Deprotonation sites of acetohydroxamic acid isomers. A theoretical and experimental study.
ABSTRACT Theoretical (ab initio calculations) and experimental (NMR, spectrophotometric, and potentiometric measurements) investigations of the isomers of acetohydroxamic acid (AHA) and their deprotonation processes have been performed. Calculations with the Gaussian 98 package, refined at the MP2(FC)/AUG-cc-pVDZ level considering the molecule isolated, indicate that the Z(cis) amide is the most stable form of the neutral molecule. This species and the less stable (Z)-imide form undergo deprotonation, giving rise to two stable anions. Upon deprotonation, the E(trans) forms give three stable anions. The ab initio calculations were performed in solution as well, regarding water as a continuous dielectric; on the basis of the relative energies of the most stable anion and neutral forms, calculated with MP2/PCM/AUG-cc-pVDZ, N-deprotonation of the amide (Z or E) structure appeared to be the most likely process in solution. NMR measurements provided evidence for the existence of (Z)- and (E)-isomers of both the neutral and anion forms in solution. Comparisons of the dynamic NMR and NOESY (one-dimensional) results obtained for the neutral species and their anions were consistent with N-deprotonation, which occurred preferentially to O-deprotonation. The (microscopic) acid dissociation constants of the two isomers determined at 25 degrees C from the pH dependence of the relevant chemical shifts, pK(E) = 9.01 and pK(Z) = 9.35, were consistent with the spectrophotometric and potentiometric evaluations (pK(HA) = 9.31).
Article: Reactivity of some products formed by the reaction of sorbic acid with sodium nitrite: decomposition of 1,4-dinitro-2-methylpyrrole and ethylnitrolic acid.[show abstract] [hide abstract]
ABSTRACT: Sorbic acid reacts with nitrite to yield mutagenic products such as 1,4-dinitro-2-methylpyrrole (NMP) and ethylnitrolic acid (ENA). In order to know the stability of these compounds, a kinetic study of their decomposition reactions was performed in the 6.0-9.5 pH range. The conclusions drawn are as follows: (i) The decomposition of NMP occurs through a nucleophilic attack by OH- ions, with the rate equation as follows: r = k(dec)NMP[OH-][NMP] with k(dec)NMP (37.5 degrees C) = 42 +/- 1 M(-1) s(-1). (ii) The rate law for the decomposition of ENA is as follows: r = k(dec)ENA[ENA]K(a)/(K(a) + [H+]), with K(a) being the ENA dissociation constant and k(dec)ENA (37.5 degrees C) = (7.11 +/- 0.04) x 10(-5) s(-1). (iii) The activation energies for NMP and ENA decomposition reactions are, respectively, E(a) = 94 +/- 3 and 94 +/- 1 kJ mol(-1). (iv) The observed values for the decomposition rate constants of NMP and ENA in the pH range of the stomach lining cells, into which these species can diffuse, are so slow that they could be the slow determining step of the alkylation mechanisms by some of the products resulting from NMP and ENA decomposition. Thus, the current kinetic results are consistent with the low mutagenicity of these species.Journal of Agricultural and Food Chemistry 12/2008; 56(24):11824-9. · 2.82 Impact Factor
[show abstract] [hide abstract]
ABSTRACT: A novel dry synthesis for the uranium(VI) dioxo-diacetohydroxamate (UAHA) complex has been developed. The complex was generated in > 80% yield by mechanically grinding solid uranyl acetate dihydrate (UAc) with solid ace-tohydroxamic acid in stoichiometric amounts. The resulting UO 2 (AHA) 2 ·4H 2 O solid is purified by washing with acetone. The stoichiometry was confirmed via colorimetric assays for U(VI) and AHA. The analogous ferric trisacetohydroxamate complex (FeAHA) was synthesized for comparison. The UAHA solid was extensively characterized by ultraviolet-visible (UV-vis), Fourier-transform infrared (FT-IR), and extended X-ray absorption fine structure (EXAFS) spectroscopies. The compound did not fluoresce after laser excitation. Proton nuclear magnetic resonance (NMR) spectra were obtained of the complex in D 2 O, acidified acetonitrile-d 3 , and DMSO-d 6 . The solubility was determined over a range of solvents. It was determined that in the purified solid, two bidentate AHA molecules bind to uranyl via the carbonyl and hydroxamate oxygen atoms, a structure analo-gous to known ferric, nickel, and lanthanum AHA complexes. In an acidic environment, binding is monodentate through the hydroxamate oxygen. And in aqueous solution, the UAHA complex assumes both binding moieties, depending on the pH. This pH-dependent speciation change is demonstrated for the first time. The easy synthesis and purification of UAHA enables researchers to strictly control reaction conditions; to eliminate interfering salts and water; and to study the complex in the solid-phase.Radiochim. Acta. 01/2007; 95958:439-450439.
