Direct Observation of Acyl Nitroso Compounds in Aqueous Solution and the Kinetics of Their Reactions with Amines, Thiols, and Hydroxamic Acids

Article abstract of DOI:10.1021/acs.jpca.8b06672

Acyl nitroso compounds or nitrosocarhonyls (RC(O)N=O) are reactive short-lived electrophiles, and their hydrolysis and reactions with nucleophiles produce HNO. Previously, direct detection of acyl nitroso species in nonaqueous media has been provided by time-resolved infrared spectroscopy demonstrating that its half-life is about 1 ms. In the present study hydroxamic acids (RC(O)NHOH) are oxidized electrochemically in buffered aqueous solutions (pH 5.9-10.2) yielding transient species characterized by their maximal absorption at 314-330 nm. These transient species decompose via a first-order reaction yielding mainly HNO and the respective carboxylic acid and therefore are ascribed to RC(O)N=O. The sufficiently long half-life of RC(O)N=O in aqueous solution allows for the first time the study of the kinetics of its reactions with various nucleophiles demonstrating that the nucleophilic reactivity follows the order thiolate > hydroxamate > amine. Metal chelates of CH₃C(O)NHOH catalyze the hydrolysis of CH₃C(O)N=O at the efficacy order of Cu^II > Zn^II > Ni^II > Co^II where only Cu^II catalyzes the hydrolysis also in the absence of the hydroxamate. Finally, oxidation of hydroxamic acids generates HNO, and the rate of this process is determined by the half-life of the respective acyl nitroso compound.

Full text of DOI:10.1021/acs.jpca.8b06672

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A: Kinetics, Dynamics, Photochemistry, and Excited States
Direct Observation of Acyl Nitroso Compounds in Aqueous Solution and
the Kinetics of Their Reactions with Amines, Thiols and Hydroxamic Acids
Eric Maimon, Ana Lerner, Amram Samuni, and Sara Goldstein
J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.8b06672 • Publication Date (Web): 16 Aug 2018
Downloaded from http://pubs.acs.org on August 16, 2018
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Direct Observation of Acyl Nitroso Compounds in Aqueous Solution and the
Kinetics of their Reactions with Amines, Thiols and Hydroxamic Acids
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Eric Maimon[a,b], Ana Lerner[b], Amram Samuni[c], Sara Goldstein*[d]
[a] Nuclear Research Centre Negev, Beer Sheva, Israel
[b] Chemistry Department, BenꢀGurion University, BeerꢀSheva 84105, Israel
[c] Institute of Medical Research ꢀ Israel Canada, Medical School, The Hebrew University of
Jerusalem, Jerusalem 91120, Israel
[d] Institute of Chemistry, The Accelerator Laboratory, the Hebrew University of Jerusalem,
Jerusalem 91904, Israel
*Corresponding Author
S. Goldstein. Eꢀmail: sara.goldstein1@mail.huji.ac.il; Tel: 972 54 7659996
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Abstract
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Acyl nitroso compounds or nitrosocarhonyls (RC(O)N=O) are reactive shortꢀlived
electrophiles, and their hydrolysis and reactions with nucleophiles produce HNO. Previously, direct
detection of acyl nitroso species in nonꢀaqueous media has been provided by timeꢀresolved infrared
spectroscopy demonstrating that its halfꢀlife is about 1 ms. In the present study hydroxamic acids
(RC(O)NHOH) are oxidized electrochemically in buffered aqueous solutions (pH 5.9 – 10.2)
yielding transient species characterized by their maximal absorption at 314 – 330 nm. These
transient species decompose via a firstꢀorder reaction yielding mainly HNO and the respective
carboxylic acid, and therefore are ascribed to RC(O)N=O. The sufficiently long halfꢀlife of
RC(O)N=O in aqueous solution allows for the first time the study of the kinetics of its reactions
with various nucleophiles demonstrating that the nucleophilic reactivity follows the order thiolate >
hydroxamate > amine. Metal chelats of CH3C(O)NHOH catalyze the hydrolysis of CH3C(O)N=O at
the efficacy order of CuII > ZnII > NiII > CoII where only CuII catalyzes the hydrolysis also in the
absence of the hydroxamate. Finally, oxidation of hydroxamic acids generates HNO and the rate of
this process is determined by the halfꢀlife of the respective acyl nitroso compound.
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Introduction
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Acyl nitroso compounds or nitrosocarhonyls (RC(O)N=O) have been proposed as transient
intermediates during periodate oxidation of hydroxamic acid (RC(O)NHOH, HX),1ꢀ4 by thermal
fragmentation of DielsꢀAlder adducts5ꢀ7 and by photolysis of various precursors including 9,10ꢀ
dimethylanthracene adducts, nitrodiazo compounds, and 1,2,4ꢀoxadiazoleꢀ4ꢀoxides8. Acyl nitroso
compounds are reactive electrophiles, and their reactions with nucleophiles (NuH) produce nitroxyl
(HNO)4,8ꢀ11, which undergoes dimerization followed by dehydration yielding nitrous oxide (N2O)12
and/or reacts with various substrates such as thiols13.
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RC(O)N=O + NuH → RC(O)Nu + HNO
(1)
No direct detection of RC(O)N=O in aqueous solution have been previously reported. In most cases
where the intermediacy of RC(O)N=O has been invoked, its generation has been indirectly
confirmed by trapping it with conjugated 1,3ꢀdienes that produces stable cycloadducts4 or by
detection of N2O4ꢀ6,8,11,14,15. Spectroscopic evidence has been reported for HC(O)N=O in the gas
phase,16 and various RC(O)N=O by timeꢀresolved infrared spectroscopy where the lifetime has
been estimated to be about 1 ms in nonꢀaqueous solutions8,17
.
