Selective Electrocatalysis of a Water-Soluble Rhenium(I) Complex for CO2 Reduction Using Water As an Electron Donor

Article abstract of DOI:10.1021/acscatal.7b03275

Reduction of CO₂ using water as an electron donor could be deemed one of the most important reactions in artificial photosynthesis. In the current study, electrochemical properties of a water-soluble rhenium(I) tricarbonyl complex with hydroxymethyl groups ([Re(CH₂OH)-OH₂]⁺) and its electrocatalysis for CO₂ reduction were investigated in aqueous solutions. In comparison with general rhenium(I) carbonyl complexes, [Re(CH₂OH)-OH₂]⁺ shows high solubility in water even in the presence of both CO₂ and OH^-. In an aqueous solution, [Re(CH₂OH)-OH₂]⁺ exhibited a catalytic current due to CO₂ reduction in a cyclic voltammogram at a positive potential of about 400 mV more than that in a DMF solution. It is suggested that two molecules of a one-electron-reduced Re complex participated in reducing one molecule of CO₂ in the aqueous solution. Electrolysis using [Re(CH₂OH)-OH₂]⁺ at -1.1 V vs NHE in a CO₂-saturated aqueous solution afforded CO and HCOOH with 95 and 4% selectivity, respectively, at pH 6.9; thermodynamically favorable H₂ evolution via water reduction was almost completely suppressed. Even at pH 4.2, the selectivity of CO₂ reduction was still high (CO 84%; HCOOH 3%). A full cell consisting of a cathode unit with [Re(CH₂OH)-OH₂]⁺ for CO₂ reduction and an IrO₂/FTO anode for water oxidation electrocatalytically produced both CO and O₂ as the main products with high efficiency and selectivity.

Full text of DOI:10.1021/acscatal.7b03275

Research Article
pubs.acs.org/acscatalysis
Cite This: ACS Catal. 2018, 8, 354−363
Selective Electrocatalysis of a Water-Soluble Rhenium(I) Complex for
CO2 Reduction Using Water As an Electron Donor
Akinobu Nakada† and Osamu Ishitani*
Department of Chemistry, School of Science, Tokyo Institute of Technology, 2-12-1-NE-1, O-okayama, Meguro-ku, Tokyo 152-8550,
Japan
S
* Supporting Information
ABSTRACT: Reduction of CO2 using water as an electron donor
could be deemed one of the most important reactions in artificial
photosynthesis. In the current study, electrochemical properties of a
water-soluble rhenium(I) tricarbonyl complex with hydroxymethyl
groups ([Re(CH2OH)−OH2]+) and its electrocatalysis for CO2
reduction were investigated in aqueous solutions. In comparison
with general rhenium(I) carbonyl complexes, [Re(CH2OH)−
OH2]+ shows high solubility in water even in the presence of
both CO2 and OH−. In an aqueous solution, [Re(CH2OH)−OH2]+
exhibited a catalytic current due to CO2 reduction in a cyclic
voltammogram at a positive potential of about 400 mV more than
that in a DMF solution. It is suggested that two molecules of a one-
electron-reduced Re complex participated in reducing one molecule
of CO2 in the aqueous solution. Electrolysis using [Re(CH2OH)−OH2]+ at −1.1 V vs NHE in a CO2-saturated aqueous solution
afforded CO and HCOOH with 95 and 4% selectivity, respectively, at pH 6.9; thermodynamically favorable H2 evolution via
water reduction was almost completely suppressed. Even at pH 4.2, the selectivity of CO2 reduction was still high (CO 84%;
HCOOH 3%). A full cell consisting of a cathode unit with [Re(CH2OH)−OH2]+ for CO2 reduction and an IrO2/FTO anode
for water oxidation electrocatalytically produced both CO and O2 as the main products with high efficiency and selectivity.
KEYWORDS: CO2 reduction, electrocatalysis, water-soluble rhenium complex, water oxidation, aqueous solution
our knowledge, selective CO production has not been achieved
in acidic aqueous solutions to date.
INTRODUCTION
■
Catalytic conversion of CO2 into useful energy-rich chemicals
using sustainable energy is one of the most promising
approaches to solve global warming and problems of fossil
fuels shortage. In this context, electrocatalytic reduction of CO2
to produce C1 chemicals has been widely investigated. CO, one
of the two-electron reduction products of CO2, is attractive as a
target product because it is gaseous and has low solubility in
aqueous media, making it easier to collect in comparison with
the liquid products; in addition, CO can be converted to
hydrocarbons by Fischer−Tropsch synthesis. For the purposes
described above, electrocatalytic CO2 reduction should proceed
with water as an abundant reductant and reaction medium.
Various metal- and carbon-based electrodes have been
investigated for CO2 reduction. Although some transition-
metal electrodes, such as Ag and Au, exhibited relatively high
selectivity for CO formation,1 H2 evolution from water also
occurred as a side-reaction in the aqueous solutions. Cu
electrodes can reduce CO2 directly to hydrocarbons,1 and a
boron-doped diamond electrode that induced CO2 reduction,
giving formaldehyde2 and methanol,3 in aqueous solutions was
recently reported. Molecular pyridine catalyst has also been
reported to electrochemically produce formic acid and
methanol.4,5 However, the selectivity of the CO2 reduction
was still not high, especially in aqueous solutions. To the best of
On the contrary, in organic solutions, selective electro-
catalytic reduction of CO2 to CO was achieved using molecular
catalysts, particularly metal complexes.6−28 However, most of
them have not been investigated in aqueous media. One of the
problems is the low solubility of these electrocatalysts in water.
