Kinetics and mechanism of hydrolysis of N-(2′-hydroxyphenyl) phthalamic acid (1) and N-(2′-methoxyphenyl)phthalamic acid (2) in a highly alkaline medium

Article abstract of DOI:10.1002/kin.20361

A kinetic study on hydrolysis of N-(2′-hydroxyphenyl)phthalamic acid (1), N-(2′-methoxyphenyl)phthalamic acid (2), and N-(2′- methoxyphenyl)benzamide (3) under a highly alkaline medium gives second-order rate constants, k_OH, for the reactions o

Full text of DOI:10.1002/kin.20361

Kinetics and Mechanism
of Hydrolysis of N-(2ꢀ-
Hydroxyphenyl)phthalamic
Acid (1) and N-(2ꢀ-
Methoxyphenyl)phthalamic
Acid (2) in a Highly Alkaline
Medium
YOKE-LENG SIM, EMMY FADHIZA DAMIT, AZHAR ARIFFIN, M. NIYAZ KHAN
Department of Chemistry, Faculty of Science, University of Malaya, 50603 Kuala Lumpur, Malaysia.
Received 16 December 2007; revised 9 April 2008; accepted 4 June 2008
DOI 10.1002/kin.20361
Published online in Wiley InterScience (www.interscience.wiley.com).
ABSTRACT: A kinetic study on hydrolysis of N-(2ꢀ-hydroxyphenyl)phthalamic acid (1), N-(2ꢀ-
methoxyphenyl)phthalamic acid (2), and N-(2ꢀ-methoxyphenyl)benzamide (3) under a highly
alkaline medium gives second-order rate constants, kOH, for the reactions of HO− with 1, 2,
and 3 as (4.73 0.36) × 10−8 at 35◦C, (2.42 0.28) × 10−6 and (5.94 0.23) × 10−5 M−1 s−1 at
65◦C, respectively. Similar values of kOH for 3, N-methylbenzanilide, N-methylbenzamide, and
N,N-dimethylbenzamide despite the difference between pKa values of aniline and ammonia
of ∼10 pK units are qualitatively explained. ꢁ 2008 Wiley Periodicals, Inc. Int J Chem Kinet 41:
C
1–11, 2009
nately, the rate of uncatalyzed hydrolysis of an amide
under physiological conditions is extremely slow (half-
life period, t1/2, >220 years) [1], and kinetic study on
slow reactions (t1/2 greater than a few days) is not an
attractive field of kinetic measurements because of the
various unavoidable practical problems. The t1/2 value
for neutral hydrolysis of amide bond of >220 years
is based upon the estimated value of the pseudo-first-
order rate constant (k0 = 8.9 × 10−11 s−1 at 35◦C) for
neutral hydrolysis of formamide (k0 = 8.0 × 10−8 s−1
at 80◦C [1]). The reported value of k0 for neutral hy-
drolysis of a peptide bond (i.e., a secondary amide
INTRODUCTION
Kinetics and mechanism of uncatalyzed and catalyzed
aqueous cleavage of amide bond may be considered
of importance partly because aqueous degradation of a
protein involves the aqueous cleavage of peptide bond,
which is essentially a secondary amide bond. Unfortu-
Correspondence to: M. Niyaz Khan; e-mail: niyaz@um.edu.my.
Contract grant sponsor: National Scientific Research and Devel-
opment Council of Malaysia for ScienceFund.
Contract grant number: 14-02-03-4014.
2008 Wiley Periodicals, Inc.
c
ꢀ

2
SIM ET AL.
as described elsewhere [6]. N-(2ꢀ-Methoxyphenyl)
benzamide (3) was synthesized using a literature pro-
cedure involving the reaction of benzoyl chloride with
2-methoxyaniline, and the observed spectroscopic data
are in complete agreement with the corresponding re-
ported data [7]. All other common chemicals used were
commercial products of the highest available purity.
Standard solutions of 1 (0.03 M) and 3 (0.005 M) were
prepared in acetonitrile, whereas the standard solution
of 2 (0.01 M) was prepared in 1,4-dioxan.
bond) at room temperature is 3 × 10−9 s−1 [2]. But a
theoretical study [3] on the mechanism of formamide
hydrolysis in water reveals two reaction steps, where
the first step, the carbonyl group of the formamide
molecule being hydrated to form a diol intermediate,
is the rate-determining step followed by the final prod-
uct formation in the fast second step. The predicted
pseudo-first-order rate constant for the rate-limiting
step (the first step) of the hydrolysis reaction at 25◦C
(3.9 × 10−10
s
−1) is claimed in excellent agreement
with the experimental data (1.1 × 10−10
s
−1). Another
and more recent theoretical study [4] on the mecha-
nism of neutral hydrolysis of formamide predicts an
observable activation free energy barrier of 48.7 kcal
mol−1 and as a consequence of such a high free energy
barrier, the authors raise important questions about the
reliability of the experimental activation free energy of
31.0 kcal mol−1 and suggest that the neutral hydrolysis
of formamide does not take place at all.
Kinetic Measurements
The rate of alkaline hydrolysis of 1 was studied spec-
trophotometrically by monitoring the disappearance of
1 at 280 nm and 35◦C in aqueous solvent containing
1% v/v CH3CN. The details of the kinetic procedure
are described elsewhere [8]. The observed absorbance
(Aobs) at different reaction times (t) was found to fit to
Eq. (1)
Kinetic studies on extremely slow reactions are gen-
erally carried out at a high temperature, as well as
high [HO−] or [H+], and these experimentally deter-
mined rate constants are used to estimate rate constants
at 37◦C by using the Eyring or Arrhenius equation
coupled with reaction mechanisms for HO−- and H+-
assisted hydrolysis [1,5]. Recently, aqueous cleavage
of 1 and 2 has been studied under mild acidic medium
at 35◦C, where an efficient intramolecular carboxyl
group assisted rate enhancement is observed [6]. But a
plausible intramolecular general acid assistance due to
the 2ꢀ-OH group in 1 involving the transition state TS1
could not be detected [6].
