Mechanistic variations in the oxidation of Piloty's acid by metal complexes

Article abstract of DOI:10.1021/ic049627z

The reactions of N-hydroxybenzenesulfonamide (Piloty's acid, PA) with a variety of metal oxidants are reported. Either nitric oxide or nitrite is the final reaction product, along with benzenesulfinate and the reduced metal compound. The nitrogen product depends on the oxidation potential of the metal oxidant and its ability to further oxidize NO to nitrite. The observation and preliminary interpretation of unusual kinetic behavior of Piloty's acid as a reductant is also described. Analogues of PA were also prepared and found to show similar reactivity.

Full text of DOI:10.1021/ic049627z

Inorg. Chem. 2004, 43, 7877−7881
Mechanistic Variations in the Oxidation of Piloty’s Acid by Metal
Complexes†
Patricia C. Wilkins, Hollie K. Jacobs, Michael D. Johnson,* and Aravamudan S. Gopalan*
Department of Chemistry and Biochemistry, MSC 3C, New Mexico State UniVersity, Box 30001,
Las Cruces, New Mexico 88003-8001
Received March 19, 2004
The reactions of N-hydroxybenzenesulfonamide (Piloty’s acid, PA) with a variety of metal oxidants are reported.
Either nitric oxide or nitrite is the final reaction product, along with benzenesulfinate and the reduced metal compound.
The nitrogen product depends on the oxidation potential of the metal oxidant and its ability to further oxidize NO
to nitrite. The observation and preliminary interpretation of unusual kinetic behavior of Piloty’s acid as a reductant
is also described. Analogues of PA were also prepared and found to show similar reactivity.
3
Introduction
potential delivery agents for NO and HNO. It is surprising,
however, that the oxidation kinetics of PA with various
Although Piloty’s acid (PA), N-hydroxybenzenesulfon-
amide, was first prepared in 1896, it was not until the 1970s
that N-hydroxysulfonamides were observed to be inhibitors
of Zn metalloenzymes such as carbonic anhydrase and matrix
5
1
oxidants has received only limited attention.
The decomposition of PA in aqueous solution has been
shown to produce benzenesulfinate, HNO/ NO , and finally
-
4
a
2
nitrous oxide according to the reactions
metalloproteinases. Piloty’s acid and some of its acylated
derivatives have also been shown to be good inhibitors of
aldehyde dehydrogenase and potent vasodilators. Recently,
-
-
C H SO NHOH + OH h C H SO NHO + H O
3
6
5
2
6
5
2
2
PA and its derivatives have been shown to decompose to
produce HNO in aqueous solution or release NO upon
oxidation.4 In view of the biomedical relevance of these
reactions, PA and its derivatives have been explored as
-
-
C H SO NHO f HNO + C H SO
6
5
2
6
5
2
a,5
2
HNO f N O + H O
2
2
These decomposition studies were carried out at high pH
*
To whom correspondence should be addressed. E-mail: agopalan@
nmsu.edu (A.S.G.), johnson@nmsu.edu (M.D.J.). Fax: 505 646 2394.
Dedicated to the memory of Michael Blake Ryan, April 18, 1971, to
September 8, 2002.
(9-13), and from the inverse proton dependence of the
reaction along with N NMR data, it was concluded that
†
15
-
the deprotonated form, C
6
H
6
SO
2
NHO , was the reactive
(
1) Piloty, O. Ber. Dtsch. Ges. 1896, 29, 1559.
(
2) (a) Blackburn, M.; Mann, B. E.; Taylor, B. F.; Worrall, A. F. Eur. J.
Biochem. 1985, 153, 553. (b) Scozzafava, A.; Supuran, C. T. J. Med.
Chem. 2000, 43, 3677. (c) Mincione, F.; Menabuoni, L.; Briganti, F.;
Mincione, G.; Scozzafava, A.; Supuran, C. T. J. Enzyme Inhib. 1998,
4
species. However, PA is relatively stable to decomposition
to HNO at physiological pH under anaerobic conditions (at
4
b
pH 7, t1/2 ) 92 h).
