Proximity Effects. Kinetics, Mechanisms and Reactivity Correlations for the Acidic and Alkaline Hydrolysis of Ortho- Substituted-N-Methylbenzohydroxamic Acids

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Dissertations Graduate College

8-1977

Irl E. Ward Western Michigan University

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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

This Dissertation-Open Access is brought to you for free and open access by the Graduate College at ScholarWorks at WMU. It has been accepted for inclusion in Dissertations by an authorized administrator of ScholarWorks at WMU. For more information, please contact [email protected]. PROXIMITY EFFECTS. KINETICS, MECHANISMS AND REACTIVITY CORRELATIONS FOR THE ACIDIC AND ALKALINE HYDROLYSIS OF o r t h q - substituted -n -methylbenzohydroxamic ACIDS

Irl E. Ward

A Dissertation Submitted to the Faculty of the Graduate College in partial fulfillment of the Degree of Doctor of Philosophy

Western Michigan University Kalamazoo, Michigan August 1977

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 E s values 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

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Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. TABLE OF CONTENTS

CHAPTER PAGE

I INTRODUCTION...... 1

Substituent Constants and the Ortho-Effect ...... I

Mechanism Background...... 20

Purp o s e ...... 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 Kinetics Procedure...... 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...... 80

V V I T A ...... 83

iii

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. LIST OF TABLES

TABLE PAGE

I Yields of Prepared Ortho-Substituted Benzoyl Chlorides...... 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 ofQ 2-Chloro-N-Methylbenzohydroxamic Acid at 90.0 C 43

V Rate Constants for Acid Catalyzed Hydrolysis of 2-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 Uncatalyzed Hydrolysis 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 Acidic Hydrolysis of Ortho-Substituted-N-Methylbenzo- hydroxamic Acids in 0.764M HC1 at 90.0 C 63

XI Comparison of Observed Rate Constants with Rate Constants 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 Alkaline Hydrolysis of Ortho-Substituted-N-Methylbenzo- hydroxamic Acids in 7.31M NaOH at 90.0 C 74

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. LIST OF FIGURES

FIGURE PAGE

1 Experimental Apparatus ...... 49

2 Dependency of k0us(sec_1) on Catalytic Acid Concentration (S)...... 54

3 Dependency of k , (sec"1) on Catalytic Base Concentration0^ ) ...... 58

4 Correlation of log with a * and Es Values for Acidic Hydrolysis...... 64

5 Correlation of log k bg with o* and Eg Values for Alkaline Hydro?ysis...... 75

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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)

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 K q 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. for all substituents and depends only on the reaction series, temperature

and solvent. 2,3 The 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 the 3 functional group undergoing reaction, or when substituent proximity, as

for ortho-substituents, introduced steric effects not accounted for by

2,3 Hammett's polar substituent constant, a.

The failure of the Hammett equation for ortho-substituents was

A 5 studied 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 polar i and 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:^’^

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 3

a* = [log (k/k0)B - log (k/ko)AJ/2.48 (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 subscripts CH 3 B 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 b e low).

The validity of equation (3) as a measure of the polar effect of

2,4 a 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 ^para

meta- 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

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

Resonance effects on the rate

constants must Option (3) is based upon

hydrolysis rat< B.ra-substituted benzoates.

Recorded value.- V i e between -0.2 and +0.5

for the acid can rrost cases, nearly zero.^

This is in contrast" W nd para) va'lues for the base

catalyzed hydrolysis whici^^ u.y lie within the range of +2.2 to 2.8.

Hammett studies have shown that, for systems which obey equations

and that differences in reactivity between (1) or (2), p}meta^ ppara

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:

log (k/kQ )A = 6Eg (4)

6 represents the reaction's susceptibility to steric effects and is

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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.,

Eg°) 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 contribution A to E . Equation (4) has been found to correlate data for several

4 8 9 reaction systems. * ’ This provides experimental evidence for the

validity of the above premise (i.e., the equality of p values) as applied

4 8 9 to 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 unimolecular 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 E s (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-substituted 4 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,

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 s 12 works 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 represents 12 the 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. of a substituents' steric effect. However, Taft also used equation (5)

to define E as a steric substituent constant from ortho-substituted s

benzoate hydrolysis as well. Charton disagreed with this definition

and represented the Eg (ortho) values in terms of a combination of

11,14 contributions from inductive, resonance and steric effects given as:

« 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

substit 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 as:^

(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,x

applicable to a variety of reaction systems.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 of 4 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-substituted 4 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*#pTT .. (para) for the described o Hammett

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 equation o 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 O

For 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 .14,20,21 (9) and represented it as:

Qx ' * ' 0 I , X + S '0R,X + 'l''7,x + h (l5)

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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, a Ko~KHammett vr s

description 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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- 22 substituted 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 set".^ 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).

