Kinetics and Mechanism of Bioactivation via S-Oxygenation of Anti-Tubercular Agent Ethionamide by Peracetic Acid

Article abstract of DOI:10.1021/acs.jpca.6b07375

The kinetics and mechanism of the oxidation of the important antitubercular agent, ethionamide, ETA (2-ethylthioisonicotinamide), by peracetic acid (PAA) have been studied. It is effectively a biphasic reaction with an initial rapid first phase of the reaction which is over in about 5 s and a second slower phase of the reaction which can run up to an hour. The first phase involves the addition of a single oxygen atom to ethionamide to form the S-oxide. The second phase involves further oxidation of the S-oxide to desulfurization of ETA to give 2-ethylisonicotinamide. In contrast to the stability of most organosulfur compounds, the S-oxide of ETA is relatively stable and can be isolated. In conditions of excess ETA, the stoichiometry of the reaction was strictly 1:1: CH₃CO₃H + Et(C₅H₄)C(=S)NH₂ → CH₃CO₂H + Et(C₅H₄)C(=NH)SOH. In this oxidation, it was apparent that only the sulfur center was the reactive site. Though ETA was ultimately desulfurized, only the S-oxide was stable. Electrospray ionization (ESI) spectral analysis did not detect any substantial formation of the sulfinic and sulfonic acids. This suggests that cleavage of the carbon-sulfur bond occurs at the sulfenic acid stage, resulting in the formation of an unstable sulfur species that can react further to form more stable sulfur species. In this oxidation, no sulfate formation was observed. ESI spectral analysis data showed a final sulfur species in the form of a dimeric sulfur monoxide species, H₃S₂O₂. We derived a bimolecular rate constant for the formation of the S-oxide of (3.08 ± 0.72) × 10² M^-1 s^-1. Oxidation of the S-oxide further to give 2-ethylisonicotinamide gave zero order kinetics.

Full text of DOI:10.1021/acs.jpca.6b07375

Article
pubs.acs.org/JPCA
Kinetics and Mechanism of Bioactivation via S‑Oxygenation of
Anti-Tubercular Agent Ethionamide by Peracetic Acid
Kudzanai Chipiso,† Isabelle E. Logan,† Matthew W. Eskew,† Benard Omondi,‡ and Reuben H. Simoyi*,†,‡
†Department of Chemistry, Portland State University, Portland, Oregon 97207-0751, United States
‡School of Chemistry and Physics, University of KwaZulu-Natal, Westville Campus, Durban 4000, South Africa
ABSTRACT: The kinetics and mechanism of the oxidation of the important antitubercular agent,
ethionamide, ETA (2-ethylthioisonicotinamide), by peracetic acid (PAA) have been studied. It is effectively a
biphasic reaction with an initial rapid first phase of the reaction which is over in about 5 s and a second
slower phase of the reaction which can run up to an hour. The first phase involves the addition of a single
oxygen atom to ethionamide to form the S-oxide. The second phase involves further oxidation of the S-oxide
to desulfurization of ETA to give 2-ethylisonicotinamide. In contrast to the stability of most organosulfur compounds, the
S-oxide of ETA is relatively stable and can be isolated. In conditions of excess ETA, the stoichiometry of the reaction was strictly
1:1: CH3CO3H + Et(C5H4)C(S)NH2 → CH3CO2H + Et(C5H4)C(NH)SOH. In this oxidation, it was apparent that only
the sulfur center was the reactive site. Though ETA was ultimately desulfurized, only the S-oxide was stable. Electrospray
ionization (ESI) spectral analysis did not detect any substantial formation of the sulfinic and sulfonic acids. This suggests that
cleavage of the carbon−sulfur bond occurs at the sulfenic acid stage, resulting in the formation of an unstable sulfur species that
can react further to form more stable sulfur species. In this oxidation, no sulfate formation was observed. ESI spectral analysis
data showed a final sulfur species in the form of a dimeric sulfur monoxide species, H3S2O2. We derived a bimolecular
rate constant for the formation of the S-oxide of (3.08
2-ethylisonicotinamide gave zero order kinetics.