Aobs = δapp[R0] exp(−kobst) + A
(1)
∞
where [R0] is the initial concentration of 1, δapp is an
apparent molar extinction coefficient of the reaction
mixture, A = Aobs at t = ∞ and kobs represents the
∞
pseudo-first-order rate constant for hydrolytic cleavage
of 1. The values of unknown parameters, kobs, δapp, and
A , were calculated from Eq. (1) using a nonlinear
∞
least-squares technique. The observed data fit to Eq. (1)
is satisfactory as evident from the standard deviations
associated with the calculated parameters as shown in
Table I, and from the plot of Aobs versus t for a typical
kinetic run as shown in Fig. 1, where a solid line is
drawn through the calculated data points using Eq. (1)
and parameters listed in Table I.
H
O
H
N
O
The rates of alkaline hydrolysis of 2, 3, and 2-
hydroxyaniline (4) were also studied by using the pro-
cedure of kinetic measurements as described above.
The plot of Aobs versus t for a typical kinetic run for
hydrolysis of 2 is also shown in Fig. 1.
O
H
TS1
The present study was initiated with an aim to dis-
cover the effect of [HO−] on the rate of hydrolysis of
1 and 2 at 35◦C and at an elevated temperature. To find
out the approximate effect of 2-CO−2 on kOH value for
2, the value of kOH for 3 was also determined. These
results as well as probable explanations are described
in this paper.
Product Characterization
The products in the respective alkaline hydrolysis
of 2 and 3 were confirmed by comparing the val-
ues of molar extinction coefficients, δ, obtained from
A
∞
values (calculated from Eq. (1)) with the corre-
sponding δ values of authentic samples of phthalic
acid or benzoic acid and 2-methoxyaniline obtained
under conditions of kinetic runs. These observations
showed that the products were 100% phthalic acid
and 2-methoxyaniline for 2 and benzoic acid and 2-
methoxyaniline for 3.
EXPERIMENTAL
Materials
N-(2ꢀ-Hydroxyphenyl)phthalamic acid (1) and N-(2ꢀ-
methoxyphenyl)phthalamic acid (2) were synthesized
International Journal of Chemical Kinetics DOI 10.1002/kin

KINETICS AND MECHANISM OF HYDROLYSIS OF PHTHALAMIC ACIDS IN AN ALKALINE MEDIUM
3
Table I Kinetic Parameters, kobs, δapp, and A , Calculated from Eq. (1) for Alkaline Hydrolysis of 1a
∞
−6 c
[NaOH] (M) 102 A∞est b 108 kobs (s−1
)
10−1
δ
app
(M−1 cm−1
)
102A
10
t
(s) Adobs
108kcealcd (s−1
)
∞
fin
1.2
8.6
4.74 0.15f
591 7f
8.8 2.0f
8.7 2.0
21.5 1.9
21.4 1.9
8.5 1.2
21.2 1.2
8.7 1.1
21.4 1.0
8.4 1.5
21.0 1.7
9.25
1.280
4.40
4.77
8.6g
4.74 0.15
5.21 0.17
5.19 0.16
5.95 0.10
6.22 0.12
6.27 0.10
6.89 0.10
7.76 0.15
8.55 0.20
591
7
7
7
4
4
4
4
6
6
21.3
549
549
597
555
590
548
605
564
21.3g
8.6
1.5
1.8
2.0
9.25
9.25
9.25
1.105
1.095
0.960
5.70
6.19
7.01
7.61
7.88
8.56
21.3
8.6
21.3
8.6
21.3
a [10] = 3.0 × 10−4 M, λ = 320 nm, ionic strength 2.0 M (by NaCl), T = 35◦C, aqueous reaction mixture for each kinetic run contained 1%
v/v CH3CN.
b Unless otherwise noted, t = 3.0 × 108 s for Ae∞st values.
c Time for the final observed absorbance (Aobs).
d Observed absorbance at tfin
.
e Calculated from Eq. (4) with k0 = 0 and kOH values shown in Table V.
f
Error limits are standard deviations.
g t = 1.5 × 108 s for Ae∞st values.
RESULTS
out within the [NaOH] range 1.2–2.0 M for the reac-
tion period of 2570 h, which is equivalent to only 0.6
and 1.0 half-lives at 1.2 and 2.0 M NaOH, respec-
tively. The nature of Eq. (1) is such that the use of the
nonlinear least-squares technique cannot yield reliable
Alkaline Hydrolysis of
N-(2ꢀ-Hydroxyphenyl)phthalamic Acid (1)
Because of the extremely slow rate of reaction of 1
with HO− at 35◦C, only a few kinetic runs were carried
values of kinetic parameters (kobs, δapp, and A ) if the
∞
2.00
2.00
1.80
1.60
1.50
1.40
1.20
1.00
0.80
0.60
0.40
0.20
0.00
1.00
0.50
0
2
4
6
8
10
0
20
40
60
80
100
120
140
160
180
200
-
10 6 t / s
Figure 1 The plots showing the absorbance (Abs) versus time (t) for alkaline hydrolysis of 1 and 2 at [NaOH] = 1.80 M,
where [10] = 3.0 × 10−4 M ( ), [2 ] = 2.0 × 10−4 M (ꢁ) and 35◦C. The solid lines are drawn through the calculated data
ꢀ
0
points using Eq. (1) as described in the text. Inset shows the observed and calculated data points until t = 1.0 × 107 s.
