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Coordination Chemistry Reviews 257 (2013) 171–186

Contents lists available at SciVerse ScienceDirect

Coordination Chemistry Reviews
journal homepage: www.elsevier.com/locate/ccr

Review

Recent advances in the photocatalytic conversion of carbon dioxide to fuels withwater and/or hydrogen using solar energy and beyond
Yasuo Izumi ∗Department of Chemistry, Graduate School of Science, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japan

Contents
1.
2.
3.

4.

5.

6.

Background . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Thermochemical conversion of CO2 to fuels . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Photon energy conversion of CO2 to fuels with water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.1.
TiO2 photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.2.
Metal-loaded TiO2 photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.3.
Highly dispersed TiO2 photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.4.
Modiﬁed/doped TiO2 photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.5.
Historical low conversion rates and misunderstandings when using TiO2 -based photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.6.
Semiconductor photocatalysts other than TiO2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.7.
Carbon-based photocatalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Photon energy conversion of CO2 to fuels using hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.1.
Photocatalytic conversion of CO2 to methane or CO using hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.2.
Photocatalytic conversion of CO2 to methanol using hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Photon energy conversion of CO2 to fuels using a new reaction system . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5.1.
Recycling of sacriﬁcial electron donors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5.2.
Photoconversion of CO2 to fuels utilizing anode oxidation and cathode reduction compartments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Concluding remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Acknowledgements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

a r t i c l e

i n f oArticle history:
Received 1 January 2012
Received in revised form 5 April 2012
Accepted 13 April 2012
Available online 20 April 2012
Keywords:
Semiconductor photocatalyst
Solar energy
Carbon dioxide
Fuel
Artiﬁcial photosynthesis
Layered-double hydroxide

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a b s t r a c t
Photocatalytic reduction of carbon dioxide to fuels using solar energy is an attractive option for simultaneously capturing this major greenhouse gas and solving the shortage of sustainable energy. Efforts todemonstrate the photocatalytic reduction of CO2 are reviewed herein. Although the photocatalytic results
depended on the reaction conditions, such as the incident/absorbing light intensity from the sun or a simulated solar light source, the performance of different systems is compared. When the reactants included
CO2 and water, it was necessary to determine whether the products were derived from CO2 and not from
impurities that accumulated on/in the catalysts as a result of washing, calcination, or pretreatment in amoist environment. Isotope labeling of 13 CO2 was effective for this evaluation using Fourier-transform
infrared (FTIR) spectroscopy and mass spectrometry (MS). Comparisons are limited to reports in which
the reaction route was veriﬁed spectroscopically, the C source was traced isotopically, or sufﬁcient kinetic
analyses were performed to verify the photocatalytic events. TiO2 photocatalytically produced methane at
the rate of ∼0.1 ␮mol h−1 gcat −1 . In aqueous solutions, formic acid, formaldehyde, and methanol were also
produced. When TiO2 was atomically dispersed in zeolites or ordered mesoporous SiO2 and doped with Pt,
Cu, N, I, CdSe, or PbS, the methane and CO formation rates were greater, reaching 1–10 ␮mol h−1 gcat −1 . As
for semiconductors other than TiO2 , CdS, SiC, InNbO4 , HNb3 O8 , Bi2 WO6 , promoted NaNbO3 , and promoted
Zn2 GeO4 produced methane or methanol at rates of 1–10 ␮mol h−1 gcat −1 , and promoted AII La4 Ti4 O15 produced CO at a rate greater than 10 ␮mol h−1 gcat −1 , in addition to the historically known ZnO and GaP(formaldehyde and methanol formation). The photocatalytic reduction of CO2 was also surveyed withhydrogen, because hydrogen can be obtained from water photosplitting by utilizing natural light. CO wasformed at a rate of ∼1 ␮mol h−1 gcat −1 using TiO2 , ZrO2 , MgO, and Ga2 O3 , whereas both CO and methanol
were formed at a rate of 0.1–1 ␮mol h−1 gcat −1 using layered-double hydroxides consisting of Zn, Cu, Al,∗ Tel.: +81 43 290 3696; fax: +81 43 290 2783.
E-mail address: yizumi@faculty.chiba-u.jp
0010-8545/$ – see front matter © 2012 Elsevier B.V. All rights reserved.
http://dx.doi.org/10.1016/j.ccr.2012.04.018

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

and Ga. When hydrogen is used, in addition to identifying the origin of the carbon, it is critical to conﬁrm
that the products are photocatalytically formed, not thermally produced via CO2 hydrogenation. The
feasibility of the strategy involving the recycling of a sacriﬁcial electron donor and the direct supply of
protons and electrons released from water oxidation catalysts to photocatalysts for the reduction of CO2 tofuels has been demonstrated. However, based on the results obtained to date, it is clear that the practicaluse of the photocatalytic reduction of CO2 as one possible solution for global warming and the world’s
energy problems requires the development of more efﬁcient photocatalysts.
© 2012 Elsevier B.V. All rights reserved.

1. BackgroundOwing to the limited amount of energy sources and the recent
effects of fossil fuel use on the global environment, the paradigm ofenergy supply is changing from one based on the use of carbon-richrocks, peat, and liquid found in the Earth to one based on renewablesources, such as energy crops, sunlight, and wind [1]. According to the Intergovernmental Panel on Climate Change (IPCC), oil(34.6%), coal (28.4%), gas (22.1%), and nuclear energy (2.0%) weremajor worldwide energy sources in 2008 (Fig. 1). The percentageof renewable energy was 12.9%; however, if the traditional simple
burning of biomass (6.3%) is excluded, renewable energy accounted
for only 6.6% of total consumption.
Carbon dioxide is one of the major greenhouse gases andis formed as a result of the consumption of fossil fuels [2].
In the CO2 emission scenarios described for 2050, bioenergy(1.6 × 1020 J y−1 ), direct solar energy (3 × 1019 J y−1 ), and windenergy (2.5 × 1019 J y−1 ) are the top three renewable technologies
that must be adopted in order to realize the ambitious target that
calls for the decrease in the atmospheric CO2 concentration toless than 440 ppm. Thus, renewable energy must be investigatedintensively, because modern biomass, direct solar energy, and windenergy supply only 1.9 × 1019 , 0.5 × 1018 , and 1 × 1018 J y−1 , respectively [1], and the application of renewable energy is increasing
very slowly.
Several methods for reducing the CO2 concentration in theatmosphere and preventing CO2 emissions due to human activity
have been investigated, such as investigating the sorption of CO2
into new/functionalized materials; increasing the quantity of green
carbon sinks (plants, phytoplankton, and algae containing chloroplasts); increasing the level of dissolved carbonate and its salts in
sea water; or capturing CO2 and transferring it to the bottom of thesea in a supercritical state [3].
It would be advantageous to capture CO2 from the atmosphereor the exhaust of factories/power stations and convert it to fuel byusing a sustainable source of energy such as sunlight. This option
solves the problems of global warming and the sustainable energy
shortage simultaneously [3–6]. It is an enormously difﬁcult task

to combine water splitting and carbon dioxide reduction [3,7,8].
Water oxidation and the subsequent reduction of CO2 are required.
This review mainly discusses the photocatalytic conversion of CO2
to fuels (Sections 3–5) using semiconductors, but also presents acomparison of the related thermochemical conversion of CO2 tofuels (Section 2) via the reduction–oxidation of metal oxides.
2. Thermochemical conversion of CO2 to fuels
The energy from the sun that reaches the Earth in 1 h is 9200
times (4.3 × 1020 J h−1 ) the energy consumed on the Earth in 1 h in2001 (4.7 × 1016 J h−1 ) [9]. In other words, all the energy consumedon the Earth in one year can be supplied from solar energy in only
1 h. To utilize the enormous energy provided by the sun, two-stepthermochemical cycles to dissociate CO2 and H2 O using metal oxide
redox reactions have been proposed [10]. Nonstoichiometric oxides
such as cerium oxide are partially reduced at higher temperatures(1873 K for cerium oxide), releasing O2 under concentrated solar
radiation, and then are oxidized again by reacting with CO2 andH2 O at lower temperatures:
2MO2 → 2MO2−ı + ıO2 (g)ıH2 O + MO2−ı → ıH2 (g) + MO2
ıCO2 + MO2−ı → ıCO(g) + MO2 ,where M is Ce, Zn, or Fe, and the above stoichiometry represents
an example for the case when M is Ce.
Compared to the redox systems consisting of ZnII O Zn0 and
FeII FeIII 2 O4 –FeII O, cerium oxide is attracting attention because
CeO2 is partially reduced via the formation of an oxygen vacancy
without signiﬁcant reorganization of the crystal lattice, andtherefore, it reversibly stores and releases lattice oxygen atoms.
Furthermore, CeO2 has a high melting point (2220 K), thermal stability, and is less susceptible to crystal-reordering phase transitions[11]. Above 1173 K under a solar ﬂux with a density of 150 W cm−2 ,the average evolution rate of O2 was 0.049 mL min−1 gcat −1 fromcerium oxide, and the obtained partially reduced CeO2−ı transformed CO2 to CO at 1173 K at an average rate of 1.8 mL min−1 gcat −1
and transformed H2 O to H2 at 1173 K at an average rate of
0.95 mL min−1 gcat −1 [10].
These thermochemical conversion rates are higher than thephotocatalytic rates discussed below, but focusing lenses for sunlight and high-temperature reactors incur high initial investment
costs.
3. Photon energy conversion of CO2 to fuels with water
3.1. TiO2 photocatalysts

Fig. 1. 2008 distribution of world energy consumption by source [1].