In addition, competition with thermodynamically favorable
water reduction makes it difficult to achieve selective formation
of CO from CO2. [Ni(cyclam)]2+ is one of the rare
electrocatalysts to be investigated in water.29−32 However, it
has exhibited relatively low selectivity for CO2 reduction, except
when a toxic mercury pool was used as a working electrode, on
which the [Ni(cyclam)]2+ got adsorbed.31,33 Although a Ru(II)
terpyridine complex was also an active homogeneous electro-
catalyst for CO2 reduction in aqueous solutions, selectivity of
CO formation was not high (up to 66% even at pH 7.2) due to
competitive H2 evolution.34 Very recently, selective CO
production in aqueous solutions has been demonstrated with
iron-porphyrin complexes as electrocatalysts.35−38 However,
high selectivity was achieved only at pH ≈ 7. In an acidic
Received: September 25, 2017
Revised: November 15, 2017
© XXXX American Chemical Society
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Figure 1. UV−vis absorption spectra of (a) [Re(H)−OH2]+ and (c) [Re(CH2OH)−OH2]+ in Britton−Robinson buffer solutions at various pHs.
The total concentration of the Re complexes was 0.1 mM. (b,d) pH dependences of absorbance at 318 (red circles) and 400 nm (blue circles), where
the solid lines are the sigmoidal fitting based on the Henderson−Hasselbalch equation in the acid−base equilibrium between (b) [Re(H)−OH2]+
and Re(H)−OH and (d) [Re(CH2OH)−OH2]+ and Re(CH2OH)−OH.
condition (pH 3.7), the selectivity of CO formation drastically
decreased because of the exclusive production of H2; it is still
challenging to develop selective and efficient electrocatalysts for
CO2 reduction to CO even in acidic aqueous solutions.
electrocatalysts and photocatalysts with Re complexes that can
be used in water. In addition, it has been reported that Re(I)
tricarbonyl complexes convert to a bicarbonato complex
Re(N^N)(CO)3{OCO(O)OH}, with very low solubility in
most solvents including water, through ligand substitution
reaction of the monodentate ligand with OH− followed by CO2
insertion.45,46 This also makes it difficult to investigate the
electrocatalytic properties of rhenium(I) complexes in an
aqueous solution.
In the current study, a new water-soluble Re(I) complex fac-
[Re(4,4′-dihydroxymethyl-2,2′-bipyridine)(CO)3(OH2)](OTf)
([Re(CH2OH)−OH2]+) was synthesized as a CO2 reduction
electrocatalyst. Interestingly, this Re(I) complex did not
convert to the bicarbonato complex even in aqueous solutions
containing OH− and CO2. It exhibited catalytic CO2 reduction
with a positive potential by about 400 mV more in an aqueous
solution than that in a DMF solution. Using [Re(CH2OH)−
OH2]+ as the molecular catalyst, electrochemical reduction of
CO2 to CO was achieved with high selectivity and high faradaic
efficiency even in acidic aqueous solution. Finally, a full
electrochemical cell was constructed with [Re(CH2OH)−
OH2]+ for CO2-to-CO conversion using water as an electron
donor.
Among the metal complex electrocatalysts, rhenium(I)
complexes fac-[Re(N^N)(CO)3L]n+ (N^N = diimine, L =
monodentate ligands) have particularly attracted attention as
efficient electrocatalysts for CO2 reduction.6,7,14,16−18,22 In
organic solutions, the Re(I) complexes exhibit excellent
selectivity for CO production over the formation of H2 and
HCOOH with high faradaic efficiency (>95%) and reasonable
turnover frequency (TOF) of up to 200 s−1.6 Addition of weak
Brønsted acid, such as water, into the organic solution
reportedly enhanced the catalytic current of CO2 reduction.14,16
An 18-electron species [Re(bpy-tBu)(CO)3]− (bpy-tBu: 4,4′-
tertialybutyl-2,2′-bipyridine), produced via two-electron reduc-
tion following ligand loss, which more efficiently reacted with
CO2 than proton, was reported.14 Furthermore, we recently
developed a binuclear complex consisting of a [Re(N^N)-
(CO)3Cl]-type catalyst and a [Ru(N^N)3]2+-type photo-
sensitizer unit connected by an ethylene chain that photo-
catalytically reduced CO2 with high selectivity even in an
aqueous solution.39,40 These outcomes strongly suggest the
possibility of using rhenium(I) complexes as a catalyst for
reducing CO2 to CO even in aqueous solutions. However, only
a few water-soluble rhenium(I) bipyridine carbonyl complexes
have been reported;41−44 this is a huge barrier for developing
RESULTS AND DISCUSSION
Synthesis of [Re(CH2OH)−OH2]+ and Its Properties in
Aqueous Solutions. The aquo complex [Re(CH2OH)−
■
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OH2]+ was synthesized by refluxing an aqueous solution
containing Re(4,4′-hydroxymethyl-2,2′-bipyridine)-
(CO)3(OTf) for 25 h; the isolated yield was 67%. The
properties and solubility of [Re(CH2OH)−OH2]+ and the
corresponding complex with a 2,2′-bipyridine ligand, i.e., fac-
[Re(bpy)(CO)3(OH2)]+ ([Re(H)−OH2]+), were investigated
in aqueous solutions.
Figure 1a shows UV−vis absorption spectra of [Re(H)−
OH2]+ in Britton−Robinson buffer solutions at various pHs;
the absorption spectra changed depending on the pH, with
isosbestic points at 306 and 366 nm. This is attributed to an
acid−base equilibrium between [Re(H)−OH2]+ and Re(H)−
OH (eq 1).
