Detailed mechanistic investigation into the S-nitrosation of cysteamine

Article abstract of DOI:10.1139/v2012-051

The nitrosation of cysteamine (H₂NCH₂CH ₂SH) to produce cysteamine-S-nitrosothiol (CANO) was studied in slightly acidic medium by using nitrous acid prepared in situ. The stoichiometry of the reaction was H₂NCH₂CH₂SH + HNO ₂ → H₂NCH₂CH₂SNO + H ₂O. On prolonged standing, the nitrosothiol decomposed quantitatively to yield the disulfide, cystamine: 2H₂NCH₂CH ₂SNO → H₂NCH₂CH₂S-SCH ₂CH₂NH₂ + 2NO. NO₂ and N ₂O₃ are not the primary nitrosating agents, since their precursor (NO) was not detected during the nitrosation process. The reaction is first order in nitrous acid, thus implicating it as the major nitrosating agent in mildly acidic pH conditions. Acid catalyzes nitrosation after nitrous acid has saturated, implicating the protonated nitrous acid species, the nitrosonium cation (NO⁺) as a contributing nitrosating species in highly acidic environments. The acid catalysis at constant nitrous acid concentrations suggests that the nitrosonium cation nitrosates at a much higher rate than nitrous acid. Bimolecular rate constants for the nitrosation of cysteamine by nitrous acid and by the nitrosonium cation were deduced to be 17.9 ± 1.5 (mol/L)^-1 s^-1 and 6.7 ×10⁴ (mol/L) ^-1 s^-1, respectively. Both Cu(I) and Cu(II) ions were effective catalysts for the formation and decomposition of the cysteamine nitrosothiol. Cu(II) ions could catalyze the nitrosation of cysteamine in neutral conditions, whereas Cu(I) could only catalyze in acidic conditions. Transnitrosation kinetics of CANO with glutathione showed the formation of cystamine and the mixed disulfide with no formation of oxidized glutathione (GSSG). The nitrosation reaction was satisfactorily simulated by a simple reaction scheme involving eight reactions.

Full text of DOI:10.1139/v2012-051

724
Detailed mechanistic investigation into the S-
nitrosation of cysteamine
Moshood K. Morakinyo, Itai Chipinda, Justin Hettick, Paul D. Siegel,
Jonathan Abramson, Robert Strongin, Bice S. Martincigh, and Reuben H. Simoyi
Abstract: The nitrosation of cysteamine (H2NCH2CH2SH) to produce cysteamine-S-nitrosothiol (CANO) was studied in
slightly acidic medium by using nitrous acid prepared in situ. The stoichiometry of the reaction was H2NCH2CH2SH +
HNO2 → H2NCH2CH2SNO + H2O. On prolonged standing, the nitrosothiol decomposed quantitatively to yield the disul-
fide, cystamine: 2H2NCH2CH2SNO → H2NCH2CH2S–SCH2CH2NH2 + 2NO. NO2 and N2O3 are not the primary nitrosating
agents, since their precursor (NO) was not detected during the nitrosation process. The reaction is first order in nitrous acid,
thus implicating it as the major nitrosating agent in mildly acidic pH conditions. Acid catalyzes nitrosation after nitrous acid
has saturated, implicating the protonated nitrous acid species, the nitrosonium cation (NO+) as a contributing nitrosating
species in highly acidic environments. The acid catalysis at constant nitrous acid concentrations suggests that the nitroso-
nium cation nitrosates at a much higher rate than nitrous acid. Bimolecular rate constants for the nitrosation of cysteamine
by nitrous acid and by the nitrosonium cation were deduced to be 17.9 1.5 (mol/L)–1 s–1 and 6.7 × 104 (mol/L)–1 s–1, re-
spectively. Both Cu(I) and Cu(II) ions were effective catalysts for the formation and decomposition of the cysteamine nitro-
sothiol. Cu(II) ions could catalyze the nitrosation of cysteamine in neutral conditions, whereas Cu(I) could only catalyze in
acidic conditions. Transnitrosation kinetics of CANO with glutathione showed the formation of cystamine and the mixed di-
sulfide with no formation of oxidized glutathione (GSSG). The nitrosation reaction was satisfactorily simulated by a simple
reaction scheme involving eight reactions.
Key words: nitric oxide, nitrosation, thiols, cysteamine, kinetics.
Résumé : Opérant en milieu légèrement acide et utilisant de l’acide nitreux préparé in situ, on a étudié la réaction de nitro-
sation de la cystéamine (H2NCH2CH2SH) conduisant à la cystéamine-S-nitrosothiol (CANO). La stoechiométrie de la réac-
tion est H2NCH2CH2SH + HNO2 → H2NCH2CH2SNO + H2O. Laissé longtemps au repos, le nitrosothiol se décompose
quantitativement avec formation du disulfure, cystamine : 2H2NCH2CH2SNO → H2NCH2CH2S–SCH2CH2NH2 + 2NO. Le
NO2 et le N2O3 ne sont pas les agents de nitrosation primaires puisque leur précurseur, le NO, n’a pas été détecté au cours
du processus de nitrosation. La réaction est du premier ordre en acide nitreux, ce qui implique qu’il est l’agent de nitrosa-
tion majeur dans des conditions de pH légèrement acides. L’acide catalyse la nitrosation après une saturation par l’acide ni-
treux et ceci implique que, dans des environnements d’acidité élevée, le cation nitrosonium (NO+) peut agir comme espèce
nitrosante. La catalyse acide, à des concentrations constantes d’acide nitreux, suggère que le cation nitrosonium provoque la
nitrosation avec une constante de vitesse beaucoup plus élevée que ce l’acide nitreux. On a déduit que les constantes de vi-
tesse bimoléculaires pour la nitrosation de la cystéamine, par l’acide nitreux et par le cation nitrosonium, sont respective-
ment de 17,9 1,5 (mol/L)–1 s–1 et 6,7 × 104 (mol/L)–1 s–1. Les ions Cu(I) ainsi que Cu(II) sont tous les deux efficaces
comme catalyseurs pour la formation et la décomposition du nitrosothiol de la cystéamine. Les ions Cu(II) catalysent la ni-
trosation de la cystéamine dans des conditions neutres alors que les ions Cu(I) ne la catalysent que dans des conditions aci-
des. La transnitrosation du CANO avec du glutathion donne lieu à la formation de cystamine et du disulfure mixte, sans
formation du produit doxydation du glutathion (GSSG). On peut simuler d’une façon satisfaisante la réaction de nitrosation
par un schéma réactionnel simple impliquant huit réactions.
Mots‐clés : oxyde nitrique, nitrosation, thiols, cystéamine, cinétique.
[Traduit par la Rédaction]
Received 2 February 2011. Accepted 16 April 2012. Published at www.nrcresearchpress.com/cjc on 22 August 2012.
M.K. Morakinyo and R. Strongin. Department of Chemistry, Portland State University, Portland, OR 97207-0751, USA.
I. Chipinda, J. Hettick, and P.D. Siegel. Allergy and Clinical Immunology Branch, Health Effects Laboratory Division, National Institute
of Occupational Safety and Health, Centers for Disease Control and Prevention, 1095 Willowdale Road, Morgantown, WV 26505, USA.
J. Abramson. Department of Physics, Portland State University, Portland, OR 97207-0751, USA.
B.S. Martincigh. School of Chemistry, University of KwaZulu-Natal Westville Campus, Private Bag X54001, Durban 4000, Republic of
South Africa.
R.H. Simoyi. Department of Chemistry, Portland State University, Portland, OR 97207-0751, USA; School of Chemistry, University of
KwaZulu-Natal Westville Campus, Private Bag X54001, Durban 4000, Republic of South Africa.
