Oxidation of Symmetric Disulfides with Hydrogen Peroxide Catalyzed by Methyltrioxorhenium(VII)

Article abstract of DOI:10.1021/jo991109w

Organic disulfides with both alkyl and aryl substituents are oxidized by hydrogen peroxide when CH₃ReO₃ (MTO) is used as a catalyst. The first step of the reaction is complete usually in about an hour, at which point the thiosulfinate, RS(O)SR, can be detected in nearly quantitative yield. The thiosulfinate is then converted, also by MTO-catalyzed oxidation under these conditions, to the thiosulfonate and, over long periods, to sulfonic acids, RSO₃H. In the absence of excess peroxide, RS(O)SR (R = p-tolyl), underwent disproportionation to RS(O)₂SR and RSSR. Kinetics studies of the first oxidation reaction established that two peroxorhenium compounds are the active forms of the catalyst, CH₃ReO₂(η²-O₂) (A) and CH₃ReO(η²-O₂)₂·(OH ₂) (B). Their reactivities are similar; typical rate constants (L mol^-1 s^-1, 25°C, aqueous acetonitrile) are k_A = 22, k_B = 150 (Bu₂S₂) and k_A = 1.4, k_B = 11 (Tol₂S₂). An analysis of the data for (P-XC₆H₄)₂S₂ by a plot of log k_B against the Hammett σ constant gave ρ = -1.89, supporting a mechanism in which the electron-rich sulfur attacks a peroxo oxygen of intermediates A and B.

Full text of DOI:10.1021/jo991109w

104
J . Org. Chem. 2000, 65, 104-107
Oxid a tion of Sym m etr ic Disu lfid es w ith Hyd r ogen P er oxid e
Ca ta lyzed by Meth yltr ioxor h en iu m (VII)
Ying Wang and J ames H. Espenson*
Ames Laboratory and Department of Chemistry, Iowa State University of Science and Technology,
Ames, Iowa 50011
Received J uly 12, 1999
Organic disulfides with both alkyl and aryl substituents are oxidized by hydrogen peroxide when
CH3ReO3 (MTO) is used as a catalyst. The first step of the reaction is complete usually in about an
hour, at which point the thiosulfinate, RS(O)SR, can be detected in nearly quantitative yield. The
thiosulfinate is then converted, also by MTO-catalyzed oxidation under these conditions, to the
thiosulfonate and, over long periods, to sulfonic acids, RSO3H. In the absence of excess peroxide,
RS(O)SR (R ) p-tolyl), underwent disproportionation to RS(O)2SR and RSSR. Kinetics studies of
the first oxidation reaction established that two peroxorhenium compounds are the active forms of
the catalyst, CH3ReO2(η2-O2) (A) and CH3ReO(η2-O2)2‚(OH2) (B). Their reactivities are similar; typical
rate constants (L mol-1 s-1, 25 °C, aqueous acetonitrile) are kA ) 22, kB ) 150 (Bu2S2) and kA ) 1.4,
kB ) 11 (Tol2S2). An analysis of the data for (p-XC6H4)2S2 by a plot of log kB against the Hammett
σ constant gave F ) -1.89, supporting a mechanism in which the electron-rich sulfur attacks a
peroxo oxygen of intermediates A and B.
Ch a r t 1. Disu lfid e a n d Oxid a tion P r od u cts; R )
In tr od u ction
Alk yl, Ar yl
Oxidation of the disulfide functional group (1, Chart
1) has been studied quite extensively because of its
important role in metabolism, via a variety of mechanistic
paths.1 With the use of different oxidants and well-
controlled situations, various intermediate compounds
have been obtained.2-6 Among these are thiosulfinates,
2, which can provide a sulfinyl group that acts as a
removable source of diasteroselectivity in asymmetric
syntheses.7 Similarly, such reactions can produce thio-
sulfonates, 3, which are powerful sulfenylating agents
useful for the temporary blocking of mercapto groups in
protein chemistry.8
Peracids are the reagents most commonly used to
oxidize disulfides, although hydrogen peroxide in acetic
acid also finds application. The hydrogen peroxide reac-
tions, in the absence of a catalyst, often require a large
excess and a higher temperature.1,3 Oxometal complexes
are catalysts for the oxidation; molybdates and tungstates
are effective, especially under phase-transfer conditions.9-11
Methyltrioxorhenium, CH3ReO3 or MTO, has been
widely used as a catalyst for reactions that use hydrogen
peroxide as a stoichiometric oxidant.12-15 One can antici-
pate that either or both of the peroxorhenium species that
exist in an equilibrium or a steady-state relation with
MTO, as shown in Scheme 1, will be the active form of
the catalyst. Since a sequence of oxidation steps will
occur, the experiments were carried out. The principal
objective of this research is the study of the reaction
kinetics, of the first stage in particular, as written in eq
1.
