Journal of Catalysis 279 (2011) 123–135

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Journal of Catalysis
journal homepage: www.elsevier.com/locate/jcat

Photocatalytic conversion of carbon dioxide into methanol using
zinc–copper–M(III) (M = aluminum, gallium) layered double hydroxides
Naveed Ahmed a, Yoshiyuki Shibata a, Tatsuo Taniguchi b, Yasuo Izumi a,⇑ab

Department of Chemistry, Graduate School of Science, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japan
Department of Applied Chemistry and Biotechnology, Graduate School of Engineering, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japana r t i c l e

i n f oa b s t r a c t

Article history:
Received 6 October 2010
Revised 7 January 2011
Accepted 9 January 2011
Available online 15 February 2011

Ordered layered double hydroxides (LDHs) consisting of zinc and/or copper hydroxides were synthesized
and combined with aluminum or gallium. These LDH compounds were then applied as photocatalysts toconvert gaseous CO2 (2.3 kPa) to methanol or CO under UV–visible light using hydrogen. ZnAAl LDH wasthe most active for CO2 photoreduction and the major product was CO formed at a rate of
1
620 nmol h g1cat , whereas methanol was the major product formed by the inclusion of Cu in the LDH
1photocatalysts, e.g., at a formation rate of 170 nmol h g1
cat using ZnACuAGa photocatalyst. The methanol selectivity improved by the inclusion of Cu from 5.9 to 26 mol% and 39 to 68 mol%, respectively, when
ZnAAl (the conversion 0.16–0.11%) and ZnAGa LDH catalysts were used (the conversion 0.02–0.03%).
Speciﬁc interaction of Cu sites with CO2 was spectroscopically suggested to enable coupling with protons
and photogenerated electrons to form methanol.
Ó 2011 Elsevier Inc. All rights reserved.

Keywords:
CO2
Photoreduction
MethanolLayered double hydroxide
Copper

1. Introduction
Carbon dioxide is one of the major greenhouse gases [1]. Several
methods for reducing CO2 concentration in the atmosphere havebeen 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 andtransferring it to the bottom of the sea in a supercritical state. It
would be advantageous to capture CO2 from the atmosphere andconvert it to fuels using sustainable energy such as sunlight, because this solves the problems of global warming and sustainableenergy shortage simultaneously.
The major obstacle to realizing this option is the thermodynamic limitation on the chemical conversion:1CO2 ðgÞ þ 2H2 OðgÞ ! CH3 OHðgÞ þ 1:5O2 ðgÞ; DGr ¼ þ689 kJ mol :Since an efﬁcient catalyst for this photoconversion to fuel is unknown [2–12], it is also important to investigate good catalysts
for CO2 conversion to methanol or other fuels using hydrogen as a
reductant, which is potentially obtained from water splitting:

2H2 OðgÞ ! 2H2 ðgÞ þ O2 ðgÞ;

1

DGr ¼ þ457:2 kJ mol :CO2 ðgÞ þ 3H2 ðgÞ ! CH3 OHðgÞ þ H2 OðgÞ;

1

DGr ¼ þ2:9 kJ mol :⇑ Corresponding author. Fax: +81 43 290 2783.
E-mail address: yizumi@faculty.chiba-u.jp (Y. Izumi).
0021-9517/$ - see front matter Ó 2011 Elsevier Inc. All rights reserved.
doi:10.1016/j.jcat.2011.01.004

The latter reaction, using hydrogen, does not essentially solve theproblems of global warming and sustainable energy shortage, because current hydrogen production processes require energy andproduce CO2. The hydrogen needs to be produced from water driven
by sunlight if the photoreduction in this paper is to be utilized in a
future environmentally benign society [13].
Catalysts for converting CO2 to methanol using hydrogen atroom temperature and utilizing UV–visible light were investigatedin this study. As described above, this option requires technologies
both to capture CO2 from the atmosphere and to convert it into
fuel. To satisfy the two requisites, layered double hydroxides(LDHs) were chosen with the expectation of (1) sorption capacity
for CO2 in the layered space and (2) tunable semiconductor properties (photocatalytically active) as a result of the choice of metal
cations. LDHs are materials based on the layered structure of brucite (Mg(OH)2), which has a hexagonal crystal structure, whereas
the MgO6 octahedra are linked at the edge to form sheets(Scheme 1A). One of the naturally occurring LDH compounds ishydrotalcite [14]. The charge of cationic sheets formulated asxþ½MII1x MIII
is compensated for by an intercalated anionx ðOHÞ2 

2


II
group, e.g., CO2
3 ; SO4 ; NO3 ; Cl ; or OH , and the M site can beIIIMg, Mn, Fe, Co, Ni, Cu, or Zn. The M site can be Al, Cr, Mn, Fe,
or Ga. x is normally variable within the range 0.17–0.33. The molar
amount of structural water intercalated between the cationic layers is about half of the total molar amount of metal cations(Scheme 1A) [15–19].
As for the ﬁrst requisite, the sorption capacity of LDHs
was 0:5—1:4 mmol g1
sorbent , which was comparable to that ofzeolites (0:95—3:2 mmol g1
sorbent ) [20,21] and greater than that of