Deprotonation Sites of Acetohydroxamic Acid Isomers.
A T heoretical and E xperimental Study
Marı ´ a L. Senent,†Alfonso Nin ˜o,‡Camelia Mun ˜oz Caro,‡Satunino Ibeas,§Begon ˜a Garcı ´ a,*,§
J ose ´M. Leal,§Fernando Secco,|and Marcella Venturini|
Departamento de Astrofı ´ sica Molecular e Infrarroja, Serrano 113b, Madrid 28006, Spain,
Grupo de Quı ´ mica Computacional, Escuela Superior de Informa ´tica, Universidad de Castilla la Mancha,
Paseo de la Universidad, 4, 13071 Ciudad Real, Spain, Departamento de Quı ´ mica, Universidad de
Burgos, Misael Ban ˜uelos s/ n, 09001 Burgos, Spain, and Dipartimento di Chimica e Chimica Industriale,
Universita ` di Pisa, Via Risorgimento 35, 56126 Pisa, Italy
Received February 4, 2003
Theoretical (ab initiocalculations) and experimental (NMR, spectrophotometric, and potentiometric
measurements) investigations of theisomers of acetohydroxamic acid (AHA) and their deprotonation
processes have been performed. Calculations with the Gaussian 98 package, refined at the MP2-
(FC)/AUG-cc-pVDZ level considering the molecule isolated, indicate that the Z(cis) amide is the
most stable form of the neutral molecule. This species and the less stable (Z)-imide form undergo
deprotonation, giving rise to two stable anions. Upon deprotonation, the E(trans) forms give three
stable anions. The ab initio calculations were performed in solution as well, regarding water as a
continuous dielectric; on the basis of the relative energies of the most stable anion and neutral
forms, calculated with MP2/PCM/AUG-cc-pVDZ, N-deprotonation of the amide (Z or E) structure
appeared to be the most likely process in solution. NMR measurements provided evidence for the
existence of (Z)- and (E)-isomers of both the neutral and anion forms in solution. Comparisons of
the dynamic NMR and NOESY (one-dimensional) results obtained for the neutral species and their
anions were consistent with N-deprotonation, which occurred preferentially to O-deprotonation.
The (microscopic) acid dissociation constants of the two isomers determined at 25 °C from the pH
dependence of the relevant chemical shifts, pKE ) 9.01 and pKZ) 9.35, were consistent with the
spectrophotometric and potentiometric evaluations (pKHA) 9.31).
Hydroxamic acids are very useful reagents with inter-
esting biological activity and medical applications. Due
to their ability to form stable chelates with a variety of
metal ions, hydroxamic acids play a key roleas bioligands
in the microbial transport of iron (siderophores); the ease
in forming metal chelates facilitates the function of
enzymes in oxygen and electron transport and in other
life-sustaining processes.1Despite their interesting prop-
erties, hydroxamic acids have remained as a poorly
characterized class of organic compounds; assignment of
the correct structure has been a major difficulty because
these compounds may adopt several forms.2Knowledge
of the preferred ionization sites and the protonation-
deprotonation mechanism is essential tolearning therole
played by hydroxamic acids in biological and complex-
Previously, we investigated acetohydroxamic acid us-
ing experimental and theoretical tools.3,4At high acidity
levels, the neutral AHA molecule can accept one proton.