Corresponding author: Moshood K. Morakinyo (e-mail: mkm2002ng@yahoo.com).
Can. J. Chem. 90: 724–738 (2012)
doi:10.1139/V2012-051
Published by NRC Research Press

Morakinyo et al.
725
Cysteamine (2-aminoethanethiol) is a very important phys-
iologically active b-aminothiol. It is produced in vivo by deg-
radation of coenzyme A. Coenzyme A is a cofactor in many
enzymatic pathways, which include fatty acid oxidation,
heme synthesis, synthesis of the neurotransmitter acetylcho-
line, and amino acid catabolism.35 Cysteamine is medically
known as Cystagon and can be used as oral therapy for the
prevention of hypothyroidism and as a growth enhancer in
patients with nephropathic cystinosis.36 Cysteamine and its
disulfide (cystamine) can also be used in topical eye drops
to dissolve corneal cystine crystals.37,38 It has also been
shown to be an effective molecule against hepatotoxicity
caused by excessive production of nitric oxide.17
Introduction
Nitric oxide (NO) is a simple gaseous molecule with fasci-
nating biological functions.1 It is one of the most versatile bi-
oactive molecules ever identified. Unlike other free radicals,
it is relatively stable, but becomes unstable in biological sys-
tems due to its reaction with oxygen, superoxide anion (O2•ˉ),
and haem proteins.2 Nitrite (NO2 ) is the major end product
–
of its reaction with oxygen. Nitrite can also release NO on
acidification and by the reductase activity of the enzyme,
xanthine oxidoreductase, at pH conditions lower than 6.0.3–5
Nitric oxide has been the focus of many research groups in
recent years, probably because of the realization that it is
synthesized by mammalian cells and can act both as a phys-
iological messenger and as a cytotoxic agent.6 It is an essen-
tial physiological signaling molecule necessary for cell-to-cell
communication and it is also involved in the regulation of
gene expression and modification of sexual and aggressive
behavior.7,8 It is actively involved in smooth muscle relaxa-
tion due to its ability to activate guanosine 3′,5′-cyclic mono-
phosphate (cGMP) in smooth muscle tissue.9,10 The role of
nitric oxide in the physiological system at any given time sig-
nificantly depends on its concentration and its environment.
It is known to perform most of its beneficial roles, such as
vasorelaxation11,12 and neurotransmission,13,14 at low concen-
trations, on the order of 5–10 nmol/L.15 At this concentration
Although a series of studies have been performed on the
metabolism of this thiol and its metabolites with respect to
oxidation by reactive oxygen species,39–41 nothing has been
done on the elucidation of its mechanism of nitrosation by
NO and (or) NO-derived metabolites (N2O3 and NO+). Its ni-
trosation is well-known and has been reported in several
manuscripts,42,43 but the exact mechanistic details are still un-
known. We report herein on the detailed kinetics and mecha-
nism of the nitrosation of cysteamine. The nitrosating agent
is produced in situ, by premixing H+ and NO2 and utilizing
–
a double-mixing stopped-flow spectrophotometer. This manu-
script also reports on detailed product analysis as well as on
transnitrosation kinetics. The most prevalent thiol-based anti-
oxidant in the physiological environment is glutathione. It is
envisaged that, in any discussion of S-nitrosation, glutathione
would be a major player.
range it cannot effectively compete with superoxide dismu-
tase (SOD) for O2
•–
.
•–
NO also readily reacts with O2 to produce the highly
damaging and genotoxic oxidant (peroxynitrite, ONOO–):
NO þ O2 ! ONOOꢂ
ꢁꢂ
½R1ꢀ
Experimental
The bimolecular rate constant for rxn. [R1] is 6.7 ×
109 (mol/L)–1 s–1.15,16 Superoxide anion radical concentra-
tions in the physiological environment are in the order of
4–10 µmol/L.15 When NO concentrations are in the micromolar
range, rxn. [R1], which produces the deleterious peroxyni-
trite, becomes competitive with the deactivating reaction
that removes superoxide anion radical.
Materials
Cysteamine hydrochloride (CA, 2-aminoethanethiol hydro-
chloride) 98%, cystamine, cysteine (99+%), sodium perchlo-
rate (Acros), glutathione, sodium nitrite, sodium chloride,
perchloric acid (72%), ferrous sulfate (Fisher Scientific),
cuprous chloride, and cupric chloride (GSF chemicals) were
used as purchased. Stock solutions of CA and nitrite were
prepared just before use. Reagent solutions were prepared
with distilled and deionized water from a Barnstead Sybron
Corporation water purification unit. Inductively coupled
plasma mass spectrometry (ICP–MS) was used to show that our
aqueous reaction medium did not contain enough transition-
metal ions to affect the overall reaction kinetics and mecha-
nism The highest metal ion concentration was cadmium at
1.5 ppb followed by lead at 0.43 ppb. Copper was virtually
undetectable.
Excessive production of nitric oxide has thus been linked
to so many cytotoxic activities such as septic shock and liver
injury.17,18 It is necessary to have processes that can reduce
NO concentrations to beneficial levels in the physiological
environment. Such processes would include its reaction with
thiols to form S-nitrosothiols (RSNOs). For example, RSNOs
have been shown to improve end-organ recovery in models
of ischemia–reperfusion injury in the heart and liver.19,20 Bio-
logical nitrosothiol formation is a post-translational modifica-
tion of biological thiols that is subsequently linked to all
those functions that are attributed to nitric oxide.21 Possible
roles of S-nitrosation include the regulation of apoptosis
through glyceraldehyde-3-phosphate dehydrogenase modifi-
cation22 as well as involvement in the pathogenesis of Parkin-
son’s disease.23 There are several studies on nitrosothiol
formation by Williams,24–33 but these studies dwell on forma-
tions and characterizations and not so much on the mecha-
nisms and kinetics. Despite the potential benefits of the
relevance of the mechanistic basis of nitrosations (both S-
and N-nitrosations),34 very few studies have been directed at
elucidating these mechanisms.
Methods
All experiments were carried out at 25.0 0.1 °C and at a
constant ionic strength of 1.0 mol/L (NaClO4). The progress
of the nitrosation reaction was monitored spectrophotometri-
cally by following the absorbance of S-nitrosocysteamine
(CANO) at its experimentally determined absorption peak of
545 nm, where an absorptivity coefficient of 16.0
0.1 (mol/L)–1 cm–1 was evaluated. Although CANO has a
second absorption peak at 333 nm, with a much higher ab-
sorptivity coefficient of 536.0 (mol/L)–1 cm–1, its maximum
absorption at 545 nm was used owing to the lack of interference
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726
Can. J. Chem. Vol. 90, 2012
Fig. 1. UV–vis spectral scan of reactants 0.01 mol/L cysteamine hy-
TOF-MS
–
drochloride (CA, a), 0.001 mol/L NaNO2 (b), 0.001 mol/L HNO2
Mass spectra were acquired on a Micromass QTOF-II
(Waters Corporation, Millford, MA) quadrupole time-of-flight
mass spectrometer. Analytes were dissolved in 50:50 aceto-
nitrile / water – 1% formic acid mixture, and analyte ions
were generated by positive-mode electrospray ionization
(ESI) at a capillary voltage of 2.8 kV and a flow rate of
5 µL/min. The source block was maintained at 80 °C and
the nitrogen desolvation gas was maintained at 150 °C and
at a flow rate of 400 L/h. Data were visualized and ana-
lyzed with the Micromass MassLynx 4.0 software suite for
Windows XP (Water Corporation, Millford, Massachusetts).