We have previously demonstrated that in aqueous solutions oneꢀelectron oxidation of
acetohydroxamic acid and suberoylanilide hydroxamic acid (SAHA) by radiolyticallyꢀborne
radicals such as •OH, •N3 and Br2•– forms the respective transient nitroxide radicals
(RC(O)NHO•/RC(O)NO•–), which decompose bimolecularly yielding transient species with a
maximum absorption at 314 nm.10,11 We have detected these transient species also upon oxidation
of RC(O)NO•– by •NO2 and during the enzymatic oxidation of HX by horseradish
peroxidase/H2O2.18 The decomposition of these transient species yields N2O, their halfꢀlife
decreases upon increasing the pH and/or [HX], and therefore they have been assigned to
RC(O)N=O.10,11,18 Yet, their halfꢀlife in aqueous media is orders of magnitude longer than that
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reported in nonꢀaqueous solutions.8,17 In the present study various HXs are oxidized
electrochemically yielding transient species with a maximum absorption at 314 – 330 nm that
decompose to yield N2O and the respective carboxylic acid. These transient species are attributable
to acyl nitroso compounds, and their sufficiently long halfꢀlife allows for the first time the study of
their kinetics with various nucleophiles in aqueous solutions.
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Experimental Section
Materials
Water for preparation of the solutions was purified using a MilliꢀQ purification system. All
chemicals purchased from SigmaꢀAldrich were of the highest available grade and were used as
received: acetohydroxamic acid (acetoꢀHX), suberohydroxamic acid (suberoꢀHX),
benzohydroxamic acid (benzoꢀHX), glycine hydroxamic acid (glyꢀHX), αꢀaminobutyric
hydroxamic acid (AABAꢀHX), Lꢀlysine hydroxamate hydrochloride (lysꢀHX), deferoxamine
mesylate salt (DFO), acetic glacial acid (99.8%), benzoic acid, dimethylamine hydrochloride,
ethylamine (70% solution), morpholine, Lꢀcysteine, glutathione (GSH), Nꢀacetylcysteine (NAC),
glycine, (3ꢀ(Nꢀmorpholino) propanesulfonic acid (MOPS), 2ꢀ(NꢀMorpholino)ethanesulfonic acid, 4ꢀ
morpholineethanesulfonic acid (MES), [tris(hydroxymethyl)methylamino]propanesulfonic acid
(TAPS) and Cu(ClO4).6H2O. FeCl3 6H2O, NiSO4 H2O and ZnSO4 7H2O were obtained from British
Drug Houses (BDH). Suberoylanilide hydroxamic acid (SAHA) was purchased from LC
Laboratories. Stock solutions of SAHA and suberoꢀHX were prepared in dimethyl sulfoxide
(DMSO). Stock solutions of amines and thiols were prepared in the same buffer and pH studied.
Most of the experiments were carried out in the presence of organic buffers to avoid the interference
of phosphate ions with the binding of metal ions to the hydroxamic acids.
.
.
.
Nitrite was assayed using the Griess reagent. The absorption at 540 nm was read 15 min after
mixing the reagent with the sample. Calibration curves were prepared using known concentrations
of nitrite.
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Bulk electrolysis
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The applied voltage was controlled using the EmStatꢀPalmSens USB powered potentiostat
electrochemical interface. Homeꢀmade electroꢀchemical cell was used for bulk electrolysis. The cell
consisted of a working electrode of graphite grains packed inside a porous Vycor glass tube (5 mm
I.D.), through which the solutions were pumped. An outer glass cylinder, with separate electrolyte
(10 mM phosphate buffer, pH 7.0) contained the platinum auxiliary electrode and Ag/AgCl (3.5 M)
as a reference electrode (205 mV vs. NHE). Solutions were pumped at 165 ꢀL sꢀ1 flow rate into a 1
cm optical cuvette positioned within a HP 8453 diode array spectrophotometer and/or into a glass
vial sealed with a rubber septum for N2O determination by gas chromatography. All experiments
were carried out at room temperature.
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Gas chromatography (GC)
GC was used to determine the yield of N2O, which is the product of HNO dimerization followed by
dehydration. Anoxic solutions of HXs (6.5 – 7.5 mL) were pumped through the electrolysis cell into
anoxic glass vials (10.7 mL). A gas aliquot of the reaction headspace (2 mL) was taken and 1 mL
was injected at room pressure and temperature onto a 5890 HewlettꢀPackard gas chromatograph
equipped with a thermal conductivity detector, a 10 ft ꢀ 1/8 inch Porapak Q column at an operating
oven temperature of 70 °C (injector and detector 150 °C) and a flow rate of 20 mL/min (He, carrier
gas). The yields of N2O were calculated on the basis of a standard curve prepared by injecting 1 mL
of known amounts of N2O gas (Maxima, Israel).
