.r.u:~‘\\‘~flt "
I‘ .
V , . I, ,r ‘fl'llvh-I.” IILI'I- _ ‘
'7 ‘n'n-K ‘,.~n .. ‘ .
\

a SECONDARY KINETIC DEUTERIUM
ISOTOPE EFFECTS IN THE .‘ E w
T BASE HYDROLYSIS 0F , ' » ' .
=r ALKYL 0I- NAPHT‘H‘OATES . _

   

 

\

 

' Thesis for the Degree of Ph. D.
MICHIGAN .‘STATE UNIVERSITY
{- ROBERT LAWRENCE TURNER
3 a 1972

 

 

,—
,,

_.-

a. » '

C" . .

.- .7 -
I ~

‘

 

I

. , INTERN-I

 

 

     

   
 
  

umuuufimmmmm

 
  

   

'7‘." '-';’r» -'

   
 

3&5

.1-,..

i.

 

"WW

 
 

. o
. ..,.~
”Mu-1'
. . .1.‘

   

     

LIBRARY

Michigan State
University

This is to certify that the

thesis entitled

SECONDARY KINETIC DEUTERIUM ISOTOPE EFFECTS
IN THE BASE HYDROLYSIS OF ALKYL a-NAPHTHOATES

presented by ~

Robert Lawrence Turner

has been accepted towards fulfillment
of the requirements for

Ph .1) o deyee in Organic Chemistry

 

Ewan/(ax:

hhmnpnmuun

Date July 19, 1972

0-7639

 

his-
-

i‘ am new
"HAG & SNIS’
? 800K BWDERYIHC.

LIBRAR SINGERS

I I I 1'.i.«’i§1 IVv

ABSTRACT

SECONDARY KINETIC DEUTERIUM ISOTOPE EFFECTS

IN THE BASE HYDROLYSIS OF ALKYL a-NAPHTHOATES

Robert Lawrence Turner

Since Bartell (1) first proposed that non-bonded interactions
may be important in secondary kinetic isotope effects, many efforts
have been directed toward elucidating the nature and significance
of these effects. In order to further investigate the role of
non-bonded interactions in secondary kinetic deuterium isotope
effects, we undertook a study of the temperature and solvent
dependence of the base hydrolysis of the methyl l-naphthoates I and
II. The results obtained from this study dictated a further investi-
gation of the isotope effects resulting from deuteration of the
alcohol portion of methyl 1-naphthoate (Ic) and 2,2-dimethylpr0pyl
l-naphthoate (III).

The surprisingly small isotope effects observed for Ib

(kn/kn - .979 i .006 at 26°) and 11b (kn/RD - .971 1 .009 at 85°)

Robert Lawrence Turner
Ia: R = CH3; X I H

Ib: R s CH3; X - D
Ic: R 2 CD3; X - H

X COZR 113: R - CH - x . CH
3’ 3
11b: R - CH3; X 8 CD3
IIIa: R - CH2C(CH3)3; x a H
IIIb: R - CD2C(CH3)3; x a H }
IIIc: R - cnzc(cn3)3; x . H .
IIId: R x - D

= CH2C(CH3)3;

in 48.12 aqueous methanol indicated that secondary kinetic 1
deuterium isotope effects arising from non-bonded interactions

are insignificant when other electronic factors are Operating. The
thermodynamic parameters calculated from the temperature dependence
of kH/kD demonstrated that the steric factors affected both AAH*
and AAS*. These opposing effects resulted in a very small value of

AAG*. Explanations for these isotope effects in terms of competing

steric effects are offered.

The large difference in the isotope effects for the base
hydrolysis of Te (kH/kD - 1.030 1 .008) and IIIb (RH/RD - 1.099 i
.010) in 54.1% aqueous dioxane at 25° indicated a change in the rate
determining step for the hydrolysis of the neopentyl ester. The
normal isotope effect for the hydrolysis of IIIb was explained in

terms of an inductive effect on the relative stabilities of the

developing alkoxide ion in the second step.

 

References

1. a. L. S. Bartell, J. Am. Chem. Soc., §;, 3567 (1961).
b. L. S. Bartell, Iowa State Journal of Science, 39, 137 (1961).

 

 

SECONDARY KINETIC DEUTERIUM ISOTOPE EFFECTS

IN THE BASE HYDROLYSIS OF ALKYL a-NAPHTHOATES

by

Robert Lawrence Turner

A THESIS

Submitted to

Michigan State University

in partial fulfillment of the requirements

for the degree of

DOCTOR OF PHILOSOPHY

Department of Chemistry

1972

To Beth

ii

ACKNOWLEDGMENTS

The author wishes to extend his appreciation to Professor
Gerasimos J. Karabatsos for his guidance and his assistance in the
interpretation of the data presented. Appreciation is also extended
to fellow members of the research group for many stimulating dis—
cussions both in and out of the field of organic chemistry.

The financial assistance provided by the Lubrizol Corporation
and the Department of Chemistry, Michigan State University, is also
gratefully acknowledged.

Finally, the author wishes to thank his wife, Beth, for her
patience and encouragement throughout the course of this investi-

gation, without which, this work could not have been completed.

111

LIST OF TABLES. . . .

LIST OF FIGURES . . .

INTRODUCTION. . . .

EXPERIMENTAL. . . . .

I. Kinetics . .

TABLE

Preparation of
Kinetic Apparatus.
Measurement of Time.
Temperature Control.
Temperature Determination.
Rate Determinations.
Treatment of Data.

II. Syntheses. .

OF CONTENTS

Solvents.

Methyl 8-d-1-naphthoate. .
Methyl-d 1-naphthoate .
Methyl 8—methy1—d-l-naphthoate. . .

2 ,2—Dimethy1pr0py -d1
2,2—Dimethylpropy1-1,

RESULTS

I. Mechanism for the Base Hydrolysis of

AND DISCUSSION.

1- -naphthoate.

-d_ l-naphthoate
2,2-Dimethy1propyl B-gfl-naphthoate. .

Alkyl Esters . . . . . . . . . . . . . .

II.

III.

IV.

V. Conclusion .

BIBLIOGRAPHY. . . . .

iv

Temperature Dependence . . . .

Interpretation of IsotOpe Effects.

Correlation of Thermodynamic Parameters.

Page

vii

20
20
20
21
22
22
23
23
24
25
25
31
31
32
34
34

35

36
39
54
58
64

65

10.

ll.

12.

13.

LIST OF TABLES

Isotope effects on the acidities of organic
acids in water. . . . . . . . . . . . . . . . .

Isotope effects for the solvolysis of tfbutyl
chloride in 60% aqueous ethanol at 25°. . . . .

IsotOpe effects for the solvolysis of para-
substituted 1-pheny1ethy1 chlorides in 50%
aqueous ethanol at 25°. . . . . . . . . . . . .

Isotope effects for the reaction of alkyl iodides
with methylpyridines in nitrobenzene. . . . . .

Comparison of calculated and experimental isotope
effects for some limiting solvolysis reactions.

Solvent and temperature dependence of the isotope
effects for the solvolysis of tfbutyl-d

chloride. . . . . . . . . . . . . . . :9. . . .

Solvent dependence of the isotOpe effects for the
solvolysis of p-alkylbenzhydryl chlorides . . .

Solvent dependence of the B-isotope effect for

the solvolysis of acetyl-d3 chloride. . . . . .

Temperature independence of the isotope effects
for the solvolysis of isopropyl compounds in
water 0 O O O C O O O O I O O O O O O O O O O 0

Temperature dependence of the isotope effect for
the solvolysis of acetyl-d3 chloride in 90%
aqueous acetone . . . . . . . . . . . . . . . .

Solvent mixtures used in the kinetic studies. . .

Effects of para-substituents on the rates of
alkaline hydrolysis of ethyl benzoates at 25° .

Substituent effects on the rates of alkaline
hydrolysis of alkyl esters at 25° . . . . . . .

Page

10

13

14

15

l6

17

17

21

37

38

t-.~.a‘~‘l:i-"~‘ - ‘

14.

15.

16.

17.

18.

19.

20.

21.

22.

23.

24.

25.

Rates of base hydrolysis of methyl l-naphthoate
and methyl 8-df1-naphthoate in aqueous methanol .

Rates of base hydrolysis of methyl 1—naphthoate
methyl 8fdfl-naphthoate, and methylfid
l-naphthoate in 54.1% dioxane/water (w/w) . . . .

Rates of base hydrolysis of 2,2—dimethy1propyl
1-naphthoate, 2,2-dimethylpropyl-—1,1—51__2
1-naphthoate, 2,2-dimethylpropyl—d1
1-naphthoate, and 2,2-dimethy1propy
Bfgfl-naphthoate. . . . . . . . . . . . . . . .

Rates of base hydrolysis of methyl 8—methy1-l—
naphthoate and methyl 8-methyled3-1-naphthoate
in 48.1% methanol/water (w/w) . . . . . . . . . .

Variation of / with temperature and solvent
composition or methyl Begfl—naphthoate . . . . .

Variation of / with temperature and solvent
composition or methyl l-naphthoates and
2,2-dimethylpropy1 1—naphthoates. . . . . . . . .

Thermodynamic parameters for methyl 1-naphthoate
and methyl 8-d71-naphthoate . . . . . . . . . . .

Thermodynamic parameters for methyl l-naphthoates
and 2,2—dimethy1propyl l-naphthoates in 54.12
dioxane/water (w/w) . . . . . . . . . . . . . . .

Thermodynamic parameters for methyl 8-methyl-l-
naphthoate and methyl 8-methy1-d3-1-naphthoate
in 48.12 MeOH/water (w/w) . . . . . . . . . . . .

Steric effects in the base hydrolysis of some
alky1 es ters O O O O O O O O O O O O O O O O O O I

Effects of chain branching on the rates of base
hydrolysis of some alkyl esters in 70% aqueous
dioxane at 40° C O O O O O O O O O O O O I O O O O

IsotOpe effects for the base hydrolysis of methyl
8-methyl-d3-1—naphthoate in ethanol . . . . . . .

vi

4O

42

43

45

46

47

48

50

52

54

56

60

-—.‘- .-...._. .4>-__'4.-.“M m- v"' I. 'I

LIST OF FIGURES

Figure Page

1. Vibrational potential energy function for
X-H and X-D bonds. 0 o o o o o o o o o o o o o o o o o 2

2. Relation of force constant changes to ZPE
differences. 0 O O O O O O I O O O O O O O O O O O O O 3

3. Run No. 204, Methyl l-naphthoate, 54.1% Dioxane,
.194M Sodium Hydroxide, 50.004°. . . . . . . . . . . . 26

4. Run No. 312, Methyl 8-methyl-l-naphthoate,
48.12 Methanol, .198M Sodium Hydroxide, 100.384° . . . 27

5. Solvent effects on the thermodynamic parameters

for the base hydrolysis of methyl l-naphthoate
in aqueous methanol solutions at 25° . . . . . . . . . 55

vii

INTRODUCTION

Secondary kinetic isotope effects resulting from deuterium
substitution are believed to arise from differences in vibrational
force constant changes in going from ground state to transition
state. The electronic structure and hence the forces which bind
atoms together are independent of the changes in atomic mass of
nuclei caused by deuterium substitution. Therefore, the potential
energy surfaces for X-H and X—D are invariant, i.e., the force
constants are the same. Kinetic isotope effects are then the
result of differences in nuclear mass on the motion of the nuclei
within the same potential energy surface.

The difference in the anharmonic vibrational energies of X-H
and X-D bonds is illustrated in Figure 1. From this figure we
conclude that both the zero-point vibrational energy and the mean
vibrational amplitude are greater for X—H than for X-D. This leads
to different average bond lengths and angles in deuterated molecules.

