Applied Catalysis B: Environmental 144 (2014) 561–569

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Applied Catalysis B: Environmental
journal homepage: www.elsevier.com/locate/apcatb

Photoconversion of carbon dioxide in zinc–copper–gallium layered
double hydroxides: The kinetics to hydrogen carbonate and further
to CO/methanol
Motoharu Morikawa a,1 , Naveed Ahmed b,1 , Yusuke Yoshida b , Yasuo Izumi b,∗ab

Department of Nanomaterial Science, Graduate School of Advanced Integration Science, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japan
Department of Chemistry, Graduate School of Science, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japana r t i c l e

i n f oArticle history:
Received 19 May 2013
Received in revised form 25 July 2013
Accepted 28 July 2013
Available online xxx
Keywords:
CO2
Photoreduction
Layered double hydroxide
XANES
FTIR

a b s t r a c t
Photocatalytic reaction mechanism for the conversion of CO2 into methanol and CO using layered doublehydroxides (LDHs) consisting of Zn, Cu, and Ga was investigated. X-ray absorption ﬁne structure wasapplied to determine the LDH site structures and to monitor the diffusion of photogenerated electrons to
CuII sites. Electron diffusion to Cu sites was an order of magnitude faster in the direction of the cationiclayers (580 ␮mol h−1 gcat −1 ) than in the perpendicular direction. According to Fourier-transform infrared
spectroscopy, CO2 was in equilibrium with hydrogen carbonate (1629 cm−1 for H13 CO3 ) for the reaction
with hydroxy group from the cationic layer or interlayer site and/or with interlayer water. The equilibrium reactions were faster for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (400–110 ␮mol h−1 gcat −1 ) than for[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O. Furthermore, the reductive decomposition of hydrogen carbonate
was suggested in H2 under UV–visible light, suggesting photocatalytic pathway to methanol/CO.
© 2013 Elsevier B.V. All rights reserved.

1. Introduction
Carbon dioxide is one of the major greenhouse gases, and methods to reduce its concentration in the atmosphere require urgent
attention [1–10]. It is advantageous to capture CO2 from the atmosphere or factories/power stations and convert it to fuel using asustainable source of energy such as sunlight. This option simultaneously solves the problems of global warming and sustainableenergy shortage. Recently, many investigators have researched this
topic; however, more efﬁcient, inexpensive photocatalyst needs tobe developed to achieve its practical application [2–4,8–10].
In this regard, a study of selective CO2 photoconversion tomethanol (68 mol%) was reported using H2 and layered double hydroxides (LDHs) of [Zn1.5 Cu1.5 GaIII (OH)8 ]+ 2 (CO3 )2− ·mH2 O[11]. The inclusion of Cu ions in the cationic layers of thephotocatalysts improved the selectivity for methanol formation compared to CO formation. The substitution of interlayer
carbonate anions for [Cu(OH)4 ]2− anions further boosted themethanol selectivity to 88 mol% [12]. The conversion rate washigher by 3.1 times using [Zn1.5 Cu1.5 AlIII (OH)8 ]+ 2 (CO3 )2− ·mH2 O
than that using [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O, but the
Al analogue was more selective to CO formation (74%) [11].

∗ Corresponding author. Tel.: +81 43 290 3696; fax: +81 43 290 2783.
1

E-mail address: yizumi@faculty.chiba-u.jp (Y. Izumi).
These authors contributed equally to this paper.

0926-3373/$ – see front matter © 2013 Elsevier B.V. All rights reserved.
http://dx.doi.org/10.1016/j.apcatb.2013.07.065

Photoreduction of CO2 to CO with water was also reported
using LDH compounds, for example, [Ni3 In(OH)8 ]+ 2 (CO3 )2− ·mH2 O
and [Ni3 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O [13] or the combination of[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O and WO3 (or TiO2 ) separated
by proton-conducting ﬁlm [14].
Studies of CO2 photoreduction have been intensively published[8,15,16]. In the case of CO2 photoreduction using LDHs, it also
needs to be investigated that the CO2 conversion to fuels proceeded photocatalytically rather than thermally for the intrinsic
storage of solar energy as fuels [8]. It is also important to monitor the surface species derived from CO2 under the photocatalytic
reaction conditions [15,16]. Furthermore, the photocatalytic reaction mechanism has rarely been clariﬁed except for a few examples
using TiO2 , which include a mechanistic study of formic acid formation from supercritical CO2 based on electron spin resonance (ESR)[17,18], another study describing methane formation from gaseous
CO2 and H2 O based on ESR [19], and an unexpected result from an
ESR- and density functional theory (DFT) calculations-based investigation describing methane formation via C2 -glyoxal (OHC CHO)
[20].
In this paper, an investigation of the site structure for a LDH[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O as semiconductor photocatalyst used in the conversion of CO2 to hydrogen carbonate underlight irradiation was performed by extended X-ray absorption ﬁne
structure (EXAFS). The kinetics of the conversion to hydrogen carbonate, initiated by charge separation in the LDH compounds underUV–visible light, was investigated in detail by X-ray absorption

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M. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

near-edge structure (XANES) and Fourier-transform infrared (FTIR)
spectroscopy [21,22]. The relevance to the photocatalytic conversion of CO2 into methanol is discussed.

