NEW AND
FUTUREDEVELOPMENTS
IN CATALYSISACTIVATION OF CARBONDIOXIDE
Edited by

STEVEN L. SUIB
Department of Chemistry and Chemical Engineering
and Institute of Materials Science, The University of Connecticut,
Storrs, CT 06269-3060

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C H A P T E R

20Photocatalytic Conversion
of Carbon Dioxide into Fuels
Using Layered Double
Hydroxides Coupled with
Hydrogen or Water
Naveed Ahmed, Motoharu Morikawa,and Yasuo Izumi
Department of Chemistry, Graduate School of Science, Chiba University,
Yayoi 1-33, Inage-ku, Chiba 263-8522, Japan

20.1 INTRODUCTION
Carbon dioxide is one of the major greenhouse gases. Several methods for reducing atmospheric CO2 concentration have been investigated, such as increasing green carbon sinks(plants, phytoplankton, and algae containing chloroplasts); increasing dissolved carbonate
and its salts in sea water; or capturing CO2 with transfer to the bottom of the sea in a supercritical state [1]. The global strategy for the suppression of CO2 emission based on Kyoto
Protocol is not on the right track due to conflicting interests between developed and developing countries and the uncertainty of an energy paradigm shift from fossil fuels to clean
energy. At present, the contribution of nuclear power is unstable due to the question of safety
whereas the contribution of natural energy (biomass, hydropower, solar, wind, and geothermal energy) to total energy consumption has been growing much slower than expected [2].
Thus, it is urgently important to investigate sustainable natural energy.
As one of the options of natural energy, it is advantageous to capture CO2 produced fromhuman activity and from the atmosphere and convert this gas to fuels using sustainable

Activation of Carbon Dioxide
http://dx.doi.org/10.1016/B978-0-444-53882-6.00021-8

589© 2013 Elsevier B.V. All rights reserved.

590

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS

sunlight, because this option solves the problems of global warming and sustainable energy
shortage simultaneously. Semiconductor-type photocatalysts are hopeful because they can becheaper by choosing inexpensive elements in comparison to metal complex homogenousphase photocatalysts consisting of precious noble metals [3]. Some good reviews for CO2
photoreduction have been reported focusing on TiO2-based photocatalysts [4], on the possibilityof solar fuel [5], and nanocomposite photocatalysts [6]. But the number of publications ofphotoreduction of CO2 is growing even rapidly later than 2010 and newer reviews are needed.
In this chapter, recent works for semiconductor-type photocatalysts to reduce CO2 are reviewed.

20.2 CO2 PHOTOREDUCTION WITH WATER USING METAL OXIDEAND OTHER SEMICONDUCTOR-TYPE PHOTOCATALYSTS
The photoreduction of CO2 follows a reduction to carbon monoxide or a consecutive reaction
to organic compounds as below:
2e−

CO2 → CO,
2e−

2e−

2e−

2e−

CO2 → HCO2 H → HCHO → CH3 OH → CH4 .
In contrast to most homogeneous photocatalysts that produce CO with the help of sacrificial reducing agents, e.g., alkyl amines and alcohols [6,7], various products ranging fromtwo-electron-reduced formic acid and CO to eight-electron-reduced methane were produced
using semiconductor photocatalysts often without the help of sacrificial reductant [4,8]. The
reasons why the selectivity to each product was determined and how multiple-electron photoreduction proceeds are not fully known [9].
In the absence of reducing agent, the photoreduction of CO2 to CO is a combination ofwater oxidation reaction and CO2 reduction reaction:

2H2 O + 4h+ →4H+ + O2 ,CO2 + 2H+ + 2e− →CO + H2 O.
CO was reported as a minor product in addition to major methane formation using semiconductor photocatalysts, TiO2 [10,11], and ZrO2 [10] (Table 20.1A).
The photocatalytic production of formic acid was reported using TiO2 under supercritical
CO2 [12,13]. The CO2 pressure of supercritical condition was reduced from 6.5 to 9.0 MPa,and then pure water or acidic solution was added to the reactor. In particular, formic acid−1 −1
was preferably formed (1. 8 µmol h gcat) when acidic solution was added as a protonsource (Table 20.1B) [13].
The photocatalytic formation (and detection) of formaldehyde has been reported very
rarely mostly because of the instability of formaldehyde. In spite of this problem, pioneering
work for photocatalysts reported formaldehyde formation using various semiconductors
with the formation rate order: CdS > ZnO ∼ TiO2 > GaP ∼ SiC [14] (Table 20.1C).
Photocatalytic conversion of CO2 into methanol in aqueous solution was reported in several papers, but care should be taken for the possibility that the C source of produced methanol was not CO2 but impurity C species accumulated during the catalyst preparation, typicallyderived from alkyl groups of precursor compounds. The reaction pathway to CO2 was20.2 CO2 PHOTOREDUCTION WITH WATER USING METAL OXIDE AND OTHER SEMICONDUCTOR

591

TABLE 20.1 Photocatalytic Reduction of CO2 with Water using Metal Oxide and other SemiconductorType Photocatalysts
Photocatalysts

Reactants

Light

Formation Rates−1 −1(µmol h gcat)

References

CO2(5%) + H2O(g, 95%)

(A) Product: CO
TiO2

LP Hg

0.11

[10]

CO2(5%) + H2O(g, 95%)

LP Hg

0.045

[10]