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 attempts 0.x

to 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 approach ing 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 unimolecular 22 acid ionizations and bimolecular substitutions.

Steric effects in these systems are represented by a graph,

according to McCoy and Riecke, relating the total steric effect of 22 ortho-G with increasing substituent volume:

•rl a) 4J

cd o H

Increasing Size of Substituent

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

22 the 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 16

again would change, "but probably to a lesser rate of change with

22 increasing 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 are typically of the same magnitude as other contributions to Q , a or E , leads to a misleading correlation x o,x s,x between 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)

22 methanol: 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 20 w 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 fy&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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 20

Mechanism Background

The hydrolysis mechanisms of carboxylic acid derivatives such as

esters, imidates, amides, hydroxamic acids, etc. have been reported to

23-29 vary 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 to 24 hydrolyze 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 acidity

18 23-25 30-32 has been supported by 0 -exchange studies, ’ product

23 24 27 28 27— 29 33 analyses, 5 * ’ and kinetic data * and is generally repre

sented by Schemes I and II:

0 OH r °h h k (H 0) 'jj II + K i k l 1 --- R-C-OH. R-C-G + H30 — — R-£-G + H 2° R-C-G + 0H >11J f jr o OH

R-C-&H r - c -£ h

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 36 type mechanism for this system. ’ ’ There was also some confusion

as to the position of protonation for the acidic hydrolysis of

23 35 36 18 amides. ’ ’ 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 amide.However,

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 ?R of benzimidates (i.e., R'-C = NR’’) 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 that:^’^

1. There is a change in the hydrolysis mechanism from benzamide to N-methyl- and N, N-dimethylbenzamide.

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 24 carboxylic acids. It was not inconsistent, therefore, that successive

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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. 32 In 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

T the Aq2 type mechanism by McClelland. Such a mechanism is illustrated

in scheme III:

Scheme III

OH OH + ?H 2 I I r -£-n h R-C-NH0 • R-C-NH„

+ o 1bh 2 ’ i - H 2 k ‘

OH E t - c - f o i ,

McClelland argued that this mechanism implies, depending on the size of

He noted that, although previously suggested,28’8^ ’88 if k, /k h e

large 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 0^"8-exchange. The fact that

a small amount of 0^8-exchange was observed simply implies that the

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

T benzamide 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 the

28 benzimidates showed no evidence for a change in mechanism.

38 39 Changes in mechanism also occur with changes in pH. McClelland

and Smith and Yates^ 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 products

28 40 18 exclusively. * As expected with such a mechanism, extensive 0 -

exchange was observed. At increased acidity (i.e., pH < 1), the break-

28 down pathway of the tetrahedral intermediate was reportedly changed.

Products obtained from the oxygen-labeled imidate hydrolysis included a

significant concentration of starting imidate, unlabeled carboxylic

28 acid, 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 “§ 18r cation in the

tetrahedral intermediate, becomes more pronounced yielding the above

mentioned products following intermediate decomposition. The competition

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 25

9ieR ?18r ,,k, (H_o) ^ 918 | r R'-C=NR" + H o0t k ^ R ’-C=Sh R m + H20 - - -- - R ’-C-NHR" I 5 OH I High JF Low Acidity______1 Acidity + 0 18R

-C-^H2R"I + H20 r ,'- c1-n h r " OH OH

0 0 II II R’-C + NH R" R’-C-NHR N n 1 b r i n U R + R 0 18H

38 39 At very high acidity (i.e., pH < -1, -^65% H2SO^) 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 . ^

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 26

0 18R o 18r

R-C=NR" + H_0+ ---- ^ R-C=5h R " + H.O 3

& High Acidity Super High j [•' (H20)jl Acidity

0 18R' O 18 1 + n R-C-NHR" + H„0 R-C-NHR I OH

cf. scheme III

R - C-NHR1 + R't)18H

Changes in mechanism with changing pH have also been reported for

carboxylate esters,amides^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:

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 27

Scheme VI

OH I + K R-C-NHOH + H30 '^IZ± R-C-NHOH +OH, 2 w OH I H R-C-NH„OH \ 3 'OH I 2 OH

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-

45 substituted 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 Vecera^ 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

0 0 II II Ph-C-NHOH + OH Ph-C-NHO + H20

?" ?' ------Ph-C-NHOH ■ — Ph-C-NH.OH I.. -- 1 2

Scheme VIII

(A) At lower pH (i.e., 7 to 9):

0 0 II k i Ph-C-NHOH + H„0 ------Ph-C-NHOH ^ Ph-C-fe0OH J ^---- 1 + 0 H 2 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 - - Ph-C-NH0_ +0H„ II 0 |Ph-C-SH„OHI 2

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 to— ?,

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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(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)

k obs = (k l + k 2 (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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. of benzohydroxamic acid at pH > 13 is in contrast to earlier work

mechanisms contributing to at hydroxide ion concentrations

n . k 2 R-C-NHOH + OH

(A) (D)

K - it. " NC

(E)

k 2 D + H O — Products

k 3 E + H„0 -----> Products

- k4 D + OH -----> Products

- k 5 E + OH — Products

27 Their derived rate law from this scheme is -given as equation (21): k2K2 + k3K3 + [k4K2 + k5KJ [0H"J kobs - ^ (21) l/fOH"J + K

studied, K 2 and K 3 >> 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 above

27 pH£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 is 26,27,29 more hindered than the previously studied benzohydroxamic and

45 ortho-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 both T acidic and basic media, via a tetrahedral type mechanism (i.e., A q 2)

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 33

it is of interest to determine if, as with the N-methylated OQ OO OA amides * * ’ discussed above, the hydrolysis mechanism changes

with increased hindrance, or if, as with the N-methylated benzimi-

28 36 40 dates ’ ’ the mechanism remains unchanged.

Secondly, it has been shown that the hydrolysis mechanism for

benzohydroxamic acid changes with pH.^ However, the general rate

law of Ahmad and coworkers^ is in contrast to that of earlier work 27 by 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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).

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 general

46 procedure, except for the preparation of ortho-nitrobenzoyl chloride

which will be described separately. Such a general procedure is

47 illustrated by the following synthesis of ortho-methylbenzoyl chloride

and is, therefore, also applicable to the preparation of the ortho-

48 48 49 chloro, -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 52 chloride, * 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 chloride.The 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 Chlorides n 1^ ______Substituent Yield,_Percent______

71.0 -CH3

-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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

17.6 117.0 - 118.0 “ CH3

- 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. ^Compound observed to decompose on melting.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 39

Table III

Elemental Analyses of Prepared Ortho-Substltuted- N-Methylbenzohydroxamic Acids

Substituent______Analysis3______%_C______% N______% H

-CH_ Theoretical 65.45 8.48 6.67 3 Observed 65.14 8.47 6.48

-0CH„ Theoretical 59.67 7.73 6.08 3 Observed 59.52 7.64 6.14

-Cl Theoretical 51.76 7.55 4.31 Observed 51.61 7.59 4.38

-Br Theoretical 41.76 6.09 3.50 Observed 41.97 5.89 3.36

-I Theoretical 34.68 5.05 2.91 Observed 34.87 4.93 3.07

-N0„ Theoretical 48.97 14.28 4.12 I 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 reactions

45 at 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

45 for 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,

43 Beer'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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 . a % Mean Deviation obs

7.31 1 2.53 2 2.56 3 2.61+ 4 2.59-(0.006) Average 2.57 1.07

6.58 1 2.23. 2 2.32-(0.003) Average 2.27 1.98

5.47 1 1.81 2 1.92 3 1.90-(0.008) 4 1.93 Average 1.89 2.12

4.40 1 1.52-(0.007) 2 1.42 3 1.47 4 1.54 5 1.39 Average 1.44 3.89

3.23 1 0.89 2 0.94-(0.001) Average 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. Table V

Rate Constants for Acid Catalyzed Hydrolysis of 2-Methyl-N-Methylbenzohydroxamic Acid