2. Experimental
The
LDH
compoundwith
theformula[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O was synthesized from anaqueous solution of metal nitrates and sodium carbonate ata controlled pH (pH 8). Details of the synthetic procedure
were described in Ref. [11]. LDH compounds with the formula[Zn3–x Cux Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (x = 0, 1.5) were synthesized from an aqueous solution of metal nitrates and ammonium
tetrachlorocuprate dihydrate at the same pH (pH 8). The details of
this procedure were described in Ref. [12].
The obtained precipitates were ﬁltered using a membrane ﬁlter,
washed with deionized water, and dried in ambient air at 290 K for
5 days. The purity of the deionized water used in the experiments
was <0.06 ␮S cm−1 . No C-containing chemicals were used except
for Na2 CO3 for these LDH syntheses. For the latter LDH compounds,
all synthetic steps up to the ﬁltration and washing (water) of the
LDHs were performed in an argon atmosphere.
The photocatalysts were evacuated (10−6 Pa) at 290 K for 2 hprior to kinetic tests. Photocatalytic CO2 reduction tests and control tests in the dark were performed in a closed circulating system
under CO2 at 0.20–2.1 kPa and H2 at 21.7 kPa using a quartz cell
illuminated by a 500-W xenon arc lamp (Ushio, Model UI-502Q;
42 mW cm−2 ) [23]. The distance between the lamp exit window
and the bottom of the cell was 20 mm. The products were analyzed
using an inline gas chromatograph equipped with a thermal conductivity detector (GC-TCD; Shimadzu, Model GC-8A). The details
are described in Refs. [11,12].
In situ XANES measurements were conducted in a transmission
mode at the Photon Factory on beamline 7C (KEK, Tsukuba) and at
SPring-8 on beamline 01B1 (Ako) [24] for Cu, Zn, or Ga in LDHs.
An LDH photocatalyst disk ( = 2 cm) was set in an air-tight quartz
batch cell equipped with polyethylene naphthalate windows (Q5116, Teijin) [25,26] under CO2 at 2.1 kPa and H2 at 21.7 kPa andirradiation by UV–visible light from a 500-W xenon arc lamp (UI502Q). The distance between the lamp exit window and the sample
disk was 20 mm. The light intensity was 42 mW cm−2 at the sample
position.
The EXAFS measurements were also conducted in a transmission mode for LDH photocatalyst powder pretreated in vacuum(10−6 Pa) at 290 K for 2 h. Data analysis software XDAP version 2.2.7(XAFS Services International, Woudenberg, the Netherlands) wasused. EXAFS spectra were also generated theoretically using ab initio multiple scattering calculation code FEFF version 8.4 [27] forcomparison to the experimental spectra.
The surface species over a 2-cm- self-supporting pressed
LDH disk (50 mg) placed under a CO2 atmosphere in a quartz
photoreaction cell [28] equipped with NaCl windows on bothsides was monitored by an FTIR instrument (JASCO, Model FT/IR4200) equipped with a mercury–cadmium–tellurium-M detector.
The LDH sample disk was evacuated (10−6 Pa) at 290 K for 2 hprior to taking spectral measurements. The spectrum acquired foreach freshly evacuated sample was set as a reference spectrum.
The energy resolution was set to 2 cm−1 . The in situ FTIR measurements were basically performed at 290 K. The sample wasirradiated by UV–visible light from a 500-W xenon arc lamp (Ushio,
Model SX-UID502XAM) via quartz ﬁber light guide (San-ei Electric Co., Model 5␾-2B-1000L; the length: 1075 mm). The distance
between the ﬁber light exit and sample disk was 90 mm. The lightintensity at the sample position was 88 mW cm−2 . During illumination, the temperature of sample reached as high as 310 K. The

Fig. 1. Structure of [Zn3–x Cux Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (A) and structure of[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O as determined by XAFS analyses (B). Interlayer
carbonates are not drawn in (A) for clarity.

isotope 13 CO2 (Cambridge Isotope Laboratories, Inc.; 13 C 99%,
18 O < 1%) was used.
3. Results
3.1. Copper site structures
The LDH compounds consist of cationic layers with anions
such as carbonates and structural water molecules, bothin interlayer space (Fig. 1A) [11,29]. For the compound[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O, Zn, Cu, and Ga cations
were in the MO6 octahedra linked at the edge to form a cationiclayer, and the cations distribute statistically [30] to exhibit asimilar XANES pattern at the Zn, Cu, and Ga K-edges [11].
The structure of the anions in the LDH compound[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O was studied by Cu K-edge
EXAFS [31]. The experimental Fourier transform (FT) (Fig. S1a)
was compared to theoretically generated data using FEFF version 8.4 (Fig. S1b–f). As a result, a (␮–O)3 Cu(OH)(H2 O)2 modelaccompanied by the coordination of two water molecules (Fig. 1B)[31] was most likely being covalently anchored to cationic layer:
Zn6 Ga2 (OH)13 {(␮–O)3 Cu(OH)(H2 O)2 }·mH2 O(Supplementarydata).
3.2. Rates of photogenerated electron trapping at the Cu sites
Next charge separation in the LDH compounds under CO2 + H2
and UV–visible light irradiation was monitored by Cu K-edge
XANES spectroscopy as the reduction of CuII sites (Fig. 2). Assuming
the wavelength of 354 nm, near the activity maximum as a function of excitation wavelength for the CO2 photoreduction using LDH[11], the photon number at the sample disk position was on theorder of 10 mmol-photons h−1 gcat −1 . A pre-edge peak at 8979 eV
because of a 1s–3d transition was used to evaluate the populationof CuII sites [32,33] in the LDH compounds (Fig. 2I). When the electrons separated from holes under UV–visible light diffused to the
CuII sites in the LDH, CuII sites were reduced to CuI . The pre-edge
peak does not appear for the CuI sites with d10 conﬁguration.
The reduction of inlayer sites of CuII to CuI was monitored for the
LDH compound [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O under CO2
and H2 and UV–visible light irradiation. The pre-edge peak intensity for 170 mg of incipient LDH (CuII : 0.55 mmol) monotonously
decreased by 15% within 50 min of irradiation (Fig. 2II-a andIII-a). The occurrence of reduction suggests that photogenerated electrons in the LDHs diffused and were trapped at Cu
sites. The e− trapping rate was 580 ␮mol h−1 gcat −1 (Table 1A).
This rate was lower than reported Cu2+ photoreduction rate
(2.2 mmol h−1 gTiO2 −1 ) in TiO2 suspension and sacriﬁcial reducing