13, 267. (d) Duffel, M. W.; Ing, I. S.; Segarra, T. M.; Dixon, J. A.;
Recently, it was shown that, under both anaerobic and
Barfknecht, C. F.; Schoenwald, R. D. J. Med. Chem. 1986, 29, 1488.
3) King, S. B.; Nagasawa, H. T. Chemical approaches toward generation
of nitroxyl (HNO). In Methods in Enzymology; Packer, L., Ed.;
Academic Press: New York, 1998; pp 211-220. Granik, V. G.;
Grigor’ev, N. B. Russ. Chem. Bull. Int. Ed. 2002, 51, 1375. Nagasawa,
H. T.; Kawle, S. P.; Elberling, J. A.; DeMaster, E. G.; Fukuto, J. M.
J. Med. Chem. 1995, 38, 1865. Fukuto, J. M.; Hszieh, R.; Gulati, P.;
Chiang, K. T.; Nagasawa, H. T. Biochem. Biophys. Res. Commun.
(
aerobic conditions at pH 7-9, oxidative decomposition of
3-
PA by O
2
, H
2
O
2
, or [Fe(CN)
6
]
takes place with the
5
production of nitric oxide. The production of NO in the
oxidation of PA was also reported in studies to determine
whether NO or HNO (NO ) was responsible for the
-
5
1992, 187, 1367. Lee, M. J. C.; Shoeman, D. W.; Goon, D. J. W.;
vasorelaxant properties of the sulfonamide.
Nagasawa, H. T. Nitric Oxide: Biol. Chem. 2001, 5, 278.
4) (a) Bonner, F. T.; Ko, Y. Inorg. Chem. 1992, 31, 2514. (b) Hughes,
M. N.; Cammack, R. Methods Enzymol. 1999, 301, 279.
5) (a) Grzesiok, A.; Weber, H.; Zamora Pino, R.; Feelisch, M. In The
Biology of Nitric Oxide; Moncada, S., Feelisch, M., Busse, R., Higgs,
E. A., Eds.; Portland Press; London, 1994; pp 238-241. (b) Zamora,
R.; Grzesiok, A.; Weber, H.; Feelisch, M. Biochem. J. 1995, 312, 333.
As part of an ongoing study into selective chelation of
actinides and transition metals, the interaction of aqueous
iron(III) with N-hydroxysulfonamides, which are structurally
similar to hydroxamic acids, was investigated. No complex
formation between PA and iron(III) was evidenced by
(
(
1
0.1021/ic049627z CCC: $27.50
© 2004 American Chemical Society
Inorganic Chemistry, Vol. 43, No. 24, 2004 7877
Published on Web 11/04/2004

Wilkins et al.
Table 1. Rate Constants for PA Oxidations at 25 °C
k
kH
kO
E°
kexc
oxidanta
pHb
7.0
(M-1 s-1
)
(M-1 s-1
)
(M-1 s-1
)
(V vs NHE)
(M-1 s-1
)
c
5
0.78d
8 e
[
Fe(bipy)2(CN)2]+
610
195
11
1.2 × 10
∼10
6
6
.5
.0
75
-
3
3
0.957d
0.882
5 d
[
[
[
[
[
IrCl6]2
7.1
7.1
1.2 × 10
2.3 × 10
2
-
d
IrBr6]
Fe(CN)6]
Fe(CN)5N3]
1.6 × 10
3
-
f
3 e
0.361
5 × 10
3
-
0.33e
5 e
10
4.0
3
+
3.0i
7.5
f
g
Fe(H2O)6]
145
0.771
+
-4
0.06h
Mb
∼1.4 × 10
a
Anaerobic conditions. Buffer ) 0.1M sodium phosphate. c bipy ) 2,2′-bipyridine. d Reference 11. e Reference 12. f Reference 13. g Reference 14.
b
h
i
-3
Reference 15. 10 M HCl, I ≈ 0.01 M.
changes in the UV-vis spectrum of a mixture of these two
reactants around pH 3. However, a redox reaction was
observed to form iron(II) and nitric oxide. Because of the
novel nature of this reaction and the potential therapeutic
uses of synthetic NO and HNO/NO donors, a systematic
study of the oxidation of PA and some of its derivatives was
undertaken.