HC1, M Trial Temp, °Ca 1()5kobsb % Mean Deviation

0.149 1 1.45 2 1.51-(0.009) Average 90.0 1.48 2.03

0.225 1 1.96^(0.02) 2 2.12 Average 90.0 2.04 3.92

0.451 1 3.16 2 2.94 3 2.89 4 2.97 Average 90.0 2.99 2.84

0.595 1 3.85 2 3.70-(0.04) Average 90.0 3.78 1.98

0.751 1 4.62 2 4.60 Average 90.0 4.61 0.22

0.751 1 1.81 2 1.90 Average 80.0 1.86 2.42

0.751 1 0.814 2 0.821 Average 70.0 0.818 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 45

Table VI

Rate Constants for Hydrolysis of Ortho-Substituted- N-Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C

Substituent Trial Z Mean Deviation 1()5k obsw 3 :

-CH3b 1 4 *42+ 2 4.30-(0.02) 3 4.41 Average 4.38 1.14

-OCH 1 10.91 2 10.71 3 10.95 Average 10.86 0.890

-Cl 1 2.55 2 2.62 3 2.62 Average 2.60 1.15

-Br 1 1.93 2 1.95 3 1.92 Average 1.93 0.518

-I 1 1.59-(0.02)° 2 1.62 Average 1.605 0.935

1 0-586 "N 0 2 2 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 Deviation obs

- c h 3 1 0.723 2 0.759 Average 0.741 2.43

-OCH- 1 3.21-(0.002) 2 3.13 Average 3.17 1.26

-Cl 1 2.59 2 2.61 Average 2.60 0.385

-Br 1 1.18-(0.002) 2 1.25 Average 1.215 2.88

-I 1 0.816 2 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 . - represents typical values for one standard deviation as determined by least squares.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 , a Strength, M °

0 0 1 0.360^(0.004) -°H3 - 2 0.440 Average 0.400

0 KC1 0.751 1 0.755 “CH3 2 0.782 Average 0.773

0.149 KC1 0.751 1 14.52 “CH3 2 15.14^(0.009) Average 14.83

0.150 CsCl 0.751 1 13.33 "CH3 2 13.23 Average 13.28

-Cl 0 NaCl 3.00 1 1.13 2 1.09 Average 1.11

-Cl 0 NaCl 6.31 1 2.00 2 1.89^(0.002) Average 1.94

-Cl 0 NaBr 6.31 1 1.33 2 1.34 Average 1.33

Pseudo first order rate constant in sec . - represents typical values for one standard deviation as determined by least squares.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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,

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 49

?o Figure Figure 1. Experimental Apparatus

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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, and Aq = 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 A t

Since A q 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 A t = 2.303" + 108 A o (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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

52 spectrum 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).

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

52 spectrum. 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 again

52 similar, in some aspects, to the standard spectrum, but an exact match

was not made. The precipitate, therefore, could not be conclusively

identified.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 obs

is 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 and

2 27 29 45 hydrolysis temperatures. These previous studies also showed ’ ’

that such a dependence is consistent with a bimolecular mechanism

involving a tetrahedral intermediate analogous to those supported for

32 28 38 the 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 less 27 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 amides^ 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 54

.60-

3.80-

obs

3.00-

1,40 o.io oTio O'. 50 0.70 (hciJ Figure 2. Dependency of k , (sec-'*') on catalytic acid concentration (M)

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 2 0 .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

EL Calculated from second order gate constants. - represents values for one standard deviation. From data reported in ref. 33. cCalculated from data from ref.27 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,

T for 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 of

27 benzohydroxamic 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,

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. esters and benzimidates previously discussed:

Products

Unlike the previous study of the acidic hydrolysis of benzohy-

27 droxamic 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:

(30) kobs * Kkl & 3 0+J + k2

v e rs u s jf wi l l y i el d anThis equation implies that a plot of k ^ g versus jf will yield anThis

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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 57

relate to the true value of kStep 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.

54 Specific 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-

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 7.4 LNaOHj

Figure 3. Dependency of k ^ C s e c on catalytic base concentration (M)

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 59

27 tetrahedral 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 ^11), 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

27 Fuller 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 27 which is inconsistent with present and previously observed data.