<0.0051
0.002715g

2.4g

CO formation
Methanol formation

Under UV–vis
CO2
In

13

Interlayer Cu–OH recoverIn vac

580
36
7.1f
7.7f

7.2g
32g

21.7 kPa of H2 .
b
4.3 kPa of H2 .
c
From Ref. [12].
dBased on peaks at 1647 cm−1 .
ePreheated at 423 K.
f
Based on peaks at 1629 and 3673 cm−1 .
gBased on peak at 1629 and 3664 cm−1 .

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

HCO3 decomposition

In vacCO2

CO2 → HCO3

In

13

Elementary step rate (␮mol h−1 gcat −1 )Photocatalyst

563

Fig.
2. (I)
Normalized
Cu
K-edge
XANES
spectraof[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O preheated at 423 K (a) and fresh[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (b) [(Inset) Expanded view of the 1s–3d
pre-edge peak region] and (II) under CO2 (2.1 kPa) + H2 (21.7 kPa) and UV–visible
light irradiation (a: 6–51 min; b: 13–180 min) and subsequent change after the light
was turned off (a: 6 min; b: 8 min), and (III) time course of the 1s–3d peak intensity
for the two LDH photocatalysts under CO2 (2.1 kPa) + H2 (21.7 kPa) and UV–visible
light irradiation and subsequent change after the light was turned off. The amountof the LDH photocatalyst charged in the photoreaction cell was 170 mg.

30f(A)[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
16d[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
400f
+
2−
[Zn1.5 Cu1.5 Ga(OH)8 ] 2 [Cu(OH)4 ] ·mH2 O

1.3d
110f 11f

66f

180f

Interlayer
Cu–OH recoverInterlayer
Cu–OH
decrease
CO2 → HCO3 HCO3
decomposition

Interlayer Cu–OH decrease

Catalytic rate in CO2 &H2 (␮mol h−1 gcat −1 )<0.0051
0.021
0.0086
0.18e
0.13
0.070
1600
2200

0.31e
0.30
0.49

<0.0020<0.0020
0.023

COformation
CO
Methanolformation formation
CuI + h+ → CuII
CuII + e– → CuI

Methanolformation

At 313 K, in
dark

Catalytic rate in CO2 &H2 a (␮mol h−1 gcat −1 )Under UV–visc
In 13 CO2 andH2 , light off
In 13 CO2 andH2 a underUV–visIn H2 ,b underUV–visIn vac In vac, underUV–visIn 13 CO2
In H2 ,b underUV–visIn vac In vac, underUV–visIn 13 CO2

Elementary step rate (␮mol h−1 gcat −1 )Photocatalyst

Table 1Comparison of the rates of photocatalytic/thermal conversion of CO2 with those of the elementary step measured by FTIR and Cu K-edge XANES. (A) CO2 (2.1 kPa) and/or H2 and (B) CO2 (0.20 kPa) or CO2 (0.20 kPa) + H2 (21.7 kPa).

M. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

agent HCO2− Na+ (0.1 mol L−1 ) [34], but higher than the CuII photoreduction rate (4.1 ␮mol h−1 gcat −1 ) for a Cu–ZnO photocatalyst(0.5 wt% of Cu) in the presence of CO, O2 , and H2 gas [28].
In total, only 0.38% of the trapped electrons should beavailable for photocatalytic six – e− reduction to methanol(0.31 ␮mol h−1 gcat −1 ) and two – e− reduction to CO(0.18 ␮mol h−1 gcat −1 ) [12] (Table 1A), suggesting the rate oflateral e− diffusion to Cu sites within the cationic layers was
sufﬁcient (Fig. 1A) in contrast to that of subsequent e− transferfrom Cu sites to substrates/intermediates.
When the UV–visible light was turned off, the CuI sites oxidized
back to CuII (Fig. 2II-a and III-a) at a rate of 1600 ␮mol h−1 gcat −1 ,higher than the e− trap rate (580 ␮mol h−1 gcat −1 ) at the Cu sites
under light (Table 1A).
The 1s–3d pre-edge peak intensity for 170 mg of incipient [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (CuII : 0.17 mmol) also
monotonously decreased by 11% within 180 min of irradiation(Fig. 2II-b and III-b). In contrast to that of the LDHs consisting of inlayer Cu sites, the reduction rate of CuII (Fig. 2I-b) to
CuI at interlayer Cu sites (Fig. 1B) was 36 ␮mol h−1 gcat −1 for[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (Table 1A). The rate was essentially the slower e− diffusion rate in the perpendicular direction tointerlayer Cu sites later than the faster e− diffusion in the cationiclayers. Even though the interlayer Cu sites were covalently bonded
to O atoms in the cationic layers (Fig. 1B), the e− diffusion rate in
the perpendicular direction (36 ␮mol h−1 gcat −1 ) was smaller thanthat in the parallel, inlayer direction (580 ␮mol h−1 gcat −1 ). In total,
5.7% of the trapped electrons were used for photocatalysis to form

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M. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

methanol (0.30 ␮mol h−1 gcat −1 ) and CO (0.13 ␮mol h−1 gcat −1 ) [12](Table 1A). When the UV–visible light was turned off, the CuI
sites quickly oxidized back to CuII (Fig. 2II-b and III-b) at a rateof 2200 ␮mol h−1 gcat −1 .
The
Zn
and
Ga
K-edge
XANES
spectrafor[Zn3–x Cux Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O were already reported inRef. [11]. The photogenerated electron trap was only observed in
Cu K-edge in situ XANES (Fig. 2) due to the reduction of CuII sites
to CuI . However, partially reduced state from ZnII or GaIII was not
successfully detected in the Zn or Ga K-edge in situ XANES spectra.