International Journal of Chemical Kinetics DOI 10.1002/kin

4
SIM ET AL.
2.50
0.50
0.40
0.30
0.20
0.10
0.00
2.00
1.50
1.00
0.50
0.00
0.00
0.05
0.10
0.15
0.20
0.25
0
10
20
30
40
50
60
70
80
90
-
10 4 t / s
Figure 2 The plot showing the dependence of absorbance (Abs) versus time (t) for alkaline degradation of 2-hydroxyaniline
(4) in the presence of phthalic acid, where [40] = [Phthalic acid] = 3.0 × 10−4 M, [NaOH] = 1.51 M, and 35◦C. The solid
lines are drawn through the calculated data points using Eqs. (2) and (1) for initial and final phase of degradation as described
in the text.
kinetic run was monitored for the reaction period of ≤1
half-life [9]. Thus, to get the reliable values of kinetic
parameters, the value of Aobs at t > 10 half-lives should
be known [9]. The expected hydrolysis products of 1
are 2-hydroxyaniline and phthalic acid. The authentic
samples of 2-hydroxyaniline and phthalic acid were
used to obtain δ for these compounds under the reac-
tion conditions of the kinetic runs which in turn gave
Ae∞st = 0.085 (where A∞est = Aobs at t ≥ 10 half-lives).
The use of Aobs = 0.085 at t ≥ 10 half-lives resulted in
monitored at 320 nm, are shown graphically in Fig. 2.
The plot of Fig. 2 reveals the presence of a maximum
which, in turn, implies the formation of, at least, one
moderately reactive intermediate in the alkaline degra-
dation of 2-hydroxyaniline. It is well known that 2- and
4-hydroxyanilines undergo a rather complex oxidative
degradation process under an alkaline medium [10].
The attempt was therefore made to determine only the
values of rate constants for the alkaline degradation
processes of 2-hydroxyaniline under the strictly reac-
tion conditions of the kinetics of alkaline hydrolysis
of 1 so that one could find the answer to the question:
Why did the observed data (Aobs versus t) for alkaline
hydrolysis of 1 obey strictly the first-order kinetics de-
spite the fact that the degradation of 2-hydroxyaniline
(an expected immediate stable hydrolysis product of 1)
followed complex two irreversible consecutive reac-
tion paths?
the values of kobs, δapp, and A as shown in Table I.
∞
It is also evident from Table I that the change in t (for
Ae∞st values) from 3.0 × 108 s (i.e., ∼22 half-lives)
to 1.5 × 108 s (i.e., ∼11 half-lives) for the reaction at
1.2 M NaOH did not essentially change the values of
kobs, δapp, and A .
∞
In the process of determining δ values of the mix-
ture of authentic 2-hydroxyaniline and phthalic acid at
320 nm in an alkaline medium, it was noticed that the
absorbance of the aqueous alkaline mixture changed
with time, indicating the instability of the authentic
sample mixture under the alkaline pH. To explore some
details of this observation, such as possible pseudo-
first-order rate constant(s) for the alkaline degradation
of 2-hydroxyaniline in the presence and absence of ph-
thalic acid, a few kinetic runs were carried out at 35◦C
and varying concentrations of NaOH. The values of
Aobs versus reaction time (t) for a typical kinetic run,
Aqueous Alkaline Degradation of
2-Hydroxyaniline
A few kinetic runs were carried out for the aqueous
degradation of 2-hydroxyaniline (in the absence and
presence of 3.0 × 10−4 M phthalic acid) within the
[NaOH] range 1.2–2.0 M at 35◦C and 2.0 M ionic
strength. The plot of Aobs versus t revealed, at least, one
maximum (Fig. 2). However, the rate of initial phase
International Journal of Chemical Kinetics DOI 10.1002/kin

KINETICS AND MECHANISM OF HYDROLYSIS OF PHTHALAMIC ACIDS IN AN ALKALINE MEDIUM
5
0.80
2.50
2.00
1.50
1.00
0.50
0.00
0.60
0.40
0.20
0.00
0.20
0
1
2
3
4
5
-
-
1
1
3
5
7
9
11
13
-
0.50
-
10 3 t / s
Figure 3 The plots showing the absorbance (Abs) versus time (t) for alkaline degradation of 2-hydroxyaniline (4) at different
[NaOH] in the absence of phthalic acid, where [40] = 3.0 × 10−4 M, 35◦C and [NaOH] = 1.0 M ( ) and 2.0 M (ꢁ) respectively.
ꢀ
The solid lines are drawn through the calculated data points using Eq. (1) as described in the text. Inset shows the observed and
calculated data points until t = 5.0 × 103 s.
of degradation appeared to be much faster than that of
the final phase of the degradation and consequently the
observed data (Aobs versus t) for the initial and final
phases of degradation were treated with Eqs. (2) and
(1), respectively.