The relation of force constant changes to the isotopic rate
ratio, kH/kD, is illustrated in Figure 2. If the force constant
decreases, the ZPE difference decreases, and a normal isotope
effect is observed, kH/kD > 1. If the force constant increases,
the ZPE difference increases, and an inverse isotope effect is
observed, kH/kD < 1. The direction of the force constant change
can be determined experimentally. The problem, however, is

1

['11

 

 

 

l X-H
I
ZP
an MD x1)
II
II
aDIIan
H
II
1L
r --€>

Figure 1: Vibrational potential energy function for X—H
and X-D bonds (2).

‘4 I ' :c'T-“fiiu_d "'

A. Decrease in force constant; AAZPE < 0, kH/kD > 1

 

 

AZPE A X'H
T3 1 X—D
E
X-H
'
r —>

B. Increase in force constant; AAZPE > 0, kH/kD < 1

m

 

 

 

r—->

Figure 2: Relation of force constant changes to ZPE differences (2).

 

4
interpreting the observed force constant changes in physical
organic terms.

The main reason for this difficulty is our inability to
determine the exact geometry of the transition state for a given
reaction by extrakinetic methods. Since many factors are involved
in the determination of a kinetic isotope effect, e.g., induction,

hyperconjugation, non-bonded interactions, hybridization, and

1 .Z_____ _.I__ A _‘

salvation, these factors are normally impossible to separate from

._——F___. -.

one another. The correlation of a kinetic isotope effect with a
reaction mechanism is, therefore, a rather complex problem. As a 3
result, there is little agreement as to the origin of force constant
changes observed in secondary kinetic isotope effects. The factors
generally applied in interpreting secondary kinetic isotope effects
are induction, hyperconjugation, and non-bonded interactions. These
factors can be related in a qualitative fashion to changes in the
force constants.
Measurements of dipole moments (3) and NMR studies (4,5) have
shown that deuterium is more electropositive than hydrogen. Molecu-
lar refraction (6) and optical activity data (7) have also demon-
strated that C-D is less polarizable than C-H. The principal
factors involved appear to be the anharmonicity of the vibrations
and the difference in average bond lengths and angles between
deuterated and undeuterated molecules. From these data on the
physical prOperties of deuterated compounds one concludes that the
electron density about carbon is greater in the C-D bond than in the

0-H bond. Therefore, if a positive charge is produced near the site

5
of deuteration, C—D should stabilize it better than C—H.
Consequently, an inverse isotope effect should be observed.
The most pertinent chemical evidence for the greater inductive
effect of deuterium arises from isotOpe effects on the acidities of
deuterated acids. These results are summarized in Table 1. It

should be pointed out however, that although the observation of

Table l: IsotOpe effects on the acidities of organic acids in

water 0

 

 

Acid KH/KD

 

DCOZH 1.070 t .0083, 1.084e
CD3C02H 1.032 i .0028
(CD3)3CC02H 1.042 t .003b
cuacnzcozn 1.08C
CDBCHZCOZH 1.01c
c6nscnzcozu 1.12c
0605002H 1.024 : .0065"b
C6D50H 1.12 1 .02b
C6D5NH; 1.06 i .02d
2,4,6-c H D NR+ 1.04 i .02d

6 2 3 3
3.5-66H3D2NH: .99 i .02d

 

8Reference 8. bReference 9. cReference 10. dReference 1.

eReference 13.

-. ~_ if.
. ) .

6
normal isotope effects for these dissociation reactions confirm
the greater inductive effect of deuterium, any quantitative
application of these effects is meaningless, since substituent
effects on weak carboxylic acids operate through changes in
entropy, i.e., AAHO = 0. Also, pKa differences are known to be
quite sensitive to both solvent polarity and temperature (12).
Since the inductive effect of deuterium located at least one

carbon atom removed from the reaction center is very small and

inseparable from other factors, inductive effects are seldom applied

“-‘A 1fi#.!_.»4 1__,..
l.'. .b ‘. I _

in correlating secondary kinetic isotope effects and reaction

mechanisms. e
Halevi has stated that hyperconjugation with C-D is less

effective than with C—H due to C-D being less polarizable (11).

However, Streitweiser and coworkers (16c) have pointed out that an

increase in hyperconjugation in the transition state results in a

weakening of the C-H force constant. Therefore, a normal isotope

effect would be predicted, Figure 2, without invoking differential

hyperconjugation for C-H and C-D.
Although the role of hyperconjugation in secondary kinetic

isotope effects has been questioned, substantial evidence has been

accumulated which makes this role quite prominent. Shiner and

coworkers (14) have reported significant rate retardations for the

solvolysis of deuterated tfbutyl chlorides. Their results are

summarized in Table 2. The variation of kH/kD with the amount of

deuterium substitution supports the contention that hyperconjugation

is important in this mechanistically limiting reaction. Furthermore,

the slight increase in kH/kD per deuterium atom is also consistent

7

Table 2: IsotOpe effects for the solvolysis of Efbutyl chloride
in 60% aqueous ethanol at 25° (14).

 

 

 

Compound kH/kD kH/kD per D
(CH3)2CClCH2D 1.092 1.092
(CH3)2CClCHD2 1.202 1.096
(033)20c1cn3 1.330 1.100
(003)20c1ca3 1.710 1.102
(c03)3cc1 2.327 1.103

 

with the conformational dependence of hyperconjugation.

The dependence of hyperconjugation on the spatial orientation
of the 0-H bond was elegantly demonstrated by Shiner and Humphrey
(15) in the solvolysis of the deuterated bicyclooctane derivatives

I and II in 60% aqueous ethanol at 45°. The normal isotope effect

 

ca 01
D
@. @
I II
kH/kD - 1.14 i .01 kH/kD - .986 1 .01

for compound I is consistent with hyperconjugative stabilization

in the transition state. Since such stabilization is impossible

 

-r.—. =-_ -—

8

when the C-H bond is in the nodal plane of the developing vacant
p orbital, no normal isotope effect is observed for 11. The small
inverse isotope effect presumably arises from the greater inductive
effect of the 0-D bond over the 0-H bond (16).

Shiner and Kriz (19) have also shown that hyperconjugation can
be transmitted through an unsaturated linkage. In the solvolysis
of the unsaturated compound III, the site of isotopic substitution

is too far removed from the reaction center for inductive or steric

I, .k".-f§ «A. ~n\- I

effects to be important. Consequently, the observed normal isotope 9

effect is ascribed to hyperconjugation through the n-orbitals of

the acetylenic bond. . 1
CH3 003
cn3-czc—c-cu3 CRB-czc-c-cn3
01 01
111 IV
kH/kD = 1.092 t .001 kH/kD = 1.655 1 .004

Shiner and coworkers (20) found rather large variations in
the B-isotope effect for the solvolysis of BEEEfSUbStitUtEd
l-phenylethyl chlorides, Table 3. The large a-isotope effect is
characteristic of reactions whose mechanisms are limiting. Although
the a-isotope effect remains essentially constant throughout the
series of substituents, indicating no change in the mechanism of
the reaction, the B-isotope effect varies considerably. As the
electron releasing ability of the substituent increases, the amount

of positive charge present at the reaction center decreases, thus

9
the need for hyperconjugative stabilization from the methyl group

is lowered. Consequently, kH/kD decreases.

Table 3: IsotOpe effects for the solvolysis of para-substituted
l-phenylethyl chlorides in 50% aqueous ethanol at 25° (20).

 

 

Substituent kH/kD (a) kH/kD (B)

 

p-OCH3 1.157 1.113

p-oc6HS 1.157 1.164

p-CH3 1.157 1.200 __
p-F 1.152 1.211

m-CH3 1.151 1.222

p-H 1.153 1.224

t. .I!.ua._.;“-x—-_-_nur -- I
I

 

This dependence of hyperconjugation on electron demand has
been used to support the nonclassical norbornyl cation (23). The

smaller isotope effect for the solvolysis of the exo-bromide, VI,

 

V VI

kH/kD 8 1.16 kH/kD = 1.04

 

over that of the endo-bromide, V, was attributed to charge delocal-

10

ization by o—participation which lowers the requirement for C-H

hyperconjugation from the 3—position.

Similar results were obtained

for the solvolysis of the 3,3-dideuterio-2-norbornyl brosylates (24).

Brown and McDonald (26) have reported inverse isotOpe effects

for the reaction of alkyl iodides with methylpyridines, Table 4.

Since no effect was observed for 3-methylfg3-pyridine or 4—methyl-d_
pyridine, the authors concluded that the inverse isotope effect for
2-methyled3-pyridine was steric in origin.

supported by the observation of an increase in kD/kH with the size

of the alkyl iodide and an increase in kD/kH for 2,6-dimethyl

pyridine.

Table 4:

methylpyridines in nitrobenzene (26).

This conclusion was

_-(-1-6-

Isotope effects for the reaction of alkyl iodides with

 

 

 

Compound Alkyl iodide Temp.(°C) kD/kH
4-Methyl-d_-pyridine CH3I 25 1.001 t .003
3-Methylfd3-pyridine CHBI 25 1.009 t .002
2—Methylfg3-pyridine CH3I 25 1.030 t .003

CH3CHZI 75 1.036

(CH3)2CHI 100 1.058
2,6-Dimethylfig6-pyridine CHBI 25 1.095 t .003

CH3CH21 75 1.072

CH CH I 100 1.070

3 2

 

3

_ 1.'.‘”‘wmml‘ we:

ll

Bartell, like Brown, has stated that secondary isotope effects
can be explained in terms of non-bonded interactions (28). The basis
for this argument is that repulsions involving deuterium atoms are
smaller than those for hydrogen due to the smaller mean-square
amplitude of vibration of C-D. The effective size of deuterium is,
therefore, smaller than that of hydrogen. Consequently, a release
of steric compression in the transition state will yield a normal
isotope effect, while an increase in steric compression will result
in an inverse effect. Although Bartell has demonstrated that non—
bonded interactions cannot be ignored when correlating secondary
kinetic isotope effects, it is often not known whether these inter-
actions are repulsive or attractive. Quantitative predictions are,
therefore, unjustified. This is particularly true in view of our
lack of knowledge concerning the geometry of the transition states.
Bartell, himself, has noted that it is impossible to separate
steric effects from other electronic factors.

The inverse isotope effect reported by Melander and Carter (29)
for the racemization of 2,2'—dibromo-4,4'-dicarboxy—6,6'-biphenyl-g2,
VII, in ethanol was attributed to non-bonded interactions. The

Br D Br D
D Br D Br
VII VIII

/ = 1.18 1 .03 / - 1.02 1 .02
kokH kaH

.. -_-'—a_-u—.—m_4 ‘1.- I
A
n
y

12

possibility of the isotOpe effect originating partially from the
greater inductive effect of deuterium on n-orbital overlap was
negated with the observation of no isotOpe effect outside of exper—
imental error for the racemization of the 5,5'-dideuterio
derivative, VIII.

Another example of the importance of steric interactions was
reported by Carter and Dahlgren (30) for the racemization of

2,2'-binaphthy1—d_ IX. The increase in rate resulting from i

2’
deuterium substitution was ascribed to an increase in the force
constant for the in-plane C-H bending vibration. Mislow and
coworkers (31), having found a large inverse isotope effect for the
racemization of the sterically non-planar ring system X, concluded:
It does not seem likely that steric factors are responsible
to a major extent for deuterium isotope effects observed in
SNl reactions . . . since the hydrogens are far less compressed
than in these systems studied . . . therefore, non-bonded

interactions are operative only under conditions of severe
crowding.

IX X

kD/kH - 1.186; 36° kD/kH - 1.137; 42°

13

A similar conclusion was drawn by Karabatsos and coworkers (32)
in the comparison of calculated and experimental isotope effects for
some limiting solvolysis reactions. Their results are summarized in
Table 5. Using the potential functions of Scott and Scheraga (33),
the authors employed Bartell's procedure (28) to calculate isotope
effects based on non-bonded repulsions. In those systems where
hyperconjugation was possible, their calculations greatly under-
estimated the experimental isotope effects. They concluded that
less than 10% of the observed isotope effect is due to non-bonded I

interactions when hyperconjugation is possible.