3.3. Monitoring carbonate, hydrogen carbonate, and hydroxyspecies for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
Next, the interactions of [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
with CO2 /13 CO2 (2.1–0.20 kPa) were monitored by FTIR at 290 K.
Under CO2 at 2.1 kPa for 3–62 min, two peaks (1655 and1570 cm−1 ) gradually increased, while a sharp peak at 3671 cm−1
gradually weakened (Fig. 3A). A weak shoulder peak was also
observed at 1511 cm−1 . Under 13 CO2 at 2.1 kPa for 3–62 min,three overlapping peaks (1629, 1530, and 1474 cm−1 ) gradually
increased and a sharp peak at 3673 cm−1 gradually weakened(Figs. 3B and 4B).
If we consider the isotope wavenumber shift based on the harmonic equation [35], the peaks at 1655, 1570, and 1511 cm−1 under
CO2 (Fig. 3A) shift to 1618, 1535, and 1477 cm−1 , which is consistent with the values observed in 13 CO2 (1629, 1530, and 1474 cm−1 ;Fig. 3B).
For CuI /TiO2 in CO2 , the peaks at 1663 and 1554 cm−1 wereassigned to hydrogen carbonate (or carboxylate) and carbonate,respectively [15]. For ZnO in CO2 , the peaks at 1635, 1595–1580, and1522 cm−1 were assigned to hydrogen carbonate, bidentate carbonate, and polydentate carbonate, respectively [36]. For MgO, thepeaks at 1685 and 1526 cm−1 were assigned to hydrogen carbonate
and unidentate carbonate, respectively [37]. Systematic vibration
mode assignment of hydrogen carbonate, bidentate carbonate,
and polydentate carbonate has been well summarized pictorially[38]. Thus, the peak at 1655 cm−1 and the two peaks at 1570
and 1511 cm−1 observed for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O in
CO2 (Fig. 3A) were assigned to antisymmetric vibrations of CO3[as (CO3 )] of hydrogen carbonate, as (CO3 ) of bidentate carbonate,and as (CO3 ) of polydentate carbonate, respectively.
The peak intensity ratio at 1511 cm−1 in CO2 was inconsistent
with the corresponding one at 1474 cm−1 in 13 CO2 . This wasdue to impurity carbonates included in the fresh LDH sample of[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O. On closer inspection of the asprepared, 2-h evacuated [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
sample, the carbonate peak intensity at 1513 cm−1
was 8.0% of that for the same molar amount of the[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O disk, indicating that the exact
formula was [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− 0.92 (CO3 )2− 0.08 ·mH2 O.
Note that the interlayer anions [Cu(OH)4 ]2− actually existed suchas (␮–O)3 Cu(OH)(H2 O)2 (Fig. 1B) but were denoted as the species
present before dehydration to avoid having a highly complicatedcomposition formula. The impurity should originate from exposure
to atmospheric CO2 (remaining after the 2-h evacuation) during
the drying sequence of the LDH synthesis at 290 K.
The amount of impurity carbonate in the LDH sample (4.0␮molCO3 per 50 mgLDH ) was not negligible vs. the photocatalytic rate on the order of 1–102 ␮mol h−1 starting from CO2 + H2
(Table 1). The impurity carbonates interfered with quantitative
measurement of carbonate peaks in CO2 . In contrast, new carbonate peaks consisting of 13 C could be clearly detected at 1530 and
1474 cm−1 in 13 CO2 . The change of impurity carbonate during pretreatment in vacuum (290 K) affected the FTIR data in CO2 , but theFig. 3. FTIR spectra for LDH compound [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O under
CO2 at 2.1 kPa (A) and under 13 CO2 at 2.1 kPa (B and C) for 3 min (a), 17 min (b),32 min (c), and 62 min (d), evacuated for 3 min (e) and 12 min (f), irradiated by
UV–visible light for 3 min (g) and 12 min (h), and under H2 at 4.3 kPa irradiatedby UV–visible light for 3 min (g ) and 12 min (h ). FTIR spectra for LDH compound[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O under 13 CO2 at 2.1 kPa (D) for 62 min (d). The
amount of LDH was 50 mg.

exchange between 12 CO3 and 13 CO2 seems slower than the reactions of 13 CO2 with surface. Fortunately in view of catalysis, the 8.0%
of carbonate was quite inert in later steps than the pretreatment invacuum and negligibly affected the photoreduction of CO2 .
The O atom source to form carbonate under CO2 is unclear. The
O atom formed by partial dehydration of hydroxy groups on the
cationic layer during pretreatment (290 K, vacuum) is one potential
source. Instead, O atom of (␮–O)3 Cu(OH)(H2 O)2 species could reactwith CO2 to form carbonates.
To evaluate the molar amount of hydrogen carbonate species,
CO2 uptake was volumetrically monitored for the LDH sample
after the 2-h evacuation [39]. The CO2 uptake after 1 h (Fig. 4B)
was equalized to the total peak area of hydrogen carbonate andM. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