of alkaline hydrolysis of 1. Thus, although the com-
plexity of the kinetics of the alkaline degradation of
2-hydroxyaniline in the absence and presence of ph-
thalic acid is not fully resolved, it is almost certain that
the rate constant for the slowest step of the alkaline
degradation of 2-hydroxyaniline is more than 100-fold
larger than that for alkaline hydrolysis of 1 (Table I),
and consequently the instability of the product
(2-hydroxyaniline) is not expected to cause deviation
of the observed data from the first-order kinetic plot for
the alkaline hydrolysis of 1. However, a relatively more
reliable observed data treatment to Eq. (1) requires the
use of the value of Ae∞st as 0.213 (absorbance due to final
alkaline degradation product(s) of 2-hydroxyaniline in
the presence of 3.0 × 10−4 M phthalic acid; Table II)
rather than 0.086 (absorbance due to 3.0 × 10−4 M 2-
hydroxyaniline and 3.0 × 10−4 M phthalic acid). The
observed data for alkaline hydrolysis of 1 were also
treated with Eq. (1) using Aobs (Ae∞st) = 0.213 at t =
3.0 × 108 s, and the calculated values of kobs, δapp, and
Aobs = δapp[R0][1 − exp(−kobst)] + A0
(2)
Although the nonlinear least-squares fit of the observed
data to Eq. (2) seems satisfactory in terms of standard
deviations associated with the calculated parameters
(kobs, δapp, and A0), a critical look at the plots of Figs. 2
and 3 shows small yet a regular deviations of observed
data points from the least-squares calculated solid
lines. Similar distinct deviations for the initial phase of
the reaction were observed for all kinetic runs within
the [NaOH] range 1.0–2.0 M. Thus, it seems that the
initial phase of the reaction does not represent a strictly
one-step reaction. The observed data for the final phase
of degradation showed a perfect fit to Eq. (1) (Fig. 2
and Table II), and these data reveal nearly 6-fold larger
rate constant for the initial phase than that for the final
phase degradation of 2-hydroxyaniline in the absence
and presence of 3.0 × 10−4 M phthalic acid.
A
∞
are summarized in Table I.
Effects of [NaOH] on the Rate
of Alkaline Hydrolysis of
The aim of the kinetic study for the degradation
of 2-hydroxyaniline in the absence and presence of
phthalic acid is to discover the possible effect(s) of the
product formation on the pseudo-first-order kinetics
N-(2ꢀ-Methoxyphenyl)phthalamic Acid (2)
To find the relative effects of 2-OH and 2-OMe on the
rates of hydrolysis of 1 and 2 under highly alkaline
International Journal of Chemical Kinetics DOI 10.1002/kin

6
SIM ET AL.
Table II Kinetic Parameters, kobs, δapp, and A0, Calculated from Eq. (2) for Initial Phase and Eq. (1) for Final Phase of
Alkaline Degradation of 2-Hydroxyanilinea
[NaOH] Reaction
Initial Phase
10−2
Final Phase
[NaOH] 106 kobs
δ
10−4 tfibn
106 kobs
(s−1
10−2
δ
10−4 ticni 10−4 tfibn
(s) (s)
app
app
(M)
(s−1
45.5 3.5d
)
(M−1 cm−1
)
102 A0
(s)
)
(M−1 cm−1
)
102 A
∞
1.0
110 3d 0.39 3.07d 8.06
(42.8 2.3)e (113 2)e (3.9 2.4)e (8.00)e (7.72 0.04d)e (197 1d)e (23.6 0.2d)e (12.0) (120)d
52.7 4.2 107 8.06
(42.2 2.3) (122 3) (4.7 2.0) (4.91)
64.9 6.3 101 8.06
−1.1 4.6
(51.7 3.9) (118 4) (2.0 3.3) (4.90)
(63.7 1.7) (73.9 0.7) (6.8 0.9) (10.9)
64.7 7.2 104 7.46
(48.6 4.0) (123 4)
1.2
1.5
3
−1.0 3.6
(8.76 0.05)
(210 1)
(218 3)
(22.7 0.3)
(18.4 0.4)
(12.0) (120)
3
(9.95 0.11)
(10.4 0.3)
(12.0) (120)
(17.2) (82.0)
(221 12) (21.3 0.9)
1.8
2.0
4
−2.9 5.5
(4.6 2.8) (4.10)
7.47
(10.8 0.1)
(220 2)
(16.2 0.2)
(11.2) (119)
70.0 8.5
(57.9 4.1) (115 3)
99.3 4.3 −2.8 6.0
(3.6 2.8) (4.10)
(11.4 0.1)
(221 3)
(16.0 0.3)
(11.2) (119)
a [2–Hydroxyaniline0] = 3.0 × 10−4 M, λ = 320 nm, ionic strength 2.0 M (by NaCl), T = 35◦C, aqueous reaction mixture for each kinetic
run contained 1% v/v CH3CN.
b Time for the final observed absorbance of the data treatment.
c Time for the first absorbance of the data treatment.
d Error limits are standard deviations.
e Reaction mixture contained 3.0 × 10−4 M phthalic acid.
medium, a few kinetic runs were carried out for the
hydrolysis of 2 under the experimental conditions al-
most similar to that for 1 and also at 65◦C. The values
and it is known∗ that CH3CN molecules hydrolyze to
produce acetic acid and ammonia as final hydrolysis
products. The rate of alkaline hydrolysis of CH3CN
is shown to follow a consecutive reaction scheme as
shown by Eq. (3) [11], where the k1 and k2 steps in-
volve HO− as a catalyst and a reactant (nucleophile),
respectively.
of kobs, δapp, and A , calculated from Eq. (1) within
∞
the [NaOH] range 1.0–2.0 M, are summarized in Ta-
ble III. The kinetic runs were carried out until 2.5–3.6
half-lives of the reactions at 35◦C, and the final prod-
ucts, 2-methoxyaniline and phthalic acid, were found
to be stable under such conditions.
