average 90.0
1.96^(0.02)2.122.04 3.92
0.451 1234
Average 90.0
3.162.942.892.972.99 2.84
0.595 12
Average 90.0
3.853.70-(0.04) 3.78 1.98
0.751 12
Average 90.0
4.624.604.61 0.22
0.751 12
Average 80.0
1.811.901.86 2.42
0.751 12
Average 70.0
0.8140.8210.818 0.43
aTemperature controlled to - 0.1°C. Pseudo first order rate constant in sec- . Ionic strength maintained at 0.751M with KC1. - represents typical values for one standard deviation as determined by least squares.
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45
Table VI
Rate Constants for Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C
Substituent Trial 1()5k w 3 : obs Z Mean Deviation
-CH3b 1 4 *42+2 4.30-(0.02)3 4.41
Average 4.38 1.14
-OCH 1 10.912 10.713 10.95
Average 10.86 0.890
-Cl 1 2.552 2.623 2.62
Average 2.60 1.15
-Br 1 1.932 1.953 1.92
Average 1.93 0.518
-I 1 1.59-(0.02)°2 1.62
Average 1.605 0.935
"N02 1 0-5862 0.580-(0.005)3 0.572
Average 0.579 0.864
aPseudo first order rate constant in sec- . A different set of cells were used than those used in the corresponding runs in Table V. °Represents typical values for one standard deviation as determined by least squares.
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Table VII
Rate Constants for Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids in 7.31M NaOH at 90.0°C
Substituent Trial loSi , 3 % Mean Deviationobs
- c h 3 1 0.7232 0.759
Average 0.741 2.43
-OCH- 1 3.21-(0.002)2 3.13
Average 3.17 1.26
-Cl 1 2.592 2.61
Average 2.60 0.385
-Br 1 1.18-(0.002)2 1.25
Average 1.215 2.88
-I 1 0.8162 0.800
Average 0.808 0.990
-n o 2 1 6 .86-(0.01)2 6.78
Average 6.82 1.17
aPseudo first order rate constant in sec . - representstypical values for one standard deviation as determined by least squares.
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47
Table VIII
Rate Constants for Catalyzed and Uncatalyzed Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids
at 90.0°C in the Presence of Salts
Substituent HC1, M Salt Ionic Trial 10^k , aStrength, M °
-°H3 0 - 0 12
Average
0.360^(0.004)0.4400.400
“CH3 0 KC1 0.751 12
Average
0.7550.7820.773
“CH3 0.149 KC1 0.751 12
Average
14.5215.14^(0.009)14.83
"CH3 0.150 CsCl 0.751 12
Average
13.3313.2313.28
-Cl 0 NaCl 3.00 12
Average
1.131.091.11
-Cl 0 NaCl 6.31 12
Average
2.001.89^(0.002)1.94
-Cl 0 NaBr 6.31 12
Average
1.331.34 1.33
Ps e u d o first order rate constant in sec . - representstypical values for one standard deviation as determined by least squares.
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48
The Constant Temperature Oil Bath
Reaction temperature was held constant to - 0.1°C by a constant
temperature oil bath. The basic apparatus employed is illustrated in
Figure 1.
A heating coil (B) was connected to a voltage regulator to supply
the necessary heating to keep the bath (A) at constant temperature. A
thermoregulator (C) was immersed to the same depth as the thermometer
(D), which had been calibrated in 0.1°C increments to insure accurate
temperature readings. Calibration of the thermometer was made against
a standard thermometer of known stem correction in the temperature
range studied. The thermoregulator was connected to the input terminals
of a relay (E) which was, in turn, connected to the voltage regulator.
The thermoregulator was then set at 90.0°C. To prevent extensive heat
loss through the walls of the bath, the container was insulated with
vermiculite packing. A mechanical stirrer (F) was also employed to
insure even heating throughout the bath. The apparatus, after con
struction, was tested for twenty-four hours to insure precision in
temperature control. Variation was never greater than - 0.1°C.
Determination of Rate Constants
The pseudo first order rate constants, f°r aH runs were
determined via the relationship between measured absorbance and
hydroxamic acid concentration. The following derivation illustrates
this relationship between measured absorbance, kQ^g , and hydroxamic
acid concentration.'*3' Since the hydrolysis rate is pseudo first-order,
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49
?o
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Figure
1. Ex
peri
ment
al
Appa
ratu
s
where a = initial hydroxamic acid concentration,x = concentration of acid reacted in time t, and k = first-order rate constant.
Concentration of the hydroxamic acid may be related to some
physical property, A, which is directly proportional to concentration.
For the above rate expression this may be illustrated as:
log ( ^ ) - log (24)oo
where A = measured property at time infinity, andAq = measured property at time = 0.
Since absorbance is directly proportional to the concentration of
hydroxamic acid, equation (24) may be written as:
log = log (/ ° . 4 ° -) (25)a x oo At
Since Aq o = 0 for the hydroxamic acid-ferric ion complex, because at
time infinity no hydroxamic acid remains, the rate expression is
finally represented as:
108 At = 2.303" + 108 Ao (26)
A plot of the log of measured absorbance (log Afc) versus time yields a
slope of -kQbg/2.303. A least squares treatment of the log of absor
bance versus time data was used for actual determination of the observed
pseudo first-order rate constant.