565

Fig. 4. The amount of 13 CO2 uptake at 290 K and the increase/decrease of surface hydrogen carbonate (1629–1647 cm−1 ; , square), carbonate (1530, 1474 cm−1 ; , diamond),
and isolated hydroxy (3673 cm−1 ; 䊉, circle) species in the FTIR spectra for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (50 mg; A and B) and [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (50 mg;
C). The 13 CO2 gas was introduced to the LDH samples at 0.20 kPa (A) or 2.1 kPa (B and C). Panel B was expanded from panel B. UV–visible light was irradiated in the absenceof H2 (A and B) and in H2 at 4.3 kPa (, open square, B ) at 78–112 min.

carbonate species identiﬁed by FTIR. We assumed that the molar
absorption coefﬁcients for hydrogen carbonate (1629 cm−1 ) andcarbonates (1530 and 1474 cm−1 ) were identical. The result of thisevaluation of the LDH sample of [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O(0.52 mmolCO2 gLDH −1 for 12-h sorption) was in accord with
reported CO2 uptake values (0.5–1.4 mmol gLDH −1 ) [40,41] and theamount of interlayer Cu sites (0.99 mmol gLDH −1 ) in the LDH.
In contrast to the stable, inert carbonate peaks, thehydrogen carbonate peak at 1629 cm−1 grew at a rate of
400 ␮mol h−1 gcat −1 under 2.1 kPa of 13 CO2 and decomposed
at a rate of 110 ␮mol h−1 gcat −1 under vacuum (Table 1A andFig. 4B). The peak under 13 CO2 for 62 min and under vacuumfor 14 min further decreased at the rate of 11 ␮mol h−1 gcat −1
during UV–visible irradiation (Figs. 3C-g, C-h and 5B and Table 1A).
Alternatively, when H2 at 4.3 kPa was added before UV–visible irradiation treatment, carbonate peaks still remained constant, while
the decrease in the hydrogen carbonate peak was accelerated to
66 ␮mol h−1 gcat −1 by hydrogen exposure (Figs. 3C-g , C-h and 4B
and Table 1A).
The isolated hydroxy peak at 3673 cm−1 in 13 CO2[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
(Fig.
3B)
was
for
not
observed
for[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O(Fig. 3D). Therefore, it was assigned to OH stretchingvibrations [(OH)] of (␮–O)3 Cu(OH)(H2 O)2species in[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (Fig. 1B). The peak decrease
under 13 CO2 is plotted in Fig. 4B and B (180 ␮mol h−1 gcat −1 ;Table 1A). In this plot, the decreased molar amount of hydroxygroup was assumed to be equal to the molar quantity of hydrogen carbonate formed over 62 min. The hydrogen carbonate
peak intensity (1629 cm−1 ) at 62 min corresponded to 2.8% ofthat of the (␮–O)3 Cu(OH)(H2 O)2 site after being transformed to(␮–O)3 Cu(O2 C–OH)(H2 O)2 .
As soon as the gaseous 13 CO2 was evacuated, the hydroxy peak
recovered at the rate of 30 ␮mol h−1 gcat −1 (Figs. 3B, 4B and B andTable 1A). The decrease/increase of hydroxy group did not coincide with the change observed for hydrogen carbonate, but it didcorrelate with 27–45% of the rate increase/decrease of hydrogencarbonate formation.
Then, under UV–visible light irradiation, the peak due to isolatedhydroxy groups further recovered at the rate of 7.7 ␮mol h−1 gcat −1(Fig. 3C-g and C-h and Table 1A), similar to the rate of hydrogen carbonate decomposition (11 ␮mol h−1 gcat −1 ) under this condition.
Alternatively, when H2 at 4.3 kPa was added before the UV–visible
irradiation, the recovery of the peak associated with isolatedhydroxy group did not change signiﬁcantly (7.1 ␮mol h−1 gcat −1 ;Fig. 3C-g and C-h and Table 1A), suggesting hydrogen carbonate did not totally decompose back to initial CO2 + surface hydroxygroup, but it was further photoreduced by hydrogen.
FTIR monitoring was also performed at a lower pressure of
13 CO (0.20 kPa). The rates of both hydrogen carbonate forma2
tion/decomposition and the decrease/recovery of Cu–OH species
slowed by an order of magnitude (Table 1B). However, the rates
of formation of Cu–OH species were again lower than those forhydrogen carbonate formation by a factor of 47–33% (Table 1B),
comparable to the ratios for the test run under 13 CO2 at 2.1 kPa(Table 1A).

3.4. Monitoring carbonate, hydrogen carbonate, and hydroxyspecies for [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
For [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O under 2.1 kPa of 13 CO2 ,the signals from interlayer carbonate interfered with thoseof the additional minor carbonate and/or hydrogen carbonate and the spectra were less precise compared to those for[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (Fig. 3D). A broad peak at
1647 cm−1 increased at the rate of 16 ␮mol h−1 gcat −1 under 13 CO2
and decreased at the rate of 1.3 ␮mol h−1 gcat −1 under vacuum(Fig. 4C). This peak can be assigned to hydrogen carbonate andquantiﬁed using the same criteria as that in Fig. 3B. The molar
amount of hydrogen carbonate under 13 CO2 exposure after 62 mincorresponded to 0.12% of that of carbonate anion in the LDH

566

M. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

Fig. 5. Time course of the photocatalytic reaction under CO2 (2.3 kPa) + H2
(21.7 kPa). The amount of LDH photocatalysts [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
was 100 mg. The reactor was set at 313 K in the dark (a) or below 313 K underUV–visible light from a Xe arc lamp (b). CO2 (, diamond), H2 O (, triangle), CH3 OH(, square), and CO (䊉, circle).

