Aerobic oxidation of methanol by a Ni(II)-O2 reaction

Article abstract of DOI:10.1039/b403668a

An unusual oxygen-activating Ni(II)-oximate complex oxidizes two-hydrogen atom donating substrates, including the traditionally inert alcohol, methanol, as well as ethanol, benzyl alcohol, benzylamine, and N-methylbenzylamine.

Full text of DOI:10.1039/b403668a

Aerobic oxidation of methanol by a Ni(II)-O2 reaction†
Sara E. Edison,a Richard P. Hotzb and Michael J. Baldwin*a
a Department of Chemistry, University of Cincinnati, Cincinnati OH 45221, USA.
E-mail: Michael.Baldwin@uc.edu; Fax: 513-556-9239; Tel: 513-556-9225
b Department of Chemistry, College of Mount St. Joseph, Cincinnati OH 45233, USA.
E-mail: richard_hotz@mail.msj.edu; Fax: 513-244-4597; Tel: 513-244-4833
Received (in West Lafayette, IN, USA) 10th March 2004, Accepted 26th March 2004
First published as an Advance Article on the web 23rd April 2004
An unusual oxygen-activating Ni(II)-oximate complex oxidizes
two-hydrogen atom donating substrates, including the tradi-
tionally inert alcohol, methanol, as well as ethanol, benzyl
alcohol, benzylamine, and N-methylbenzylamine.
uptake is observed when an acetonitrile solution of 3 is exposed to
air. Formaldehyde is formed in the methanol reaction, and was
monitored colorimetrically using the Hantzsch reaction.10 Fig. 2
shows that the amount of dioxygen consumed is approximately half
the amount of formaldehyde produced. It is proposed that hydrogen
peroxide is a product of this reaction in a 1:1 ratio with
formaldehyde; however, the efficient catalase-like activity of 3
prevents its detection. One equivalent of O2 is produced for every
two equivalents of H2O2 disproportionated, resulting in the
observed ~ 2:1 ratio of formaldehyde production to net dioxygen
consumption.
The catalase efficiency of 3 was confirmed in separate
manometry and colorimetry experiments. During the colorimetry
experiment, 400 equivalents of H2O2 were added to a methanolic
solution of 1. Addition of hydroxide to form 3 resulted in the
expected color change from purple to brown and was accompanied
by a vigorous bubbling of the solution as the H2O2 was
disproportionated. After approximately fifteen minutes the bub-
bling ceased and an aliquot of the solution was added to a Ti(SO4)2
reagent to test for H2O2.11 No H2O2 remained. As a control,
Zn(TRISOXH3)Cl2 was treated in an identical fashion to 1 and the
H2O2 remained after the same time period. For the manometry
experiment, 200 equivalents of H2O2 were added to a solution of 1.
Three equivalents of hydroxide were added to form 3 and the
amount of O2 generated from the disproportionation of H2O2 was
monitored. All of the H2O2 was disproportionated within 15
minutes, much faster than its formation during the methanol
oxidation reaction.
The chemistry of transition metals with dioxygen is important both
biologically and commercially. Many biological oxidations using
O2 are catalyzed by enzymes that contain transition metals in their
active sites.1 These processes inspire development of new transition
metal catalysts for industrial substrate oxidations by dioxygen, an
inexpensive and environmentally friendly oxidant.2 Understanding
the mechanism of atypical M–O2 reactions may lead to the rational
design of oxygen activation catalysts that are capable of unusual
substrate oxidations.
Reactions of Ni(II) complexes with O2 are uncommon, but not
unprecedented. However, they generally require irreversible ligand
oxidation. Some Ni(II)–amidate complexes (including complexes
of peptides) react with O2, inevitably undergoing ligand oxidation
by hydrogen atom transfer,3 rendering the complex unsuitable for
catalysis. Ni(II) thiolates react with O2 at the sulfur rather than the
nickel, resulting in various sulfur oxygenates.4 The recently
reported reaction of a Ni(II)–carbene complex with O2 appears to be
driven by hydrogen atom abstraction from a p-allyl co-ligand.5
6
Reactions of Ni(
I
) with O2 and Ni(II) with H2O2,7 (systems
containing either a reduced metal or oxygen species relative to
Ni(II) or O2),form [LNi(III)(m-O)]2 products without requiring
ligand oxidation.
We previously reported that deprotonation of a single oxime in
Ni(II)(TRISOXH3)(NO3)2 ([TRISOXH3 = tris(2-hydroxyimino-
propyl)amine]), (1), results in formation of a structurally charac-
The observation of methanol oxidation led us to investigate other
possible substrates. A 2 mM solution of 3 was reacted with O2 using
either the substrate as the solvent (methanol, ethanol, benzyl
2+
terized oximate-bridged dimer, [Ni(II)(TRISOXH2)(CH3CN)]2
(2). Further deprotonation of the complex forms another putative
oximate-bridged dimer, (3), Fig. 1. 3 undergoes multiple turnovers
during reaction with O2 without significant ligand decomposition,
making this the first homogeneous Ni(II) + O2 reaction that is not
driven by ligand oxidation.8 Here we report this reaction with
several oxidation substrates (including relatively inert metha-
nol9).
Adding three equivalents of hydroxide to a methanolic solution
of 1 to form 3 under air or O2 results in a rapid color change from
purple to brown that is not observed under anaerobic conditions.
This color change is accompanied by consumption of O2 as
determined by manometry. In contrast, no color change or O2
Fig. 1 Formation of the oxygen active species, 3.
† Electronic supplementary information (ESI) available: experimental
details and data for the O2 reaction with 3 and its oxidation substrates,
manometry data for the disproportionation of H2O2, and calibration curves
for chemical detection of products. See http://www.rsc.org/suppdata/cc/b4/
b403668a/
Fig. 2 Production of formaldehyde and consumption of dioxygen from the
reaction of 1 mM 3 in methanol and acetonitrile (CH2O formation in MeOH
=
squares, O2 uptake in MeOH = circles, O2 uptake in MeCN =
diamonds). Each data point is the average of three trials and the error bars
are the standard deviations.
1212
C h e m . C o m m u n . , 2 0 0 4 , 1 2 1 2 – 1 2 1 3
T h i s j o u r n a l i s © T h e R o y a l S o c i e t y o f C h e m i s t r y 2 0 0 4