[Re(N^N)(CO)3(OH2)]+
⇄ Re(N^N)(CO)3(OH) + H+
Figure 2. UV−vis absorption spectra of [Re(CH2OH)−OH2]+ (0.1
mM) in aqueous Na2CO3 solutions (0.1 M) before (blue, pH 11.0)
and after (red, pH 7.1) bubbling with CO2. The broken and dotted
lines show expected spectra at pH 7.1 and 11.0 based on the acid−base
equilibrium between [Re(CH2OH)−OH2]+ and Re(CH2OH)−OH.
(1)
Figure 1b shows the pH dependence of the absorption at 318
and 400 nm; their sigmoidal fitting using the Henderson−
Hasselbalch equation provided a pKa value of 8.5.
[Re(CH2OH)−OH2]+ also showed similar spectral changes
with pKa = 8.5 (Figure 1c,d). This clearly indicates that these
spectral changes can also be attributed to deprotonation of the
aqua ligand (eq 1). The hydroxyl groups of the 4,4′-
dihydroxymethyl-2,2′-bipyridine ligand should not be deproto-
nated in this pH range because of the high pKa values of the
hydroxymethyl groups; for example, pKa is 15.4 in the case of
benzyl alcohol.47
OCO2H cite because Re(H)−OCO2H is removed from the
solution by precipitation. The hydroxymethyl group should
increase the solubility of the Re(CH2OH)−OCO2H in an
aqueous solution. The high solubility of both [Re(CH2OH)−
OH2]+ and Re(CH2OH)−OH makes it possible to investigate
the electrochemical properties of the Re complex in various
aqueous solutions in the presence of CO2.
On the other hand, under a CO2 atmosphere, these two
complexes exhibited very different solubility in aqueous
solutions. Bubbling in an aqueous Na2CO3 (0.1 M) solution
containing [Re(H)−OH2]+ rapidly formed a yellow precipitate,
which was assigned to the corresponding bicarbonato complex
Re(bpy)(CO)3{OC(O)OH} (eq 2, Re(H)−OCO2H) by its
FT-IR spectrum (Figure S1).45,46 It has been reported that the
solubility of this bicarbonato complex was very low not only in
aqueous solutions but also in various organic solvents.
Electrochemistry. Figure 3a shows the cyclic voltammo-
grams (CVs) of [Re(CH2OH)−OH2]+ (0.5 mM) measured in
an aqueous solution containing Na2SO4 (0.1 M) as the
electrolyte. Two irreversible waves were observed at Ep = −1.15
and −1.31 V vs NHE in a scan to the negative direction under
an Ar atmosphere, and a reverse scan to the positive direction
showed one irreversible peak at Ep = +0.06 V, which is
assignable to oxidation of a Re0−Re0 dimer complex according
to previous reports on electrochemistry of [Re(N^N)-
(CO)3(Sol)]+ (Sol = solvent ligand) in organic solutions.22
The first-reduction wave was attributed to a diimine-ligand-
based reduction (eq 3). The irreversibility of the first-reduction
wave and the appearance of the oxidation wave of the Re0−Re0
dimer complex strongly indicated that the aqua ligand was
rapidly dissociated from the one-electron-reduced complex (i.e.,
[Re(N^N)(CO)3(OH2)]•), and the 17-electron species [Re-
(N^N)(CO)3]• was produced (eq 4). The 17-electron species
underwent dimerization (eq 5).48
Re(bpy)(CO)3(OH) + CO2
→ Re(bpy)(CO)3{OC(O)OH}↓
(2)
In the current study, [Re(CH2OH)−OH2]+ was found to be
still dissolved in the Na2CO3 solution even after bubbling with
CO2, and as described in the electrochemical measurements,
even 0.5 mM of total concentration of the complex could be
dissolved in aqueous solutions at pH 4.2 and 6.9 under a CO2
atmosphere. Figure 2 shows the absorption spectra of
[Re(CH2OH)−OH2]+ (0.1 mM) in an aqueous Na2CO3
solution with and without CO2 bubbling, where the pHs of
the solutions were 7.1 and 11.0, respectively. On the basis of
the acid−base equilibrium, both spectra could be reproduced
by linear combination of the spectra of [Re(CH2OH)−OH2]+
and Re(CH2OH)−OH, which were obtained in the absence of
CO2 at pH 4.1 and 11.5, respectively (red and blue lines in
Figure 1c). These results indicate that [Re(CH2OH)−OH2]+
and/or Re(CH2OH)−OH were the main species even in an
aqueous solution containing Na2CO3 under a CO2 atmosphere;
in other words, the equilibrium constant of formation of
Re(CH2OH)−OCO2H, probably Re(H)−OCO2H as well, is
small in the solution. The difference in formation of the
bicarbonato complex can possibly be owed to the difference in
their solubility. The equilibria with [Re(H)−OH2]+ and
Re(H)−OH should be continuously moved to the Re(H)−
[Re(N^N)(CO)3(OH2)]+ + e−
→ [Re(N^N)(CO)3(OH2)]•
(3)
[Re(N^N)(CO)3(OH2)]•
rapid
⎯⎯⎯⎯⎯⎯→ [Re(N^N)(CO)3]• + H2O
(4)
2[Re(N^N)(CO)3]•
→ [(CO)3(N^N)Re−Re(N^N)(CO)3]
(5)
This discussion is supported by the fact that the dimer
oxidation wave was similarly observed even when the scan
reversed at the first reduction potential with −1.2 V (Figure
S2). The second-reduction wave was assignable to the
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Figure 3. CVs of [Re(CH2OH)−OH2]+ (0.5 mM) in (a) a Na2SO4 (0.1 M) aqueous solution at pH 4.2 and (b) an Et4NBF4 (0.1 M) DMF solution
measured under CO2 (red) and Ar (blue) atmospheres with a scan rate of 0.1 V s−1. The black dotted lines show background scans without
[Re(CH2OH)−OH2]+ under a CO2 atmosphere.