Results and Discussion
The oxidation of PA by aqueous iron(III) in acid produced
iron(II) and NO; the latter was identified by scavenging with
+
+
metmyoglobin, Mb . Addition of Mb initially gave the
-
6
+
metmyoglobin-NO (Mb NO) adduct that underwent a slow
base-catalyzed reductive nitrosylation to form the deoxy-
myoglobin-NO adduct (Mb°NO). Both adducts were identi-
fied by their visible absorption spectra. The scavenging of
NO occurred much faster than the reduction of iron(III) by
PA.
8
Experimental Section
Oxidants were either purchased from Aldrich (K
2
IrCl
(CN)
5 3
Fe(CN) N ] ) by literature methods. Myoglobin (equine) was
6
, K
2 6
IrBr )
or Fluka (K Fe(CN) ) or prepared ([Fe(bipy)
3
6
2
2
]+ and
The iron(III)/PA reaction at I ≈ 0.01 M is relatively rapid
3-
7
[
(
Table 1), and a plot of kobs vs the PA concentration is linear
purchased from Sigma. Piloty’s acid was purchased from Acros
and recrystallized from EtOH/water before use. The other N-
hydroxysulfonamides were synthesized by the method of Blackburn.2
All other chemicals were of the highest purity commercially
-
3
-1
with a very small intercept (∼6 × 10 s ) compared to
-
1
observed rate constants between 0.1 and 1.3 s ; however,
it is unlikely that this is an equilibrium given that the reverse
reaction is thermodynamically unfavorable. Because the
a
3-
2-
available. Solutions of [Fe(CN)
5 3
N ]
and [IrCl
6
]
were used
immediately after preparation to limit possible hydrolysis. Solutions
were prepared fresh daily and were maintained under an Ar
atmosphere.
reaction was carried out at pH 3 and the pK
a
of PA is 9.29,
3+
the electron-transfer reaction must occur between [Fe(H
2
6
O) ]
2
+
or [Fe(H O) (OH)] and C H SO NHOH, as very little
2
5
6
5
2
The kinetics data were collected on a Dionex stopped-flow
instrument interfaced with an OLIS data acquisition/data reduction
system or on a Hewlett-Packard model 8452A diode array spec-
trophotometer followed by data reduction using OLIS fitting
routines. These routines fit the absorbance change data to the
deprotonated PA is present at pH 3.
At an ionic strength of 0.2 M the kobs vs PA concentration
plot for the iron(III) oxidation of PA exhibits saturation
behavior. In contrast to the observation at low ionic strength,
this indicates formation of a weak Fe -PA “intermediate”
at high ionic strength. The data conform to the expression
3
+
1
exponential function ∆Abs ) exp(-k t) for pseudo-first-order data
sets and ∆Abs ) exp(-k t) + k t for mixed zeroth- and first-order
1
o
data sets. The fitting routine documentation along with the kinetic
data are available in the Supporting Information. All reactions were
studied at the absorption maximum of the metal complex and were
carried out at 25.0 °C and in 0.1 M sodium phosphate to control
pH. For aqueous iron(III), the reaction was studied using a 20-fold
excess of phenanthroline. This rapidly produces a red species at
k K [PA]
et eq
k
)
obs
1
+ K [PA]
eq
where ket is the first-order rate constant for electron transfer
and Keq is the formation constant for the Fe -PA interme-
3+
diate. At pH 3, the values for ket and Keq are 0.85 and 1.0
5
10 nm. The phosphate buffer also set the ionic strength to ∼0.2
-1
-1
-
-
s
and 290 and 250 M for 0.2 M Cl and ClO
4
solutions,
M over the limited pH range for this study. For the aquairon(III)
study, the pH was set using HCl, and the ionic strength was
approximately 0.01 M in the absence of added salts or 0.2 M in
3+
respectively. This indicates that observation of the Fe -
PA intermediate is a consequence of the high ionic strength
and not a specific ion effect.
4
the presence of added KCl or NaClO . The kinetic data were fitted
using Origin linear fitting routines. The first-order rate constants
are typically (10%. The errors on the zeroth-order rate constants
vary widely and are typically (20-30%. Concentration ranges of
the oxidants are given in the Supporting Information and are
typically a 10- to 100-fold excess over the Piloty’s acid to comply
with pseudo-first-order conditions.