View Article Online
Table 1 Bond dissociation energies (kcal mol21) for hydrogen atom donors and hydrogen peroxide.12 Product quantities are the average of three experiments.
(NA = not available, NR = no reaction)
Eq. of Product
after 1 hour
Eq. of Product
after 24 hours
Substrate
1st H·
2nd H·
Total
Methanol
Ethanol
Benzyl alcohol
2-Propanol
93
90
87.5
91
31
26
18
26
14
NA
23
NA
46
124
116
105.5
117
102
NA
123
NA
134.5
3.95 ± 0.19
2.90 ± 1.23
0.85 ± 0.02
NR
10.37 ± 0.45
8.54 ± 3.14
3.02 ± 0.57
NR
D,L
-1-Phenylethanol
88
NR
NR
N-Methylbenzyl amine
Benzyl amine
a-Methylbenzyl amine
H2O2
NA
100
NA
88.5
0.24 ± 0.04
0.93 ± 0.30
NR
0.64 ± 0.10
4.95 ± 0.95
NR
alcohol), or by using acetonitrile as the solvent and adding 200
equivalents of substrate (benzylamine, N-methylbenzylamine).
Both protocols resulted in the expected color change and produced
multiple equivalents of most oxidized products (Table 1). The
primary alcohols produced aldehydes, and the amines produced N-
benzylidene benzylamine. The oxidation of benzylamine may
follow one of two pathways. It may undergo oxidative deamination
to form benzaldehyde and ammonia (ammonia was detected as one
of the products of this reaction) with the benzaldehyde further
reacting with excess amine to form the Schiff base product, N-
benzylidene benzylamine. Alternatively, two hydrogen atoms may
be abstracted to form the imine, which then reacts with excess
amine to form ammonia and the same Schiff base product. N-
Methylbenzylamine likely undergoes reversible addition–elimina-
tion through an aminal intermediate, also forming N-benzylidene
benzylamine, as previously reported by Murahashi, et al.13
All of the substrates studied were oxidized by the abstraction of
two hydrogen atoms according to reaction (1):
coworkers’ Cu complex that oxidized less than one equivalent of
methanol.9
Funding was provided by the donors of the Petroleum Research
Fund administered by the American Chemical Society (ACS-PRF
33960-G3) and the University of Cincinnati. We would also like to
thank Prof. Michael J. Goldcamp (Wilmington College) for helpful
discussions.
Notes and references
1 L. Que, Jr. and Y. Watanabe, Science, 2001, 292, 651–653.
2 R. A. Sheldon, Biocatalyic and biomemetic oxidations from an
industrial perspective; B. Meunier, Ed., Imperial College Press,
London, 2000, pp. 613–662.
3 (a) F. P. Bossu, E. B. Paniago, D. W. Margerum, S. T. J. Kirskey and J.
L. Kurtz, Inorg. Chem., 1978, 17, 1034–1042; (b) W. Bal, M. I. Djuran,
D. W. Margerum, E. T. J. Gray, M. A. Mazid, R. T. Tome, E. Nieboer
and P. J. Sadler, J. Chem. Soc., Chem. Commun., 1994, 1889–1890; (c)
E. Kimura, M. Sasada, M. Shionoya, T. Koike, H. Kurosaki and M.
Shiro, J. Bioinorg. Chem, 1997, 2, 74–82; (d) D. Chen, R. J. Motekaitis
and A. E. Martell, Inorg. Chem., 1991, 30, 1396–1402; (e) C.-C. Cheng,
J. Gulia, S. E. Rokita and C. J. Burrows, J. Mol. Catal. A., 1996, 113,
379–391.
4 C. A. Grapperhaus and M. Y. Darensbourg, Acc. Chem. Res., 1998, 31,
451–459.
5 B. R. Dibble and M. S. Sigman, J. Am. Chem. Soc., 2003, 125,
872–873.
6 B. S. Mandimutsira, J. L. Yamarik, T. C. Brunold, W. Gu, S. P. Cramer
and C. G. Riordan, J. Am. Chem. Soc., 2001, 123, 9194–9195.
7 (a) K. Shiren, S. Ogo, S. Fujinami, H. Hayashi, M. Suzuki, A. Uehara,
Y. Watanabe and Y. Moro-oka, J. Am. Chem. Soc., 2000, 122, 254–262;
(b) S. Itoh, H. Bendoh, M. Nakagawa, S. Nagatomo, T. Kitagawa, K. D.
Karlin and S. Fukuzumi, J. Am. Chem. Soc., 2001, 123,
11168–11178.
(1)
RH2 + O2 ? R + H2O2.
