Mechanism of 1,4,5,8-naphthalene tetracarboxylic acid dianhydride hydrolysis and formation in aqueous solution

Article abstract of DOI:10.1039/b512187f

The study of highly conjugated, carbonyl-containing molecules such as 1,4,5,8-naphthalene tetracarboxylic dianhydride, III, is of interest since reactivity differences and transmission of electronic effects through the conjugated framework can be evidenced. The kinetics of hydrolysis of III in aqueous solution were determined from 5 M acid to pH 10. In basic solution hydrolysis of III yields, sequentially, 1,4,5,8-naphthalene diacid monoanhydride, II, and 1,4,5,8-naphthalene tetracarboxylic acid, I. The second order rate constant for alkaline hydrolysis is 200 fold higher for the first ring opening. The water-catalyzed hydrolysis of III yields a pH-dependent mixture of ionic forms of I and II. The rate constant for water-catalyzed hydrolysis of III is 25 fold higher than that for II. In concentrated acid the rates for reaching equilibrium (I, II and III) increase and III is the major product. The pK_as of I (3.24, 5.13 and 6.25) and II (3.05, 5.90) were determined by potentiometric, fluorescence and UV spectroscopy titrations and by quantitative fit of the kinetic and equilibrium data. The apparent, pH-dependent, equilibrium constants, K_EqII, for anhydride formation between I and II were obtained from the UV spectra. The quantitative fit of kinetic and equilibrium data are consistent with the assumption that anhydride formation only proceeds with the fully protonated species for both I and II and permitted the estimation of the equilibrium constants for anhydride formation, K_EqII. The value of K_EqII (I II) between pH 1 and 6 was ca. 5. Geometry optimization calculations in the gas phase of the reactions of III in alkaline, neutral and acid conditions, at the DFT level of theory, gave electronic distributions that were qualitatively consistent with the experimental results. The Royal Society of Chemistry 2006.

Full text of DOI:10.1039/b512187f

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www.rsc.org/obc | Organic & Biomolecular Chemistry
Mechanism of 1,4,5,8-naphthalene tetracarboxylic acid dianhydride
hydrolysis and formation in aqueous solution
T. C. Barros,a I. M. Cuccovia,*c J. P. S. Farah,b J. C. Masini,b H. Chaimovichc and M. J. Politic
Received 30th August 2005, Accepted 1st November 2005
First published as an Advance Article on the web 18th November 2005
DOI: 10.1039/b512187f
The study of highly conjugated, carbonyl-containing molecules such as 1,4,5,8-naphthalene
tetracarboxylic dianhydride, III, is of interest since reactivity differences and transmission of electronic
effects through the conjugated framework can be evidenced. The kinetics of hydrolysis of III in aqueous
solution were determined from 5 M acid to pH 10. In basic solution hydrolysis of III yields,
sequentially, 1,4,5,8-naphthalene diacid monoanhydride, II, and 1,4,5,8-naphthalene tetracarboxylic
acid, I. The second order rate constant for alkaline hydrolysis is 200 fold higher for the first ring
opening. The water-catalyzed hydrolysis of III yields a pH-dependent mixture of ionic forms of I and
II. The rate constant for water-catalyzed hydrolysis of III is 25 fold higher than that for II. In
concentrated acid the rates for reaching equilibrium (I, II and III) increase and III is the major product.
The pKas of I (3.24, 5.13 and 6.25) and II (3.05, 5.90) were determined by potentiometric, fluorescence
and UV spectroscopy titrations and by quantitative fit of the kinetic and equilibrium data. The
apparent, pH-dependent, equilibrium constants, KEqII, for anhydride formation between I and II were
obtained from the UV spectra. The quantitative fit of kinetic and equilibrium data are consistent with
the assumption that anhydride formation only proceeds with the fully protonated species for both I and
II and permitted the estimation of the equilibrium constants for anhydride formation, KEqII. The value
of KEqII (I ꢀ II) between pH 1 and 6 was ca. 5. Geometry optimization calculations in the gas phase of
the reactions of III in alkaline, neutral and acid conditions, at the DFT level of theory, gave electronic
distributions that were qualitatively consistent with the experimental results.