0.72) × 102 M−1 s−1. Oxidation of the S-oxide further to give
Sulfenic acid is, however, unstable and disproportionate to form a
stable S-oxide,18 which is also a reactive intermediate, and has the
capacity to also bind to cells.
INTRODUCTION
■
Ethionamide, (ETA), an antimycobacterial drug is used as a
second line therapeutic agent in treatment of multidrug resistant
tuberculosis.1 ETA is effective against M. tuberculosis, however,
its use has been restricted due to its toxicity and low therapeutic
index, a comparison of the amount of a therapeutic agent
that causes the therapeutic effect to the amount that causes
toxicity.2 As a prodrug, ETA requires bioactivation to exert its
pharmacological activity.3 It is oxidatively converted to its active
form by the Flavin-containing monooxygenase (FMO) EtaA of
Mycobacterium tuberculosis.4 This bioactivation results in reactive
S-oxide intermediate,5 which bind to NAD+/NADH; forming
adducts, thereby inhibiting the enoyl CoA reductase InhA.6,7
Mammalian flavin-containing monooxygenase (FMO) is also
active toward ETA, resulting in its bioactivation.8
A large number of antituberculosis drugs are known to be
hepatotoxic. ETA’s bioactivation by the host’s system has been
implicated in its hepatotoxicity and gastrointestinal toxicity due
to ethionamide.19 Despite great strides being made in the field of
toxicology, failure of drugs due to unfavorable kinetics properties
as well as unsafe toxic profiles remains a challenge. Here we
report on the kinetics and mechanistic information regarding
oxidation of ethionamide to aid the understanding of bioactiva-
tion of similar class of xenobiotics.
Bioactivation of drugs to reactive metabolites is an unfavorable
process during drug metabolism.9 Precautions are always taken
during drug development to eliminate possibility of reactive
metabolites.10,11 This bioactivation appears to be the prerequisite
for drug induced toxicity due to thiono-sulfur containing com-
pounds.12 Drugs containing a heteroatom such as sulfur, nitrogen
or phosphate, which are soft nucleophiles, are easily oxygenated.13
S-oxygenation is crucial in metabolism of drugs containing
the sulfur atom. However, the sulfur atom is bioactivated to
conceivably reactive metabolites such as the S-oxide and sulfenic
acid.14 Thionamides, thioureas, thiocarbamates, and thiones are
S-oxygenated by FMO1 and FMO3,15,16 to their higher states
S-oxides through sulfenic acid,17 which can react with macro-
molecular cellular proteins, forming protein-electrophile
adducts and subsequently triggering adverse drug reactions.
We report, in this manuscript, on the oxidation of ETA by
peracetic acid. ETA bioactivation studies have been performed
using liver microsomes; without a clear delineation of the
products of this oxidation. Microsomes are usually nonspecific;
with over 50 CYP450 genes and 15 subgenes as well as about ten
Flavin-containing monooxygenases.20−22 Product identification
after bioactivation in such environments is not trivial. Peracetic
acid (and in some cases, persulfate), on the other hand, is a very
“clean” oxidant; it bioactivates via oxygen addition to give acetic
Received: July 22, 2016
Revised: September 28, 2016
© XXXX American Chemical Society
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Scheme 1. Sulfenic Acid Formed Is Unstable with Respect to the S-Oxide
acid and an oxidized substrate; which is exactly what the
confirmed ETA bioactivating monooxygenase does. Oxidation
mechanisms of peracetic acid have been intensely studied and
have been well-characterized. Peracetic acid oxidizes as an
electrophile (as in Br2, HOBr, etc.). The reactivity of the
peracid is determined by the electron withdrawing character
of the substituents. The stronger the parent acid, the more
reactive is the derived peracid. Thus, trifluoroperacetic and
2,4-dinitroperbenzoic acids are stronger oxidants than peracetic
and m-chloroperbenzoic (MCPBA) acids. The peracetic acid we
have utilized in this study is the weakest of all peracetic acids, thus
ensuring a smaller subset of possible oxidation products. The vast
majority of peracetic acids are used in three specific oxidations:
the Baeyer−Villiger oxidation of ketones to esters, oxidation of
alkenes to epoxides and oxidation of heteroatoms to their oxides:
amine to amine oxides, sulfides to sulfoxides, and phosphine
to phosphine oxides.23,24 Peracids are generally inert and do not
oxidize alcohols, esters and ethers. Thus, they are the best way to
mimic bioactivation of ETA by Flavin monooxygenase, EtaA,25
while avoiding complicating issues associated with microsomal
oxidations.