A/B + RSSR f MTO/A + RS(O)SR + H2O (1)
The first reaction proved to be faster than the subse-
quent oxidations. This allowed, when hydrogen peroxide
was taken in only 1:1 ratio to the disulfide, the detection
of the thiosulfinates, otherwise difficult to realize owing
to their transient nature. We also report, qualitatively,
on the formation of thiosulfonates and sulfonic acids.
(1) Organic Sulfur Chemistry: Structure and Mechanism; Oae, S.,
Ed.; CRC Press: Boca Raton, 1991; Vol. 1, pp 213-216.
(2) Bhattacharya, A. K.; Hortman, A. G. J . Org. Chem. 1978, 43,
2728.
Exp er im en ta l Section
(3) Allen, P., J r; Brook, J . W. J . Org. Chem. 1962, 27, 1019.
(4) Gu, D.; Harpp, D. N. Tetrahedron Lett. 1993, 34, 67.
(5) Folkins, P. L.; Harpp, D. N. J . Am. Chem. Soc. 1993, 115, 3066.
(6) Freeman, F. Chem. Rev. 1984, 84, 117-135.
(7) Cogan, D. A.; Liu, G.; Kim, K.; Backes, B. J .; Ellman, J . A. J .
Am. Chem. Soc. 1998, 120, 8011.
Rea gen ts. Aqueous acetonitrile was used as the reaction
solvent, the water being purified by a Millipore-Q water
purification system. The disulfides were used as received from
Aldrich, save for p-chlorophenyl disulfide, which was prepared
(8) Trost, B. M. Chem. Rev. 1978, 78, 8.
(9) Anisimov, A. V.; Mokhammad, R.; Akhmad, S. A.; Borisenkova,
A. V.; Tarakanova, S. V.; Barkhanova, S. V.; Dergacheva, V. M.;
Negrimovskii, V. M.; Kaliya, O. L.; Luk’yanets, E. A. Neftekhimiya
1994, 34, 358.
(10) Anisimov, A. V.; Mokhammad, R. A.; Tarakanova, A. V.;
Borisenkova, S. A. Neftekhimiya 1994, 34, 421.
(12) Espenson, J . H.; Abu-Omar, M. M. Adv. Chem. Ser. 1997, 253,
99-134.
(13) Gable, K. P. Adv. Organomet. Chem. 1997, 41, 127-161.
(14) Herrmann, W. A.; Ku¨hn, F. E. Acc. Chem. Res. 1997, 30, 169-
180.
(15) Espenson, J . H. J . Chem. Soc., Chem. Commun. 1999, 479-
488.
(11) Singh, P. K.; Field, L. J . J . Org. Chem. 1988, 53, 2608.
10.1021/jo991109w CCC: $19.00 © 2000 American Chemical Society
Published on Web 12/09/1999

Hydrogen Peroxide Oxidation of Symmetric Disulfides
J . Org. Chem., Vol. 65, No. 1, 2000 105
Ta ble 1. Ra te Con sta n ts a t 298 K for th e Oxid a tion of
Disu lfid es by P er oxor h en iu m Com p lexes A (k3) a n d B
(k4) in Aqu eou s Aceton itr ilea -c Con ta in in g 0.100 M
CF 3SO3H
Sch em e 1
by the literature method16 and purified by recrystallization
from ethanol-water, and for tert-butyl disulfide, which was
vacuum distilled.
1
Kin etics. Both UV/visible and H NMR methods were used.
Quantitative measurements were made at 25 °C in 4:1 v/v
acetonitrile-water at pH 1 (to stabilize MTO in peroxide-
containing media)17 maintained by trifluoromethanesulfonic
acid. Owing to the large molar absorptivities of the diaryl
disulfides, it was often necessary to use cuvettes of 0.01-0.05
cm optical path. One or more wavelengths were chosen for each
compound to allow quantitative monitoring of the reaction
progress. The conditions and wavelengths selected were such
that the loss of ArSSAr was followed, although at other
wavelengths not used for kinetics, a rise in absorbance
accompanying the buildup of 2 could be seen. Under the
conditions of the spectrophotometric determinations, princi-
pally that [ArSSAr] was quite low and [H2O2] effectively
constant, the reactions followed pseudo-first-order kinetics.