and 6, to 100 in concentrated acid. Ab initio calculations for
the pathway connecting the undissociated 1,8-Acid to the 1,8-
An are consistent with a mechanism involving a rate determining
intramolecular proton transfer concerted with oxygen alignment
towards the carbonyl center forming a transition state with a
neutral and planar intermediate with an sp3 carbon. The second
step is dehydration, through another transition state, yielding a
complex between water and 1,8-An.10
Tetra-substituted naphthalene acid, namely 1,4,5,8-naphtha-
lene tetracarboxylic acid, I, Scheme 1, and its mono II and
dianhydrides, III, are intermediates in the synthesis of polymers
and dyes.11
Fully protonated I yields 1,4,5,8-naphthalene tetracarboxylic
dianhydride, III, upon drying.11 Hydrolysis of III in aque-
ous base was reported to yield I without the formation of
1,4,5,8-naphthalene diacid monoanhydride, II, as an intermediate,
Scheme 1.11 Compound II was obtained by acidification of I
with mineral acid at temperatures above 60 ◦C. Kofman and
coworkers prepared I, and its derivatives, II and III, in acid at
90 ◦C and 140–150 ◦C, respectively.11 Acid dissociation constants
and alkaline hydrolysis in aqueous DMSO were also analyzed.11
Compounds I, II and III showed 4, 3 and 2 dissociable groups,
respectively, and the authors rationalized the data by assuming an
intermediate, formed after OH− attack at carbonyl groups of III,
leads to the appearance of a dissociable group and the opening
of the two anhydrides to yield I. The formation of II in aqueous
solution upon heating was suggested from NMR spectra and it
was proposed that, in moist DMSO, an equilibrium between I and
1
Introduction
The experimental and theoretical analysis of anhydride formation
in intramolecular reactions is a paradigmatic model chemical
system, extensively used to analyze ground and transition state
contributions to the rates of intramolecular and enzymatic
reactions.1
Hydrolysis of carboxylic anhydrides in aqueous solution can
be spontaneous, acid or base-catalyzed.1–6 Dicarboxylic acids can
be in equilibrium with the correspondent anhydrides in aqueous
solutions and, in general, the equilibrium is displaced towards
the diacid.7,8 The equilibrium constant may vary from 10−7 to 25
depending on structure and the unionized diacid is the kinetically
significant species for anhydride formation.7–9
The equilibrium between 1,8-naphthalic anhydride, 1,8-An,
and its hydrolysis product 1,8-naphthalene dicarboxylic acid, 1,8-
Acid, in aqueous solution has been analyzed previously.10 The
pH dependence of hydrolysis kinetics of 1,8-An is complex and
indicates an equilibrium reaction between 1,8-An and 1,8-Acid
at pHs lower than 5.0. Kinetics and equilibrium results were
rationalized by assuming that protonated forms of the 1,8-Acid
are in equilibrium with 1,8-An. The values of the equilibrium
constants for 1,8-An formation range from 4, between pH 0
aUniversidade Paulista, Instituto de Cieˆncias da Sau´de, Bauru, SP, Brazil
bDepartamento de Qu´ımica Fundamental and Instituto de Qu´ımica da,
Universidade de Sa˜o Paulo, SP, Brazil
cDepartamento de Bioqu´ımica do, Instituto de Qu´ımica da, Universidade de
Sa˜o Paulo, SP, Brazil. E-mail: imcuccov@quim.iq.usp.br
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Scheme 1 Reactions involved in the hydrolysis of III to II and I.
III is possible.12 Crystal structures of compounds I, II and III are
known.12–16
mono and dianhydrides in the acid, water and base-catalyzed
reactions.
1,4,5,8-Naphthalene tetracarboxylic acid, I, Scheme 1, is an
interesting model where short-range electronic and steric interac-
tions, as well as protonation states, can modulate chemical reactiv-
ity. Intramolecular reactions of I exhibit short range interactions
that are typical of those encountered in an enzyme active site.