Because of the ambiguity with respect to the nature of
the oxidant, i.e., the extent of deprotonation, and the
(6) Shirota, F. N.; DeMaster, E. G.; Lee, M. J. C.; Nagasawa, H. Nitric
Oxide 1999, 3, 445.
(
7) Schilt, A. A. J. Am. Chem. Soc. 1960, 80, 3000. Brauer, G., Ed.
Handbook of PreparatiVe Inorganic Chemistry; Academic Press: New
York, 1963; Vol. 2.
Ionic reaction products were identified chromatographically on
2-
-
(8) Hoshino, M.; Maeda, M.; Konishi, R.; Seki, H.; Ford, P. C. J. Am.
Chem. Soc. 1996, 118, 5702. Laverman, L. E.; Wanat, A.; Oszajca,
J.; Stochel, G.; Ford, P. C.; van Eldik, R. J. Am. Chem. Soc. 2001,
123, 285.
a Dionex model 2020i ion chromatograph using a CO
3
/HCO
3
eluent solution. Details of this separation are available in the
Supporting Information.
7878 Inorganic Chemistry, Vol. 43, No. 24, 2004

Oxidation of Piloty’s Acid by Metal Complexes
possibility of both inner- and outer-sphere electron transfer
with iron(III), the investigation of known outer-sphere
oxidants was undertaken. The results of the oxidation of PA
by a variety of oxidants under anaerobic conditions and over
a range of pHs were examined. The evolution of nitric oxide
was directly observed with some oxidants, whereas other
oxidants led to the formation of nitrite, presumably by rapid
further oxidation of the initially produced NO.
Nitric Oxide Producing Oxidants. Oxidation of PA by
3-
3-
[Fe(CN)
6
]
or [Fe(CN)
5 3
N ]
produced the reduced metal
complex, benzenesulfinate, and NO via a single electron
transfer
3
-
-
[
Fe(CN)6] + C H SO NHO f
6 5 2
Fe(CN)6] - + C H SO + NO + H
4
-
+
[
6 5 2
The gas bubbles observed in the product solutions were again
shown to be nitric oxide by scavenging with metmyoglobin
+
(
Mb )
OH-
+
+
0
Mb + NO f Mb NO
8 Mb NO
+
0
Although Mb can oxidize PA directly to form Mb NO, this
reaction is very slow (Table 1) and does not interfere with
+
the scavenging studies. The rapid formation of the Mb NO
3
-
-4
adduct therefore shows that NO production is from the
Figure 1. (a) Zeroth-order trace for reaction of [Fe(CN)6] (3.5 × 10
-3
+
M) with PA (1.0 × 10 M) at pH 7.5, T ) 25 °C, [phosphate] ) 0.1 M,
primary reaction and not from the added Mb (t1/2 for
+
and 420 nm. (b) First-order trace for reaction of [Fe(bipy)2(CN)2] (5.7 ×
+
8
Mb NO formation <10 s, pH 7 and 7.5).
-5
-3
10
M) with PA (1.2 × 10 M) at pH 7, T ) 25 °C, [phosphate] ) 0.1
The kinetic experiments were performed under pseudo-
first-order conditions with PA in excess. Reactions were
followed by absorbance decreases for the formation of the
iron(II) products. The rate data are reported in Table 1 along
with relevant thermodynamic and kinetic information for
M, and 522 nm.
drolysis reaction is too slow around pH 7 (t1/2 ≈ 90 h) and
hydroxylamine was not found as a decomposition product
5
a
by other investigators. Because it is also well-known that
aqueous metal ions can catalyze hydrolysis, this pathway
must be considered. Hence, we have briefly studied the
oxidation of hydroxylamine by ferricyanide and showed with
the Mb scavenging technique that NO was not produced at
pH 7.5. Therefore, any mechanism involving the initial
formation of hydroxylamine can be discounted.
each oxidant.