Table 5: Comparison of calculated and experimental isotope effects

for some limiting solvolysis reactions (32).

 

 

 

Compound kH/kD (calc.) kH/kD (exp.)
Izt-nalahthylcarbinyl-OI,0I--_c_l__2 chloride 1.30 1.30
acetyl-jg3 chloride .98 1.68
_t_;-buty1-d9 chloride 1.10 2.39

 

Hakka and coworkers (34) demonstrated that the isotope effect
for the solvolysis of EfbutYIfgg chloride, Table 6, was solvent
independent. They concluded that there was no significant inter-
action of solvent with the developing cation in this solvolysis
reaction. These findings were confirmed by Frisone and Thornton

(35) who studied the effect of a series of solvents with the same

ionizing power,

14

Table 6: Solvent and temperature dependence of the isotope effects

for the solvolysis of tfbutyl—d chloride (34).

 

 

 

_9
Temp.(°C) kH/kD (water) kH/kD (50% ethanol)

5 2.55 2.542

10 2.53 2.505 3
15 A 2.48 2.465
20 2.45 2.419 I
25 2.387 I
30 2.345 5

 

In the solvolysis of p-alkylbenzhydryl chlorides, Shiner and
Verbanic (21) have reported significant changes in the isotope
effects with solvent, Table 7. The observed normal isotope effect
was attributed to hyperconjugation through the phenyl ring. They
suggested that the decrease in kH/kD with increasing solvent polarity
was due to a greater dispersal of the developing charge through
solvation, resulting in a decrease in the effectiveness of hyper-
conjugation. Similarly, the decrease in kH/kD with greater chain
branching was ascribed to an increase in steric hindrance to solvent

participation at the site of isotopic substitution.

Schubert prefers to explain these results in terms of steric
hindrance to solvation at the para-position of the phenyl ring (22).
The decrease in kH/kD with increasing solvent polarity is ascribed

to an increase in solvent stabilization of the delocalized charge,

15

thereby, lowering the demand for hyperconjugative stabilization.

Table 7: Solvent dependence of the isotope effects for the
solvolysis of p-alkylbenzhydryl chlorides (21).

 

 

 

Com d 90% 80% 70% 67%
poun Ethanol Acetone Acetone Acetone
p-CD3C6H4CH¢C1 1.025 1.058 1.038 1.021 i.
p-CH3CD2C6H4CH0C1 1.009 1.025 1.019 1.012 {
p-(CH3)2CDC6H4CH¢C1 .998 1.006.

 

Evans (2) found that the B-isotope effect for the solvolysis
of acetylug3 chloride depended strongly on solvent, Table 8.
Although the change in kH/kD with solvent composition may be related
to a change in mechanism, the data definitely indicate that solvent
may play an important role in determining the magnitude of a
secondary kinetic isotope effect.

The role of solvation in secondary kinetic isotope effects is
still generally ignored. This is not due to the unimportance of
solvation, but to the lack of quantitative data on this subject.

It is, therefore, important that the possible dependence of an
isotope effect on salvation be considered when interpreting isotope
effects. In Halevi's words (11): "It is an oversimplification

to disregard solvent when dealing with small effects in highly

polar solutions."

16

Table 8: Solvent dependence of the B-isotope effect for the

solvolysis of acetylfd_ chloride (2).

3

 

 

 

(% Ac:::::7aater) Temp.(°C) kH/kD
75 —3l.0 1.135 1 .020
80 -3l.0 1.122 t .010
85 -31.0 1.122 t .013
90 -28.9 1.044 i .009
95 -25.5 1.004 1 .004

 

The temperature dependence of the isotope effect for the
solvolysis oft-butyl-d9 chloride reported by Hakka and coworkers,
Table 6, conforms closely to the normal approximation that AAG* =
AAH*. Leffek and coworkers (36), however, found an unusual tem—
perature independence of the isotope effect for the solvolysis of
isopropyl tosylate, isopropyl methanesulfonate, and isopropyl
bromide, Table 9. This temperature independence was found to arise
entirely from entropy differences dominating AAG* and, therefore,
could not be explained in terms of zero-point energy differences
alone. Calculations indicated that differences in the six-fold
barriers to internal rotation between deuterated and undeuterated
substrates would tend to cancel the enthalpy difference and create
an entropy difference of the magnitude observed. The temperature

independence was then attributed to a fortuitous cancellation of

the AAH* arising from zero-point energy sources and internal

17
rotation effects. This type of rationalization is supported by
the theoretical analysis of Wolfsberg and Stern (37) who demon-
strated that compensating effects in force constant changes may

lead to temperature independent kH/kD values.

Table 9: Temperature independence of the isotope effects for the

solvolysis of iSOpropyl compounds in water (36).

 

 

 

Compound Temp. Range kH/kD
Isopropyl-d6 tosylate 6° to 30° ' 1.54
Isopropyl-d6 methanesulfonate 5° to 30° 1.55
Isopropyl—g6 bromide 40° to 70° 1.32

 

An unusual temperature dependence was reported by Evans (2)

for the solvolysis of acetylfd_ chloride in 90% aqueous acetone.

3

Table 10: Temperature dependence of the isotope effect for the

so 1volysis of acetyl-g

3 chloride in 90% aqueous acetone (2).

 

 

 

Temperature (°C) kH/kD
- 9.54 1.070 t .010
-15.72 1.072 t .015
-22.01 1.059 t .002
~28.88 1.044 t .009
-33.68 1.026 t .010

 

18

This temperature dependence, an increase in kH/kD with increasing

temperature, was found to be solvent dependent, since it was absent
in more polar solvents, Table 10. A similar temperature dependence

was reported by Smith (38) for the solvolysis of the endo-norbornyl

p-nitrobenzoates, XI and X11, in 80% aqueous ethanol. This
temperature dependence, which predicts an inversion in kH/kD at
lower temperatures, was attributed to large entropy contributions

to AAG*.

 

I I
c. I. I.
3 CH3 CD3 CH3

CD3 OPNB CH3 OPNB
XI XII

Tem . °C) EH/Ea (XI) kg/kg (XII)
100 .997 t .015 1.011 1 .013
125 1.003 t .014 1.017 t .014
142 1.027 t .013 1.043 t .013
150 1.026 t .010 1.049 t .011

It would appear that a precise evaluation of secondary kinetic
isotOpe effects demands an accurate identification of the force
constant changes in the activation process. Furthermore, any such
evaluation must include both the temperature and solvent dependence

of the isotope effects measured.

In order to further investigate the role of non—bonded interac-

tions in secondary kinetic isotope effects, we undertook to study

6 r‘-“ ..r--.-x~‘:-.
.

19
the temperature and solvent dependence of the base hydrolysis of
methyl l-naphthoate, methyl 8-methy1-l-naphthoate, and their
8—deuterated analogs. The results obtained from this study dictated
a further investigation of the isotOpe effects resulting from
deuteration of the alcohol portion of methyl 1—naphthoate and

2,2-dimethy1propy1 l-naphthoate.

.A__._-__.‘p-
;‘

1.~_.. -__.__._

EXPERIMENTAL

I. Kinetics

A. Preparation of Solvents

Methanol: Methanol was purified in two steps. First, three
liters of methanol (Baker Reagent Grade) was refluxed for 12 hours
with a mixture of 150 m1 furfural and 375 ml 10% aqueous sodium
hydroxide. The resulting solution was fractionally distilled
through a 30 cm Vigreaux column. 0f the distilled methanol, 150 ml
was reacted with 30 g of powdered magnesium until all of the
magnesium had dissolved. Then, two liters methanol was added,
refluxed for four hours, and fractionally distilled through a

50 cm Vigreaux column (39).

Water: Water was refluxed for 12 hours while bubbling
nitrogen through the liquid and fractionally distilled through a

30 cm Vigreaux column.

1,4-Dioxane: Dioxane was purified according to the procedure

outlined by L. F. Fieser in Experiments in Organic Chemistry (39).

 

Solvent Mixtures: The various solvent mixtures were prepared
by weight ratios using a Torbal balance (Torsion Balance Co.,
Model PL-2). Sodium hydroxide (Baker Reagent Grade) was added to
the solvent mixtures and the base concentrations determined by

20

J"h*-. -Q..L-“.A..-
r

21
titration against Fisher Primary Standard potassium hydrogen
phthalate by using phenolphthalein as indicator. The concentrations
of sodium hydroxide in the various solvent mixtures are reported

in Table 11.

Table 11: Solvent mixtures used in the kinetic studies.

 

 

 

Solvent “M “iii?” (molifii‘ii‘ler)
10.0% Methanol/Water .947 .1990 i .0004
30.0% Methanol/Water .806 .2029 i .0004
48.1% Methanol/Water .659 .2031 i .0004
48.1% Methanol/Watera .659 .1980 1 .0008
80.0% Methanol/Water .308 .2007 i .0008
54.1% Dioxane/Water .806 .1943 i .0003

 

aSolvent used for the base hydrolysis of methyl 8-methyl-1-
naphthoates.

B. Kinetic Apparatus

A Unicam SP.800 ultraviolet spectrophotometer equipped with an
SP.825 program controller, a mains relay, and a Sargent Recorder
(Model SR) was used for all absorbance measurements. Except for the
hydrolyses of the methyl 8—methy1-l-naphthoates, all kinetic runs

were followed continuously by using the automatic program controller.

22

C. Measurement of Time

A Precision Scientific electric digital timer (Model No. 69235)

accurate to 0.01 minute was used for the measurement of time.

D. Temperature Control

The constant temperature bath consisted of a 16 liter glass jar
insulated on the sides and bottom by at least one inch of vermilite
packing supported within a plywood box. The top of the bath was

covered with a sheet of 1/4 inch asbestos. The bath was equipped

- \.\. .._.-___~.‘ .. _—4A‘-4A.
. n.

with a mechanical stirrer (Talboys Instrument Corp., Model No. 104),
a Teel pump (Dayton Electrical Mfg. Co., Model No. 1P731), a Beckman
differential thermometer, a 125 watt heating blade (Cenco), an
electronic relay (Precision Scientific Co., Model No. 62690), and a
mercury micro-set thermoregulator (Precision Scientific Co.,
Model No. 62541). At temperatures below 35°, cooling was provided
by the circulation of cold tap water through two feet of 1/4 inch
copper tubing supported in the bath. A continuous heating coil
was used for temperatures above 35°. The continuous cooling or
heating was regulated in conjunction with the intermittent heating
provided by the electronic relay in order to obtain the desired
temperature. Bath temperatures constant to 1 .006° were obtained
over the range of 15° to 69°. The temperatures of the cell
compartment in the spectrophotometer were constant to 1 .012°.

The constant temperature bath used for temperatures :_70° was
constructed in the same manner as outlined above except the Teel

pump was excluded and a 3 x 5 inch hole was made in the asbestos

23

for the insertion of a wire basket to hold the quartz ampules.

E. Temperature Determination

The temperatures of the bath and cell compartment were
measured by using a Hewlett Packard quartz thermometer (Model No.

2801A).

F. Rate Determinations

1. Methyl and 2,2-Dimethy1propyl l-naphthoates:

r nu...- -.—.__.___.___._>

Hydrogen and deuterium kinetic runs were performed randomly.
Solvent (3.3 ml) was placed in the UV cell and allowed to thermally
equilibrate within the cell compartment for two hours. Approximately
6 ul of a 0.10M ester solution was injected into the cell, shaken,
and readings began immediately. Twenty five to sixty absorbance
readings were made at 325 mu over the first 2.5 to 3 half—lives,
depending on the rate of the reaction. The infinite value was
measured after ten half-lives. The absorbance range measured was

.40 to .01 absorbance unit.