k1[HO−
]
CH3CN + H2O −−−−−−→− CH3CONH2
k2[HO−
]
−−−−−−→− CH3CO−2 + NH3
(3)
Effects of [NaOH] on the Rate
of Alkaline Hydrolysis of
This implies that the k1 step does not consume HO−,
whereas the k2 step consumes HO−. The values of k2/k1
= 19 [11] and k2 = 3.8 × 10−5 M−1 s−1 (25◦C), 8.6
× 10−5 M−1 s−1 (35◦C), 3.1 × 10−4 M−1 s−1 (50◦C)
[12], and 3.7 × 10−5 M−1 s−1 (25◦C) [11]. In view of
these reported data, the value of k1 is 95-fold and 0.8-
N-(2ꢀ-Methoxyphenyl)benzamide (3)
The rate of alkaline hydrolysis of 3 was studied within
the [NaOH] range 0.5–2.0 M at 35 and 65◦C. The
observed data(Aobs versus t) showedareasonablygood
fit to Eq. (1), and the least-squares calculated values of
fold larger than the second-order rate constant (kOH
)
kobs, δapp, and A are shown in Table IV. An attempt
∞
for hydroxide ion-catalyzed hydrolysis of 1 and 3, re-
spectively, provided k2/k1 remained unchanged with
the increase of the temperature from 25 to 35◦C. Thus,
the hydrolysis products of 1% v/v CH3CN are expected
to consume ∼0.19 M NaOH and as a consequence this
[HO−] (=0.19 M) correction was considered in the ob-
served data analysis for alkaline hydrolysis of 1 and 3.
to obtain kobs at [NaOH] < 0.5 M was unsuccessful
because the reaction rates became too slow to monitor
conveniently.
DISCUSSION
The rates of alkaline hydrolysis of 1 and 3 were stud-
ied in an aqueous solvent containing 1% v/v CH3CN,
∗We thank one of the reviewers for informing us about this.
International Journal of Chemical Kinetics DOI 10.1002/kin

KINETICS AND MECHANISM OF HYDROLYSIS OF PHTHALAMIC ACIDS IN AN ALKALINE MEDIUM
7
Table III Kinetic Parameters, kobs, δapp, and A , Calculated from Eq. (1) for Alkaline Hydrolysis of 2a
∞
b
fin
108kcdalcd
(s−1
108kcealcd
(s−1
)
−5
[NaOH]
(M)
Temperature
(◦C)
108kobs
(s−1
10−1
(M−1 cm−1
δ
10
t
app
)
)
102A
(s)
102Acobs
)
∞
1.0
1.2
1.5
1.8
2.0
0.5
1.0
1.2
1.5
1.8
2.0
35
27.0 0.6f
29.7 0.8
27.9 1.6
36.8 0.7
38.3 1.2
93.7 3.5
246 12
294 22
378 12
453 18
535 18
290 3f
72.5 0.7f
68.7 0.8
56.6 2.0
59.4 0.6
55.4 1.0
65.0 0.7
65.0 1.2
65.4 1.8
65.3 0.7
66.3 0.9
65.5 0.7
64.9
64.9
65.6
64.9
64.9
82.6
77.6
69.0
66.2
61.3
98.8
81.4
79.5
74.5
73.5
69.6
26.2
28.5
31.9
35.4
37.7
95.7
238
295
381
466
523
22.0
26.4
33.0
39.6
44.0
121
242
290
362
435
483
303
4
339 10
348
361
269
268
268
262
264
265
3
5
4
6
9
4
5
4
65g
4.83
4.83
4.82
4.80
4.79
4.78
a [20] = 2.0 × 10−4 M, λ = 280 nm, ionic strength 2.0 M (by NaCl), aqueous reaction mixture for each kinetic run contained 2% v/v
1,4-dioxan.
b Reaction time for final observed data point.
c Value of Aobs at t = tfin
.
d Calculated from Eq. (4) with k0 and kOH values shown in Table V.
e Calculated from Eq. (4) with k0 = 0 and kOH value summarized in Table V.
f
Error limits are standard deviations.
g Data treatment for all kinetic runs was carried out by including a data point at t ≥10 half-lives with Aobs = Ae∞st (0.650).
Pseudo-first-order rate constants (kobs) for alkaline
hydrolysis of 1 showed an apparent good fit to Eq. (4)
where k0 and kOH represent the pseudo-first-order rate
constant for the reaction of 1 with H2O and the second-
order rate constant for the reaction of 1 with HO−,
respectively. The least-squares calculated values of k0
kobs = k0 + kOH[HO−]
(4)
Table IV Kinetic Parameters, kobs, δapp and A , Calculated from Eq. (1) for Alkaline Hydrolysis of 3a
∞
[NaOH]
(M)
Temperature
(◦C)
107kobs
(s−1
10−1
(M−1 cm−1
δ
107kcdalcd
(s−1
app
b
fin
102Acobs
)
−4
)
)
102A
10
t
/(s)
∞
0.5
1.0
1.2
1.5
1.8
2.0
0.5
35
19.2 1.0 e
50.0 3.4
63.1 2.2
79.7 5.2
89.6 6.1
531 10 e
584 20
590 10
612 18
626 18
613 11
14.7 0.5 e
11.5 0.9
12.0 0.5
11.9 0.9
11.5 0.9
12.2 0.6
14.7 0.4
15.3 0.6
11.7 0.1
15.7 0.7
12.1 0.2
14.0 0.4
12.5 0.3
16.5 0.3
12.1 0.3
16.4 0.6
12.7 0.4
15.7 0.4
19.2
19.1
18.6
18.6
18.6
18.6
12.9
12.8
12.8
12.6
10.8
12.7
7.51
12.6
3.91
7.00
3.89
7.00
33.1
20.8
20.1
16.5
15.4
14.9
16.3
17.1
11.8
16.0
12.3
13.7
12.8
16.5
12.8
16.5
13.1
16.2
20.9
49.3
60.7
77.7
94.8
106
177
188
476
500
595
625
774
812
954
999
1070
1120
108
200
5
8
65
487
501 11
566
530 17
559
567 10
8
202 13
456
1.0
1.2
1.5
1.8
2.0
7
3
491 37
571 12
626 25
795 21
759 24
928 23
1070 64
1100 40
1100 50
5
565
526
584
6
8
6
558 14
584
552 10
7
a [30] = 5.0 × 10−5 M, λ = 260 nm, ionic strength 2.0 M (by NaCl), aqueous reaction mixture for each kinetic run contained 1% v/v CH3CN.
b Reaction time for final observed data point.
c Value of Aobs at t = tfin
.
d Calculated from Eq. (4) with k0 and kOH as shown in Table V.
e Error limits are standard deviations.