Increased hindrance in the present system, caused by N-methylation,

27 over 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 previously

27 studied benzohydroxamic acid series, lack of N-hydrogens in the

present case prohibits the tautomerization equilibria proposed in the

27 earlier 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)

- (34);

Cl Cl

(32)

Products (33)

Products (34)

Under the strongly alkaline conditions used in the kinetic studies,

JL will be almost completely converted to 2 . (The pKa's of N-tert.-

butylbenzohydroxamic and N-phenylbenzohydroxamic acids are 10.1

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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

27 earlier work, is derived as follows (water concentrations are

included in the constants):

-dS/dt = f2j(COH7 k2 + k3)

= [ l j K l O H ^ ( [0H7k2 + k 3) (35)

S = [l](l + [OH J K)

Therefore:

-dS/dt = SKfOHj (COH^k2 + k3)/l + fOHjf K (36)

Under the reaction conditions employed, KpDHj >> 1 and:

kobs = k2C°H3 + k 3 (37) where k , is the pseudo first-order rate constant. The form of obs

equation (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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 various obs

ortho-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 by

2 57 computer 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 63

Table X

Observed and Calculated Rate Constants for Acidic Hydrolysis of ortho-Substituted-N-Methylbenzohydroxamic Acids in 0.764M HC1 at 90.0°C

Ortho- , , log k, (sec 1) log k, (sec Substituent a * Eg (observed)3 (calculated)

0 0 -4.359 -4.403 "CH3

-°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 58 the 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-test^’^ 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),

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 64

-5.00-

80-

■4.60-

-4.40-

-4.20 -0.30 - 0,10 0.10 0.30 0.50 0.70 0.90 1.10 0* Figure 4. Correlation of log kQbs with a* and Es values for acidic hydrolysis

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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

-N 0 2 -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 from calculated 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 66

a comparison of residuals (i.e., | log k ^ g - log kca^cl) 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 equation obs 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- s

ficance 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 not s

reproduce 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 equilibrium 4 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-

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 67

hydroxamic acids with the polar and steric substituent parameters

(i.e., a* and Eg respectively) in equation (12), taken together with

2 45 a similar correlation reported earlier for the acid catalyzed

hydrolysis of ortho-substituted benzohydroxamic acids, lend support to

22 the 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

aci d s . ^ 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

22 by 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.

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 ratio

2,45 for 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

system.Yet, 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 45 hindered, N-H system. * The differing values for these ratios are,

nonetheless, consistent with the qualitative conclusions of McCoy and

22 Riecke. 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. 69

The 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 form:^ Hammett

* q L *------^ X = 0 = i - G (38)

Exner,^ 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, s

Lastly, 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 71

Table XII

Comparison of o ° and a Values for Para-Substituents of Benzenl Derivltives in Aqueous Media^

Substituent a ° a P 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, for 4 the saponification of ethyl para-dimethylaminobenzoate, that only

part of the resonance interaction of the p a r a - (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 when 4 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 the

22 ortho-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 the

29 general hydrolysis mechanism for hydroxamic acids by Buglass et. al.

27 and by Berndt and Fuller discussed earlier. The difference between

2,45 the 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 various obs

ortho-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 57 ment * 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 74

Table XIII

Observed and Calculated Rate Constants for Alkaline Hydrolysis of Ortho-Substituted-N-Methylbenzohydroxamic Acids in 7.31M NaOH at 90.0°

Ortho- log k Substituent a * Eg (observed) ’ log k (calc.) *

-CH3 0 0 -6.130 -6.178

- 0.22 0.99 -5.499 -5.466 -°CH3

-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).

4,59 limits 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"5^’58 showed this correlation to be significant

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. without prohibited reproduction Further owner. copyright the of permission with Reproduced -5.60 -6.4C -6.80- Figure 5. Correlation of log kQ^ s with with s kQ^ log of Correlation 5. Figure log (k - 0.2 0.0 0.2 0.4 * a and E g values for alkaline hydrolysis alkaline for values g E and E. s 0.6 1.0 1.2 75 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 to

4 59 a 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 59 studied ’ * ’ ’ 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^ C oH^ ) (see equation (37)).

Contributions to kQbg from k2 and k3 will vary with each ortho

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 that

27 proposed 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

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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 s larger 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.,

~~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~~

~~Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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.~~

~~Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. BIBLIOGRAPHY~~

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~~Reproduced with permission of the copyright owner. Further reproduction prohibited without permission. 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.~~

~~Reproduced with permission of the copyright owner. Further reproduction prohibited without permission.~~