CO2(sat) + H2O(l)HP Hg arc

0.35

[11]

TiO2(anatase)CO2(6.5 MPa), H2O(l)

Xe

1.0

[12]

TiO2(anatase)CO2(9.0 MPa), H2O(l)

Xe

1.8

[13]

CO2(saturated) + H2O(l)

ZrO2
TiO2
(B) Product: HCO2H(C) Product: HCHO
CdS

Xe/HP Hg

29

[14]

CO2(saturated) + H2O(l)Xe/HP Hg

17

[14]

CO2(saturated) + H2O(l)Xe/HP Hg

16

[14]

CO2(saturated) + H2O(l)Xe/HP Hg

14

[14]

CO2(saturated) + H2O(l)Xe/HP Hg

14

[14]

Ti–Y-zeolite

CO2(24 µmol) + moisture(120 µmol)

HP Hg

0.031

[16]

Ti–MCM-48

CO2(24 µmol) + moisture(120 µmol)

HP Hg

3.1

[17]

Ti–SBA-15

CO2(36 µmol) + moisture(180 µmol)

HP Hg

0.081

[18]

TiO2(d = 14 nm)

CO2(saturated) + 0.2 N
NaOH

Hg

0.047

[19]

Ag–TiO2

CO2(saturated) + 0.2 N
NaOH

Hg

0.094

[20]

CO2 + moisture (111 kPa)

UV

0.014

[21]

CO2(24 µmol) + moisture(120 µmol)

HP Hg

0.046

[16]

Ti–MCM-48

CO2(24 µmol) + moisture(120 µmol)

HP Hg

7.5

[17]

Ti–SBA-15

CO2(36 µmol) + moisture(180 µmol)

HP Hg

0.31

[18]

Ti–FSM-16

CO2(36 µmol) + moisture(180 µmol)

HP Hg

0.15

[22]

HP Hg

0.084

[23]

ZnO
TiO2
GaP
SiC
(D) Product: CH3OH(E) Product: CH4
TiO2(JRC-TIO-4)
Ti–Y-zeoliteTi–porous SiO2 film

CO2(36 µmol) + moisture(180 µmol)

592

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS

monitored by Fourier-transform infrared (FTIR) spectroscopy [15] and kinetic study forthoroughly washed TiO2 (P25, Degussa) samples [11]. CO was mainly produced instead of
methane or methanol in the CO2 photoreduction test using thoroughly washed TiO2 (P25)with an Hg lamp (Table 20.1A).
For the mixture of CO2 gas and moisture, methanol was formed using atomically dispersed
Ti species on molecular sieves (Ti–Y-zeolite, Ti–MCM-48, and Ti–SBA-15) with UV-visible
light [16–18] (Table 20.1D), but the major product was methane in these reports (Table 20.1E).
The average particle size of 14 nm was optimized for TiO2 for the photoreduction of CO2 tomethanol [19]. The additive effects of metal nanoparticles on photocatalytic methanol formation were also reported [20].
TiO2 has been most frequently reported for the photoreduction of CO2 to produce methane.
TiO2 consisting of major anatase phase and minor rutile phase formed methane very slowly(0.014 µmol h−1 g−1) [21]. The photocatalytic activity per unit amount of Ti was significantly
catenhanced when Ti oxide species were anchored within the pores of zeolite [16], orderedmesoporous SiO2 [22], and porous SiO2 thin films [23], but the activity per unit amount of
catalyst modestly increased because the dispersion of Ti site increased [16–18,22,23] (Table
20.1E). Promoting effects of CdSe and Pt were discussed for the photocatalytic CO2 reduction
into methane [24].
In summary, semiconductors, typically TiO2, were active for CO2 photoreduction in water.
The photocatalytic rates using anatase/rutile-type TiO2 were mostly less than 1 µmol h−1 g−1cat,
but the photoactivity was increased to 1 − 10 µmol h−1 g−1
bydispersing
Ti
sites
on
molecularcatsieves, by the addition of promoters, e.g., CdSe and Pt, and under supercritical reaction
conditions.

20.3 CO2 PHOTOREDUCTION WITH HYDROGEN USING METALOXIDE-BASED AND OTHER SEMICONDUCTOR-TYPE
PHOTOCATALYSTS
The photoreduction of CO2 using hydrogen is an alternative option to reduce the amount of
CO2 sustainably, if the hydrogen was produced using a sustainable energy source, e.g., water
splitting using sunlight [1].
The photoreduction of CO2 using hydrogen as a reductant has been reported to produce
CO using Rh/TiO2 (5.1 µmol h−1 g−1) [25], ZrO2 (0.56–0.62 µmol h−1 g−1) [10,26], MgO
cat
cat
−1 −1
(1.6 µmol h gcat) [27], and β-Ga2O3 (0.76 µmol h−1 g−1
)
[28],
and
to
producemethane usingcat−1
Ti–SBA-15 (0.012 µmol h−1 gcat
) [29] (Table 20.2). Thus, metal oxide semiconductors, e.g.,
TiO2, ZrO2, MgO, and Ga2O3, photoreduced CO2 in the presence of H2 at the rate of∼1 µmol h−1 g−1cat. Rhodium was an effective promoter.