The photon energy of sunlight can be converted to electricenergy using solar cells [12] and to chemical energy using photocatalysts. The development of photocatalysts to convert solar energy
to chemical energy is an indispensable option for storing energy
and for mobile use, especially when using sustainable, cheaper

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

173

Table 1A
Reported CO2 photoreduction catalysts, reaction conditions, and the formation rates in water/with moisture using TiO2 .
Photocatalyst

Reactants

T (K)Brand nameAmount (g)

CO2

H2 O

TiO2

1

Saturated

Liq (100 mL)

TiO2
TiO2
TiO2 (P25)
Anatase TiO2
TiO2 (JRC-TIO-4)
Anatase TiO2
TiO2 (P25)
TiO2 pellet
Ti-PVG
Ti-PVG

0.3
0.15
0.15
0.05
–
–
–
–
–
–

Saturated
Saturated
Saturated
9.0 MPa (sc)
111 kPa (gas total)
Gas (0.04–0.15 mmol)
Gas (17%)111 kPa (gas total)
111 kPa (gas total)
Gas (0.04–0.15 mmol)

Liq (120 mL)
100 mL, 0.1N NaOH
Liq (1.5 mL)
Liq (5 mL, added later)
Saturated gas
Gas (0.04–0.25 mmol)
Gas (83%)
Saturated gas
Saturated gas
Gas (0.04–0.25 mmol)

a

333
278
308
323

323

Light source

Major product {formation rate(␮mol h−1 gcat −1 )}Reference

500 W Xe/HP Hg

Formaldehyde (16)Methanol (3.3)
Formaldehyde (2.0a )Formic acid (22a )CO (0.35)
Formic acid (1.8)
Methane (0.014)
Methane (0.17)
Methane (15a )Methane (0.0021)
Methane (0.0025)
Methane (0.020)[14]500 W HP Xe
500 W W-halogen
500 W HP Hg arc (>310 nm)
990 W Xe (>340 nm)
UV (254 nm)
75 W HP Hg (>280 nm)
UV (365 nm)
UV (254 nm)
UV (254 nm)
75 W HP Hg (>280 nm)

[19]
[40]
[39]
[18]
[15]
[16]
[45]
[15]
[15]
[16]

Spectroscopic veriﬁcation of reaction route, isotope tracing to identify the C source, or sufﬁcient control kinetic tests to verify the photocatalytic event is required.

semiconductors [3,5,9,13]. Photocatalytic conversions of CO2 tofuels with water using semiconductors are summarized in Table 1.
In a pioneering paper in 1979, the photoreduction of CO2 toformaldehyde and methanol in puriﬁed water using the semiconductors TiO2 , ZnO, CdS, GaP, SiC, and WO3 was reported [14]. Basedon the correlation between the conduction band energy potential
and the yield of methanol, it was suggested that the photoreductionof CO2 proceeded by the photoexcited electrons in the conduction
band moved to CO2 . The conduction band energy minimum washigher than that for CO2 photoreductions.
CO3 2− + 6H+ + 4e− → H2 CO(aq) + 2H2 O, E ◦ (298 K) = +0.197 V
CO3 2− + 8H+ + 6e− → CH3 OH(aq) + 2H2 O, E ◦ (298 K) = + 0.209 V,
in which the electric potentials are referenced to a standard hydrogen electrode (SHE). The photocatalytic results in this review are
explained on the basis of this band gap electron excitation mechanism.
The photoreduction of CO2 using TiO2 has been most extensively
investigated (Table 1A). The CO2 photoreduction in watersaturated CO2 gas was compared at 111 kPa using TiO2 powder(reference catalyst JRC-TIO-4, Catalysis Society of Japan), anchoredTi oxide on porous Vycor glass (PVG), and 4-mm pelletizedTiO2 under UV (ultraviolet) light for 48 h. The methane formation rates were 14 nmol h−1 gcat −1 , 2.5 nmol h−1 gcat −1 , and
2.1 nmol h−1 gcat −1 , respectively [15], suggesting that the speciﬁc
surface area of the catalyst was of prime importance. CO and H2
were minor products.
The photocatalytic performance was also compared on a single crystal surface of rutile-type TiO2 (1 0 0) (Ti terminating) andTiO2 (1 1 0) (O terminating) (Fig. 2). Careful kinetic measurements
were performed, because the surface area of a single crystal surface
is limited in comparison to that of porous powders. The TiO2 (1 0 0)

Fig. 2. Comparison of the (1 0 0) and (1 1 0) crystal faces in the crystal structure ofrutile TiO2 and their performance in photocatalytic CO2 reduction [16].

surface was more photocatalytically active than the TiO2 (1 1 0) surface by a factor of 7.4, suggesting that a reductive Ti-terminating
surface is favorable for reducing CO2 [16].
The reactant CO2 was condensed to a supercritical (sc) condition (Tcritical = 304.1 K, Pcritical = 7.375 MPa) [17] to boost thereaction frequency. When anatase-type TiO2 powder was set insupercritical CO2 (308 K, 9.0 MPa) and irradiated under UV–visible
light (>340 nm) followed by the addition of water, formic acid wasformed at a rate of 1.8 ␮mol h−1 gcat −1 [18]. The two-step synthesisof formic acid starting from CO2 was accelerated by the additionof phosphoric, nitric, or hydrochloric acid in the second step. In
particular, an H3 PO4 solution worked best for protonation to form
formic acid at a pH of 2.2 and a rate of 2.9 ␮mol h−1 gcat −1 .
This result reported in Ref. [18] is one of the best rates in Table 1A
for conditions where the photocatalytic reaction route was supported by spectroscopy. The reaction mechanism was proposed
to proceed through a one-electron reduced radical intermediate
species (·CO2 − ) whose presence was conﬁrmed by electron spinresonance (ESR) measurements:hv

TiO2 −→e− (as Ti3 + ) + h+hv

Ti4+ –O2− −→Ti3+ –O−
CO2 + e− → • CO2 −• CO − + 2H+ (of water or acid) + e− → HCO H
2
2• CO − + • CO − + 2H+ (of water or acid) → HCO H + CO2
2
2
2

3.2. Metal-loaded TiO2 photocatalysts
Various studies on the use of an added metal to enhance thephotocatalytic activity of TiO2 have also been reported. An aqueous
suspension of TiO2 , Rh/TiO2 , or Rh/WO3 –TiO2 at 333 K was bubbled
with a CO2 ﬂow and illuminated by a xenon lamp, and the exit gas
was trapped at 273 K and analyzed [19]. The major product whenusing the TiO2 suspension was formaldehyde, which was formed
at a rate of 2.0 ␮mol h−1 gcat −1 (Table 1A). A minor amount offormic acid was also detected. However, the major product detected
when using the air-calcined Rh/WO3 –TiO2 suspension was formic
acid, which formed at a rate of 1.6 ␮mol h−1 gcat −1 (Table 1B).
Minor amounts of methanol and formaldehyde were also detected.
Furthermore, the product distribution clearly changed again
when H2 -reduced (473 K) Rh/WO3 –TiO2 was used. In this case,methanol was synthesized exclusively at a rate greater than
4.0 ␮mol h−1 gcat −1 .

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Table 1B
Reported CO2 photoreduction catalysts, reaction conditions, and the formation rates in water/with moisture using metal-loaded, highly-dispersed, and modiﬁed TiO2 .
Photocatalyst

Reactants

T (K)Light source

Major product {formation
rate (␮mol h−1 gcat −1 )}Reference

UV (365 nm)

Ethanol (30a ) formic acid(19a ), methane (12a )Methane (0.046), methanol(0.031)
Methane (0.080)
Methane (7.5), methanol(3.1)
Methane (12)[45]

Brand nameAmount (g)

CO2

H2 O

TiO2 –MWCNT–

Gas (17%)

Gas (83%)

Ti-Y-zeolite

0.15

Gas (24 ␮mol)

Gas (120 ␮mol)328

HP Hg (>280 nm)

Pt Ti-Y-zeolite
Ti-MCM-48 (Si/Ti = 80)

0.15–

Gas (24 ␮mol)
Gas (24 ␮mol)

Gas (120 ␮mol)
Gas (120 ␮mol)

328
328

HP Hg (>280 nm)
HP Hg (>280 nm)

Pt Ti-MCM-48 (1.0 wt% Pt,
Si/Ti = 80)
Ti-SBA-15 (0.29 wt% Ti,
Si/Ti = 270)
Ti-SBA-15 (0.05 wt% Ti)–

Gas (24 ␮mol)

Gas (120 ␮mol)328

HP Hg (>280 nm)

0.05

Gas (36 ␮mol)

Gas (180 ␮mol)

323

0.05

Gas (38 ␮mol)

Gas (76 ␮mol)313

100 W HP Hg(>250 nm)
120 W HP Hg

CdSe-Pt/TiO2
PbS-Cu/TiO2
Rh(1 wt%)/WO3 (2 wt%)-TiO2
Rh/WO3 –TiO2 -reduced
CoII -Pc(0.7 wt%)/TiO2
ZnII -Pc(1 wt%)/TiO2
NdIII (0.2 wt%)/TiO2
Pd(1 wt%)-TiO2

0.3
0.3
0.3
0.15
0.15
–
0.15

Gas (40 Pa)
Gas
Saturated
Saturated
Saturated
Saturated (2.3 g)
Saturated
Saturated

Gas (400 Pa)
Saturated gas
Liq (120 mL)
Liq (120 mL)
100 mL, 0.1N NaOH
100 mL, 0.1N NaOH
Liq
Liq (1.5 mL)

Cu-TiO2
Cu-N-TiO2 NT
Cu-N-TiO2 NT
Pt N-TiO2 NT
No catalyst
I-TiO2

0.15–0.6
–
–
–
0
0.2

Saturated
108 kPa (gas total)
108 kPa (gas total)
108 kPa (gas total)
108 kPa (gas total)
99 kPa300 mL, 0.2 N NaOH
Saturated
Saturated
Saturated
Saturated
2.3 kPa

a

333
333

278
323
317
317
317
317

300 W Xe (>420 nm)
300 W Xe (UV cut)
500 W HP Xe
500 W HP Xe
500 W W-halogen
500 W W-halogen
UV
500 W HP Hg arc(>310 nm)Hg (254 nm)
Sun
No light
Sun
Sun
Xe

[28]
[28]
[29]
[29]

Methane (0.31), methanol(0.081)Ethane (0.020), methane(0.016), ethene (0.007)
Methane, methanol
CO (0.82), methane (0.58)
Formic acid (1.6a )Methanol (4.0a )Formic acid (150a )Formic acid (98a )Methanol (23a )Methane (0.37)

[26]
[27]
[19]
[19]
[40]
[41]
[44]
[39]

Methanol (20a )Methane (4.4)
Methane (0.13a )Methane (2.9)
Methane (0.10a )CO (2.4a )[42,43]
[13]
[13]
[13]
[13]
[37]

[30]
[31]

Spectroscopic veriﬁcation of reaction route, isotope tracing to identify the C source, or sufﬁcient control kinetic tests to verify the photocatalytic event is required.