The absorbance-time data from each experiment were ana-
lyzed according to eq 2, from which a value of kψ was obtained
literature values.7,19-21 Adding excess hydrogen peroxide into
the above solution eventually afforded the final product,
p-toluenesulfonic acid. The reaction was also monitored by 1H
NMR, during which both the thiosulfinate 2 and thiosulfonate
3 were detected. No R,R′-disulfonate 4 was detected for any of
the substrates used in this study.
Abst ) Abs∞ + (Abs0 - Abs∞) exp-k
(2)
ψt
by nonlinear least-squares fitting. In the case of dialkyl
disulfides, where there are no useful UV absorptions, the
reaction kinetics was evaluated from NMR intensities. These
values were converted by computer program18 to concentra-
tions based on the total intensity of the starting material.
Certain data sets were analyzed by first-order kinetics accord-
ing to eq 3. That was allowable at high concentrations of
Resu lts
P r elim in a r y Exp er im en ts. No significant interaction
was found between the disulfides and MTO on the basis
of 1H NMR results. This finding implies that the reaction
does not proceed by way of a prior complex between these
two. The result further suggests that the reaction, when
it does occur as hydrogen peroxide is added, does not
feature attack of the sulfur at the rhenium center, for
there is no reason to believe that this interaction would
occur then, when it did not do so on its own. All six
substrates (see Table 1) were initially examined in the
absence of a catalyst. Without MTO, there was no
evidence for a reaction, even with excess hydrogen
peroxide, as indicated either by a UV spectrum that
remained nearly unchanged for several hours or by the
absence of new peaks growing in the NMR spectrum.
Once MTO had been added, however, the spectra
immediately began to show the buildup of the product.
These determinations were also performed with equal
initial concentrations of the disulfide and hydrogen
peroxide. A certain small amount of the thiosulfonate was
detected along with the major thiosulfinate for R ) Ph,
p-tolyl, and p-chlorophenyl. With three of the compounds,
ψt
Ct ) C0 e-k
(3)
hydrogen peroxide, where (see later) the rate becomes peroxide-
independent. However, at lower peroxide concentrations the
data conform to the Michaelis-Menten form for catalytic
reactions, and the data cannot be analyzed by an integrated
rate law. In these circumstances the method of initial rates
was used. To determine the value of the initial rate, vi, the
concentration was expressed as a polynomial function (eq 4)
Ct ) C0 - a1‚t - a2‚t2 - a3‚t3 - ...
(4)
by least-squares fitting, from which it can be seen that vi )
a1. For one compound, with R ) p-tolyl, it was confirmed that
both UV and NMR methods gave the same ultimate parameter
value.
P r od u cts. The procedure used to identify the reaction
products was based on chromatography and spectroscopy. A
solution containing MTO (2.5 µmol) and hydrogen peroxide (1
mmol) was slowly added to 1 mmol p-ditolyl disulfide (for
example) in 5.0 mL of acetonitrile. The reaction was allowed
to run for 2 h, during which time its progress was monitored
periodically by TLC. When this test showed that the reaction
progress had nearly stopped, the mixture was separated by
preparative TLC, using cyclohexanes-ethyl acetate (95:5) as
the eluting agent. This gave p-tolyl p-toluenethiosulfinate, 2,
as the major product, along with a barely detectable amount
of p-tolyl p-toluenethiosulfonate, 3. These products were
identified either by 1H NMR and GC-MS techniques, in
comparison with the values of the authentic compounds, or
t
those with R ) Me, Bu, and p-nitrophenyl, the thiosul-
finate was very nearly the only product. Isosbestic points
were maintained during these reactions, even with a
small excess of hydrogen peroxide. Such a series of
repetitive scans is shown in Figure 1.
For this second group of disulfides, the further steps
of oxidation beyond the thiosulfinate are a great deal
slower, and for all of the disulfides studied there is a
(18) We are grateful to Dr. M. Englehardt of Bruker Corp. for
supplying this program.
(16) Overman, L. E.; Matzinger, D.; O’Connor, E. M.; Overman, J .
D. J . Am. Chem. Soc. 1974, 96, 6089.
(17) Abu-Omar, M.; Hansen, P. J .; Espenson, J . H. J . Am. Chem.
Soc. 1996, 118, 4966-4974.
(19) Kozzuka, S.; Takahashi, H.; Oae, S. Bull. Chem. Soc. J pn. 1970,
43, 129 and references therein.
(20) Chemla, F. Synlett 1998, 894.