compound. In the (OH) region, no peak due to isolated hydroxygroup appeared (Fig. 3D, left).
3.5. CO2 conversion tests under light and in the dark at 313 K
Hydrogen carbonate observed in the FTIR spectra under
CO2 can be reduced by the electrons trapped at the Cu site,
as monitored by XANES. CO was formed (0.13 ␮mol h−1 gcat −1 )using [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O [12] as resulting fromthe reaction of a reduced hydrogen carbonate with a proton
and an e− to form CO and water at the interlayer sites including the Cu sites [11]. On the other hand, methanol formation(0.30 ␮mol h−1 gcat −1 ) from a reduced hydrogen carbonate stillrequires multiple reduction steps. To obtain insight into the steps,
kinetic tests under UV–visible light, and control tests in the dark
at 313 K (maximum accessible temperature at the reactor under
the test conditions [11]) were performed under CO2 and H2 using[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O.
No products above the detection limit of GC–TCD weredetected over a 5-h period in the dark at 313 K using[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (Table 1A and Fig. 5a). WhenUV–visible light was applied, the reactor was below 313 K, and amajor amount of methanol and minor amount of CO were formed(Fig. 5b). Thus, hydrogen carbonate, which formed bound to inlayer
Cu ions (Fig. 1A), should be reduced by multiple e− transfer via theforbidden direct e− transition from O 2p to M 3d (M = Zn, Cu, or Ga)
in LDHs [11,12] and H+ transfer.
By
contrast,
in
the
dark
at
313 K
using
[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O, CO was formed at the rateof 21 nmol h−1 gcat −1 , but no methanol was found (Table 1A).
The rate was 16% of that observed under UV–visible light [12].
Using [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O, trace amounts ofmethanol and CO formed at the rates of 23 and 8.6 nmol h−1 gcat −1 ,respectively (Table 1A). The rates were 4.7% and 12%, respectively,of those observed under UV–visible light [12].
In addition, a photocatalytic test under CO2 at 0.20 kPa + H2 at
21.7 kPa was performed for comparison to the reported results
under CO2 at 2.1 kPa + H2 at 21.7 kPa (Table 1A) [12] using[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O. The methanol and CO formation rates were only 0.90% and less than 3.9%, respectively, of thoseunder CO2 at 2.1 kPa (Table 1B).
4. Discussion
4.1. Photogenerated electron trap
The Cu K-edge EXAFS analysis (Supplementary data) indicated that the structure of the interlayer Cu sites for thecompound[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
transLDH
formed into (␮–O)3 Cu(OH)(H2 O)2 by reaction with hydroxygroups in the cationic layer (Fig. 1B). As determined by
Cu K-edge in situ XANES monitoring, electron trapping(580 ␮mol h−1 gcat −1 ) at the Cu sites was 18–260 times faster thanthat observed for catalytic CO2 conversion into methanol + CO using[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (2.2 ␮mol-e− h−1 gcat −1 ) and[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (2.1 ␮mol-e− h−1 gcat −1 ) underlight (Table 1A). Furthermore, e− trapping was 16 times faster for
inlayer Cu sites in [Zn1.5 Cu1.5 Ga(OH)8 ]+ than for interlayer Cu sites
as anchored (␮–O)3 Cu(OH)(H2 O)2 (36 ␮mol h−1 gcat −1 ; Table 1A).
In summary, photogenerated e− trapping at the Cu sites of
LDH samples in this study was fast enough compared to photocatalysis to afford CO2 conversion. Speciﬁcally, lateral e− diffusionin the cationic sheets (Fig. 1A) was faster than diffusion in the
direction perpendicular to the attached (␮–O)3 Cu(OH)(H2 O)2 sites(Fig. 1B).
4.2. Dynamic equilibrium of CO2 with hydrogen carbonate
In the FTIR measurements of [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
under 13 CO2 , carbonate as (CO3 ) peaks were observed at 1530
and 1474 cm−1 (Fig. 3B). The peaks were stable in vacuum or
an atmosphere of H2 and/or under irradiation with UV–visible
light, suggesting bidentate/polydentate carbonate species were
not incorporated during catalytic CO2 reduction under UV–visible
light.
By contrast, for the LDHs, the hydrogen carbonate as (CO3 ) peak
was weaker (1629 cm−1 under 13 CO2 , Fig. 3B) and the species werein equilibrium with CO2 and hydroxy groups. The rates of increase
and decrease of the hydrogen carbonate peak under 13 CO2 and vacuum (Fig. 3B) were 400 and 110 ␮mol h−1 gcat −1 , 930- and 260-foldhigher, respectively, than those of catalytic CO2 conversion andslightly greater than the e− trapping rates at the Cu sites (Table 1A).
In this case, CO2 is in equilibrium with hydroxy groups in the cation
layers or interlayer (␮–O)3 Cu(OH)(H2 O)2 species.
CO2 + HO(cationic layer)  HO–CO2 (cationic layer)
CO2 + (␮–O)3 Cu(OH)(H2 O)2  (␮–O)3 Cu(O2 C–OH)(H2 O)2The equilibrium of CO2 with Zn–OH sites to form hydrogen carbonate was calculated on the basis of Hartree–Fock and DFT to be
exothermic by 33 kJ mol−1 [42]. As in Ref. [42], hydrogen carbonate
species in this study may also form hydrogen bonds with neighboring hydroxy groups. The equilibrium of CO2 (and water) with
Cu2 O(1 1 1) surface species, as observed by DFT calculations, wasalso reported. Hydrogen carbonate was the major surface species
and was suggested as the intermediate in the reduction of CO2[43].
All the LDH compounds in this study were precipitated at pH 8
during the synthesis and have a basic nature. Near the basic surface
in the interlayer space of the LDHs [44], the following equilibriumis also plausible.
CO2 + H2 O(interlayer)  HO–CO2− (interlayer) + H+ (interlayer)
The sharp peaks that appeared at 3671–3673 cm−1 in CO2(Fig. 3A) or 13 CO2 (Panel B) were assigned to OH for interlayer hydroxy group at (␮–O)3 Cu(OH)(H2 O)2 sites (Fig. 1B). These
were separated from the OH peaks of interlayer water molecules(3348 cm−1 ) and those of hydroxy groups in the cationic layers(3468 cm−1 ) [45]. The rates of decrease and increase of the hydroxy
peak were not equivalent to the rates of increase and decrease of
the hydrogen carbonate peak in 13 CO2 ; rather, the latter rates weregreater by 2.2–3.7-fold. The situation was similar when a lower
pressure (0.20 kPa) of 13 CO2 was used (2.1–3.0-fold).
In summary, the hydroxy groups of the cationic layers and/orinterlayer water reacted in the early stage of 13 CO2 exposure, andM. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569

