Much enhanced catalytic reactivity of cobalt chlorin derivatives on two-electron reduction of dioxygen to produce hydrogen peroxide

Article abstract of DOI:10.1021/ic502678k

Effects of changes in the redox potential or configuration of cobalt chlorin derivatives (Co^II(Ch_n) (n = 1-3)) on the catalytic mechanism and the activity of two-electron reduction of dioxygen (O₂) were investigated based on the detailed kinetic study by spectroscopic and electrochemical measurements. Nonsubstituted cobalt chlorin complex (Co^II(Ch₁)) efficiently and selectively catalyzed two-electron reduction of dioxygen (O₂) by a one-electron reductant (1,1′-dimethylferrocene) to produce hydrogen peroxide (H₂O₂) in the presence of perchloric acid (HClO₄) in benzonitrile (PhCN) at 298 K. The detailed kinetic studies have revealed that the rate-determining step in the catalytic cycle is the proton-coupled electron transfer reduction of O₂ with the protonated Co^II(Ch₁) complex ([Co^II(Ch₁H)]⁺), where one-electron reduction potential of [Co^III(Ch₁)]⁺ was changed from 0.37 V (vs SCE) to 0.48 V by the addition of HClO₄ due to the protonation of [Co^III(Ch₁)]⁺. The introduction of electron-withdrawing aldehyde group (position C-3) (Co^II(Ch₃)) and both methoxycarbonyl group (position C-13²) and aldehyde group (position C-3) (Co^II(Ch₂)) on the chlorin ligand resulted in the positive shifts of redox potential for Co(III/II) from 0.37 V to 0.45 and 0.40 V, respectively, whereas, in the presence of HClO₄, no positive shifts of those redox potentials for [Co^III(Ch_n)]⁺/Co^II(Ch_n) (n = 2, 3) were observed due to lower acceptability of protonation. As a result, such a change in redox property resulted in the enhancement of the catalytic reactivity, where the observed rate constant (k_obs) value of Co^II(Ch₃) was 36-fold larger than that of Co^II(Ch₁).

Full text of DOI:10.1021/ic502678k

Article
pubs.acs.org/IC
Much Enhanced Catalytic Reactivity of Cobalt Chlorin Derivatives on
Two-Electron Reduction of Dioxygen to Produce Hydrogen Peroxide
Kentaro Mase, Kei Ohkubo, and Shunichi Fukuzumi*
Department of Material and Life Science, Graduate School of Engineering, ALCA, Japan Science and Technology Agency (JST),
Osaka University, Suita, Osaka 565-0871, Japan
S
* Supporting Information
ABSTRACT: Effects of changes in the redox potential or
configuration of cobalt chlorin derivatives (CoII(Chn) (n = 1−
3)) on the catalytic mechanism and the activity of two-electron
reduction of dioxygen (O2) were investigated based on the
detailed kinetic study by spectroscopic and electrochemical
measurements. Nonsubstituted cobalt chlorin complex
(CoII(Ch1)) efficiently and selectively catalyzed two-electron
reduction of dioxygen (O2) by a one-electron reductant (1,1′-
dimethylferrocene) to produce hydrogen peroxide (H2O2) in
the presence of perchloric acid (HClO4) in benzonitrile
(PhCN) at 298 K. The detailed kinetic studies have revealed
that the rate-determining step in the catalytic cycle is the
proton-coupled electron transfer reduction of O2 with the protonated CoII(Ch1) complex ([CoII(Ch1H)]+), where one-electron
reduction potential of [CoIII(Ch1)]+ was changed from 0.37 V (vs SCE) to 0.48 V by the addition of HClO4 due to the
protonation of [CoIII(Ch1)]+. The introduction of electron-withdrawing aldehyde group (position C-3) (CoII(Ch3)) and both
methoxycarbonyl group (position C-132) and aldehyde group (position C-3) (CoII(Ch2)) on the chlorin ligand resulted in the
positive shifts of redox potential for Co(III/II) from 0.37 V to 0.45 and 0.40 V, respectively, whereas, in the presence of HClO4,
no positive shifts of those redox potentials for [CoIII(Chn)]+/CoII(Chn) (n = 2, 3) were observed due to lower acceptability of
protonation. As a result, such a change in redox property resulted in the enhancement of the catalytic reactivity, where the
observed rate constant (kobs) value of CoII(Ch3) was 36-fold larger than that of CoII(Ch1).