Scheme 1. Proposed Mechanisms of CO2 Reduction via the One-Electron46,50 and Two-Electron Reduction51−53 Pathways in
Organic Solutions
generation of two-electron-reduced species [Re(N^N)(CO)3]−
through reduction of the Re dimer (eq 6).
clearly indicate that the substitution of the aqua ligand with a
DMF molecule took place in a rapid manner. Figure 3b shows
CV measured in a DMF solution containing NEt4BF4 (0.1 M)
as the electrolyte. Under an Ar atmosphere, although two
reduction waves were observed in the CV measured in the
DMF solution as well as the case using the aqueous solution,
the shape was very different from that measured in the aqueous
solution (compare the blue lines in Figure 3a,b). In DMF, the
first-reduction wave was reversible with E1/2 = −1.09 V vs
NHE, and the second-reduction wave was irreversible with Ep =
−1.26 V. It is noteworthy that the peak attributed to the
oxidation of the Re0−Re0 dimer was not observed in DMF even
under the Ar atmosphere. The reversibility of the first-reduction
wave implies that the coordination of the DMF ligand to the
one-electron-reduced species of the Re complex was more
stable than that of the aqua ligand (eqs 7 and 8). The stability
of one-electron-reduced species of the Re complexes against
dissociation of a monodentate ligand has also been reported in
[(CO)3(N^N)Re−Re(N^N)(CO)3] + 2e−
→ 2[Re(N^N)(CO)3]−
(6)
Similar electrochemical measurements were carried out in a
DMF solution, which has often been used as a solvent for
electrochemical CO2 reduction. Before investigating the
electrochemistry, we confirmed the stability of [Re-
(CH2OH)−OH2]+ in DMF. The FT-IR spectrum of [Re-
(CH2OH)−OH2]+ (0.5 mM), just after being dissolved into
DMF, was the same as the DMF complex [Re(CH2OH)−
DMF]+, and this spectrum did not change after several hours
(Figure S3). ESI-MS spectrum of [Re(CH2OH)−OH2]+ in
DMF solution showed a strong signal at m/z = 560 assignable
to [Re(CH2OH)−DMF]+ with a small signal of [Re-
(CH2OH)−MeOH]+ of which the MeOH ligand should be
supplied by the mobile phase (Figure S4). These observations
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Figure 4. Time courses of CO, HCOOH, and H2, along with passed electrons during electrolysis at −1.1 V vs NHE in an aqueous Na2SO4 (0.1 M)
solution under a CO2 atmosphere. (a) With [Re(CH2OH)−OH2]+ (0.5 mM) and without NaOH (pH 4.2). (b) Without [Re(CH2OH)−OH2]+
and without NaOH (pH 4.2). (c) With [Re(CH2OH)−OH2]+ (0.5 mM) and with 0.1 M NaOH (pH 6.9). (d) Without [Re(CH2OH)−OH2]+ and
with 0.1 M NaOH (pH 6.9). (e) Same condition as (c), except for using D2O and NaOD instead of H2O and NaOH, respectively.
the cases of other [Re(N^N)(CO)3L]n+-type complexes (L =
Cl−, MeCN).22,46,48,49
solution, which was observed under the Ar atmosphere at +0.06
V (Figures 3a and S3). This strongly suggests that rapid CO2
binding with the catalyst proceeded before dimerization of
[Re(N^N)(CO)3]•.
[Re(N^N)(CO)3(DMF)]+ + e−
→ [Re(N^N)(CO)3(DMF)]•
Two mechanistic pathways have been proposed for electro-
chemical CO2 reduction using Re complexes as an electro-
catalyst, which proceed through one- and two-electron
reduction, respectively.7 In the former case, formation of a Re
dimer in which one CO2 molecule bridges two molecules of
[Re(N^N)(CO)3]• has been reported to be a key intermediate
(left-hand cycle in Scheme 1).46,51−53 However, the one-
electron reduction pathway would be slow in organic solutions,
if it occurs, because the first-reduction wave is usually reversible
even under the CO2 atmosphere, as mentioned above. In fact,
most of the Re(I) tricarbonyl complexes exhibited catalytic
current enhancement of CO2 reduction only at the second-
reduction potential in organic solutions. In the latter case
(right-hand cycle in Scheme 1), formation of a carboxylate
complex has been proposed as a key intermediate.46,50 We
cannot identify the exact reaction mechanism(s) of electro-
catalytic CO2 reduction in aqueous solutions at this stage.
However, it is reasonable to conclude that the remarkable
positive shift (ca. 400 mV) of the onset potential of the catalytic
CO2 reduction could take place because of the rapid
dissociation of the aqua ligand just after one-electron reduction
of the Re complex.
(7)
[Re(N^N)(CO)3(DMF)]•
slow
⎯⎯⎯⎯⎯→ [Re(N^N)(CO)3]• + DMF
(8)
The destabilization of the aqua ligand in the one-electron-
reduced species is one of the unique properties of the aquo
Re(I) complex.