20.4 SYNTHESES AND CHARACTERIZATION OF LAYERED DOUBLEHYDROXIDES FOR CO2 PHOTOREDUCTION WITH HYDROGEN
Layered double hydroxide (LDH) compounds are a family of clay-like materials. Their layers
are positively charged in contrast to clay families in which the layers are negatively charged

20.4 SYNTHESIS AND CHARACTERZATION OF LAYERED DOUBLE HYDROXIDES FOR CO2 PHOTOREDUCTION

593

TABLE 20.2 Photocatalytic Reduction of CO2 with Molecular Hydrogen Using Metal Oxide and Other
Semiconductor-type Photocatalysts
CO2

H2

Light

Formation Rates−1 −1(µmol h gcat)

References

Rh/TiO2

150 µmol

50 µmol

HP Hg

5.1

[25]

ZrO2

150 µmol

50 µmol

HP Hg

0.56

[26]

ZrO2

5.6 kPa

106 kPa

UV

0.62

[10]

MgO

150 µmol

50 µmol

HP Hg

1.6

[27]

β-Ga2O3150 µmol

50 µmol

HP Hg

0.76

[28]

38 µmol

76 µmol

HP Hg

0.012

[29]

Photocatalysts(A) Product: CO(B) Product: CH4
Ti–SBA-15[1,30–32]. LDHs are the materials based on the layered structure of brucite Mg(OH)2, which
has a hexagonal crystal structure and in which the MgO6 octahedra are linked at the edge toIII
x+ is compensatedform sheets. The charge of cationic sheets formulated as [MII
1−x Mx (OH)2 ]
2−
2−
−
−
−
with an intercalated anion group, e.g., CO3 , SO4 , 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 x varies
normally within the range of 0.17–0.33. The molar amount of structural water intercalated
between the cationic layers is about a half of the total molar amount of metal cations.
Ordered LDHs consisting of zinc and/or copper hydroxides were reportedly combinedwith gallium. The interlayer anions were carbonates or [Cu(OH)4]2− [33]. The BET surface
area (SBET) values were in the range between 83 and 35 m2 g−1 (Table 20.3). The SBET value for2− · mH O was 23% greater than that for [Zn Cu Ga(OH) ]+[Zn3 Ga(OH)8 ]+
2
1.5
1.5
8 2
2 (CO3 )+2−(CO3 ) · mH2 O. The SBET values for [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ]2− · mH2 O decreased by 50%
compared to the value for [Zn3 Ga(OH)8 ]+2 (CO3 )2− · mH2 O because of the substitution of carbonate ions with [Cu(OH)4]2− ions. On the other hand, when compared with the SBET
2− · mH O, the values for [Zn Cu Ga(OH) ]+value for [Zn1.5 Cu1.5 Ga(OH)8 ]+
1.5
1.5
8 2
2
2 (CO3 )2−[Cu(OH)4 ] · mH2 O did not change significantly (+8.8%).
The X-ray diffraction (XRD) (003), (006), (009), (104), (015), (107), (018), (1010), (0111), (110),2− · mH O (Figure 20.1a)(113), and (116) peaks were observed for [Zn3 Ga(OH)8 ]+22 (CO3 )2− · mH O[Cu(OH)[34,35]. There was a downward shift in these peaks for [Zn3 Ga(OH)8 ]+
4]
2
2
+
2−
(b) and [Zn1.5 Cu1.5 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O (d). Weaker peaks derived from the CuO(+ mark) and Cu(OH)2 impurity phase (*mark) appeared in the spectra of
2− · mH O and [Zn Cu Ga(OH) ]+ [Cu(OH) ]2− · mH O (Fig[Zn3 Ga(OH)8 ]+
1.5
1.5
8 2
4
2
2
2 [Cu(OH)4 ]ure 20.1b and d). Diffraction peaks due to impurity phases were not found for
2− · mH O (a) and [Zn Cu Ga(OH) ]+ (CO )2− · mH O (c).
[Zn3 Ga(OH)8 ]+
1.5
1.5
8 2
3
2
2
2 (CO3 )The interlayer interval was 0.751 and 0.753 nm on the basis of the (003) diffraction angle
2− · mH O and [Zn Cu Ga(OH) ]+ (CO )2− · mH O (Table 20.3A
for [Zn3 Ga(OH)8 ]+
1.5
1.5
8 2
3
2
2
2 (CO3 )2− · mH O(COand B) [30]. The interlayer interval increased from 0.751 nm for [Zn3 Ga(OH)8 ]+
3)
2
2
+
2−
to 0.792 nm for [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O (Table 20.3A and C), reflecting the594

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS

2− · mH O (x = 0, 1.5;
TABLE 20.3 Physicochemical Characterization of [Zn3−x Cux Ga(OH)8 ]+22 [A]A = CO3, Cu(OH)4) LDHs

Eg (eV)

Entry
A

SampleSBET(m2 g−1)aInterlattice
Distance(nm)

Fit toα × hv ∝ (hv − Eg)n1
Extrapolated n = 2

3
2

2

3

(0 0 3)b

70 (83 )

5.6

5.9

5.4

5.2

5.0

0.751

B

2− · mH O[Zn3 Ga(OH)8 ]+22 (CO3 )+2−[Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O

57a (67b)

3.5

4.2

3.2

3.0

2.6

0.753

C

2− · mH O[Zn3 Ga(OH)8 ]+2
2 [Cu(OH)4 ]

35a

4.2

4.8

3.6

3.4

2.7

0.792

(impurity phase(s))

3.2

4.6

3.0

2.8

2.5

D

2− · mH2O[Zn1.5 Cu1.5 Ga(OH)8 ]+62a
2 [Cu(OH)4 ]

3.0

3.2

2.8

2.7

2.6

0.782

a

Preheated at 383 K for 2 h under vacuum. bPreheated at 423 K for 1 h under vacuum.