(c), and product 0.001 mol/L S-nitrosocysteamine (CANO, d).
EPR Measurements
A Bruker e-scan EPR spectrometer (X-band) was used for
the detection of radicals and to confirm the release of NO
upon the decomposition of CANO. Detection of NO was
through the use of the NO-specific spin trap, aci-nitroethane.44
The following settings were used for a typical run: microwave
power, 19.91 mW; modulation amplitude, 1.41 G; receiver
gain, 448; time constant, 10.24 ms; sweep width, 100 G; and
frequency, 9.78 GHz. All measurements were taken at room
temperature.
from the absorption peaks of nitrogen species (NO2 , NO+,
–
and HNO2) at this absorption wavelength (see Fig. 1).
Product identification and verification was achieved by
complementary techniques: UV–vis spectrophotometry,
HPLC, quadrupole time-of-flight mass spectrometry (TOF-
MS), and electron paramagnetic resonance (EPR) spectrome-
try. Kinetics measurements for fast reactions were performed
on a Hi-Tech Scientific double-mixing SF61-DX2 stopped-
flow spectrophotometer. For slower reactions, a PerkinElmer
Lambda 25 UV–vis spectrophotometer was used.
Stoichiometric determinations were performed by mixing
excess perchloric acid and nitrite with a fixed amount of CA
such that CA was the limiting reagent. The reaction was al-
lowed to proceed to completion while rapidly scanning the
reacting solution between 200 and 700 nm. The concentra-
tion of CANO was deduced from its maximum absorption at
545 nm.
Results
The major reaction studied was a simple nitrosation of cys-
teamine through in situ production of nitrous acid as the ni-
trosating agent. The yield of CANO varied, based on the
order of mixing reagents. The highest yield, which gave
quantitative formation of CANO, was obtained by premixing
nitrite with acid and incubating the solution for approxi-
mately 1 s before subsequently reacting this nitrous acid
with cysteamine in a third reagent stream of the double-mixing
stopped-flow spectrophotometer. The value of the absorptiv-
ity coefficient of CANO at 545 nm adopted for this study
of 16.0 mol/L–1 cm–1 was derived from this mixing order.
Premixing CA with acid before the addition of nitrite gave
a lower yield of CANO. The available protons are reduced
by the protonation of CA.
HPLC technique
The HPLC system utilized a Shimadzu Prominence (Co-
lumbia, Maryland) SPD-M10A VP diode array detector, dual
LC 600 pumps, LPT-6B LC–PC interface controller, SIL-10AD
auto injector, and Class-VP chromatography data system
software. All samples were loaded on a reversed phase Dis-
covery 5 µm C18 column (Supelco, Bellefonte, Pennsylva-
nia). They were run isocractically at 0.15% trifluoroacetic
acid (TFA) / H2O to 10% methanol / H2O, and CANO was
detected at 545 nm (for selective scans) with maximum spec-
tral scan mode employed to detect all eluents. A flow rate of
1 mL/min was maintained. All solutions for HPLC analysis
were made with Milli-Q Millipore purified water and fil-
tered with Whatman Polypropylene 0.45 µm pore-size filter
devices before injection (10 µL) into the column using the
auto injector. To curb the interaction of the protonated
amines on the analytes with the silanol groups on the sta-
tionary phase (which was causing tailing of peaks), the so-
dium salt of 1-octanesulfonic acid (0.005 mol/L) was
incorporated into the aqueous mobile phase as an HPLC
ion-pairing agent. This was sufficient to neutralize the pro-
tonated amines and produce good resolution while eliminat-
ing peak tailing.
½R2ꢀ
H2NCH2CH2SH þ Hþ $ ½H3NCH2CH2SHꢀþ
This subsequently reduces the amount of the immediately
available nitrosating agent, HNO2.
Stoichiometry and product analysis
The stoichiometry, in conditions of excess acid and nitrite
over CA, was strictly 1:1.
½R3ꢀ
H2NCH2CH2SH þ HNO2 !
H2NCH2CH2SꢂNO þ H2O
Stoichiometry rxn. [R3] is attained within a minute of mix-
ing the reagents at high acid concentrations, and a little lon-
ger in lower acid concentrations. High acid concentrations
rapidly form the product (CANO), which also rapidly decom-
poses. A resolvable mass spectrum is taken at low acid con-
centrations, and this generally will show both the reagent
(CA) and the single product (CANO). This mass spectrum is
shown in Fig. 2A. There are no other peaks that can be dis-
cerned from the reaction mixture, apart from those belonging
to the reagent and the product. Incubation of the product of
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Morakinyo et al.
727
Fig. 2. Mass spectrometry analysis of the product of nitrosation of cysteamine hydrochloride (CA) showing (A) formation of S-nitrosocysteamine
(CANO), m/z = 107.9999; (B) product of decomposition of CANO after 24 h. Cystamine, a disulfide of CA was produced with a strong
peak, m/z = 153.0608.
nitrosation overnight shows a final conversion of the CANO
into the disulfide, cystamine (H2NCH2CH2S–SCH2CH2NH2,
Fig. 2B).
cess concentrations of the trapping reagent, heights of the
spectral peaks can be used, after a standardization, as a quan-
titative measure of the available NO. NO has a very poor sol-
ubility in water, and so the technique is limited only to low
NO concentrations.
½R4ꢀ
2H2NCH2CH2SꢂNO !
H2NCH2CH2SꢂSCH2CH2NH2 þ 2NO
Nitric oxide can be trapped and observed by EPR techni-
ques through the use of a nitroethane trap.44 The aci-anion
of this trap, formed in basic environments, generates a long-
lived spin adduct with NO, which is detectable and quantifi-
able by its distinct EPR spectrum pattern and g value.
Reaction kinetics
The reaction has a very simple dependence on cysteamine
concentrations (see Figs. 4A and 4B). There is a linear first-
order dependence on the rate of nitrosation with CA concen-
trations with an intercept kinetically indistinguishable from
zero. The reaction kinetics are more reproducible in the for-
mat utilized in Fig. 4 in which concentrations of acid and ni-
trite are equal and in overwhelming excess over CA
concentrations. The effect of nitrite is also simple and first-
order (see Figs. 5) when nitrite concentrations do not exceed
the initial acid concentrations. Effectively, under these condi-
tions, all N(III) species in the reaction environment are in the
form of protonated HNO2. Acid effects are much more com-
plex and are dependent on the ratio of initial acid to nitrite
concentrations. The nitrosation response to acid differs, de-
pending on whether initial proton (H3O+) concentrations ex-
ceed initial nitrite concentrations. The most important
parameter affecting the rate of nitrosation is the concentration
of HNO2. A set of kinetics traces that encompass a proton
concentration variation, which ranges from concentrations
ꢂ
ꢁ
½R5ꢀ
CH3CH ¼ NO2 þ NO þ OHꢂ !
½CH3CðNOÞðNO2Þꢀꢁ2ꢂ þ H2O
This adduct ([CH3C(NO)(NO2)]•2–) is EPR-active. No
peaks were obtained in the EPR spectrum when the nitro-
ethane trap was utilized during the formation of CANO, indi-
cating that NO is not released during the S-nitrosation and
that it is not a major nitrosation agent in this mechanism.