International Journal of Chemical Kinetics DOI 10.1002/kin

8
SIM ET AL.
Table V Values of k0 and kOH Calculated from Eq. (3)
Amides
Temperature (◦C)
108k0 (s−1
)
108kOH (M−1 s−1
)
1
35
35
35
35
35
35
65
65
35
35
65
65
65
65
1.2 1.1a,b
3.86 0.76a,b
4.73 0.36
3.40 0.72c
4.35 0.35
11.4 3.2
0
1.3 1.1c
0
2
3
14.8 4.9
0
22.0 3.7
−47 11
285
8
0
242 28
568 27
604 25
(59.7 2.3) × 102
(59.4 2.3) × 102
(62.4 3.8) × 102
(62.2 3.2) × 102
33 33
0
(−81 292) × 102
0
(−0.5 4.7) × 102
0
a The values of k0 and kOH were calculated from kobs values obtained with Ae∞st = 0.213 (Table I).
b Error limits are standard deviations.
c The values of k0 and kOH were calculated from kobs values obtained with Ae∞st = 0.086 (Table I).
and kOH are summarized in Table V. The values of k0
with standard deviations of more than 100% are con-
sidered to be highly unreliable. The maximum contri-
bution of k0 compared to kOH [HO−] is <15% within
the experimental conditions of the study, and therefore
k0 may be neglected compared to kOH [HO−]. The val-
ues of kOH were also calculated from Eq. (4) with k0
= 0, and these calculated values, as shown in Table V,
increased by <12% compared to the corresponding
values obtained with k0 = 0. The extent of reliability
of the observed data fit to Eq. (4) is evident from kcalcd
values summarized in Table I.
The respective values of pKa1 and pKa2 of 1 are
3.34 and 9.52 at 35◦C [13]. The value of pKa3, which
corresponds to the ionization of amide group of 1,
should be >15 because the pKa of benzamide is 14–15
[14]. Thus, a plausible mechanism for hydrolysis of
1 under the present experimental conditions may be
expressed by Scheme 1, where pKa of conjugate acid
of the leaving group is larger in the k3 and k4 steps
absolute residual errors (ARE = (kobsi – kcalcdi)/kobsi
,
where kobsi and kcalcdi represent respective experimen-
tally determined and the least-squares calculated rate
constant at the ith value of [NaOH]; Table III). Al-
though the minimum contribution of k0 compared with
kOH[HO−] is ∼40% at 35◦C, the calculated value of
k0 (14.8 × 10−8
s
−1) is unexpectedly large for the fol-
lowing reasons: The reported respective values of k0,
kOH, and the hydronium ion-catalyzed second-order
rate constant (kH) for hydrolysis of formamide are
8.4 × 10−8
s
−1, 0.211 and 1.78 × 10−2 M−1 s−1 at
80◦C, respectively [1], and the values of kOH for hy-
drolysis ofbenzamide, N-methylbenzamide, andN,N-
dimethylbenzamide are 14.2 × 10−4, 6.40 × 10−4, and
10.3 × 10−4 M−1
s
−1, respectively, at 100.4◦C [15].
Similarly, the values of kH for hydrolysis of re-
spective benzamide, N-methylbenzamide, and N,N-
dimethylbenzamide are 39 × 10−5, 6.4 × 10−5, and
10 × 10−5 M−1 s−1 at 100.4◦C [16]. These reported
results demonstrate that the intrinsic reactivity toward
hydrolysis is significantly larger for formamide than
that for benzamide and N-substituted benzamide, and
consequently the value of k0 for benzamide and 2
should be significantly smaller than that for formamide
(k0 = 8.4 × 10−8 s−1 at 80◦C). Thus, the apparent
satisfactory observed data fit to Eq. (4) in terms of
ARE at 35◦C is considered to be unreliable because
the calculated value of k0 is unexpectedly large. Simi-
larly, the least-squares calculated negative value of k0
than that in the k and k steps, respectively. These
−1
−2
−1
conclusions reveal that k ꢃ k3 and k ꢃ k4 and
−2
consequently k3 and k4 steps are the rate–determining
steps. The observed rate law, rate = kobs [1]T, and
Scheme 1 can lead to Eq. (5), which is similar to Eq. (4)
with k0 = k3K1 [H2O] and kOH = k4K2. In Eq. (5),
K1 = k1/k−1, and K2 = k2/k
.