High Performance Liquid Chromatography (HPLC)
HPLC was used to determine the yields of benzoic acid and acetic acid as endꢀproducts of the
electroꢀoxidation of benzoꢀHX and acetoHX, respectively. The HPLC system (Knauer, Germany)
was equipped with an Azura P6.1L pump and an Azura UVD 2.1L detector. Samples (20 ꢀL) were
injected onto a YMC C18 column (150 x 4.6 mm I.D, 3 ꢀm). Chromatographic separations were
performed at room temperature using an isocratic mobile phase consisting of a 50:50 solution of
acetonitrile:phosphate buffer pH 2.6 in the case of benzoꢀHX and benzoic acid, and a 10:90 solution
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of the same solvents for samples containing acetoꢀHX and acetic acid. The flow rate was 1.0 ml
minꢀ1. In order to reduce solvent noise in the chromatogram, samples were first evaporated in a
rotorꢀevaporator and then dissolved in the same solvent of the mobile phase. The analyte
concentration remained constant throughout the process. Solutions of benxoꢀHX and benzoic acid
were used as standards. Detection was at 215 nm. Peak identification was based on retention times
compared with those of selected standards. Quantitation was based on calibration curves obtained
by plotting peak area vs. concentration (0 ꢀ 500 ꢀM).
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AcetoꢀHX and acetic acid were analyzed using the standard addition method, and detected at 195
nm. Acetic acid and acetoꢀHX at different concentrations ranging between 30% and 250% of the
analyte concentration were added to a known volume of analyte solution. The addition volume did
not exceed 1% of the total sample volume.
Results and Discussion
The structures of HXs used in this study and their pKa(ꢀNHOH) are given in Fig. 1.
R = CH3, acetoꢀHX (9.0)10,19
R = C6H5, benzoꢀHX (8.54)19
R = CH2NH2, glyꢀHX (7.4 ꢀ 7.8)20ꢀ22
R = CH(NH2)(CH2)3CH2NH2, lysꢀHX (7.9 – 8.0)20,23
R = CH(NH2)CH2CH3, AABAꢀHX
suberoꢀHX (8.86, 9.42)24
SAHA (9.6)11
DFO (8.7, 9.0, 10.1)25
Figure 1. Structures of HXs and their pKa(ꢀNHOH) given in paranthesis.
Electroꢀoxidation of HXs in 0.1 M KCl aqueous media buffered at pH 5.9 – 10.2 generates
relatively longꢀlived transient species having a maximum absorption at 314 – 330 nm. Since there is
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a limited data on E1/2 and onset potential values of acetoꢀHX in aqueous media,26ꢀ28 the applied
voltage was varied until a maximal absorption was achieved. Figure 2 shows the spectra of the
transient species derived from electroꢀoxidation of acetoꢀHX, benzoꢀHX, SAHA and suberoꢀHX at
pH 7.3. These transient species decay via a first order reaction (Figure 2, inset, τ1/2 > 1.5 min)
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implying that any depletion during the measurements (ca. 10 s) is negligible. Same kd and ε314
values were obtained under anoxic conditions.
1.0
A
B
0.2
1.0
A
0.1
0.2
0.5
A
0.8
0.0
0
4
8
12
Time (min)
0.0
0.6
0.4
0.2
0.0
0
10
20
30
40
Time (min)
0.1
0.0
270
300
330
360
390
420
450
250
300
350
400
450
λ
(nm)
λ
(nm)
0.3
0.2
0.1
0.0
0.3
0.2
0.1
0.3
0.2
0.1
C
D
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0.2
0.1
0.0
A
0.0
0
0.0
0
10
20
30
5
10
15
Time (min)
Time (min)
250
300
350
400
450
280
320
360
400
440
λ
(nm)
λ
(nm)
Figure 2. Spectra of transient species formed upon electroꢀoxidation of HXs in aerated solutions containing
0.1 M KCl and 20 mM MOPS buffer at pH 7.3. Inset: first order kinetics of the decay monitored at λmax. (A)
500 ꢀM acetꢀHX, kd = 0.19 ± 0.01 minꢀ1; (B) 150 ꢀM benzoꢀHX, kd = 0.46 ± 0.02 minꢀ1; (C) 230 ꢀM SAHA,
kd = 0.14 ± 0.01 minꢀ1; (D) 95 ꢀM suberoꢀHX, kd = 0.21 ± 0.01 minꢀ1.
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The εmaxꢀvalues have been calculated from the dependence of Amax on [HX]o at the concentrations
range where Amax increases proportionally with [HX]o (Figure 3). These calculations are based on
the assumption that no appreciable depletion of the transient species thus formed takes place until
measurement, and that one molecule of HX yields one molecule of a transient species.
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subero-HX, 314 nm
1.0
benzo-HX, 330 nm
SAHA, 314 nm
aceto-HX, 314 nm
0.8
0.6
0.4
0.2
0.0
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200
[HX]o (
300
M)
400
500
ꢀ
Figure 3. Dependence of Amax of the transient species derived from electroꢀoxidation of HX on [HX]o in 0.1
M KCl and 20 mM MOPS buffer at pH 7.3.
In the case of alkylamino hydroxamic acids, which are less basic than alkyl hydroxamic acids
(Figure 1), the halfꢀlife of the transient species are relatively short (Figure 4, inset) and highly
dependent on [HX]o. In these cases a lower limit for εmax is set at relatively low [HX]o. The εmaxꢀ
values are summarized in Table 1.
0.03
0.03
A
B
0.03
0.02
0.01
0.00
0.03
0.02
0.01
0.00
0.02
0.01
0.00
0.02
0.01
0.00
0
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0
1
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3
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5
Time (min)
Time (min)
250
300
350
400
(nm)
450
500
250
300
350
400
450
500
λ
λ
(nm)
Figure 4. Spectra of transient species formed by electroꢀoxidation of 15 ꢀM AABAꢀHX (A) or glyꢀHX (B)
in 0.1 M KCl and 20 mM MOPS buffer at pH 7.3. Inset: firstꢀorder decay kinetics of the transient species
monitored at 325 nm, kd = 1.56 ± 0.08 minꢀ1 (A) and at 317 nm, kd = 0.96 ± 0.05 minꢀ1(B).