and scanning them spectrophotometrically for depletion of ETA
over periods of up to 24 h in unstirred vessels. Excess peracetic
acid could be determined by adding excess acidified iodide and
back-titrating with standard thiosulfate using freshly prepared
starch as indicator. Mass spectra of oxidation products, were
acquired on a high-resolution (m/Δm = 30 000) Thermo Scientific
LTQ-Orbitrap Discovery mass spectrometer (San Jose, CA)
equipped with an electrospray ionization source. The electrospray
ionization mass spectrometry (ESI-MS) source parameters were
set as follows: spray voltage (kV), 2.5 in negative mode and 4.5 in
positive mode; spray current (μA), 1.96; sheath gas flow rate, 20;
auxiliary gas flow rate, 0.01; capillary voltage (V), −16; capillary
temperature (°C), 300; and tube lens (V), − 115. Detection was
carried out in both the negative ionization mode and positive
(−ESI) . The detection parameters were set up as follows: analyzer;
FTMS, positive and negative polarity; mass range; normal,
resolution; 30 000, scan type; centroid.
RESULTS
■
Stoichiometry. The stoichiometry of this reaction was
complex. Attained stoichiometry was dependent on the time
of incubation as well as on the ratio of oxidant to reductant.
In excess ETA of peracetic acid, a stoichiometry of 1:1 was
quickly established, within 5 s, in which a single oxygen atom
was inserted on to the sulfur center to form a sulfenic acid (R1)
which immediately transformed into the more stable zwitterionic
S-oxide (R2, see Scheme 1).
EXPERIMENTAL SECTION
■
Materials. Reagent grade Ethionamide (ETA), peracetic acid,
sodium chloride, perchloric acid and barium chloride were
purchased from Sigma-Aldrich (USA) and were used without
further purification. Water solutions for reactions were purified
using a Barnstead Sybron Corp. water purification unit capable
of producing both distilled and deionized water (Nanopure).
Inductively coupled plasma mass spectrometry (ICPMS) was
used to evaluate concentrations of metal ions in the reagent
water. ICPMS results showed negligible amounts (<0.1 ppb) of
copper, iron and silver ions with approximately 1.5 ppb of cadmium
and 0.43 ppb in lead as the highest metal ion concentrations.
In previous experiments from our lab, no discernible differences in
kinetics data had been obtained between experiments run with
chelators (EDTA, deferroxamine) and those run without, and so all
experiments were carried out without the use of chelators.
CH3CO3H + Et(C5H4)C(S)NH2 → CH3CO2H
+ Et(C5H4)C(NH)SOH
(R1)
Et(C5H4)C(NH)SOH ↔ Et(C5H4)C(NH2)S → O
(R2)
On prolonged standing, the S-Oxide formed in reaction 2 slowly
decomposes to yield the final product 2-ethylisonicotinamide,
which represents complete desulfurization of ETA (see Figure 1
and Scheme 2).
Figure 2 shows the ESI spectrum of a 1:1 mixture of ETA
and peracetic acid before the reaction proceeds to completion.
It shows a significant peak for the S-oxide as a major oxidation
product. Sulfenic acids and S-oxides are difficult to stabilize in the
absence of bulky groups surrounding the sulfur center. ETA is
one of a few organosulfur compounds that can stabilize the sulfur
center in the S(0) state when not in the polymeric S8 state.
The other major peak observed is the one for the product,
2-ethylisonicotinamide. Thus, the reaction is not discriminating
enough to halt at the S-oxide, which is the equivalent of a
2-electron oxidation.
