Nitric Oxide Reacts Very Fast with Hydrogen Sulfide, Alcohols, and Thiols to Produce HNO: Revised Rate Constants

Article abstract of DOI:10.1021/acs.inorgchem.1c01061

The chemical reactivity of NO and its role in several biological processes seem well established. Despite this, the chemical reduction of ?NO toward HNO has been historically discarded, mainly because of the negative reduction potential of NO. However, this value and its implications are nowadays under revision. The last reported redox potential, E′(NO,H+/HNO), at micromolar and picomolar concentrations of ?NO and HNO, respectively, is between -0.3 and 0 V at pH 7.4. This potential implies that the one-electron-reduction process for NO is feasible under biological conditions and could be promoted by well-known biological reductants with reduction potentials of around -0.3 to -0.5 V. Moreover, the biologically compatible chemical reduction of ?NO (nonenzymatic), like direct routes to HNO by alkylamines, aromatic and pseudoaromatic alcohols, thiols, and hydrogen sulfide, has been extensively explored by our group during the past decade. The aim of this work is to use a kinetic modeling approach to analyze electrochemical HNO measurements and to report for the first-time direct reaction rate constants between ?NO and moderate reducing agents, producing HNO. These values are between 5 and 30 times higher than the previously reported keff values. On the other hand, we also showed that reaction through successive attack by two NO molecules to biologically compatible compounds could produce HNO. After over 3 decades of intense research, the ?NO chemistry is still there, ready to be discovered.

Full text of DOI:10.1021/acs.inorgchem.1c01061

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Nitric Oxide Reacts Very Fast with Hydrogen Sulfide, Alcohols, and
Thiols to Produce HNO: Revised Rate Constants
†
†
̂
Nicolas I. Neuman, Mateus F. Venancio, Willian R. Rocha, Damian E. Bikiel, Sebastián A. Suárez,*
and Fabio Doctorovich*
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ABSTRACT: The chemical reactivity of NO and its role in several
biological processes seem well established. Despite this, the
•
chemical reduction of NO toward HNO has been historically
discarded, mainly because of the negative reduction potential of
NO. However, this value and its implications are nowadays under
revision. The last reported redox potential, E′(NO,H+/HNO), at
micromolar and picomolar concentrations of •NO and HNO,
respectively, is between −0.3 and 0 V at pH 7.4. This potential
implies that the one-electron-reduction process for NO is feasible
under biological conditions and could be promoted by well-known
biological reductants with reduction potentials of around −0.3 to
•
−0.5 V. Moreover, the biologically compatible chemical reduction of NO (nonenzymatic), like direct routes to HNO by
alkylamines, aromatic and pseudoaromatic alcohols, thiols, and hydrogen sulfide, has been extensively explored by our group during
the past decade. The aim of this work is to use a kinetic modeling approach to analyze electrochemical HNO measurements and to
•
report for the first-time direct reaction rate constants between NO and moderate reducing agents, producing HNO. These values
are between 5 and 30 times higher than the previously reported keff values. On the other hand, we also showed that reaction through
successive attack by two NO molecules to biologically compatible compounds could produce HNO. After over 3 decades of intense
•
research, the NO chemistry is still there, ready to be discovered.
•
reduction of NO at a Hg electrode was shown to occur at a
potential of −1.0 V versus saturated calomel electrode, and
N2O was the only detected product.30,31
1. INTRODUCTION
•
Since the discovery of the physiological role of NO in the
1980s, the chemical and biochemical reactivity of this small
molecule was studied in detail over the following decades, and
it still continues to be reviewed.1−9 Azanone (HNO/NO−) or
However, at physiological pH, HNO is the main species with
an estimated E°(•NO,H+/HNO) ≈ −0.11 V, which becomes
−0.55 V at pH 7.12 In contrast to the well-accepted value for
•
nitroxyl is the one-electron-reduction product of NO, and its
the NO/NO− reduction potential, the NO/HNO reduction
•
reactivity and biological relevance are currently under intense
debate.10−13 HNO dimerizes very fast (kdim = 8 × 106 M−1
potential is not completely established. Shafirovich and Lymar
derived in 2002 the value of −0.55 V at pH 7 by using values
for ΔfG°(•NOaq) = 102 kJ mol−1 and the gas-phase enthalpy of
formation for 1HNO of 107 kJ mol−1 (as recommended by W.
R. Andersson),32 as well as the tabulated entropy S°(HNOgas)
= 220.7 J mol−1 K−1 and approximations for the free energy of
formation of HNO in aqueous solution, ΔfG°(1HNOaq), which
were calculated by estimating the HNO hydration energy as
similar to the hydration energy of HCN and H2CO. This value
has been revised in 2017 by Rocha et al.29 The authors
s−1),14 which limits its concentration and lifetime in solution.
•
Moreover, HNO reacts fast with its sibling NO (kNO,HNO
=
5.6 × 106 M−1 s−1)15 and at a moderate rate (kHNO,O = 3 × 103
2
M−1 s−1) with oxygen.14,16,17 Interestingly, despite the
overlapping reactivity of both compounds in biological
media18,19 with oxygen, thiols, ascorbic acid, transition
metals,20,21 and hemoproteins,22−25 the biochemical pathways
turn out to be very different.26,27
The proposed reduction potential for the •NO/3NO−
couple is −0.8 V.12 This was obtained by Bartberger et al. in
2002 with a combination of experimental and computational
methods and has been the most reliable value in the literature
since then.28 Our calculation of this reduction potential
through TPSSh-D3(BJ)/def2-QZVP/COSMO(SMD) yielded
a reduction potential of −0.81 V, which is in excellent
agreement with the accepted value.29 Experimental irreversible
Special Issue: Renaissance in NO Chemistry
Received: April 6, 2021
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computed the reduction potential at pH 7 in aqueous solution
using quantum-mechanical calculations and Monte Carlo
statistical mechanical simulations. It was derived as an efficient
protocol to obtain the solvation free energy of HNO using
three different computational approaches. A combination of
linear response theory and continuum solvation approaches
(COSMO/SMD) for more than a hundred uncorrelated
structures obtained from Monte Carlo simulations was used
to obtain −4.62 kcal mol−1 as the solvation free energy of
HNO, which is at least 2 times greater than the value of −5 kJ
mol−1 previously estimated by Shafirovich and Lymar in 2002,
on the basis of chemically related compounds. Considering the
In this work, first of all, we developed a kinetic modeling
approach to analyze electrochemical HNO measurements, in
order to determine, for the first time, the direct kRXH,NO rate
•
constant for the reaction between NO and ascorbic acid,
cysteine, and H2S, as a proof of concept. These values are up to
6 times higher than the previously reported keff values. On the
other hand, we also showed that the reaction of the successive
attack by two •NO molecules to biologically compatible
compounds could produce HNO.
