Borax-catalyzed and pH-Controlled selective oxidation of organic sulfides by H2O2: An environmentally clean protocol

Article abstract of DOI:10.1002/ejoc.200900265

The selective oxidation of sulfides to sulfoxides and sulfones was achieved in high yields at: room temperature with borax as a recyclable catalyst and H₂O₂ as the terminal oxidant by varying the pH of the reaction medium. The borax/H₂O₂ system can chemoselectively oxidize alkyl and aryl sulfides in the presence of oxidation-prone functional, groups such as C=C, -CN, and -OH.

Full text of DOI:10.1002/ejoc.200900265

SHORT COMMUNICATION
DOI: 10.1002/ejoc.200900265
Borax-Catalyzed and pH-Controlled Selective Oxidation of Organic Sulfides
by H2O2: An Environmentally Clean Protocol
Sahid Hussain,[a,b] Saitanya K. Bharadwaj,[a] Ravindra Pandey,[a] and
Mihir K. Chaudhuri*[a,c]
Keywords: Sustainable chemistry / Acidity / Chemoselectivity / Oxidation / Sulfur
The selective oxidation of sulfides to sulfoxides and sulfones
was achieved in high yields at room temperature with borax
as a recyclable catalyst and H2O2 as the terminal oxidant by
varying the pH of the reaction medium. The borax/H2O2 sys-
tem can chemoselectively oxidize alkyl and aryl sulfides in
the presence of oxidation-prone functional groups such as
C=C, –CN, and –OH.
(© Wiley-VCH Verlag GmbH & Co. KGaA, 69451 Weinheim,
Germany, 2009)
charge on the sulfur atom in the transition state with the
sulfide behaving as nucleophile, yet there is a succinct differ-
ence. The difference lies in mono and diperoxborates having
the transition state earlier along the reaction coordinate[2]
with relatively less positive charge being developed on the
sulfur atom. This in turn emphasizes that proton transfer
would play a significant role in the peroxoborate oxidation
of organic sulfides. Indeed, it was proposed[2] that the lower
extent of positive charge development on the sulfur atom
for the reaction of peroxoborates was due to the greater
importance of the proton transfer in stabilizing the transi-
tion state by peroxoborates. These results and those of Pizer
and Tihal[9] on the effect of pH on the formation of dif-
ferent peroxoborates and the elegant reviews by McKillop
and Sanderson[7b] on sodium perborate (SPB) oxidations in
addition to our earlier experience in the synthetic peroxo–
boron chemistry[10] caused us to perceive the possibility of
borax-activated and pH-controlled selective oxidation of
organic sulfides in a protic medium like MeOH or MeOH/
H2O. In a very recent study, Gomez et al.[11] suggested that
SPB could be used for nucleophilic oxidation, whereas so-
dium percarbonate (SPC) could be used for electrophilic
oxidation. It was further suggested that the influence of
support might be relatively more pronounced than the oxi-
dant.