The detailed pathways of anhydride formation and hydrolysis
in the 1,4,5,8-naphthalene tetracarboxylic system have not been
analyzed. It seemed unlikely that the hydrolysis of III yielded I
without the intervening formation of II, as described previously.11
Here we present kinetic and equilibrium data that contribute to
the understanding of the mechanism of formation and hydrolysis
of III in aqueous solutions. Cyclization of I to II and further
cyclization to III, in acid, were also studied. The pKs of I and
II were determined using UV-Vis and fluorescence spectroscopy,
potentiometric titration and kinetic data. Cyclization of both I and
II depends on the protonation of the carboxylic groups. Geometry
optimization of III, selected intermediates and monoanhydrides
was performed to define differences in reactivity between the
2
Experimental
2.1 Materials
1,4,5,8-Naphthalene tetracarboxylic dianhydride, III, (Aldrich)
was used without further purification. The tetra sodium salt of
1,4,5,8-naphthalene tetracarboxylic acid,17 I-Base, (Scheme 1) was
prepared by alkaline hydrolysis of III, adding III to yield a final
concentration of 0.001 M in aqueous solutions of NaOH and
maintaining the sample at 30 ◦C for 30 min (see results). At the end
of the reaction, absorbance measurement at 368 nm guaranteed
complete hydrolysis of III to I-Base (see results).
1,4,5,8-Naphthalene diacid monoanhydride, II: 12.5 mL of an
aqueous solution of I-Base (1 × 10−3 M) in NaOH (0.005 M)
was added to a 25 mL volumetric flask containing water and HCl
sufficient to give pH 1.0 and 5 × 10−4 M of protonated I-Base,
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I-H4. The solution was maintained at 30 ◦C until I-H4 was almost
completely cyclized (ca. 80%, see results) to II-H2.
Observed first-order rate constants, kw, calculated by the SV18
MV program, are the mean values of at least 8 runs.
Other reagents (Merck, Aldrich or Sigma) were P.A. or spec-
troscopic grade and used as supplied. All aqueous solutions and
buffers were prepared in freshly glass bi-distilled water. The buffers
used were the sodium salts of borate (pH 9.0 to 10.0), phosphate
(pH 6.4 to 8.0), acetate (pH 3.8 to 5.6), citrate (pH 2.2 to 6.0),
4-morpholineethanesulfonic acid, MES (pH 5.5 to 6.5), and [2-
(hydroxymethyl) aminomethane], TRIS (pH 7.5 to 8.5). Buffer
concentrations, shown in the figure legends, varied from 0.01 M
to 0.05 M. HCl or HClO4 were used for pHs lower than 2.0,
and in the H0 region.18 NaClO4 was obtained by neutralization of
concentrated HClO4 with sodium hydroxide.
Potentiometric titrations. I-Base was titrated with 0.1148 M
HCl, under N2, using 50 mL of a 0.001 M solution of I-Base
prepared in NaOH (pH 11.5). Titrant HCl was added to the
solution with a Hamilton micro-syringe in volumes from 25 to
500 lL. All potentiometric titrations of I-Base were completed in
ca. 15 min in order to minimize cyclization (see results). Data (pH
vs. acid volume) were linearized with modified Gran functions to
obtain pKas values.19
Spectrophotometric titrations. The pKII-1 and pKII-2 of II were
determined from the absorbance (365 nm) measured at several
pHs, by adding 50 lL of 1 × 10−3 M solution of II, prepared at
pH 1.0, to 2.0 mL of 0.05 M buffer at several pHs.
2.2 Methods
pKI-H1 and pKI-H2 of I were determined from measurement of
the absorbance (at 250 nm), at several pHs, by adding 40 lL of
1 × 10−3 M of I-Base, prepared at pH 11.5, to 2.0 mL of 0.05 M
buffer. The absorbances were measured within intervals of 30 s to
minimize cyclization.
All kinetic reactions, titrations and equilibrium measurements
were performed at 30 ◦C. I was also titrated by fluorescence
(emission 454 nm, excitation 290 nm).
UV-Vis absorption spectra and kinetic measurements were
recorded using Hitachi U-2000, Cary 3E, Shimadzu UV-2401 PC
or Beckman DU-7 spectrophotometers. Kinetic measurements
were done at 30 ◦C
0.1. Fast reactions, above pH 7, were
recorded in an Applied Photophysics Model SX-18 MV Stopped
Flow system. Steady-state fluorescence spectra were recorded
either in SPEX-1681 or Hitachi-F-2000 spectrofluorimeters. The
excitation and emission slit widths were set to 1 mm, band-pass
of 13 nm. Fluorescence emission spectra obtained with the SPEX
fluorometer were corrected using the fabricant software. pHs were
measured with a Beckman model U71 pH meter. I was titrated
using a Radiometer PHM 82 Standard pH Meter equipped with
a glass electrode calibrated with standard buffers at 30 ◦C.