Reaction Product Analysis of Selected Hydroxamic Acids in Alkaline Solutions
Product analyses for the alkaline hydrolysis of ortho-nitro-,
ortho-chloro- and ortho-methoxy-N-methylbenzohydroxamic acids were
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51
performed at base concentrations approximating those of the kinetic
runs. Product identifications were obtained by melting points and by
comparison of the product infrared spectrum with a standard infrared 52spectrum of the expected ortho-substituted benzoic acid product.
ortho-Chloro-N-methylbenzohydroxamic acid (0.37 g) was added to
7.5M sodium hydroxide solution (40 ml) prepared via dilution of the
previously described saturated sodium hydroxide solution with nitrogen
purged distilled water. The reaction, carried out in 50 ml nalgene
cells, was run at 90.0° - 0.1°C for 20 days which, according to its
half-life calculated from the kinetic runs, corresponded to 98%
reaction. On acidification of the reaction solution with concentrated
hydrochloric acid in the cold over a period of 40 minutes a white
precipitate formed (0.28 g, 84.3% yield, m.p. 137-139°C) and was
separated by suction filtration. The precipitate was recrystallized
from hot water to yield 0.20 grams of ortho-chlorobenzoic acid (60.2%
yield, mp.p 140.5-141.5°C, lit.53 140-141°C).
ortho-Methoxy-N-methylbenzohydroxamic acid (0.37 grams) was added
to '-'7.5H sodium hydroxide solution, and the reaction carried out for
21 days (98% completion) as described above. The reaction mixture was
then acidified as described above and extracted three times with 15 ml
portions of absolute ether. The ether layer was dried with calcium
chloride and evaporated with air to yield an off-white precipitate
(0.283 grams, 85.5% yield, m.p. 94-96°C). The precipitate was recry
stallized from hot water to yield 0.202 grams of ortho-methoxybenzoic
acid (62.2% yield, m.p. 98-99.5°C, lit.53 100-101°C).
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ortho-Nitro-N-methylbenzohydroxamic acid (0.437 grams was added to
/^7.5M sodium hydroxide solution (10 ml), and the reaction carried out
for 14 days as described above. The dark orange solution was then
acidified in the cold with concentrated hydrochloric acid over a period
of 30 minutes. The resulting tarry-black precipitate (^0.8 grams,
liquifies ^190°C) was separated by suction filtration. The mother
liquor was then allowed to stand for one week after which time a black-
gritty precipitate formed and was separated. The two precipitates were
left to dry for 7 days to yield a dark brown powder when crushed
(^0.09 grams, m.p. 110-114(d)°C). The precipitate, along with the
mother liquor, was then extracted with absolute ether. The ether layer
was dried with calcium chloride and air evaporated to yield a light
brown precipitate (0.08 grams, m.p. l.'50-152(d)°C, lit.33 for ortho-
nitrobenzoic acid, 146-147°C). An infrared spectrum of the precipitate
showed some similarity with that of a standard ortho-nitrobenzoic acid 52spectrum. However, an exact match was not made. Recrystallization
from hot water yielded a similar light brown precipitate (0.04 grams,
m.p. 149-151(d)°C). An infrared spectrum of this compound was again52similar, in some aspects, to the standard spectrum, but an exact match
was not made. The precipitate, therefore, could not be conclusively
identified.
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RESULTS AND DISCUSSION
Acidic Hydrolysis Mechanism
The first-order dependence of the observed first-order rate
constant (i.e., k , ) on hydronium ion concentration given in Table V obsis illustrated in Figure 2. Such a dependence is consistent with that
observed in previous studies for less hindered benzohydroxamic acid
hydrolyses over comparable hydronium ion concentration ranges and2 27 29 45hydrolysis temperatures. These previous studies also showed ’ ’
that such a dependence is consistent with a bimolecular mechanism
involving a tetrahedral intermediate analogous to those supported for32 28 38the acidic hydrolysis of benzamide, benzimidates ’ and aryl
esters.^ The general acidic hydrolysis mechanism for the previously
studied benzohydroxamic acids is illustrated in scheme VI.
Values for the activation parameters, AS* and AH*, have been
calculated for the presently studied system from the rate data given
in Table V. Table IX compares these values of AS* and AH* for the
present system with those obtained from previous studies of less27 2(hindered benzohydroxamic acid hydrolyses under similar conditions. ’
The values listed in this table are in the usual range for the bimole-
cular-tetrahedral mechanism for ami d e s ^ and benzimidates^ lending
further support for such a hydrolysis mechanism for the ortho-
substituted-N-methylbenzohydroxamic acid series. Furthermore, the
enthalpy of activation is higher and the entropy of activation is more
negative for the ortho-methyl-N-methylbenzohydroxamic acid hydrolysis
53
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54
.60-
3.80-
obs
3.00-
1,40 o.io oTio(hciJ
Figure 2. Dependency of k , (sec-'*') on catalytic acid concentration (M)
O'. 50 0.70
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55
Table IX
Activation Parameters3 for Acidic Hydrolysis ° f R.j -<0V CO-N-OH
Rx R2 R3 AH*, (Kcal/mole) AS*(e.u.)