In order to resolve these matters, an alternative method to
produce H2O2 has been extensively studied by photochemical
or thermal means without the use of noble metal catalysts.10−17
Extensive efforts have been implemented based on homoge-
neous transition metal complexes such as cobalt,18−28 iron,29−32
and copper,33−36 and other metal complexes17 have been
employed for the two-electron and four-electron reduction of
O2. So far, monomeric cobalt-centered complexes have been
reported to show high affinity for the O2 reduction and act as
efficient catalysts for the selective two-electron reduction of O2
to produce H2O2.18−20 Nevertheless, the stability of those
catalysts under the acidic conditions was not well taken into
account. Decomposition of metal complexes with a porphyr-
inoid ligand is mainly caused by demetalation of metal center in
an acidic medium, which depends on two factors.37−40 One is
the core size and rigidity of a macrocyclic ligand, because a low-
valent metal ion, which is produced in a catalytic cycle, is
difficult to be accommodated in a macrocyclic ligand due to its
large ionic radius and low electrostatic interaction. The other is
the nucleophilicity of core nitrogen atoms, which is responsible
for the replacement of the metal ion by incoming proton.
INTRODUCTION
■
Efficient conversion of electrical energy produced by solar cells
into easily storable energy has attracted increasing attention as a
key technology for the practical use of solar energy.1,2
Hydrogen peroxide (H2O2) produced by two-electron
reduction of earth abundant dioxygen (O2) is a promising
candidate as a high chemical energy carrier, because it has high
energy density owing to its liquid state at ambient temper-
ature.3 H2O2 stored as high chemical energy can be transformed
into electrical energy through fuel cells.3−8 H2O2-fuel cell
supplies theoretical output voltage of 1.09 V by utilizing H2O2
as both a fuel at cathode and an oxidant at anode, emitting only
water and oxygen after power generation, and its cell
performance has been improved significantly.3−8 Currently,
H2O2 is produced in industry mainly by the anthraquinone
method, in which H2O2 is produced in the process of the
autoxidation of a 2-alkylanthrahydroquinone to the correspond-
ing anthraquinone.9 However, this process requires hydrogen as
a reducing reagent and a noble metal such as palladium to
regenerate the anthrahydroquinone. The cost of the anthra-
quinone process depends heavily on effective recycling of the
anthraquinone, extraction solvents, and noble metal catalysts
for the hydrogenation, which are all quite expensive.
Received: November 11, 2014
© XXXX American Chemical Society
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and [Co(ChnH)]+ was determined by using the Hill equation, which
analyzes changes of absorption spectra during the titration as a
function of concentration of an added acid.46 The amount of hydrogen
peroxide (H2O2) formed was determined by titration with iodide ion:
A diluted CH3CN solution (2.0 mL) of the product mixture (40 μL)
We have previously reported that a cobalt chlorin complex
(CoII(Ch)) efficiently and selectively catalyzes two-electron
reduction of O2 by using one-electron molecular reductants
(ferrocene derivatives) with high catalyst stability and small
overpotential in the presence of perchloric acid (HClO4) in
benzonitrile (PhCN).18 The remarkable stability and efficiency
are attributed to the nature of chlorin ligand; the core size of
the ligand is larger and more flexible than that of porphyrin
ligand, and lower basicity of the core nitrogen atoms is induced
by the protonation of the carbonyl group (position C-131) to
prevent protonation of core nitrogen atoms. In addition, the
chemical reduction approach under homogeneous conditions
allows the survey of the catalytic mechanism, where the rate-
determining step was determined to be proton-coupled
electron transfer (PCET) from CoII(Ch1) to O2.18 However,
effects of changes in the redox potential and configuration of
chlorin ligand on the catalytic activity have yet to be examined
to further improve the catalytic activity.