2. MATERIALS AND METHODS
2.1. Experimental Procedures. Reagents. Cobalt(II)
5,10,15,20-tetrakis[3-(p-acetylthiopropoxy)phenyl]porphyrin was
purchased from Frontiers Scientific. All other reagents were purchased
from Sigma-Aldrich and used as received. NO was generated
anaerobically by the dropwise addition of degassed water to a mixture
of 4 g of NaNO2, 8.5 g of FeSO4, and 8.5 g of NaBr. The produced
NO was passed through a NaOH solution to remove higher oxides
and bubbled into degassed water in order to get a saturated solution
of NO ([NO] = 2 mM). This solution was kept under a strictly
anaerobic NO atmosphere and injected into the electrochemical cell
using a gastight syringe. The solutions of H2S were prepared by
bubbling with H2S(g) and dissolving Na2S. The reaction rate obtained
was similar using both methods, so it does not seem to be influenced
by the typical impurities present in NaHS (it is enriched with
polysulfides).
•
2017 protocol, it seems reasonable to take for the NO/HNO
reduction potential in aqueous solution, the value obtained by
Rocha and co-workers, E°′ = −0.16 V at pH 7,29 showing that
it is a viable process under physiological conditions.
Liochev and Fridovich had discussed the possibility of the
production of NO/HNO within living cells.17 The equilibrium
•
redox potential for NO/HNO in the cell can be much less
negative because the [•NO]/[HNO] ratio can be high, and
according to the Nernst equation, its logarithm influences the
actual redox potential. In this context, using the E°(NO,H+/
HNO) ≈ −0.11 V, the calculated redox potential E′ at
micromolar and picomolar concentrations of NO and HNO,
respectively, is between −0.30 and 0 V at pH 7.4. These redox
values are compatible with the observed reactions of NO with
alcohols, thiols, and HS−, as will be described in more detail
below.
Amperometric measurements of the HNO concentration were carried
out as previously described.43−45 A three-electrode system consisting
of a Pt counter electrode, a Ag/AgCl reference electrode, and a
working electrode consisting of Au modified with a monolayer of
cobalt porphyrin with 1-decanethiol which was covalently attached.
The method was demonstrated to be specific for HNO, showing no
Pryor, De Master, and co-workers found that anaerobic
aqueous solutions of thiols and thiol-containing proteins
exposed to NO result in quantitative formation of the
corresponding disulfides (RSSR) or sulfenic acids RS−
OH.33,34 These reactions, which took place in minutes, were
shown to produce N2O. The authors disregarded HNO as a
reaction intermediate, possibly because of the lack of selective
detection methods for HNO, with all being indirect at the
moment (1980s and 1990s). In this context, during the past
decade, the biologically compatible chemical reduction of •NO
(nonenzymatic) as direct routes to HNO was explored by our
group.35−40 The reduction of •NO to HNO has to be
necessarily coupled to other reactions that produce com-
pounds such as N2O, driving the reaction forward and
overcoming an unfavorable thermodynamic barrier. This
chemical reductive route recently received key support,
showing that •NO can be converted to HNO by the
aforementioned thiols,37 but also by hydrogen sulfide,35,40,41
aromatic and pseudoaromatic alcohols (i.e., ascorbic acid,
tyrosine, and salicylic acid),36,39 and alkylamines.38 In all cases,
the final concentrations of nitrite in solution and the amount of
N2O in the reaction chamber headspace were determined.
Both were produced in approximately a 1:1 ratio, as was
expected to occur due to the reaction between HNO and •NO.
−
interference or spurious signal due to the presence of NO, O2, NO2 ,
and other RNOS. Signal recording was performed with a TEQ 03
potentiostat. For each case, we also confirmed that the maximum
employed concentrations of NO and all organic reductants produced
negligible signals by themselves.
•
2.2. Kinetic Modeling of Reactions of NO with Reductants
to Produce HNO. In previous work, we have utilized an HNO-
specific electrode for sensing HNO concentrations as functions of
time and applied it to the analysis of several reactions that were shown
to produce HNO, as well as reactions of HNO with different chemical
and biological targets.44−46 In order to retrieve detailed kinetic
information related to HNO generation and consumption, a fairly
complicated kinetic model is necessary, and we have developed and
used this model on occasions.44,46 In many other cases, we used, for
simplicity, the electric current traces obtained with the electrode in a
semiquantitative way. The electric current traces for HNO sensing
usually show a relatively fast increase, which peaks and then decays
with different rates. In a full kinetic analysis, the slope of the initial
increase, the peak current, and the shape and rate of the decrease of
the current after the peak all contain relevant information. In the
simplified approach, only the peak current and position in time are
used. The peak current is transformed to a maximum concentration of
HNO, [HNO]max, by calibration with a standard curve derived from
the addition of Angeli’s salt (AS) to the cell, which has a known
decomposition constant. Then an effective rate velocity for each
experiment is derived as
−
2•NO + HNO → → → NO2 + N2O + H+
In previous work, we reported the effective rate constants keff
for the reactions of NO with the mentioned reductants in
anoxic conditions.42 These estimated values are a result of the
combined generation and consumption pathways. Because the
decay of HNO is extremely fast (see above), these values are
underestimations and imply that the “real” rate constants have
to be larger. A more detailed discussion of the relationship
between the “effective” and “real” rate constants is given in
Section 2.2 and the Supporting Information (SI).