The following reaction mechanism, which again includes a oneelectron reduced species (• CO2 − ), was proposed [19] and is basedon a probable photo-assisted thermal methanol synthesis at 333 K,
starting from CO2 and H2 O.
CO2 + e− → • CO2 −• CO − + H O → HC(O)O• + OH−22

oxidation of the CdSe quantum dots, the photocatalyst was deactivated. PbS quantum dots (4 nm) were also effective for boosting
the photocatalytic conversion of CO2 to CO and methane (Table 1B)[27].
3.3. Highly dispersed TiO2 photocatalysts
The photocatalytic enhancement of TiO2 was also attemptedusing highly dispersed active Ti ion species. The performance of

HC(O)O• + e− → HCO2 −HCO2 − + H+ → HCO2 H
It is also plausible that water photosplitting ﬁrst occurred photocatalytically, and then the thermal reaction of CO2 with H2(or a H species present on the Rh surface) proceeded, because
Rh/WO3 –TiO2 is a typical thermal catalyst for CO/CO2 hydrogenation [20–24].
The addition of Cu+ to TiO2 powder photocatalysts was reported
to lead to the formation of methanol [16]. The effects of doped-d10 conﬁguration Cu+ , Ag+ , and Pb2+ ions on d0 -conﬁguration metals,
such as Ti, Nb, and Ta oxides, have been discussed for photocatalytic
water splitting reactions, and hopefully can be applied also to thecarbon dioxide photoreduction [25].
CdSe quantum dots (2.5 and 6.0 nm) were mixed with a Pt/TiO2
photocatalyst, and the photocatalyst was sensitized for the photoreduction of CO2 (40 Pa) with moisture (400 Pa) under visible
light exposure [26]. The photoexcitation of the electrons between
the band gap of TiO2 by visible light (>420 nm) is hardly possible, but the photoexcitation of electrons between the band gapof the CdSe quantum dots is possible using visible light, and theexcited electrons would be injected to the conduction band of TiO2
(Fig. 3). With this system, the conversion of CO2 to methane (major)
and methanol (minor) was observed for 4–6 h, but because of theFig. 3. Energy diagram for bulk and nano CdSe with TiO2 , and the redox potentials
for the CO2 reduction and water oxidation reaction steps.
Reprinted with permission from Wang et al. [26]. Copyright (2010) American Chemical Society.

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Ti species anchored on PVG prepared from TiCl4 and highly dispersed TiO2 powder in the photocatalytic reduction of CO2 wascompared [16]. The Ti sites of Ti–PVG were surface-isolated tetrahedral TiO4 based on the intense Ti 1s–3d electronic transition peak
and a weaker signal ascribed to Ti O Ti (or Ti O Si) bonding in
the Ti K-edge extended X-ray absorption ﬁne structure (EXAFS)
compared to highly dispersed anatase-type TiO2 .
Under CO2 gas and water vapor, the photocatalysts were irradiated with UV light through water and a cut-off ﬁlter ( > 280 nm)
at 323 K. The anatase-type TiO2 exclusively produced methanewith a formation rate of 0.17 ␮mol h−1 gcat −1 , in contrast to the
Ti–PVG, which led to the formation of methane, methanol, and
CO. The methane formation rates for the Ti–PVG photocatalyst
increased from 1.6 nmol h−1 gcat −1 to 5.0 nmol h−1 gcat −1 when themolar ratio of H2 O and CO2 increased from 0 to 3 and then toincreased to 20 nmol h−1 gcat −1 when the molar ratio of H2 O and
CO2 increased to 5.
In addition, the total yield increased when this UV irradiation
test was conducted at 323 K compared to that conducted at 275 K.
Based on the temperature dependence, the thermal assist mechanism may be applicable in Ref. [16], but the charge separation
between Ti and O in the Ti3+ O− bond was monitored by ESR spectroscopy, and the quenching of photoluminescence in the presenceof CO2 was also shown. Essentially, the charge separation between
Ti and O in the Ti3+ O− bond played both reductive and oxidative
photocatalytic roles [16].
The photocatalytic reduction of CO2 (24 ␮mol, 0.73 kPa) byatomically dispersed Ti sites in Y-zeolite and its derivatives
in the presence of water (120 ␮mol, 3.7 kPa) using UV lightfrom a high-pressure Hg lamp (>280 nm) at 328 K [28,29] wasalso reported. Atomically dispersed Ti sites in Y-zeolite prepared via an ion-exchange method produced both methane(46 nmol h−1 gcat −1 ) and methanol (31 nmol h−1 gcat −1 ), whereas a
Pt-loaded ion-exchanged Ti-Y-zeolite selectively formed methane(80 nmol h−1 gcat −1 ). Impregnated Ti oxide species in Y-zeolite
exclusively formed methane.
The selectivity difference was explained as follows. The Ti sites
in the ion-exchanged Y-zeolite are atomically dispersed on thebasis of X-ray absorption near-edge structure (XANES) and EXAFS
analyses. The charge separation between Ti and O in the Ti3+ O−
bond was evaluated on the basis of the photoluminescence spectra. The lifetime of the excited Ti3+ O− was determined to be
54 ␮s, which is substantially higher than that for TiO2 powders(nanosecond order). Owing to quantum size effects, the dispersed
Ti3+ sites were more negative, and accordingly exhibited a greater
reducing potential for the CO2 substrate, enabling methanol formation based on the band-gap values given from the UV–visible
spectra. The E◦ values were −0.32 and −0.244 V for methanol
and methane formation, respectively (Fig. 3). CO seemed to be anintermediate for the reduction of CO2 , but H2 was not an intermediate. The water could have been oxidized to OH and H+ , and
the Pt sites probably worked to suppress charge recombination[28].
Ti-MCM-48
producedbothmethaneSimilarly,(7.5 ␮mol h−1 gcat −1 )andmethanol(3.1 ␮mol h−1 gcat −1 ),while Pt Ti-MCM-48 was the most active and selectively
formed methane (12 ␮mol h−1 gcat −1 ) [29]. Ti-SBA-15 also
produced methane and methanol with formation rates of
0.31 and 0.081 ␮mol h−1 gcat −1 , respectively (Table 1B) [30].
In contrast, when the Ti-SBA-15 was illuminated in humid
helium to remove any C residues, the rates of photoreductionof CO2 to ethane, methane, and ethylene were quite small(0.020–0.007 ␮mol h−1 gcat −1 ) [31]. Minor carbonaceous impurities in the Ti-MCM-41 were detected by Fourier-transform infrared(FTIR) spectroscopy using isotopes [32]. Based on these results, it
is possible that the C residues remaining from the original organic

175

Fig. 4. (a) Photograph of the reaction chambers set in the sunlight during summer
in Pennsylvania, USA for photocatalytic CO2 conversion under CO2 saturated withmoisture (total pressure <108 kPa). (b) Spectral irradiation for the photocatalytic
tests.
Reprinted with permission from Varghese et al. [13]. Copyright (2009) American
Chemical Society.

template compounds used in the syntheses of Y-zeolite, MCM-48,
and SBA-15 were the (partial) source of the products reported inRefs. [28–30], but the very low Ti loading level in Ref. [31] may berelated to the lower photocatalytic activity.
Exclusive formation of CO and O2 was reported for Ti-MCM41 under CO2 and H2 O [32]. The light intensity from a 266-nm
pulsed Nd:Yag laser was varied between 80 and 160 mW cm−2 , anda proportional increase in CO production was observed, which indicated a single photon process accompanying the charge separation
between Ti and O in the Ti3+ O− bond. The difference in the major
product (CO) compared to that formed in Refs. [29–31] (methane)is due to the different light source: a UV-pulsed laser versus a Hg
lamp.
3.4. Modiﬁed/doped TiO2 photocatalysts
Nitrogen-doped TiO2 nanotube arrays were synthesized by theanodization of Ti metal in an ethylene glycol aqueous solution(2 vol%) with 0.3 M ammonium ﬂuoride as a the N source. The
Cu- or Pt-loaded N-doped TiO2 nanotube photocatalysts were
set outside on a sunny summer day in Pennsylvania, USA underwater-vapor-saturated CO2 (total pressure <108 kPa) for 2.5–3.5 h
at 317 K (Fig. 4) [13]. The major product using the Cu- and Pt-loaded

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176

Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

N-doped TiO2 nanotubes was methane, which was produced at a
rate of 4.4 ␮mol h−1 gcat −1 and 2.9 ␮mol h−1 gcat −1 , respectively, in
addition to minor amounts of CO, ethane, propane, butane, pentane, and hexane. The major advantages of using N-doped TiO2
nanotubes for the CO2 photoreduction were considered to be (1)a minimum thickness of the N-doped TiO2 nanotubes of ∼10 nm
versus a hole diffusion length of ∼10 nm [33] and an electron diffusion length of ∼10 ␮m in TiO2 ; [34] and (2) the electron trapat the Cu or Pt to donate electrons to CO2 . A later paper on theCO2 photoreduction using Pt-TiO2 nanotubes reported the exclusive formation of methane in addition to trace amounts of ethane[35]. Based on a comparison of the two studies, it is possible to conclude that (at least part of) the hydrocarbons produced in Ref. [13]originated from the tape used in the reactor.
When the fraction of sunlight with wavelengths shorter than
400 nm was removed, the photocatalytic activity of CO2 reduction was reduced to 3% of that under the entire solar spectrum.
The control reaction in the dark (using the Cu-loaded N-doped
TiO2 nanotubes) and in the absence of catalyst produced methaneat rates of 0.13 ␮mol h−1 gcat −1 and 0.10 ␮mol h−1 gcat −1 , respectively. The reaction mechanism operating in response to the UV
light was proposed as follows [13].
H2 O + h+ → • OH + H+
H+ + e− → • H•H + •H → H