CH3 °H3 H 20.8-(0.8) -21.2-(1.3)
c h 3 H ch3 19.4-(0.3) -17.8b-(0.4)
H H H 19.4 -20.7C
H H H 20.2 -17.9d
ELCalculated from second order gate constants. - represents values for one standard deviation. From data reported in ref. 33. cCalculated from data from ref. 2 7 at two temperatures, 0.485M HC1, ionic strength 0.577M (KC1). dRef. 29 at 1.00M HC104 , five temperatures, ionic strength 1.00M.
than for the corresponding para-methyl compound. These results are as
expected, although not conclusive due to uncertainties in AH* values,Tfor a bimolecular-tetrahedral mechanism (i.e., A 2) in which the more
hindered compound shows a greater amount of conformational restriction
within the tetrahedral intermediate and larger activation energy for
the formation of the intermediate.
The rate law for such a bimolecular-tetrahedral mechanism,
analogous to that in scheme VI, which is consistent with the above data,
is similar to that previously derived for the acidic hydrolysis of27benzohydroxamic acid under similar conditions. It may be derived
from equations (27) - (29) in which the tetrahedral intermediate is
not shown, but may be involved analogously to the hydrolyses of amides,
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esters and benzimidates previously discussed:
Products
Unlike the previous study of the acidic hydrolysis of benzohy-27droxamic acid, a small but measurable rate constant, reported in
Table VIII, was obtained for hydrolysis of ortho-methyl-N-methylbenzo-
hydroxamic acid in the absence of added salt or hydrochloric acid,
necessitating the addition of equation (29) to the mechanism. Although
not shown, this step depicts the formation of the tetrahedral inter
mediate via attack of water on the unprotonated hydroxamic acid. From
this mechanism, the rate law (which is consistent with the observed
data) may be given as:
intercept of k2 depicting the observed rate constant at ionic strength
0.751M in the absence of added hydrochloric acid. In reality, the
value of this rate constant is far below the extrapolated intercept of
Figure 2 (see Table VIII). Extrapolation of the data in Figure 2 to
zero hydrochloric acid concentration is unwarranted due to specific
ion effects'^ (see below) caused by a change in the cations present
from hydronium and potassium cations in the vicinity of the observed
data, to only potassium cations at the extrapolated intercept. A plot
of equation (30), therefore, will yield an intercept which does not
kobs * Kkl & 3 0+J + k2 (30)
versus jf will yield anThis equation implies that a plot of k ^ g
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57
relate to the true value of k S t e p represents the contribution
of equation (29) to the overall rate law resulting from hydrolysis
involving different charge types than that within the concentration
range of hydrochloric acid studied, although some small contribution
from equation (29) might exist within this concentration range. The
value of k£ at zero hydrochloric acid concentration can, therefore,
only be determined by direct measurement.
The first-order reaction of ortho-methyl-N-methylbenzohydroxamic
acid occurs, as discussed above, in the absence of added hydrochloric
acid, but in the presence of 0.751M potassium chloride. The pseudo
first-order rate constant for this reaction is about 6% of the observed
rate constant at 0.150M hydrochloric acid at ionic strength 0.751M.
These observed rate constants and those used in Figure 2 are reported
in Tables V and VIII.54Specific ion effects, analogous to those proposed for k£ in
equation (30), as well as ionic strength effects on reaction rates are
expected^ and were observed at moderate acid and salt concentrations.
Table VIII illustrates the effect of varying salt concentration and
salt ions on kobg. The data indicates that ionic strength effects,
specific cation and specific anion effects occur outside experimental
error.
Alkaline Hydrolysis Mechanism
The first-order dependence of the observed rate constant (i.e.,
kobs) on hydroxide ion concentration given in Table IV is illustrated
in Figure 3. Such a dependence is consistent with the bimolecular-
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2.60'
2.40-
2.20-
2.00-
1.80-
obs
1.60-
1.40-
1.20-
1.00-
3.4 4.2 5.0 5.8 6.6 7.4LNaOHj
Figure 3. Dependency of k ^ C s e c on catalytic base concentration (M)
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59
27tetrahedral mechanism proposed by Berndt and Fuller for the alkaline
hydrolysis of benzohydroxamic acid at more moderate basicity discussed
earlier. The first-order dependence is also consistent with the obser
vations of Ahmad and coworkers,^ previously discussed, at low basicity
(i.e., pH = ^ 8 to ^ 1 1), but is in sharp contrast to their conclusion
that kQks is independent of hydroxide ion concentration at bascities
greater than pHs?12. The observation of a first-order dependence of
kobs on hydroxide ion concentration for the hydrolysis of ortho-chloro-
N-methylbenzohydroxamic acid at high basicity and high ionic strength
provides additional support for the rate law proposed by Berndt and 27Fuller which has been previously discussed (i.e., see "Mechanisms
Background").
The reactions for the hydrolysis of ortho-substituted-N-methyl-
benzohydroxamic acids at high basicity and ionic strength are more
complex than in moderate acidic media. Pseudo first-order rates are
observed according to the equation:
“ dt kobsCSJ (31)where [Vj is the total stoichiometric amount of hydroxamic acid at
any time.
The consistency of the observed data with that obtained by Berndt
and Fuller at more moderate basicities (i.e., ^ 0 . 1 - 2M) suggests a
hydrolysis mechanism, for the presently studied, more hindered system,
in which formation of a tetrahedral intermediate may occur from
nucleophilic attack of either hydroxide ion or water on the hydroxamate
anion (i.e., analogous to scheme IX). However, the mechanism proposed
by Ahmad and coworkers at high pH^^ (i.e., mechanism B - scheme VIII)
neglects the potential of hydroxide ion as the attacking nucleophile
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27which is inconsistent with present and previously observed data.