124

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

Scheme 1. (A) Structure of Zn57xy Cux MIIIy ðOHÞ72  38H2 O cluster models: (i) x = y = 0; (ii) x = 0, y = 12, M = Al; (iii) x = 0, y = 1, M = Ga; (iv) x = 1, y = 0; and (v) x = 1, y = 12,
M = Al. (B) Hydrogen carbonate species formed from the reaction of CO2 with the surface hydroxy group of Zn57xy Cux MIIIy ðOHÞ72  38H2 O.

montmorillonites (0:4—0:5 mmol g1
sorbent ) [22,23]. The proportionof carbonate ions in ½Zn0:75 Ga0:25 ðOHÞ2 0:25þ ðCO3 Þ2
0:125  mH2 O is
1:1 mmol g1
LDH . The tests for CO2 sorption were reported for previously degassed LDH compounds in which the MII site was Mg, Co,
Ni, Cu, and Zn and the MIII site was Al and Ga [15,24–26]. In contrast to the fact that various negatively charged sheet layeredmaterials such as montmorillonites, smectites, and kaolinites havebeen studied in detail, the number of positively charged sheet layered materials that can be applied directly to carbonate anion exchange is limited and they are not well studied.
As for the second requisite, ZnII and GaIII were chosen as the cations of LDHs to form the semiconductor samples. Zinc and copper
have an afﬁnity for CO2 and they form carbonate and formate species in methanol synthesis, typically at 523–573 K [27]. In addition,
Zn and Cu are associated with the trivalent elements Ga, Al, or Cr in
the synthesis of industrial methanol [28]. In this context, LDH comxþ
2
pounds ½ZnII1xy CuIIy MIII
consisting of Zn and/or Cu
x ðOHÞ2 2=x ðCO3 ÞIIIIII
associated with Ga or Al were chosen to behave as semiconductors and potentially produce methanol via a photo process.
Photoreduction of CO2 using hydrogen as a reductant has beenreported to produce carbon monoxide using Rh/TiO2
1
1 1(5:1 lmol h g1
gcat ) [30,31],
cat ) [29], ZrO2 (0:51—0:70 lmol h
1 11
MgO (1:6 lmol h gcat ) [32], and b-Ga2O3 (0:76 lmol h g1cat )1[33], and to produce methane using TiO2 (0:30 lmol h g1
cat )[8,34]. In this study, new catalysts for producing methanol were
found for the ﬁrst time in the photoreaction of CO2 using hydrogenas a reductant. The synthesis, photocatalytic performance, andspectroscopic monitoring of the interaction of metal sites with
CO2, and bandgap energies vs each reaction step of CO2 reduction
and that of H2 oxidation to H+ were reported.

(Wako Pure Chemical, >99.9%), sodium carbonate (Wako Pure
Chemical, >99.8%), and sodium hydroxide (Kanto Chemical,>97.0%) were used as received. Deionized water (<1.0 lS cm1)
was used throughout the syntheses.
2
An LDH compound formulated as ½Zn3 GaIII ðOHÞ8 þ
2 ðCO3 Þ 
mH2 O was synthesized [35]. A mixed solution (20 ml) of 0.75 M
Zn(NO3)26H2O and 0.25 M GaðNO3 Þ3  nH2 O was dropped at a rateof 0.6 ml min1 into 100 ml of a 0.075 M Na2CO3 solution in a ﬂask
at 290 K with stirring at a rate of 900 rpm. The pH was adjusted to
8 by adding 1.0 M NaOH (20 ml total). After the addition wascomplete, the mixture was continuously stirred at the same rate
at 290 K for 2 h. The pH was maintained at 8 by adding 1.0 M NaOH
(0.5 ml total). Then the temperature of the mixture was raised to353 K and continuously stirred for an additional 22 h. The pH of thesolution did not change from 8 throughout the 22 h. The obtained
precipitates were ﬁltered using a polytetraﬂuoroethene-based
membrane ﬁlter (Omnipore JGWP04700, Millipore) with a poresize of 0.2 lm, washed with deionized water, and dried in ambient
air at 290 K for 5 days. Following the same procedure, an LDH comIII2
pound formulated as ½Zn3 Al ðOHÞ8 þ mH2 O was synthe2 ðCO3 Þ
sized using Al(NO3)39H2O (0.25 M) instead of GaðNO3 Þ3  nH2 O(0.25 M).
Cu atoms were substituted at the ZnII sites in the LDH com2
pounds formulated as ½Zn3 MIII ðOHÞ8 þ mH2 O (M = Ga or
2 ðCO3 ÞAl). At the beginning of the synthesis, a mixed solution (20 ml)
was prepared as 0.375 M Zn(NO3)26H2O, 0.375 M Cu(NO3)23H2O,
and 0.25 M Al(NO3)39H2O (or 0.25 M Ga(NO3)3nH2O) to set themolar ratio of ZnII, CuII, and MIII (Ga, Al) ions to 3:3:2. Henceforth,
the steps are essentially identical to those described above for the2
syntheses of ½Zn3 MIII ðOHÞ8 þ mH2 O.
2 ðCO3 Þ