The stable cations and neutral AHA forms have been
studied with ab initiomethods of good quality;3,4calcula-
tions predict bonding of the proton tothe hydroxylamine
oxygen group, which produces a third stable cation in the
presence of solvent, the CO oxygen being the most basic
site. Calculations in the gas phase suggest different
amidic and imidic structures that become stabilized by
The most stable AHA forms in the solid state were
found to be those prone to intermolecular H-bonding;5
Z-E isomerism was observed for AHA in DMSO6and
for monoalkylhydroxamic acids in several solvents.7On
thebasis of theoretical calculations and comparativedata
†Departamento de Astrofı ´ sica Molecular e Infrarroja.
‡Universidad de Castilla la Mancha.
§Universidad de Burgos.
|Universita `di Pisa.
(1) Brown, D. A.; Chidambaram, M. V. In Metal Ions in Biological
Systems; Marcel Dekker: New York, 1982; Vol. 14.
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Ventura, O. N.; Rama, J . B.; Turi, L.; Dannenberg, J . J . J . Am. Chem.
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Caro, C. Chem-Eur. J . 2000, 6, 2644.
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10.1021/jo0341564 CCC: $25.00 © 2003 American Chemical Society
Published on Web 07/18/2003
J . Org. Chem. 2003, 68, 6535-65426535
obtained with amides,8it has been reported that the
Z(cis) conformation of monohydroxamate group (CON-
HOH) is preferentially stabilized by intramolecular H-
bonding in nonpolar solvents9or by H-bonding in water.10
NMR and UV studies performed on the structure and
acid-base behavior, respectively, of N-phenylbenzohy-
droxamic acid have shown that the Z(cis):E(trans) ratio
strongly depends on the solvent used.11A property that
has been paid some attention but is the origin for certain
controversy is the acid-base behavior of hydroxamic
acids. O- versus N-deprotonation has received different
interpretations. First, it was accepted that hydroxamic
acids are O-acids,6,12-14but extensive IR and UV mea-
surements in dioxaneand aqueous alcohol15indicatethat
hydroxamic acids are N-acids, a conclusion supported by
17O NMR and FT-IRC studies of benzohydroxamate ion
in MeOH.12,16Potentiometric measurements compatible
with O- and N-deprotonation of hydroxamic acids have
been reported2,17-19assuming that the acid dissociation
constant Kexp ≈ KOMe + KNMe; however, irregularities
attributed to solvent effects have been observed.2,14b,19,20
The most probable deprotonation processes of AHA
were also investigated with ab initio methods. Previous
ab initio calculations of the isolated molecules confirm
the NH-acidity on the basis of the anions relative
energies.21-23Most of the ab initio studies on acid-base
properties of hydroxamic acids concern formoacetohy-
droxamic acid, which appears todissociatemost probably
by N-deprotonation.24-27Toget additional insight intothe
deprotonation sites of AHA, this work undertakes a
theoretical and experimental study of this acid at differ-
ent temperatures; the solvent effect was included con-
sidering water as a continuous dielectric. Relevant NMR
information on the deprotonation process and on confor-
mational isomerism is compared with thethermodynamic
results. The microscopic constants determined by NMR
are consistent with the macroscopic value obtained from
potentiometric and UV-vis measurements.
The ORIGIN 3800 IRIX 64 processor computer of the
CTI of C.S.I.C. (Madrid) was used for computation. All
ab initiocalculations were performed with the Gaussian
98 (Revision A.1x)28program at the MP2/AUG-cc-pVDZ
level. Original Fortran codes have been designed for the
treatment of the data. The zero-point vibrational energy
and thermodynamic properties of the species in the gas
phase have been calculated at the MP2/AUG-cc-pVDZ
level using the harmonic approximation and the rigid
rotor approximation for all the conformers.
In the calculations performed in solution, water was
regarded as a continuous dielectric, characterized by a
constant permittivity. The Polarized Continuum Model
(PCM) was used, implemented in the Gaussian 98
package29using the same basis at the MP2 level. The
cavity radii are those recommended for the UAHF
model.29The solvent dielectric constant has been set at
78.5 for T ) 298.15 K, 74.8 for T ) 308.15 K, and 71.5
for T ) 318.15 K.30The changes of free energy in solution
have been computed according tothe recipes given in ref
R esults and Discussion
T heoretical R esults: Neutral F orms. The starting
species for the quantum mechanical analysis were the
neutral stable structures already described (Figure 1),
which werefound with theMP2/cc-pVDZ method and full
geometry optimization.3The calculations performed with
the Gaussian 98 package were refined using the AUG-
cc-pVDZ basis set, which is more appropriate for the
treatment of anion species.28,32Searching for the stable
anions and simulation of the chemical dissociation pro-
cesses were performed starting from the reoptimized
(8) (a) Hupe, D. J .; Wu, D. J . Am. Chem. Soc 1977, 99, 7653. (b)
Yamaki, R. T.; Paniago, E. B.; Carvalho, S.; Howarth, O. W.; Kam, W.