Although the oxidation of NO to NO
-
2
is extremely facile
9
3-
(
E° ) -0.46 V vs NHE in base ), neither [Fe(CN)
6
]
nor
+
3-
[Fe(CN)
5 3
N ]
rapidly oxidizes NO further under the condi-
3-
tions of this study. However, when [Fe(CN)
6
]
was used
3
-
in excess, a second slower conversion of [Fe(CN)
6
]
to
4-
[Fe(CN)
6
]
was observed, presumably because of further
The second possibility is a rate-determining breakdown
of an “aggregate” form of PA (dimer, trimer, etc.) into a
monomeric form that reacts with the metal complex. If true,
then at high pH (>9.5), where the PA is deprotonated,
electrostatic repulsion should augment aggregate breakdown,
and the reaction profile should “reduce” to simple first-order
behavior. This is what is observed. The amount of first- and
zeroth-order behavior should also depend on the PA con-
centration under investigation. At the lowest PA concentra-
tions studied, where more aggregate dissociation occurs, the
amount of zeroth-order character for the disappearance of
ferricyanide is significantly less than at higher concentrations.
This observation is consistent with PA existing in an
equilibrium involving an associated form in solution. It is
tempting to ascribe this behavior to hydrogen bonding, but
such an interaction seems unlikely to be strong in aqueous
media. However, the crystal structure of PA has been solved,
and the authors point out that there is “extensive hydrogen
bonding” both within and between the layers of PA
oxidation of the NO produced in the first electron-transfer
step. No further studies of this reaction were undertaken.
With excess PA, only benzenesulfinate and NO are produced.
3-
3-
Surprisingly, when [Fe(CN)
oxidant, the reaction is mixed zeroth- and first-order in
Fe(III)] complex, and the order is somewhat dependent on
6
]
5 3
or [Fe(CN) N ]
was the
[
pH and ionic strength; see Figure 1a. Because it is difficult
to ascribe zeroth-order behavior to either oxidant in this
reaction, the effect likely arises from a rate-determining
production of an “active form” of PA.
There are three possibilities for the observed zeroth-order
behavior. The first is the hydrolysis of PA into hydroxyl-
amine and benzenesulfinate followed by rapid reduction of
Fe(III) by the free hydroxylamine. This pathway can be
excluded from consideration, however, given that the hy-
(
9) CRC Handbook of Chemistry and Physics, 51st ed.; Weast, R., Ed.;
CRC Press: Boca Raton, FL, 1970; p D112. deMaine, M. M.;
Stanbury, D. M. Inorg. Chem. 1991, 30, 2104.
Inorganic Chemistry, Vol. 43, No. 24, 2004 7879

Wilkins et al.
molecules in the solid state.10 This hydrogen-bonding
network might not apply to the solution state. As an initial
probe toward better defining this aggregation, the reaction
was carried out in 0.25 M urea (pH 7.5) with ferricyanide.
Neither the zeroth- nor the first-order portion of the reaction
was affected.
The final possibility involves a rate-determining proton
migration from nitrogen to oxygen of the anion of PA prior
to the oxidation. To explore this possibility, we examined
above could react directly with the oxidant rather than just
the “active” PA. This could be due to the higher reduction
potentials of these oxidants increasing the driving force for
the oxidation. Similar behavior has been observed for the
1
7
ferricyanide oxidation of dithionite ion.
Plots of the rate constants for all of the oxidants vs PA
concentration were linear over a 10-fold concentration range
of the sulfonamide and exhibited a zero y intercept. The data
conform to the rate expression
the ferricyanide oxidation of PA in D
2
O at a pD equivalent
-d[oxidant]/dt ) k[oxidant][PA]
to pH 7.6. The ratio k /k from these studies was 6.9, which
H
D
is surprisingly high. This suggests that proton migration
might be the cause of the zeroth-order behavior. To confirm
that deuterium exchange was rapid, a sample of PA was
The second-order rate constant is inversely proportional to
+
+
the H concentration for [Fe(bipy)
2
(CN)
2
] and conforms
to the following rate expression
2
dissolved in deuterated chloroform, and a drop of D O was
k ) k + k K
AH[H+]-1
H O
added. Deuterium exchange was complete within 5 min, the
time required to acquire a new spectrum. Similar zeroth order
behavior has been observed for the oxidation of phosphorous
+
From the k vs inverse H concentration data, the following
-
1 -1
1
6
values were obtained: k
H
) 11 M
s
and k
O
K
AH ) 6.0 ×
acid. This was ascribed to an active and inactive form of
the P(III) center. The zeroth-order component was ascribed
to migration of a hydrogen from the phosphorus to oxygen.