Increased hindrance in the present system, caused by N-methylation,27over that in the previously studied benzohydroxamic acid series will
cause, however, considerable variations in the proposed tetrahedral-
bimolecular mechanism over that proposed by Berndt and Fuller. Although
observed data for the basic hydrolysis of ortho-chloro-N-methylbenzo-
hydroxamic acid is consistent with that observed for the previously27studied benzohydroxamic acid series, lack of N-hydrogens in the
present case prohibits the tautomerization equilibria proposed in the 27earlier case. Therefore, although analogous, the proposed tetrahedral-
bimolecular mechanism for the basic hydrolysis of ortho-chloro-N-
methylbenzohydroxamic acid will differ from that proposed in scheme IX
by the exclusion of three steps. The resulting proposed mechanism
which is consistent with observed data is illustrated in equations (32)
Under the strongly alkaline conditions used in the kinetic studies,
JL will be almost completely converted to 2 . (The pKa 's of N-tert.-
and 9.15 respectively.) The proposed mechanism illustrated in equations
(32) - C34), therefore, indicates nucleophilic attack by either
hydroxide ion or water on only the ionized form of the hydroxamic acid.
- (34);
Cl Cl
Products
Products
(32)
(33)
(34)
butylbenzohydroxamic and N-phenylbenzohydroxamic acids are 10.1
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61
Although not shown in equations (33) and (34), the involvement of
the tetrahedral intermediates consistent with observed data is probably
similar to that discussed for amides, esters (i.e., scheme II), and
for benzohydroxamic acid (i.e., scheme VIII) under alkaline conditions.
The rate law for this mechanism and for the hydroxide ion concen
trations employed, consistent with the observed data and with that in 27earlier work, is derived as follows (water concentrations are
included in the constants):
-dS/dt = f2j(COH7 k2 + k3)
= [ljKlOH^ ( [0H7k2 + k3) (35)
S = [l](l + [OH J K)
Therefore:
-dS/dt = SKfOHj ( C O H ^ k 2 + k3)/l + fOHjf K (36)
Under the reaction conditions employed, KpDHj >> 1 and:
kobs = k2C°H3 + k3 (37)where k , is the pseudo first-order rate constant. The form of obsequation (37) is consistent with Figure 3.
Specific salt effects'^ are expected at the high concentrations
employed to maintain constant ionic strength in the alkaline hydrolyses.
Table VIII illustrates the effect of varying salt concentration and
salt ions on at ionic strengths comparable to those employed in
the kinetic runs. The data indicates that ionic strength effects and
specific ion effects occur outside experimental error. Note that the
rate constants reported in Table VIII are for reactions in the absence
of any added hydroxide. In these cases, the reaction involved the
hydroxamic acid reacting with water rather than the conjugate base
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62
reacting with hydroxide ion or water and, therefore, involves different
charge types and thus a different mechanism. However, at the concen
trations of catalytic base employed in this study, there will be
specific salt and specific ion effects for all charge t y p e s . T h e r e
fore, any extrapolation of Figure 3 beyond the extremes of the observed
data is unwarranted due to both ionic strength and specific ion effects.
Proximity Effects for Acidic Hydrolysis
The values of the observed rate constants (i.e., k , ) for variousobsortho-substituents used in the correlation with the Taft-Pavelich
equation (i.e., equation (12)) are given in Table VI. The values of
p*, 6 and log kQ in equation (12) were determined by computer using the
method of multiple linear regression. Calculated values for log k for
each ortho-substituent were also determined by computer using this
method. Table X illustrates the comparison between the log of the
observed rate constant, used in determining values for p*, 6 and
log kQ , and the log of the calculated rate constant, determined by2 57computer from a least squares treatment ’ of the observed data, for
each ortho-substituent.
The correlation of reactivity in the hydrolysis reaction by
equation (12) is illustrated in Figure 4. The values for the reaction
constants resulting from the multiple regression analysis are p* =
-0.688 and 6 = 0,278. The statistical significance of the correlation
of the observed data by equation (12) is measured by the coefficient
of multiple regression (i.e., correlation coefficient). ^ ^ This
value is the square root of the ratio of the explained variation to
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63
Table X
Observed and Calculated Rate Constants for Acidic Hydrolysis ofortho-Substituted-N-Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C
Ortho- , , log k, (sec 1) log k, (secSubstituent a * Eg (observed)3 (calculated)
"CH3 0 0 -4.359 -4.403
-°ch3 -0.22 0,99 -3.964 -3.977
-no2 0.97 -0.75 -5.237 -5.279
-Cl 0.37 0.18 .-4.585 -4.608
-Br 0.38 0 -4.714 -4.665
-I 0.38 -0.20 -4.795 -4.720
aAverage of two to three runs from Table VI. ^Determined from equation (12) and calculated values of p* and 6 (see below).
57 58the total variation of the experimental data. 5 For the present
system studied, the correlation coefficient (R) was calculated to be
0.989 (for a perfect correlation, R = 1.00).
The reliability of the correlation coefficient as a measure of
statistical significance depends on the number of data sets used in
the multiple linear regression analysis and on the number of variables
calculated. The F - t e s t ^ ’^ allows for these factors and, in the
present case, indicated the correlation to be significant within the
1% level (i.e., a very high level of significance can be attributed to
the correlation).