2

2CO2 + 4e− → 2CO + O2
CO + 6e− + 6H+ → CH4 + H2 O
Iodine-doped TiO2 is known as a visible-light-responsive photocatalyst, typically for photooxidation [36]. Essentially, doped I−anions form impurity energy levels above the valence band of TiO2
and extend the availability of the visible light. This strategy wasapplied to the CO2 photoreduction in the gas phase. The CO formation rate (2.4 ␮mol h−1 gcat −1 ; Table 1B) [37] is comparable toveriﬁed methane formation rates using other dispersed or doped
TiO2 photocatalysts. Linear O2 formation was also monitored, but
to further discuss the following proposed mechanism compared tocompetitive H2 formation, quantitative kinetic data are needed.
TiO2 → h+ + e−

This signiﬁcant up-hill free-energy change is due to the greater
enthalpies for CO2 and H2 O. The free-energy change for water splitting is lesser by 34%.
2H2 O(g) → 2H2 (g) + O2 (g),Therefore, in view of the free-energy change, the CO2 conversion
is extremely unfavorable, and the reduction of impurity carbon tomethanol could be misinterpreted as the reaction of CO2 [38,39].
In one example, TiO2 was synthesized by a sol–gel route startingfrom titanium(IV) n-butoxide and polyethylene glycol (PEG) [38] or
n-butanol plus acetic acid [40–43]. Methanol (20 ␮mol h−1 gcat −1 )[42,43] and formic acid formation (150–98 ␮mol h−1 gcat −1 ) [40,41]
were reported using a suspended, Cu-loaded, sol–gel-derived TiO2
in NaOH solution saturated with CO2 under UV light from a
mercury lamp [42,43] or using CoII -phthalocyanine(Pc)/TiO2 andZnII -Pc/TiO2 under a W halogen lamp [40,41]. In an aqueous
solution saturated with CO2 using NdIII /TiO2 under UV irradiation,
the formation rate of methanol was as much as 23 ␮mol h−1 gcat −1[44]. The TiO2 powder used in this study was also prepared via asol–gel route.
TiO2 has also been synthesized via a sol–gel route startingfrom Ti(IV) tetrachloride and hydrochloric acid, followed by theaddition of ammonia until the pH reached 7. Nanocomposites ofTiO2 and multi-walled carbon nanotubes (MWCNTs) produced
ethanol, formic acid, and methane with formation rates of 30,
19, and 12 ␮mol h−1 gcat −1 , respectively, (Table 1B) in H2 O andCO2 (molar ratio 5:1) under a UV lamp (365 nm) for 5 h [45]. In
this study, TiO2 (P25, Degussa) and TiO2 synthesized by a sol–gel
method were also tested for the CO2 photoreduction. Unmodiﬁed
TiO2 (P25) produced ethanol, formic acid, and methane withformation rates of 1.0, 19, and 15 ␮mol h−1 gcat −1 , respectively,
and the sol–gel-derived TiO2 was even more active. These results
were in contradiction to those reported in Refs. [15,16] in whichanatase-type (or anatase-type major) TiO2 produced only methaneat rates of 0.014–0.17 ␮mol h−1 gcat −1 under similar reaction
conditions (Table 1A).
To investigate the surface reaction mechanism starting from CO2
to fuels, the isotope distribution of adsorbed CO (12 CO at 2115 cm−1
and 13 CO at 2069 cm−1 ) over CuI /TiO2 was monitored using diffuse
reﬂectance infrared Fourier-transform spectroscopy in 13 CO2 gas(Fig. 5) [38]. Adsorbed 12 CO species were the primary products,
indicating that a reverse disproportionation reaction of impurity
carbon with CO2 proceeded to form CO.
12

2H2 O + 4 h+ → 4H+ + O2

Gr ◦ = +457.2 kJ mol−1C(impurityin/oncatalyst) + 13 CO2 → 12 CO + 13 CO

CO2 + 2H+ + 2e− → CO + H2 O

If the carbon residues deposited/buried on the catalyst surface
during the catalyst synthesis were involved in reactions with photocatalytically activated surface-adsorbed water via

To summarize Sections 3.2–3.4, when TiO2 was atomically dispersed on zeolites or ordered mesoporous SiO2 or doped with Pt,
Cu, N, I, CdSe, or PbS, photocatalytic methane or CO formation rates
increased to 1–10 ␮mol h−1 gcat −1 .

CO + 2H2 O → CH3 OH + O2 ,3.5. Historical low conversion rates and misunderstandings whenusing TiO2 -based photocatalysts
The reaction routes for the CO2 reduction with water should
be carefully checked to determine whether the carbon source wasactually the CO2 reactant. The major obstacle for the reduction of
CO2 is the thermodynamic limitation [3]:CO2 (g) + 2H2 O(g) → CH3 OH(g) + 1.5O2 (g),Gr ◦ = +689 kJ mol−1Gr ◦ = +432 kJ mol−1

in which the free-energy change is lesser by 37% compared to the
reduction of CO2 to methanol, the thermodynamic limitation canbe partially compensated for by the impurity carbon.
Recently, based on detailed gas chromatography (GC) and
GC–MS (Mass spectrometry) analyses, the C source was carefully
investigated with respect to the production of CO, methane, ethane,acetic acid, formic acid, and methanol from CO2 [39]. The tests wereconducted in CO2 -saturated water using TiO2 (P25) at 278 K underUV–visible light (>310 nm). Using as-received TiO2 , methane wasformed at a rate 150 nmol h−1 gcat −1 , but the molar amount corresponded well to that of acetic acid desorbed by calcination at 623 K.
The reaction was not considered to be CO2 photoreduction, but thephoto-Kolbe reaction [46]:Author's personal copy
Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

177

Fig. 5. Isotope distribution of the adsorbed CO (12 CO at 2115 cm−1 and 13 CO at 2069 cm−1 ) over CuI /TiO2 monitored by diffuse reﬂectance infrared Fourier-transform
spectroscopy in 13 CO.
Reprinted with permission from Yang et al. [38]. Copyright (2010) American Chemical Society.

CH3 CO2 H + h+ → • CH3 + CO2 + H+
• CH

3 + CH3 CO2 H → CH4 +• CH CO H
22

After calcination at 623 K and further thorough washingof the TiO2 with deionized water, the reaction was repeated,
and the organics were not detected; the actual CO andmethane formation rates under UV–visible light were 350
and 25 nmol h−1 gcat −1 , respectively. Similarly, photochemicallydeposited Pd–TiO2 was washed several times and set in CO2 saturated water under UV–visible light. Methane was the major
product (370 nmol h−1 gcat −1 ), and ethane and CO were formedin minor amounts. The methane source was tested using 13 CO2 ,and the GC–MS peak at m/e = 17 conﬁrmed the formation of 13 CH4directly from 13 CO2 [39].
In summary, based on the studies presented in Refs. [38,39], itis essential to verify the C source for the products reported in Refs.
[40–45]. The products were very likely derived from an alkyl groupof the catalyst precursor compounds or the organic solvent.
3.6. Semiconductor photocatalysts other than TiO2
In the pioneering report in 1979, the photoreduction of CO2
to formaldehyde and methanol in water was reported using TiO2 ,ZnO, CdS, GaP, SiC, and WO3 (Table 1C). The rates of formaldehyde
production were in the order:
CdS > ZnO ∼ TiO2 > GaP ∼ SiC
However, the rates of methanol formation followed quite a different order [14]:SiC  CdS ∼ GaP > ZnO > TiO2
CdS nanoparticles were supported on montmorillonite, and thephotoreduction of CO2 was tested in an alkaline solution under
UV light (254 nm) [47]. Although water photosplitting proceeded
predominantly (H2 formation rate: 8.7 ␮mol h−1 gcat −1 ), the secondary hydrogenation of CO2 was suggested with a formation
rate of 0.93 ␮mol h−1 gcat −1 . p-SiC powder and copper powder
were dispersed in a potassium hydrogen carbonate solution saturated with CO2 at a pH of 5 at 313 K. When the suspension was
illuminated with UV light from a mercury lamp (7 mW cm−2 at
365 nm), methane was the major product, which had a formation rate of 0.63 ␮mol h−1 gcat −1 , and ethylene and ethane wereminor products [48]. The Cu site was suggested to be an acceptor ofphotogenerated electrons, but the Cu-p-SiC photocatalyst deactivated in 1–2 h.
InTaO4 that was previously synthesized by solid-state synthesis starting from In2 O3 and Ta2 O5 was impregnated with 0–1.0 wt%
NiO using an aqueous nickel nitrate solution [49], calcined at 623 K,
reduced at 773 K, oxidized at 473 K, and then used as a catalyst
for the photcatalytic reduction of CO2 . After the photocatalytic test
in a CO2 -saturated potassium bicarbonate solution was complete,
the aqueous phase was analyzed, and methanol was formed at a
maximum rate of 1.4 ␮mol h−1 gcat −1 using NiO/InTaO4 . Unfortunately, control reactions (in the dark, in the absence of catalyst,
in the absence of CO2 and/or KHCO3 ) and other product information, including analysis of the gas phase were not reported. The
NiO/InTaO4 photocatalysts were also tested for the gas-phase CO2
reduction using a monolith reactor [50]. The major product was
acetaldehyde, which was formed at a rate of 0.21 ␮mol h−1 gcat −1(Table 1C), rather than minor methanol formation.
InNbO4 was synthesized by solid-state reactions and then tested
for the CO2 photoreduction in a KHCO3 solution [51]. Based on
the control tests and the fact that the catalyst was prepared via
a carbon-free synthetic method, the methanol formation, which
occurred at a rate of 1.3 ␮mol h−1 gcat −1 under UV–visible light,
should be photocatalytic. Hydrothermally synthesized HNb3 O8nanobelts and Bi2 WO6 nanoplates also underwent the photoreduction of CO2 to methane with rates of 3.6 and 1.1 ␮mol h−1 gcat −1 ,respectively, (Table 1C) [52,53] which are comparable to TiO2 based photocatalysts. The [0 0 1] face of the nanoplates was
suggested to be responsible. Pt-promoted perovskite NaNbO3 wasreported to produce methane from CO2 + H2 O at a rate comparable to the levels of the other best semiconductor photocatalysts:
4.9 ␮mol h−1 gcat −1 [54].
BiVO4 was synthesized by a hydrothermal method starting from
acidic bismuth nitrate and alkaline ammonium vanadate solutions[55]. Cetyltrimethyl ammonium (CTMA) bromide (to obtain monoclinic BiVO4 crystals) or PEG (to obtain tetragonal BiVO4 crystals)
was added to both solutions, and the two solutions were mixed.
The mixed solution was heated at 473 K in an autoclave under
microwave irradiation. In the cooled water saturated with CO2 gas,
the monoclinic BiVO4 was reported to produce ethanol with rates
of 2000 and 110 ␮mol h−1 gcat −1 under UV–visible and visible light,
respectively.
A control reaction was performed in nitrogen gas instead of CO2 ,and no ethanol production was detected. However, this fact did notAuthor's personal copy
178

Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Table 1C
Reported CO2 photoreduction catalysts, reaction conditions, and the formation rates in water/with moisture using semiconductor photoatalysts other than TiO2 .
Photocatalyst

Reactants

T (K)Light source

Major product{formation rate(␮mol h−1 gcat −1 )}Reference

500 W Xe/HP
Hg
500 W Xe/HP
Hg
500 W Xe/HP
Hg
500 W Xe/HP
Hg
8 W HgHg (>275 nm)
500 W halogen
300 W Xe
300 W Xe arc(>420 nm)
300 W Xe arc
300 W Xe arc(>400 nm)
500 W halogen
350 W Xe
300 W Xe arc
Light
Light
Light
400 W HP Hg
400 W HP Hg
400 W HP Hg

Formaldehyde (17),methanol (5.0)
Formaldehyde (29),methanol (17)Formaldehyde (14),methanol (16)Formaldehyde (14),methanol (76)Methane (0.93)
Methane (0.63)
Methanol (1.4a )Acetaldehyde (0.21a )Methane (1.1)

[14]

[47]
[48]
[49]
[50]
[53]

Ethanol (2000a )Ethanol (110a )

[55]
[55]

Methanol (1.3)
Methane (3.6)
Methane (4.9)
Methane (0.41)
Methane (1.9)
Methane (6.5)
CO (2)CO (14)
CO (7.7), formic acid
(4.3)
CO (6.0), formic acid
(1.7)
CO (73), formic acid
(2.3)
CO (31)
CO (24), formic acid
(2.7)
CO (42)

[51]
[52]
[54]
[57]
[57]
[57]
[60]
[60]
[60]

Brand nameAmount (g)

CO2

H2 O

ZnO

1

Saturated

Liq (100 mL)

CdS

1Saturated

Liq (100 mL)

GaP

1Saturated

Liq (100 mL)

SiC

1Saturated

Liq (100 mL)

CdS-montmorillonite
Cu-p-SiC
NiO(1.0 wt%)/InTaO4
NiO(2.6 wt%)/InTaO4
BiWO6

1 g/L
0.10
0.14
0.1

Saturated
Saturated
Saturated
Gas
Saturated

0.2 mM NaOH
0.5 M KHCO3
0.2 M KHCO3
Gas
1 mL

BiVO4 , monoclinic
BiVO4 , monoclinic

0.20
0.20

Saturated
Saturated

100 mL
100 mLInNbO4
HNb3 O8
Pt-NaNbO3
Zn2 GeO4
RuO2 –Zn2 GeO4
RuO2 -Pt-Zn2 GeO4
Cu-BaLa4 Ti4 O15 b
Ag-BaLa4 Ti4 O15 b
Ag-CaLa4 Ti4 O15 b

0.14
0.1
0.1
–
–
–
0.3
0.3
0.3

Saturated
Gas (94 kPa)
Gas (80 kPa)
Gas
Gas
Gas
Saturated
Saturated
Saturated

0.2 M KHCO3
Gas (7 kPa)
3 mL
Gas
Gas
Gas
360 mL
360 mL
360 mL

Ag-SrLa4 Ti4 O15 b

0.3

Saturated

360 mL400 W HP Hg

Ag-BaLa4 Ti4 O15 c

0.3

Saturated

360 mL400 W HP Hg

cAg-CaLa4 Ti4 O15
Ag-SrLa4 Ti4 O15 c

0.3
0.3

Saturated
Saturated

360 mL
360 mL400 W HP Hg
400 W HP Hg[ReI (CO)3 (dcbpy)Cl]-Zr6 O4 (OH)4 (bpdc)∼0.02

Saturated

Acetonitrile(2 mL), TEA(0.1 mL)

450 W Xe(>300 nm)

313
303

273
273

318

[14]
[14]
[14]

[60]
[60]
[60]
[60]
[61]

a

Spectroscopic veriﬁcation of reaction route, isotope tracing to identify the C source, or sufﬁcient control kinetic tests to verify the photocatalytic event is required.
Cu or Ag was photodeposited during catalyst preparation.
c
Ag was reduced in liquid phase during catalyst preparation.

brule out the presence of carbon impurity intermediates derived
from CTMA+ cations. In the FTIR spectrum for the as-prepared
monoclinic BiVO4 , very weak peaks appear in the 3000–2900 cm−1
region [55], which are likely due to the presence of C H bonds
derived from CTMA+ or PEG used during the catalyst preparation.
Delafossite CuGa1−x Fex O2 (x = 0.15) was tested in an ambientpressure of CO2 -saturated H2 O under a Xe arc lamp. The formation rate of the major product CO was reported to the 19% of thecorresponding rate to form methane obtained with CdSe-Pt/TiO2
under visible light [26,56], indicating the superiority of TiO2 -based
photocatalysts.
Single crystalline Zn2 GeO4 nanoribbons with a thickness as
small as 7 nm, a width of 20–50 nm, and a length of hundredsof micrometers were applied in the photoreduction of CO2 in
the presence of water [57]. The band-gap value of the nanoribbons was estimated to 4.5 eV, and their speciﬁc surface area was
28 m2 g−1 . The methane formation rate (0.41 ␮mol h−1 gcat −1 ) for
13–16 h under light using the Zn2 GeO4 nanoribbons increased to
1.9 ␮mol h−1 gcat −1 with the addition of 1 wt% RuO2 , and further to
6.5 ␮mol h−1 gcat −1 with the further addition of 1 wt% Pt (Table 1C).
The crystalline nanoribbons were suggested to suppress charge
recombination and facilitate charge transport from the bulk to thesurface-active sites. Zinc germanium oxynitride was prepared from
Zn2 GeO4 with ammonia at 1073 K [58]. The band gap increased to2.70 eV, and the rate of CO2 photoreduction to methane increased.