After the confirmed 1:1 stoichiometry is achieved in the first
5 s of the reaction, the second part of the stoichiometry could not
be conclusively determined (see Figure 1). The most surprising
aspect of this stoichiometry is the lack of sulfate production. With
ionic strength maintained by NaCl, there was no other possibility
of precipitation with barium chloride apart from sulfate. Figure 3
shows the ESI mass spectrum of a reaction solution in which
Methods. The rapid reactions of ethionamide with peracetic
acid were followed on a Hi-Tech Scientific SF61- DX2 double-
mixing stopped-flow spectrophotometer. ETA and peracetic
acid were in separate vessels, but both reagent vessels were
maintained at the requisite ionic strength. Slower reactions
involving the reactions of the S-oxide and decomposition of
reaction products following oxidation of ETA were monitored on
a conventional PerkinElmer Lambda 2S UV−vis spectropho-
tometer. All kinetics experiments were performed at 25.0
0.5 °C and an ionic strength of 1.0 M (NaCl). ETA is sparingly
soluble in water at neutral pH conditions, but its solubility
increased in highly acidic environments. Stock solutions of
ETA were prepared by first dissolving a known sample of ETA in
0.1 M perchloric acid followed by serial dilutions with water
to attain the desired strength. ETA has an intense yellow color.
Stoichiometric determinations were carried out by mixing various
ratios of ETA and peracetic acid in tightly sealed volumetric flasks
B
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Total concentration of ETA and ETAS-O is the initial ETA
concentration, [ETA]0
[ETA]0 = [ETA](t) + [ETAS‐O](t)
(1)
[ETA](t) and [ETAS-O](t) are the instantaneous concen-
tration of ETA and the S-oxide, respectively. If we assume that ε1
and ε2 are the absorptivity coefficients of ETA and ETAS-O
respectively, then instantaneous absorbance at any time, t is
given by
A(t) = ε1[ETA]0 + [ETAS‐O](t){ε2 − ε1}
(2)
Equation 2 is correct, irrespective of whether ε1 > ε2 or
vice versa. From eq 2, we can write a differential equation
for the rate of change of absorbance with respect to time
which can be qualified by the appropriate absorptivity co-
efficients to find rate of depletion/formation of ETA/ETAS-O.
This differential equation can be integrated to evaluate the
instantaneous concentrations of ETA or ETAS-O in terms
of the observed overall absorbance. This gives the following
analytical solution of instantaneous concentration of ETAS-O
at any time with respect to the experimentally observed
absorbance:
Figure 1. Absorbance trace for the reaction at 380 nm showing the
biphasic nature of the reaction. The first phase of the reaction, showing
formation of the S-oxide, ETAS-O, is over in 5 s. The second, slower
reaction takes up to an hour for completion. This section involves
further oxidation of the S-Oxide to yield the desulfurized analogue of
ETA, 2-ethylisonicotinamide.
peracetic acid is in high excess over ETA, taken after pro-
longed standing to ensure complete consumption of ETA.
The spectrum shows a single strong peak for the product,
2-ethylisonicotinamide at m/z = 151.08682. In a previous study,
a similar compound, thionicotinamide, was oxidized by peracetic
acid only as far as the sulfinic acid. Only a trace amount of the
sulfinic acid of ETA is observed in this spectrum at m/z =
191.16957. The expected sulfate peak, even trace amounts, was
not observed. Barium chloride did not afford BaSO4 precip-
itation, either.
Kinetics. The kinetics of the reaction were complicated by the
biphasic nature of the reaction as well as the fact that both ETA
and the S-oxide absorb strongly at the wavelength of observation,
380 nm. The absorptivity coefficients of ETA and its S-oxide,
ETAS-O at 380 nm are 656 and 4737 M−1 s−1 respectively.
Figure 4 shows spectra of ETA and ETAS-O superimposed. Even
though absorptivity of ETA is much lower than that of ETAS-O,
it cannot be ignored since it is the starting material and all
observed absorbance at the commencement of the reaction is
derived solely from it. As the reaction proceeds, the observed
increase in absorbance is due to a combination of depletion of
ETA and formation of ETAS-O. An increase in absorbance is
observed because the absorptivity coefficient of the product
exceeds that of the reactant. Ultimately, the effective absorptivity
coefficient is the difference in these two values, but to obtain
unambiguous values of initial rates and rate constants for the
reaction, absorbance data had to be deconvoluted and displayed
with respect to each component singly and solely on its
concentration variation based on Beer Law. To this effect, we are
aided by the mass balance equation that says that ETA is either in
its native form or in the S-Oxide form. The derivation is set up as
follows:
A(t) − ε1[ETA]0
[ETAS‐O](t) =
ε2 − ε1
(3)
The instantaneous concentrations of ETA can thus be
derived from eqs 1 and 3. The only variables in eq 3 are the
concentration of ETAS-O and the observed absorbance.