Correlations of reactivity in the hydrolysis reaction were
attempted by both equation (10) and (11). Table XI illustrates the
comparison of the log of the observed rate constant (i.e., log k0jjS),
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64
-5.00-
80-
■4.60-
-4.40-
-4.20 1.100.900.700.500.30- 0,10 0.10-0.300*
Figure 4. Correlation of log kQbs with a* and Es values for acidic hydrolysis
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65
used in determining the value of p* for a correlation with equation (11)
or 6 for a correlation with equation (10), with the log of the
calculated rate constant (i.e., log ^ca^c ), determined by linear
regression from the correlation of the observed data with either a *
(i.e., equation (11)) or Eg (i.e., equation (10)), for each ortho
substituent.
Table XX
Comparison of Observed Rate Constants with Rate Constants Calculated from Equations (10) and (11) for the Acidic Hydrolysis of Ortho- Substituted-N-Methylbenzohydroxaraic Acids in 0.764 M HC1 at 90.0°C
Ortho- a log ka (calculated) ^Substituent log k (observed) equation (10) equation (11)
-CH3 -4.359 -4.635 -4.285
-och3 -3.964 -3.926 -4.058
-N02 -5.237 -5.172 -5.287
-Cl -4.585 -4.506 -4.668
-Br -4.714 -4.635 -4.678
-I -4.795 -4.778. -4.678
aPseudo first-order rate constant in sec ^Determined fromcalculated value of 6 = 0.716. determined from calculated value of p* = -1.032.
For the correlation of the observed rate constants with a * values
alone (i.e., equation (11)), it was found that R = 0.974. The F-test
showed this value of R to be significant at the 1% level. However,
this correlation is poorer than that with a * and Eg values together
(i.e., equation (12) and Table X) as illustrated by not only a com
parison of R-values and their significance levels, but also by either
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66
a comparison of residuals (i.e., | log k ^ g - log kca^c l) for each
ortho-substituent, or a comparison of the average residual for each
correlation (i.e., 0.075 for equation (11) and 0.040 for equation (12)).
Since k , = k , for a perfect correlation, how well the equationobs calc.correlates the data can be measured by a determination of residuals in
addition to the correlation coefficient tests.
For the correlation of the observed rate constants with Eg values
alone (i.e., equation (10)), it was found that R = 0.934. The F-test
showed this correlation to be significant at the 1% level. However,
as with the correlation with a * values alone, this correlation is poorer
than that with a* and E values together (i.e., R = 0.989 with signi- sficance at the 1% level). Examination of values of log kca^c ,
determined by a least squares treatment of the observed data illustrated
in Table XI, shows that a correlation with E values alone does notsreproduce the value of log kQ within acceptable limits. Reproduction
of the value of log kQ (i.e., one of the constants which is calculated
directly in the least squares treatment) within acceptable limits
(i.e., typically - 15% as a median precision of rate or equilibrium4 59
constants for the Hammett equation) is necessary if a correlation
between observed rate data and Taft substituent parameters is to be
considered valid or significant. It is clear, therefore, that the
correlation which best relates the observed rate data for the acidic
hydrolysis of ortho-substituted-N-methylbenzohydroxamic acids is that
involving both a* and Eg values (i.e., equation (12)).
The results reported here for the correlation of the rate data
for the acid catalyzed hydrolysis of ortho-substituted-N-methylbenzo-
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67
hydroxamic acids with the polar and steric substituent parameters
(i.e., a* and Eg respectively) in equation (12), taken together with2 45a similar correlation reported earlier for the acid catalyzed
hydrolysis of ortho-substituted benzohydroxamic acids, lend support to22the qualitative description by McCoy and Riecke for the steric effect
of ortho-substituents and to the usefulness of equation (12) as a
first approximation to a quantitative approach for the general descrip
tion of the ortho-effect, which was discussed earlier.
The contention that steric effects do exist in the present system
(i.e., N-CH^) is supported by three pieces of evidence. First, X-ray
data, reviewed by M. Charton, showed a "twisting" of the “CC^H group
out of planarity with the phenyl ring in ortho-substituted benzoic
acids . ^ Charton suggested that the values for the interplanar angle
between the ortho-substituted phenyl ring and the carboxyl group (n)
is a function of the van der Waals radius of the ortho-substituent.
This implies the presence of a steric inhibition to resonance (i.e.,
the secondary steric effect) caused by the ortho-substituent as proposed 22by McCoy and Riecke, which was discussed earlier. Secondly, the
correlation obtained for the acid catalyzed hydrolysis of ortho
substituted benzohydroxamic acids showed nearly equal dependence of
the reaction system on polar effects (i.e., a * ) and steric effects
(i.e., Eg) as defined by Taft (i.e., p* = -0.87 and 6 = 0.76).2 *̂ "*
Thirdly, the construction of space-filling models shows that the ortho-
substituted-N-methylbenzohydroxamic acids are more conformationally
restricted than the corresponding ortho-substituted benzohydroxamic
acids implying the presence of steric effects in the presently studied system.