Although catalytic current enhancement was observed in
both solutions under the CO2 atmosphere (red lines in Figure
3), the potentials when the catalytic current reached 0.25 mA
cm−2 were significantly different for each solution. In the
aqueous solution, it was ∼−1.0 V, which corresponds to the
first reduction of [Re(CH2OH)−OH2]+. On the other hand, a
two-electron reduction of [Re(CH2OH)−DMF]+ was required
for the appearance of the catalytic wave, where the potential
when the catalytic current reached 0.25 mA cm−2 was ∼−1.4 V
in the DMF solution. Surprisingly, the onset potential for CO2
reduction electrocatalysis was positively shifted by about 400
mV using the aquo complex rather than the DMF complex. It is
noteworthy that, under the CO2 atmosphere, the oxidation
wave of the Re0−Re0 dimer disappeared in the aqueous
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Bulk Electrolysis. Constant-potential electrolysis of a
Na2SO4 (0.1 M) aqueous solution containing [Re-
(CH2OH)−OH2]+ (0.5 mM) was conducted under the CO2
atmosphere at E = −1.1 V vs NHE, where the pH of the
solution was 4.2. Continuous production of CO was observed
as the main product during electrolysis for 3 h, accompanied by
small amounts of H2 and HCOOH (Figure 4a). The
selectivities of CO formation and CO2 reduction (i.e., CO +
HCOOH) were 84 and 87% even under acidic conditions. CO
generation was not observed during electrolysis in the absence
of the Re(I) catalyst, but H2 was directly produced from the
glassy carbon electrode (Figure 4b). This suggests that the
actual selectivity of [Re(CH2OH)−OH2]+ for CO2 reduction
should be higher than the as-observed result in Figure 4a.
Higher pH of the reaction solution (pH 6.9), which was
achieved by the addition of NaOH (0.1 M), improved the
selectivity of CO formation and CO2 reduction to 95 and 99%
(Figure 4c) because of the suppression of H2 evolution from
the glassy carbon electrode (Figure 4d). The faradaic efficiency
of the total reduction products was close to 100% in all cases.
Because the turnover number of the evolved CO was larger
than 5, on the basis of the total amount of [Re(CH2OH)−
OH2]+ used, it is clearly demonstrated that this Re complex can
work as a selective catalyst for CO2 reduction even in acidic
aqueous solutions.
It was reported that electrocatalytic reduction of CO2 to CO
proceeded with more than 90% selectivity using Ag,54 Au,55 and
Cu−In56 metal electrodes, as well as with a Fe-complex
molecular catalyst35 in aqueous solutions. However, this high
selectivity was achieved only at pH ≈ 7. Although the Fe-
complex systems were tested in acidic conditions (pH = 3.7),
the selectivity of the CO formation drastically decreased due to
the preferential hydrogen production. Therefore, to the best of
our knowledge, this is the first electrocatalytic system that
proceeds via highly selective CO formation from CO2 even in
acidic aqueous solutions.
Although a few full electrochemical cells with water oxidation
were reported, the water oxidation product (i.e., O2) was not
confirmed in most cases,57,58 except for a report by Saveant et
al.38 Therefore, we attempted to combine the electrocatalytic
CO2 reduction cathodic system using [Re(CH2OH)−OH2]+
with an anodic cell for water oxidation. A fluorine-doped tin
oxide film coated with IrO2 (IrO2/FTO) was employed as the
anode; IrO2 is a well-known water oxidation catalyst.59 Figure 5
Figure 5. Linear sweep voltammograms of IrO2/FTO and Pt wire
electrodes in an aqueous solution containing [Re(CH2OH)−OH2]+
(0.5 mM), Na2SO4 (0.1 M), and NaOH (0.1 M) under a CO2
atmosphere (pH 6.9).
shows the linear sweep voltammogram of the IrO2/FTO in a
CO2-saturated aqueous solution (pH 6.9). Large anodic current
attributable to water oxidation was observed at around 1.1 V vs
NHE, which was lower by 300 mV than that using a Pt-wire
electrode.
Figure 6 illustrates the full-cell configuration where two
compartments were separated by a Nafion membrane. These
It has been reported that addition of proton sources into
reaction solutions enhances the rate of electrocatalytic CO2
reduction using Re(I) complexes in organic solutions.14,16 To
obtain information on the roles of protons, the kinetic isotope
effect (KIE) on the electrocatalysis of [Re(CH2OH)−OH2]+
was investigated in an aqueous solution at pH = 6.9. In a D2O
solution containing Na2SO4 (0.1 M), NaOD (0.1 M), and
[Re(CH2OH)−OH2]+ (0.5 mM), electrolysis at −1.1 V also
produced CO as the main product with high selectivity of 95%
(Figure 4e). However, the rate of CO formation drastically
decreased in comparison with that of the corresponding
experiment that uses H2O and NaOH instead of D2O and
NaOD (Figure 4c,e). The KIE factor iH2O/iD2O was ∼6, where
iH2O and iD2O are the stable catalytic currents during electrolysis
in the H2O and D2O solutions, respectively (Figure S5).
Interestingly, lowering the pH of the reaction solution from 6.9
to 4.2 (i.e., increasing the proton concentration) did not
increase the rate of CO2 reduction (Figure 4a,c). These results
suggest that a solvent isotope effect not related to proton
activity is involved in the rate-determining step following CO2
activation.
Figure 6. Schematic illustration of the full-cell configuration for
reduction of CO2 using water as the electron donor: the cathodic
chamber consists of [Re(CH2OH)−OH2]+ and a glassy carbon
electrode, and the IrO2/FTO anode is used in the anodic chamber.
Both chambers were filled with an aqueous solution containing
Na2SO4 (0.1 M), NaOH (0.1 M), and CO2.
two compartments were filled with the same aqueous solution
containing Na2SO4 (0.1 M) and NaOH (0.1 M) under a CO2
atmosphere (pH 6.9), except that [Re(CH2OH)−OH2]+ (0.5
mM) was added only into the cathodic compartment;
therefore, there was no chemical bias between the two
electrodes. A stable current of ∼0.3 mA cm−2 was observed
at a cell voltage of 2.5 V. At the same time, both the reduction
Electrocatalytic CO2 Reduction Coupled with Oxida-
tion of Water. Electrochemical CO2 reduction systems as a
countermeasure to global warming should be combined with
renewable energy sources such as solar batteries and anodes
that can oxidize water. However, most reports on electro-
catalytic CO2 reduction describe only the cathodic half reaction.