2. Methyl 8-methy1-l-naphthoates:

Hydrogen and deuterium kinetic runs were performed alternately.
Approximately 3.8 ml of the reaction mixture was placed in each of
eleven quartz ampules. The ampules were sealed and simultaneously
placed in the constant temperature bath. After allowing seventy
minutes for thermal equilibration, the ampules were removed at
selected time intervals over a range of 0.6 half-life and cooled

quickly in an ice bath. The ampules were stored at 0° until the

24
completion of each run (38).
0f each ampule, 3.00 ml was extracted with 10.00 ml
cyclohexane (Baker GC-Spectrophotometric Quality Solvent). The
cyclohexane solution was dried over anhydrous potassium carbonate
(Mallinckrodt Analytical Reagent) and the absorbance measured

against a blank solution at 310 mu. There was no absorbance by

“—4“,-
a

a reaction mixture treated in this manner after sixteen half-lives.
The absorbance range measured was .42 to .25 absorbance unit.

The absorbance was found to follow Beer's Law over the

- .*.h_v___.—.

concentration ranges measured for all the esters studied. Examples

of typical runs are illustrated in Figures 3 and 4. -1
G. Treatment of Data

First order rate constants were determined by solving the rate
expression
(At - A“) = Aoexp(-kt) - Abexp(-kt)
At - absorbance at time t
A“ = absorbance at infinite time
by using a least squares curvefitting program, KINFIT (40).

The average rate constants reported are the mean average of the
independent determinations. If an individual rate determination
appeared questionable, then the 4d rule was applied. That is, if a
given determination deviated from the mean by more than four times

the mean deviation of the remaining determinations, it was discarded.

The rate ratios, kH/kD, are the ratios of the corresponding

means. The reported deviations in kH/kD were calculated by using the

25

expression (41)

_ 2 2 2 2 4 1/2
(RH/ED - ton/k. + k. (ID/Isl
Thermodynamic parameters were determined from a solution of
In k/T versus l/T by using a least squares program, AKTIV. A solu-
tion of 1n kH/kD versus l/T with the least squares program HANDS and

the least squares curvefitting program KINFIT yielded values for

AAH¢ and AAS*.

~ '. "11‘4K_ null: 15”.. u. --I’

II. Syntheses

A. Methyl 8-deuterio-l-naphthoate -.

1. Anhydro-8-hydroxymercuri-1-naphthoic acid (42)

To a five liter, three-necked round-bottomed flask (equipped
with mechanical stirrer, reflux condenser, and addition funnel) was
added 142.4 g (.72 mole) 1,8-naphthalic anhydride in a solution of
100.8 g (2.52 mole) sodium hydroxide and 3500 ml water. After
heating the mixture slowly to reflux, a solution of 252.0 g
(.79 mole) mercuric acetate, 108 m1 glacial acetic acid, and 540 ml
water was added dropwise. The resulting mixture was refluxed for
forty hours. The hot mixture was filtered and the precipitate
washed with six 1000 ml portions water, 1000 m1 ethanol, and 500 m1
ether. The white solid was dried in an oven to constant weight;

yield = 265 g (992).

2. 8-Bromo-l-naphthoic acid (42)
A mixture of 130.1 g (.35 mole) anhydro-8-hydroxymercuri-

1-naphthoic acid, 565 ml glacial acetic acid, and 95 ml water was

26

-2.2 -

1n (At - Am)

 

Time (minutes)

Figure 3: Run No. 204, Methyl l-naphthoate, 54.1% Dioxane,
.l94M Sodium Hydroxide, 50.004°

.. An)

In (At

27

 

l l L

0 100 200 300

Time (minutes)

Figure 4: Run No. 312, Methyl 8—methy1-l-naphthoate,
48.1% Methanol, .l98M Sodium Hydroxide, 100.384°

1. -1-._‘A_1___- f--._..7.

28
added to a two liter, three-necked round-bottomed flask (equipped
with mechanical stirrer, addition funnel, and thermometer). While
cooling in an ice bath, a solution of 64.0 g (.40 mole) bromine in
285 m1 concentrated aqueous sodium bromide was added dropwise. The
mixture was heated slowly to 90° and the resulting yellow solution
poured onto 900 g ice/900 ml water. The aqueous mixture was

extracted with three 600 m1 portions ether. The combined ether

1}; twig-x. . an

solutions were extracted with three 480 m1 portions 10% aqueous
sodium hydroxide. The combined aqueous extracts were acidified
with concentrated hydrochloric acid and filtered. The white

precipitate was washed with six 500 ml portions cold water and -—:‘
recrystallized from 30% aqueous ethanol; yield - 61.5 g (70%),

mp l75-6°.

3. Silver 8-bromo-1-naphthoate

To a one liter, round-bottomed flask (equipped with mag-
netic stirrer and addition funnel) was added a solution of 30.0 g
(.12 mole) 8-bromo-l-naphthoic acid, 6.6 g (.062 mole) sodium
carbonate, and 600 ml water. A solution of 20.8 g (.122 mole)
silver nitrate in 180 ml water was added dropwise and the resulting
mixture stirred at room temperature for one hour. The white silver
salt was filtered and washed with two 150 ml portions water, 150 ml
methanol, and 150 ml acetone. The solid was dried in a vacuum

desiccator over phosphorous pentoxide; yield - 40.7 g (95%).

4. 8-Bromo-1-iodonaphthalene
A suspension of 38.0 g (.106 mole) silver 8-bromo—1-

naphthoate in 800 ml dry carbon tetrachloride was added to a three

29
liter, three-necked round-bottomed flask (equipped with mechanical
stirrer, reflux condenser, addition funnel, and calcium sulfate
drying tube). A solution of 27.9 g (.11 mole) iodine in 1700 ml
dry carbon tetrachloride was added drOpwise and the resulting
mixture refluxed for ten hours. The solution was filtered and the
filtrate washed with two 500 ml portions water, 300 ml saturated
aqueous sodium bicarbonate, 400 ml 20% aqueous sodium thiosulfate,
and 300 ml water. After drying the carbon tetrachloride solution
over anhydrous magnesium sulfate, the solvent was removed by
distillation and the residue recrystallized from 90% aqueous

ethanol; yield 8 20.0 g (60%), mp 99-100°.

5. 8-Deuterio-1-bromonaphthalene (43)

To a 250 ml, three-necked round-bottomed flask (equipped
with mechanical stirrer, reflux condenser, addition funnel, and
calcium sulfate drying tube) was added a mixture of 1.25 g (.051
mole) powdered magnesium, 60 ml anhydrous ether, 16.5 g (.050 mole)
8—bromo-l-iodonaphthalene, and one crystal iodine. The mixture was
heated gently until the reaction began, then stirred at room teme
perature until all of the magnesium was dissolved. The reaction was
cooled in an ice bath and 6.5 ml D20 was added dropwise. After
stirring the mixture overnight at room temperature, the mixture was
filtered and the magnesium salts washed with two 30 ml portions
ether. The ether was removed from the combined ether solutions by

distillation and the residue fractionally distilled under vacuum;

yield - 6.4 g (62%), bp 90° (.05 mm).

flaw.“ v! u'fl.

30
6. 8-Deuterio-1—naphthoic acid (42)
The acid was synthesized according to the procedure outlined
by C. G. Papaioannou. The crude product was recrystallized three

times from 30% aqueous ethanol.

7. N,N-Nitrosomethyl urea (44)
N,N-Nitrosomethyl urea was synthesized according to the

procedure outlined by F. Arndt in Organic Syntheses; yield - 68%.

 

8. Methyl 8-deuterio-1-naphthoate (45)

A mixture of 5.8 ml 40% aqueous potassium hydroxide and
20 ml ether was added to a 50 ml erlenmeyer flask. The flask was
cooled in an ice bath to <5° and the temperature maintained while
2.0 g N,N-nitrosomethyl urea was added slowly. The ether layer was
decanted and added to a solution of 1.66 g (.0096 mole) 8-deuterio—
1-naphthoic acid in 60 m1 ether. The mixture was stirred at <5° for
one hour, then at room temperature for two hours. The excess diazo—
methane was destroyed by the dropwise addition of glacial acetic
acid. The ether solution was washed with two 20 m1 portions 8%
aqueous potassium bicarbonate, 25 ml water, and dried over anhydrous
magnesium sulfate. The ether was removed by distillation and the
crude product fractionally distilled under vacuum; yield a 1.22 g
(68%), bp 124° (1.3 mm). The ester was purified by vapor phase
chromatography on an Aerograph A 90-P3 by using a six foot x 1/4

inch carbowax column. NMR showed the ester to be >98% d1.

. 3"

*_ '..- m 427‘ -I.h

31

B. Methyl-d_ 1-naphthoate

3

1. Silver l-naphthoate
Silver l-naphthoate was synthesized by the procedure

outlined previously for silver 8-bromo-1-naphthoate; yield = 89%.

2. Methyl—d_ 1-naphthoate

3
A mixture of 2.44 g (.0087 mole) silver l-naphthoate,

1.31 g (.0090 mole) methy15g_ iodide and 30 m1 ethyl acetate was

3
added to a 50 m1 round-bottomed flask (equipped with magnetic

stirrer and reflux condenser). The mixture was refluxed for 24

hours and filtered. The yellow precipitate was washed with two

15 ml portions of ethyl acetate. The solvent was removed from the
combined solutions by distillation. The crude yield was 1.32 g (80%).
The ester was purified by vapor phase chromatography on an

Aerograph A 90-P3 by using a six foot x 1/4 inch carbowax column.

NMR showed the ester to be >98% d3.

C. Methyl 8-methyl-d3-l-naphthoate (42)

The samples of methyl 8-methyl-l-naphthoate and methyl
8—methyl-d3-l-naphthoate used in this study were synthesized by
C. G. Papaioannou. The esters were purified by vapor phase chroma-
tography on a Hewlett Packard F & M Scientific 700 Laboratory
Chromatograph by using a six foot x 1/4 inch carbowax column. NMR

showed the deuterated ester to be >98% g3.

32

D. 2,2-Dimethylpropy1-g_ 1-naphthoate

ll
1. Pinacol-fi12 hydrate (46)
Pinacol-d hydrate was synthesized according to the

-12
procedure outlined by R. Adams in Organic Syntheses by using

 

acetoneed6; yield = 41%.

2. Pinacolone-gi12 (46)

Pinacolone-g_ was synthesized according to the procedure

12
outlined by R. Adams in Organic Syntheses; yield - 70%.

 

3. Pivalic-d9 acid (46)

To a 50 m1, three-necked round-bottomed flask (equipped
with magnetic stirrer, addition funnel, fractionating column, and
thermometer) was added a solution of 3.1 g (.0775 mole) sodium
hydroxide in 25 ml water. The solution was cooled to 0° in an ice-
salt bath and 4.45 g (.0279 mole) bromine was added at such a rate
that the temperature never exceeded 5°. After cooling again to 0°,
0.91 g (.0084 mole) pinacolone-—_d__12 was added and the reaction
mixture stirred at 0° until it was decolorized. Then, 3.7 ml concen-
trated sulfuric acid was added dropwise and the acidic solution
extracted with four 10 ml portions ether. The combined ether solu-
tions were washed with three 10 ml portions 10% aqueous sodium
hydroxide. The combined aqueous extracts were then acidified with
concentrated hydrochloric acid and extracted with four 10 ml portions
ether. The combined ether extracts were washed with 10 ml 20% aqueous

sodium thiosulfate and 10 ml water. After drying the ether solution

over anhydrous magnesium sulfate, the ether was removed by

 

33

distillation; yield - 0.67 g (72%).