The positive shift of the top of the valence band was considered toenhance water photooxidation and effectively supply the produced
protons to the CO2 .
Layered-perovskite photocatalysts AII La4 Ti4 O15 (A = Ca, Sr, andBa) that have been reported for complete water splitting [59] wererecently applied to the CO2 photoreduction [60]. In a series of photocatalytic tests in CO2 -saturated water under a 400-W Hg lamp,
CO was formed in a competitive reaction to produce H2 and a minor
amount of formic acid. The reactivity order to form CO wasAg(2 wt%)-BaLa4 Ti4 O15 -LPR (73)
> Ag(1 wt%)-CaLa4 Ti4 O15 -LPR (31)
> Ag(1 wt%)-SrLa4 Ti4 O15 -LPR (24)
> Ag(1 wt%)-BaLa4 Ti4 O15 -PD (14)> Ag(1 wt%)-CaLa4 Ti4 O15 -PD (7.7)> Ag(1 wt%)-SrLa4 Ti4 O15 -PD (6.0),
where LPR and PD indicate Ag doping to photocatalysts by liquidphase reduction and photodeposition, respectively, and the values
in the parentheses are the CO formation rates in the unit of␮mol h−1 gcat −1 (Table 1C). The formation rates of the by-productformic acid were 4.3–1.0 ␮mol h−1 gcat −1 . Silver nanoparticles ofAuthor's personal copy
Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186∼10 nm on the edge sites of the layered BaLa4 Ti4 O15 were suggested to be favorable for the reduction of CO2 to CO.
2H2 O + 4 h+ (basal plane of BaLa4 Ti4 O15 ) → 4H+ + O2 ,E ◦ (298 K) = +1.229 V
CO2 + 2H+ + 2e− (Ag nanoparticles on the edge)→ CO + H2 O, E ◦ (298 K) = −0.11 V
It was also discussed that the spatial separation of the oxidation
and reduction sites suppressed the reverse reaction of CO to CO2 .
The chemistry of metal-organic frameworks (MOFs) has
recently attracted attention owing to the freedom of element
choice and the size of the micro/mesopore for realizing designed
chemical functions, such as sorption and catalysis. The immobilization of homogeneous photocatalysts for the CO2 photoreduction
to an MOF [61] can be compared to the semiconductor-type
photocatalysts in this review. Among the various homogeneous
photocatalysts for the CO2 photoreduction, including cobalt, nickel,
iron, and rhenium complexes [6], ReI (CO)3 (dcbpy)Cl (4.2 wt%;
dcbpy = 2,2 -bipyridine-5,5 -dicarboxylic acid) was incorporated
to a Zr6 O4 (OH)4 (bpdc) (bpdc = para-biphenyl-dicarboxylic acid)framework. When the catalyst was placed in CO2 -saturated acetonitrile including triethylamine as a sacriﬁcial electron donor
under UV–visible light from a 450-W xenon lamp (>300 nm),the turnover number to CO was 10.9 in 20 h, corresponding to
42 ␮mol gcat −1 h−1 . The amount of catalyst includes both the Re
complex and the MOF. After the 20-h reaction, 43.6% of the Releached into the supernatant, indicating the detachment of the Re
carbonyl species from the MOF.
When ReI (CO)3 (dcbpy)Cl was used as a homogeneous photocatalyst under similar conditions, the turnover number to CO was
3.5 (corresponding to 760 ␮mol gcat −1 h−1 ) in 6 h, but the catalyst deactivated after 6 h, even when triethylamine was added to
the reaction system at 6 h. Control reactions were also reported
in the absence of CO2 , in the dark, and in solvent labeled withisotopic CD3 CN, and the lack of formic acid/methanol formation
was conﬁrmed by proton nuclear magnetic resonance (1 H NMR)
spectroscopy. These results strongly suggested the photocatalytic
formation of CO, but recyclability tests indicated that catalyst deactivation occurred after two 6-h reaction runs [61].
If the scope of this review is widened to include photocatalysts that require a sacriﬁcial electron donor in order to reduce theCO2 , colloidal ZnS [62,63] and CdS [64,65] nanoparticles should beconsidered. These catalysts formed formic acid and CO with highquantum yields when excited under UV–visible light [62,63,65] orvisible light only (>400 nm) [64].
In summary, CdS, SiC, InNbO4 , HNb3 O8 , BiWO6 , promoted
NaNbO3 , and promoted Zn2 GeO4 produced methane or methanol
with rates of 1–10 ␮mol h−1 gcat −1 , and promoted AII La4 Ti4 O15 produced CO with a rate greater than 10 ␮mol h−1 gcat −1 . The morenegative conduction band energy for CdS, SiC, Nb, and Ta seems
to be one of the key factors for effective CO2 photoreduction[14,59,66]. A homogeneous Re complex supported on an MOF
improved the life of the homogeneous photocatalyst.
3.7. Carbon-based photocatalysts
Small carbon nanoparticles of less than 10 nm can be covalently
functionalized, e.g., with PEG, to give them strong absorption andemission properties in the visible light region. Gold was photoreduced onto functionalized C nanoparticles, and the photocatalytic
conversion of CO2 to formic acid in an aqueous solution wasreported with a quantum yield of 0.3% based on 1 H NMR analysis[67], but detailed quantitative kinetic data were not shown in the179

paper. Semiconductor-like charge separation was suggested, andthe advantage of the Au-functionalized C was claimed due to theaqueous solubility. Separately, efﬁcient CO formation was reported
using graphitic carbon nitride under visible light (>420 nm); however, it is essential to verify whether the CO was derived from CO2 ,and not from the surface functional groups or the carbon nitride
itself, before any further discussion is warranted [68].
A solvent-exfoliated graphene (SEG) dispersion, obtained via
the ultrasonic treatment and centrifugation of natural graphite in
N,N-dimethylformamide, was mixed with TiO2 (P25) and ethyl cellulose to form ﬁlms that were calcined at 673 K [69]. The rate ofphotoreduction of CO2 to methane using the SEG (0.27 wt%)–TiO2(8.3 ␮mol h−1 m−2 ) was 4.5 times greater than that using TiO2 (P25)
under UV light. The less-defective SEG was a better photocatalytic
promoter than reduced graphene oxide, demonstrating an electric
diffusion effect, rather than the presence of reactive defect sites.
A disorder-engineered black TiO2 photocatalyst prepared byhydrogenation at 473 K was reported to be effective for the water
reduction reaction with the use of a sacriﬁcial electron donor. A
mid-gap electronic state resulting from the introduction of thestructural disorder was assumed to change the powder color andshorten the band gap. The color was stable for more than a year[70]. Unfortunately, the black color was later reported to be dueto the presence of ∼0.6 ␮mol of chromium contained per unit
gram of the TiO2 nanotubes, which originated from the passivating Cr oxide layer of the stainless steel autoclave [71]. The visible
light response using metal-cation-doped TiO2 photocatalysts has
already been established as being due to the impurity energy
level below the conduction band [36,72–74]. Although dark blue
TiO2 with structural disorder that was free from Cr was instantly
oxidized back to white TiO2 at room temperature in air, the application of TiO2 with structural disorder to the CO2 photoreduction is
expected.
4. Photon energy conversion of CO2 to fuels using hydrogen4.1. Photocatalytic conversion of CO2 to methane or CO usinghydrogenThe reduction of CO2 with hydrogen is thermodynamically favorable compared to the reduction with water(Gr ◦ = +689 kJ mol−1 ; Section 3.5) as follows.
CO2 (g) + 3H2 (g) → CH3 OH(g) + H2 O(g), Gr ◦ = +2.9 kJ mol−1
CO2 (g) + 4H2 (g) → CH4 (g) + 2H2 O(g), Gr ◦ = −113.6 kJ mol−1
In this case, we assume that hydrogen is supplied via sustainable ways, e.g., photocatalytic water reduction utilizing sunlight[59,75,76].
The reduction of CO2 to methane with H2 gas was ﬁrst reported
using Ru/RuOx /TiO2 [77]. The initial formation rate of methane was
49 ␮mol h−1 gcat −1 at 319 K using light from a solar simulator witha total intensity of 80 mW cm−2 (Table 2). The CO2 conversion was
believed to be due to thermal and UV light effects. However, a later
report concluded that it was solely a thermal effect [78]. Comparable or slower methane formation to that reported in Ref. [77]was reproduced at 295 and 373 K in the dark. The enhancement ofmethane formation by light was completely suppressed by placing a 25-mm water ﬁlter on the 150-W xenon lamp. Even when
the irradiation power was varied between 2 and 350 mW cm−2 , nochange in the kinetic results was observed as long as the water ﬁlter
was used.
The photoreduction of CO2 using hydrogen as a reductant has
also been reported to produce carbon monoxide using Rh/TiO2(5.1 ␮mol h−1 gcat −1 ) [79], ZrO2 (0.56 ␮mol h−1 gcat −1 ) [80], andAuthor's personal copy

[3]
[3]
[3]
[3]
[4]
[4]
[3]
[3]
CO (0.37), methanol (0.13)
CO (0.58), methanol (0.20)
CO (0.080), methanol (0.051)
Methanol (0.17), CO (0.079)
Methanol (0.30), CO (0.13)
Methanol (0.49), CO (0.070)
CO (0.030)
CO (0.047)
500 W Xe arc (42 mW cm−2 )500 W Xe arc (106 mW cm−2 )500 W Xe arc (42 mW cm−2 )500 W Xe arc (42 mW cm−2 )500 W Xe arc (42 mW cm−2 )500 W Xe arc (42 mW cm−2 )500 W Xe arc (42 mW cm−2 )500 W Xe arc (42 mW cm−2 )4.2. Photocatalytic conversion of CO2 to methanol using hydrogen

305–313
305–313
305–313
305–313
305–313
305–313
305–313
305–313[3]500 W Xe arc (42 mW cm−2 )313

305–313[77] a
[79]
[80]
[81]
[82]
[83]
[83]
[31]

Methane (49)
CO (5.1)
CO (0.56)
CO (1.6)
CO (0.76)
Methane (4.1)
CO (0.62)
Methane (0.012), ethene(0.008), ethane (0.005)
CO (0.62), methanol (0.039)

MgO (1.6 ␮mol h−1 gcat −1 ) [81] under CO2 + H2 (total 25 kPa) illuminated with a Hg lamp and ␤-Ga2 O3 (0.76 ␮mol h−1 gcat −1 )under CO2 + H2 (total 3.2 kPa) illuminated with a Hg–Xe lamp [82](Table 2). When TiO2 and ZrO2 in CO2 + H2 + H2 O and CO2 + H2 were
illuminated with UV light of 365 and 254 nm from a near-UV ﬂuorescent black lamp, methane and CO were produced at rates of
4.1 ␮mol h−1 gcat −1 and 0.62 ␮mol h−1 gcat −1 , respectively [83].
In summary, rhodium was an effective additive, but only simple
metal (Ti, Zr, Mg, and Ga) oxide semiconductors photoreduced CO2
with molecular H2 at a rate of ∼1 ␮mol h−1 gcat −1 .

a

0.10
0.10
0.10
0.10
0.10
0.10
0.10
0.10
[Zn1.5 Cu1.5 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O[Zn1.5 Cu1.5 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O[Zn3 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
Cu-ZnO
␣-Ga2 O3

2.3 kPa(0.18 mmol)
2.3 kPa
2.3 kPa
2.3 kPa
2.3 kPa
2.3 kPa
2.3 kPa
2.3 kPa
2.3 kPa
0.10[Zn3 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O

Sufﬁcient control kinetic tests to deny the photocatalytic event were reported in Ref. [78].

22 kPa
22 kPa
22 kPa
22 kPa
22 kPa
22 kPa
22 kPa
22 kPa

22 kPa (1.7 mmol)

H2
61 kPa50 ␮mol
50 ␮mol
50 ␮mol
50 ␮mol
5.6 kPa 100 kPa
106 kPa
76 ␮molCO2
5.1 kPa150 ␮mol150 ␮mol150 ␮mol150 ␮mol
5.6 kPa5.6 kPa38 ␮mol
Amount (g)
0.1
0.3
0.3
0.3
1.0
0.25
0.25
0.05
Brand name
Ru/RuOx /TiO2
Rh/TiO2
ZrO2
MgO
␤-Ga2 O3
TiO2
ZrO2
Ti-SBA-15 (0.05 wt% Ti)

H2 O
Reactants (gas)Photocatalyst

Table 2
Reported CO2 photoreduction catalysts, reaction conditions, and the formation rates with hydrogen [and moisture].