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and oxidation products (i.e., CO, HCOOH, and H2 in the
cathodic compartment and O2 in the anodic compartment)
were generated (Figure 7); 13.8 μmol of CO, 1.5 μmol of
1
CO2 + H2O → HCOOH +
O2
(10)
2
1
H2O → H2 +
O2
(11)
2
ΔH0 × F
x
x
εx(%) = 100 ×
Ebias
(12)
where ΔH0x is the enthalpy change of each reaction (ΔH0
=
CO
1.47 eV; ΔH0
= 1.32 eV; ΔH0 = 1.48 eV) and Ebias is
HCOOH
H2
the cell voltage (2.5 V).
We can now compare this result with the rarely reported
electrochemical CO2 reduction systems with a full-cell
configuration. Two full-cell systems consisting of other
molecular catalysts, which were already mentioned in the
Introduction, were also reported. One of the systems uses
Ru(II) terpyridine complex as both CO2 reduction and water
oxidation electrocatalysts; this system generates CO at a cell
voltage of 2.8 V with εCO = 17% (the selectivity of CO
formation and total ε were 33 and 53%, respectively).34
Recently, high εCO and FCO (50 and 90%, respectively) were
reported at pH 7.3 in aqueous solutions using an iron-
porphyrin catalyst combined with phosphate cobalt oxide as a
water oxidation anode at a cell voltage of 2.5 V.38 However, the
selectivity of CO2 reduction using this complex drastically
decreased in the acidic aqueous solution, while the selectivity
can be retained using [Re(CH2OH)−OH2]+ even in acidic
conditions. Another study reports a full cell consisting of metal
electrodes (i.e., a Ag cathode and a Pt/Ir anode); the full cell
was reported to exhibit relatively high selectivity for CO
Figure 7. Time course of produced CO, HCOOH, H2, and O2 with
passed electron by electrolysis at a cell voltage of 2.5 V in the full-cell
configuration illustrated in Figure 6.
HCOOH, 0.4 μmol of H2, and 9.6 μmol of O2 were generated
after 2 h of electrolysis followed by waiting for 1 h after the
external voltage was removed, which was necessary due to
diffusion of the gaseous products into the gas phase of the
compartments. The selectivity of the CO formation was 88%
based on the reduction products (the selectivity of CO2
reduction was 98%). The turnover number of the CO
formation based on the total amount of [Re(CH2OH)−
OH2]+ used was 2.3 after the 2 h of electrolysis. The faradaic
efficiency of O2 formation was almost quantitative, while the
total faradaic efficiency of the reduction products (i.e., CO,
HCOOH, and H2) was 81%. This difference might be caused
by air contamination of the cathodic cell. We actually observed
contamination of about 5 μmol of air, which was determined by
measuring N2 in the cathodic compartment at the end of the
electrolysis. On the other hand, air contamination was much
lower in the anodic compartment (<0.6 μmol), and the amount
of air was stable during electrolysis; we subtracted the
contaminated air from the detected O2 to evaluate the actual
O2 generation by electrolysis. Additional current by reduction
of the contaminated O2 as a side-reaction probably lowered the
“apparent” faradaic efficiency in the cathodic compartment.
From these experiments, we concluded that the full cell
consisting of both the cathodic system with the [Re-
(CH2OH)−OH2]+ catalyst and the IrO2/FTO anode can
effectively work for electrochemical CO2 reduction selectively
producing CO using water as the electron donor.
formation (∼80%) and energy conversion efficiency (εCO
≈
40%) in 0.5 M of KHCO3 aqueous solution during the initial
stage of electrolysis. However, this full-cell system was gradually
poisoned just after the electrolysis started; this caused a
decrease in the selectivity of CO formation because of
increasing H2 production.57 In contrast, the [Re(CH2OH)−
OH2]+ in the current study showed a constant rate and
selectivity for CO2 reduction for at least ∼2−3 h (see Figures 4
and 7). In the case of a Ni-modified carbon fiber cathode,58 the
selectivity for CO formation was ∼50% at the cell voltage of
3.054 V. Due to the lower selectivity (the faradaic efficiency for
each product was not mentioned), the energy conversion
efficiency of CO formation was much lower (εCO = ∼24%)
compared with the system using [Re(CH2OH)−OH2]+ (εCO
=
41%). The smaller cell voltage (2.5 V) is also one of the reasons
for the better result in the full cell using [Re(CH2OH)−OH2]+
as the molecular catalyst and IrO2/FTO as the anode. This
clearly indicates the large potential of the molecular catalyst
[Re(CH2OH)−OH2]+ for selective and durable CO2 reduction
in aqueous media.
CONCLUSIONS
■
The water-soluble Re(I) complex, [Re(CH2OH)−OH2]+, was
successfully developed as an electrocatalyst for CO2 reduction
in water. This complex was highly soluble in the aqueous
solution even under a CO2 atmosphere, and it worked as an
efficient selective electrocatalyst for the reduction of CO2 to
CO. This is the first molecular catalyst for CO2 reduction with
high selectivity even in the acidic aqueous solution. The
catalytic reaction proceeded at a much lower overpotential in
the aqueous solution than that in an organic solution because
rapid dissociation of the aqua ligand proceeds from one-
electron-reduced species of [Re(CH2OH)−OH2]+, resulting in
The overall reactions for formation of CO, HCOOH, and H2
accompanied by the formation of O2 in the full-cell system are
represented as eqs 9−11, respectively. Efficiencies of conversion
from electrical energy to chemical energy in each reaction (εx)
were calculated as εCO = 41%, εHCOOH = 4%, and εH2 = 1%
using eq 12 with faradaic efficiencies of FCO = 0.71, FHCOOH
0.08, and FH2 = 0.02, respectively.