567

4.3. Photocatalytic reaction mechanism from CO2

Scheme 1. Proposed photocatalytic cycle of CO2 reduction to methanol or CO using
LDH catalysts consisting of Zn, Ga, and interlayer Cu sites.

the slower diffusion of CO2 into the interlayer space later led to
the formation of (␮–O)3 Cu(O2 C–OH)(H2 O)2 . If the isolated hydroxy
peak was because of an impurity (e.g., Cu(OH)2 ) powder thatcoexisted with [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O) [12], this timedelay cannot be rationalized.
Hydrogen carbonate formation was difﬁcult to monitor for[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O under CO2 or 13 CO2 owing to
the strong IR absorption of the enormous amount of interlayer
carbonates. Even so, the rates of formation and decomposition (16–1.3 ␮mol h−1 gcat −1 ; Table 1A) were signiﬁcantly lower(4.0–1.2% of those for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O). The
smaller interlayer spacing for [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
compared to that for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (0.753 nm[11] vs. 0.792 nm [12]) may be the reason.
The hydrogen carbonate peaks at 1629 cm−1
for
[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O increased under 13 CO2 at
2.1 kPa and quickly decreased under vacuum or UV–visible light(Fig. 4B and B ). The peak decrease under UV–visible light was
further accelerated 6.0-fold in the presence of H2 (Fig. 4B ). In
clear contrast, the recovery of the isolated Cu–OH peak was notaffected by the presence of H2 (7.7 and 7.1 ␮mol h−1 gcat −1 , Fig. 3C
and Table 1A). This difference suggested that hydrogen carbonate
dominantly decomposed back to initial CO2 + surface hydroxy
under vacuum but was further reduced under H2 and UV–visible
light irradiation (Scheme 1).
The trend in the FTIR measurements under CO2 and vacuum
was similar to that under 13 CO2 at 0.20 kPa (Fig. 4A). As theamount of hydrogen carbonate in equilibrium decreased (from28 to 2.3 ␮mol gcat −1 ), the rates of formation/decomposition werereduced by an order of magnitude by decreasing the 13 CO2 pressure
from 2.1 to 0.20 kPa (Table 1A and B).
For[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O
under
CO2
at
0.20 kPa + H2 at 21.7 kPa and UV–visible light, the photocatalytic rates of methanol and CO formation decreased by two ordersof magnitude (Table 1A and B). Furthermore, when H2 at 21.7 kPa
was introduced to the LDH sample treated with CO2 for 62 minfollowed by a 14-min evacuation and the sample was irradiated,methanol formation was conﬁrmed by GC-TCD in addition to CO2
and CO formation [11].
In summary, the rapid decomposition of the hydrogen carbonate
species under vacuum and UV–visible light and its further promotion by the addition of H2 as well as the critical catalytic ratedependence on the equilibrium concentration of the hydrogen carbonate species strongly suggested that hydrogen carbonate was theintermediate in the photocatalytic formation of methanol and CO
from CO2 using the LDH compounds.