−2
kobs = k3K1[H2O] + k4K2[HO−]
(5)
(−4.7 × 10−7
s
−1) at 65◦C is meaningless and conse-
quently k0 may be considered as negligible compared
with kOH[HO−] in Eq. (4). Thus, a reliable value of
kOH was obtained from Eq. (4) with k0 = 0, and such
The values of kobs, obtained within the [NaOH] range
1.0–2.0 M at 35 and 65◦C, for hydrolytic cleavage of
2 show apparently satisfactory fit to Eq. (4) in terms of
International Journal of Chemical Kinetics DOI 10.1002/kin

KINETICS AND MECHANISM OF HYDROLYSIS OF PHTHALAMIC ACIDS IN AN ALKALINE MEDIUM
9
O
-
-
X = O , Y = CO2 for 1
H
N
-
X = OMe, Y = CO2 for 2
X = OMe, Y = H for 3
Y
X
SH
-
k2 [HO ]
k
k1 [H2O]
k
-2
-1
OH
O
NH2
H
N
O
H
N
X
k4
O
OH
+
OH
X
X
Y
Y
Y
-
-
T2
T
-
very fast
H2O
HO
NH2
O
-
X
k3
OH
+
Y
Scheme 1
a calculated value of kOH is shown in Table V. There-
fore, the observed data were treated with Eq. (4) and
the least-squares calculated values of kOH with k0 =
0 are summarized in Table V. The plausible reaction
mechanism for hydrolysis of 2 is shown in Scheme 1,
where k3 and k4 steps are rate determining.
The rate constants kobs, obtained for hydrolysis of 3
within the [NaOH] range 0.5–2.0 M, showed a good fit
to Eq. (4). However, the least-square calculated values
of k0 and kOH revealed a negative values of k0 (Table V)
at both 35 and 65◦C, which are meaningless. Hence,
the values of kOH were also calculated from Eq. (4)
with k0 = 0 and such values are also shown in Table V.
The negative values of k0 show that the contribution
of k0 is nearly insignificant compared to kOH[HO−] in
Eq. (4) within the [NaOH] range 0.5–2.0 M.
down of this reactive tetrahedral intermediate to form
the hydrolysis products is a fast step. This finding is
apparently contrary to what one believes in terms of
relative leaving abilities of the leaving groups in the
partitioning of the reactive tetrahedral intermediate. In
view of these theoretical studies, the rate law for al-
kaline hydrolysis of amides or, at least, formamide,
N-methylacetamide, DMF and DMA should be as
rate = kOH[HO−][amide]. But the alkaline hydroly-
sis of several tertiary amides [22–26] and closely re-
lated compound [27] revealed the rate law as rate =
(kOH[HO−] + kOꢀ H[HO−]2)[amide], which cannot be
explained convincingly in terms of these theoretical
studies [17–21]. Even the practical application of the
concept of dianionic tetrahedral intermediate, derived
from the kinetic term kOꢀ H[HO−]2 in the rate law, has
been used in the synthesis of amines [28].
The value of kOH for 1 is nearly 5-fold smaller
than that for 2, which may be attributed to unfavor-
able electrostatic effect of ionized phenolic group of 1.
Nearly 9-fold smaller value of kOH for the reaction of
HO− with anionic N-(2ꢀ-hydroxyphenyl)phthalimide
than that with N-(2ꢀ-methoxyphenyl)phthalimide was
attributed to similar electrostatic effect [29]. The value
The reaction mechanism for alkaline hydrolysis of 3
may be considered similar to the one shown in Scheme
1 with the replacement of 2ꢀ-CO2− by 2-H in SH. The
leaving groups (HO−) in k and k steps are much
−1
−2
better than the leaving groups in k3 and k4 steps in terms
of pKa of the conjugate acids of leaving groups. Conse-
quently, k−1 ꢃ k3 and k−2 ꢃ k4 and these inequalities
lead to the k3 and k4 steps as the rate determining.
Perhaps it is interesting to note that almost
all theoretical studies on the mechanism of aque-
ous alkaline hydrolysis of formamide [17–21], N-
methylacetamide, N,N-dimethylformamide (DMF),
and N,N-dimethylacetamide(DMA) [21] revealedthat
hydroxide ion nucleophilic attack at the carbonyl car-
bon in the formation of reactive tetrahedral interme-
diate is the rate-determining step, whereas the break-
of kOH (6.1 × 10−5 M−1
s
−1) for 3 may be compared
with kOH value (11 × 10−5 M−1
s
−1) [5] for the re-
action of HO− with N-methylbenzanilide obtained at
65◦C. The values of kOH for 2 and 3 reveal nearly
20-fold decrease in kOH caused by the presence of a
2-CO−2 group of 2. The replacement of 2-H by 2-CO2−
decreased the value of kOH by nearly 18-fold for the
alkaline hydrolysis of methyl benzoate [30].
International Journal of Chemical Kinetics DOI 10.1002/kin

10
SIM ET AL.
H
H
O
O
O
H
N
N
Z
Z
k
KT
O
+
ZNH2
O
OH
Y
Y
Y
-
-
Z = 2
= 2
= 2
-
-
-
O C6H4, Y = CO2 for 1
-
OMeC6H4, Y = CO2 for 2
OMeC6H4, Y = H for 3
Scheme 2
It is perhaps noteworthy that the values of kOH
for 3 and N-methylbenzanilide are similar to kOH for
N-methylbenzamide and N,N-dimethylbenzamide at
100.4◦C [16], despite the fact that pKa of ammonia is
larger than pKa of aniline by 10 pK units [31]. A qual-
itative explanation of these results may be described
as follows: To avoid the formation of highly unstable
anion, 2-XC6H5NH−, in the k4 step (Scheme 1) under
aqueous medium, the k4 step is proposed to involve ad-
ditional steps as shown in Scheme 2, where the k step
is the rate determining and pKa of the ZNH+2 group is
significantly larger for Z = R than that for Ar [32].
It is evident from Schemes 1 and 2 that k4 = kKT.
It can be easily realized from Schemes 1 and 2 that
(a) k2 (for Z = Ar, aryl group) > k2 (for Z = R, alkyl
group), (b) k (for Z = Ar) ≤ k (for Z = R), (c)
6. Sim, Y.-L.; Ariffin, A.; Khan, M. N. Int J Chem Kinet
2006, 38, 746.