+
2− · mH O
FIGURE 20.1 XRD spectra of as-synthesized samples of [Zn3 Ga(OH)8 ]+(a), [Zn3 Ga(OH)8 ]222 (CO3 )++2−
2−[Cu(OH)4 ]2− · mH2 O (b), [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O (c), and [Zn1.5 Cu1.5 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O(d). (*) and (+) marks indicate peaks derived from Cu(OH)2 and CuO powders, respectively.

greater size of [Cu(OH)4]2− ions when compared to that of (CO3)2− [36,37]. In a similar manner, by substituting (CO3)2− with [Cu(OH)4]2−, the interlayer interval increased from 0.753 nm+
2−
for [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O to 0.782 nm (Table 20.3B and D).
In the UV-visible absorption spectra for LDH compounds (Figure 20.2), the Eg values for+
2− · mH O
2−[Zn3 Ga(OH)8 ]+and [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O were 5.6 and22 (CO3 )4.2 eV, respectively (Table 20.3A and C). Based on the corresponding n value (1/2 or 3/2) by
the fitting to the Davis-Mott equation, the electronic transition was direct from oxygen 2p to+2−
Cu 3d, 4s, or 4p levels and Zn/Ga 4s or 4p levels for [Zn3 Ga(OH)8 ]2 (CO3 ) · mH2 O and+2−[Zn3 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O. In addition, a shoulder peak corresponded to Eg values+
2−of 3.2 eV for [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O (Figure 20.2b) originated from the Cu(OH)2
and/or CuO impurity phase(s) that was evident in XRD pattern (Figure 20.1b). By fitting to
the Davis–Mott equation (n = 3/2; Table 20.3C), the electronic transition was forbidden direct

20.4 SYNTHESIS AND CHARACTERZATION OF LAYERED DOUBLE HYDROXIDES FOR CO2 PHOTOREDUCTION

595

FIGURE 20.2 Diffuse reflectance UV-visible absorption spectra of as-synthesized samples of [Zn3 Ga(OH)8 ]+
2

+
+
2−
2−
(CO3 )2− · mH2 O (a), [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O (b), [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O (c), and
2− · mH O[Zn1.5 Cu1.5 Ga(OH)8 ]+[Cu(OH)
]
(d).
4
2
2

from oxygen 2p to Cu 3d, 4s, or 4p levels of Cu(OH)2 and/or CuO phase(s). The UV absorp+
tion edge shifted toward the lower energy side for [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 )2− · mH2 O
2− · mH O (a). The E value was 3.5 eV(Figure 20.2c) when compared to [Zn3 Ga(OH)8 ]+g2
2 (CO3 )2−(Table 20.3B). Upon further replacing interlayer (CO3) ions with [Cu(OH)4]2− ions, the band
gap further decreased to 3.0 eV (Table 20.3D). This band-gap value corresponds to the boundary wavelength between UV and visible light.
Based on the Cu K-edge extended X-ray absorption fine structure (EXAFS) spectrum for
2− · mH O (Figure 20.3), two intense peaks inthe LDH sample of [Zn3 Ga(OH)8 ]+2
2 [Cu(OH)4 ]the Fourier transform at 0.16 and 0.27 nm (phase shift uncorrected) derived from Cu!O andCu … Zn (or Cu … Ga) pairs, respectively, appeared, but a peak derived from Cu!Cl bond(s)at ∼0.195 nm [38] (phase shift uncorrected) was not present (Figure 20.3b). No Cu!Cl peaks

2− · mH O sample. (a) k 3-weighted
FIGURE 20.3 Cu K-edge EXAFS spectra for [Zn3 Ga(OH)8 ]+2
2 [Cu(OH)4 ]EXAFS χ-function and (b) its associated Fourier transform. The solid and dotted lines represent the magnitude andthe imaginary part in (b).

596

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS

2− · mH O
FIGURE 20.4 Model structure of [Zn3 Ga(OH)8 ]+2 based on XANES and EXAFS analyses.
2 [Cu(OH)4 ]were found in the Fourier transform for this LDH photocatalyst, demonstrating the completehydrolysis of [CuCl4]2− into [Cu(OH)4]2− during the synthesis:[CuCl4 ]2− + 4NaOH → [Cu(OH)4 ]2− + 4NaCl.