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The relative dependence of the present system (i.e., N-CH^) on
steric effects in comparison to that for the previously studied
ortho-substituted benzohydroxamic acid system^’ "̂* (i.e., N-H) may be
illustrated by a comparison of the absolute values of the ratio of
6/p* for the two reaction systems. The absolute value of this ratio2,45for the less hindered N-H system was found to be 0.874, while that
for the more hindered present system is 0.404. These values imply
that the steric effects, as measured by 6Eg (see below), are smaller
in magnitude in the N-CH^ system than in the less hindered, N-H
s y s t e m . Y e t , as discussed above, steric effects, as a function of
conformational restriction and steric inhibition to resonance in the
reactant state, are greater in the N-CH^ system than in the less 2 45hindered, N-H system. * The differing values for these ratios are,
nonetheless, consistent with the qualitative conclusions of McCoy and 22Riecke. Their graphical illustration of the total ortho-substituent
steric effect as a function of substituent bulk and as a combination
of primary and secondary effects (i.e., equation (19)) is a function
of the basic structural skeleton of the reaction system in addition to
being a function of the substituent, solvent and the reaction type
(i.e., nucleophilic substitution versus acid ionization, for example).
Therefore, for the presently studied N-CH^ system, the extent of con
tributions from terms in equation (19) as well as the shape and slopes
of graphical representation of the total steric effect of an ortho
substituent are expected to differ from those for the similar, but
structurally different N-H system under similar reaction, solvent and
temperature conditions.
Reproduced with permission of the copyright owner. Further reproduction prohibited without permission.
69The consistency of the differing values of 6/p* for the N-CH^
and N-H systems with the conclusions of McCoy and Riecke above does
not, however, explain the observation that, for reactions performed
using identical temperatures and solvents at similar ionic strength
and hydrochloric acid concentrations, the value of 6 is much lower
for the more hindered N-CH^ system than for the less hindered N-H
system (i.e., SCN-CH^) = 0.278 and 6 (N-H) = 0.76). However, the logic
used by McCoy and Riecke in their description of the total steric
effect of an ortho-substituent as the combination of primary and
secondary effects does suggest a possible explanation for this obser
vation. The reaction's susceptibility to steric effects, 6 , may be
considered as a combination of the susceptibility to primary steric
effects and to steric hindrance to resonance in the reactant state 22(i.e., the secondary effect). The lower value of 6 for the N-CH^
system may be due to a lower susceptibility of the reaction system to
the secondary steric effect. This may occur in the N~CH^ system from
an increase in the interplanar angle between the carbonyl and ortho
substituted phenyl groups (p), as discussed by Charton,^ over that in
the N-H system due to the presence of the larger N-methyl group in the
reactant state. The larger value of p for the N-CH^ system in the
reactant state lessens the amount of resonance interaction between the
carbonyl group and the ortho-substituted phenyl ring which, in turn,
lessens the susceptibility of system to resonance interaction effects
of ortho-substituents.
Three pieces of evidence indicate that, for the presently studied
N-CHj system, Eg is a nearly true measure of a substituent's steric
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70
effect and 6Eg is a good measure of the total relative steric effect
of an ortho-substituent on the hydrolysis rate constant. First, as
discussed earlier in relation to equation (5), Taft found that, for
ortho-substituted aromatic systems, the Eg values for symmetrical top
substituents (i.e., X, CH^, t-Bu.) roughly paralleled their van der
Waals radii. Secondly, in Taft's definition of Eg values for substi
tuents (i.e., equation (4)), as applied to systems similar to that
presently studied, there is a resonance contribution to Eg from those
substituents which can, by electron release, exhibit direct resonance
interaction with the carbonyl reaction site. The extent of such a
resonance contribution to Eg values for the ortho-substituents employed
in this work can be determined by the extent of a similar resonance
contribution to crT, ... (para) values which is of general f o r m : ^ Hammett
* q L * ------- ^ X = 0 = i - G (38)
E x n e r , ^ Jaffe"^ and Taft^’^ have recorded values of a . (para)’ Hammett(i.e., Op) and a° (i.e., "insulated" values in which resonance
interaction with the carbonyl reaction site is prevented by inter
position of a -CH2 group) for some para-substituents. These values
are given in Table XII. As the table indicates, the apparent resonance
contribution (i.e., - a°) of the type shown in equation (38) to
values by all para-substituents, save -OCH^, is extremely small imply
ing a similarly negligible resonance contribution by ortho-substituents
to their E values, sLastly, the resonance contribution of the ortho-OCHj group to its
Eg value, which is similar to the resonance contribution of the para-
OCHg group to its op value given in Table XII, may, in reality, be
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71
Table XII
Comparison of o ° and a Values for Para-Substituents of Benzenl Derivltives in Aqueous Media^
Substituent a ° aP P
-CH3 -0.15 -0.17
-OCH3 -0.12 -0.27
-N°2 0.82 0.78
-Cl 0.27 0.23
-Br 0.26 0.23
-I 0.27 0.27
smaller in magnitude than indicated in Table XII. Taft has shown, for4
the saponification of ethyl para-dimethylaminobenzoate, that only
part of the resonance interaction of the para-(CHj) group with the
carbonyl group is lost in going from the ester to the saponification
transition state (i.e., "... the resonance energy for the interaction
of the para-dimethylamino and carbethoxyphenyl groups has been cal
culated to be about 8 Kcal/mole.^ Yet, the activation energy for the
saponification rate was decreased by only about 2.5 Kcal/mole when4
this resonance was nearly completely destroyed by steric inhibition.") .