Introduction
The discovery of hydrogen carbonate catalyzed oxidation
–
of organic sulfides[1] through peroxocarbonate (HCO4 ) and
the subsequent seminal contributions of Davies and co-
workers[2] showing that borate is a better catalyst than
–
–
HCO4 with rapidly formed peroxoborates, HOOB(OH)3
–
–
and (HOO)2B(OH)2 , being more reactive than HCO4 , and
extending the pH functional range to the 8–12 region are
very important developments in the realm of oxidation
chemistry of organic sulfur. The examples to bear with this
contention include oxidants like H2O2,[3] periodate,[4] perac-
ids,[5] permonosulfate,[6] and peroxymonocarbonate.[1,7] It
was suggested that a protic solvent facilitates the nucleo-
–
philic displacement of the carbonate in the case of HCO4
oxidations. Also, it was shown recently from ab initio calcu-
lations on the oxidation of (CH3)2S by H2O2 in aqueous
solutions[8] that the transition state involves O–O bond
breaking with concurrent S–O bond formation, and that
hydrogen transfer occurs after the system has passed the
transition state. Solvent molecules or H2O2 (with higher
concentration) can efficiently lower the activation barrier
for the oxidative transformations. Insofar as the oxidation
of organic sulfides by peroxoborate is concerned, the reac-
tions generally proceed in a manner similar to other peroxo-
based oxidants, for instance, by development of a positive
Given the environmental acceptability, cost effectiveness,
and ease of handling, the borax/H2O2 system offers an ideal
combination for the chosen oxidations. A number of advan-
tages for the use of peroxoborates,[7b] the vast chemistry and
practical importance of sulfoxides and sulfones,[12] and the
plethora of oxidizing agents[3–7,13] and catalysts[14,15] are
well documented in the literature, thereby making any elab-
oration of these aspects redundant. Our main concern for
the present work was the selective oxidation of organic sul-
fides in an environmentally clean and catalytic manner with
[a] Department of Chemistry, Indian Institute of Technology Gu-
wahati,
Guwahati-781039, India
[b] Department of Chemistry, Indian Institute of Technology
Patna,
Patna-800013, India
[c] Tezpur University,
Tezpur-784028, India
Fax: +91-3712-267006
E-mail: mkc@iitg.ernet.in
mkc@tezu.ernet.in
Eur. J. Org. Chem. 2009, 3319–3322
© 2009 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
3319

S. Hussain, S. K. Bharadwaj, R. Pandey, M. K. Chaudhuri
SHORT COMMUNICATION
H2O2 being the terminal oxidant. Herein, we report the re- Table 1. Optimization of pH for selective oxidation of methyl
phenyl sulfide.
sults of our endeavor in the borax-catalyzed selective oxi-
dation of organic sulfides by H2O2 in methanol at pH 6–7
and at pH 10–11.
Entry
pH
Time [h] Product(s)
1
2
3
4
5
6
7
8
6.0
6.0
7.0
7.5
8.0
9.0
10.0
11.0
5
6
6
Sulfoxide (80%)
Sulfoxide (90%)
Sulfoxide (92%)
6.5
6
6
2.5
2.5
Sulfoxide (85%)/Sulfone (15%)
Sulfoxide (75%)/Sulfone (25%)
Sulfoxide (35%)/Sulfone (65%)
Sulfone (93%)
Results and Discussion
In order to ascertain the optimum conditions, several re-
action runs were carried out on methyl phenyl sulfide as the
model substrate, each time with the substrate (2 mmol),
H2O2 (35% aqueous solution, 6 mmol), and borax
Sulfone (93%)
(0.2 mmol) in MeOH (2 mL) at different pH, as shown in much higher. This indicates that a higher concentration of
Table 1. The pH of the reaction solution was adjusted by diperoxoborates favors sulfone formation over the corre-
the careful addition of a 0.1  NaOH solution. The control sponding sulfoxide.
experiment without catalyst showed 65% conversion after
In order to generalize the scope of the reaction, a series
12 h to 85% sulfoxide and 15% sulfone. Use of four or five of structurally diverse sulfides were subjected to oxidation
molar equivalents of H2O2, instead of three molar equiva- under the optimized reaction conditions, and the results are
lents, at pH 6–7 reduced the reaction time by ca. 1 h; how- presented in Table 2. The reactions went on well to afford
ever, 15–20% of sulfone was formed along with the sulfox- products in high yields. It is notable that the sulfides were
ide, thereby reducing the selectivity of the oxidation. The chemoselectively oxidized in the presence of some oxi-
selective oxidation to sulfone, if desired, was best performed dation-prone functional groups such as C=C, –CN, and
at pH 10–11 (Table 1).