(1)

2−
2[Zn3 Ga(OH)8 ]+ + [Cu(OH)4 ]2− + mH2 O ! [Zn3 Ga(OH)8 ]+· mH2 O.
2 [Cu(OH)4 ](2)

Based on the Cu K-edge X-ray absorption near-edge structure (XANES) spectra and thetheoretical simulation, complex formation between cationic layer and [Cu(OH)4]2− anion was
suggested [39] (Figure 20.4):[Zn3 Ga(OH)8 ]+2 [Cu(OH)4 ]2− · mH2 O ! Zn6 Ga2 (OH)13 {(µ−O)3 Cu(OH)}· mH2 O + 3H2 O. (3)20.5 CO2 PHOTOREDUCTION WITH HYDROGEN USING LAYEREDDOUBLE HYDROXIDESLDH photocatalysts were reported to produce not only CO but methanol in the CO2 photore+
2−
duction with hydrogen. While [Zn3 Ga(OH)8 ]2 (CO3 ) · mH2 O was CO selective (61 mol%),+2−[Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O was methanol selective (68 mol%; Figure 20.5a, b andTable 20.4A, B).
+
When [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 )2− · mH2 O was preheated at 423 K (Figure 20.5d), the−1 −1methanol formation rate increased by a factor of 1.8 (0. 31 µmol h gcat). Due to the preheat+2−
ing at 423 K, the SBET value for [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O increased by 18%, but
the increase did not fully explain the increase of photocatalytic methanol and CO formation
by a factor of 1.8 and 2.3, respectively (Table 20.4D). The available interlayer space that was

597

20.5 CO2 PHOTOREDUCTION WITH HYDROGEN USING LAYERED DOUBLE HYDROXIDESTABLE 20.4 Rates of Photocatalytic Conversion of CO2 with H2 into CH3OH and CO Over LDH
Photocatalystsa
Formation rate (µmol h

−1

Entry

PhotocatalystA

2− · mH O[Zn3 Ga(OH)8 ]+22 (CO3 )B

2− · mH O[Zn1.5 Cu1.5 Ga(OH)8 ]+22 (CO3 )C

2− · mH Ob[Zn3 Ga(OH)8 ]+22 (CO3 )D

2− · mH Ob[Zn1.5 Cu1.5 Ga(OH)8 ]+22 (CO3 )E

2− · mH O[Zn3 Ga(OH)8 ]+2
2 [Cu(OH)4 ]F[Zn1.5 Cu1.5 Ga(OH) ] [Cu(OH)4 ]
+
8 2

2−

· mH2 O

g−1
cat)

CH3OH

CO

Σ

Conversion(%, C-base)

0.051

0.080

0.13

0.02

(±0.004)(±0.006)

0.17

0.079

0.25

0.03

(±0.014)(±0.006)

0.050

0.074

(±0.004)(±0.006)

0.31

0.18

(±0.009)(±0.002)

0.30

0.13

(±0.009)(±0.010)

0.49

0.070

(±0.015)(±0.006)

Selectivity
to CH3OH(mol%)
39
(±4)
68
(±4)

0.12

0.02

0.50

0.07

40
(±4)
63
(±1)

0.43

0.04

0.56

0.05

71
(±2)
88
(±2)

a

The catalyst amount was 100 mg. 2.3 kPa of CO2 and 21.7 kPa of H2. Values in the parentheses are experimental errors for evaluation.
bPreheated at 423 K for 1 h under vacuum.

FIGURE 20.5 Time course of photocatalytic reactions in CO2 (2.3 kPa) + H2 (21.7 kPa). Hundred milligram of the++LDH catalyst was charged: fresh [Zn3 Ga(OH)8 ]2 (CO3 )2− · mH2 O (a), [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 )2− · mH2 O (b),
+
+
2−
2−
[Zn3 Ga(OH)8 ]2 (CO3 ) · mH2 O preheated at 423 K (c), [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O preheated at 423 K (d),++fresh [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ]2− · mH2 O (e), and [Zn1.5 Cu1.5 Ga(OH)8 ]2 [Cu(OH)4 ]2− · mH2 O (f). The reactor was
illuminated from a 500-W Xe arc lamp. CO2 (!; diamond), H2O (▲; triangle), CH3OH (■ÆÊÃµÕ>Ài®]Ê>`Ê "Ê­UÆÊVÀVi®°