Upon use of the trap, however, during the decomposition of
CANO, a strong EPR spectrum indicating the presence of
NO was obtained, proving stoichiometry [R4]. Figure 3
shows a series of EPR spectra that proves the presence of
NO in the decomposition of CANO. The first three spectra
–
show that, on their own, nitroethane, CA, and NO2 cannot
generate any active peaks that can be attributed to NO. In ex-
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728
Can. J. Chem. Vol. 90, 2012
Fig. 3. Electron paramagnetic resonance (EPR) spectra of the NO radical generated during the decomposition of S-nitrosocysteamine (CANO)
–
using 0.5 mol/L nitroethane (NE) in 1.0 mol/L NaOH solution as the spin trap. (A) 0.5 mol/L NE, (B) 0.01 mol/L NO2 , (C) 0.04 mol/L
cysteamine hydrochloride (CA), (D) [CANO] = 0.01 mol/L, (E) [CANO] = 0.02 mol/L, (F) [CANO] = 0.03 mol/L, and (G) [CANO] =
0.04 mol/L.
that are lower than nitrite to those that exceed, was performed.
A simple plot of added acid to the rate of nitrosation shown
in Fig. 6A clearly displays a discontinuity at the point
where acid exceeds nitrite concentrations. By utilizing the
dissociation constant of HNO2, the final concentrations of
Effect of copper
Copper has been implicated in many reactions involving
organosulfur compounds.45–47 Cu(I) is known to be much
more active than Cu(II). It is easier to assay for Cu(II) and
the addition of Cu(II) to organosulfur reactions involves its
initial conversion to the more active Cu(I). Figure 7A shows
that addition of Cu(II) rapidly increases the nitrosation reac-
tion. Cu(II) ions, however, also catalyze the decomposition of
CANO. A plot of the rate of nitrosation versus Cu(II) concen-
trations shows the sigmoidal rate increase followed by an
abrupt saturation, which is typical of catalytic mechanisms
(figure not shown). No nitrosation occurs in the presence of ni-
trite without acid (which is necessary to generate the HNO2).
Figure 7B, however, shows that, even in the absence of acid,
Cu(II) ions can still effect the nitrosation of CA. Cu(I) ions,
however, could not effect nitrosation in the absence of acid.
Figure 7C shows the catalytic effect of Cu(II) ions on the
decomposition of CANO. CANO was prepared in situ from
two feed streams of the double-mixing stopped flow ensem-
ble. Cu2+ ions derived from cupric chloride were next
added from a third feed stream after complete production
of CANO. All feed streams were buffered at pH 7.4. The
two control experiments (traces a and b) show the slow de-
composition expected from CANO solutions at pH 7.4.
Trace b was treated with EDTA to sequester all metal ions
in the reaction solution. This trace is identical to the one
without EDTA (trace a), confirming the absence of metal
ions in reagent solution, which can affect the decomposition
kinetics. Addition of micromolar quantities of Cu(II) ions
effects a big acceleration in the rate of decomposition.
Higher Cu(II) concentrations give a strong autocatalytic de-
cay of the nitrosothiol (trace f in Fig. 7C). The catalytic ef-
fect of Cu(II) ions on the decomposition can also be
H3O+, NO2 , and HNO2 can be calculated for each set of
–
initial concentrations. These data are shown in Table 1 for
the input concentrations used for the set of kinetics experi-
ments. Close examination of the data in Table 1 shows that
the kinetics of the reaction are vastly different depending on
whether initial acid concentrations exceed initial nitrite con-
centrations. When nitrite is in excess (the first four read-
ings), there is a sharp nonlinear increase in the rate of
nitrosation with nitrous acid concentrations (Fig. 6B), and a
saturation in rate with respect to the actual proton concen-
trations despite the recalculation of the concentrations
undertaken in Table 1. Figure 6B would strongly indicate
second-order kinetics in nitrous acid, and a plot of nitrosa-
tion rate versus nitrous acid to the second power is linear
(plot not shown). Table 1 shows that, in the series of data
plotted for Fig. 6B, nitrous acid concentrations and excess
proton concentrations increase while nitrite concentrations
decrease, and the final observed rate of nitrosation is de-
rived from these three species concentrations and not iso-
lated to nitrous acid. At acid concentrations that exceed
initial nitrite concentrations, there is a linear dependence in
rate on acid concentrations (Fig. 6C). Note that a saturation
would have been achieved in nitrous acid concentrations as
would be expected in this format in which the nitrous acid
saturates at 50 mmol/L. The increase in nitrosation rate with
an increase in acid after saturation of HNO2 indicates that
other nitrosants apart from HNO2 exist in the reaction envi-
ronment.
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Morakinyo et al.
729
Fig. 4. (A) Absorbance traces showing the effect of varying cystea-
Fig. 5. Initial rate plot from the data derived from varying nitrite ion
mine hydrochloride (CA) concentrations. The reaction displays first-
concentrations in excess acid over nitrite. [CA]0 = 0.10 mol/L;
order kinetics in CA. [NO2 ]0 = 0.10 mol/L; [H+]0 = 0.10 mol/L;
[H+]0 = 0.10 mol/L; and [NO2 ]0 was varied from 0.005 to
–
–
and (a) [CA]0 = 0.005 mol/L; (b) [CA]0 = 0.006 mol/L; (c) [CA]0 =
0.007 mol/L; (d) [CA]0 = 0.008 mol/L; (e) [CA]0 = 0.009 mol/L;
and (f) [CA]0 = 0.010 mol/L. (B) Initial rate plot of the data in
Fig. 4A. The plot shows the strong first-order dependence of the
rate of formation of S-nitrosocysteamine (CANO) on CA, with an
intercept kinetically indistinguishable from zero.
0.010 mol/L. The plot shows strong first-order dependence on initial
rate of formation of S-nitrosocysteamine (CANO) on nitrite. CA,
cysteamine hydrochloride.
absorb around the same wavelength and with similar absorp-
tivity coefficients. Figure 8A shows a series of superimposed
spectra of four well-known nitrosothiols, which show nearly
identical spectra. Figure 8B shows the spectrophotometric
characterization of possible transnitrosation in mixtures of
CANO and GSH. GSNO has a higher absorptivity coefficient
than CANO at the same wavelength,48 and this explains the
initial increase in absorbance on addition of GSH to CANO.
LC–MS data have shown that the reaction mixture can transi-
ently (before attainment of the equilibrium mixture) contain
any of at least six components at varying times and concen-
trations: CANO, S-nitrosoglutathione (GSNO), GSH, CA–
CA (cystamine), GSSG (oxidized glutathione), and CA–GS
(mixed disulfide). The rate of transnitrosation will be difficult
to evaluate owing to the presence of the mixed disulfide. Fig-
ure 8B shows that, relatively, CANO is more stable than
GSNO because of the rapid rate of consumption of GSNO
upon its formation as compared with the much slower
CANO decomposition with or without EDTA. Figure 9
shows a GC–MS spectrum of mixtures of CANO and GSH.
This involved a 3:5 ratio of CANO to GSH, which is also
trace f in Fig. 8B. The remarkable aspect of this spectrum is
the absence of the GSSG, whose precursor is GSNO. Thus,
GSNO, when formed, rapidly decomposes by reacting with
CA to form the mixed disulfide and releasing NO. The cyst-
amine formed and observed in the spectrum is derived from
the normal decomposition of CANO (see rxn. [R10], vide in-
fra).
evaluated through the rate of production of nitric oxide,
which is known to be the major product, together with the
disulfide, of nitrosothiol decompositions. Figure 7D shows
EPR spectra using the nitroethane trap, of the enhanced
production of NO with addition of Cu(II) ions. All three
spectra were taken after the same time lapse from CANO
incubation, and thus the peak heights of the NO-induced
spectra can be correlated with the extent of reaction.