Since the absorbance value is known at any instantaneous
time value, the corresponding concentration and absorbance
can be easily derived. Each absorbance−time plot could then be
deconvoluted into two separate plots; one representing the
ETAS-O absorbance/concentrations, and the other ETA. Initial
rates can thus be unambiguously determined. The same kinetics
constants were obtained if pseudo-first-order kinetics were
assumed and rate constants determined from half-lives. Figure 5
shows the multiple scans for the ETA−peracetic acid reaction.
The peak at 380 nm is quickly established, and, since the scans
were taken every 120 s, Figure 5 only shows the decrease of the
peak at 380 nm.
Figure 6a shows the raw experimental absorbance data
for the reaction while varying initial concentrations of ETA.
The nonzero initial absorbance represents contribution from
the initial concentrations of ETA. Using this raw experimental
data would deliver an erroneously low rate of reaction and rate
constant. Qualitatively, the reaction shows first order kinetics
in [ETA]0. The deconvoluted absorbance data, separating the
absorbance contribution from ETAS-O, is obtained by applying
eq 3 to every data point. This can be accomplished easily in
Microsoft Excel to produce two new absorbance−time spread-
sheets: one for ETAS-O and the other for ETA. Figure 6b
shows the deconvoluted absorbance data derived from the
ETA concentrations alone. It shows the expected decay in ETA
a
Scheme 2. Oxidation Scheme of ETA to Final Product
a
Only the S-oxide is observed.
C
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Figure 2. ESI mass spectrum of a 1:1 mixture of peracetic acid and ETA before reaction goes to full completion. It shows a strong peak for the substrate,
as expected, at m/z = 167.06389. The S-oxide shows up at m/z = 183.05898. The desulfurized product, the amide, 2-ethylisonicotinamide is at m/z =
151.08602. No higher oxidation states of sulfur are observed.
concentrations as ETAS-O is concomitantly formed. Rate of
formation of ETAS-O and rate of consumption of ETA are
equivalent, based on the mass balance relation, eq 1.
a decreased rate of reaction with acid and a lower product
formation.
Mechanism. The mechanism is complicated by the lack of
sulfate production and the lack of any other oxidation products
apart from the S-oxide. Figure 9 shows spectrum of a reaction
mixture with excess oxidant and before full completion. It
showed no sulfinic nor sulfonic acids, but nearly quantitative
formation of the S-oxide before its slow decomposition to the
amide. At this moment, only a small peak is observed for the
amide.
In Figure 7a, peracetic acid was varied while it was still
maintained in high excess over ETA. Figure 7b shows the
isolation of the consumption of ETA. Again, due to the mass
balance equation, these two plots mirror each other. As expected,
pseudo-first order kinetics were observed. To determine reaction
order with respect to peracetic acid, the variation of the observed
or apparent rate constant with respect to initial peracetic acid
concentration was evaluated.
These experimental results show that oxidation of ETA does
not proceed past the S-oxide. The cleavage of the C−S bond in
ETA occurs at ETAS-O. The leaving group is the radical species
SO, sulfur monoxide. Concomitantly, with the extrusion of
the SO moiety, is the formation of the final organic product,
2-ethylisonicotinamide. Through such well-known synthesis
reactions such as the Baeyer−Villiger reaction,23 the mechanism
of peracid oxidations is well-known. The initial oxidation is
an electrophilic attack by peracetic acid on the carbon−sulfur
double bond. Peracids are electrophilic at the oxygen next to the
acidic proton due to a strong H-bond interaction between the
acidic proton and the carbonyl group of the carboxylic acid. Thus,
initial oxidation of ETA involves oxygen atom addition to the
nucleophilic sulfur atom via an initial epoxidation. In Scheme 3,
the oxygen atom that eventually ends up on the S-oxide is the one
next to the acidic proton.