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Table 1. εmaxꢀvalues of the transient species formed by electroꢀoxidation of HX.*
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HX
λ
max (nm)
314
314
330
314
ε (Mꢀ1cmꢀ1)
3050 ± 150
2150 ± 100
1650 ± 100
1450 ± 100
≥ 2200
suberoꢀHX
acetoꢀHX
benzoꢀHX
SAHA
glyꢀHX
317
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AABAꢀHX
325
≥ 2200
*Assuming that electroꢀoxidation of HX is complete and that each molecule of HX yields one molecule of a
transient species. All solution contained 0.1 M KCl and 20 mM MOPS buffer at pH 7.3.
The decomposition of the transient species produces N2O formed via HNO dimerization
followed by dehydration. Accumulation of N2O was determined following electroꢀoxidation of HX
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in 0.1 M KCl and 20 mM MOPS buffer at pH 7.3 under anoxia to avoid HNO reaction with O2 .
Anoxic solutions (6.5 – 7.5 mL) were pumped through the electrolysis cell into an anoxic glass vial
and 1 mL into an optical cuvette for the determination of the oxidation yield (Amax). A gas aliquot of
the reaction headspace was taken either immediately following the electroꢀoxidation or 40 min later.
Accumulation of N2O immediately following the electroꢀoxidation was below the detection limit
whereas that measured when the decay of the transient species was completed depends on the initial
concentration of the transient species (Table 2).
Table 2. N2O yields formed upon electroꢀoxidation of HX at pH 7.3 (20 mM MOPS buffer).
HX
acetoꢀHX
Amax
0.21a
0.37
1.08
1.61
HNO yield, %
[transient]o, ꢀM
[N2O], ꢀM
98
35
63
182
260
71
73
72
69
172
500
750
benzoꢀHX
SAHA
0.23
0.29
0.33
139
176
200
61
87
91
88
99
91
0.20
0.29
0.20
0.29
30
45
44
45
The experimental error is ±10%.
a – 10 mM phosphate buffer.
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Accumulation of nitrite at the end of the electroꢀoxidation process was the same as that measured
when the decay of the transient species was completed. In the case of SAHA (where N2O yield is
relatively low) and in the case of benzoꢀHX (where N2O yield is maximal), the accumulation of
nitrite was less than 4% implying that there is no clear reciprocal ratio between the yields of HNO
and nitrite.
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HPLC measurements were performed to identify and quantify acetic acid and benzoic acid as endꢀ
products of the electroꢀoxidation of acetoꢀHX and benzoꢀHX, respectively, at pH 7.3. The results
demonstrate that each molecule of the transient species decomposes to yield one molecule of the
respective carboxylic acid (Table 3). The residual concentration of benzoꢀHX was negligible, and it
was significant (114 ꢀM) in the case where oxidation of 1 mM acetoꢀHX was incomplete resulted
in 837 ꢀM transient species.
Table 3. Carboxylic acid yields formed upon electroꢀoxidation of HX at pH 7.3.
HX
acetoꢀHX
Amax
1.08
1.61
1.80
RC(O)OH yield, %
[transient]o, ꢀM [RC(O)OH], ꢀM
500
750
837
479
773
818
96
103
98
benzoꢀHX
0.073
0.125
0.147
0.21
0.256
0.27
44
76
89
127
155
164
260
41
85
83
115
160
146
262
93
112
93
91
103
89
0.43
101
The experimental error is ±10%.
The transient species is assigned to RC(O)N=O since upon hydrolysis it generates HNO and
the respective carboxylic acid. In addition, the rate of its decomposition increases as the pH
increases and as shown below it reacts also with various nucleophiles. RC(O)N=O can be formed
via two consecutive oneꢀelectron oxidation processes (reactions 2 and 3) or a single oneꢀelectron
oxidation process followed by dismutation of the nitroxide radical (reactions 2 and 4).
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RC(O)NHOH → RC(O)NHO• + H+ + e–
(2)
(3)
(4)
RC(O)NHO• → RC(O)N=O + H+ + e–
2 RC(O)NHO• → RC(O)NHOH + RC(O)N=O
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Both processes lead to the oxidation of RC(O)NHOH to RC(O)N=O, which decomposes via
hydrolysis yielding the respective carboxylic acid and HNO (reaction 5) that subsequently forms
N2O (equation 6).
RC(O)N=O + H2O → HNO + RC(O)O– + H+
2 HNO → HONNOH → N2O + H2O
(5)
(6)
HNO yields lower than 100% imply that while oneꢀelectron oxidation takes place, RC(O)NHO•
might also decompose via dimerization yielding products other than HNO (reaction 7).
2 RC(O)NHO• → unidentified products
(7)
The low yields of nitrite formed only during the electrolysis process could be attributed to
RC(O)N=O oxidation to NO (reaction 8), which in aerated solutions yields nitrite.
RC(O)N=O + H2O → RC(O)O– + NO + 2H+ + e–
(8)
Irrespective of the detailed mechanism, the relative long halfꢀlife of RC(O)N=O in aqueous
solutions allows the study of its reactions with various nucleophiles as demonstrated below.
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Reaction with amines
3
4
The decay of RC(O)N=O upon its reaction with various amines obeys firstꢀorder kinetics
5
6
7
8
9
and the observed first order rate constant (kobs) increases linearly upon increasing the concentrations
+
+
of dimethylamine (pKa(ꢀNH3 ) = 10.73), ethylamine (pKa(ꢀNH3 ) = 10.7) or morpholine (pKa(ꢀ
+
10
11
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60
NH3 ) = 8.36) at constant pH (Figure 5A), and increases also as the pH increases (Figure 5B)
implying that the reacting species is the basic form of the amine.