T (K)

319

Light source

Solar simulator
500 W HP Hg
500 W HP Hg
500 W HP Hg
200 W Hg–Xe
UV (365 nm)
UV (254 nm)
120 W HP Hg

Reference

Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Major product {formation rate(␮mol h−1 gcat −1 )}180

Layered-double hydroxide (LDH) compounds are a family ofclays, but their layers are positively charged in contrast to moregeneral clays, in which the layers are negatively charged, e.g., thesmectite and vermiculite families [3,84–86]. LDHs are materials
based on the layered structure of brucite (Mg(OH)2 ), which has ahexagonal crystal structure in which the MgO6 octahedra are linked
at the edge to form sheets. One of the naturally occurring LDH
compounds is hydrotalcite. The charge of cationic sheets formulated as [MII 1−x MIII x (OH)2 ]x+ is compensated with an intercalatedanion group, e.g., CO3 2− , SO4 2− , NO3 − , Cl− , or OH− . The MII site
can be Mg, Mn, Fe, Co, Ni, Cu, or Zn, and the MIII site can be Al,
Cr, Mn, Fe, or Ga. The value of the variable x is normally within therange 0.17–0.33. The molar amount of structural water intercalated
between the cationic layers is about half of the total molar amountof metal cations.
LDH compounds have already been applied to photocatalytic water oxidation [87,88]. Zn4 TiIV x , Zn4 CeIII x , and Zn4 CrIII x(0.25 < x < 2) LDH compounds exhibited good quantum yields forphotocatalytic water oxidation in the presence of silver nitrate as asacriﬁcial oxidizing reagent under visible light (>400 nm) [87]. The
LDH compound [Zn0.69 CrIII 0.31 (OH)2 ]0.31+ 0.31NO3 − ·0.6H2 O and thenanohybrid material Zn Cr LDH and layered titanate were also
effective for photocatalytic water oxidation in the presence of
0.01 M AgNO3 as a sacriﬁcial oxidizing reagent under visible light(>420 nm) [88].
Ordered LDHs consisting of zinc and/or copper hydroxides combined with aluminum have been reported to begood photocatalysts for the conversion of gaseous CO2 tomethanol or CO under UV–visible light using hydrogen [3].
The LDH compound [Zn3 AlIII (OH)8 ]+ 2 (CO3 )2− ·mH2 O (band gap:
5.7 eV) was the most active for the CO2 photoreduction,and the major product was CO, which formed at a rate of
0.62 ␮mol h−1 gcat −1 , along with minor amounts of methanol(Table 2). The methanol selectivity (5.9 mol%) modestly increased
to 26%, but the total photoactivity decreased, when Cu sites wereintroduced in the LDH, i.e., [Zn1.5 Cu1.5 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O(band gap: 4.1 eV). The CO2 conversion was 0.16–0.11% using[Zn3−x Cux Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O (0 ≤ x ≤ 1.5) photocatalysts.
Ordered LDHs consisting of zinc and/or copper hydroxides combined with gallium were synthesized, and the methanol formation
rate and selectivity using [Zn1.5 Cu1.5 GaIII (OH)8 ]+ 2 (CO3 )2− ·mH2 O
LDH (band gap: 3.5 eV) were 0.17 ␮mol h−1 gcat −1 and 68
mol%, respectively. The CO2 conversion was 0.03–0.02% using[Zn3−x Cux Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (0 ≤ x ≤ 1.5) photocatalysts.
The speciﬁc interaction of the Cu sites with CO2 was spectroscopically suggested (based on XANES analysis) to enable coupling with
protons and photogenerated electrons to form the methanol. The
participation of Cu sites in the redox process in the CO2 photoreduction was also suggested [5].
The rate of photoconversion of CO2 to methanol was improved
to 0.49 ␮mol h−1 gcat −1 by replacing the interlayer carbonate
anions of [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O with [Cu(OH)4 ]2− ,Author's personal copy
Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

and the methanol selectivity increased to 88 mol% [4]. The band-gap
value of the improved LDH catalyst was 3.0 eV. A direct electronictransition from the O 2p to the metal 3d, 4 s, or 4p was suggested
for the photocatalysis excited largely by UV.
No chemicals containing carbon atoms were usedthroughout
thesynthesis
of
the
LDH
compounds[Zn3−x Cux A(OH)8 ]+ 2 X2− ·mH2 O (A = Al, Ga; X = CO3 , Cu(OH)4 ;0 ≤ x ≤ 1.5) except for sodium carbonate, or during the pretreatment prior to the catalytic tests. After the synthesis of theLDH, the powder was thoroughly washed using deionized water(<1.0 ␮S cm−1 ) before drying. The sodium carbonate was thesource of the interlayer anions, and the products produced in thereaction were essentially identical with the species derived fromthe photoconversion of gas-phase CO2 . Furthermore, no products
were detected when the [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
compound was suspended in deionized water under UV–visible
light for 5 h [3]. Thus, the possibility [38,39] that carbon impurities
that accumulated during the LDH catalyst preparation and/or
activation before the kinetic tests were conducted was converted
rather than CO2 is not plausible in Refs. [3,4].
thepossibilityofathermalreaction
However,should also be seriously considered [77,78]. The inprocess
spectrafor[Zn3 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
and[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O during the photoreaction in CO2 + H2 were reported (Fig. 6). The formation ofboth CO and methanol reached a maximum at a wavelength of 400 nm both for [Zn3 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O and[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (Fig. 6, inset).
This trend is in clear contrast to methane formation using
Ru/RuOx /TiO2 . At ﬁxed temperature (319 K) in CO2 + H2 , themethane formation rate increased by 2.8 times when the reaction
mixture was irradiated with light from a solar simulator [77], but
a second study found that the light effect was due to an increase in
the temperature from 298 to 314 K and suggested the temperature
control in Ref. [77] was not exact. When a water ﬁlter was inserted
between the catalyst and a 150-W xenon lamp, no effect of the light
was observed [78]. In contrast, when a water ﬁlter was inserted
for the [Zn1.5 Cu1.5 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O catalyst (Fig. 7B), themethanol and CO formation rates increased (Table 2) owing toan increase in the UV + visible light intensity from 42 mW cm−2

181

Fig. 6. Dependence of the formation rates of methanol and CO on the cutoff wavelength for photoreactions in CO2 (2.3 kPa) + H2 (22 kPa) using 0.10 g of[Zn3 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O or [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O. Sharp cutoff ﬁlters UV-32, L-37, L-42, or Y-48 were used at the exit of the Xe arc lamp. (Inset)
In-process spectrum of product formation rates versus wavelength [3].

(Fig. 7A) to 106 mW cm−2 (Fig. 7B). The maximum temperature during these photocatalytic reactions was 313 K, but a control reaction
using [Zn1.5 Cu1.5 Al(OH)8 ]+ 2 (CO3 )2− ·mH2 O in the dark at 296–313 K[3,89] under CO2 + H2 exhibited no conversion above the detection
limit of the GC. Thus, the catalysis using LDH compounds underCO2 + H2 proceeded photocatalytically.
In summary, using LDH photocatalysts consisting of Zn, Cu,
Al, and Ga, both CO and methanol were formed at rates of
0.1–1 ␮mol h−1 gcat −1 .
5. Photon energy conversion of CO2 to fuels using a newreaction system
5.1. Recycling of sacriﬁcial electron donors
The regeneration of a sacriﬁcial electron donor in the photoreduction of CO2 was reported [8,90] to make the overall catalytic
reaction system cyclic. The tricyclic tertiary amine 10-endo-anti11-aza-10-methoxy-11-methyltricyclo [4.3.1.12,5 ] undecane wasFig. 7. Quartz photocatalytic reactor illuminated using a Xe arc lamp Model UI-502Q, Ushio (a) and Model SX-UID502XAM, Ushio (b) used in Refs. [3,4].

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182

Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Fig. 8. Recyclable photocatalytic CO2 reduction to formic acid using a sacriﬁcial electron donor amine that can be regenerated by hydrogenation using a Pd catalyst.
Reprinted by permission from Macmillan Publishers Ltd: Michl [8]. Copyright (2011).

selected as the sacriﬁcial electron donor (Fig. 8). An electron,a proton, and a hydrogen atom were donated by the amine,and the CO2 was converted to formic acid using a p-terphenyl
photocatalyst under irradiation from a Hg lamp at 254 nm. The
amine was transformed to the stable alkene 10-endo-anti-11aza-10-methoxy-11-methyltricyclo [4.3.1.12,5 ] undec-7-ene. The
alkene was successfully hydrogenated back to the original alkane
using a Pd/C catalyst. In Ref. [90], the hydrogen was assumed
to be generated from the photosplitting of water, similar to thephotocatalytic CO2 reduction using H2 in Refs. [3,4].
5.2. Photoconversion of CO2 to fuels utilizing anode oxidation andcathode reduction compartments
Various photocatalysts for the photooxidation of water havebeen reported [59,75,91,92]. The idea of an artiﬁcial dark reaction
following water photooxidation was also proposed [7,9,93].
In 2006, a solar fuel cell was proposed on the basis of the concept of a polymer electrolyte fuel cell (PEFC) and a phosphoric acidfuel cell consisting of an electrolyte for the transfer of protons fromthe anode side to the cathode side. A solar photovoltaic assembly was proposed to separate the holes and electrons, an anodecatalyst to oxidize water using the holes, and a cathode catalyst
to reduce the protons to hydrogen using the electrons (Fig. 9). In
photosynthesis, protons are reduced by Photosystem I to hydrogenequivalents. In other words, electrons are stored through the conversion of NADP to NADPH. Thus, if the photon conversion device
is based on a single-bandgap absorber consisting of a semiconductor, the theoretical thermodynamic conversion efﬁciency is 32% in
unconcentrated sunlight [9].
A dye-sensitized solar cell utilizing a polymer membrane and amanganese catalyst was proposed in 2009 [93]. Inspired by Photosystem II, in which a cube-shaped Ca Mn oxide catalyst oxidizes
water to O2 , protons, and electrons, the synthetic model Mn oxide
cluster complex Mn4 O4 was incorporated into a proton-conducting
membrane, e.g., Naﬁon. The complex and an H+ -conducting membrane were integrated into a solar cell (Fig. 10). An organic
ruthenium dye captured sunlight, and the excited electrons were
injected into the neighboring TiO2 nanoparticles. The manganese
catalyst also captured sunlight and oxidized water. The grabbed

Fig. 9. Concept of a fuel cell supplied with hydrogen and oxygen (or air) that converts chemical energy to electricity. The solar fuel cell uses light to separate holes andelectrons that oxidize water and reduce protons, respectively.
Source: Lewis and Nocera [9]. Copyright (2007) National Academy of Sciences, USA.