−
was treated with an excess amount of NaI, and the amount of I3
formed was determined by the absorption spectrum (λmax = 361 nm, ε
= 2.8 × 104 M−1 cm−1).47
Kinetic Measurements. Rate constants of oxidation of ferrocene
derivatives by O2 in the presence of a catalytic amount of CoII(Chn) (n
= 1−3) and an excess amount of HClO4 in PhCN at 298 K were
determined by monitoring the appearance of an absorption band due
to 1,1′-dimethylferrocenium ion (Me2Fc+: λmax = 650 nm, εmax = 290
M−1 cm−1).15 At the wavelengths monitored, spectral overlap was
observed with [CoII(Ch1H)]+ [λ = 650 nm (1.4 × 104 M−1 cm−1), 700
nm (7.0 × 103 M−1 cm−1)], [CoII(Ch2H)]+ [λ = 650 nm (ε = 8.5 ×
103 M−1 cm−1), 700 nm (ε = 1.0 × 104 M−1 cm−1)], and
[CoII(Ch3H)]+ [λ = 650 nm (ε = 1.2 × 104 M−1 cm−1), 700 nm (ε
= 5.0 × 103 M−1 cm−1)]. An air-saturated PhCN solution was used for
the catalytic reduction of O2 by Me2Fc. The concentration of O2 in an
air-saturated PhCN solution (1.7 × 10−3 M) was determined as
reported previously.48 The concentrations of Me2Fc employed for the
catalytic reduction of O2 were much larger than that of O2, when O2 is
the reaction-limiting reagent in the reaction solution. The catalytic
turnover numbers (TON = mol of H2O2 as the product of the two-
electron reduction of O2/mol of CoII(Chn), n = 1−3) were determined
based on the initial concentration of CoII(Chn), when the
concentrations of CoII(Chn), Me2Fc, and HClO4 in O2-saturated
PhCN were 1.0 × 10−7 M, 5.0 × 10−2 M, and 5.0 × 10−2 M,
respectively. The formation of Me2Fc+ was monitored by a rise in
absorbance at 700 nm to avoid excess absorption at 650 nm.
Electrochemical Measurements. Cyclic voltammetry (CV)
measurements were performed on an ALS 630B electrochemical
analyzer, and voltammograms were measured in deaerated PhCN
containing TBAPF6 (0.10 M) as a supporting electrolyte at room
temperature. A conventional three-electrode cell was used with a glassy
carbon working electrode (surface area of 0.3 mm2) and a platinum
wire as the counter electrode. The glassy carbon working electrode
(BAS) was routinely polished with BAS polishing alumina suspension
and rinsed with acetone before use. The potentials were measured with
respect to the Ag/AgNO3 (1.0 × 10−2 M) reference electrode. All
potentials (vs Ag/AgNO3) were converted to values vs saturated
calomel electrode (SCE) by adding 0.29 V.49 Redox potentials were
determined using the relation E1/2 = (Epa + Epc)/2.
We report much enhanced catalytic activity of cobalt chlorin
complexes for selective two-electron reduction of O2 to H2O2
by examining a series of cobalt chlorin derivatives, CoII(Chn) (n
= 1−3), shown in Chart 1. The detailed study based on the
Chart 1. Cobalt Chlorin Derivatives
kinetics and thermodynamics have provided valuable insight
into the development of a more efficient electrocatalyst for two-
electron reduction of O2 to H2O2, which is essential to realize a
H2O2-energy society.
EPR Measurements. The EPR spectra were performed on a JEOL
X-band EPR spectrometer (JES-ME-LX) using a quartz EPR tube
containing a deaerated frozen sample solution at 4 K. The internal
diameter of the EPR tube is 4.5 mm, which is small enough to fill the
EPR cavity but large enough to obtain good signal-to-noise ratios
during the EPR measurements at low temperature. The EPR spectra
were measured under nonsaturating microwave power conditions. The
amplitude of modulation was chosen to optimize the resolution and
the signal-to-noise (S/N) ratio of the observed spectra. The g values
were calibrated with a Mn2+ marker.
EXPERIMENTAL SECTION
■
General Procedure. Chemicals were purchased from commercial
sources and used without further purification, unless otherwise noted.