Transnitrosation
The most abundant thiol in the physiological environment
is glutathione (GSH). A pertinent reaction to study would be
the interaction of any nitrosothiol formed in the physiological
environment with glutathione. There is a crucial question
about the lifetimes of the nitrosothiols formed. If thiols are
carriers of NO from the point of production to the point of
its usage, the nitrosothiols formed should be stable enough
and have a long enough lifetime to enable them to perform
this activity. Spectrophotometric study of transnitrosation is
rendered complex by the fact that many nitrosothiols seem to
Mechanism
The overall rate of nitrosation is dependent on the rate of
formation and accumulation of the nitrosating agents. A num-
ber of possibilities have been suggested, and these include
HNO2, NO, NO2, N2O3, and the nitrosonium cation, +NO.
Some nitrosants cannot exist under certain conditions; for
Published by NRC Research Press

730
Can. J. Chem. Vol. 90, 2012
Fig. 6. (A) Plot of the raw kinetics data for the effect of added acid
–
on the rate of nitrosation.[CA]0 = 0.05 mol/L; [NO2 ]0 = 0.05 mol/L;
and [H+]0 was varied between 0.01 and 0.08 mol/L. For the first
four data points, [NO2 ]0 > [H+]0. The linear potion of the plot
–
represents excess proton concentrations over nitrite. Acid concen-
trations plotted here are those experimentally added and not cal-
culated. (B) Plot of initial rate versus nitrous acid concentrations
of the experimental data in Fig. 6A and calculated in Table 1.
This plot involves the region at which initial nitrite concentrations
exceed acid concentrations. (C) Plot showing the linear depen-
dence of the rate of formation of S-nitrosocysteamine (CANO) on
[H3O+] in high acid concentrations, where initial acid concentra-
tions exceed nitrite concentrations. Nitrous acid concentrations are
nearly constant in this range (see Table 1).
example, the +NO cannot exist at high pH49 conditions and
NO2 cannot be produced in strictly anaerobic conditions. Fig-
ure 10 shows two nitrosation reactions at equivalent condi-
tions except that one was run in solutions that had been
degassed with argon and immediately capped. Aerobic envi-
ronments can form NO2 from NO and subsequently N2O3
will also be formed.50
½R6ꢀ
NO þ 1=2O2 ! NO2
½R7ꢀ
NO þ NO2 ! N2O3
Both NO2 and N2O3 are potent nitrosating agents.29 There
is a slightly higher rate of nitrosation in aerobic environments
when compared with those in anaerobic environments. This is
a reflection of the low dissolved oxygen concentrations avail-
able in aqueous solutions at approximately 2 × 10–4 mol/L.
Thus, although formation for these nitrosants is favored,
their viability is hampered by the low concentrations of
oxygen as well as NO, which is only produced during the
decomposition of the nitrosothiol. NO was not detected dur-
ing the formation of CANO. In the aqueous environment
utilized for these experiments, reaction of aqueous N2O3 to
nitrous acid would be overwhelmingly dominant. This is a
rapid equilibrium reaction that favors the weak acid, HNO2.
½R8ꢀ
N2O3ðaqÞ þ H2OðlÞ $ 2HNO2ðaqÞ; KH
Early studies by Bunton and Steadman51 had erroneously es-
timated a value for KH = 0.20 (mol/L)–1, which would indi-
–1
cate appreciable amounts of N2O3(aq) over nitrous acid. A
more accurate evaluation by Markovits et al.52 gave a more real-
istic value of KH = 3.03 × 10–3 (SD 0.23 × 10–3) (mol/L)–1.
–1
According to Turney,53 appreciable amounts of N2O3 can
only exist in highly acidic environments, in the region of
5 mol/L perchloric acid. Such conditions are not applicable
for the physiological environment. Most nitrosation studies
in aerobic NO environments neglect the contribution to ni-
trosation by NO2 and only concentrate on NO and N2O3.
NO2 concentration will always be negligibly low because
of rxn. [R6], which is two orders of magnitude greater
than the rate of nitrosation by NO2 (1.1 × 109 (mol/L)–1 s–1
vs 2.0 × 107 (mol/L)–1 s–1, respectively),54,55 is hampered
by the low concentrations of oxygen available. In the pres-
ence of NO, it is more unstable with respect to N2O3.
All nitrosation kinetics show first-order dependence in
both thiol and nitrite with a complex dependence on acid.
Data in Figs. 4 and 5 show this strong first-order dependence
on CA and on nitrite. In conditions in which acid concentra-
tions are less than initial nitrite concentrations, the effect of
nitrite is more difficult to interpret, since a change in its con-
centrations concomitantly induces a change in both acid and
HNO2 concentrations, and thus its sole effect cannot be iso-
lated. First-order kinetics in nitrite strongly suggest that the
major nitrosant at the beginning of the reaction is HNO2,
since there is a direct correlation between nitrite concentra-
tions and HNO2, especially when there is consistently excess
Published by NRC Research Press

Morakinyo et al.
731
Table 1. Input species concentrations and final reagent concentrations for acid dependence ex-
periments (units of mol/L).
[H+]added
0.01
0.02
0.03
0.04
0.05
0.06
0.07
0.08
[NO2 ]added
[HNO2]initial
0.01
0.02
0.03
0.04
0.05
0.05
0.05
0.05
[NO2 ]final
[HNO2]final
9.86×10–3
1.97×10–2
2.92×10–2
3.82×10–2
4.50×10–2
4.78×10–2
4.89×10–2
4.91×10–2
[H3O+]final
1.38×10–4
3.63×10–4
7.90×10–4
1.82×10–3
5.03×10–3
1.22×10–2
2.11×10–2
3.09×10–2
–
–
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
4.01×10–2
3.04×10–2
2.08×10–2
1.18×10–2
5.03×10–3
2.22×10–3
1.07×10–3
8.93×10–4
Fig. 7. (A) Absorbance traces showing the effect of varying Cu2+ concentrations on the rate of formation of CANO. There is a progressive increase
–
in the rate of S-nitrosocysteamine (CANO) formation with an increase in Cu2+ concentration. [CA]0 = [NO2 ]0 = [H+]0 = 0.05 mol/L; and (a)
[Cu2+]0 = 0.00 mol/L, (b) [Cu2+]0 = 5 µmol/L, (c) [Cu2+]0 = 15 µmol/L, (d) [Cu2+]0 = 100 µmol/L, (e) [Cu2+]0 = 1 mmol/L, and (f) [Cu2+]0 =
2.5 mmol/L. CA, cysteamine hydrochloride. (B) Absorbance traces showing the formation of CANO via a reaction of CA, nitrite, and copper(II)
without acid. The amount of CANO formed increases with an increase in Cu2+ concentration. No formation of CANO was observed with Cu+. The
result shows the catalytic effect of Cu2+ on the rate of formation of S-nitrosothiol (RSNO). [CA]0 = [NO2 ]0 = [H+]0 = 0.05 mol/L; and (a)
–
[Cu2+] = 0.0025 mol/L, (b) [Cu2+] = 0.0030 mol/L, (c) [Cu2+] = 0.0035 mol/L, (d) [Cu2+] = 0.0040 mol/L, (e) [Cu2+] = 0.0045 mol/L, and (f)
[Cu2+] = 0.0050 mol/L. (C) Effect of Cu2+ on the stability of CANO in a pH 7.4 phosphate buffer. This reaction involves the reduction of Cu2+ to
Cu+, which causes the decomposition of CANO. [CANO]0 = 0.028 mol/L; (a) [Cu2+]0 = 0.00 mol/L (pure CANO), (b) [Cu2+]0 = 0.00 mol/L
(pure CANO in 0.00001 mol/L EDTA), (c) [Cu2+]0 = 5 µmol/L, (d) [Cu2+]0 = 10 µmol/L, (e) [Cu2+]0 = 20 µmol/L, and (f) [Cu2+]0 = 30 µmol/L.