Δ[HNO]max
v
≈
eff
Δtmax
The order of the reaction is also derived from a plot of veff as a
function of the reactant concentrations and a rate law is proposed,
from which an “effective” rate constant keff can be derived. This is an
empirical methodology that allows a comparison of several related
reactions and allows one to obtain kinetic information from an
otherwise very complicated reaction system.
B
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•
Scheme 1. General Mechanism for HNO Generation from the Reaction of NO with a Reducing RXH Molecule and
Electrocatalytic Oxidation by the HNO-Specific Electrode
d[CoIIIPNO−]
However, in certain situations, it is necessary, in order to gain more
insight, to utilize a more complete kinetic model including the
interaction of HNO with the electrode and several of the known
reactions of HNO with other species. Although this model provides
more information than the simplified “effective” rate model, it still
requires many assumptions and simplifications. The numerical
= kon[CoIIIP][HNO] − kox[CoIIIPNO−]
(4)
(5)
dt
d[CoIIIPNO]
= −koff,NO[CoIIIPNO] + kox[CoIIIPNO−]
dt
relationship between the “effective” and model-derived constants
depends on which pathways for HNO generation and consumption
are included in the model. Two possible cases are discussed further in
the SI.
Additionally, two decay routes for HNO are always included, which
involve the reaction with •NO and HNO dimerization. The values for
these rate constants were fixed from literature values (kHNO,NO = 5.8 ×
106 M−1 s−1 and kdim = 8 × 106 M−1 s−1). The suffix “decay” refers to
the contribution of these two reactions to the change in the HNO
concentration.
Determination of the kinetic rate constants for HNO generation
and consumption using our previously reported HNO-specific
electrode36,37,44,46 involves simulations of current traces derived
from the binding of HNO to an electrode-anchored cobalt(III)
porphyrin, followed by its electrocatalytic oxidation to •NO and
subsequent release. The simulations require the solution of a large set
of coupled differential equations that include the catalytic cycle
occurring at the surface of the electrode, as well as different HNO
generation and consumption routes that occur in the bulk. The
general mechanistic model is shown in Scheme 1, where the rate
constants involved in the model are written in red.
i d[HNO]y
= −kNO,HNO[HNO][NO] − 2kdim[HNO]2
j
j
j
z
z
z
dt
k
{
decay
(6)
(7)
d[NO]
dt
= −kNO,HNO[HNO][NO]
Finally, for each particular system, there will be specific HNO
generation and decay pathways that will generate changes in the
For example, while the formation of a semistable RX(HNO)
adduct (we do not specify the charge or spin state of this adduct) may
be modeled by a simple bimolecular reaction step, with a single
constant kRXH,NO,adduct, its subsequent decay to release HNO may
occur through several different pathways. This is most likely the case
•
HNO and NO concentrations in the bulk solution. These terms are
labeled by the suffix “generation”. The total change in the HNO
concentration can be written by the addition of the elec, decay, and
generation terms in the differential equations, and similar expressions
•
can be written for the NO concentration.
•
in the reaction of NO with thiols, which will be discussed in the
following sections. Therefore, the rate constant kRX(HNO),decay must be
seen as a phenomenological unimolecular reaction rate, which does
not imply a particular mechanism. The necessity of inclusion of the
adduct in the kinetic modeling will become clear below.
The set of differential equations that describe the reactions
occurring at the surface of the electrode are shown below. The suffix
“elec” (short for electrochemical) in the differential equation for
HNO indicates that these terms include only the reactions at the
electrode surface. F is Faraday’s constant.
i d[HNO]y
i d[HNO]y
i d[HNO]y
j
j
j
z
z
z
j
j
j
z
z
z
j
j
j
z
z
z
=
+
dt
dt
dt
k
{
k
{
k
{
total
elec
decay
i
d[HNO]
dt
y
j
j
j
z
z
z
+
k
{
generation
(8)
The sets of differential equations were numerically solved, and the
current versus time traces were simulated, using the ordinary
differential equation (ode) functions in the GNU Octave program.47
We have previously validated the model by simulating the current
profile when different concentrations of AS (with a known decay
constant of 9.1 × 10−4 s−1)48 were added to the electrochemical cell.
The obtained kinetic parameters are shown in Table 1. These are the
same ones previously reported by us,46 and we assume that they
remained unchanged for the experiments with ascorbic acid, cysteine,
or hydrogen sulfide.
2.3. Computational Details. Reaction Mechanisms. All geo-
metries involved in the reaction mechanisms proposed in this work
were obtained at the TPSSh-D3(BJ)/def2-TZVPP level of
theory.49−51 In order to account for the effects of the solvent
environment, the conductor-like polarizable-continuum model (C-
PCM)52 in association with the solvation model based on density
(SMD)53 was employed on geometry optimizations. In order to
dI
dt
d[CoIIIPNO−]
dt
= koxF
(1)
(2)
i d[HNO]y
= −kon[HNO][CoIIIP] + koff[CoIIIPNO−][H+]
j
j
j
z
z
z
dt
k
{
elec
d[CoIIIP]
dt
= −kon[HNO][CoIIIP] + koff[CoIIIPNO−][H+]
+ koff,NO[CoIIIPNO]
(3)
C
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of magnitude lower than those in the AAH− experiments, and
the ones used in the H2S experiments are almost 1 order of
magnitude lower still; however, similar peak currents are
Table 1. Reactions and Rate Constants for Generation and
Decay of HNO, Together with the Electrochemical Model
of the HNO-Specific Electrode
a
•
obtained, which indicate that the reaction rates between NO
reaction
rate constants
and RXH decrease in the order RXH = H2S ≫ CysH ≫
AAH−.