J . Chem. Soc., Dalton Trans. 1977, 24, 4817. (c) Keeffe, J . R.; J encks,
W. P. J . Am. Chem. Soc. 1983, 105, 265.
(9) Liczynska-K ochany E.; Iwamura, H. J . Org. Chem. 1982, 47,
(10) Brown, D. A.; Coogan, R. A.; Fitzpatrick, N. J .; Glass, W. K.;
Abukshima, D. A.; Shiels, L.; Ahlgre ´n, M.; Smolander, K.; Pakkanen,
T. T.; Pera ¨kyla ¨, M. J . Chem. Soc., Perkin Trans. II 1996, 2673.
(11) Garcı ´ a, B.; Ibeas, S.; Mun ˜oz, A.; Leal, J . M.; Ghinami, C.; Secco,
F.; Venturini, M. Inorg. Chem. 2003, in press.
(12) Palm, V. A., Ed. Tables of Rate and Equilibrium Constants of
Heterocyclic Organic Reactions; VINITI: Moscow, 1975.
(13) Monzky, B.; Crumbliss, A. L. J . Org. Chem. 1980, 45, 4670.
(14) (a) Bagno, A.; Comuzzi, C.; Scorrano, G. J . Am. Chem. Soc.
1994, 116, 916. (b) Bagno, A.; Scorrrano, G. J . Phys. Chem. 1996, 100,
(15) Exner O.; Kaka ´c, B. Collect. Czech. Chem. Commun. 1963, 28,
(16) Exner, O.; Holubek, J . Collect. Czech. Chem. Commun. 1965,
(17) (a) Bordwell, F. G.; Fried, H. E.; Hughes, D. L.; Lynch, T. S.;
Satish, A. V.; Whang, Y. E. J . Org. Chem. 1990, 55, 3330. (b) Bordwell,
F. G.; Liu, W. Z. J . Am. Chem. Soc. 1996, 118, 8777.
(18) Decouzon M.; Exner, O.; Gal, J . F.; Maria P. C. J . Org. Chem.
1990, 55, 3980.
(19) Exner, O.; Hradil, M.; Mollin, J . Collect. Czech. Chem. Commun.
1993, 58, 1109.
(20) (a) Monzky, B.; Crumbliss, A. L. J . Org. Chem. 1980, 45, 4670.
(b) Brink, C. P.; Crumbliss, A. L. J . Org. Chem. 1982, 47, 1171. (c)
Brink, C. P.; Fish, L. L..; Crumbliss, A. L. J . Org. Chem. 1985, 50,
(21) Ventura, O. N.; Rama, J . B.; Turi, L.; Dannenberg, J . J . J . Am.
Chem. Soc. 1993, 115, 5754.
(22) Yamin, L. J .; Ponce, C. A.; Estrada, M. R.; Vert, F. Tomas J .
Mol. Struct. (THEOCHEM) 1996, 360, 109.
(23) Remko, M. J . Phys. Chem. A 2002, 106, 5005.
(24) Remko, M.; Mach, P.; Scheleyer, P. R.; Exner, O. J . Mol. Struct.
(THEOCHEM) 1993, 279, 139.
(25) Stinchcomb, D. M.; Pranata, J . J . Mol. Struct. (THEOCHEM)
1996, 370, 25.
(26) Remko, M. Phys. Chem. Chem. Phys. 2000, 2, 1113.