Benzonitrile (PhCN) used for spectroscopic and electrochemical
measurements was distilled over phosphorus pentoxide prior to use.41
Cobalt chlorin [CoII(Chn), n = 1−3] was synthesized by the published
method (see ref 18 for the details).42−45 1,1′-Dimethylferrocene
(Me2Fc) were purchased from Aldrich Chemical Co. and purified by
sublimation or recrystallization from ethanol. Tetra-n-butylammonium
hexafluorophosphate (TBAPF6) purchased from Wako Pure Chemical
Industries, Ltd., was twice recrystallized from ethanol and dried in
RESULTS AND DISCUSSION
■
1
vacuo prior to use. H NMR spectra (300 MHz) were recorded on a
Protonation of CoII(Chn), n = 1−3. To examine the
proton acceptability of CoII(Chn), acid titration was performed
in PhCN by means of UV−vis spectroscopy. The protonation
equilibrium constants (K) of CoII(Chn) with HClO4 were
determined by monitoring absorption spectral change at Q-
bands of CoII(Chn) upon addition of HClO4 to deaerated
PhCN solutions of CoII(Chn) at 298 K. The Q bands of
CoII(Ch2) at 675 nm decreased with increasing concentration
of HClO4, followed by appearance of a broaden absorption
band at 678 nm as shown Figure 1. Such a UV−vis absorption
spectral change with clean isosbestic points has previously been
analyzed as monoprotonated at the carbonyl group (position
JEOL AL-300 spectrometer at room temperature, and chemical shifts
(ppm) were determined relative to tetramethylsilane (TMS). MALDI-
TOF-MS measurements were performed on a Kratos Compact
MALDI I (Shimadzu) using freshly prepared dithranol as a matrix in
the reflectron-positive mode. UV−vis spectroscopy was carried out on
a Hewlett-Packard 8453 diode array spectrophotometer at room
temperature using a quartz cell (light path length = 1 cm).
Spectroscopic Measurements. Titrations were carried out by
adding known quantities of a solution of HClO4 in high concentration
into a solution of CoII(Chn). The solution of HClO4 used to effect the
titration contained the same concentration of CoII(Chn) as used in the
titration so as to avoid the need to account for dilution effects during
the titration. The protonation equilibrium constant between Co(Chn)
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bulky ClO4− counter anion to the carbonyl group (position C-
131).
The protonation of CoII(Chn) (n = 1−3) at the carbonyl
group (position C-131) was previously evidenced by the
changes in EPR spectra, where the g|| value remained nearly the
same, whereas the g⊥ value changed slightly upon addition of
HClO4 to deaerated PhCN solutions of CoII(Chn) (n = 1−3) at
4 K.18 The larger changes in g values were observed for the
protonation of CoII(Ch1) as compared to those of CoII(Chn) (n
= 2, 3), as shown in Figure 2.
Figure 1. (a) Absorption spectral changes in the protonation of
CoII(Ch2) (2.0 × 10−5 M) upon addition of HClO4 to a deaerated
PhCN solution of CoII(Ch2) at 298 K. Inset shows absorbance change
at 675 nm upon addition of HClO4. (b) Hill plot of absorbance change
at 675 nm upon addition of HClO4, where α is (A − A0)/(A∞ − A0).
C-131) to form [CoII(ChnH)]+ (n = 1−3), as shown in Scheme
1.18
Scheme 1. Protonation of CoII(Chn), n = 1−3
The K value for the monoprotonation of CoII(Ch2) was
determined to be 2.2 × 102 M−1 in deaerated PhCN at 298 K.
Likewise, the addition of HClO4 to a solution of CoII(Ch3)
afforded similar spectral changes as shown in Figure S4 in the
Supporting Information. However, higher concentration of
HClO4 was required to obtain monoprotonated CoII(Ch3) as
compared with that of Co(Ch2). The relatively smaller K values
of CoII(Ch3) were determined to be 70 M−1 as listed in Table
1. These K values of CoII(Chn) (n = 2, 3) are much smaller
Table 1. Protonation Equilibrium Constants of CoII(Chn)
and [CoIII(Chn)]+ (n = 1−3) Measured at 298 K
K, M−1
cobalt chlorin
Co(II)
Co(III)
a
a
CoII(Ch1)
CoII(Ch2)
CoII(Ch3)
2.2 × 104
2.2 × 102
7.0 × 10
1.2 × 102
b
b
a
b
Data taken from ref 18. No protonation.
than the K value of CoII(Ch1) (K = 2.2 × 104 M−1)18 because
of a decrease in the electron density of macrocyclic ligands
caused by the introduction of electron-withdrawing substitu-
ents, such as methoxycarbonyl group (position C-132) and
aldehyde group (position C-3), as expected from the weaker
basicity of macrocyclic ligands with the electron-withdrawing
substituents as compared to the basicity of unsubstituted
macrocyclic ligand. The methoxycarbonyl group (position C-
132) near by the carbonyl group (position C-131) may also
cause steric hindrance, preventing the approach of proton with
Figure 2. EPR spectra of (a) CoII(Ch1) (1.0 × 10−3 M) in deaerated
PhCN at 4 K (black) and CoII(Ch1) (1.0 × 10−3 M) upon addition of
HClO4 (1.0 × 10−2 M) in deaerated PhCN at 4 K (red), (b)
CoII(Ch2) (1.0 × 10−3 M) in deaerated PhCN at 4 K (black) and
CoII(Ch2) (1.0 × 10−3 M) upon addition of HClO4 (1.0 × 10−2 M) in
deaerated PhCN at 4 K (red), and (c) CoII(Ch3) (1.0 × 10−3 M) in
deaerated PhCN at 4 K (black) and CoII(Ch3) (1.0 × 10−3 M) upon
addition of HClO4 (1.0 × 10−2 M) in deaerated PhCN at 4 K (red).