RC(O)N=O + R'NH2 → RC(O)NHR' + HNO
(9)
0.12
0.16
0.12
0.08
0.04
0.00
B
A
benzo
subero
pH 10.1
0.08
0.04
0.00
SAHA
k
k
pH 9.5
0
4
8
12
16
0
2
4
6
[dimethylamine] (mM)
[ethylamine] (mM)
Figure 5. (A) Dependence of kobs of the decay of RC(O)N=O derived from benzoꢀHX (
and SAHA (
g
), suberoꢀHX (ꢀ)
n
) on [dimethylamine] at pH 9.4; (B) Dependence of kobs of the decay of CH3C(O)N=O on
[ethylamine]o at pH 9.5 and 10.1. All solutions contained 0.1 M KCl and 50 mM borate buffer.
The apparent bimolecular rate constant (kapp) calculated from the slope of the lines such as those
presented in Figure 5 are summarized in Table 4.
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Table 4. The apparent bimolecular rate constant (Mꢀ1sꢀ1) of the reaction of RC(O)N=O with amines.
5
6
7
8
9
amine
pH
C5H6C(O)N=O CH3C(O)N=O
RC(O)N=O
from SAHA
RC(O)N=O
from suberoꢀHX
dimethylamine
7.2
7.8
8.4
9.2
9.4
10.2
9.5
10.1
10.1
0.28 ± 0.01
13.2 ± 0.8
0.25 ± 0.01
0.39 ± 0.02
1.32 ± 0.08
5.9 ± 0.3
0.32 ± 0.02
6.1 ± 0.2
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4.7 ± 0.3
31.7 ± 1.8
8.4 ± 0.6
25.4 ± 0.9
< 0.1
ethylamine
triethylamine
morpholine
7.5
9.4
1.9 ± 0.1
12.0 ± 0.3
10.0 ± 0.6
4.4 ± 0.2
All solutions contained 0.1 M KCl and either 20 mM MOPS (pH 7.2 – 7.8), 20 mM TAPS (pH 8.4) or 50
mM borate (pH 9.2 – 10.2) buffers.
In the case of CH3C(O)N=O reaction with dimethylamine, a fit of kapp (Table 4) to a pKꢀcurve using
pKa(ꢀNH3 ) = 10.73 resulted in k9 = 154
±
5 Mꢀ1sꢀ1 assuming that the reaction with (CH3)2NH3 is
+
+
negligible. This value is similar to that obtained with ethylamine and is more than an order of
magnitude higher than that with morpholine.
The rate of CH3C(O)N=O decay at pH 10.1 (50 mM borate buffer) was unaffected by10 – 60 mM
+
triethylamine (pKa(ꢀNH3 = 11.0) implying that the reaction is too slow most probably due to steric
hindrance.
Reaction with thiols
We have studied the reaction of RC(O)N=O with various thiols (R'SH) including cysteine
+
+
(pKa(ꢀSH) = 8.18, pKa(ꢀNH3 ) = 10.36)30, GSH (pKa(ꢀSH) = 8.7, pKa(ꢀNH3 ) = 10.1)30 and NAC
(pKa(ꢀSH) = 9.51)30. The decay of RC(O)N=O obeys first order kinetics in the presence of R'SH and
kobs increases linearly upon increasing [R'SH] and increases also as the pH increases (Figure 6)
implying that the reaction proceeds with R'S– (reaction 10).
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RC(O)N=O + R'S– → RC(O)SR' + NO–
(10)
3
4
5
6
7
8
0.20
0.15
0.10
0.05
0.00
0.16
0.12
0.08
0.04
0.00
B
A
benzo
pH 7.3
9
10
11
12
13
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59
60
k
k
aceto
pH 6.6
0.0
0.5
1.0
1.5
2.0
2.5
0
1
2
3
[cysteine] (mM)
[GSH] (mM)
Figure 6. (A) Dependence of kobs of the decay of CH3C(O)N=O on [GSH] at pH 6.6 and 7.3; (B)
Dependence of kobs of the decay of C5H6C(O)N=O or CH3C(O)N=O on [cysteine] at pH 6.6. All solutions
contained 0.1 M KCl and 20 mM MOPS buffer.
The experiments were carried out at pH << pKa(ꢀSH) since at higher pHs the rates become too fast
for measurements under pseudoꢀfirstꢀorder conditions using the current experimental setup. The
kappꢀvalues derived from the slope of the lines such as those in Figure 6 are summarized in Table 5.
Table 5. The apparent bimolecular rate constant (Mꢀ1sꢀ1) of the reaction of RC(O)N=O with thiols in 0.1 M
KCl and 20 mM MOPS buffer.
RC(O)N=O
formed from formed from
RC(O)N=O
RC(O)N=O
formed from
AABAꢀHX
thiol
pH C5H6C(O)N=O CH3C(O)N=O
SAHA
glyꢀHX
cysteine
6.6
7.3
165 ± 9
56 ± 3
250 ± 15
GSH
NAC
6.4
6.5
6.6
7.3
1150 ± 70
660 ± 70
145 ± 10
59 ± 2
240 ± 10
27 ± 1
91 ± 3
21 ± 1
77 ± 4
6.4
6.5
7.3
7.9
172 ± 10
535 ± 20
22 ± 1
54 ± 3
9.5 ± 0.2
37 ± 2
7.8 ± 0.1
23 ± 1
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7
8
Assuming that the rate of R'SH reaction with RC(O)N=O is negligible, one calculates k10 = (6 – 8) x
103 Mꢀ1sꢀ1 for C5H6C(O)N=O reaction with the thiolate ions of GSH, cysteine and NAC. In the case
of CH3C(O)N=O and the acyl nitroso derived from SAHA, k10 is about 2 ꢀ 3 folds lower. The acyl
nitroso derived from glyꢀHX and AABAꢀHX at pH < 9 are positively charged, and this could be the
reason for the relatively higher rate constant, i.e., k10 = (1 – 5) x 105 Mꢀ1sꢀ1.