Structural calculations. Geometry optimization was per-
formed by DFT level of theory using perturbed BP; DN**
(Spartan v5).20
3
Results and discussion
3.1 Spectra
Absorbance and fluorescence measurements. Absorption and
fluorescence measurements were made using 3 mL quartz cuvettes
with 10 mm optical path. Final concentrations of III or I, usually
1 × 10−5 M, were obtained by adding 1% (v/v) of the stock
solution into the appropriate solvent or buffer. Stock solutions
of III and I (1 × 10−3 M) were prepared in CH3CN or DMSO
and in aqueous NaOH respectively. The standard errors of the
extinction coefficients were less than 5%.
Typical UV-Vis absorption and fluorescence emission spectra
of III are presented in Fig. 1. The UV maximum absorbance
wavelengths, kmax, and corresponding molar absorptivities, e, of III
in carbon tetrachloride and acetonitrile are collected in Table 1.
The value of e of III in CH3CN (363.5 nm), is 30 618 M−1 cm−1,
a typical value of p, p* transitions, indicating a highly conjugated
aromatic system. Spectral resolution of III and high e values were
similar to that observed for 1,4,5,8-naphthalene derivatives.21
Emission and excitation spectra of III in CH3CN showed fine
vibrational structure and were symmetric, demonstrating that the
ground and excited states have the same geometry (Fig. 1A). Fluo-
rescence quantum yields, Uf, and maximum emission wavelength,
kemmax, of III in CH3CN and CCl4 are presented in Table 2. The low
Uf value of III in CCl4 can be attributed to a solvent quenching
effect.22
Kinetic measurements. The kinetics of hydrolysis and cycliza-
tion were followed at the maximum absorbance wavelength as
follows: 368 nm for III (e = 23 737 M−1cm−1) and 308 nm for I (e =
10 000 M−1cm−1). Reactions were initiated by adding 0.025 mL of
stock solutions of the substrates to 2.5 mL of buffer giving a final
concentration of 1.0 × 10−5 M. Rate constants were calculated
from first-order kinetics with at least 4–5 half-lives in all cases and
are the averages of at least three independent experiments differing
by no more than 5%.
The fluorescence spectra (emission and excitation) of III in a
mixture of 50% CH3CN–water changed with time (Fig. 1A). This
effect is assigned to a hydrolysis reaction which leads to opening of
the anhydride moiety and to a decrease in the aromatic conjugation
Fast kinetics. At each shot of the stopped flow system 0.1 mL
of III stock solution in CH3CN was mixed with 2.5 mL of buffer.
Table 1 Absorption maximum wavelengths (kmax) and molar absorptivity coefficients (e) of III in several solvents
Solvent
k1/nm
e1/M−1cm−1
k2/nm
e2/M−1 cm−1
k3/nm
e3/M−1cm−1
CH3CN
CCl4
HClO4 0.1 M
HClO4 0.1 M + NaClO4 4 M
HClO4 12 M
329
329
332
333
336
13 214
10 189
11 708
12 215
11 700
345.5
346
348
351
354
23 410
18 600
19 699
18 607
20 751
363.5
364
368
370
375
30 618
25 294
23 737
17 784
25 552
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(Fig. 1B) shows a rapid decrease, with half-life of ca. 24 s followed
by a much slower process with a half-life of ca. 2 h which leads to
further decrease in the absorbance at 368 nm and an increase in
absorbance at 308 nm. The initial fast decrease in the absorbance
at 368 nm is attributed to hydrolysis of III yielding II (Scheme 1).
Compound II at pH 7 is almost completely dissociated (see below)
and the fully dissociated ionic form of II will be referred as II-Base,
(Scheme 1, and see below). Note that, in this initial reaction, an
isosbestic point at 338 nm is evident (Fig. 1B). Above 260 nm the
UV spectrum of II-Base shows a single kmax at 348 nm (Fig. 1B,
line 7). The slower process can be attributed to base catalyzed
hydrolysis of II-Base yielding I (Fig. 1B, line 15), which at pH 7.0 is
also fully dissociated, I-Base, see below (Scheme 1). The spectrum
of I-Base exhibits a kmax at 308 nm. In the alkaline hydrolysis of
II → I an isosbestic point at 326 nm is clear.
In the following section a kinetic study of the reactions III →
II and II ꢀ I in a wide range of acidities are presented.