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The same set of spectral titrations was performed on one-
electron oxidized species, [CoIII(Chn)]+ (n = 1−3), produced
by aerial oxidation of CoII(Chn) in the presence of a small
amount of HClO4. As can be seen in the previous report,18 the
protonation equilibrium constants (K) of [CoIII(Chn)]+ (n =
1−3), as given by eq 1, exhibited much smaller values due to
lower basicity of chlorin ligand induced by the higher-valent
Co(III) ion center, as shown in Figure S5 in the Supporting
Information.
sweeps for CoII(Ch2) and CoII(Ch3) show two-reduction
waves. On the basis of redox potentials of previously
characterized Co(II) porphyrins,50 the first reduction processes
are assigned as due to formation of the Co(I) complex,
([CoI(Ch1)]−). Additionally, CoII(Ch2) and CoII(Ch3) undergo
the second reduction ascribed to one-electron ligand reduction
to form Co(I) radical anion complex ([CoI(Chn•−)]2−, n = 2,
3). Each redox couple was positively shifted from 0.37 V (vs
SCE) for [CoIII(Ch1)]+/CoII(Ch1) to 0.40 V for [CoIII(Ch3)]+/
CoII(Ch3) and 0.45 V for [CoIII(Ch2)]+/CoII(Ch2), respec-
tively. Those are explained by the introduction of electron-
withdrawing methoxycarbonyl group (position C-132) or both
methoxycarbonyl group (position C-132) and aldehyde group
(position C-3).
K
[CoIII(Chn)]+ + H+ ⇄ [CoIII(ChnH)]2+
(1)
The K value for the monoprotonation of [CoIII(Ch1)]+ was
previously determined to be 1.2 × 102 M−1 in air-saturated
PhCN at 298 K. However, the addition of HClO4 to an air-
saturated PhCN solution of [CoIII(Chn)]+ (n = 2, 3) resulted in
virtually no spectral change, where K values could not be
determined. The protonation equilibrium constants between
CoII(Chn) and [CoII(ChnH)]+ together with [CoIII(Chn)]+ and
[CoIII(ChnH)]2+ as given by Scheme 1 and eq 1 are listed in
Table 1.
Upon addition of HClO4 to the deaerated [CoIII(Chn)]+
solutions, the redox potential of [CoIII(Ch1)]+/CoII(Ch1) was
positively shifted from 0.37 to 0.48 V, whereas the redox
potential of [CoIII(Chn)]+/CoII(Chn) (n = 2, 3) remained the
same as that in the absence of HClO4 (Figure 4).
Redox Properties of CoII(Chn), n = 1−3. Electrochemical
measurements of cobalt(II) chlorin derivatives were measured
to examine the influence of the substituents of chlorin
macrocyclic ligands on the catalytic activity. The cyclic
voltammograms for all the CoII(Chn) in deaerated PhCN
solutions containing TBAPF6 (0.10 M) exhibited reversible
two-oxidation waves for conversion of CoII(Chn) to
[CoIII(Chn)]+ and [CoIII(Chn)]+ to [CoIII(Chn•+)]2+, respec-
tively, as shown in Figure 3.
Figure 4. Cyclic voltammograms of deaerated PhCN solutions of (a)
CoII(Ch1) (2.0 × 10−3 M), (b) CoII(Ch2) (2.0 × 10−3 M), and (c)
CoII(Ch3) (2.0 × 10−3 M) with HClO4 (2.5 × 10−2 M) recorded in
the presence of TBAPF6 (0.10 M) at 298 K, respectively; sweep rate:
0.10 V s−1.
These results are consistent with the observation that the
ligand protonation hardly occurred on CoII(Chn) (n = 2, 3) due
to diminished basicity and steric hindrance on chlorin ligand
caused by the introduction of substituents. Table 2 summarizes
all the redox potentials of [CoIII(Chn)]+/CoII(Chn) (n = 1−3)
Table 2. Redox Potentials (V vs SCE) of CoIII/CoII in the
Absence and Presence of HClO4 Measured at 298 K
E1/2 of CoIII/CoII, V
Figure 3. Cyclic voltammograms of deaerated PhCN solutions of (a)
CoII(Ch1) (2.0 × 10−3 M), (b) CoII(Ch2) (2.0 × 10−3 M), and (c)
CoII(Ch3) (2.0 × 10−3 M) recorded in the presence of TBAPF6 (0.10
M) at 298 K, respectively; sweep rate: 0.10 V s−1.