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Reaction with hydroxamic acids
The rate of the first order decay of RC(O)N=O (kobs) increases linearly upon increasing
[HX] added following the oxidation and increases also as the pH increases (Figure 7). These results
imply that RC(O)N=O reacts with X– (reaction 11) and at constant pH the rate depends on the pKa(ꢀ
NHOH) of the reacting HX.
RC(O)N=O + RC(O)NHO– → RC(O)ONHC(O)R + NO–
(11)
0.28
0.20
0.15
0.10
0.05
0.00
A
B
pH 7.8
pH 8.5
0.21
0.14
0.07
0.00
pH 7.8
pH 7.0
pH 6.6
k
pH 5.9
k
pH 7.3
0.0
0.5
1.0
1.5
2.0
2.5
0.0
0.5
1.0
1.5
2.0
2.5
[gly-HX] (mM)
[benzo-HX] (mM)
Figure 7. (A) Dependence of kobs of the decay of CH3C(O)N=O on [glyꢀHX] at pH 5.9, 6.6, 7.0 and 7.8; (B)
Dependence of kobs of the decay of C5H6C(O)N=O on [benzoꢀHX] at pH 7.3, 7.8 and 8.5. Solutions contained
0.1 M KCl and 20 mM of either MES (pH 5.9), MOPS (pH 6.6 – 7.8) or TAPS (pH 8.5) buffers.
The kappꢀvalues calculated from the slope of the lines such as those presented in Figures 7 are
summarized in Table 6.
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Table 6. The apparent bimolecular rate constant (Mꢀ1sꢀ1) of RC(O)N=O reaction with HX at various pHs. In
parenthesis are the pKa(ꢀNHOH) of the reacting HX.
4
5
6
7
8
9
pH
acetoꢀHX
(9.0)
benzoꢀHX
(8.54)
glyꢀHX
(7.4 ꢀ 7.8)
65 ± 5
lysꢀHX
(7.9 – 8.0)
DFO
(8.7, 9.0, 10.1)
CH3C(O)N=O
5.9
6.2
6.6
7.0
7.3
7.8
7.9
8.5
69 ± 4
101 ± 8
169 ± 12
270 ± 18
414 ± 21
510 ± 18
733 ± 36
10
11
12
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19 ± 2
90 ± 7
49 ± 4
356 ± 20
39 ± 2
303 ± 17
21 ± 1
352 ± 21
C6H5C(O)N=O 7.3
36 ± 2
73 ± 5
268 ± 14
7.8
8.5
7.3
RC(O)N=O
derived from
suberoꢀHX
9.5 ± 1.0
All solutions contained 0.1 M KCl and 20 mM of either MES (pH 5.9 – 6.2), MOPS (pH 6.6 – 7.9) or TAPS
(pH 8.5) buffers.
Both glyꢀHX and DFO contain an amine group (Figure 1) where pKa(ꢀNH3 ) = 9.222 and 8.425,
+
+
respectively. In the case of glyꢀHX the highest studied pH 8.5 is lower than pKa(ꢀNH3 ) = 9.2 thus
eliminating the contribution of the amine to the reaction rate. A fit of kapp to a pKꢀcurve resulted in
pKa(ꢀNHOH)) = 7.8 ± 0.1 and k11 = 50 ± 5 and 860 ± 40 Mꢀ1sꢀ1 for the reaction of CH3C(O)N=O
with +H3NCH2C(O)NHOH and +H3NCH2C(O)NHO–, respectively. In the case of DFO the rate
increases about an order of magnitude as the pH increases from 7.3 to 8.5, thus excluding the
contribution of the amine moiety of DFO.
The decomposition of RC(O)N=O in the presence of excess acetoꢀHX yields N2O. In fact,
under the experimental conditions of Table 2 the yields of N2O were unaffected by the presence of
2 – 3 mM acetoꢀHX. Since the reaction of RC(O)N=O with 2 – 3 mM acetoꢀHX efficiently
competes with its hydrolysis, it is concluded that the reaction of RC(O)N=O with acetoꢀHX forms
HNO (reaction 11).
HXs avidly chelate transition metal ions through the hydroxamate moiety,22,31 and therefore
the question arises as to the effect of these chelates on the halfꢀlife of RC(O)N=O. The effect of
metal ions including CuII, NiII, ZnII and CoII (denoted MeII) on the decay of CH3C(O)N=O has been
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studied in the absence and presence of added acetoꢀHX at pH 7.3 (20 mM MOPS buffer). The most
effective transition metal was CuII, which catalyzes efficiently the hydrolysis of CH3C(O)N=O even
5
6
7
in the absence of acetoꢀHX. Figure 8A shows that the catalysis rate constant (kcat = (1.5
±
0.7) x 104
8
9
Mꢀ1sꢀ1 is practically the same whether electroꢀoxidation proceeds with 25 or 100 ꢀM acetoꢀHX
implying that the effect of residual unoxidized acetoꢀHX is negligible. This value decreased by
about 50% in the presence of 2 ꢀ 8 ꢀM CuII and 100 ꢀM acetoꢀHX. Since at [acetoꢀHX]/[CuII] >10
the predominant complex at pH 7.3 is CuX2,32 it is concluded that CuII alone is a better catalyst than
CuX2. In the case of all other transition metals studied, only the HXꢀmetal complex has a significant
effect on the decay rate of CH3C(O)N=O (Figure 8B).