The difference in these two numbers, therefore, represents that part of
the total resonance energy which is the same in the transition state
as in the reactant state. On this basis the resonance contribution of
the ortho-OCH^ group to its Eg value is probably considerably smaller
than that indicated for the polar constants in Table XII. Therefore,
it is clear from the above evidence, that Eg values for the ortho
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72
substituents employed in the presently studied N-CH^ system are good
measures of the actual steric effect. In addition, the fact that a
poorer correlation exists for equation (10) than for equation (12) for
the N-CH^ system is further support for Taft's contention that 6Eg and
thus E values contain no polar contributions. This contention, s -----although in sharp contrast to Charton's claim as illustrated in
equation (7), is consistent with the qualitative description of the22ortho-steric effect by McCoy and Riecke.
The value of p* obtained in this study (i.e., p* = -0.688) is an
overall value (i.e., a composite of p* values for equations (27) - (29)
in the mechanism although, as previously discussed, the contribution
to the overall rate of equation (29) is probably very small). Since
p* < 0 in equation (12), electron donating groups accelerate the rate
compared to that of the reference compound, ortho-methyl-N-methylbenzo-
hydroxamic acid. This overall value of p* is consistent with the
bimolecular mechanism proposed in equations (27) - (29) and with the29general hydrolysis mechanism for hydroxamic acids by Buglass et. al.
27and by Berndt and Fuller discussed earlier. The difference between2,45the p* values for the N-CH^ system and the N-H system (i.e.,
p*(N-H) = -0.868) can be compared to the difference in the p* values
for the N-H system2,^ and the ortho-substituted benzamide system 4(i.e., p*(amide)^r-0) under similar temperature and reaction conditions.
For the latter difference, the negative p* value is consistent with
the greater electronegativity of N-hydroxyl compared to N-hydrogen in
changing from amides to hydroxamic acids, provided that the polar effect
on the protonation step in equation (27), which is enhanced by electron
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73
donating groups, is greater than that for the nucleophilic attack by
water on the protonated intermediate (i.e., equation (28)). For the
former difference, substitution of an electron donating -CH^ group,
relative to -H, at the nitrogen will reverse, to some extent, the
polar effect for the protonation step observed when the N-hydrogen of
benzamide was replaced by the electronegative N-hydroxyl to yield the
hydroxamic acid. The overall value of p*, therefore, should be less
negative for the N-CH^ system than for the corresponding N-H system as
a result of a decrease in sensitivity toward polar effects in the
protonation step (i.e., equation (27)).
Proximity Effects for Alkaline Hydrolysis
The values of the observed rate constants, k , , for variousobsortho-substituents used in the correlation with the substituent para
meters a * and Eg (i.e., equation (12)) are given in Table VII. The
values of p*, 6 and log kQ in equation (12) were determined by the
method of multiple linear regression . ^ Calculated values of log k
for each ortho-substituent were also determined by this method. Table
XIII illustrates the comparison between the log of the observed rate
constant, used in determining values for p*, 6 and log kQ , and the log
of the calculated rate constant, determined by a least squares treat- 2 57ment * of the observed data, for each ortho-substituent. In this
table, log kQks for ortho-NO^ is omitted from the correlation. The
reasons for this omission are two-fold. First, inclusion of the value
of the log kQbg for ortho-NO^ results in a very poor correlation which
does not reproduce the values of log kQbg within the acceptable
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74
Table XIII
Observed and Calculated Rate Constants for Alkaline Hydrolysis ofOrtho-Substituted-N-Methylbenzohydroxamic Acids in 7.31M NaOH at 90.0°
Ortho- log kSubstituent a * Eg (observed) ’ log k (calc.) *
-CH3 0 0 -6.130 -6.178
-°CH3 -0.22 0.99 -5.499 -5.466
-Cl 0.37 0.18 -5.585 -5.695
-Br 0.38 0 -5.915 -5.850
-I 0.38 -0.20 -6.093 -6.032
aPseudo first-order rate constant in sec . Average of two runs from Table VII. Determined from equation (12) and calculated values of p* and 6 (see below).
4,59limits discussed earlier. Secondly, as described in the Experi
mental Method, Apparatus and Syntheses" section, product analysis was
inconclusive and could not help in establishing the mode of reaction
for the alkaline hydrolysis of ortho-nitro-N-methylbenzohydroxamic
acid. It may well be that a different mechanism or combination of
mechanisms from that governing the hydrolyses of the other compounds
is occurring. It is not, therefore, inconsistent or arbitrary to omit
the value of log k ^ g for the hydrolysis of this compound from the
correlation.
The correlation of reactivity by equation (12) is illustrated in
Figure 5. The values of the reaction constants resulting from the
multiple regression analyses are p* = 0.863 and S = 0.911. For this
system, the correlation coefficient (R)^,^ ,‘*8 was calculated to be
0.928. The F-test"5̂ ’58 showed this correlation to be significant
Reproduced with permission of the copyright owner. Further reproduction prohibited without permission.
log
(k75
-6.80-
-6.4C
-5.60
0.4 0.6 1.2-0.2 0.0 0.2 1.0E.s
Figure 5. Correlation of log kQ^s with a * and Eg values for alkaline hydrolysis
Reproduced with permission of the copyright owner. Further reproduction prohibited without permission.