Rate = k[PAA]x[ETA]
We can then set kapp = k[PAA]x, where kapp is the “apparent”
rate constant observed. A plot of kapp versus [PAA]0 gave a
straight line. This means the order with respect to peracetic
acid is 1.
(4)
We thus determined that
Rate = −d[ETA]/dt = k[ETA][PAA]
with k = (3.08 0.72) × 102 M−1 s−1.
Effect of Acid. Since ETA can only dissolve in an acidic
environment, at any time, we never ran the reaction in higher
than pH 3.0 even if no further acid was added. Figure 8
shows the inhibitory effect of acid on the reaction. One observes
(5)
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Figure 3. ESI mass spectrum obtained in high excess of peracetic acid over ETA. Only two viable peaks are observed: one for product
2-ethylisonicotinamide, and the other for the unreacted peracetic acid (not shown on this scale).
Oxidation of the thioamide bond by peracetic acid leading to
the formation of the S-Oxide has been studied by our lab.26 This
oxidation proceeds via epoxidation and subsequent hydrolysis
and rearrangement from the sulfenic acid to the S-oxide. Further
research from our lab showed that the zwitterionic form of the
sulfinic acid was vulnerable to cleavage of the carbon−sulfur bond,
releasing a sulfoxyl anion radical that dimerizes to form dithionite.27
Dithionite is distinguished by its characteristic absorbance at
315 nm as well as m/z = 129 ESI peak in the negative mode.
Dithionite is also easily oxidized, even by atmospheric oxygen,
to sulfate. None of these dithionite spectral characteristics were
observed, and, since sulfate was not observed, it appears that
cleavage of the carbon sulfur bond was effected from the S-Oxide.
ESI spectra were obtained which focused on m/z fragments
between 60 and 120 in order to determine the final sulfur species.
Figure 10 shows that the major and predominant peak in this
region was at m/z = 98.95501. Though this is close to sulfate,
we were aware that is was not sulfate since no precipitate was
observed with barium sulfate. This peak is attributable to the
Figure 4. Spectra of (a) [ETA] = 0.0001 M, (b) product of [ETA] =
0.0001 M and per acetic acid [PAA] = 0.00001 M, and (c) [PAA] =
0.0001 M.
dimeric sulfur monoxide species, H2S2O2(aq). Sulfur monoxide,
SO, and its dimer, S2O2, have been observed in low tempera-
tures in the gas phase.28,29 Sulfur monoxide has also been
formed and detected in the activation of frustrated Lewis pair
of N-sulfinyamine.30 Thus, mechanistically, the oxidation of the
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Figure 5. Multiple scan spectra showing reaction of [ETA] = 0.0001 M,
[PAA] = 0.005 M. Spectra acquired every 2 min. This shows a slow
decrease in the peak at 380 nm.
Figure 7. (a) Variation of CH3O3H in ETA oxidation by per acetic acid.
Fixed [ETA] = 0.0001 M varied [CH3O3H] = (a) 0.00020 M, (b)
0.00030 M (c) 0.00045 M (d) 0,00055 M, (e) 0.00065 M. (b) Variation
of [CH3O3H] showing consumption of ETA. The conditions are the
same as in panel a.
Figure 6. (a) Variation of ETA in its oxidation by per acetic acid showing
total absorbance from ETA and ETAS-O at 380 nm. [CH3CO3H] =
0.005 M, [H+] = 0.0001 M, INaCl = 1.0 M ; and varied [ETA]0
(a) = 0.00005 M, (b) 0.00010 M, (c) 0.00015 M, (d) 0.00020 M
(e) 0.00025 M. (b) Concentration variations of ETA in its oxidation by
per acetic acid showing appearance of ETAS-O at 380 nm. These traces
are obtained from subtracting at each data point the absorbance in panel
b from that in panel a.
Figure 8. Variation of [HClO4]. Fixed [ETA] = 0.0001 M, [PAA] =
0.005 M. Varied [HClO4] (a) 0.1 (b) 0.25 M, (c) 0.5 M, (d) 0.75 M,
(e) 1.0 M. Traces a−e are sequentially from top to bottom.