Generation
−
kAS = 9.1 ×
AS(N2O32−) + H+ → HNO + NO2
Initially, the only HNO source considered in the
(9)
10−4 s−1
•
mechanistic model was the direct reaction NO + RXH →
Decay
HNO + RX•, where the RX• radical could be the ascorbyl
radical,36 or it could be a thiyl radical RS•. Although these
radicals would react further and be eliminated from the
solution (by reaction with •NO, disproportion of ascorbyl into
dehydroascorbic and ascorbic acid, or disulfide bond
formation), the fate of the radicals was not explicitly included
in the models. This followed from the need to reduce the
number of independent reaction rates included in the models.
A set of initial simulations of the experimental data in Figure 1
reproduced very well the first few seconds of each experiment
(fast rise of the current), but inevitably the models predicted a
decay of the current, which was much faster than the
experimental one. The reason for this fast decay was the
extremely large kNO,HNO and kdim rate constants. To explain this
observation, we proposed the formation of an adduct between
RXH and •NO, which could act as an HNO “reservoir”,
releasing HNO more slowly or directly reacting with the
electrode and donating NO− to the cobalt(III) porphyrin. This
adduct is in agreement with experimental observations: in the
case of the ascorbic acid reaction with NO, we observed an
intermediate of the formula AA(H)NO• by cryo-MS36 and in
the case of thiols, an [RSNOH]• adduct has been proposed
before (Suarez et al. 2017 and references cited therein).37 The
inclusion of adduct formation and decay rate constants in the
models allowed us to reproduce quite well the “rise” and
“decay” regions of the current traces. The rate constants
derived from simulation of the current traces in Figure 1 are
shown in Table 2, and additional kinetic constants are written
in Table S1. The SI also contains an analysis of the sensitivity
of the simulations to different values of the rate constants
(Figures S2−S4 and the surrounding text). Table 2 also
includes effective rate constants between RXH and •NO,
reported in earlier works. These effective rate constants reflect
reactions that were written in a minimal or simplified form. In
some of these reactions, very unstable products are generated
(such as S•− in eq 23) that make the reaction, as written, highly
energetically unfavorable. The true mechanisms must involve
•−
k
NO,HNO = 5.8 ×
HNO + •NO → N2O2 + H+
(10)
(11)
106 M−1 s−1
kdim = 8 × 106
HNO + HNO → N2O + H2O
M−1 s−1
Electrochemical
kon = 3.1 × 104
CoIII(Por) + HNO → CoIII(Por)(NO−) + H+
(12)
M−1 s−1
koff = 0 s−1
kox = 0.089
CoIII(Por)(NO−) → CoIII(Por)(NO) + e−
(13)
b
s−1
koff,NO = 0.11
CoIII(Por)(NO) → CoIII(Por) + •NO
(14)
s−1
a
The amount of cobalt(III) porphyrin chemisorbed on the electrode
surface is included in the model through an empirical concentration of
b
10−9 M. The oxidation rate constant depends on the applied
potential. On our initial optimization of the HNO-specific electrode,
we found that a potential of 800 mV gave the best response and
always work at this potential.
characterize the species involved on each mechanism as a local
minimum (absence of imaginary frequencies) or a transition state
(only one imaginary frequency on the reaction coordinate mode),
vibrational frequencies were also calculated at the same level of
theory.
All of these calculations were also performed using the ORCA 4.2.1
software.54
3. RESULTS AND DISCUSSION
3.1. Kinetic Analysis of Electrochemical HNO Sensing
Experiments. We performed simulations of the HNO-specific
electrode current responses for the reactions of •NO with three
organic reductants: ascorbic acid (AA stands for ascorbic acid,
while AAH− stands for an ascorbate anion, which is the form at
pH 7.5), L-cysteine (CysH), and H2S. The experimental and
simulated current traces are shown in Figure 1. The
concentrations of reactants are given in the legend to Figure
1; the peak current values for the different traces increase
approximately linearly with the concentrations of the reagents.
The concentrations used in the CysH experiments are 1 order
Figure 1. Experimental (black) and simulated (red) currents (I, nA) versus time (t, s) for reactions of AAH− (0.4, 0.6, and 1 μM) with •NO (200
μM) (left), CysH (10 μM) with •NO (1, 2, and 3.8 μM) (middle), and H2S (2 μM) with •NO (0.2, 0.4, and 0.6 μM) (right). The reactions were
performed anaerobically at pH 7.4 (buffer phosphate) and with an applied potential of 800 mV. The arrows indicate the order of increasing
concentration.
D
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Table 2. Rate Constants for Adduct Formation and Decay or for Direct Reaction between RXH and NO, as Well as Effective
a
Rate Constants Previously Reported
reaction
•NO + AAH− → AAH(HNO)•−
AA(HNO)•− → HNO + AAH•−
NO• + AAH− → HNO + AAH•−
kRXH,NO,adduct or kRXH,NO,direct
kRX(HNO),decay
keff
ref
36
37
(15)
(16)
(17)
28(4) M−1 s−1
1.8(2) s−1
b
43(15) M−1 s−1
8(1) M−1 s−1
•NO + CysH → Cys(HNO)•
Cys(HNO)• → HNO + Cys•
•NO + CysH → HNO + Cys•
•NO + HS− → S(HNO)•−
(18)
(19)
(20)
(21)
(22)
180(20) M−1 s−1
0.5(1) s−1
8(1) × 102 M−1 s−1
8(1) × 103 M−1 s−1
25(1) M−1 s−1
S(HNO)•− → HNO + S•−
0.16(1) s−1
•NO + HS− → HNO + “S•−
”
6(1) × 103 M−1 s−1
35, 42
•NO + HS− + X/X• → HNO + SY•−/SY−
c
(23)
a
b
Errors are shown in parentheses. Measured by simulating the rate of generation and decay of the ascorbyl radical signal, followed by time-
c
resolved EPR. The species in quotation marks is highly unstable, making the reaction as written very unfavorable. Species such as this will likely
react very rapidly with other radicals, such as •NO, or other sulfide-derived species, leading to polysulfides, and thus the species in quotation marks
must be considered transient or a very short-lived intermediate. The second reaction could become thermodynamically favorable if another reactant
X/X• reacts in a concerted or sequential fashion with a transient S•− to produce a SY•− or SY− species. X/X• could be a sulfur-derived species,
•NO, or another sulfur-centered radical.