76
nearly within the 5% level. These results imply that a fair correlation
exists between values of log k ^ g and the corresponding substituent
parameters a * and Eg . The statistical level of significance of this
correlation may be considered good.
Although poorer than the corresponding correlation for the acidic
hydrolysis at a slightly less signigicance level, the correlation of
log f°r this system by equation (12) is significant. As Figure 5
illustrates, the correlation does reproduce the trend of the data to4 59a median precision within acceptable limits. * Further, extremely
high ionic strength, large salt and specific ion effects may contribute,
as Tables VIII and XIII indicate, to marked deviations in the hydrolysis
rate constants for various ortho-substituents. Much larger catalytic
hydroxide ion and salt concentrations than those previously 2 3 26-29 45 59studied ’ * ’ ’ may also produce different charge types (i.e.,
the neutral substrate reacting with water) from those in earlier studies
which may adversely affect certain steps in the proposed hydrolysis
mechanism described in equations (32) - (34) decreasing the signifi
cance or validity of the correlation. The determinate reaction of
ortho-nitro-N-methylbenzohydroxamic acid under these conditions may be
an example of such adverse affects.
That a correlation for the alkaline hydrolysis of the N-CH^
system does exist suggests that the sensitivity of steps k2 and k3 in
equations (33) and (34) to polar and steric substituent effects are
proportional. This arises since the correlation with log kQbg is
actually a correlation with log (k2 + k^ CoH^ ) (see equation (37)).
Contributions to kQbg from k 2 and k3 will vary with each ortho
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77
substituent due to varying substituent effects. If the susceptibility
of each step to substituent effects is proportional, then a correlation
with log k0fcs is possible. However, if the sensitivity to substituent
effects were to vary randomly from steps k2 to k^, then, coupled with
the variance in contribution to kobg from steps k2 and k^, a correlation
between log kQbs and the substituent parameters in equation (12) may
not be possible.
Correlations of reactivity in the hydrolysis reaction were attempted
by both equations (10) and (11). An extremely poor correlation was
obtained with equation (10) yielding a value of the correlation co
efficient R = 0.743. With equation (11), no correlation at all was
obtained.
The positive value of p* obtained for the correlation of log kQbg
with o * and Eg values (i.e., equation (12)) is consistent with the
bimolecular mechanism proposed in equations (32) - (34) and with that27proposed by Berndt and Fuller in an earlier work. Since p* is
actually a composite of p values for the steps in the mechanism, it is
consistent with equations (32) - (34) that electron withdrawing sub
stituents accelerate the rate for each step in the mechanism relative
to that of the reference compound, ortho-methyl-N-methylbenzohydroxamic
acid.
The positive value of 6 means that the rate of hydrolysis is
decelerated as E becomes smaller. Decreasing values of E for ortho- s s ------substituents presumably correspond to increasingly effective steric
bulk,^,10,'I''1',1^ ’22,^0 although, for the present system, a very small
resonance contribution may be present. This assessment of the steric
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78
effect of an ortho-substituent for the presently studied system, as
6E , is consistent with the observed data in Table XIII. It shows the slarger substituents, the electron withdrawing effects being constant,
exhibiting slower hydrolysis rates. The proportional decrease in
log kQks with increasing substituent bulk for these ortho-halo sub
stituents is also qualitatively consistent with the conclusions of
Taft,^ McCoy and Riecke22 that Eg values for ortho-substituents contain
no polar contributions.
Finally, it should be noted that although the relative dependence
of the N-CHg system on steric effects (i.e., <S/p*) for the alkaline
hydrolysis is greater than that for the acidic hydrolysis, a direct
comparison of these dependencies is not possible since they each involve
different mechanisms, charge types, catalytic solution, solvent concen
trations and ionic strength effects. However, both hydrolysis systems
do lend further support to some contentions about some heretofore
debated points. First, the success of the application of the Taft-
Pavelich equation to both hydrolysis systems provides further evidence
for Taft’s assumption that substituent effects may be treated as an
independent sum of steric, resonance and polar contributions. Further,
the successful application of a * and Eg to these hydrolyses and other
reaction types is evidence that substituent effects are functions of
only the substituent and do not vary with the reaction system. This
contrasts the conclusion of Charton as illustrated in equations (7)
and (9) in which, he claims, the value for "h" varies with the
reaction system. Lastly, as discussed in this and the previous
section, the present study has provided evidence that Eg values, in
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79
agreement with the conclusions of Taft^ and McCoy and Riecke,^ are
good measures of the actual steric effect of ortho-substituents and,
at least for the substituents studied (except perhaps -OCH^),
contain only a very small resonance contribution.
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BIBLIOGRAPHY
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VITA
The author was born in Rushville, Indiana on August 20, 1949. He
graduated from Maine Township High School South in 1967 and entered
Rose Polytechnic Institute that same year. He received his B. S. degree
in chemistry in 1971 and entered Western Michigan University as a grad
uate student later that year with an appointment as a graduate teaching
assistant. While completing the masters program in organic chemistry,
the author was also employed as a part time gas chromatographic tech
nician and research chemist at the A. M. Todd Company. He worked as a
film chemist for the E. I. DuPont Company during 1974 upon completion
of the M. A. degree. He returned to W. M. U. in 1975 and entered the
Ph.D. program as a graduate teaching assistant. In 1976, he was awarded
a graduate associateship for support while completing the Ph.D. degree
program.
The author is married with no children.
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