3.2 Kinetics
The first order rate constant, kw, for the reaction III → II was
calculated from the absorbance change followed at 368 nm (where
III absorbs) using only the fast change process. The variation of
the observed rate constants, kw, for the alkaline hydrolysis of III
yielding II-Base with pH is shown in Fig. 2A. The solid line in
Fig. 2A was obtained by fitting the data with eqn (1), where kw is
the first order rate constant for the water-catalyzed reaction and
kOH the second-order rate constant for hydroxide attack (III →
II-Base, Scheme 1). The best-fit parameters were: kw = 0.005 s−1
III
and kOH = 8.5 × 104 M−1s−1.
III
Fig. 1 (A) Fluorescence excitation (a, b, c) and emission (aꢀ, bꢀ, cꢀ) spectra
of III, 1 × 10−5 M, in acetonitrile–water (1:1 v/v) at different times, in
minutes: zero (a, aꢀ), 8 (b, bꢀ) and 45 (c, cꢀ). (B) UV-Vis spectra of III, 1 ×
10−5 M, in phosphate buffer, 0.02 M, pH 7, as a function of time: (1) t =
zero; (2) t = 6 s; (3) t = 12 s; (4) t = 18 s; (5) t = 30 s; (6) t = 54 s; (7) t =
5 min; (8) t = 9 min (9) t = 20 min; (10) t = 40 min; (11) t = 1 h; (12) t =
1.5 h; (13) t = 2 h; (14) t = 3 h; (15) t = 4 h.
kw = kw + kOH[OH]
(1)
The kw for the water-catalyzed hydrolysis and hydroxide attack
on II-Base yielding I-Base, were obtained starting the reaction by
addition of II-Base to solutions of different pHs (insert, Fig. 2A).
II
Using eqn (1) the best-fit parameters were kw = 2.0 × 10−4 s−1
and kOH = 4.5 × 102 M−1 s−1. The kwIII/kwII and kOHIII/kOHII ratios
II
(25 and 190, respectively) show the difference in reactivity between
the dianhydride and monoanhydride towards water and OH−.
Addition of III to buffers of pHs between 1 and 6, also led to
a time dependent changes in kmax from 368 nm to 348 nm and a
slower change from 348 to 308 nm (not shown). The fast process
can be attributed to hydrolysis of III to II as analyzed above. The
slower process, however, is more complex, since the UV spectra of
the products correspond to II ꢀ I equilibrium mixture and the
concentration of I at equilibrium was pH-dependent, decreasing
with acidity (see below).
Table 2 Fluorescence parameters for III in CH3CN and CCl4
Solvent
k1/nma
k2/nm
k3/nm
Ufb
CH3CN
CCl4
372
369
390
389
414
415
0.130
0.015
a k1, k2, k3 are the emission maximum wavelengths (excitation wavelength
k was 330 nm). b Uf = Quantum yield.
extent and consequently in a decrease in Uf. Using methanol or
ethanol as solvents, a time dependent shift was observed in III
absorption kmax from 368 nm to 346 nm, also indicative of solvolysis
(not shown).
Like in other solvents (Table 1) compound III, as its diimide
homologues,21 showed a well-resolved UV-Vis absorbance spec-
trum in aqueous solution (phosphate buffer 0.02 M, pH 7.0)
(Fig. 1B, line 1). The UV-Vis spectrum of III in water exhibits
a kmax at 368 nm and a second peak at 348 nm, similar to that
obtained for films of III.23 These features show that, as dianhydride
or diimide, the chromophoric group remains relatively unsolvated.
Adding III to aqueous buffer (pH 7.0) resulted in a time-
dependent spectral change (Fig. 1B). The absorbance at 368 nm
The time-dependent spectrum of a reaction mixture in 0.1 M
HClO4, with III as the initial reagent, is shown in Fig. 2B. Two
consecutive isosbestic points can be differentiated: one at 341 nm
during the fast first absorbance decrease and another at 320 nm
in the slower phase. The existence of two consecutive isosbestic
points lends credence to our assumption that the observed kinetics
can be best explained by two consecutive reactions also at this
pH, i.e. III → II and II ꢀ I. The peak at kmax of III disappears
after ca. 10 min (Fig. 2B) therefore, at equilibrium the spectrum
corresponds to that of II, with a small contribution of I. It is
also clear from the spectra that in the 1–6 pH range the final
concentration of III is negligible.
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