cobalt chlorin
without HClO4
with 25 mM of HClO4
CoII(Ch1)
CoII(Ch2)
CoII(Ch3)
0.37
0.45
0.40
0.48
0.45
0.40
In contrast to oxidation sweep, reduction sweep for
CoII(Ch1) shows one-reduction waves whereas reduction
a
Data taken from ref 18.
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in the absence and presence of HClO4. Consequently, the
oxidation potential of CoII(Ch1), where the Co(II) complexes
are responsible for the rate-determining O2 reduction step in
the catalytic cycle as reported previously, turned out to be the
highest value in the presence of HClO4.
Catalytic Two-Electron Reduction of O2 by Me2Fc with
CoII(Chn) in the Presence of HClO4. Catalytic O2 reduction
under homogeneous conditions was conducted by the addition
of 1,1′-dimethylferrocene (Me2Fc) as a molecular reductant to
an air-saturated PhCN solution, containing CoII(Chn) (n = 1−
3) as a catalyst and HClO4 as a proton source. The efficient
oxidation of Me2Fc by O2 occurred in the presence of
CoII(Chn) (n = 1−3) to afford 1,1′-dimethylferrocenium ion
(Me2Fc+) as shown in Figure 5, whereas the oxidation has
Figure 6. Plots of (a) kobs vs [CoII(Ch1)] (closed circle), (b) kobs vs
[CoII(Ch2)] (closed square), and (c) kobs vs [CoII(Ch3)] (closed
triangle) for the two-electron reduction of O2 (1.7 × 10−3 M) by
Me2Fc (2.5 × 10−2 M) in the presence of HClO4 (2.5 × 10−2 M) in
PhCN at 298 K.
conditions for the formation of Me2Fc+ on concentrations of a
series of catalysts (CoII(Chn), n = 1−3). The kobs values
increased linearly with increasing concentrations of catalysts
(CoII(Chn), n = 1−3). The higher activity was achieved when
CoII(Chn) (n = 2, 3) were employed as catalysts, where the kobs
value of CoII(Ch2) was 9-fold larger than that of CoII(Ch1) and
the kobs value of CoII(Ch3) was 36-fold larger than that of
CoII(Ch1). The turnover numbers (TON) of CoII(Chn) (n =
1−3) were determined from the produced concentration of
Me2Fc+ to be more than 60,000, 79,000, and 83,000
respectively, when small concentrations of catalysts were
employed, as shown in Figure S6 in the Supporting
Information.
Figure 5. (a) Absorption spectral change in the two-electron reduction
of O2 (1.7 × 10−3 M) by Me2Fc (2.5 × 10−2 M) with CoII(Ch2) (4.0 ×
10−6 M) in the presence of HClO4 (2.5 × 10−2 M) in air-saturated
PhCN at 298 K. The black and red lines show the spectra before and
after addition of HClO4, respectively. The dotted line is the
absorbance at 650 nm due to 3.4 × 10−3 M of Me2Fc+. (b) Time
profile of absorbance at 650 nm due to Me2Fc+. Inset: first-order plot.
hardly occurred in the absence of CoII(Chn) (n = 1−3) under
the same experimental conditions. In order to determine the
stoichiometry of the catalytic O2 reduction with CoII(Chn) (n =
1−3), the oxidized amount of Me2Fc by O2 was quantitatively
determined under the reaction conditions with an excess
amount of Me2Fc and HClO4 relative to the concentration of
O2 (i.e., [O2] ≪ [Me2Fc], [HClO4]). A representative example
of the UV−vis spectral change in the catalytic reduction of O2
by Me2Fc with CoII(Ch2) is shown in Figure 5, where the
formation of Me2Fc+ was monitored by a rise in absorbance at
650 nm. The time profile of the formation of Me2Fc+ in the
reduction of limited concentration of O2 (1.7 × 10−3 M) is
shown in Figure 5b, where two-equivalent of Me2Fc+ (3.4 ×
10−3 M) to the concentration of O2 dissolved in PhCN was
formed at the end of catalytic reaction. This result indicates that
two-electron reduction of O2 to produce H2O2 catalyzed by
CoII(Chn) (n = 1−3) was successfully achieved with nearly
100% selectivity in the PhCN solution, as given by eq 2. The
formation of H2O2 was confirmed by the iodometric titration
(see Experimental Section).