10
11
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0.24
0.16
0.08
0.00
0.16
0.12
0.08
0.04
0.00
ZnII
B
A
NiII
k
CoII
0
20
40
60
80
0
4
8
12
16
[MeII] (
ꢀM)
[CuII] (
ꢀM)
Figure 8. (A) Dependence of kobs of the decay of CH3C(O)N=O formed by electroꢀoxidation of 25 ꢀM
acetoꢀHX (empty symbol) or 100 ꢀM acetoꢀHX (filled symbol) on [CuII] added following the electroꢀ
oxidation; (B) Dependence of kobs of the decay of CH3C(O)N=O formed by electroꢀoxidation of 100 ꢀM
acetoꢀHX on [MeII] and 150 ꢀM acetoꢀHX added following the electroꢀoxidation. All solutions contained 0.1
M KCl and 20 mM MOPS buffer at pH 7.3.
At high ratio of [MeII]/[HX], where the predominant species is MeX+, the catalysis rate was about
–
twice higher than that of MeX2 and/or MeX3 formed at high [HX]/[MeII] ratios (results not shown),
implying that MeX+ is a better catalyst. No rate increase of CH3C(O)N=O decomposition by 0.5
mM or 1 mM DFOꢀFeIII was observed, whereas DFO alone does act as a nucleophile (Table 6).
Thus, the metal binding to HX might affect the reaction rate either by catalyzing the hydrolysis of
RC(O)N=O or by blocking HX nucleophilic activity.
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Conclusions
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5
6
7
8
9
Acyl nitroso species are formed in aqueous solutions via electroꢀoxidation of HXs, and
undergo hydrolysis yielding mainly HNO and the respective carboxylic acid. Acyl nitroso species
are characterized by their typical absorption at 314 – 330 nm and their relatively long halfꢀlives
allow for the first time the study of the kinetics of their reactions with various nucleophiles where
the nucleophilic reactivity follows the order thiolate > hydroxamate > amine. Metal complexes of
acetoꢀHX catalyze the hydrolysis of CH3C(O)N=O at the efficacy order of CuII > ZnII > NiII > CoII
where only CuII catalyzes the hydrolysis also in the absence of the hydroxamate Finally, oxidation
of HXs can be used as a method for HNO generation where the halfꢀlife of this process is
determined by that of the respective acyl nitroso species.
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ACKNOWLEDGEMENT
This work has been supported by the Pazy Foundation Grant number 276/18.
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References
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Amides. J. Chem. Soc. 1964, 1318ꢀ1320.
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King, S. B.; Nagasawa, H. T. Chemical Approaches Toward Generation of Nitroxyl.
(HNO) at Biologically Relevant Temperatures from the RetroꢀDielsꢀAlder Reaction of Nꢀ
HydroxyureaꢀDerived Acyl Nitrosoꢀ9,10ꢀDimethylanthracene Cycloadducts. Tetrahedron Lett.
2000, 41, 4265ꢀ4269.
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Zeng, B.ꢀB.; Huang, J.; Wright, M. W.; King, S. B. Nitroxyl (HNO) Release from New
Functionalized NꢀHydroxyureaꢀDerived Acyl Nitrosoꢀ9,10ꢀDimethylanthracene Cycloadducts.
Bioorg. Med. Chem. Lett. 2004, 14, 5565ꢀ5568.
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Evans, A. S.; Cohen, A. D.; GurardꢀLevin, Z. A.; Kebede, N.; Celius, T. C.; Miceli, A. P.;
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(12) Shafirovich, V.; Lymar, S. V. Nitroxyl and Its Anion in Aqueous Solutions: Spin States,
Protic Equilibria, and Reactivities Toward Oxygen and Nitric Oxide. Proc. Natl. Acad. Sci. USA
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(13) Paolocci, N.; Jackson, M. I.; Lopez, B. E.; Miranda, K.; Tocchetti, C. G.; Wink, D. A.;
Hobbs, A. J.; Fukuto, J. M. The Pharmacology of Nitroxyl (HNO) and Its Therapeutic Potential:
Not Just the Janus Face of NO. Pharmacol. Therapeut. 2007, 113, 442ꢀ458.
(14) Huang, J. M.; Sommers, E. M.; KimꢀShapiro, D. B.; King, S. B. Horseradish Peroxidase
Catalyzed Nitric Oxide Formation From Hydroxyurea. J. Am. Chem. Soc. 2002, 124, 3473ꢀ3480.
(15) Samuni, Y.; Samuni, U.; Goldstein, S. The Mechanism Underlying Nitroxyl and Nitric
Oxide Formation From Hydroxamic Acids. Biochim. Biophys. Acta 2012, 1820, 1560ꢀ1566.
(16) O'Bannon, P. E.; Suelzle, D.; Dailey, W. P.; Schwarz, H. Spectroscopic Detection of Two
Neutral CHNO2 Isomers: Nitrocarbene and Nitrosoformaldehyde. J. Am. Chem. Soc. 1992, 114,
344ꢀ345.
(17) Cohen, A. D.; Zeng, B.ꢀB.; King, S. B.; Toscano, J. P. Direct Observation of an Acyl
Nitroso Species in Solution by TimeꢀResolved IR Spectrocopy. J. Am. Chem. Soc. 2003, 125, 1444ꢀ
1445.