(27) Yen, S. J .; Lin, C. Y.; Ho, J . J . J . Phys. Chem. A 2000, 104,
(28) Frisch, M. J .; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.;
Robb, M. A.; Cheeseman, J . R.; Zakrzewski, V. G.; Montgomery, J . A.,
J r.; Stratmann, R. E.; Burant, J . C.; Dapprich, S.; Millam, J . M.;
Daniels, A. D.; Kudin, K. N.; Strain, M. C.; Farkas, O.; Tomasi, J .;
Barone, V.; Cossi, M.; Cammi, R.; Mennucci, B.; Pomelli, C.; Adamo,
C.; Clifford, S.; Ochterski, J .; Petersson, G. A.; Ayala, P. Y.; Cui, Q.;
Morokuma, K.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.;
Foresman, J . B.; Cioslowski, J .; Ortiz, J . V.; Stefanov, B. B.; Liu, G.;
Liashenko, A.; Piskorz, P.; Komaromi, I.; Gomperts, R.; Martin, R. L.;
Fox, D. J .; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.;
Gonzalez, C.; Challacombe, M.; Gill, P. M. W.; J ohnson, B. G.; Chen,
W.; Wong, M. W.; Andres, J . L.; Head-Gordon, M.; Replogle, E. S.;
Pople, J . A. Gaussian 98, revision A.1x; Gaussian, Inc.: Pittsburgh,
(29) Barone, M.; Cossi, B.; Mennucci, B.; Tomasi, J . J . Chem. Phys.
1997, 107, 3210.
(30) Riddik, J . A.; Bunger, W. B.; Sakano, T. K. Organic Solvents,
4th ed.; J ohn Wiley and Sons: New York, 1987; Vol. II, p 75.
(31) Tomasi J .; Persico M. Chem. Rev. 1994, 94, 2027.
(32) Kendall, R. A.; Dunning, T. H., J r.; Harrison, R. J . J . Chem.
Phys. 1992, 96, 6796.
Senent et al.
6536 J . Org. Chem., Vol. 68, No. 17, 2003
AHA presents different stable structures and shows
amide-imide tautomerism; conformational analysis pro-
vides stable acidic forms. The internal rotation of the
C-N bond that enables interconversion of the Z(cis) h
E(trans) forms is hindered by barriers higher than 25
kcal/mol;4hence, thetwoforms can betreated as different
species. The conformers generated by OH rotation and
CH3torsion can be classified as (Z)-amidic, (E)-amidic,
(Z)-imidic, and (E)-imidic forms. The(Z)- and (E)-isomers
are assumed to coexist, giving rise to different deproto-
nation processes leading toformation of independent (Z)-
and (E)-anions, respectively. Thus, the probability of the
process involving the (E)-form must be taken into con-
sideration for a proper interpretation of the mechanisms.
The minima for the MP2/cc-pVDZ and MP2/AUG-cc-
pVDZ potential energy surfaces agree well. As expected,
the diffuse functions modify slightly the relative energies
of the neutral species; the energy change ranged from
2.8% (in the most stable (E)-imide tautomer) to 15.3%
(in the most stable (Z)-imide tautomer). Table 1 lists the
relative energies of the most stable neutral forms with
respect tothe (Z)-amide tautomer (E ) -283.673447 au).
The amide forms display nonplanar geometry (Figure 1),
whereas the imide forms (not shown) adopt planar
structures with formation of a CdN double bond. The
intramolecular O3...H6-O5 H-bond and the steric hin-
drance of the methyl group cause the relative stabilities.4
Another significant feature is the difference in dipole
moment between the amide and imide forms, since the
polar solvent effect is more significant on the amide
structure; moreover, although the energy levels are
separated by relatively high barriers, the (Z)-amide, the
(E)-amide, and the (Z)-imide tautomers show similar
stability. The (E)-amide tautomer lies only 1.6 kcal/mol
abovethe(Z)-geometry; themost stable(Z)- and (E)-imide
tautomers (termed (Z)-II imide tautomer and (E)-III
imide tautomer) lie 1.1 and 5.4 kcal/mol above the (Z)-
amide form, respectively.
Totest the validity of the second-order Mo ¨ller-Plesset
approximation for determining relative energies, the
calculations of the isolated molecule have been repeated
with the MP4 method. With the exception of the (E)-
amide, whose MP4 relative energy drops from 1.6 to 0.8
kcal/mol, there are not significant changes with respect
to the data shown in Table 1.