4. 2,2-Dimethy1propanol-g11
A slurry of 0.35 g (.0083 mole) lithium aluminum deuteride
and 15 ml anhydrous ether was added to a 50 m1, three-necked round—
bottomed flask (equipped with reflux condenser, addition funnel,
magnetic stirrer, and calcium Sulfate drying tube). A solution of
0.67 g (.0059 mole) pivalic-d9 acid in 10 ml anhydrous ether was
added dropwise at such a rate so as to maintain a gentle reflux.
After the addition was complete, the mixture was refluxed for four
hours. While cooling the reaction in an ice bath, 1.45 ml water
was added dropwise. The mixture was then stirred at room tempera-
ture until white crystalline salts formed. The lithium salts were
filtered and washed with two 10 m1 portions ether. The combined

ether solutions were dried over anhydrous magnesium sulfate and the

ether removed by distillation; yield - 0.48 g (82%).

5. l-Naphthoyl chloride

l-Naphthoic acid was refluxed for 20 hours with a two-fold
excess of thionyl chloride. The excess thionyl chloride was removed
by distillation and the crude product was fractionally distilled

under vacuum; yield - 91%, bp 203-5° (65 mm).

6. 2,2-Dimethylpropylfdll

To a 25 ml, three-necked round-bottomed flask (equipped

l-naphthoate

with reflux condenser, addition funnel, and magnetic stirrer) was
added a solution of 0.93 g (.0049 mole) l-naphthoyl chloride in

10 ml dry pyridine. A solution of 0.49 g (.0050 mole)

34

2,2—dimethy1propanol-g11

The mixture was heated slowly to reflux, then cooled to room

in 10 ml dry pyridine was added dropwise.

temperature and the solution poured into 50 ml water. The aqueous
mixture was extracted with three 25 ml portions ether. The combined
ether extracts were washed with three 25 ml portions 10% hydrochloric
acid, two 15 m1 portions saturated aqueous sodium bicarbonate, and

15 ml water. The ether solution was dried over anhydrous magnesium
sulfate and the ether removed by distillation. The crude yield was
1.13 g (90%). The crude product was recrystallized four times from
70% aqueous ethanol; mp 51-51.5°. The hydrogen compound melted at

52.5—53°. NMR showed the ester to be >96% Q11.

E. 2,2--Dimethy1propyl-l,lug2 l-naphthoate

The ester was synthesized by the same procedure outlined for

the synthesis of 2,2—dimethylpropyl-d

-1l l-naphthoate; mp 52.5-53°.

NMR showed the ester to be >98% d2.

F. 2,2-Dimethylpropyl Segfl-naphthoate

The ester was synthesized by the same procedure outlined for

the synthesis of 2,2-dimethylpropylfg_ l-naphthoate by using

11
8-deuterio—l-naphthoyl chloride, mp 52.5-53°. NMR showed the

ester to be >98% d1.

RESULTS AND DISCUSSION

A serious problem is encountered when attempting to interpret
secondary kinetic isotope effects in terms of reaction mechanisms.
This problem is partially due to our lack of understanding concerning

the effects of deuterium substitution on enthalpies and entropies

.9144 6 .vxncJ’ ‘MvmwA1;f
.4

of activation. It was shown previously that secondary kinetic
isotope effects may vary considerably with temperature and solvent.
The data reported by Evans (2) and Smith (38) where kH/kD increases
with increasing temperature cogently illustrated this point. It is
obvious, therefore, that mechanistic interpretations based on an
isotope effect at a single temperature are meaningless and may
often lead to the wrong conclusions. Furthermore, the possible
solvent dependence of a secondary kinetic isotOpe effect must also
be investigated if any realistic conclusions are to be drawn
concerning the mechanism of the reaction under investigation.

Before secondary kinetic isotope effects may be classified
as a useful mechanistic tool, their origin must be well understood.
Due to the considerable controversy over the nature of secondary
kinetic isotope effects, one is left with a serious doubt as to its
usefulness in mechanistic interpretations. In order to probe further
into the origin of secondary isotope effects, reactions must be
studied in which the mechanism is fully established by other methods
and does not rely on secondary kinetic isotope effects.

35

 

36

I. Mechanism for the Base Hydrolysis of Alkyl Esters

Unlike many reactions in organic chemistry, the mechanism for
the alkaline hydrolysis of alkyl oxygen esters is well established.
The reaction is believed to proceed via rate determining attack of
hydroxide at the carbonyl carbon of the ester function. Thus, in
the rate determining step, a neutral substrate is converted to a
tetrahedral anion in which the steric interactions should increase.
This reaction is, therefore, of considerable interest in terms of
secondary kinetic isotope effects originating from non-bonded
interactions.

Based on the second order kinetic behavior first reported by
Warder (49) and the oxygen labeling experiments of Bender (47) and
of Polanyi and Szabo (48), Moffat and Hunt (50) have proposed the

following reaction scheme for the base hydrolysis of alkyl esters.

180 I' 180- — 180
II _ k] I k ll _
R-c-ocu3 + OH ‘__ R-C-OCH3 .1, R—C—OH + 00113
k
2 OH
.11..
180H
I
R-C-OCH
. 3
OH
kSle4
18
II 1 _ .1 i“ .3 II _
R-c-ocn3 + 80H *— R-C'J-OCHa -—-> R-cmon + ocn3
k .-
2 0

 

 

37
Since oxygen-18 exchange occurs for these reactions, the life-
time of the tetrahedral anion must be significantly longer than the
time for a molecular vibration. The tetrahedral anion is, therefore,
classified as a true intermediate. The kinetic importance of the
tetrahedral intermediate is further supported by the observation of
a change in the rate determining step with varying pH for the base

hydrolysis of several amides (51).

Table 12: Effects of para-substituents on the rates of alkaline
hydrolysis of ethyl benzoates at 25° (53).

 

 

 

pfsubstituent Relative Rate EA
NH2 0.023 20.00
OCH3 0.21 18.65
CH3 0.45 18.20
H 1.00 17.70
Br 5.25 16.80
NO2 100.3 14.50

 

Another consequence arises from the proposed mechanism for the
base hydrolysis of alkyl esters. In forming the transition state
leading to the tetrahedral intermediate, the electron density at
the carbonyl carbon is increased considerably. As a result, the
hydrolysis reaction should be facilitated by electron withdrawing

substituents. The effects of para-substituents on the rates of

 

38
base hydrolysis of ethyl benzoate, Table 12, clearly illustrates
the sensitivity of these reactions to electronic interactions,
0 - +2.56. This is also reflected in the relative rates of methyl

acetate, methyl chloroacetate, and methyl dichloroacetate, Table 13.

Table 13: Substituent effects on the rates of alkaline hydrolysis
of alkyl esters at 25° (52). i

 

 

 

Compound . k/k(MeOAc) k/k(EtOAc)
CH3COZCH3 1.0 1.67
CICH2C020H3 761
CIZCHCOZCH3 16,000
CH302CC02CH3 170,000
CH3COZCH2CH3 1.00
CH3COZCHZCH(CH3)2 0.70
CH3002CH2C(CH3)3 0.18
CH3CH2C02CH2CH3 0.47
(CH3)ZCHC02CH2CH3 0.10
(CH3)3CC020H2CH3 0.011

 

The rate retardations of alkyl substituents on the acid portion
of the ester, Table 13, may be safely attributed to a combination
of inductive and steric effects. The effects of alkyl substitution
on the alcohol moiety, however, are generally ascribed to steric

factors alone.

 

39
The absence of hyperconjugative structures involving the

8-substituents for alkyl l-naphthoates makes these compounds ideal

O R

for studying the magnitude of secondary kinetic isotope effects
arising from non—bonded interactions. The rates of hydrolysis of
methyl 1-naphthoate, methyl 8-methyl-1-naphthoate, 2,2-dimethyl-
propyl l-naphthoate, and their deuterated analogs are reported in
Tables 14-17. The variations of kH/kD with solvent and temperature
are summarized in Tables 18 and 19. The thermodynamic parameters
calculated by the least-squares programs AKTIV, HANDS, and KINFIT
are reported in Tables 20—22. The errors reported for the
thermodynamic parameters are two standard deviations yielding

95% confidence limits.

II. Correlation of Thermodynamic Parameters

Following the analysis of Gore and coworkers (54), we will
discuss separately each effect of substitution on the enthalpies
and entropies of activation, AH* and AS*. The following factors
Operate in the base hydrolysis of alkyl esters of aromatic acids.
(a) Inductive effects (+I) that donate electrons to the carbonyl
carbon thus stabilizing the ground state and increasing AH*.

(b) Polar effects that increase the solvation of the transition

state, thus decreasing AH*. These effects are usually small when

 

40

Table 14: Rates of base hydrolysis of methyl 1-naphthoate and
methyl 8fgfl-naphthoate in aqueous methanol.

 

 

10.0% Methanol/Water (w/w)

 

3 -1 1

 

 

 

 

 

Temp.(°C) kH x 10 sec kD x 103 sec- kH/kD
15.571 i .006 2.9548 i .0385 2.9897 i .0285 .988 i .016
35.728 i .008 12.113 i .084 12.257 i .148 .988 i .014
50.021 1 .009 28.691 i .289 28.815 1 .376 .996 i .016

30.0% Methanol/Water (w/w)

Temp.(°C) kH x 104 sec.1 kD x 104 sec”1 kH/kD
15.565 i .006 6.1733 i .0515 6.3842 1 .0361 .967 i .010
35.722 1 .008 32.108 i .421 32.471 i .244 .989 i .015
50.011 i .010 85.663 1 .947 85.576 1 .402 1.001 t .012

 

- n'.’?5t!l ‘— Iflu Arazmuum
‘ .I

'r
it“,

 

41

Table 14: (cont'd.)

 

 

48.1% Methanol/Water (w/w)

 

 

 

 

 

 

Temp.(°C) kH x 104 sec”1 kD x 104 sec.”1 kH/kD
15.565 i .006 1.4176 i .0136 1.4654 1 .0030 .965 i .009
26.361 i .008 4.3323 1 .0194 4.4232 1 .0201 .979 i .006
35.729 1 .008 9.7586 i .0375 9.9154 1 .0946 .984 i .010
41.141 1 .010 14.447 1 .060 14.622 1 .036 .988 i .006
50.007 1 .012 28.875 1 .112 28.941 1 .522 .998 i .018
57.963 i .010 58.216 1 .194 57.914 1 .700 1.005 t .013
66.629 1 .012 100.85 i .89 100.66 1 1.92 1.002 t .021

80.0% Methanol/water (w/w)
o 5 -1 5 -1

Temp.( C) kH x 10 sec kD x 10 sec kH/kD
15.577 i .006 1.5212 1 .0156 1.5293 i .0147 .995 i .014
35.722 1 .008 11.726 i .110 11.745 i .051 .998 i .010
50.014 1 .012 47.108 1 .473 47.180 1 .804 .998 i .020

 

42

 

 

 

 

OOO. N NNO.N OOO. N NOO. ONN. N ONO.NO OON. N OON.ON NNN. N ONO.ON ONO. N NO0.0m
ONO. N mNO.N OOO. N NOO. NNN. N ONN.ON NON. N ONN.ON NNN. N www.mN ONO. N ONO.mm
OOO. N OmO.N OOO. N NON. OONO. N ONON.O NOmO. N mOmN.O OONO. N ONOm.O OOO. N NNN.ON
mOOO O O N 6m. N 666 N 666
I I I I I .986
1\Ox s\mx ON x OOON ON 1 O ON ON x xx NO.V a
O O N
.A3\3V amum3\msmxofic NH.qm 6H oumonunamalfl mmlflanuma

was

.oumosunamcINLwIm Hugues .OumonunnmaIH stume mo mNmzaoupzs ommn mo amumm umH canny

43

Table 16: Rates of base hydrolysis of 2,2-dimethy1propyl

1-naphthoate, 2,2-dimethylpropyl-l,l-g_ l-naphthoate,

2,2-dimethy1propyl—Qn11

8jdfl-naphthoate.