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

electrons from the water were passed from the Mn catalyst to thedye molecules. The electrons were transmitted to the cathode via
an external circuit, while the protons transmitted through the H+ conducting membrane were reduced to H2 on the cathode catalyst
with the electrons. A solar cell (or solar fuel cell using water as thefuel) integrating an H+ -conducting membrane and an iridium oxide
catalyst was also reported.
The concept of the combined photocatalytic reduction of CO2
was also proposed in 2010 [7]. In this system, a water oxidation
catalyst releases protons and electrons, followed by competitive
reduction reactions of the protons to form H2 and react with CO2
to form fuel(s) (Fig. 11). If good ligands are chosen for the homogeneous water oxidation catalyst complexes, O2 molecule formation
is synchronized with the removal of the H+ and the electrons. However, the drawback of rapid deactivation often exists because of thedecomposition of the ﬁnely tuned ligands, e.g., heterocyclic organic
compounds. Heterogeneous catalysts for water oxidation may be
more stable and self-reparable.
The feasibility of using methanol oxidation as a source of
H2 in a PEFC [94] was also investigated. The fuel cell catalysts
consisted of a TiO2 photocatalyst for the oxidation of methanol to
CO2 and a Pt catalyst to reduce the protons to H2 . The two catalysts were separated by an H+ -conducting polymer (Fig. 12). The
amount of TiO2 was optimized to 3.0 mg cm−2 on carbon paper,and the Pt C at 0.2 mg cm−2 on C paper, depending on the catalytic
rate of each. Under UV–visible light illumination (100 mW cm−2 )using a Cu sulfate ﬁlter for the anode (TiO2 ) in a 1 M methanol solution, 0.1 M H2 SO4 as the electrolyte, and (Pt C) as the cathode in a

183

Fig. 10. Dye-sensitized solar cell utilizing a polymer membrane and a manganese
catalyst [93].
Copyright permitted by Dr. Gerhard Swiegers, University of Wollongong.

Fig. 11. Combination of water oxidation catalyst and the reduction of protons to H2 or CO2 to fuels. The reduction is promoted as a result of charge separation under lightirradiation. The reduction catalyst might be assembled on an electrode, a suspension of nanoparticles, or conducting organic or inorganic membranes capable of directing
charge transport.
From Hurst [7]. Reprinted with permission from AAAS.

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184

Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

Fig. 12. Feasibility test of a fuel cell consisting of a TiO2 photocatalyst for methanol oxidation, a Pt catalyst for the production of H2 , and an H+ -conducting polymer between
them.
Reprinted with permission from Seger and Kamat [94]. Copyright (2009) American Chemical Society.

0.1 M H2 SO4 solution, a short circuit current of 0.34 mA cm−2 was
obtained. Under similar reaction conditions, hydrogen was produced at a rate of 3.5 ␮mol h−1 cm−2 (1.2 mmol h−1 gTiO2 −1 ). The
photocatalytic reaction at the anode was proposed to proceed asfollows.
hv

TiO2 −→h+ + e−
CH3 OH + h+ → • CH3 O + H+

• CH O + H O → CO + 5H+ + 5e−
3
2
2

2H+ + 2e− → H2
In total, hydrogen was produced by the photocatalytic decomposition of methanol, and a stoichiometric electric current
was obtained for the overall reaction: CH3 OH + H2 O → CO2 + 3H2 .
This study described the reverse reaction of the photocatalytic

Fig. 13. Proposed reaction mechanism using a PEFC starting from H2 and CO2 (a) and from H2 O and CO2 (b) [100].

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Y. Izumi / Coordination Chemistry Reviews 257 (2013) 171–186

conversion of CO2 , but also utilized a Pt electrocatalyst to reduce
protons to H2 using a PEFC.
It is well known that plants and cyanobacteria use the reducing
power generated by light-driven water oxidation in Photosystem II
to produce NADPH by reducing the protons in Photosystem I [95].
2H2 O → O2 + 4H+ + 4e− (Photosystem II),E ◦ (pH 7) = +0.815 V

NADP+ + H+ + 2e−→ NADPH(Photosystem I), E ◦ (pH 7) = −0.320 V
The weak reductant formed by light energy in Photosystem II
reduces the weak oxidant formed by different light energy in Photosystem I. This linked mechanistic model is called the Z-scheme.
In 1982, artiﬁcial H2 formation starting from water was reported
using a photodiode consisting of a Mg-doped p-type iron oxide anda Si-doped n-type iron oxide [96]. Using the combination of PtWO3 for water oxidation and Pt-SrTiO3 for water reduction, thestoichiometric H2 and O2 photocatalytic formation from water was
also reported by applying the Z-scheme [97]. Furthermore, the separation of a TiO2 catalyst for water oxidation and a Pt-TiO2 catalyst
for proton reduction by Naﬁon was reported to separately form O2
and H2 gas [98].
The concept of the Z-scheme was also applied to the photoreduction of CO2 using semiconductor catalysts for the wateroxidation and homogeneous metal complex catalysts for the CO2
reduction. These studies have been previously reviewed as anextension of homogeneous photocatalysts of the photoreductionof CO2 [99].
The direct transfer of protons and electrons released fromwater oxidation catalysts to photocatalysts for the reduction of
CO2 must be advantageous for achieving efﬁcient rates for thephotoconversion of CO2 to fuels. A Pt C catalyst and an LDH
photocatalyst [Zn1.5 Cu1.5 GaIII (OH)8 ]+ 2 (CO3 )2− ·mH2 O (see Section
4.2) were directly attached to a 50-␮m-thick proton-conducting
membrane. The CO2 and water, supplied separately to eachcatalyst in the PEFC cell (Fig. 13B), were converted to methanol
at a rate of 5.1 ␮mol h−1 gcat(LDH) −1 . It has also been found thatheating at 413 K makes the membranes used in Refs. [4,100] more
H+ -conductive. The development of a total photocatalytic system
consisting of WO3 and Zn Cu Ga LDH photocatalysts for the
direct transfer of protons and electrons is underway [100].
6. Concluding remarks
The dramatic growth of renewable solar energy is needed.
The initial investigation of the renewable energy options utilizing the thermochemical and photocatalytic conversion of CO2 tofuels has been reviewed. Semiconductor photocatalysts are advantageous if they are cheap and sustainable. Starting from CO2
and water, TiO2 photocatalytically produced methane at a rateof ∼0.1 ␮mol h−1 gcat −1 . When TiO2 was atomically dispersed inY-zeolite, MCM-48, or SBA-15, or doped TiO2 was used, the rateof methane or CO formation increased to 1–10 ␮mol h−1 gcat −1 .
CdS, SiC, InNbO4 , HNb3 O8 , Bi2 WO6 , Pt-NaNbO3 , and RuO2 -PtZn2 GeO4 produced methane or methanol at rates greater than
1 ␮mol h−1 gcat −1 , and Ag-BaLa4 Ti4 O15 produced CO at a rate of73 ␮mol h−1 gcat −1 .
The photocatalytic reduction of CO2 with molecular hydrogen
was also surveyed. CO was formed at a rate of ∼1 ␮mol h−1 gcat −1
using TiO2 , ZrO2 , MgO, and Ga2 O3 , while CO and methanol wereformed at rates of 0.62 ␮mol h−1 gcat −1 and 0.49 ␮mol h−1 gcat −1 ,respectively, using LDHs consisting of Zn, Cu, Al, and Ga.
Spectroscopic monitoring using ESR, EXAFS, XANES, photoluminescence, diffuse reﬂectance UV–visible, FTIR, and NMR techniques,
and isotope monitoring using GC–MS and electron microscopy are

185

essential for the veriﬁcation of photocatalytic events and the sourceof C used to produce the fuels. The reaction mechanism of photocatalysis is still not well understood; e.g., the preferable formationof methane from CO2 + H2 O using TiO2 has not been explained. It
should be more complex than the consecutive reduction from CO2
to formic acid (or CO), the formyl group, formaldehyde, methoxy,
methanol, and then methane [3] based on recent reports using
ESR, which suggest formyl dimerization to glyoxal (OHC–CHO) as
an intermediate of an efﬁcient electron acceptor [101]. Theoreticalunderstanding based on the comparison between band energy andthe red-ox potential is in progress, but not enough to explain the
product selectivity using TiO2 [102].
Further efﬁciency improvements for the CO2 conversion arehighly expected by the combination of different photocatalysts(Section 4) and the use of new reaction systems (Section 5). In particular, the combination of water photosplitting to form hydrogen(or protons and electrons) and CO2 photoreduction with the formedhydrogen (or the formed protons and electrons) can potentiallyboost the efﬁciency of CO2 conversion.
Acknowledgements
The author thanks the Monbusho (MEXT) Research Fellowshipfor making it possible to do research as a Visiting Scientist in
the Groups of Professor Solomon and Professor Hodgson during
1996–1997. The author is also grateful for ﬁnancial support fromthe Grant-in-Aid for Scientiﬁc Research C (Proposal No. 2255 0117)
from MEXT (2010–2012) and the Feasibility Study Stage of A-STEP(Proposal No. AS231Z01459C) from the Japan Science and Technology Agency for research related to this review.
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