Anion F orms. (Z)-AHA displays the three (Z)-Ia, (Z)-
Ib, and (Z)-II anions (Figure 2), the most stable of them
being (Z)-Ia. (E)-AHA gives rise to three less stable
anions. In the case of Ib-type anions, (E)-Ib represents a
conformer, whereas the geometry of (Z)-Ib is unstable,
although a rotamer of (Z)-Ib may exist. In these struc-
tures, the remaining H atom is bonded to O5, O3, and
N4, respectively; their MP2/AUG-cc-pVDZ relative ener-
gies are listed in Table 1. Extension of the basis set
causes an energy change of some 45%, this feature
indicating that calculations of the AHA anions require
the use of basis sets containing diffuse orbitals.
The (Z)-Ia and (Z)-II anions arise from geometry
optimization of the neutral (Z)-amide after elimination
of one proton, either H7+ or H6+ (see Figure 3, A ) (Z)-
amide; I ) (Z)-imide form II); in the amide tautomer, H6
is bound to O5, and H7 to N4. After elimination of H7+
(R1 process), optimization affords the most stable Ia
anion (E ) -283.121064 au). Loss of H6 (R2 process)
gives rise tothe (Z)-II anion (E ) -283.0936 au). On the
T ABL E 1.
AHA Considering the Molecule Isolated and in the Presence of Solventa
MP2/AUG-cc-pVDZ Calculated R elative E nergies, ER (kcal/mol), of the Stable Neutral and Anionic F orms of
aSolvation free energies,∆Gs, are given in kcal/mol. Dipole moments, µ, are given in Debyes.bEa) -283.673447 au.cEb) -283.688362
au.dEc) -283.121064 au.eEd) -283.221794 au.
F IGUR E 1. Neutral (Z)- and (E)-amidic forms of AHA.
F IGUR E 2. (Z)- and (E)-anions of AHA.
Deprotonation Sites of Acetohydroxamic Acid Isomers
J . Org. Chem, Vol. 68, No. 17, 2003 6537
other hand, H7 is bound toO3 in the (Z)-imide tautomer;
in this case, elimination of H6 (R3 process) produces the
(Z)-Ia anion, and loss of H7+ gives oneof the(Z)-Ia anion
rotamers. The conformers generated by the OH and CH3
internal rotations produce rotamers of the (Z)-Ia, (Z)-Ib,
and (Z)-II anions. Loss of H atoms from the neutral (E)-
forms gives risetothreedifferent anions (Figure4). After
elimination of H6+ or H7+ (RE2 and RE1 processes), the
(E)-amide tautomer can produce the most stable (E)-II
anion (E ) -283.101532 au) and (E)-Ia anion (E )
-283.100375 au). The(E)-imideform can producethe(E)-
Ia and (E)-Ib anions (E ) -283.075842 au) after elimina-
tion of H7+ or H6+ (RE3 or RE4 processes).
Theproduct of theR3 process is the(Z)-Ia anion rather
than the (Z)-Ib form, since this was found tobe unstable
at the MP2 level. Figure 2 shows two(Z)-structures, (Z)-
Ia and (Z)-Ib anions, connected by the interconversions
O3‚‚‚H6-O5 and O3-H6‚‚‚O5 (see Figure 1). The stable
(Z)-Ia anion has a planar structure, with the methyl
group antieclipsed with respect toO3, and H6 connected
toO5. The H-bond closes a cycle with a 2.5448 Å distance
between the two oxygens. The O3‚‚‚H6 equilibrium
intramolecular H-bond distance was evaluated to be
1.8170 Å and the MP2/AUG-cc-pVDZ dipole moment to
be 3.2959 D. However, (E)-Ia anion exhibits minimum
energy conformations. Some rotamers of the (Z)-Ib anion
shown in Figure 2 display some stability, although this
anion, denoted as 4 A2 by Yazal and Pang,33is reported
to be unstable with the Density Functional Theory.
There is no experimental evidence of the (Z)-Ib anion.
The (Z)-Ia anion T (Z)-Ib anion interconversion was
simulated searching for the minimum energy path.