7. Hibbert, F.; Mills, J. F.; Nyburg, S. C.; Parkins, A. W. J
Chem Soc, Perkin Trans 2 1998, 629.
8. Cheong, M. Y.; Ariffin, A.; Khan, M. N. J Phys Chem B
2007, 111, 12185.
9. Khan, M. N. Micellar Catalysis, Surfactant Science Se-
ries; CRC Press, Taylor & Francis Group: Boca Raton,
FL, 2006; Vol 133, Chap. 7.
10. Brown, K. C.; Corbett, J. F. J Chem Soc, Perkin Trans 2
1979, 308.
11. Widequist, S. Arkiv Kemi 1956, 10, 265.
12. Elsemongy, M. M.; Abu Elamayem, M. S.; Moussa,
M. N. H.; Gouda, M. M. J Indian Chem Soc 1975, LII,
1130.
13. Sim, Y.-L.; W. Ahmad, W. H.; Ariffin, A.; Khan, M. N.
Indian J Chem A 2008, 47, 240.
14. Barlin, G. B.; Perrin, D. D. Quart Rev Chem Soc 1966,
20, 75.
15. Bunton, C. A.; Noyak, B.; O’Connor, C. J. J Org Chem
1968, 33, 572.
16. Bunton, C. A.; Farber, S. J.; Milbank, A. J. G.; O’Connor,
C. J.; Turney, T. A. J Chem Soc, Perkin Trans 2 1972,
1869.
17. Bakowies, D.; Kollman, P. A. J Am Chem Soc 1999,
121, 5712.
18. Pliego, J. R. Chem Phys 2004, 306, 273.
19. Zahn, D. Chem Phys Lett 2004, 383,
134.
−2
−2
k (for Z = Ar) > k (for Z = R), and (d) KT (for Z =
Ar) < KT (for Z = R). The inequalities (a)–(d) show
that K2KTk (for Z = Ar) may not be expected to be
significantly different from K2KTk (for Z = R). The
values of kOH for 2 and 3 are ∼10- to 11-fold larger
at 65◦C than those at 35◦C (Table V). This shows
that the rate is increased by ∼2.2-fold for each in-
crease of 10 K. Application of this empirical rule gives
∼20-fold rate change with the change in temperature
from 65 to 101◦C. However, electrophilic sites of 3
and N-methylbenzanilide are sterically more hindered
compared to that of N,N-dimethylbenzamide and this
could reduce kOH values for amides with relatively
more sterically hindered electrophilic sites.
20. Xiong, Y.; Zhan, C.-G. Huazhong Shifan Daxue Xuebao
Zirankexueban 2004, 38, 344.
21. Xiong, Y.; Zhan, C.-G. J Phys Chem A 2006, 110,
12644.
22. Biechler, S. S.; Taft, R. W., Jr. J Am Chem Soc 1957,
79, 4927.
BIBLIOGRAPHY
23. Schowen, R. L.; Jayaraman, H.; Kershner, L. J Am Chem
Soc 1966, 88, 3373.
24. Menger, F. M.; Danohue, J. A. J Am Chem Soc 1973,
95, 432.
25. Young, J. K.; Pazhanisamy, S.; Schowen, R. L. J Org
Chem 1984, 49, 4148.
1. Hine, J.; King, R. S.-M.; Midden, R.; Sinha, A. J Org
Chem 1981, 46, 3186.
2. Kahne, D.; Still, W. C. J Am Chem Soc 1988, 110, 7529.
3. Gorb, L.; Asensio, A.; Tunon, I.; Ruiz-Lopez, M. F.
Chem—A Eur J 2005, 11, 6743.
4. Almerindo, G. I.; Pliego, J. R., Jr. J Braz Chem Soc
2007, 18, 696.
26. Sim, Y.-L.; Ariffin, A.; Khan, M. N. J Org Chem 2008,
73, 3730.
5. Broxton, T. J.; Duddy, N. W. Aust J Chem, 1979, 32,
1717.
27. Baxter, N. J.; Rigoreau, L. J. M.; Laws, A. P.; Page,
M. I. J Am Chem Soc 2000, 122, 3375.
International Journal of Chemical Kinetics DOI 10.1002/kin

KINETICS AND MECHANISM OF HYDROLYSIS OF PHTHALAMIC ACIDS IN AN ALKALINE MEDIUM
11
28. Gassman, P. G.; Hodgson, P. K. G.; Balchunis, R. J. J
Am Chem Soc 1976, 98, 1275.
29. Sim, Y.-L.; Ariffin, A.; Khan, M. N. J Org Chem 2007,
72, 2392.
30. Khan, M. N. Indian J Chem A 1986, 25, 831.
31. March, J. Advanced Organic Chemistry: Reactions,
Mechanisms and Structure, 2nd ed.; McGraw-Hill
Kogakusha; Tokyo; 1977; p 229 and references cited
therein.
32. (a) Fox, J. P.; Jencks, W. P. J Am Chem Soc 1974,
96, 1436 and references cited therein; (b) Page, M. I.;
Jencks, W. P. J Am Chem Soc 1972, 94, 8828;
(c) Satterthwait, A. G.; Jencks, W. P. J Am Chem Soc
1974, 96, 7031; (d) Cox, M. M.; Jencks, W. P. J Am
Chem Soc 1981, 103, 572; (e) Gresser, M. J.; Jencks,
W. P. J Am Chem Soc 1977, 99, 6970; (f ) Morris,
J. J.; Page, M. I. J Chem Soc, Perkin Trans 2 1980,
212.
International Journal of Chemical Kinetics DOI 10.1002/kin