(D) Electron paramagnetic resonance (EPR) spectra of the NO radical showing the catalytic effect of Cu2+ on the rate of decomposition of CANO.
The intensity of the spectra increases with an increase in Cu2+ concentration. [CANO]0 = 0.01 mol/L; (a) [Cu2+]0 = 0.00, (b) [Cu2+]0 = 1.0 ×
10–4 mol/L, and (c) [Cu2+]0 = 2.0 × 10–4 mol/L.
Published by NRC Research Press

732
Can. J. Chem. Vol. 90, 2012
Fig. 8. (A) Superimposed spectra of four well-known nitrosothiols
(a) CysNO, (b) S-nitrosocysteamine (CANO), (c) S-nitrosoglu-
tathione (GSNO), and (d) nitrosothiol of penicillamine (SNAP).
Three of them nearly have the same ɛ at 545 nm. The observed se-
paration in the absorbances of these three at 545 nm is due to stag-
gering of the time lag before acquiring the spectrum. (B) The effect
of glutathione (GSH) on the stability of CANO in a pH 7.4 phos-
phate buffer. All experimental traces have [CA]0 = 0.03 mol/L,
observed in rate is not solely attributable to changes in
HNO2 concentrations. When acid is in excess over nitrite,
the amount of HNO2 saturates, but as can be seen in
Fig. 6C, the rate keeps increasing linearly with an increase
in acid, indicating the presence of multiple nitrosating
agents apart from HNO2. The simplest rate law that can ex-
plain this nitrosation rate involves +NO and HNO2. +NO
can be formed from the protonation of HNO2:
[NO2 ]0 = 0.03 mol/L, and [H+]0 = 0.05 mol/L. (a) [GSH] = 0, (b)
–
½R9ꢀ
HNO2 þ H3Oþ $ HðOHÞN¼Oþ þ H2O; Kb
[GSH] = 0, (c) [GSH] = 0.01 mol/L, (d) [GSH] = 0.02 mol/L, (e)
[GSH] = 0.03 mol/L, and (f) [GSH] = 0.05 mol/L. Traces a and
b are controls. Trace a has no EDTA. Trace b is the same as
trace a with 10 µmol/L of EDTA. There is no significant differ-
ence between traces a and b, indicating that the water used for
preparing reagent solutions did not contain enough trace metal
ions to affect the kinetics. Ten microlitres of EDTA was added to
(H(OH)N=O+ is a hydrated nitrosonium cation, +NO + H2O).
The overall rate of reaction, using these two nitrosants
gives
ꢀ
ꢁ
½Hþꢀ½NðIIIÞꢀT
½1ꢀ
Rate ¼
½CAꢀðk1 þ k2Kb½HþꢀÞ
½Hþꢀ þ Ka
where [N(III)]T is the total concentrations of nitrous acid and
nitrite as calculated in Table 1 (columns 4 and 5) and [H+] is
the calculated acid concentration in column 6 of Table 1. Ka is
the acid dissociation constant of nitrous acid, Kb is the equili-
brium for the protonation of nitrous acid (rxn. [R9]), k1 is
the bimolecular rate constant for the nitrosation of CA by
HNO2, and k2 is the bimolecular rate constant for the nitro-
+
sation by NO. Ka was taken as 5.62 × 10–4 (mol/L)–1 and,
thus, at high acid concentrations, a plot of Rate/[N(III)]T
[CA] vs [H+] should give a straight line with slope k2Kb
and intercept k1. This type of plot only applies to conditions
of excess acid over nitrite (Fig. 6C). One can derive values
for k1 and product k2Kb from these series of plots. A value
of k1 = 17.9
1.5 (mol/L)–1 s–1 was derived from a series
of several complementary experiments. k2 was deduced as
6.70 × 104 (mol/L)–1 s–1 after assuming a Kb value of
3.16 ×10–2 (mol/L)–1.
Decomposition of CANO
RSNOs differ greatly in their stabilities, and most of them
have never been isolated in solid form. Homolysis of the S–N
bond occurs both thermally and photochemically, but, gener-
ally, these processes are very slow at room temperature and
in the absence of radiation of the appropriate wavelength. In
vivo decomposition of RSNOs is likely to proceed through
other pathways apart from these two. The presence of trace
amounts of Cu(II) or Cu(I) can efficiently catalyze RSNOs
to disulfide and nitric oxide (see Figs. 7A–7D).
½R10ꢀ
2RSNO ! RSSR þ 2NO
The characterization of the S–N bond is crucial in deter-
mining the stability of RSNOs, and, unfortunately, very little
work has been performed in this area. The little crystallo-
graphic data available on isolable RSNOs indicate that the
S–N bond is weak, sterically unhindered, elongated at be-
tween 0.17 and 0.189 nm,56 and exists as either the cis (syn)
or trans (anti) conformers. These conformers are nearly isoe-
nergetic but separated by a relatively high activation barrier
of ~11 kcal mol–1 (1 cal = 4.184 J). CANO is a primary S-
nitrosothiol, with primary and secondary nitrosothiols known
to be kinetically unstable to decomposition when compared
with tertiary nitrosothiols.57,58 S-Nitroso-N-acetylpenicillamine
is the most well-known stable tertiary nitrosothiol.59 The
mechanism of decomposition of nitrosothiols was believed to
acid over nitrite. After performing calculations for the relevant
reactive species in the reaction medium (Table 1), the plot
of the initial rate of nitrosation versus nitrous acid concen-
trations shows a parabolic dependence, which erroneously
suggests second-order kinetics in HNO2 (Fig. 6B) for the
region in which initial nitrite concentrations exceed acid.
As mentioned in the Results section, the data in Table 1
show that HNO2 is not changing in isolation. Both nitrite
and acid are also concomitantly changing, and so the effect
Published by NRC Research Press

Morakinyo et al.
733
Fig. 9. A GC–MS spectrum of a 3:5 ratio of S-nitrosocysteamine (CANO) to glutathione (GSH); trace f in Fig. 8B. Final products were
predominantly the mixed disulfide and cysteamine with no evidence of oxidized glutathione (GSSG). CA, cysteamine hydrochloride.
proceed through the homolytic cleavage of this weak S–N
bond. Computed S–N bond dissociation energies for nitroso-
thiols show very little variation, however, between primary,
secondary, and tertiary nitrosothiols, resulting in very similar
homolysis rates of reaction at elevated temperatures for all ni-
trosothiols. The computed activation parameters for thermal
homolysis to occur are prohibitively high for this to be a dom-
inant pathway under normal biological conditions. For exam-
ple, assuming an approximate S–N bond energy of 31 kcal,58
a rough calculation will indicate that if decomposition was to
proceed solely through homolysis, the half-life of a typical ni-
trosothiol would be in the region of years instead of the mi-
nutes to hours that are observed.60 This analysis does not
support the work reported by Roy et al.61 in which they estab-
lished homolysis as the major decomposition pathway. They
also managed to trap the thiyl radical, RS·. Close examination
of their reaction conditions show high trap concentrations
(0.2 mol/L DMPO) and nitrosothiol (0.1 mol/L). Such condi-
tions would catch even very low concentrations of thiyl radical,
but would not necessarily indicate that this was the dominant
pathway and whether it was derived solely from homolytic
cleavage of the S–N bond. Nitrosothiol decomposition rates in
the absence of metal ions are known to be mostly zero order57
and this is inconsistent with an S–N bond homolysis-driven
decomposition,62 which would have been expected to be first
order. In the presence of excess thiol, however, decomposition
of nitrosothiols proceeds via a first-order decay process.59 The
absence of first-order decay kinetics in the decomposition of
CANO also rules out homolysis as a dominant pathway. For a
series of nitrosothiols, the homolysis-dependent decomposition
rates are not dependent on the bulkiness of the substituent
groups attached to the nitrosothiols, which seems to enhance
the calculations that showed no discernible difference in the
S–N bond energies irrespective of whether the nitrosothiol
was primary, secondary, or tertiary.63 In aerated solutions, the
decay was autocatalytic, thus implicating N2O3 as the autocata-
lytic propagating species (see rxns. [R6] and [R7]). Nitroso-
thiol decompositions were much slower in solutions in which
NO was not allowed to escape, implicating NO as a possible
retardant of nitrosothiol decomposition.25 This can be ex-
plained by the recombination of the thiyl radical with NO in
the reaction cage after initial cleavage. Elimination of NO
from the medium by bubbling air forces rxn. [R11] to the
right.