Figure 5 plots the change in energy as a function of R (R
) O3H6); theremaining coordinates werefully optimized
for each selected structure. The (Z)-Ia anion (R ) 1.8170
Å) represents the unique minimum of the MP2/AUG-cc-
pVDZ path, although thegradient decreases in theregion
where the RHF curve (dashed line) shows a second
minimum (R ) 0.97 Å). RHF calculations led to two
equilibrium structures separated by a transition state
(RTS) 1.07 Å). The stability of the (Z)-Ib anion decreases
according to a correlation already observed.2
The two isomers of anion II, (E)- and Z, lie 12.3 and
17.2 kcal/mol above the (Z)-Ia anion (Table 1). The
interconversion (Z)-II anion T (E)-II anion requires that
a potential barrier of 50.3 kcal/mol be overcome. In the
(Z)-II anion, themethyl group is antieclipsed with respect
toO3. TheO3C2N4O5 frameis distorted tominimizethe
steric interactions of the two oxygens (O3O5 ) 2.9295
Å). For instance, the O5N4C2 angle changed from 107.6°
in the (Z)-Ia anion to 128.2° in the (Z)-II anion. In the
(E)-II anion, the methyl group eclipses O3; the two
oxygens are disengaged, but the charge distribution
blocks theintramolecular H-bond stabilization. TheGibbs
energy change of the interconversion (Z)-II anion T (E)-
II anion was evaluated tobe -5.7 kcal/mol, a value that
also decreases with temperature.
Solvent E ffect. In ab initiocalculations performed in
solution, water was regarded as a continuous dielectric,
characterized by a constant permittivity. The Polarizing
Continuum Model (PCM) was used, implemented in the
Gaussian 98 package.29Table1 lists therelativeenergies
(ER) of the most stable anions and neutral forms calcu-
lated with MP2/PCM/AUG-cc-pVDZ; also listed are the
solvation energies (∆GS(MP2)), defined as the difference
between the MP2 energies of the isolated species and the
corresponding values in solution and the solvation ener-
(33) El Yazal, J .; Pang, Y. P. J . Chem. Phys. 1999, 103, 8346.
F IGUR E 3. Deprotonation process of (Z)-AHA.
F IGUR E 4. Deprotonation of (E)-AHA.
F IGUR E 5.
interconversion as a function of the O3H6 bond distance.
Energy change of the O3H6‚‚‚O5O3‚‚‚H6O5
Senent et al.
6538 J . Org. Chem., Vol. 68, No. 17, 2003
gies calculated at the RHF level (∆Gs(RHF)). The two
parameters ∆GS(MP2) and ∆Gs(RHF) involve three non-
electrostatic contributions (cavitation, dispersion, and
repulsion energies). As expected, the solvation effect on
the neutral forms is negligible and somewhat larger for
amides compared to imides. This effect, however, is
important for the anions. The most stable (Z)-form is the
(Z)-Ia anion (E ) -283.221794 au), though the relative
energy of the (Z)-II anion drops from 17.2 to 6.5 kcal/
mol. For the (E)-forms, the solvent reverts the relative
order of stability of (E)-II and (E)-Ia anions and favors
(E)-N-deprotonation. The solvation free energies calcu-
lated at the Hartree-Fock level are shown in the last
column of Table 1.
The presence of solvent increases all dipole moments.
The dipole moments of the (Z)-Ia anion and the (Z)-II
anion produce strong N- and O-deprotonation, respec-
tively. A simple Mulliken analysis shows that the prox-
imity of the two oxygens in the (Z)-II anion raises both
the negative charge around the O5 and the electric
moment. As expected, thelargest solvation effect appears
in the polar species (Z)-II, which becomes stabilized by
74.0 kcal/mol upon solvent addition. Solvation augments
the Z-O-deprotonation probability.