Figure 2. TPSSh-D3(BJ)/def2-TZVPP/CPCM(SMD)-optimized geometries for the first step proposed.
coupled reactions, which lead to more stable products and thus
drive the reactions forward. However, the observed HNO
production is compatible with a first-order rate on both NO
and RXH, and therefore we write these reactions in the
simplest possible way, which does not capture the complete
mechanism, only the dependence of HNO production on the
reactant concentrations.
For the ascorbic acid and H2S experiments, the traces could
be simulated by considering only the formation and decay of
= 43(15) M−1 s−1, measured by following the generation and
decay of the ascorbyl radical signal by EPR,36 is in good
agreement with the value obtained in this work. Inclusion of
the adduct formation reactions made it necessary to include
the direct reaction of the adducts with the electrode, with
forward and reverse rate constants provided in Table S1. It is
important to note that the information derived from the
electrochemical measurements does not allow one to
determine the actual identity of the RX(HNO)• adducts, and
especially the exact mechanism for its decomposition. The
•
the adduct, without the inclusion of a direct reaction (kH S,NO
2
and kAsc,NO,direct were both set to 0 M−1 s−1). As can be seen
from Figure 1 (left and right), the adduct model reproduces
quite well the experimental traces. The adduct formation rates
were 28 M−1 s−1 for the reaction between •NO and AAH− and
8 × 103 M−1 s−1 for the reaction between H2S and •NO. In the
case of CysH, both pathways were considered, obtaining 180
M−1 s−1 for the direct route and 700−900 M−1 s−1 for the
adduct formation route. The rate constants for AAH− and
CysH are already quite a bit larger than the previously
determined “effective” rate constants. However, the direct rate
constant for the reaction between NO and AAH−, kAsc,NO,direct
k
RX(HNO),decay rate constants were obtained from the choice of a
unimolecular decay route for the adducts, which was the
simplest possible model. In the case of the ascorbic acid-
derived adduct with HNO, a unimolecular decay route is likely,
based on the experimental and mechanistic results previously
reported.36
On the other hand, it is clear that the reaction does not
involve a simple outer-sphere reduction coupled to proton
release/uptake, which is thermodynamically unfavorable as
evidenced by the reduction potentials. E′(NO,H+/HNO), at
•
micromolar and picomolar concentrations of NO and HNO,
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respectively, is between −0.3 and 0 V at pH 7.4, and
E°′(RO•,H+/ROH) for alcohols is between 0.10 and 0.97 V at
pH 7.36 In the case of the CysH- and H2S-derived adducts with
HNO, HNO release may result from a more complex
mechanism. In the following sections, we show through
AIMD and DFT calculations that the formation of an RSNO•−
adduct, as well as the RSN(O)NO− adduct, is thermodynami-
cally favored. E°′(RSSR,2H+/2RSH) is close to −0.25 V, and
thus an outer-sphere electron transfer is, in principle,
thermodynamically favorable.
It is possible to observe that the prereactive complex
significantly stabilizes the system by approximately 6 kcal
mol−1. From this stabilization, a low barrier of 3 kcal mol−1 is
required for the formation of CH3SNO•−. This product, as
well as the product from the next step, involving the reaction
with another •NO molecule, is in agreement with the
formation of an adduct, as suggested by the kinetic data of
structurally related thiols such as cysteine. Despite the fact of
the intermediates being more stable than the product, this ΔG
is not large and may be interpreted as the prereactive complex,
with the product being in an equilibrium slightly shifted to the
prereactive side. Another remark should be added regarding
the N−O distance; Figure 2 shows that the formation of an
intermediate and the following structures leads to a slight
increase of the bond distance from 1.149 to 1.235 Å. This
increase may be due to repulsion between the negative charge
•
3.2. Mechanism for the Reaction of 2 equiv of NO
−
with RS Species. We propose a two-step mechanism for
55−57
•
reaction of the RS− species with two NO molecules,
sequentially, as follows:
CH3S− + •NO → CH3SNO•−
CH3SNO•− + •NO → CH3SN(O)N(O)−
(24)
(25)
•
from CH3S− and the unpaired electron of NO.
Figures 4 and 5 show the optimized geometries and the ΔG
profile for the second step of the proposed sequential
The optimized geometries were characterized as a local
minimum by the absence of imaginary frequencies, while
transition states were confirmed by the presence of a single
imaginary frequency. Figure 2 shows the geometries for the
•
first step of the sequential reaction of RS− with two NO
molecules. In the figure, “Reac” stands for the isolated CH3S−
•
and NO molecules, “Int” represents the prereactive complex,
“TS” corresponds to the transition state, and “Prod” stands for
the product of the reaction. Figure 3 shows the Gibbs free
energy profile for the reactants, precomplex, transition state,
and product.