598

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS

created at 423 K should facilitate the diffusion of CO2 into the reaction space and the reaction
with the surface hydroxy groups that are bound to the Cu sites.
By switching the interlayer anion group from carbonate to [Cu(OH)4]2− for LDH compounds consisting of [Zn3Ga(OH)8]+ cations, the methanol formation rate increased by afactor of 5.9 (Table 20.4E and A) while the methanol selectivity did not change (71–68 mol%).
By the substitution of anions, the SBET values even decreased by 50% (Section 20.4). For the
2− · mH O, the formation rates of methanol and CO for specific[Zn1.5 Cu1.5 Ga(OH)8 ]+22 (CO3 )quantities of Cu were 0.051 and 0. 024 µmol h−1 mmol−1
Cu, respectively. These values increased
to 0.097 and 0. 056 µmol h−1 mmol−1,
respectively,
when
preheated at 423 K in a vacuum. ForCu+
the [Zn3 Ga(OH)8 ]2 [Cu(OH)4 ]2− · mH2 O, the formation rates of methanol and CO were 0.30
and 0. 13 µmol h−1 mmol−1, respectively. Therefore, the interlayer Cu sites were 5.3–5.9 times
Cu
more effective than Cu sites in cationic layers, if we assume that the Cu atoms are primaryactive sites.
2− · mH O (Figure 20.4f), the methanol formation
Using [Zn1.5 Cu1.5 Ga(OH)8 ]+2
2 [Cu(OH)4 ]rate and selectivity further increased to 0. 49 µmol h−1 g−1
cat and 88 mol% (Table 20.4F). The
methanol formation rate per unit amount of Cu was 0. 12 µmol h−1 mmol−1(Table 20.5F),
Cu−1−1
which is in close agreement with the calculated value of 0. 11 µmol h mmolCu, which wascontributed by the in-layer Cu sites (3/4 × 0. 051 µmol h−1 mmol−1Cu) and the interlayer Cu
sites (1/4 × 0. 30 µmolh−1 mmol−1).
Thestericavailability(accessibility)of {(µ–O)3Cu(OH)}Cusites (Figure 20.4) may be related to the reactivity difference.
In this context, the Cu site plays a photocatalytic role by binding CO2 and coupling CO2
with protons and photogenerated electrons, utilizing the CuI and CuII redox couple [1,33].
This binding between CO2 and Cu was spectroscopically monitored in the multiple scattering+
region of Cu K-edge XANES for [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 )2− · mH2 O. The binding suggested the reaction mechanism via the CO2 insertion to hydroxy group to form hydrogencarbonate (Scheme 20.1a→b). The one-electron reduction steps b→c→d in Scheme 20.1
should be promoted utilizing the CuI and CuII redox couple. Subsequent one-electron reduction steps e→f, f→g, and g→h should also be coupled with the CuI and CuII redox. Final
one-electron reduction in the step h→a produces methanol. If the species e is dehydrated,
another catalytic reaction cycle to form CO is proposed regenerating species a. The role of
the CuI sites was also suggested in the CO2 photoreduction in alkaline aqueous solution [40].
The photocatalytic energetic diagram is depicted in Scheme 20.2. The band-gap values
2− · mH O and
2− · mH O(5.6–3.5 eV) for [Zn3 Ga(OH)8 ]+[Zn1.5 Cu1.5 Ga(OH)8 ]+222 (CO3 )
2 (CO3 )TABLE 20.5 Rates of Photocatalytic Formation of CH3OH and CO per the Quantity of Cu in the LDH
Photocatalystsa
Formation Rate Per Amount of Cu (mmol h−1 mmol−1
Cu )
Entry

CH3OH

CO

Σ

0.051

0.024

0.077

D

2− · mH O[Zn1.5 Cu1.5 Ga(OH)8 ]+22 (CO3 )2− · mH Ob[Zn1.5 Cu1.5 Ga(OH)8 ]+
(CO
)
3
2
2

0.097

0.056

0.15

E

2− · mH O[Zn3 Ga(OH)8 ]+2
2 [Cu(OH)4 ]

0.30

0.13

0.43

F

[Zn1.5 Cu1.5 Ga(OH) ] [Cu(OH)4 ]

0.12

0.017

0.14

B

a

Photocatalyst

+
8 2

2−

· mH2 O

The catalyst amount was 100 mg. Values in the parentheses are experimental errors for evaluation.
bPreheated at 423 K for 1 h under vacuum.

20.5 CO2 PHOTOREDUCTION WITH HYDROGEN USING LAYERED DOUBLE HYDROXIDES599

SCHEME 20.1 Proposed photocatalytic catalytic cycle of CO2 reduction to methanol or CO over Zn–Cu–Ga LDH
catalysts.

+
2− · mH O
SCHEME 20.2 Energy levels at the band gap for ZnO, [Zn1.5 Cu1.5 Ga(OH)8 ]+2 , and [Zn3 Ga(OH)8 ]22 (CO3 )(CO3 )2− · mH2 O versus the reaction potentials for each step of CO2 reduction to methanol/CO and H2 oxidation to
protons.

were greater than that for ZnO (3.2 eV), enabling the CO2 photoreduction at relatively higher
potentials at −0.32 V to form methanol and to form CO at −0.11 V by supplying photoexcitedelectrons in conduction band to CO2. In Scheme 20.2, the center of the conduction band
minimum energy and the valence band top energy was assumed to be constant, but theseenergy levels should be exactly determined later. On the other hand, photogenerated holes
at the valence band were neutralized with electrons formed in the H2 oxidation (0 V). The
selectivity to produce methanol versus CO formation increased by adding Cu to
2− · mH O (Table 20.4). The band gap became narrower with the addi[Zn3 Ga(OH)8 ]+22 (CO3 )tion of Cu; however, the potential for the methanol formation step (CO2 + 6H+ + 6e− →CH3OH + H2O) [24] was 0.21 V higher than that for the CO formation step (Scheme 20.2).
Thus, the selectivity of CO2 photoreduction was not controlled by energetics of reaction steps
in the viewpoint of electron flow [41–45], but the adsorption of CO2 and the stability of the
reduced species is critical [1].
The durability of LDH catalysts and the verification of photocatalytic reaction mechanism
2− · mH O photocatalyst waswere also reported. When the [Zn1.5 Cu1.5 Ga(OH)8 ]+2
2 [Cu(OH)4 ]600

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELSFIGURE 20.6 Time course of photoreactions in CO2 (2.3 kPa) + H2 (21.7 kPa) using 100 mg of [Zn1.5 Cu1.5 Ga(OH)8 ]+2(CO3 )2− · mH2 O for 20 h. Other reaction conditions and legends are the same as those for Figure 20.4.