It is possible that analogous behavior might be observed in
the ferricyanide oxidation of PA.
-
5
-1
1
0
s . As in the previous section, deprotonation of PA
5
-1 -1
was assumed, and k
O
is calculated to be 1.2 × 10 M s .
Although the direct detection of NO was not possible with
these oxidants, from a mechanistic standpoint, one can infer
that the initial one-electron oxidation of PA produces nitric
Nitrite/Nitrous Acid Producing Oxidants. Oxidation of
+
+
2-
2-
oxide, which is then rapidly oxidized to NO by a second
2
PA by [Fe(bipy) (CN)
2
] , [IrBr
6
] , or [IrCl
6
]
was rapid
mole of oxidant. The lifetime of the nitrosonium ion in water
and formed 2 mol equiv of reduced metal complex, one of
benzenesulfinate, and one of nitrous acid as shown below.
The benzenesulfinate and nitrite products were identified and
quantified by capillary ion chromatography.
-
10
+
is very short (∼ 3 × 10 s), and NO rapidly hydrolyzes
1
8
+
to nitrite as shown below for [Fe(bipy)
2 2
(CN) ]
+
-
[
Fe(bipy) (CN) ] + C H SO NHO f
2 2 6 5 2
-
+
+
[Fe(bipy) (CN) ] + C H SO + NO + H
2
[Fe(bipy) (CN) ] + C H SO NHOH + H O f
2
2
6
5
2
2
2
6
5
2
2
-
+
2
[Fe(bipy) (CN) ] + C H SO + HNO + 3H
+
+
2
2
6
5
2
2
[Fe(bipy) (CN) ] + NO f [Fe(bipy) (CN) ] + NO
2 2 2 2
2
2
[IrCl6] - + C H SO NHOH + H O f
+
+
6
5
2
2
NO + H O f HNO + H
2 2
[IrCl6] - + C H SO + HNO + 3H
3
-
+
2
6 5 2 2
Although all of the metal complexes used in this study
are outer-sphere oxidants, i.e., they exchange their ligands
very slowly with respect to the electron-transfer rates, there
is no consistent relationship among the oxidation rates of
PA and the charge of the oxidant, the reduction potentials
of the oxidants (E°), or the self-exchange rates (kexc) of the
redox couples (Table 1). This lack of correlation suggests
that, rather than a “simple” outer-sphere electron-transfer
process, the PA reactions are complicated by rapid intra-
molecular steps within the initial oxidation product, followed
by an irreversible rapid breakdown to products. Because there
are structural differences between the oxidized and reduced
forms of PA, treatment with the Marcus cross relationship
is inappropriate.
The 2:1 oxidant-to-PA stoichiometry in these reactions was
established by the spectral changes observed for the
+
2-
3-
[
Fe(bipy)
2
(CN)
2
] to [Fe(bipy)
2
(CN)
2
], [IrBr
6
]
6
to [IrBr ] ,
2
-
3-
and [IrCl
6
]
to [IrCl
6
]
conversions.
As in the earlier systems, all kinetic experiments were
performed under pseudo-first-order conditions with PA in
excess over oxidant, and the electron-transfer rate constants
were determined by observing the absorbance change as-
sociated with formation of the Fe(II) complex or loss of the
Ir(IV) complex. The data for loss of oxidant at each PA
concentration were fit by a single exponential, and a typical
trace is shown in Figure 1b. No zeroth-order behavior was
observed. This indicates that the “inactive PA” described
Other Substrates. To examine the effect of modifying
the benzene ring in Piloty’s acid, three related N-hydroxy-
sulfonamides were synthesized, and their oxidation with
ferricyanide was studied. Two compounds in which the para
(
10) Scholtz, J. N.; Engel, P. S.; Glidewell, C.; Whitmire, K. H. Tetrahedron
1989, 45, 7695.