(21) Sas, W. J . Chem. Res. 1993, 160.

106 J . Org. Chem., Vol. 65, No. 1, 2000
Wang and Espenson
F igu r e 2. Variation of initial rate of reaction of dimethyl
disulfide with the concentration of hydrogen peroxide. The
curve is a fit to eq 5.This is applicable when [B] is negligible
compared to [A]. The values of k1 and k-1 (Scheme 1) were
fixed at 15.5 L mol-1 s-1 and 0.17 s-1. Conditions used: 0.5
mM MTO, 5 mM Me2S2 in 4:1CH3CN-H2O at 25.0 °C and pH
1.
F igu r e 1. Spectral changes at 5.0 min intervals for the
oxidation of p-nitrophenyl disulfide (0.5 mM) with H2O2 (0.65
M) in the presence of 0.1 mM MTO at pH 1 in 4:1 CH3CN-
H2O.
data analysis, based on the same chemical mechanism,
considerable rate advantage found in the first stage of
oxidation. This information was used to plan the quan-
titative studies of the kinetics, so that only the first
oxidation stage was important during the data acquisi-
tion time.
was used for di-tert-butyl disulfide. The concentration-
t
time data for all of the experiments with Bu2S2 were fit
globally with the use of the program FitSim.23
At high hydrogen peroxide, when [B] . [A], the
absorbance-time (UV-vis) or concentration-time (NMR)
data fit first-order kinetics. Because [B] = [Re]T, the rate
law under this condition becomes
Kin etics Stu d ies. The first part of understanding the
reaction scheme lies in the two kinetic steps shown in
Scheme 1. In this solvent system, these are the relevant
parameters at 298 K: k1 ) 15.5 L mol-1 s-1, k-1 ) 0.17
v ) k4[Re]T‚[R2S2] ) kψ[R2S2]
(6)
s-1, K1 ) 91 L mol-1, K2 ) 347 L mol-1 22
It is useful to
.
differentiate studies at “low” and “high” concentrations
of hydrogen peroxide, by which we mean experiments at
<7 mM and >0.5 M, respectively. We start by presuming
that A and B are separate participants in the oxidation
and denote their respective rate constants as k3 and k4.
We further presumed, and later confirmed for two
compounds, that k3 and k4 are of the same order of
magnitude. If so, then the term in the rate law for the
reaction of B with disulfide makes e10% contribution in
experiments with “low” concentrations of hydrogen per-
oxide. Under those conditions, the reaction proceeds
largely with the A form of the catalyst, and its velocity
is given by
For these determinations, a series of experiments was
carried out with constant [H2O2] and [R2S2] in which [Re]T
was varied, typically 0.1-0.5 mM. The plot of kψ against
[Re]T was linear, as depicted for three compounds in
Figure 3. The values of k4 given by the slopes of these
lines are summarized in Table 1.
It should be noted that the various aspects of the data
analysis presented are all different components of the
same reaction scheme. The differences arise from the
different ranges of concentration chosen for different sets
of experiments and also because the relative rate con-
stants are different from one disulfide to the next.
F a te of th e Th iosu lfin a tes. Compounds of the for-
mula RS(O)SR tend to be unstable and usually cannot
be isolated. Careful NMR measurements were made
during the reaction between 20 mM di(p-tolyl) disulfide
and a limited (10 mM) concentration of hydrogen perox-
ide. Once the thiosulfinate had been formed, it slowly
decomposed with the formation of the thiosulfonate, RS-
(O)2SR, and partial regeneration of the disulfide. This is
compatible with a disproportionation process:
k1k3[Re]T‚[R2S2]‚[H2O2]
v )
(5)
k-1 + k3[R2S2] + k1[H2O2]
For purposes of data analysis, the values of k1 and k-1
were set at their known values.22 With that, the only
unknown in eq 5 is k3. Its value could be determined for
each experiment, whether the data were in the form of
vi or kψ. Under those conditions the variation in the initial
rate with hydrogen peroxide concentration, at fixed
concentrations of the disulfide and MTO, defines a
rectangular hyperbola. Data from experiments with
Me2S2 are presented in Figure 2. A second method for
2RS(O)SR f RS(O)2SR + RSSR
(7)
A companion experiment with di(tert-butyl) disulfide gave
a thiosulfinate that seemed stable for a much longer
(22) Huang, R.; Espenson, J . H. J . Org. Chem. 1999, 64, 6374.
(23) Zimmerle, C. T.; Frieden, C. Biochem. J . 1989, 258, 381-387.