2

l-naphthoate, and 2,2—dimethy1propyl

 

 

54.1% Dioxane/Water (w/w); T 8 24.721 i .010°

 

-1

 

Compound k x 10 kH/kD
2,2—Dimethy1propyl 6.2152 1 .0312
l-naphthoate
2,2-Dimethy1propyl-l,l—d2 5.6563 1 .0414 1.099 t .010
l-naphthoate
2,2—Dimethylpropyl-d 5.3868 1 .0284 1.154 t .008
-ll
l-naphthoate
2,2-Dimethylpropyl 6.2646 1 .0332 .992 i .007

8fdfl-naphthoate

 

 

54.1% Dioxane/Water (w/W);

T - 35.404 1 .008°

 

 

Compound k x 104 -1 kH/kD
2,2-Dimethylpropyl 1.3128 1 .0057
l-naphthoate
2.2-Diuwzthlen'opyl-l.l-g2 1.2091 1 .0057 1.086 1 .007
l—naphthoate
2,2-Dimethylpr0pyl-g_ 1.1224 1 .0048 1.170 t .007
11
l-naphthoate
2.2-Dimethylpr09y1 1.3238 1 .0053 .992 1 .006

8fidfl-naphthoate

 

 

44

Table 16: (cont'd.)

 

 

54.1% Dioxane/Water (w/w); T 8 50.016 1 .010°

 

4 -1

 

 

 

 

 

Compound k x 10 sec kH/kD
2,2-Dimethy1propyl 3.2965 1 .0120
l-naphthoate
2,2—Dimethylpropyl-l,l-g_2 2.9851 1 .0214 1.104 r .009
1-naphthoate
2,2-Dimethylpropyl-d 2.8846 1 .0126 1.143 t .006
—11
1-naphthoate
2,2-Dimethylpropy1 3.3178 1 .0146 .994 i .006
88drl-naphthoate
30.0% Methanol/Water (w/w); T 8 50.006 1 .010
4 -1
Compound k x 10 sec kH/RD
2,2-Dimethylpropyl 6.8492 1 .0356
l-naphthoate
2,2-Dimethylpropyl-l,l-lc_l2 6.2258 1 .0753 1.100 t .014
l-naphthoate
2,2-Dimethylpropyl-d 5.8134 i .0458 1.178 i .011

l-naphthoate —11

 

 

45

Table 16: (cont'd.)

 

 

48.1% Methanol/Water (w/w); T 8 50.016 i .010°

 

4 -1

 

Compound k x 10 sec kH/kD
2,2-Dimethy1propyl 3.2241 1 .0182
l-naphthoate
2,2-Dimethylpropyl-l,l-d 2.8396 1 .0126 1.135 1 .008
-2
l-naphthoate
2,2-Dimethy1propylfd_ 2.6961 i .0155 1.196 i .010

l-naphthoate 11

 

Table 17: Rates of base hydrolysis of methyl 8-methy1-l-naphthoate
and methyl 8-methy18d3-l-naphthoate in 48.1% Methanol/water (w/w).

 

 

 

Temp.(°C) kH x 106 sec.1 kD x 106 sec.1 kH/kD
70.201 1 .042 2.1663 1 .0115 2.2035 1 .0302 .983 i .014
84.985 1 .030 7.8550 1 .0598 8.0862 1 .0463 .971 i .009

100.384 1 .020 25.864 1 .204 26.752 1 .186 .967 i .010

 

.umums was uam>aom poems use no mowumu uswwoa mum muco>aom ones

 

46

 

NNO. N NOO.N O0.00
NNO. N NOO.N OO.NO
OOO. N OOO. ONO. N OOO. ONO. N OOO. NNO. N NOO.N ONO. N OOO. N0.0m
OOO. N OOO. NN.NN
OOO. N NOO. ONO. N OOO. ONO. N OOO. ONO. N OOO. ONO. N OOO. ON.OO
AOO. N OOO. OOO. N ONO. O0.0N
NNO. N OOO. OOO. N OOO. ONO. N NOO. ONO. N OOO. NO.mN

NOONONO NN.OO mom: N0.00 mom: NN.OO moo: N0.00 NOON: N0.0N NO.O.Oame

 

 

.mumonunamcIHLmrw

assume you aoNuNmoaaoo uco>Nom was ousumumaaou saws nx\=x mo cowumwum> ”ma manna

47

.o<.w~ u unaumumaaman

.Nmums was uao>aom poem: onu mo mowumu uswaos mum muso>Nom mafia

 

oumonusawsINLmrm

 

NOO. N NOO. OOO. N NOO. OOO. N OOO. OOOOONO NN.NO NOOONONONNNONOIN.N
OOO. N OON.N NOO. N ONN.N OOO. N OON.N NONNONO ON.NO

ONO. N OON.N OONz NN.ON NNI. NNOoeNnOOaIN

NNO. N ONN.N moo: N0.00 OINOOONONONNNNNOIN.N
ONO. N OOO.N NOO. N OOO.N OOO. N OON.N NOONONO ON.OO

OOO. N OON.N mom: NN.OO OI NNOOONOONOIN

ONO. N OON.N OONz NO.OO OIN.N-NOO6NONONNNaNOIN.N
OOO. N OOO.N ONO. N ONO.N OOO. N ONO.N NONNONO NN.OO NOOONNOOOOIN OmrNOeNmz
OOO. N OOO. OOO. N NOO. OOO. N OOO. NOONONO ON.OO
OOOO. N ONO. ONO. N OOO. ONO. N OOO. moo: NN.ON I.

ONO. N OOO. NNO. N NOO.N moo: N0.00 NOOOONNOOOINIOIO NOON»:

N.N.ONO ON\ON N.N.OOO ON\ON No0.000 ON\ON NNON>NOO OOOOOaOO

 

 

.mmumozusamala Humouaaunuoaamlm.~
was mmumonunamsIN assume you coNuNmanoo usm>Nom mam ousumumneou nua3 nx\mx mo soaumfium> "ma nanny

48

.Noumz paw uso>Nom poem: one no mowumu unwaos mum muco>Hom use

a

.OuNaNN oosovaaoo Nmm mawvaoNz maoNumN>mm mumwsmum o3» mum maouum vouuoand

 

 

ON. N ON.OOI OO. N OO.NN NO. N NO.OOI NO. N OO.NN OOOOONO NN.OO
OO.N N OO.ONI NO. N NO.NN NO.N N ON.ONI OO. N OO.NN NON: N0.00
ON.N N NO.NNI OO. N OO.NN ON.N N OO.NN- OO. N OO.NN mom: NN.OO

OO. N NO.ONI ON. N OO.NN NO. N OO.ONI ON. N OO.NN moo: N0.00

NO. N ON.ONI ON. N OO.NN OO. N NO.ONI ON. N OO.NN NON: N0.0N

MOO MOO MOO MOO ONON>NOO

 

 

n.0umo5unmmaIHLwrw Nmsuos was oumonusamcIH Nanuoa Now muouoamuma oNamshvoauony "om manna

49

.muNaNN moaovasoo Nmm manNoNO msoaumN>ov unavamum

1.10.1m 1

.uouma was uno>Hom means can mo mowumu

 

.HHmzHM amuwoumv

.mnz<x amuwoumo

unmwma mum muao>HoO one

a

o3» mum mnouuo wouuoqomw

 

 

OO. N O0.0 ONN N O ON. N N0.0 NN N NN NOOOONO NN.OO
OO. N O0.0 ON N ON NO. N O0.0 O N NN NON: N0.00
ON. N O0.0 ON N NON OO. N N0.0 ON N OON OOOz_NN.OO
OO. N O0.0 ON N NON OO. N O0.0 ON N NON mom: N0.0N
ON. N NN.O NO N ON NN. N NN.O Om N ON moms N0.0N
ONOOO ONOOO ONOOO ONOOO Ouao>NoO

 

 

NN.O.NOOOO uON ONOON

 

50

 

1.1....NI ill-l.

.ONNsNN NOONONNOOO NOO OONONONN

maoNuwN>mv vumvsmum oBu mum muouuo wouuonomm

 

 

OO. N NO.NN- NN. N OO.NN NNOONNOOOOINLer NNOONONNNNNaNOIN.N
ON. N NO.NN- OO. N OO.NN NNOONNOOOOIN NNOrNNOONONNONNaNOIN.N
OO. N OO.NN- ON. N OO.NN «NOONNOOOOIN NmrN.NINNOONONNONNaNOIN.N
Om. N OO.NN- ON. N OO.NN NNOOONOOOOIN NNOONONNONOaNOIN.N
ON. N NN.OO- NN. N NO.NN NuOoeuNOOaIN OmrNNONoz
ON. N OO.ON- OO. N OO.NN ONOOONOOOOINLer NNONNO
NO. N NN.OOI NO. N NO.NN NNOOONNONOIN NNONNz
+m< +m< vaaoaaoo

 

 

O.A3\3v noum3\osmonv NN.On aN moumonunauana
NOOONQNOLNOENOIN.N was Ooumonunnmsua Nasuos now Ououmsmuma oNsmczvosNonh “Hm oNan

51

 

.ONNsNN ONONONOOOO NOO OONONNNN

.NNNzNN sONOONOu

.mnz<m Emuwoumn

OnowuwN>ov eumpamum osu one Ououuo vuuuoaomu

 

mumosunnmaIHLmrm

 

OO. N OO.O ON N NN NO. N O0.0 O N ON NNOONONNNNNENOIN.N
NH oumonunnmcIH
OO. N NO.O- OON N NO- OO. N NO.O OON N OO- OINNOONONNONNaNOIN.N
fil oumonunmmaIH
NO. N NN.O OON N OO OO. N NO.O OON N NO OIN.NINNO6NONNONNNNOIN.N
OO. N ON.O- OON N OO- OO. N ON.OI ONN N OO- NNOOON;ONOIN OmrNNONNz
OO. N OO.O ONN N O ON. N NO.O NN N NN ONOOONNOOOINLOIO NNsuoz
anoaao
ONOOO ONOOO ONOOO ONOOO O O

 

 

NN.O.NNONO .NN NNONN

52

Table 22: Thermodynamic parameters for methyl 8-methyl-l-naphthoate
and methyl 8-methyl-d_-l-naphthoate in 48.1% MeOH/water (w/w).a

 

 

 

 

 

 

3
Compound AH* AS*
Methyl 8-methyl-1—naphthoate 20.23 i .24 -25.80 i .70
Methyl 8-methyl-d3-1-naphthoate 20.37 i .28 -25.35 i .80
Program AAH¢ AAS*
HANDS -139 i 44 -O.44 i .12
KINFIT -135 i 82 -0.43 i .24

 

aReported errors are two standard deviations yielding 95%

confidence limits.

 

53

the substituent is alkyl. (c) Primary steric effects that result

 

from an increase in non-bonded interactions in forming the tetra-
hedral intermediate, thus increasing AH7. These effects should
make AS* more negative by decreasing the number of available energy
levels of rotation in the transition state, i.e., by hindering

internal rotation. (d) Secondary steric effects that force the

 

carbonyl group out of coplanarity with a ring or center of unsatura-
tion, thus decreasing the electron density at the carbonyl carbon

and raising the energy of the ground state. This effect will operate
primarily in the ground state and results in a smaller AH* and a

more negative AS*. The importance of this effect is reflected in

the observed thermodynamic parameters in Table 23. When the ester
function is in a position where it must interact with a perifhydrogen,
the rate is significantly retarded. However, the decrease in rate
results from a more negative entropy and not from an increase in
enthalpy. This is partially attributed to the opposing effects (c)
and (d). Bulky substituents may also reduce the stability of the
transition state by inhibiting formation of the solvent shell

around the transition state, thus leading to (e) effects that

 

sterically hinder salvation. Since the solvent is less efficient

in stabilizing the developing charge in the transition state, AH¢

will increase while A87 will become more positive.

 

These factors are also expected to influence the secondary
isotOpe effects observed for the base hydrolysis of alkyl esters.
Due to the complexity of the interaction mechanisms, it is evident

that only a qualitative interpretation can be offered at this point.