–OH (Table 2, Entries 6–8). Sulfides containing a methyl es-
From perusal of the distribution diagram of the peroxo– ter group gave sulfone, along with its hydrolysis product in
borate system presented by Pizer and Tihal[9] it is evident basic medium (Table 2, Entry 9).
that at pH 6–7 mono and diperoxoborates, (HO)3BOOH–
Dibenzothiophene (DBT) and substituted DBT oxi-
–
and (HO)2B(OOH)2 , respectively, occur in equal but rela- dations are rather difficult with standard oxidation pro-
tively low concentration along with a very minute amount cedures.[3b] However, upon treatment with the borax/H2O2
of inactive peroxoboric acid, (HO)2BOOH. With the in- system some of these substrates were converted into the cor-
crease in pH to 10–11, peroxoboric acid disappears while responding sulfoxides and sulfones (Table 2, Entries 12 and
both the peroxoborates occur in relatively higher concentra- 13). Although we succeeded in oxidizing DBT and 4-
–
tions, with the concentration of (HO)2B(OOH)2 being methyl-DBT, our protocol did not work well for 4,6-di-
Table 2. Borax-catalyzed (0.1 equiv.) sulfide oxidation in MeOH by H2O2 at room temperature.
Entry
Substrate
pH 6–7
pH 10–11
Time [h]
Sulfoxide[a]
Time [h]
Sulfone[a]
1
2
PhSCH3
PhSC4H9
6
8
92,83,[b] 94[c]
2.5
4
93, 85,[b] 95[c]
82
90
3
4
PhSC6H13
C2H5SC4H9
24
5
65
82
24
3
87
84
5
6
7
8
C4H9SC4H9
8
8
7
5
5
8
6
10
24
–
78
85
85
90
3.5
6
6
–
5
3
2.5
6
24
24
78
90
88
PhSCH2CH=CH2
PhSCH2CH2CN
PhCH2SCH2CH2OH
PhSCH2CH2CO2CH3
PhSCH2Ph
p-NO2PhSCH2Ph
Dibenzothiophene
4-Methyldibenzothiophene
4,6-Diethyldibenzothiophene
–
9
92
90, 80,[b] 92[c]
88
62 + 25[e]
10
11
12
13
14
94,78,[b] 97[c]
95
75
75
10
55 + 42[d]
45 + 58[d]
–
[a] Percent isolated yield. [b] Yield after fifth cycle. [c] Yield at 5-g scale. [d] Percent of sulfone. [e] Hydrolyzed product.
3320 Eur. J. Org. Chem. 2009, 3319–3322
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© 2009 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim

Environmentally Clean Protocol for the Oxidation of Organic Sulfides
solvent of the pure fraction was evaporated, the product was fully
dried, and the isolated yield was calculated. The aqueous phase
containing borax, obtained after extraction of the reaction mixture
with ethyl acetate, was reused with the addition of fresh substrate
and H2O2 (3 equiv.) with adjustment of the pH value of the solu-
tion.
methyl-DBT. Only 10% oxidation was achieved after 24 h,
possibly as a result of steric crowding by the methyl groups
on DBT, which renders approach of the oxidant to the sul-
fur difficult, thereby causing problems (Table 2, Entry 14).
It may be mentioned that with an increase in alkyl chain
length of the sulfides, the rate of the reaction decreases. This
may be due to orientation of the hydrophobic alkyl chain
around the sulfur atom.
Acknowledgments
Recyclability of the catalyst was examined through a
series of reactions with methyl phenyl sulfide by using the
aqueous phase containing borax, obtained after extraction
of the reaction mixture with ethyl acetate. This was charged
with fresh substrate and H2O2 (3 equiv.) with adjustment of
the pH of the solution. The catalyst could be reused for at
least five reaction cycles with consistent activity. Import-
antly, the reaction can be performed on a relatively large
scale (5 g) to give good yields (Table 2, Entries 1 and 10),
which shows its potential for scale-up applications.
S. H. (2003–07) and S. K. B. thank the Council of Scientific and
Industrial Research, New Delhi for research fellowships.