FIGURE 20.7 Time course of photocatalytic reaction in CO (2.3 kPa) + H2 (21.7 kPa). 0.10 g of the LDH photocata2− · mH O (a1, a2) and [Zn Cu Ga(OH) ]+ [Cu(OH) ]2− · mH O (b1, b2) was
lysts [Zn1.5 Cu1.5 Ga(OH)8 ]+
1.5
1.5
8 2
4
2
2
2 (CO3 )charged. The reactor was at 313 K without Xe arc lamp irradiation (a1, b1) while it was at 313 K under UV-visible
light from a Xe arc lamp (a2, b2). CO2 (!; diamond), H2O (▲; triangle), CH3OH (■ÆÊÃµÕ>Ài®]Ê>`Ê "Ê­UÆÊVÀVi®°tested for 20 h illuminated with UV-visible light, methanol formation continued and theselectivity was 76–84 mol% (Figure 20.6). Throughout the test, the only products that wereidentified were methanol, CO, and water.
+2−
In the control tests of CO2 reduction at 313 K, [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O+
2−produced no products (Figure 20.7a1) whereas [Zn1.5 Cu1.5 Ga(OH)8 ]2 [Cu(OH)4 ] · mH2 O
produced trace amounts of methanol and CO at the formation rates of 0.023 and−1 −1
0.0086 µmol h−1 g−1gcat
cat (Figure 20.7b1). Then, the methanol and CO formation (0.25 µmol h+2−
in total) using [Zn1.5 Cu1.5 Ga(OH)8 ]2 (CO3 ) · mH2 O under UV-visible irradiation(Figure 20.7a2) were proved to proceed via a band-gap excitation mechanism (Scheme 20.2).
2− · mH O
The methanol and CO formation rates using [Zn1.5 Cu1.5 Ga(OH)8 ]+2
2 [Cu(OH)4 ]without light were only 4.7% and 12%, respectively, but not zero compared to those underUV-visible light.

REFERENCES

601

20.6 CONCLUSIONSThe photoreduction of CO2 using semiconductors is one of the new methods to reduce theCO2 concentration in the atmosphere. If the photoreduction proceeds to form fuels (methanol,
formic acid, or methane) by sunlight, this option will solve the problem of sustainable energy
shortage simultaneously. The photoreduction of CO2 in water was reported typically using
−1 −1
−1 −1
TiO2 at rates less than 1 µmol h gcat. The rates were improved to 1 − 10 µmol h gcat bydispersing Ti sites on microporous/mesoporous molecular sieves, by the addition of promoters, and under supercritical reaction conditions.
When we assume photocatalytic water splitting to form hydrogen, this is an alternative
option to photoreduce CO2 with hydrogen. Using TiO2, ZrO2, MgO, and Ga2O3, the photoreduction rates to CO were ∼ 1 µmol h−1 g−1cat. As a new category of semiconductor photocatalysts, layered double hydroxides (LDHs) consisting of Zn, Cu, Ga, or Al photoreduced
CO2 with H2 to methanol at rates of 0.05–0.5 µmol h−1 g−1
and to CO at rates ofcat
0.07–0.18 µmol h−1 g−1
.
The
rates
are
not
sufficient
and
less
than
the
photoreduction of CO2
cat
in water using TiO2-based photocatalysts, but the combination of divalent and trivalentelements for semiconductor LDHs will improve the efficiency. The major product methanol
can be directly used as fuel, and the blank reaction test (no product) of CO2 in H2 at 313 K
2− · mH O demonstrated actual photocatalytic
in the dark using [Zn1.5 Cu1.5 Ga(OH)8 ]+22 (CO3 )formation of methanol and CO based on band-gap excitation mechanism [1,33].

Acknowledgments
The authors are thankful for financial support from a Grant-in-Aid for Scientific Research C (2255 0117, YI)
from MEXT (2010–2012, YI), from the Iwatani Naoji Foundation (2011–2012, YI), from Asahi Glass Foundation(2009–2010, YI), and from the Promotion Section of Technology Innovation, Daikin Industries (2008–2009, YI). NA
thanks Visiting Graduate Students Scholarship from Japanese Government (2008–2012).

References
[1]
[2]
[3]
[4]

N. Ahmed, Y. Shibata, T. Taniguchi, Y. Izumi, J. Catal. 279 (2011) 123–135.
J. Tollefson, Nature 473 (2011) 134–135.
A.J. Morris, G.J. Meyer, E. Fujita, Acc. Chem. Resear. 42 (2009) 1983–1994.
V.P. Indrakanti, J.D. Kubicki, H.H. Schobert, Energy Environ. Sci. 2 (2009) 745–758; Y. Izumi, Coor. Chem. Rev.
257 (2013) 171–186.
[5] N.S. Lewis, D.G. Nocera, Proc. Natl. Acad. Sci. USA 103 (2006) 15729–15735.
[6] R.D. Richardson, E.J. Holland, B.K. Carpenter, Nat. Chem. 3 (2011) 301–303.
[7] A.M. Khenkin, I. Efremenko, L. Weiner, J.M.L. Martin, R. Neumann, Chem. Eur. J. 16 (2010) 1356–1364.
[8] S.C. Roy, O.K. Varghese, M. Paulose, C.A. Grimes, Nano 4 (2010) 1259–1278.
[9] I.A. Shkrob, T.W. Marin, H. He, P. Zapol, in: J. Phys. Chem. C 116 (2012) 9450–9460.
[10] C.C. Lo, C.H. Hung, C.S. Yuan, J.F. Wu, Sol. Energy Mater. Sol. Cells 91 (2007) 1765–1774.
[11] T. Yui, A. Kan, C. Saitoh, K. Koike, T. Ibusuki, O. Ishitani, Appl. Mater. Interfaces 3 (2011) 2594–2600.
[12] S. Kaneco, H. Kurimoto, K. Ohta, T. Mizuno, A. Saji, J. Photochem. Photobiol. A 109 (1997) 59–63.
[13] S. Kanoco, H. Kurimoto, Y. Shimizu, K. Ohta, T. Mizuno, Energy 24 (1999) 21–30.
[14] T. Inoue, A. Fujishima, S. Konishi, K. Honda, Nature 277 (1979) 637–638.
[15] C.C. Yang, Y.H. Yu, B. van der Linden, J.C.S. Wu, G. Mul, J. Am. Chem. Soc. 132 (2010) 8398–8406.
[16] M. Anpo, H. Yamashita, Y. Ichihashi, Y. Fujii, M. Honda, J. Phys. Chem. B 101 (1997) 2632–2636.
[17] M. Anpo, H. Yamashita, K. Ikeue, Y. Fujii, S.G. Zhang, Y. Ichihashi, et al., Catal. Today 44 (1998) 327–332.