54

Table 23: Steric effects in the base hydrolysis of some alkyl

 

 

 

esters .
Ester Relative AH* AS*
Rates
Ethyl 2-naphthoatea 1.0 16.4 -l7.1
Ethyl l-naphthaatea 0.32 16.3 -19.8
Ethyl 2-phenanthroatea 1.0 17.2 -l4.5
Ethyl 9-phenanthroatea 0.54 15.6 -21.2
Ethyl benzoateb 1.0 17.1 -15.9
Ethyl 2-methylbenzoateb 0.12 17.3 -19.2
Ethyl 3—methylbenzoateb 0.70 17.4 -15.5
Methyl 2-anthroquinone 1.0 12.7 -22
carboxylatec
Methyl l-anthrgquinone 0.0011 9.8 -45
carboxylate

 

8Reference 55. bReference 56. cReference 57.

III. Temperature Dependence

Examination of AH* and A37 for methyl l-naphthoate and methyl
8-methyl-1-naphthoate in 48.1% aqueous methanol indicates that the
difference in rates (~5000) is due to both an increase in AH* and
a decrease in AS*. This is consistent with the primary steric effect
discussed previously, which is taken here as the major reason for
the rate difference, as well as the secondary steric effect and the

effect of steric hindrance to salvation. The importance of solva-

tion in these hydrolyses is illustrated by examining the solvent

6
1nd”.

 

 

 

 

 

55
l T I I
24 - 1
20 t N
16 I- -
”a?
.-I
o
E
2
3‘2
>~ fl: 23.
00
H
0)
t5
8 - N
4 N- -I
0 l I I l
O .2 .4 .6 .8 100

Ma le Fraction Water

Figure 5: Solvent effects on the thermodynamic parameters for the

base hydrolysis of methyl l-naphthoate in aqueous methanol

solutions at 25°. 0 8 AG*, A 8 AH*, CI 8 -TAS*

‘A—

 

56

 

 

AH* AS*
Methyl l—naphthoate 15.59 -21.98
Methyl 8-methyl-l—naphthoate 20.23 -25.80

dependence of the thermodynamic parameters for methyl l—naphthoate,
Figure 5. AH* and TAS* change dramatically with increasing solvent
polarity. However, since they change in apposite directions, they
result in a much smaller change in A07.

The difference in rates of hydrolysis of methyl l-naphthoate
and 2,2-dimethylpropyl l-naphthoate in 54.1% aqueous dioxane (N14)
is not due to an increase in the enthalpy of activation. AH* actually
decreases for the bulkier ester. The rate difference is entirely

1

due to a difference of 7 e.u. in AS . This observation is not

unprecedented and has been observed in similar systems, Table 24.

Table 24: Effects of chain branching on the rates of base hydrolysis
of some alkyl esters in 70% aqueous dioxane at 40° (54).

 

 

 

Ester Relative AH* AS*
Rates

Methyl l-naphthoate 1.0 14.5 -21.2
Ethyl l-naphthoate 0.36 12.5 -29.7
IsOpropyl l-naphthoate 0.059 12.5 -33.3
Methyl benzoate 1.0 12.5 -25.9
Ethyl benzoate 0.37 13.0 -26.2
IsoprOpyl benzoate 0.064 13.8 -27.1

thutyl benzoate 0.0029 15.7 -27.3

 

 

57

In the l-naphthyl system where the ester function must interact
with a pggifhydrogen, the effect of chain branching on the alcohol
moiety is to greatly decrease AS*, while slightly decreasing AH*.
Although these effects may be explained in terms of the opposing
steric effects, i.e., primary and secondary, such an explanation is
inconsistent with the parameters observed for methyl 8-methyl-l-
naphthoate. It seems reasonable that the secondary steric effect
should operate to a greater extent in the latter compound, since
the 8-methyl substituent is in a better position to inhibit conju-
gation of the carbonyl entity with the naphthalene ring system.
Yet, AH* increases for the 8~methy1 ester and AS* decreases
moderately. It appears that some other factor may be influencing
the thermodynamic parameters for 2,2-dimethylpropyl l-naphthoate.

The large difference in the isotope effects for methyl-d

-3

l—naphthoate, XIII, and 2,2-dimethylpropy1-1,l-£l_2 l-naphthoate, XIV,

is inconsistent with the mechanistic assumptions for this reaction.

Based upon the mechanism outlined by Moffat and Hunt (50), it was

COZCD3 COZCDZCCCH3)3

XIII XIV

kH/kD = 1.030; 24.7° kH/kD = 1.099; 24.7°

expected that kH/kD should be greater for XIII, due to the greater

number of deuterium atoms. This was not observed experimentally.

58

This anomaly can be easily rationalized if there is a change in
the rate determining step for the hydrolysis of XIV. Then, the
greater inductive effect of deuterium will operate so as to create
a difference in the stabilities of the developing anions, resulting
in a significant normal isotope effect. The origin of the isotope
effect will be discussed in more detail below.

The proposal of a change in the rate determining step for the
hydrolysis of the neopentyl ester is not unrealistic in view of the

predicted difference in stabilities of the two corresponding anions,

‘s.-‘ ,1-.-_,.- . u!__u-m-
. _ -

CH3O and (CH3)3CCH2

not been determined, it seems realistic that it should be signifi-

0-. Although the pKa for neopentyl alcohol has

cantly greater than that of ethanol, pKa - 17, which is ten times
less acidic than methanol, pKa - 16 (58). Furthermore, changes in
the rate determining step have been observed previously for the
base hydrolysis of several amides (51) and for the aminolysis of
substituted phenyl benzoates (59). All of these reactions are

known to proceed through a tetrahedral intermediate.
IV. Interpretation of Isotope Effects
A. Methyl Bedfl-naphthoate

A previous study of the secondary isotope effect for the base
hydrolysis of methyl Bagrl-naphthoate in 56% aqueous acetone reported
kH/kD - 1.00 i .01 at 25° (32). The present investigation is con-
sistent with that result. From Table 18, we see that kH/kD depends
on the solvent composition. Unfortunately, because of the magnitude

of the isotope effects and their errors, any speculation as to the

59
nature of this dependence is meaningless.

From the variation of kH/kD with temperature in 48.12 aqueous
methanol, we find that the isotope effect decreases with increasing
temperature. This trend is what is normally observed, however,
the change in kH/kD is more pronounced than the exponential drift
predicted if AAG¢ = AAH*. Indeed, an examination of the thermo-
dynamic parameters reveals that AABat and AAS* oppose each other,
157 cal/mole and 0.48 e.u., respectively, to result in a relatively
small value of AAG*. The isotope effect predicted on the basis
of AAH* alone is kH/kD . .77 at 25°. The factors determining this
isotOpe effect, therefore, affect both the enthalpy and entropy of
activation.

An inverse isotope effect would be predicted based on the
primary steric effect, that is, the difference in steric hindrance
to formation of the transition state. However, if this were the
origin of the effect, then we would expect AAH* > 0 and AAS¢ < 0.
Although AAH* is in the predicted direction, AAS* is not. Following
the analysis discussed previously, one may expect that there may be
a difference in the dihedral angle between the carbonyl function and
the plane of the naphthalene ring system for the deuterated and the
undeuterated compounds in the ground state, a secondary steric
effect. This would tend to decrease AAH* and make AAS* even more
negative. Therefore, an explanation of the isotope effect, which
is consistent with the observed thermodynamic parameters, must
entail some other factor.

Based on substituent effects on the aromatic proton chemical

shifts of naphthalene, W. B. Smith and coworkers (60) have recently

._-‘.‘..
I

.—

-—~4-—-4 a 4» Aw-:-l—r"

60
reported that there is an attractive interaction between the carbonyl
oxygen and the 8-hydrogen in l-formyl naphthalene. Assuming that
this type of interaction is also operating in methyl l—naphthoate,
there should be a greater attraction for the 8-hydrogen than for the
8—deuterium due to the lower polarizability of deuterium. Since the
attractive interaction should be diminished in the transition state,
the enthalpy of activation ought to be greater for the hydrogen

compound, as observed. Similarly, the freedom of rotation of the

Lb

ester function in the ground state should be greater for the deuter— a

ium compound, resulting in a more negative AS*. Although this

explanation is consistent with the calculated thermodynamic parameters, ;:9

the author wishes to point out that this is only a possible explana-

tion and is subject to controversy.
B. Methyl 8-methy1fid3-1-naphthoate

The results obtained for methyl 8-methy15g3—l-naphthoate are
inconsistent with those reported by Karabatsos and Papaioannou,
Table 25 (61). This may be partially due to the different methods

and solvent used for the measurement of kH/kD.

Table 25: Isotope effects for the base hydrolysis of methyl
8-methylgg3-l-naphthoate in ethanol (61).

 

 

 

Temp.( c) kH/kD
61.5 i 0.5 .84 i .02
79.5 i 0.5 .89 i .02

 

 

 

61

Like methyl Begfl-naphthoate, the small inverse isotope effect
for methyl 8-methyleg3-l-naphthoate is the result of opposing
enthalpy and entropy effects. However, the values of AAH* and AAS*
are opposite in sign, —135 cal/mole and -0.44 e.u., respectively,
from those observed with methyl 8fdfl-naphthoate. The small experi-
mental effect arises from the dominance of the entropy term in AAG*.
This inverse effect may be attributed to a combination of the various
steric factors discussed in part II of this section. The primary
steric effect, steric hindrance to formation of the transition
state, is counterbalanced by the secondary steric effect and the
difference in steric hindrance to salvation. The latter two effects
would tend to make AAH* < 0 while having little effect on AAS*.

The significance of the isotope effect for this ester is the
fact that it is small because of the opposing contributions of
enthalpy and entropy to AAG*. A similar observation was reported
by Smith (38) for the solvolysis of the egggrnorbornyl p-nitro-
benzoate XII, AAH* - 157 cal/mole and AAS* 8 0.44 e.u. A large
effect was predicted originating from non-bonded interactions.
However, steric factors apparently affected both AAH¢ and AAS*,
resulting in a rather small isotope effect.

As previously reported, Mislow and coworkers (31) found that
AAS* made a considerable contribution to AAG*, Opposing the effect

4;

of AAH¢ (AAH - 240 cal/mole; AAS=|= = 0.53 e.u.), for the racemi-

zation of 4,5-dimethylfig6—9,lO-dihydrophenanthrene, X, in benzene.
Similar results were reported by Carter and Dahlgren (30) for the

racemization of 2,2'-binaphthy1fg , IX, in dimethylformamide

¢ 4:

(AAH = 270 cal/mole; AAS - 0.54 e.u.). Non-bonded interactions,

ugh fi-myvu-L.-9r. 1* 1°
.1

 

62
therefore, appear to have a significant effect upon the zero-point
energy differences for secondary kinetic isotope effects. However,
these interactions also affect the entropy of activation so as to
nearly cancel the zero-point energy differences.

These results indicate that secondary kinetic isotope effects
originating from non-bonded interactions will be insignificant when
other electronic factors operate. Only when severe crowding is
occurring in the transition state, as in the racemization reactions

mentioned, will non-bonded interactions be important.

C. Methyl-g3 1-naphthoate

Due to the magnitude of the errors in the thermodynamic

parameters for methyl-1d3

be drawn regarding the origin of the small normal isotope effect.

1-naphthoate, no definite conclusions can

It is reasonable to assume, however, that the major factor involved
would be the greater inductive effect of deuterium. Thus, the
electron density on the carbonyl carbon atom in the ground state
would be greater for the deuterated compound and a normal isotope

effect would be predicted.