Figure 5. Reaction Gibbs free energy, in kcal mol−1, for the second
step in the proposed mechanism. The imaginary frequency for the
transition state was evaluated as −552.71 cm−1.
mechanism. It is possible to observe that the interaction
between the reactants yields a species in which the nitric oxide
molecules are in trans-like orientation. This structure is about 1
kcal mol−1 less stable than the reactants and transition state,
which involves rotation of the NO moieties to reach a cis-like
orientation. Despite the relatively high barrier, the energy cost
for this rotational rearrangement is compensated for by the
approximate 20 kcal mol−1 stabilization provided to the system.
Figure 3. Reaction Gibbs free energy, in kcal mol−1, for the first step
of the proposed mechanism. The imaginary frequency for the
transition state was evaluated as −85.42 cm−1.
Figure 4. Optimized geometries for the second step in the proposed mechanism.
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It is important to point out that the results shown above
indicate that the second step of the proposed mechanism leads
to a RSN(O)NO− species in which the N−N bond is 1.6 Å.
This value is significantly smaller than the N−N distance in the
NO dimer of 2.1 Å obtained from the gas-phase MP2/def2-
QZVP reference geometry. This effect probably occurs because
of donation of the RS− electron density to the ON−NO
moiety of the products. Table S2 shows the CHELPG-
calculated charges for the Int, TS, and Prod species. It is
possible to observe that nitrogen atoms in the products have
Figure 6 shows the PES as a function of the rN−S distance for
the approach of the HS− (A) and H2S (B) species, respectively,
•
to a NO molecule. From these two figures, it is possible to
observe that the interaction between these species is
predominantly repulsive, reaching more than 150 kcal mol−1
at 1.2 Å.
These results reveal that the two-step •NO consecutive
reaction with H2S/HS− may lead to high activation energies in
the first step, and at this moment, we cannot provide a suitable
reaction mechanism. Similar to the investigation on the
reactivity of (NO)2 and RS− species, we studied the reaction
of the NO dimer with HS− through the following reaction:
charges of opposite sign (qN−S‑bonded = 0.837 and qN−N‑bonded
=
−0.397), leading to an attractive electrostatic interaction and,
consequently, the reduction of the N−N bond length.
HS− + (NO)2 → SNO− + HNO
•
(29)
3.4. Mechanism for the Reaction of 2 equiv of NO
with HS− Species. In order to evaluate the Gibbs free
energies and propose a molecular mechanism for the two-step
reaction of •NO with HS−, described by the chemical
equations
Figure 7 shows the optimized geometries of the reactants,
intermediate, transition state, and products for the reaction of
the NO dimer with HS−. It can be seen that the N−N distance
on (NO)2 is reduced by the formation of an entrance complex.
After the formation of this intermediate, HS− donates a proton
HS− + •NO → HSNO•−
HSNO•− + •NO → SNO− + HNO
(26)
•
to one NO, yielding HNO and a negatively charged SNO−
(27)
species.
we performed DFT calculations at the same level of theory of
the above-shown RS− species to find the reactants (Reac),
intermediates (Int), transition states (TS), and products
(Prod). After several unsuccessful attempts to obtain the TS
for
Figure 8 shows the Gibbs free energy for the reaction
coordinate. It is possible to observe that formation of the
entrance complex, which shortens the N−N distance, stabilizes
the reactants by approximately 6 kcal mol−1. Once stabilized, a
significant activation energy needs to be achieved in order to
form the products. We speculate that, despite the low
equilibrium constant for nitric oxide dimerization in the gas
phase,58−63 this stable intermediate might be acting as the
(NO)2 source because of the “accumulation” of this stable
species. On the other hand, it is also possible that the reactants
may not form the entrance complex, following a direct Reac →
TS path in order to find an effective activation barrier of 12.26
kcal mol−1 and yield the products, 3.46 kcal mol−1 more stable
than the reactants.
HS− + •NO → [HS ··· NO]‡•−
(28)
we decided to perform a systematic procedure to obtain a
reliable guess structure for this activated complex. We
performed a set of 20 TPSSh-D3(BJ)def2-TZVPP/CPCM-
(SMD) geometry optimizations by varying the N−S distance
from 3.2 to 1.2 Å. This potential energy surface (PES)
corresponds to the reaction coordinate for the HS−/H2S
•
reaction with NO, as shown in Scheme 2.
A summary of the energetics for the formation of several
nitrosylated thiol species, as well as other products of the
reactions of •NO with H2S/HS− or CH3S, is shown in Table 3,
together with previously reported energetics for the formation
of AAH(HNO)•−, PhSNHO•, and CH3SNHO•. The
previously reported calculations were performed at levels of
theory different from that used in this work. However,
qualitative comparisons can be made for certain steps. For
Scheme 2. Representation of the Reaction Coordinate for
•
Construction of the PES for HS− + NO → HSNO•− (the
Procedure for H2S Was Similar)
•
example, the activation energy for the reaction of NO with
CH3SH is much larger than that for the reaction with CH3S−,
and the latter reaction is also more favorable. This difference
likely arises from the concerted proton rearrangement
•
Figure 6. TPSSh-D3(BJ)/def2/TZVP/CPCM(SMD) PES surface as a function of the rN−S distance for NO and (A) HS− or (B) H2S.
G
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Figure 7. TPSSh-D3(BJ)/def2-TZVPP/CPSM(SMD)-optimized geometries for the direct reaction of HS− with the NO dimer.
•
(detected by EPR). Because HNO rapidly reacts with NO or
HNO and the ascorbyl radical disproportionates (less rapidly)
to ascorbate and dehydroascorbate, the relatively high energy
of the intermediate is compensated for.