[1] a) D. E. Richardson, H. Yao, K. M. Frank, D. A. Bennett, J.
Am. Chem. Soc. 2000, 122, 1729; b) D. A. Bennett, H. Yao,
D. E. Richardson, Inorg. Chem. 2001, 40, 2996 and references
cited therein.
[2] D. M. Davies, M. E. Deary, K. Quill, R. A. Smith, Chem. Eur.
J. 2005, 11, 3552.
[3] a) W. C. Schumb, C. N. Satterfield, R. L. Wentworth, Hydrogen
Peroxide, Reinhold, New York, 1955; b) R. Noyori, M. Aoki,
K. Sato, Chem. Commun. 2003, 1977.
[4] F. Ruff, A. Kuscman, J. Chem. Soc. Perkin Trans. 2 1985, 683.
[5] a) D. Swern in Organic Peroxides (Ed.: D. Swern), Wiley, New
York, 1970, vol. 1, pp. 313; b) J. Sanchez, T. N. Myers in Kirk-
Othmer Encyclopedia of Chemical Technology (Ed.: J. I.
Kroschwitz), Wiley-Interscience, New York, 1997, vol. 18, pp.
231.
Conclusions
In conclusion, the present study presents a metal-free cat-
alyst for the oxidation of sulfides to sulfoxides or sulfones
at different pH values by using H2O2 as the oxidant in
MeOH at room temperature. The borax/H2O2 system can
be applied to the oxidation of alkyl and aryl sulfides in the
presence of oxidation-prone functional groups such as
C=C, –CN, and –OH. Refractory sulfides, viz., DBT and
4-methyl-DBT, are also capable of being oxidized quite ef-
fectively, though with less selectivity insofar as sulfoxidation
is concerned. The oxidations of DBT substrates are espe-
cially important in the context of transportation fuel chem-
istry research targeting desulfurization of diesel and gaso-
line, for instance.
[6] C. A. Bunton, H. J. Foroudian, A. Kumar, J. Chem. Soc. Perkin
Trans. 2 1995, 33.
[7] a) Y. C. Yang, J. A. Baker, J. R. Ward, Chem. Rev. 1992, 92,
1729; b) A. McKillop, W. R. Sanderson, Tetrahedron 1995, 51,
6145; c) J. Muzart, Synthesis 1995, 1325; d) A. McKillop, W. R.
Sanderson, J. Chem. Soc. Perkin Trans. 1 2000, 471.
[8] J. W. Chu, B. L. Trout, J. Am. Chem. Soc. 2004, 126, 900.
[9] R. Pizer, C. Tihal, Inorg. Chem. 1987, 26, 3639.
[10] M. K. Chaudhuri, B. Das, Inorg. Chem. 1985, 24, 2580.
[11] M. V. Gomez, R. Caballero, E. Vazquez, A. Moreno, A. Hoz,
A. Diaz-Ortiz, Green Chem. 2007, 9, 331.
[12] a) P. C. B. Page (Ed.), Topics in Current Chemistry Vol. 204:
Organosulfur Chemistry I, Springer, Berlin, 1999; b) P. Metzner,
A. Thuillier in Sulfur Reagents in Organic Synthesis, Academic
Press, London, 1994; c) S. Patai, Z. Rappoport, Synthesis of
Sulfones, Sulfoxides, and Cyclic Sulfides, Wiley, Chichester,
1994.
[13] a) For reviews on reagents for sulfoxidation, see: P. Kowalski,
K. Mitka, K. Ossowskab, Z. Kolarska, Tetrahedron 2005, 61,
1933; b) G. Kar, A. K. Sakia, U. Bora, S. K. Dehury, M. K.