(
(
(
11) Pelizzetti, E.; Mentasti, E.; Pramauro, E. Inorg. Chem. 1978, 17, 1181.
12) Stasiw, R.; Wilkins, R. G. Inorg. Chem. 1969, 8, 156.
13) Greenwood, N. N.; Earnshaw, A. In Chemistry of the Elements, 2nd
ed.; Butterworth-Heinemann: Oxford, U.K., 1997; p 1093.
(17) Scaife, C.; Wilkins, R. G. Inorg. Chem. 1980, 19, 2242 and references
(
(
(
14) Chou, M.; Creutz, C.; Sutin, N. J. Am. Chem. Soc. 1977, 99, 5615.
15) Behlke, J.; Scheler, W. Acta Biol. Med. Ger. 1962, 8, 88.
16) Vista, A.; Holm, D.; Wang, P. L.; Veith, G. D. Inorg. Chem. 1971,
therein.
(18) Bonner, F. T.; Stedman, G. In Methods in Nitric Oxide Research;
Feelisch, M., Stamler, J. S., Eds.; Wiley: Chichester, U.K., 1996; pp
3-18.
10, 631. Mehrotra, R. N. J. Chem. Soc., Dalton Trans. 1978, 681.
7880 Inorganic Chemistry, Vol. 43, No. 24, 2004

Oxidation of Piloty’s Acid by Metal Complexes
Table 2. Rate Constants for Other N-Hydroxysulfonamide Oxidations
at 25 °C
jugation between the nitrogen center and the R group attached
to the sulfur.
oxidanta
pHb
k (M-1 s-1
)
Further attempts to examine the effects of structural
changes on the oxidation processes were undertaken by
preparation of N-methylated PA. Although the compound
could be synthesized, it proved to be too unstable in aqueous
media for study of its oxidation. Presumably, rapid hydrolysis
of the S-N bond occurred to form N-methylhydroxylamine
and benzenesulfinate.
p-bromoPA
[
Fe(bipy)2(CN)2]+
Fe(bipy)2(CN)2]+
7.0
1100
p-methoxyPA
[
7.0
740
N-hydroxy-n-butylsulfonamide
Fe(bipy)2(CN)2]+
[
7.0
.0
7.0
230
53
290
6
IrCl6]2
-
[
a
Anaerobic conditions. b Buffer ) 0.1M sodium phosphate.
Conclusions
As seen from this work, the nature of the oxidant is critical
in the generation of NO from Piloty’s acid. These preliminary
kinetic studies show that these “simple” reactions actually
occur via a series of steps that need to be systematically
unraveled. Such investigations could drive the development
of pharmaceutically useful NO precursors.
position in the benzene ring was substituted with either a Br
or OCH group showed little variation in rate and maintained
3
a mixed first- and zeroth-order kinetic profile. No correlation
with respect to electron-withdrawing or -donating character
was observed. Replacement of the phenyl ring with an
n-butyl group again showed little effect on the oxidation rate
and an analogous mixed order behavior was observed.
Judging from the similarities in rates of these widely different
compounds, it is possible that the outer-sphere oxidation
takes place directly at the nitrogen center and that conjugation
between the N and the R group, through the sulfur, is not
important in this system.
Acknowledgment. We acknowledge financial support
from the National Institutes of Health under PHS Grant SO6
GM08136. Thanks also go to Dr. Dennis Johnson for help
with the ion chromatography and to Dr. Ralph G. Wilkins
for his helpful discussions. We would also like to thank Ms.
Gloria Martinez for some technical assistance.
2
The oxidations of the same substrates with [Fe(bipy) -
+
2-
Supporting Information Available: Listing of details for ion
chromatography and tables of all kinetic data. This material is
available free of charge via the Internet at http://pubs.acs.org.
(CN)
2
] and [IrCl
6
]
were also studied; see Table 2. Each
reaction was nicely first-order with no evidence of a zeroth-
order component. The rates varied by only 5-fold among
the substituted analogues, which again suggests little con-
IC049627Z
Inorganic Chemistry, Vol. 43, No. 24, 2004 7881