From the entries in Table 3, we can observe some striking
•
tendencies regarding the reactivity of RSH with NO. The
reaction seems to become less favorable in the order R = Me,
•
Ph, and H, with the reaction between H2S and NO being
extremely unfavorable. The pKa of PhSH is 6.62, while that of
CH3SH is 10.4, and that of H2S is 6.90.64 There does not seem
to be a correlation between the activation energies and pKa of
the thiols. A similar tendency can be observed when CH3S−
•
and HS− reactivity is compared with that of NO. In these
Figure 8. TPSSh-D3(BJ)/def2-TZVPP/CPCM(SMD) Gibbs free
energy, in kcal mol−1, for the reaction of HS− with (NO)2. The
imaginary frequency for the transition state was evaluated as
−1179.24 cm−1.
cases, acidity is not a factor, so perhaps the dominating factor
is the stabilization of the sulfur-centered radical during the
reaction. In this regard, methyl or phenyl substituents could
play a role in the stabilization of a radical compared to H.
However, experimentally, the reaction between •NO and H2S/
HS− shows the fastest reaction constant, followed by PhSH
and cysteine (which could be considered the most similar to
CH3SH).37 This trend is the inverse of that predicted by the
calculations of the transition states obtained by the reaction of
•
occurring in the reaction of NO with CH3SH. A similarly
large activation energy was found for the reaction of PhSH
•
•
with NO. The reaction of AAH− with NO was found to
produce an AA(HNO)•− intermediate (ascorbic acid has two
vicinal acidic hydrogen atoms). This intermediate (detected by
MS)36 was calculated to be +16 kcal mol−1 above the reactants,
but the reaction was found to move forward by a series of
exergonic steps, such as HNO dissociation to produce AA•−
•
these sulfur compounds with NO. It is observed that the
•
activation energies for direct H2S/HS− with NO are much
higher compared to those of RSH (R = CH3, Ph). Therefore, a
•
Table 3. Summary of Energetics (Gibbs Free Energies) Obtained for the Reactions of Thiols/Thiolates with NO and (NO)2
a
and Previously Reported Reactions of Ascorbyl Anion and Thiophenol with NO
Int (kcal mol−1
)
TS (kcal mol−1
)
ΔG⧧ (kcal mol−1
)
Prod (kcal mol−1
)
ref
•
AAH− + NO → AA(HNO)•−
n.d.
n.d.
+16
−2.51
−0.68
−4.96
−18.45
−11.46
9.69
b, 36
c, 37
c, 37
•
PhSH + NO → PhSNHO•
0
0
19.67
23.36
−3.06
10.35
n.d
n.d
n.d
85.18
40.99
12.26
19.67
23.36
3.01
9.32
n.d
n.d
n.d
30.11
36.13
18.6
CH3SH + NO → CH3SNHO•
•
•
CH3S− + NO → CH3SNO•−
−6.07
1.03
d
d
d
d
d
d
d
d
CH3SNO•− + NO → CH3SN(O)NO−
•
CH3SN(H)O• + NO → CH3SN(O)NOH
•
CH3SN(O)NOH → CH3SN(O)N(H)O
CH3SN(O)N(H)O → CH3SNO + HNO
4.18
HS− + NO → SNOH•−
55.07
4.86
−6.34
63.48
19.23
−3.46
•
H2S + NO → [HSNHO]•
•
HS− + (NO)2 → SNO− + HNO
a
b
c
ΔG⧧ = G(TS) − G(Int). Calculated at the B3LYP/6-31G(d,p)/CPCM level. Calculated at the TPSSh/def2-TZVP/CPCM(SMD) level plus
d
two explicit water molecules. Calculated at the TPSSh-D3(BJ)/def2-TZVPP/CPCM(SMD) level of theory. The geometries and ΔG profiles are
shown in Figures S8−S15.
H
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more complicated mechanism than a simple bimolecular
reaction between •NO and H2S/HS− seems to be taking
place. Figures S10−S15 show that, in contrast to the reaction
mechanism for methyl and phenyl groups, the reactions of
•NO with H2S/HS−, in principle, do not occur through N−S
bonding but by proton transfer, which leads to N−S
coordination. It is also important to highlight that the energy
barriers for the formation of intermediates and transition
states, despite being higher than those for RSH, are lower for
H2S in comparison to HS− by approximately 40 kcal mol−1. On
the other hand, the reaction of HS− with (NO)2 is quite
favorable and could be taking place. Previously carried out
experiments show that this reaction is not first-order in NO
under all experimental conditions but shows from a kinetic
standpoint an intermediate order between 1 and 2, particularly
at higher concentrations of NO. This will be further
investigated from theoretical and experimental points of view
and informed in a near future.
reported keff values. These new higher values extend the idea
that reducing or hypoxic environments are expected to
generate HNO from nitric oxide. Nitroxyl is becoming
acknowledged as an endogenously produced messenger that
mediates specific physiological responses, many of which were
•
attributed yet to direct NO effects. For example, it is known
that H2S may transform endogenous NO• to HNO, which in
cells activates the HNO/transient receptor potential channel/
calcitonin gene-related peptide cascade, via the formation of
disulfide bonds, which results in continuous calcium influx.
This cascade regulates blood pressure and control of cardiac
contractility (unlike NO•), suggesting broad physiological
relevance of the findings. We have shown that the rate of HNO
production through an anaerobic reaction between NO• and
H2S/HS− is orders of magnitude faster than that for any
known HNO donor and could be taking place because of the
stabilization of (NO)2 by the negatively charged HS−. The
intermediate shown in Figure 7 has an N−N bond distance
that is 0.25 Å shorter than that in free (NO)2. On the other
hand, we also showed that the reaction of the successive attack
by two NO molecules to biologically compatible compounds
could produce HNO. The results summarized in Table 3
suggest deprotonated adducts of the type RXNO•− to be more
stable and to require lower activation energies to form
compared to protonated adducts of the type RX(HNO)•,
with the exception being R = H. Therefore, detailed
mechanistic reinvestigations of the reactions of NO with
aromatic alcohols are warranted.