=
1
CO2 → CO +
O2
(9)
2
360
DOI: 10.1021/acscatal.7b03275
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ACS Catalysis
Research Article
rapid reaction with CO2. The cathodic system using [Re-
(CH2OH)−OH2]+ was successfully coupled to the IrO2/FTO
anodic system for water oxidation. Selective reduction of CO2
to CO using water as the electron source was demonstrated
using this full-cell configuration with a cell voltage of 2.5 V.
These results clearly indicate the high potential of Re(I)
complex catalysts for CO2 reduction even in aqueous solutions.
4.85 (m, 4H). FT-IR (DMF) νCO/cm−1: 2036, 1934. ESI-MS
(eluent: MeCN) m/z: 528 ([M − PF6]+).
fac-[Re(4,4′-hydroxymethyl-2,2′-bipyridine)(CO)3(DMF)]-
(PF6). This material was synthesized in DMF solution through a
solvent ligand substitution reaction.62 [Re(4,4′-hydroxymethyl-
2,2′-bipyridine)(CO)3(MeCN)](PF6) (0.5 mM) was dissolved
in dry DMF and placed 20 h at room temperature to give the
purpose product. FT-IR (DMF) νCO/cm−1: 2028, 1919, 1908.
ESI-MS (eluent: MeOH) m/z: 560 ([M − PF6]+).
EXPERIMENTAL SECTION
■
General. 1H NMR spectra were recorded on a JEOL
ECA400II (400 MHz) system. Electrospray ionization mass
spectroscopy (ESI-MS) was performed with a Shimadzu
LCMS-2010A system using HPLC-grade MeCN as the mobile
phase. IR and UV−vis absorption spectra were recorded using a
JASCO FT/IR-4600 spectrophotometer with 1 cm−1 resolution
and a JASCO V-565 spectrophotometer.
IrO2/FTO Electrode. A colloidal IrO2 solution was prepared
according to the reported method.63 Na2IrCl6·6H2O (16 mg)
was dissolved in water, and NaOH solution was added to adjust
the pH to 12. The solution was heated to 343 K with stirring
for 30 min and then cooled to room temperature. The pH of
the solution was adjusted to 9 by adding a HNO3 solution. The
solution was further heated to 348 K with stirring for 20 min
and again cooled to room temperature. The obtained blue
solution containing colloidal IrO2 was diluted with water to a
total volume of 100 mL. This IrO2 colloidal solution was again
twice diluted with water just before the following processes.
The diluted solution (50 μL) was casted on a fluorine-doped
tin oxide (FTO) glass film and dried at room temperature. The
obtained film was heated to 573 K in air for 4 h.
Electrochemical Measurements. CVs were measured in a
glass cell containing aqueous solution of Na2SO4 (0.1 M) as a
supporting electrolyte. The voltammograms were recorded
using an ALS/CHI 760ES electrochemical analyzer with a
glassy carbon disk working electrode (3 mm diameter), a Pt
wire counter electrode, and a reference electrode of Ag/AgCl in
saturated KCl solution (+0.198 V vs NHE). For measurements
in DMF solution, an Et4NBF4 (0.1 M) supporting electrolyte
and a Ag/AgNO3 (0.01 M, +0.548 V vs NHE)64 reference
electrode were employed. The electrolyte solution containing a
Re complex catalyst was purged with Ar or CO2 for 20 min
before measuring.
Bulk Electrolysis of CO2 Reduction. Bulk electrolysis for
CO2 reduction was carried out in an O-ring-sealed H-shaped
Pyrex cell, in which two chambers were separated by a Nafion
membrane. A plate of glassy carbon as a working electrode and
Ag/AgCl (sat. KCl) as the reference electrode were installed
into one of the chambers containing an aqueous Na2SO4
solution (13 mL) of [Re(CH2OH)−OH2]+ (0.5 mM). The
geometrically effective surface area of the working electrode was
3.0 cm2. In another chamber, a Pt wire counter electrode was
placed in an aqueous Na2SO4 solution (13 mL). The working
and counter electrode chambers were respectively saturated
with CO2 and Ar by bubbling and sealing. The electrical
potential was applied by a Hokuto Denko HZ-7000
potentiostat.
Synthesis. fac-Re(4,4′-hydroxymethyl-2,2′-bipyridine)-
(CO)3Br. This was prepared according to the reported method
with some modification.60,61 A toluene (10 mL) solution
containing Re(CO)5Br (408 mg, 1.00 mmol) and 4,4′-
hydroxymethyl-2,2′-bipyridine (222 mg, 1.00 mmol) was
refluxed for 3 h. After the obtained yellow precipitate was
collected by filtration, the target complex was isolated by
column chromatography on silica gel (eluent: CH2Cl2 solution
containing 5−10% methanol). After evaporation, the yellow
solid was washed with diethyl ether. Yield: 295 mg (0.52 mmol,
52%). Elemental Anal. Calcd for C15H12N2O5BrRe: C, 31.81;
1
H, 2.14; N, 4.95. Found: C, 31.65; H, 1.97 N, 4.91. H NMR
(400 MHz, acetone-d6) δ/ppm: 9.03 (d, 2H, J = 5.6 Hz), 8.63
(d, 2H, J = 1.2 Hz), 7.75 (dd, 2H, J = 5.6, 1.2 Hz), 4.96 (s, 4H).