Published by NRC Research Press

734
Can. J. Chem. Vol. 90, 2012
Fig. 10. Evaluating the effect of oxygen on the rate of nitrosation of
cysteamine. The solutions purged with argon (trace a) give a very
slightly lower rate of nitrosation, indicating a small and insignifant
contribution of nitrosation by N2O3. [CA]0 = 0.05 mol/L; [H+]0 =
Fig. 11. (A) Evaluation of unambigous nitrous acid dependence by
employing [H+]0 = [NO2 ]0 and varying both at the same equimolar
–
concentrations. [CA]0 fixed at 0.25 mol/L. (a) [H+]0 = [NO2 ]0 =
–
0.005 mol/L, (b) [H+]0 = [NO2 ]0 = 0.010 mol/L, (c) [H+]0 =
–
[NO2 ]0 = 0.015 mol/L, (d) [H+]0 = [NO2 ]0 = 0.020 mol/L, (e)
–
–
–
0.08 mol/L; and [NO2 ]0 = 0.07 mol/L. CA, cysteamine hydrochloride.
[H+]0 = [NO2 ]0 = 0.025 mol/L, (f) [H+]0 = [NO2 ]0 = 0.030 mol/L,
–
–
and (g) [H+]0 = [NO2 ]0 = 0.035 mol/L. (B) Plot of initial rate of
–
nitrosation versus initial nitrous acid concentrations as calculated in
Table 2 and derived from the data in Fig. 11A. This plot is linear, as
opposed to Fig. 6C.
½R11ꢀ
RSNO $ RSꢁ þ NO
No mechanism has yet been proposed for the implication
of N2O3 in the autocatalytic mechanism. One can assume
that N2O3 should assist in the cleavage of the S–N bond,
thus accelerating the rate of homolytic decomposition of
RSNOs.
In aerated solutions, NO then undergoes rxns. [R6] and
[R7] to yield N2O3.
½R12ꢀ
RSNO þ N2O3 ! RSNðNO2ÞOꢁ þ NO
–
The incorporation of the NO2 group in the nitrosothiol
would weaken the S–N bond, leading to its easy cleavage
and regeneration of the N2O3 molecule, which will further
catalyze the decomposition.
½R13ꢀ
RSNðNO2ÞOꢁ ! RSꢁ þ N2O3
The sequence of rxns. [R11]–[R13] will deliver autocataly-
sis in N2O3, since rxn. [R13] produces N2O3, which recycles
back into rxn. [R12].
with the maximum metal ion concentrations of 0.43 ppb in
Pb2+, small differences can be observed in EDTA-laced reac-
tion solutions and those that are not (see traces a and b in
Fig. 7C). Nitrosothiol decompositions have also implicated
other metal ions and not just specifically copper.66 Although
Cu(II) ions are used to catalyze nitrosothiol decomposition,
the active reagent is Cu(I). In this catalytic mechanism, Cu
redox cycles between the +2 and +1 oxidation states. Only
trace amounts of Cu(I) are needed to effect the catalysis, and
these can be generated from the reduction of Cu(II) ions from
minute amounts of thiolate anions that exist in equilibrium
with the nitrosothiol.
Recent results have suggested a novel pathway for the de-
composition of nitrosothiols that involves an initial internal re-
arrangement to form an N-nitrosation product.64 Typical
nitrosothiol decomposition products after such an internal rear-
rangement would include thiiranes and 2-hydroxy-mercaptans.
This mechanism has been used successfully to explain the
production of a small but significant quantity of 2-hydroxy-
3-mercaptopropionic acid (apart from the dominant product
cystine) in the decomposition of S-nitrosocysteine. The GC–
MS spectrum shown in Fig. 2, however, gives a dominant
signal for dimeric cystamine, with very little evidence of
any other significant products, indicating very little initial
N-transnitrosation.
½R14ꢀ
½R15ꢀ
½R16ꢀ
RSNO $ RSꢂ þ NOþ
The copper-catalyzed decomposition of nitrosothiols has
been well-documented.65–67 The varying rates of nitrosothiol
decompositions reported by different research groups can be
traced possibly to adventitious metal ion catalysis from the
water used in preparing reagent solutions. Even in our case,
2Cu2þ þ 2RSꢂ ! RSSR þ 2Cuþ
Cuþ þ RSNO ! ½RSðCuÞNOꢀþ
Published by NRC Research Press

Morakinyo et al.
735
Table 2. Input species concentration and final reagent concentrations for nitrous acid depen-
dence experiments with [H+]0 = [NO2 ]0.
–
[H+]added
[NO2 ]added
[HNO2]initial
5.00×10–3
1.00×10–2
1.50×10–2
2.00×10–2
2.50×10–2
3.00×10–2
3.50×10–2
[H3O+]final
1.42×10–3
2.11×10–3
2.64×10–3
3.08×10–3
3.48×10–3
3.83×10–3
4.16×10–3
[NO2 ]final
[HNO2]final
3.58×10–3
7.89×10–3
1.24×10–2
1.69×10–2
2.15×10–2
2.62×10–2
3.08×10–2
–
–
5.00×10–3
1.00×10–2
1.50×10–2
2.00×10–2
2.50×10–2
3.00×10–2
3.50×10–2
5.00×10–3
1.00×10–2
1.50×10–2
2.00×10–2
2.50×10–2
3.00×10–2
3.50×10–2
1.42×10–3
2.11×10–3
2.64×10–3
3.08×10–3
3.48×10–3
3.83×10–3
4.16×10–3
Table 3. The set of reactions used for simulating the nitrosation of cysteamine.
RSH, cysteamine.
Reaction
No.
M1
M2
M3
M4
M5
M6
M7
M8
Reaction
kf
kr
H+ + NO2 ⇌ HNO2
1.10×109
1.0×102
3.16×105
17.9
6.18×105
1.0×108
1.00×107
~0
–
2HNO2 ⇌ N2O3 + H2O
HNO2 + H+
⇌
+N=O + H2O
RSH + HNO2 → RSNO + H2O
RSH + +N=O ⇌ RSNO + H+
N2O3 ⇌ NO + NO2
2RSNO ⇌ RSSR + 2NO
RSH + N2O3 → RSNO + HNO2
6.7×104
8.0×109
4.0×10–3
5.0×103
1.0×10–3
6.0×108
~0
~0
Note: kf, forward rate constant; kr, reverse rate constant.