The experimental bimolecular rate constants between RXH
•
and NO allowed us to calculate the activation free energies
ΔG⧧, using the Eyring equation, as described in the SI. The
equilibrium constant Keq, for formation of the reactant complex
from the individual reactants, was assumed to be equal to 1
M−1 because this value was not available for AAH− or CysH,
and for H2S or HS−, the energies of formation of the
intermediate differ considerably. The obtained activation free
energies were 15.6 kcal mol−1 (AAH−), 14.5 kcal mol−1
(CysH), and 12.2 kcal mol−1 (HS−). These values show a
relatively low dispersion, but because the Keq values likely differ
between reactions and from the value of 1 M−1, we cannot at
this point directly compare the activation energies derived
from the experimental bimolecular constants and theoretical
values.
ASSOCIATED CONTENT
■
sı
* Supporting Information
The Supporting Information is available free of charge at
https://pubs.acs.org/doi/10.1021/acs.inorgchem.1c01061.
Finally, the previous calculations show that the RSNO•− (R
= Ph, Me) and RSN(O)NO− or SN(O)NO2− adducts are
clearly stable with respect to the reactants, supporting inclusion
of the adduct hypothesis in the electrochemical kinetic
modeling. There are various ways in which these adducts can
decay and release HNO. RSNO•− can be protonated to give
HNO and RS•, which would rapidly dimerize or react further.
As a representative example, we have calculated one possible
pathway by which RSN(O)NOH can also release HNO and
RSNO through a series of intramolecular proton-transfer
reactions, as shown in Table 3 for R = Me. At this point, we
have only calculated the free energies of reaction for each of
these steps. The first steps, leading from CH3SH to
CH3SN(O)NOH through a sequential reaction with two
•NO molecules, are exergonic, while the following steps,
leading to CH3SNO and HNO, are slightly endergonic. This is
entirely compatible with the experimentally inferred formation
Description of the procedures for kinetic simulations, set
of differential equations that describe the reactions
occurring at the surface of the electrode, electronic
structure calculation details, as-optimized geometries,
Gibbs free energy, and CHELPG charges calculated for
intermediates, transition states, and products (PDF)
AUTHOR INFORMATION
■
Corresponding Authors
Sebastián A. Suárez − Departamento de Química Inorgánica,
Analítica y Química Física, Facultad de Ciencias Exactas y
Naturales, Universidad de Buenos Aires, Buenos Aires C1053,
Argentina; INQUIMAE-CONICET, Ciudad Universitaria,
Buenos Aires C1428EHA, Argentina; orcid.org/0000-
0003-0236-5743; Email: seba@qi.fcen.uba.ar
Fabio Doctorovich − Departamento de Química Inorgánica,
Analítica y Química Física, Facultad de Ciencias Exactas y
Naturales, Universidad de Buenos Aires, Buenos Aires C1053,
Argentina; INQUIMAE-CONICET, Ciudad Universitaria,
Buenos Aires C1428EHA, Argentina; orcid.org/0000-
0003-1088-2089; Email: doctorovich@qi.fcen.uba.ar
•
of an adduct between RXH and NO, followed by a slow
release of HNO. In the context of the electrochemical
measurements, both adducts could react with cobalt(III)
porphyrin, directly generating CoIII(Por)(NO−) and RS• or
RSNO.
4. CONCLUSIONS
Authors
•
It is noted that NO and HNO can be interconverted in
Nicolas I. Neuman − Instituto de Desarrollo Tecnológico para
la Industria Química, INTEC, UNL-CONICET, Paraje El
Pozo, Santa Fe 3000, Argentina; Institut fur Anorganische
Chemie, Universität Stuttgart, Stuttgart D-70569, Germany;
orcid.org/0000-0003-3368-0228
biologically compatible media, depending on the redox state of
the environment. We have shown that the direct reaction rate
constants between NO and moderate reducing agents
producing HNO are up to 30 times higher than the previously
̈
I
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of the Nitric Oxide Story. Trends Pharmacol. Sci. 2008, 29 (12), 601−
608.
̂
Mateus F. Venancio − Departamento de Físico-Química,
Instituto de Química, Universidade Federal da Bahia,
Salvador, Bahia 40170-110, Brazil
(12) Flores-Santana, W.; Salmon, D. J.; Donzelli, S.; Switzer, C. H.;
Basudhar, D.; Ridnour, L.; Cheng, R.; Glynn, S. A.; Paolocci, N.;
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Chemistry Is Unique Among Nitrogen Oxides in Biological Systems.
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Willian R. Rocha − Departamento de Química, Instituto de
Ciencias Exatas, Universidade Federal de Minas Gerais, Belo
̂
Horizonte, Minas Gerais 31270-901, Brazil; orcid.org/
0000-0002-0025-2158
Damian E. Bikiel − Departamento de Química Inorgánica,
Analítica y Química Física, Facultad de Ciencias Exactas y
Naturales, Universidad de Buenos Aires, Buenos Aires C1053,
Argentina; INQUIMAE-CONICET, Ciudad Universitaria,
Buenos Aires C1428EHA, Argentina
Complete contact information is available at:
https://pubs.acs.org/10.1021/acs.inorgchem.1c01061
Author Contributions
†These authors contributed equally. The manuscript was
written through contributions of all authors. All authors have
given approval to the final version of the manuscript.
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
The research reported in this publication was supported by the
Ministerio de Ciencia, Tecnología e Innovación Productiva
(Grants PICT-2015-3854 and PICT-2017-1930) and the
Universidad de Buenos Aires (UBACYT) Project
20020170100595BA. N.I.N., D.E.B., S.A.S., and F.D. are
members of CONICET. M.F.V. and W.R.R . thank the
Conselho Nacional de Desenvolvimento Científico e Tecno-
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̀
lógico, INCT-Catálise, and Fundaca̧ o de Amparo a Pesquisa
̃
do Estado de Minas Gerais for financial support.
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