Kinetics and Mechanism of Two-Electron Reduction
of O2 Catalyzed by CoII(Chn). The mechanism of catalytic O2
reduction with CoII(Chn) (n = 2, 3) was determined by
investigating the dependence of the observed rate constant for
the formation of Me2Fc+ on concentrations of CoII(Chn) (n =
2, 3), Me2Fc, HClO4, and O2.18 The kobs values increased
linearly with increasing concentration of CoII(Chn) (n = 2, 3)
(Figure 6) and O2 (Figures 7c and S9c in the Supporting
Information). In contrast, the kobs values remained constant
irrespective of the change in the concentration of Me2Fc
(Figures 7b and S9b in the Supporting Information). However,
in the case of CoII(Chn) (n = 2, 3), plots of kobs values vs the
concentration of HClO4 exhibited a saturation behavior
(Figures 7a and S9a in the Supporting Information).51 Such a
saturation behavior of kobs values on the concentration of
HClO4 together with the first-order dependence of kobs values
on the concentration of CoII(Chn) (n = 2, 3) and O2 is
explained by a change in the rate-determining step from
proton-coupled electron transfer (PCET) from CoII(Chn) (n =
2, 3) to O2 to produce [CoIII(Chn)]+ (n = 2, 3) and HO2 to
•
deprotonation of [CoII(ChnH)]+ to form CoII(Chn) at large
concentrations of HClO4 as shown in Scheme 2.
O2 + 2H+ + 2Me2Fc → H2O2 + 2Me2Fc+
(2)
The rate of the formation of Me2Fc+ obeyed pseudo-first-
order kinetics under the conditions of [O2] ≪ [Me2Fc],
[HClO4], suggesting that the oxidation of CoII(Chn) by O2 is
rate-determining step and the catalytic rate has first-order
dependence on the concentration of O2.
Thus, the kinetic equation for the reduction of O2 catalyzed
by both CoII(Ch2) and CoII(Ch3) is given by eq 3,
d[Me2Fc+]/dt = kcat(n)[CoII(Chn)][O2][HClO ]
4
(3)
where kcat(n) (n = 2, 3) is catalytic rate constant with CoII(Chn)
(n = 2, 3). From the protonation equilibrium between
CoII(Chn) and [CoII(ChnH)]+ (eq 4),
Figure 6 presents the comparison of the dependence of
observed rate constant (kobs) under the same pseudo-first-order
E
DOI: 10.1021/ic502678k
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
−1
which predicts a linear correlation between kobs(n) and
[HClO4]−1 as shown in Figure S12 in the Supporting
Information. The kcat(n) and K values were determined from
−1
the slope and intercept of the linear plot of kobs(n) vs
[HClO4]−1 to be 9.6 × 106 M−2 s−1 and 1.8 × 102 M−1 for
CoII(Ch2) and 2.2 × 107 M−2 s−1 and 78 M−1 for CoII(Ch3),
respectively. These K values agree with the results determined
from the protonation of CoII(Chn) titrated with HClO4 as
shown in Figure S4 in the Supporting Information. Such
agreement of the K values determined from the titration of
CoII(Chn) (n = 2, 3) with HClO4 and catalytic conditions
provides evidence of the deprotonation of [CoII(ChnH)]+ to
afford CoII(Chn) (n = 2, 3) as an intermediate in a pre-
equilibrium step that is followed by rate-determining inner-
sphere electron transfer from CoII(Chn) to O2 coupled with
proton transfer. This is also consistent with the conclusion that
[CoIII(Chn)]+ (n = 2, 3) hardly undergo the protonation as
observed in Figure S5 in the Supporting Information, and
hence, no change of the redox potential of [CoIII(Chn)]+/
CoII(Chn) (n = 2, 3) was observed as mentioned in Figure 4.52
Thus, the introduction of electron-withdrawing substituents on
chlorin ligand resulted in the positive shift of redox potential,
resulting in the rejection of protonation in the presence of
HClO4 due to lower basicity. In contrast to CoII(Chn) (n = 2,
3), [CoII(Ch1H)]+ formed by the protonation of CoII(Ch1) in
the presence of a large concentration of HClO4 also acts as a
reactive intermediate in catalytic O2 reduction. However, one-
electron oxidation potential of CoII(Ch1) was changed from
0.37 V (vs SCE) to 0.48 V due to the protonation, leading
slower catalysis. As a result, such a change in redox property
resulted in the enhancement of the catalytic reactivity, where
the observed rate constant (kobs) value of CoII(Ch3) was 36-fold
larger than that of CoII(Ch1).53
Figure 7. Plots of kobs vs (a) [HClO4], (b) [Me2Fc], and (c) [O2] for
the two-electron reduction of O2 in PhCN at 298 K. Typical
conditions: CoII(Ch2) (4.0 × 10−6 M), O2 (1.7 × 10−3 M), Me2Fc (2.5
× 10−2 M), and HClO4 (2.5 × 10−2 M).