602

20. PHOTOCATALYTIC CONVERSION OF CARBON DIOXIDE INTO FUELS[18] J.S. Hwang, J.S. Chang, S.E. Park, K. Ikeue, M. Anpo, Topics Catal. 35 (2005) 311–319.
[19] K. Kočí, L. Obalová, L. Matějová, D. Plachá, Z. Lacný, J. Jirkovský, O. Šolcová, Appl. Catal. B 89 (2009)494–502.
[20] K. Kočí, K. Matějů, L. Obalová, S. Krejčíková, Z. Lacný, D. Plachá, L. Čapek, A. Hospodková, O. Šolcová, Appl.
Catal. B 96 (2010) 239–244.
[21] S.S. Tan, L. Zou, E. Hu, Catal. Today 115 (2006) 269–273.
[22] K. Ikeue, H. Yamashita, M. Anpo, T. Takewaki, J. Phys. Chem. B 105 (2001) 8350–8355.
[23] K. Ikeue, S. Nozaki, M. Ogawa, M. Anpo, Catal. Today 74 (2002) 241–248.
[24] C. Wang, R.L. Thompson, J. Baltrus, C. Matranga, J. Phys. Chem. Lett. 1 (2010) 48–53.
[25] Y. Kohno, H. Hayashi, S. Takenaka, T. Tanaka, T. Funabiki, S. Yoshida, J. Photochem. Photobiol. A 12 (1999)
117–123.
[26] Y. Kohno, T. Tanaka, T. Funabiki, S. Yoshida, Chem. Commun. (1997) 841–842.
[27] K. Teramura, T. Tanaka, H. Ishikawa, Y. Kohno, T. Funabiki, J. Phys. Chem. B 108 (2004) 346–354.
[28] K. Teramura, H. Tsuneoka, T. Shishido, T. Tanaka, Chem. Phys. Lett. 467 (2008) 191–194.
[29] C.C. Yang, J. Vernimmen, V. Meynen, P. Cool, G. Mul, J. Catal. 284 (2011) 1–8.
[30] F. Cavani, F. Trifirò, A. Vaccari, Catal. Today 11 (1991) 173–301.
[31] Y. Izumi, F. Kiyotaki, T. Minato, Y. Seida, Anal. Chem. 74 (2002) 3819–3823.
[32] Y. Izumi, F. Kiyotaki, Y. Seida, J. Phys. Chem. B 106 (2002) 1518–1520.
[33] N. Ahmed, M. Morikawa, Y. Izumi, Catal. Today, 185 (2012) 263–269.
[34] Y. Sun, Y. Zhou, Z. Wang, X. Ye, Appl. Surf. Sci. 255 (2009) 6372–6377.
[35] A. Ennadi, A. Legrouri, A. De Roy, J.P. Besse, J. Solid State Chem. 152 (2000) 568–572.
[36] G.W. Brindley, S. Kikkawa, Clay. Clay. Miner. 28 (1980) 87–91.
[37] D. Yan, J. Lu, J. Ma, M. Wei, S. Qin, L. Chen, D.G. Evans, X. Duan, J. Meter. Chem. 20 (2010) 5016–5024.
[38] G. Leofanti, M. Padovan, M. Garilli, D. Carmello, A. Zecchina, G. Spoto, S. Bordiga, G. Turnes Palomino, C.
Lamberti, J. Catal. 189 (2000) 91–104.
[39] M. Morikawa, N. Ahmed, Y.Yoshida, Y. Izumi, submitted to J. Catal. on April 16, 2013.
[40] I.H. Tseng, J.C.S. Wu, H.Y. Chou, J. Catal. 221 (2004) 432–440.
[41] Y. Izumi, T. Itoi, S. Peng, K. Oka, Y. Shibata, J. Phys. Chem. C 113 (2009) 6706–6718.
[42] Y. Izumi, K. Konishi, H. Yoshitake, Bull. Chem. Soc. Jpn. 81 (2008) 1241–1249.
[43] Y. Izumi, Y. Shibata, Chem. Lett. 38 (2009) 912–913.
[44] Y. Izumi, K. Konishi, D. Obaid, T. Miyajima, H. Yoshitake, Anal. Chem. 79 (2007) 6933–6940.
[45] D. Masih, H. Yoshitake, Y. Izumi, Appl. Catal. A 325 (2007) 276–282.