Chaudhuri, Tetrahedron Lett. 2003, 44, 4503; c) A. Ozanne-
Beaudenon, S. Quideau, Tetrahedron Lett. 2006, 47, 5869; d)
J. J. P. Selvam, V. Suresh, K. Rajesh, D. C. Babu, N. Suryaki-
ran, Y. Venkateswarlu, Tetrahedron Lett. 2008, 49, 3463.
[14] For a review on other metal catalysts for sulfoxidation, see: K.
Katarzyna, Z. Kolarska, K. Mitka, P. Kowalski, Tetrahedron
2005, 61, 8315.
Experimental Section
General: Reagent-grade chemicals such as borax (E. Merck, India)
and H2O2 were used as purchased. However, H2O2 was estimated
by permanganometry before use. DBT, 4-methyl DBT, and 4,6-di-
methyl DBT were purchased from Sigma Aldrich, India. Other or-
ganic sulfides were prepared by literature procedures. The products
were characterized by comparing their spectroscopic data recorded
with a Nicolet Impact-410 Fourier transform infrared spectropho-
tometer and a Varian-400 FT NMR spectrometer.
[15] a) V. Khanna, G. C. Maikap, J. Iqbal, Tetrahedron Lett. 1996,
37, 3367; b) K. Sato, M. Hyodo, M. Aoki, X. Q. Zheng, R.
Noyori, Tetrahedron 2001, 57, 2469; c) M. L. Kantam, B. Neel-
ima, C. V. Reddy, M. K. Chaudhuri, S. K. Dehury, Catal. Lett.
2004, 95, 19; d) S. Velusamy, A. V. Kumar, R. Saini, T. Punni-
yamurthy, Tetrahedron Lett. 2005, 46, 3819; e) K. Bahrami, Tet-
rahedron Lett. 2005, 46, 2009; f) H. Firouzabadi, N. Iranpoor,
A. A. Jafari, E. Riazymontazer, Adv. Synth. Catal. 2006, 348,
434; g) M. D. Rosa, M. Lamberti, C. Pellecchia, A. Scettri, R.
Villano, A. Soriente, Tetrahedron Lett. 2006, 47, 7233; h) M. A.
Hashimi, E. Fisset, A. C. Sullivan, J. R. H. Wilson, Tetrahedron
Lett. 2006, 47, 8017; i) I. Khedher, A. Ghorbel, J. M. Fraile,
J. A. Mayoral, J. Mol. Catal. A 2006, 255, 92; j) M. R. Maurya,
Oxidation of Organic Sulfides by H2O2/Borax: In a typical experi-
ment, to a 25-mL flask equipped with a magnetic stirrer and H2O2
(6.0 equiv.) was added borax (0.2 equiv.) and MeOH (2 mL) fol-
lowed by thioanisole (2 equiv.). To the resulting solution was added
0.1  NaOH to maintain the pH of the solution at 10, and the
mixture was stirred at room temperature for the time indicated in
Table 2. The reaction was monitored by TLC (n-hexane/ethyl ace-
tate, 9:1). After complete disappearance of the reactant, the prod-
uct was extracted with ethyl acetate. To the ethyl acetate extract
was added Na2S2O5 to destroy the excess amount of H2O2, and the
solution was then filtered and dried with anhydrous Na2SO4. The
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S. Hussain, S. K. Bharadwaj, R. Pandey, M. K. Chaudhuri
SHORT COMMUNICATION
A. K. Chandrakar, S. Chand, J. Mol. Catal. A 2007, 263, 227;
k) M. Mba, L. J. Prins, G. Licini, Org. Lett. 2007, 9, 21; l)
W. A. Maksoud, S. Daniele, A. B. Sorokin, Green Chem. 2008,
10, 447 and references cited therein; m) A. P. D. Costa, P. M.
Reis, C. Gamelas, C. C. Roma, B. Royo, Inorg. Chim. Acta
2008, 361, 1915.
Received: March 12, 2009
Published Online: May 27, 2009
3322
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Eur. J. Org. Chem. 2009, 3319–3322