The bond dissociation energies in Table 1 show that transfer of
the first hydrogen to O2 is unfavorable. Transfer of the second H-
atom to form H2O2 makes the overall reaction exothermic. Thus,
there is a thermodynamic requirement for a two H-atom reaction.
However, several potential substrates whose oxidations are thermo-
dynamically favorable (2-propanol,
D
, -1-phenylethanol, a-me-
L
thylbenzylamine) produce no color change with 3, and no ketones
were formed from the alcohols. These unreactive substrates share
the common feature that they are branched at the a-carbon.
The reactivity of 3 is reminiscent of several enzymatic processes.
Galactose oxidase catalyzes the aerobic oxidation of a primary
alcohol to an aldehyde with concurrent H2O2 formation,14 similar
to the oxidation of methanol, ethanol, and benzyl alcohol by 3.
Copper amine oxidases catalyze the aerobic oxidation of a primary
amine to form an aldehyde, NH3, and H2O2,15 related to the reaction
of 3 with benzylamine to form N-benzylidene benzylamine and
ammonia. Both heme and dinuclear Mn catalase enzymes catalyze
H2O2 disproportionation.16
8 M. J. Goldcamp, S. E. Robison, J. A. Krause Bauer and M. J. Baldwin,
Inorg. Chem., 2002, 41, 2307–2309.
9 (a) P. Chaudhuri, M. Hess, J. Muller, K. Hildenbrand, E. Bill, T.
Weyhermuller and K. Wieghardt, J. Am. Chem. Soc., 1999, 121,
9599–9610; (b) C. Xin Zhang, H. Liang, E. Kim, J. Sherare, M. E.
Helton, E. Kim, S. Kaderli, C. D. Incarvito, A. D. Zuberbuhler, A. L.
Rheingold and D. K. Karlin, J. Am. Chem. Soc., 2003, 125, 634–635.
10 T. Nash, Biochem. J., 1953, 55, 416–421.
11 P. A. Clapp, D. F. Evans and T. S. S. Sheriff, Anal. Chim. Acta, 1989,
218, 331–334.
The oxidation of methanol or benzylamine was also investigated
in aqueous solution. The aqueous environment inhibited formal-
dehyde production (1.88 ± 0.03 equiv. after 24 h, versus 5.07 ± 0.31
in acetonitrile, both experiments containing 50% methanol) but
does not significantly affect benzylamine oxidation (5.44 ± 0.54
equiv. after 24 h).
In summary, we have reported substrate oxidation by the first
Ni(II) + O2 reaction that does not proceed via irreversible ligand
oxidation. It instead requires an exogenous source of hydrogen
atoms. This can come from a relatively inert substrate such as
methanol, provided a second low energy H-atom dissociation is
available. The aerobic oxidation of methanol by discrete transition
metal complexes has been reported in only a few instances, most
notably in Wieghardt and coworkers’ Zn complex with a redox
active ligand that displayed catalytic oxidation, and Karlin and
12 (a) J. A. Dean, Lange’s Handbook of Chemistry, 13th Ed., McGraw-Hill
Book Co., New York, 1985; (b) V. I. Vedeneyev, L. V. Gurvich, V. N.
Kondrat’yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies,
Ionization Potentials, and Electron Affinities, Edward Arnold LTD.,
London,1966; (c) Y.-R. Luo, Bond Dissociation Energies in Organic
Compounds, CRC Press, Boca Raton, 2003.
13 S.-I. Murahashi, N. Yoshimura, T. Tsumiyama and T. Kojima, J. Am.
Chem. Soc., 1983, 105, 5002–5011.
14 K. Clark, J. E. Penner-Hahn, M. M. Whittaker and J. W. Whittaker, J.
Am. Chem. Soc., 1990, 112, 6433–6434.
15 M. Mure and J. P. Klinman, J. Am. Chem. Soc., 1995, 117,
8698–8706.
16 (a) C. D. Putnam, A. S. Arvai, Y. Bourne and J. A. Trainer, J. Mol. Biol.,
2000, 296, 295–309; (b) A. J. Wu, J. E. Penner-Hahn and V. L. Pecararo,
Chem. Rev., 2004, 104, 903–938.
C h e m . C o m m u n . , 2 0 0 4 , 1 2 1 2 – 1 2 1 3
1213