Hydrogen Peroxide Oxidation of Symmetric Disulfides
J . Org. Chem., Vol. 65, No. 1, 2000 107
sulfur atom) shows a buildup of positive charge relative
to that in the free molecule. This effect, although in the
same direction, is considerably larger than that found
for RSAr (F ) -0.98).24 We argue that the enhanced F
value arises principally from the conjugative effect caused
by the para-conjugated phenylmercapto group. In this
case, the electronic effect of this group seems comparable
to that of a phenyl group. The electronic effect of the para
substituent on the sulfur atom that is not being oxidized
appears muted.
When these results are combined with the information
already available about organic sulfides24 and thioke-
tones,22 comparisons can be made to gain an understand-
ing of these results. The disulfides are some 103-104 less
reactive than either of those families of compounds. This
comparison does not require a precise one-to-one match
of substituents, because the rate effects are quite large.
This is one series that is reasonable for comparison: for
Et2S, thiocamphor, and Me2S2, values of k3 are 2 × 104,
3 × 103, and 6.4 L mol-1 s-1. This large difference can be
ascribed to the substitution of an SR group for an R
group. Despite the electron-releasing conjugative effect,
the SMe group is electron-withdrawing; its Hammett σ
value is +0.07 and the Taft σ1 value is +0.23. This effect
may lower the nucleophilicity of the sulfur atom to which
it is bound in the disulfide, thus lowering, relative to the
sulfides and thioketones, the rate at which it attacks the
peroxorhenium compound.
F igu r e 3. The pseudo-first-order rate constants (4:1 CH3CN-
H2O at 25.0 °C and pH 1) for the oxidation of diaryl disulfides
(0.5 mM) by H2O2 (0.65 M) vary linearly with the total catalyst
concentration, as shown in order of increasing slope for the
para substituents NO2, Cl, and H. The slope of this line affords
the values of k4.
The reactivity of B as compared to A is given by the
ratio k4/k3. With but the one exception of allyl alcohols,25
clearly a special case, the reactivity of these two are
comparable and often in favor of A. In the two cases
examined here (and by extrapolation, likely all six) B is
somewhat more reactive than A by a factor of ca. 7.
From the results obtained here and the general pattern
set by other reagents with the MTO-H2O2 system, we
propose a transition state in which one sulfur atom of
the substrate nucleophically attacks the peroxo group of
A or B. Probing more closely, we note that certain
reactions show virtually no steric effect, such as observed
for R2S2: the values of k3 are 6.4 (R ) Me) and 5.6 (R )
tBu) L mol-1 s-1. This insensitivity of rate to steric bulk
is in accord with values for PAr3,26 R2S,24 and Br-.27 In
contrast, the epoxidation of olefins show an important
steric effect.28 Oxidation of MeSStBu by peracetic acid
forms MeS(O)StBu and MeSS(O)tBu in 1:2 yield,29 show-
ing that these reactions also have small steric effects.
F igu r e 4. Hammett plot for the oxidation of substituted
phenyl disulfides by the peroxorhenium compound B. The
slope of the line gives F ) -1.89.
Ack n ow led gm en t. This research was supported by
the U. S. Department of Energy, Office of Basic Energy
Sciences, Division of Chemical Sciences under contract
W-7405-Eng-82.
period, suggesting that disproportionation is sterically
impeded by the bulky R group.
J O991109W
Discu ssion
To assess the extent to which electronic factors play a
role in the transition state, the rate constants for the
diaryl disulfides were examined by the Hammett cor-
relation. The plot of log(k4) vs σ, the Hammett substituent
constant, was constructed (Figure 4. The four points
define a reasonable straight line; its slope is the reaction
constant, F ) -1.89 (correlation coefficient 0.997). The
negative value of F implies that the reaction center (the
(24) Vassell, K. A.; Espenson, J . H. Inorg. Chem. 1994, 33, 5491.
(25) Tetzlaff, H. A. R.; Espenson, J . H. Inorg. Chem. 1999, 38, 881-
885.
(26) Abu-Omar, M. M.; Espenson, J . H. J . Am. Chem. Soc. 1995,
117, 272.
(27) Espenson, J . H.; Pestovsky, O.; Huston, P.; Staudt, S. J . Am.
Chem. Soc. 1994, 116, 2869.
(28) Al-Ajlouni, A.; Espenson, J . H. J . Org. Chem. 1996, 61, 3969-
3976.
(29) Block, E.; O’Connor, J . J . J . Am. Chem. Soc. 1974, 96, 3921.