FT-IR (acetone) νCO/cm−1: 2020, 1918, 1897.
fac-[Re(4,4′-hydroxymethyl-2,2′-bipyridine)(CO)3(OH2)]-
(OTf)·3H2O. An acetone solution (50 mL) of Re(4,4′-
hydroxymethyl-2,2′-bipyridine)(CO)3Br (295 mg, 0.52 mmol)
and AgOTf (168 mg, 0.65 mmol) was refluxed for 15 h. After
the AgBr precipitate was filtered off, the filtrate was evaporated.
The residual solid was added to water, and the solution was
refluxed for 15 h. After removal of the insoluble solid by
filtration, the solvent was removed by evaporation. Yield: 229
mg (0.35 mmol, 67%). Elemental Anal. Calcd for
C16H20F3N2O12ReS: C, 27.16; H, 2.85; N, 3.96. Found: C,
1
26.84; H, 2.52; N, 3.89. H NMR (400 MHz, D2O) δ/ppm:
9.07 (d, 2H, J = 5.8 Hz), 8.48 (s, 2H), 7.68 (d, 2H, J = 5.8 Hz),
4.90 (s, 4H). ESI-MS (eluent: MeCN) m/z: 528 ([M − H2O +
MeCN − OTf]+).
fac-[Re(2,2′-bipyridine)(CO)3(OH2)](OTf). Similar proce-
dures for fac-[Re(4,4′-hydroxymethyl-2,2′-bipyridine)-
(CO)3(OH2)](OTf) were applied to the synthesis of this
complex. Yield: 64%. Elemental Anal. Calcd for
C14H10F3N2O7ReS: C, 28.33; H, 1.70; N, 4.72. Found: C,
1
28.21; H, 1.81; N, 4.66. H NMR (400 MHz, D2O) δ/ppm:
In the experiments using a two-electrode configuration
without a reference electrode, an IrO2/FTO anode was
employed as a counter electrode instead of the Pt wire. In
this case, both the chambers were purged with CO2. The
gaseous products (i.e., CO, H2, and O2) were measured using
gas chromatography (Agilent Technologies 490 Micro GC or
Inficon MGC3000A); the instrument was directly connected to
the reaction cell. The gaseous products in cathodic and anodic
chambers were alternately measured by the same micro-GC.
Formic acid in the liquid phase was analyzed after electrolysis
by an Otsuka Capi-3300I electronics capillary electrophoresis
system.
9.12 (d, 2H, J = 5.2 Hz), 8.51 (d, 2H, J = 8.0), 8.29 (dd, 2H, J =
8.0, 5.2 Hz), 7.72 (dd, 2H, J = 8.0, 5.2 Hz). ESI-MS (eluent:
MeCN) m/z: 444 ([M − H2O + MeCN − OTf]+).
fac-[Re(4,4′-hydroxymethyl-2,2′-bipyridine)(CO)3(MeCN)]-
(PF6). A dry MeCN solution (10 mL) containing Re(4,4′-
hydroxymethyl-2,2′-bipyridine)(CO)3Br (98 mg, 0.17 mmol)
and AgPF6 (45 mg, 0.18 mmol) was refluxed for 15 h. After the
AgBr precipitate was filtered off, the filtrate was evaporated.
Yield: 99 mg (0.15 mmol, 86%). Elemental Anal. Calcd for
C17H15F6N3O5PRe: C, 30.36; H, 2.25; N, 6.25. Found: C,
30.36; H, 2.49; N, 6.18. 1H NMR (400 MHz, CD3CN) δ/ppm:
8.93 (d, 2H, J = 5.6 Hz), 8.47 (s, 2H), 7.66 (d, 2H, J = 5.6 Hz),
361
DOI: 10.1021/acscatal.7b03275
ACS Catal. 2018, 8, 354−363

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Research Article
(18) Portenkirchner, E.; Schlager, S.; Apaydin, D.; Oppelt, K.;
ASSOCIATED CONTENT
■
Himmelsbach, M.; Egbe, D. M.; Neugebauer, H.; Knor, G.; Yoshida,
̈
S
* Supporting Information
The Supporting Information is available free of charge on the
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FT-IR spectra, 13C NMR spectrum, cyclic voltammo-
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AUTHOR INFORMATION
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Corresponding Author
*E-mail: ishitani@chem.titech.ac.jp.
ORCID
(24) Cheung, K.-C.; Guo, P.; So, M.-H.; Lee, L. Y. S.; Ho, K.-P.;
Wong, W.-L.; Lee, K.-H.; Wong, W.-T.; Zhou, Z.-Y.; Wong, K.-Y. J.
Organomet. Chem. 2009, 694, 2842−2845.
Osamu Ishitani: 0000-0001-9557-7854
(25) Portenkirchner, E.; Gasiorowski, J.; Oppelt, K.; Schlager, S.;
Present Address
Schwarzinger, C.; Neugebauer, H.; Knor, G.; Sariciftci, N. S.
̈
†A.N.: Department of Energy and Hydrocarbon Chemistry,
Graduate School of Engineering, Kyoto University, Nishikyo-
ku, Kyoto 615-8510, Japan.
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Anxolabehere-Mallart, E.; Lau, K.-C.; Lau, T.-C.; Robert, M. J. Am.
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS
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(29) Jarzebinska, A.; Rowinski, P.; Zawisza, I.; Bilewicz, R.; Siegfried,
This work was supported by JST CREST Grant Number
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