D. 2,2-Dimethylpropyl-l,ljd l-naphthoate

2

From the difference in kH/kD for 2,2-dimethylpropy1-l,l--‘c_i_2

l-naphthoate, XIV, and methyl-mg3 l-naphthoate, XIII, we conclude
that there is a change in the rate determining step in going from
the methyl to the neopentyl ester. We attribute the isotope

effect for the neopentyl compound to the difference in the induc—

tive effects of deuterium and hydrogen. Since deuterium is more

.. --_._..._.-____ ...___.._ _ _...
‘ Z

63
electropositive than hydrogen, it will donate electrons more
effectively, thus, destabilizing the developing anion in the tran-
sition state. This rationalization is supported by the significant
isotOpe effect on the acidity of formic acid, KH/KD - 1.070, which
is attributed to the greater inductive effect of deuterium, Table 1.
The thermodynamic parameters determined for this reaction are useless

in supporting any further conclusions as to the origin of the effect.

E. 2,2-Dimethylpropyl-d

_11 l-naphthoate

The isotope effect for 2,2-dimethy1propyl-‘<_i11

may be ascribed primarily to the greater inductive effect of deuter-

l-naphthoate

ium and its effect on the stability of the developing anion in the

transition state, as was the case with 2,2-dimethylpropyl-l,1fid2

l-naphthoate. The additional effect of 4-6Z may be attributed to

a combination of the difference in the inductive effects of the

deuterated and undeuterated Eybutyl groups and on the relief of

non-bonded interactions in the transition state. Again, the large
*

errors in AAH and AAS* make any conclusion based on them meaning-

less.
F. 2,2-Dimethylpropyl 8-g71-naphthoate

The isotope effect reported for 2,2-dimethy1propyl 8fgfl-
naphthoate are on the edge of experimental error and for all prac-

tical purposes may be considered nonexistent.

‘1." f. .71:
.

, .' ‘. 33.2-. ,‘Qn

 

64

V. Conclusion

In conclusion, the small inverse isotope effects observed for
the base hydrolysis of methyl Bedfl-naphthoate and methyl 8-methy1-
ga-l-naphthoate were found to result from opposing values of AAH$
and AAS*. The origin of these effects was attributed to a combina-
tion of steric interactions that affect both the enthalpies and
entropies of activation. From the relative magnitude of these
isotope effects in comparison with other systems, it was concluded
that secondary isotope effects arising from non—bonded interactions
will be insignificant when other electronic factors operate, e.g.,
hyperconjugation and induction. Only under conditions of severe
crowding in the transition state will non-bonded interactions be
important.

Furthermore, the difference in the isotope effects for methyl-gm3
l-naphthoate and 2,2-dimethy1propyl-1,l--_c1_2 l-naphthoate was attri-
buted to a change in the rate determining step for the latter
compound. The origin of the normal isotope effect for 2,2-dimethy1-
propyl-1,17d. l-naphthoate was attributed to an inductive effect on

2
the stability of the developing alkoxide ion.

'got ~l9 '

 

 

BIBLIOGRAPHY

(1)

(2)
(3)
(4)

(5)
(6)

(7)

(8)

(9)

(10)
(11)
(12)

(13)

(14)

(15)

(16)

BIBLIOGRAPHY

E. A. Halevi, "Secondary Isotope Effects," Progress in Phys.
Org. Chem., Vol. 1, S. G. Cohen, A. Streitweiser, and R. W.
Taft, ed., John Wiley and Sons, New York (1963), p. 109.

 

 

T. A. Evans, Ph.D. Thesis, Michigan State University, 1968.

D. R. Lide, J. Chem. Phys., 21, 343 (1957).

 

J. W. Simmons and J. H. Goldstein, J. Chem. Phys., 29, 1804
(1952).

 

G. V. D. Tiers, J. Am. Chem. Soc., 12, 5585 (1957).

 

J. A. Dixon and R. W. Schiessler, J. Am. Chem. Soc., 26,
2197 (1954).

 

A. Streitweiser, Jr., J. R. Wo1fe, and W. D. Schaeffer, Tet.,
g, 338 (1959).

A. Streitweiser, Jr. and H. S. Klein, J. Am. Chem. Soc., 85,
2759 (1963).

 

H. S. Klein and A. Streitweiser, Jr., Chem. and Ind. (London),
180 (1961).

 

E. A. Halevi, M. Nussim, and A. Ron, J. Chem. Soc., 866 (1963).

 

E. A. Halevi and B. Ravid, Pure and App. Chem., 8, 339 (1964).

G. V. Calder and T. J. Barton, J. Chem. Ed., 48, 338 (1971).

 

R. P. Bell and W. B. T. Miller, Trans. Far. Soc., 52, 1147
(1963).

 

V. J. Shiner, Jr., B. L. Murr, and G. Heinemann, J. Am. Chem.
Soc., 85, 2413 (1963).

 

V. J. Shiner, Jr. and J. S. Humphrey, Jr., J. Am. Chem. Soc.,
.85, 2416 (1963).

 

a. V. J. Shiner, Jr., J. Am. Chem. Soc., 83, 240 (1961).

b. V. J. Shiner, Jr., J. Am. Chem. Soc., 18, 2653 (1956).
c. A. Streitweiser, Jr., R. H. Jagow, R. C. Fahey, and S.
Suzuki, J. Am. Chem. Soc., 89, 2326 (1958).

 

 

 

65

 

(17)

(18)

(19)

(20)

(21)

(22)

(23)

(24)

(25)

(26)

(27)

(28)

(29)

(30)

(31)

(32)

66

E. M. Arnett, T. Cohen, A. A. Bothner-By, R. D. Bushick, and
G. Sowinski, Chem. and Ind. (London), 473 (1961).

 

E. S. Lewis and G. M. Coppinger, J. Am. Chem. Soc., 18, 4495
(1954).

 

V. J. Shiner, Jr. and G. S. Kriz, Jr., J. Am. Chem. Soc.,_88,
2643 (1964).

V. J. Shiner, Jr., W. E. Buddenbaum, B. L. Murr, and G. Lamaty,
J. Am. Chem. Soc., 88, 418 (1968).

 

V. J. Shiner, Jr. and C. J. Verbanic, J. Am. Chem. Soc., 28,
373 (1957).

 

a. W. A. Sweeney and W. M. Schubert, J. Am. Chem. Soc., 28,

4625 (1954).

b. W. M. Schubert, J. M. Craven, R. G. Minton, and R. B. Murphy,
Tet., 8, 194 (1959).

J. P. Schaefer, M. J. Dagani, and D. S. Weinberg, J. Am. Chem.
Soc., 82, 6938 (1967).

J. P. Schaefer, M. J. Dagani, and D. S. Weinberg, J. Am. Chem.
Soc., 88, 6938 (1968).

 

E. M. Arnett, P. M. Duggleby, and J. J. Burke, J. Am. Chem.
Soc., 88, 1350 (1963).

 

H. C. Brown and G. J. McDonald, J. Am. Chem. Soc., 88, 2514
(1966).

 

H. C. Brown, M. E. Azzaro, J. C. Koelling, and G. J. McDonald,
J. Am. Chem. Soc., 88, 2520 (1966).

 

a. L. S. Bartell, J. Am. Chem. Soc., 88, 3567 (1961).
b. L. S. Bartell, Iowa State J. Sci., 88, 137 (1961).

 

 

a. L. Melander and R. E. Carter, J. Am. Chem. Soc., 88, 295
(1964).

b. L. Melander and R. E. Carter, Acta Chem. Scand., 18, 1138
(1964).

 

 

R. E. Carter and L. Dahlgren, Acta Chem. Scand., 88, 504 (1969).

 

K. Mislow, R. Graeve, A. J. Gordon, and G. H. Wahl, Jr.,
J. Am. Chem. Soc., 88, 1733 (1964).

 

G. J. Karabatsos, G. C. Sonnichsen, C. G. Papaioannou, S. E.
Schepple, and R. L. Shone, J. Am. Chem. Soc., 88, 463 (1967).

 

 

(33)

(34)

(35)

(36)

(37)
(38)

(39)

(40)

(41)

(42)
(43)

(44)

(45)

(46)

(47)
(48)
(49)
(50)

(51)

(52)

67

R. A. Scott and H. A. Scheraga, J. Chem. Phys., 88, 2209 (1965).

 

L. Hakka, A. Queen, and R. E. Robertson, J. Am. Chem. Soc.,
‘81, 161 (1965).

 

G. J. Frisone and E. R. Thornton, J. Am. Chem. Soc., 88, 1211
(1968).

 

K. T. Leffek, R. E. Robertson, and S. E. Suganori, Can. J. Chem.,
88, 1989 (1961).

 

M. Wolfsberg and M. J. Stern, Pure and App. Chem., 8, 325 (1964).

 

v. F. Smith, Ph.D. Thesis, Michigan State University, 1972. E

L. F. Fieser, "Experiments in Organic Chemistry," 3rd edition, 9
D. C. Heath and Co., Boston (1957), p. 285.

J. L. Dye and V. A. Nicely, J. Chem. Ed., 88, 443 (1971).

 

F. Daniels, J. W. Williams, P. Bender, R. A. Alberty, and t
C. D. Cornwell, "Experimental Physical Chemistry," 6th edition,

McGraw—Hill Book Co., Inc., New York (1962), p. 402.

C. G. Papaioannou, Ph.D. Thesis, Michigan State University, 1967.

S. E. Scheppele, Ph.D. Thesis, Michigan State University, 1964.

F. Arndt, "Organic Syntheses," Coll. Vol. 2, A. H. Blatt, ed.,
John Wiley and Sons, Inc., New York (1946), p. 461.

 

R. Shone, Ph.D. Thesis, Michigan State University, 1965.

R. Adams and E. W. Adams, "Organic Syntheses," Coll. Vol. 1,
A. H. Blatt, ed., John Wiley and Sons, Inc., New York (1946),
pp. 459, 462.

 

M. L. Bender, J. Am. Chem. Soc., 18, 1626 (1951).

 

M. Polanyi and A. L. Szabo, Tran. Far. Soc., 88, 508 (1934).

 

R. Warder, Chem. Ber., 85, 1361 (1881).

 

A. Moffat and H. Hunt, J. Am. Chem. Soc., 88, 2082 (1959).

 

S. L. Johnson, "Advances in Physical Organic Chemistry,"
V. Gold, ed., Academic Press, Inc., London (1967), p. 237.

a. E. 8. Could, "Mechanism and Structure in Organic Chemistry,"
Holt, Rinehart, and Winston, New York (1959), pp. 317-318.

b. S. Sarel, L. Tsai, and M. S. Newman, J. Am. Chem. Soc.,

18, 5420 (1956).

 

(53)

(54)

(55)

(56)

(57)

(58)

(59)

(60)

(61)

68

a. R. W. Taft, Jr., J. Am. Chem. Soc., 28, 3120 (1952).
b. C. K. Ingold and W. S. Nathan, J. Chem. Soc., 222 (1936).

 

 

J. F. Corbett, A. Feinstein, P. H. Gore, G. L. Reed, and E. C.
Vignes, J. Chem. Soc. (B), 974 (1969).

 

M. Adam-Briers. P. J. C. Fierens, and R. H. Martin, Helvitica
Chimica Acta, 88, 2021 (1955).

 

D. P. Evans, J. J. Gordon, and H. B. Watson, J. Chem. Soc.,
1430 (1937).

 

P. H. Gore, A. Rahim, and D. N. Waters, J. Chem. Soc. (B),
202 (1971).

 

J. B. Hendrickson, D. J. Cram, and G. S. Hammond, "Organic
Chemistry," 3rd edition, McGraw—Hill Book Co., New York
(1970), p. 304.

F. M. Menger and J. H. Smith, J. Am. Chem. Soc., 88, 3824
(1972).

 

W. B. Smith, D. L. Deavenport, and A. M. Ihrig, J. Am. Chem.
Soc., 88, 1959 (1972).

 

G. J. Karabatsos and C. G. Papaioannou, Tet. Let., 2629 (1968).

. w.’ . (Lb . ' ux'hfl—‘m—“Jlm tk—m I1
,.

V