Deprotonation Processes. The deprotonation site of
AHA still is a matter of controversy, and generally a
pronounced influence of the molecule environment is
accepted.34,35For an isolated molecule (ideal gas model),
the variation of the thermodynamic properties corre-
sponding tothe R1, R2, and R3 processes of the (Z)-forms
(Figure 3) and the four RE1, RE2, RE3, and RE4
processes (Figure 4) at 298.15, 308.15, and 318.15 K are
listed in Table 2 along with the properties calculated in
the presence of solvent. For an isolated molecule, we
consider the process AH h A-+ H+and treat the system
as an ideal gas. The entropy values of the neutral and
anionic species were taken from the Gaussian package
output. In the presence of water, the Gibbs energies can
be evaluated with the Born-Haber cycle:
where∆Gs(AH), ∆Gs(A-), and ∆Gs(H+) arethesolvation
Gibbs energies of the acid species, anion, and proton,
respectively (Table 1); ∆Gs(H+) was calculated as ∆Gs
(H3O+). Its value was calculated to be -107.5 kcal/mol
with MP2/AUG-cc-pVDZ, far from theexperimental value
(-264.61 kcal/mol) employed in ref 36; ∆Gisolated moleculewas
evaluated with the gas ideal model and from the De
dissociation energies (∆H ) De+ p∆V and ∆G ) ∆H -
T∆S). Thedetermination of Derequires calculation of the
total electronic energies and the zero-point energies
performed in a harmonic analysis (ZPVE ) 48.9 ((Z)-A),
49.0 ((E)-A), 48.9 ((Z)-I), 48.6 ((E)-I), 40.6 ((Z)-Ia), 40.0
((E)-Ia), 39.6 ((Z)-Ib), 40.7 ((Z)-II), and 40.6 ((E)-II) (kcal/
mol)). In the case of the solvent, we did not include the
solvent effect on thevibrational analysis. In addition, the
variation of the nonelectrostatic contributions to the
solvation energy with the temperature has been ne-
glected given the range of the temperature variation.
If the molecule is isolated, then the most likely process
is N-deprotonation of the (Z)-amide, which produces (Z)-
Ia anion (or formation of (Z)-Ia from the imide, in an
equilibrium mixture of amide and imide). In this case,
∆G298.15 has been calculated to be 332.0 kcal/mol, in
agreement with theexperimental data of ref 18 measured
in the gas phase (339.1 kcal/mol). This result supports
previous ab initiocalculations21,23and coincides with that
determined for formohydroxamic acid.2,24The next prob-
able processes are N and O deprotonation of the (E)-
amide; from thermodynamic considerations, (Z)-O-depro-
tonation appears tobeunlikely. Also, in solution themost
probable process is the formation of the (Z)-Ia anion from
the (Z)-amide and the (E)-Ia anion from the (E)-amide
(∆G298.15) 170.7 kcal/mol and ∆G298.15) 175.4 kcal/mol,
respectively). However, (Z)-N-deprotonation is favored.
The twoprocesses, (Z)-O-deprotonation (∆G298.15) 177.1
kcal/mol) and (E)-O-deprotonation (∆G298.15) 178.5 kcal/
mol), have a similar probability.
∆G298.15 obtained from the experimental pKHA value
(see below) was smaller than the theoretical value in
solution. The causes of the differences between experi-
mental and calculated pKHA values are discussed in
Spectrophotometric Measurements: Acid Dis-
sociation Constant. Within the ordinary pH range,
AHA behaves as a simple monoprotic acid whose dis-
sociation, represented by reaction 2,
(34) Eigen, M.; Kruse, W. Z. Naturforsch. B 1963, 18, 857.
(35) Diebler, H.; Secco, F.; Venturini, M. J . Phys. Chem. 1984, 88,
(36) Liptak, M. D.; Shields, G. S. J . Am. Chem. Soc. 2001, 123, 7314.
(37) Oliveira Silva, C.; Nascimento, M. A. C. Adv. Chem. Phys. 2002,
T ABL E 2.
T hermodynamic Properties (kcal/mol) Corresponding to the Most Probable deprotonation Processes of AHA
isolated molecule (AH h A-+ H+)PCM (AHsh A-s+ H+s)
R1: amide h (Z)-Ia anion
R2: amide h (Z)-II anion
R3: imide h (Z)-Ia anion
RE1: amide h (Z)-Ia anion
RE2: amide h (Z)-II anion
RE3: imide h (Z)-Ia anion
RE4: imide h (Z)-Ib anion
∆Gsolution) ∆Gisolated molecule- ∆Gs(AH) + ∆Gs(A-) +
HA h A-+ H+
Deprotonation Sites of Acetohydroxamic Acid Isomers
J . Org. Chem, Vol. 68, No. 17, 2003 6539