F
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Figure 9. ESI(+) spectrum of a 1:3 mixture of ETA to peracetic acid after 5 min and before full completion of the reaction. The strongest peak is for the
S-oxide and the unreacted ETA. Very little of the amide product, at m/z = 151.08658 is observed at this time.
Scheme 3. Further Oxidation of the S-Oxide through Epoxidation to Form the Product and Release a Reactive Leaving Species
That Further Dimerizes to Give the Major Sulfur Species
S-oxide is accompanied by the formation sulfur monoxide and
the product amide (Scheme 3). Peracetic acid, a weak oxidant
that oxidizes by addition of oxygen, is inert to sulfur monoxide
and its dimers. Figure 10 also shows a strong peak at m/z =
82.9541, which is exactly 16 mass units lower than the pre-
dominant peak. This is attributable to the S2O species which
could be derived from the reaction of SO and elemental sulfur in
S6 or S7 state before formation of the S8 state, which is inert.
Possible sulfur-based radicals, if present could be detected by
EPR. Our X-Band EPR spectrometer could not detect any radical
species. For these types of radicals, one needs a Q-band EPR
spectrometer which was unavailable.
production of the amide product without observation of a sulfinic
acid indicates that formation of the product is after the formation
of the S-Oxide.
Effect of Acid. Figure 8 shows the inhibitory effect of acid on
this reaction. Acid reduces the rate of reaction and amount of
S-Oxide formed before further desulfurization of the thioamide.
The observed acid retardation occurs at very high acid
conditions, at pH values between 1 and 0. At these highly acidic
conditions, the thiocarbonyl group of ETA will be come
appreciably protonated. The proton is known to reside between
the amine and the sulfur center, but is more biased more toward
the more nucleophilic sulfur center. Representing ETA as
R(NH2)CS, we can then write the following equilibrium:
Scheme 3 shows the major mechanistic pathway for the oxida-
tion of ETA. Certainly there are other minor pathways, from
observation of the S2O species. The observation of almost instant
R(NH2)CS + H+ ⇄ [R(NH2)CS − H]+ Kb
(R3)
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Figure 10. ESI mass spectra focusing on the sulfur species. Dimeric sulfur monoxide species is the dominant sulfur species. No sulfate was detected.
Since this is an electrophilic attack by PAA on ETA, we can
consider the protonated amide to be inert, and if it does react,
that this rate is exceedingly small, compared to the unprotonated
thioamide. We can thus derive a more accurate rate law (eq 5) of
the form
ACKNOWLEDGMENTS
■
This work was funded by Grant Number CHE 1056366 from the
National Science Foundation and a partial research professor
vote from the University of KwaZulu-Natal.
Rate = −d[ETA]/dt = k[ETA]0[PAA]0 /(1 + Kb−1[H+])
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Equation 6 clearly quantifies observed acid retardation effect.
The second term in the denominator is negligible at low acid
concentration, and the rate reverts to eq 5. In our reaction
system, acid is always present, from the PAA reactant itself, and
from the initial acid amounts used to dissolve PAA. Thus, the
derived rate constant represents an upper limit value in the limit
very low acid concentrations.
CONCLUSIONS
■
This short kinetics study shows that ethionamide is oxidized by
peracetic acid almost quantitatively to the S-oxide. Ethionamide,
itself, is a prodrug. It is known to be metabolized to the S-oxide
before interacting with its cellular target. Thus, peracetic acid can
adequately mimic the microsomal bioactivation of ethionamide.
AUTHOR INFORMATION
■
Corresponding Author
*E-mail address: rsimoyi@pdx.edu; phone number: 503-725-
3895.
(7) Vale, N.; Gomes, P.; Santos, H. A. Metabolism of the
Antituberculosis Drug Ethionamide. Curr. Drug Metab. 2013, 14,
151−158.
(8) Qian, L.; Ortiz de Montellano, P. R. Oxidative Activation of
Thiacetazone by the Mycobacterium Tuberculosis Flavin Monoox-
Notes
The authors declare no competing financial interest.
H
DOI: 10.1021/acs.jpca.6b07375
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DOI: 10.1021/acs.jpca.6b07375
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