Further experiments on the effect of nitrous acid
½R17ꢀ
½RSðCuÞNOꢀþ ! RSꢂ þ NO þ Cu2þ
Since nitrous acid is the major nitrosation species, another
series of experiments were undertaken in which the effect of
nitrous acid could be unambiguously determined. For the
data shown in Fig. 11A, nitrite and acid were kept at equimo-
lar concentrations and varied with that constraint. Final pro-
ton and nitrite concentrations after dissociation are also
equal for electroneutrality (see Table 2). Thus, if our simple
mechanism is plausible, a plot of nitrous acid versus initial
nitrosation rate should be linear with the expected slight tail-
ing at either extreme of the nitrous acid concentrations owing
to nonlinear generation of the free acid derived from the
quadratic equation. Using the same rate law as eq. [1], a
value of k1 can be derived and compared with values ob-
tained from the Fig. 6C plot. Kinetics constants derived from
the Fig. 11B plot are not expected to be as accurate as those
from Fig. 6C, but should deliver values close to those from
high acid environments. For the case of nitrous acid as the
nitrosating agent (too low a concentration of excess H3O+ to
effect equilibrium of rxn. [R9]), data derived from Fig. 11B
delivered an upper limit rate constant for k1 that was higher
than that derived from Fig. 6C-type data, but well within the
error expected for such a crude extrapolation. Other sets of
kinetics data were derived in which CA concentrations were
varied while keeping the nitrous acid concentrations constant.
Varying nitrous acid concentrations were used and, for each,
a CA dependence was plotted (data not shown). These were
straight lines through the origin and with the slope deter-
mined by the concentration of nitrous acid and the free pro-
tons existing in solution after the dissociation of nitrous acid.
For these kinetics data, eq. [1] had to be evaluated com-
pletely, since Ka ≈ [H3O+] and no valid approximation could
be made. The values of k1 and k2Kb were utilized in eq. [1]
to recalculate initial rates and slopes with respect to CA
The cascade of rxns. [R14]–[R17] continues until the ni-
trosothiol is depleted and converted to the disulfide. The
complex formed in rxn. [R16] has been justified through den-
sity functional theory studies that report that Cu(I) binds
more strongly to the S center than to the N center of the ni-
trosothiol.65 This preferential binding weakens and lengthens
the S–N bond, thus facilitating its cleavage. It is difficult to
extrapolate this proposed mechanism into the physiological
environment, but some experimental results have shown that
protein-bound Cu2+ catalyzed nitric oxide generation from ni-
trosothiols, but not with the same efficiency as the free hy-
drated ion.68 Further studies are needed to evaluate the
efficiency of this catalysis, especially in copper-based en-
zymes.
Effect of Cu(II) on the nitrosation of cysteamine (Fig. 7B)
The fact that Cu(II) ions can assist in nitrosating CA from
nitrite in the absence of acid while Cu(I) ions are unable to
effect this can be explained by recognizing that nitrosation is
an electrophilic process, where NO effectively attaches as the
(formally) positive charge, while eliminating a positively
charged proton. Cu(II), in this case, acts by complexing to
the nitrite and forming a nitrosating species that will release
the Cu(II) ions after nitrosation.
Published by NRC Research Press

736
Can. J. Chem. Vol. 90, 2012
Fig. 12. (A) Simulation of the short set of reactions shown in Ta-
ble 3. Solid lines are experimental data and the symbols indicate si-
mulations. The mechanism is dominated by the value of the
bimolecular rate constant for the nitrosation of cysteamine hydro-
centration ensured that the simulations were insensitive to
value of the bimolecular nitrosation constant for N2O3
(kM8). Figure 12B clearly shows that the concentration of
N2O3 never rises to a level where it would be significant.
The rate constants adopted for rxn. [M1] (Table 3) were
not crucial for as long as they were fast (in both directions)
and were connected by Ka. Increasing the forward and re-
verse rate constants while maintaining Ka only slowed the
simulations, but did not alter the simulation results. Reac-
tion [M2] (Table 3) was also a rapid hydration reaction
that strongly favored HNO2, thus rapidly depleting any
N2O3 formed or merely keeping these concentrations de-
pressed should the tandem rxns. [R6] + [R7] have been
added to the mechanism. These were not utilized for this
simulation because EPR experiments had not detected NO
at the beginning of the reaction and had only detected NO
during the decomposition of the nitrosothiol. There are no
reliable thermodynamics parameters for rxn. [M3] (Table 3),
since this would be heavily dependent on the pH of the me-
dium. For the purposes of this simulation and some of our
previous estimates, the equilibrium constant for rxn. [M3]
was set at 3.16 × 10–2 (mol/L)–1. Kinetics constants for
rxns. [M4] and [M5] (Table 3) were derived from this
study. In the absence of significant concentrations of N2O3,
the simulations were insensitive to the kinetics constants
utilized for rxns. [M6] and [M8] (Table 3). Reaction [M7]
(Table 3) was assumed to be slow and, in the absence of
Cu(II) ions or other adventitious ions, this is indeed so (see
Figs. 7A and 7C). The simulations are shown in Fig. 12A,
which gives a reasonably good fit from such a crude mech-
anism. Figure 12B is provided to show the concentration
variation of some of the species that could not be experi-
mentally determined. Of note is that N2O3 concentrations
are depressed throughout the lifetime of the nitrosation.
Even though +NO is also depressed in concentration, this
can be explained by its high rate of nitrosation as soon as
it is formed and an equilibrium that favors its hydration
back to HNO2 once the thiol is depleted.
–
chloride (CA) by HNO2 (k1). [CA]0 = [NO2 ]0 = 0.05 mol/L. (B)
Results from the modeling that delivered simulations of the two
traces shown in Fig. 12A. These simulations show the concentration
variations of those species that could not be experimentally moni-
tored. Neither N2O3 nor NO+ rise to any significant concentration
levels.
Conclusion
This detailed kinetics study has shown that, even though
there are several possible nitrosating agents in solution, in
the physiological environment, with its slightly basic pH and
low oxygen concentrations, the major nitrosant is HNO2. Fur-
thermore, the lifetime of CANO is not long enough to be
able to be effectively used as a carrier of NO. GSH, which
has been touted as a possible carrier of NO, appears to have
a shorter lifetime than CANO. However, because of its over-
whelming concentration in the physiological environment it
might still be a major carrier of NO when it is compared
with other thiols. Protein thiols, it appears, would be the best
carriers by either intra- or inter-molecular transfer of NO
through protein nitrosothionylation through exposed and ac-
cessible cysteine groups.69
concentrations. These values were utilized to check the
accuracy of the kinetics constants derived from high-acid-
dependent experiments. It was noted that, within experimental
error, these values checked and were generally robust.
Simulations
A small set of simple reactions was utilized to simulate the
nitrosation kinetics. This set is shown in Table 3. At low acid
concentrations, the nitrosation is dominated by HNO2 as the
nitrosant, and as acid concentrations are increased, both nitrous
acid and the nitrosonium cation are involved in the nitrosation.
Even though Table 3 includes nitrosation by N2O3, the lack of
oxygen (approximately kept constant at 2.0 × 10–4 mol/L) and
its most favored and rapid reaction to nitrous acid made the
nitrosation by N2O3 insignificant. N2O3 is the dehydrated
form of HNO2 (rxn. [R8]), while reaction conditions were
in an overwhelming excess of water. Its insignificant con-
Acknowledgements
This work was supported by Grant No. CHE 1053611
from the National Science Foundation (R.H.S.) and partly
from Grant No. R01 EB2044 from the National Institutes of
Health (R.S.).The findings and conclusions in this report are
Published by NRC Research Press

Morakinyo et al.
737
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