Scheme 2. Catalytic Mechanism of O2 Reduction with
CoII(Chn), n = 2, 3
CONCLUSION
■
We have conclusively demonstrated effects of changes in redox
potential or configuration of cobalt chlorin complexes
(CoII(Chn) (n = 1−3)) on the catalytic mechanism and activity
of two-electron reduction of dioxygen (O2). All of the cobalt
chlorin complexes (CoII(Chn) (n = 1−3)) efficiently and
selectively catalyzed two-electron reduction of O2 by Me2Fc to
produce H2O2 in the presence of HClO4 in PhCN at 298 K.
The detailed kinetic studies have revealed that the rate-
determining step in the catalytic cycle is the proton-coupled
electron transfer from CoII(Chn) (n = 2, 3), which is in
protonation equilibrium between [CoII(ChnH)]+, to O2 to
[CoII(Chn)] = [CoII(Chn)]t /K[HClO ]
the concentration of CoII(Chn) in eq 3 is rewritten by eq 5,
(4)
4
d[Me2Fc+]/dt = kcat(n)[CoII(Chn)]t[O2][HClO ]
4
/(1 + K[HClO ])
(5)
4
where CoII(Chn)t are the total concentrations of CoII(Chn) and
[CoII(ChnH)]+. Thus, by combining eqs 3 and 4, we obtain eq
5. The kobs(n) (n = 2, 3) value is given by eq 6,
produce [CoIII(Chn)]+ (n = 2, 3) and HO2 . The introduction
•
of electron-withdrawing substituents on chlorin ligand resulted
in the enhancement of catalytic reactivity, which is attributed to
the more negative redox potential of [CoIII(Chn)]+/CoII(Chn)
(n = 2, 3) in the presence of HClO4 due to lower acceptability
of proton. In summary, the present study provides valuable
insights toward the development of more efficient catalysts for
the selective two-electron reduction of O2 to produce H2O2,
which is a promising candidate for a renewable energy source.
kobs(n) = kcat(n)[CoII(Chn)]t[HClO ]/(1 + K[HClO ])
4
4
(6)
which is consistent with saturation behaviors as a function of
the concentration of HClO4 as shown in Figures 7a and S9a in
the Supporting Information.
In order to determine the kcat(n) values, eq 6 is rewritten by
eq 7,
ASSOCIATED CONTENT
■
kobs(n)−1 = 1/kcat(n)[CoII(Chn)]t[HClO ]
S
* Supporting Information
Spectroscopic and kinetic data. This material is available free of
charge via the Internet at http://pubs.acs.org.
4
+ K/kcat(n)[CoII(Chn)]t
(7)
F
DOI: 10.1021/ic502678k
Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry
Article
(17) Liu, S.; Mase, K.; Bougher, C.; Hicks, S. D.; Abu-Omar, M. M.;
Fukuzumi, S. Inorg. Chem. 2014, 53, 7780−7788.
AUTHOR INFORMATION
Corresponding Author
*E-mail: fukuzumi@chem.eng.osaka-u.ac.jp.
Notes
■
(18) Mase, K.; Ohkubo, K.; Fukuzumi, S. J. Am. Chem. Soc. 2013,
135, 2800−2808.
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The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
This work was supported by an Advanced Low Carbon
Technology Research and Development (ALCA) program
from Japan Science Technology Agency (JST) to S.F. and by
Japan Society for the Promotion of Science (JSPS: 26620154
and 26288037 to K.O. and 25·727 to K.M.).
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