(18) Maimon, E.; Samuni, A.; Goldstein, S. Nitrogen Dioxide Reaction with Nitroxide Radical
Derived from Hydroxamic Acids: The Intermediacy of Acyl Nitroso and Nitroxyl (HNO). J. Phy.
Chem. A 2018, 122, 3747ꢀ3753.
(19) Monzyk, B.; Crumbliss, A. L. Acid DissociationꢀConstants (Ka) and Their Temperature
Dependencies (ꢁHa, ꢁSa) for a Series of CarbonꢀSubstituted and NitrogenꢀSubstituted Hydroxamic
Acids in AqueousꢀSolution. J. Org. Chem. 1980, 45, 4670ꢀ4675.
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(20) Agrawal, Y. K. Dissociation Constants of Some Hydroxamic Acids. Z. Naturforsch. 1976,
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(21) Paniago, E. B.; Carvalho, S. The Formation Constants and Structures of Copper(II)ꢀ
Glycinehydroxamic Acid Complexes. Inorg. Chim. Acta 1984, 92, 253ꢀ257.
(22) Kurzak, B.; Kozlowski, H.; Farkas, E. Hydroxamic and Aminohydroxamic Acids and Their
Complexes with Metal Ions. Coord. Chem. Rev. 1992, 114, 169ꢀ200.
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(24) Caudle, M. T.; Caldwell, C. D.; Crumbliss, A. L. Dihydroxamic Acid Complexes of
Iron(III): Ligand pKa and Coordinated Water Hydrolysis Constants. Inorg. Chim. Acta 1995, 240,
519ꢀ525.
(25) Borgias, B.; Hugi, A. D.; Raymond, K. N. Isomerization and Solution Structures of
Desferrioxamine B Complexes of Aluminum(3+) and Gallium(3+). Inorg. Chem. 1989, 28, 3538ꢀ
3545.
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Acids. J. Electroanal. Chem. 1977, 80, 387ꢀ393.
(27) Arrigan, D. W. M.; Deasy, B.; Glennon, J. D.; Johnston, B.; Svehla, G. Incorporation of
Hydroxamic Acid Ligands into Nafion Film Electrodes. Analyst 1993, 118, 355ꢀ359.
(28) Taylor, R. J.; May, I. The Reduction of Actidine Ions by Hydroxamic Acids. Czech. J. Phys.
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(29) Smulik, R.; Dabski, D.; Zielonka, J.; Michalowski, B.; Adamus, J.; Marcinek, A.;
Kalyanaraman, B.; Sikora, A. Nitroxyl (HNO) Reacts with Molecular Oxygen and Forms
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(30) Perrin, D. D., Ed. Stability Constants of Metal Ion Complexes, Part B: Organic Ligands;
Pergamon Press: OxfordꢀNew YorkꢀTorontoꢀSydneyꢀParisꢀFrankfurt, 1979.
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(31) Codd, R. Traversing the Coordination Chemistry and Chemical Biology of Hydroxamic
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TOC Graphic
4
hydroxamic acid
acyl nitroso
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7
O
O
O
Nu:
HNO +
R
NHOH
R
Nu
R
N
O
H+
-
2e , 2H+
8
9
0.2
CH3C(O)N=O
0.2
CH3C(O)N=O
C6H5C(O)N=O
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pH 7.3
0.1
0.0
0.1
0.0
k
pH 6.6
300
360
(nm)
420
0
1
2
3
λ
[GSH] (mM)
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Figure 1. Structures of HXs and their pKa(-NHOH) given in paranthesis.
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Figure 2. Spectra of transient species formed upon electroꢀoxidation of HXs in aerated solutions containing
0.1 M KCl and 20 mM MOPS buffer at pH 7.3. Inset: first order kinetics of the decay monitored at λmax. (A)
500 ꢁM acetꢀHX, kd = 0.19 ± 0.01 minꢀ1; (B) 150 ꢁM benzoꢀHX, kd = 0.46 ± 0.02 minꢀ1; (C) 230 ꢁM SAHA,
kd = 0.14 ± 0.01 minꢀ1; (D) 95 ꢁM suberoꢀHX, kd = 0.21 ± 0.01 minꢀ1.
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Figure 3. Dependence of Amax of the transient species derived from electro-oxidation of HX on [HX]o in 0.1 M
KCl and 20 mM MOPS buffer at pH 7.3.
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Figure 4. Spectra of transient species formed by electro-oxidation of 15 µM AABA-HX (A) or gly-HX (B) in
0.1 M KCl and 20 mM MOPS buffer at pH 7.3. Inset: first-order decay kinetics of the transient species
monitored at 325 nm, kd = 1.56 ± 0.08 min-1 (A) and at 317 nm, kd = 0.96 ± 0.05 min-1(B).
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Figure 5. (A) Dependence of kobs of the decay of RC(O)N=O derived from benzoꢀHX (■), suberoꢀHX (△) and
SAHA (●) on [dimethylamine] at pH 9.4; (B) Dependence of kobs of the decay of CH3C(O)N=O on
[ethylamine]o at pH 9.5 and 10.1. All solutions contained 0.1 M KCl and 50 mM borate buffer.
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Figure 6. (A) Dependence of kobs of the decay of CH3C(O)N=O on [GSH] at pH 6.6 and 7.3; (B) Dependence
of kobs of the decay of C5H6C(O)N=O or CH3C(O)N=O on [cysteine] at pH 6.6. All solutions contained 0.1 M
KCl and 20 mM MOPS buffer.
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