On the basis of our study by Cu K-edge in situ XANES and in situ
FTIR, we postulate the following initial reaction steps for the CO2
photoreduction at the LDH interlayer sites including the interlayer
Cu sites (Scheme 1): (Step 1) the reaction of CO2 with hydroxy groupbound to interlayer CuII to form hydrogen carbonate, (Step 2) the
forbidden direct e− transition from O 2p to M 3d (M = Zn, Ga) [11],(Step 3) the diffusion of electrons and holes in the LDHs to CuII sites
and H atom, respectively, at the surface, and (Step 4) the transfer of trapped electrons from CuI to hydrogen carbonate forming
anionic hydrogen carbonate. Steps 1 and 2 occurred at a sufﬁcient
rate (Table 1), and Step 4 (or possibly an event later than Step 4)
should be rate-determining.
The direct e− transition of LDHs in this study was considered asforbidden, based on the ﬁt [11,12] of UV–visible absorption data
to the equation of Davis and Mott. This is forbidden due to thesymmetry of the Bloch functions of valence band and conduction
band. Even so, minor term in the perturbation Hamiltonian does
not diminish due to the symmetry, and weaker absorption wasobserved experimentally [46].
In this study, no C-containing intermediate/adsorbed species
was observed in the FTIR spectra except for carbonates andhydrogen carbonate (Fig. 3), and the bidentate and polydentate
carbonates were inert. For TiO2 , the photooxidation of carbonate
to CO3 – radicals was detected by ESR at 4.5 K, but the photocatalytic reaction mechanism was proposed to proceed via a formate
species [CO2 + H+ + 2e− → HCO2 − (surface)] followed by conversion
to methane [47]. The surface concentration of hydroxy groups
was 1.7–4.5 groups nm−1 over TiO2 vs. 1.7 × 105 groups nm−1
for [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O (speciﬁc surface area,
57 m2 g−1 ) [11]. Because of this difference, hydrogen carbonate was preferentially observed for the LDH compounds in thisstudy.
Although the e− diffusion rate to CuII sites (Step 3) was
16 times greater for [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O
than for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O as determined by
Cu K-edge in situ XANES (Fig. 2 and Table 1A), the molar
ratio of hydrogen carbonate species in 13 CO2 for 62 min
was 0.12% and 2.8%, respectively, to the anion {(CO3 )2−or [Cu(OH)4 ]2− } of LDHs. In addition to hydrogen carbonate population (23 times greater for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O)
effects, e– transfer at the interface (Step 4) may be more
effective between a Cu atom and hydrogen carbonate species(␮–O)3 Cu(O2 C–OH)(H2 O)2 formed at the interlayer Cu–OH sites(Fig. 1B) than between a Cu atom and the hydrogen carbonate
species formed at inlayer (M–)3 OH sites (M = Cu, Zn, or Ga; Fig. 1A).
By these effects, with [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O, a signiﬁcant improvement in the utilization efﬁciency of trapped electrons
at the Cu sites is observed (5.7% vs. 0.38% at inlayer (M–)3 OH sites).
Because of the compensating effects of Steps 3 and 4, the photocatalytic CO2 conversion rates into methanol and CO were similar for[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O [0.49 ␮mol(CO2 ) h−1 gcat −1 ]and[Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O[0.43 ␮mol(CO2 )h−1 gcat −1 ; Table 1A].
4.4. Veriﬁcation of photocatalysis from CO2 + H2 using LDHs
Using the LDH compound [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O,
no products were detected after 5 h at 313 K in the dark under CO2
and H2 (Fig. 5a). The Zn–Cu–Al LDH was reported to be stable as
high as at 573 K [48].
Unfortunately, the purity of [Zn3–x Cux Ga(OH)8 ]+ 2 [Cu(OH)4 ]2−·mH2 O(x = 0,
1.5)
was
lower
than
that
of
[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O in this study; minor
CuO and Cu(OH)2 were included for the LDH compounds

568

M. Morikawa et al. / Applied Catalysis B: Environmental 144 (2014) 561–569[Zn3–x Cux Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O (x = 0, 1.5) [12,31]. In thetime course at 313 K using them in the dark, minor CO/methanol
were formed later than the induction period of 1–2 h (Table 1A).
As one of the plausible reasons, reduced CuI sites [49,50] from the
CuII impurity during the induction period may form CO/methanol
similar to the process of conventional methanol synthesis[51–53].

Appendix A. Supplementary data
Supplementary data associated with this article can befound, in the online version, at http://dx.doi.org/10.1016/j.apcatb.
2013.07.065.

References
5. Conclusions
For the LDH [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O, photogenerated e– diffusion from the bulk to the interlayer (␮–O)3 Cu(OH)(H2 O)2sites was monitored by Cu K-edge in situ XANES and compared with that of [Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O. The rateof e– diffusion was 16 times faster for inlayer Cu sites of[Zn1.5 Cu1.5 Ga(OH)8 ]+ than for interlayer (␮–O)3 Cu(OH)(H2 O)2sites.
The reaction with CO2 was monitored by in situ FTIR. Hydrogencarbonate (1629 cm−1 ) was formed under 13 CO2 at the rate of
400 ␮mol h−1 gcat −1 , and it decomposed under vacuum at the rateof 110 ␮mol h−1 gcat −1 for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O.
The peak intensity change was synchronized with thedecrease/increase (180–30 ␮mol h−1 gcat −1 ) of an isolated
Cu–OH peak (3673 cm−1 ). The time delay was considered
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possibly due to differences in the interlayer spacing (0.792
vs. 0.753 nm).
The decomposition of hydrogen carbonate proceeded underUV–visible light, and for [Zn3 Ga(OH)8 ]+ 2 [Cu(OH)4 ]2− ·mH2 O, it wasenhanced 6-fold by the addition of H2 . The rate of recovery of
the Cu–OH peak changed only negligibly in the presence of H2
under UV–visible irradiation, suggesting further photoreductionof hydrogen carbonate by H2 . Furthermore, by decreasing theCO2 pressure, the rates of formation/decomposition of hydrogen carbonate and the rates of photocatalyzed methanol + CO
formation from CO2 + H2 decreased by one to two orders of magnitude, respectively. Thus, hydrogen carbonate was conﬁrmed
to be an intermediate in the photocatalysis of CO2 to methanol
and CO.
The
LDH
compoundwith
theformula[Zn1.5 Cu1.5 Ga(OH)8 ]+ 2 (CO3 )2− ·mH2 O was synthesized withadequate purity, and no products were detected in the control reaction under CO2 + H2 at 313 K in the dark. The actionof the LDH compounds was conﬁrmed to originate from initialcharge separation caused by photons in the photocatalysis underCO2 + H2 .

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Acknowledgments
The authors are grateful for the ﬁnancial supports received fromthe Grant-in-Aid for Scientiﬁc Research C (2255 0117) from Ministry of Education, Culture, Sports, Science, and Technology (MEXT,
2010–2012) and the Asahi Glass Foundation (2009–2010). The Xray absorption experiments were performed with the approvalof the Photon Factory Review Committee (Nos. 2011G033 and2009G552) and the grant of the Priority Program for DisasterAffected Quantum Beam Facilities (2011A1977, SPring-8 & KEK).
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569

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