2. Experimental

2.2. Characterization

2.1. Syntheses

Nitrogen adsorption isotherm measurements were performed
at 77 K within the pressure range 1.0–90 kPa in a vacuum system
connected to diffusion and rotary pumps (106 Pa) and equippedwith a capacitance manometer (Models CCMT-1000A and GM2001, ULVAC). The Brunauer–Emmett–Teller (BET) surface area
(SBET) was calculated on the basis of eight-point measurements

For the synthesis of LDH compounds, zinc nitrate hexahydrate(Wako Pure Chemical, >99%), copper nitrate trihydrate (Wako Pure
Chemical, >99%), gallium nitrate hydrate (Wako Pure Chemical,>99.9%; n = 7–9 of GaðNO3 Þ3  nH2 O), aluminum nitrate nonahydrate

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

between 10 and 46 kPa (P/P0 = 0.10–0.45) on the adsorption isotherm. The samples were evacuated at 383 K for 2 h before themeasurements.
X-ray diffraction (XRD) data were obtained using a MiniFlex diffractometer (Rigaku) at a Bragg angle of 2hB = 5–70° with a scanstep of 0.01° and a scan rate of 7 s per step for the sample powders.
The measurements were performed at 30 kV and 15 mA using Cu
Ka emission and a nickel ﬁlter. Scanning electron microscopy(SEM) measurements were performed using JEOL Model JSM6510A at the Chemical Analysis Center, Chiba University. The samples were mounted on a conducting carbon tape and coated withgold. The electron accelerating voltage was between 15 and25 kV and the magniﬁcation was between 20,000 and 50,000
times.
Optical spectroscopic measurements were performed using a
UV–visible spectrophotometer (JASCO, Model V-650). D2 and halogen lamps for wavelengths below and above 340 nm, respectively,
and an integrating sphere (JASCO, Model ISV-469) were used forthe diffuse reﬂectance measurements. Measurements were performed at 290 K within the wavelength range 200–900 nm using
100 mg of fresh samples. Diffuse reﬂectance spectra were converted to absorption spectra on the basis of the Kubelka–Munkequation [36,37]. The bandgap value was evaluated on the basisof either simple extrapolation of the absorption edge or the ﬁt to
the equation of Davis and Mott [37],a  hm / ðhm  Eg Þn ;in which a, h, and m are the absorption coefﬁcient, Planck’s constant,and the wavenumber, respectively, and n is 1/2, 3/2, 2, and 3 forallowed direct, forbidden direct, allowed indirect, and forbidden
indirect transitions, respectively.
Cu, Zn, and Ga K-edge X-ray absorption ﬁne structure (XAFS)
spectra were measured at 30–290 K in transmission mode in the
Photon Factory at the High Energy Accelerator Research Organization (Tsukuba) on beamlines 7C and 9C. The storage-ring energy
was 2.5 GeV and the top-up ring current was 450 mA. A Si(1 1 1)
double-crystal monochromator was inserted into the X-ray beampath. The X-ray intensity was maintained at 65% of the maximum
ﬂux using a piezo translator set to the crystal. The slit opening size
was 1 mm (vertical)  (1–3) mm (horizontal) in front of the I0 ionization chamber. The I0 and Itransmit ionization chambers were
purged with N2 and Ar, respectively. The scan steps were 9,0.3, and 2.5 eV in the pre-edge, edge, and postedge regions,
respectively. The data accumulation time was 1–5 s for each data
point. The Zn K, Cu K, and Ga K-edge absorption energy values
were calibrated to 9660.7, 8980.3, and 10,368.2 eV for the spectraof Zn, Cu, and Ga metals, respectively [38,39]. The energy positionof the monochromator was reproduced within ±0.1 eV.
2.3. Analyses
The XAFS data were analyzed using an XDAP package [40]. The
pre-edge background was approximated by a modiﬁed Victoreen
function, C2/E2 + C1/E + C0. The background of the postedge oscillation was approximated by a smoothing spline function, calculatedby an equation for the number of data points, where k is the wavenumber of photoelectrons:
Data
PointsXðlxi  BGi Þ2

i¼1

expð0:075ki Þ

2

6 smoothing factor:

Multiple-shell curve-ﬁt analyses were performed for the
Fourier-ﬁltered k3-weighted extended X-ray absorption ﬁne structure (EXAFS) data in k- and R-space using empirical amplitude andphase-shift parameters extracted from the EXAFS data for CuO and
ZnO. The interatomic distance (R) and its associated coordination

125

number (N) for the CuAO and ZnAO pairs were set to
0.19521 nm with the N value of 4 [41] and to 0.1978 nm with
the N value of 4, respectively, on the basis of their crystal structure[42]. The many-body reduction factor S20 was assumed to be equal
for the sample and the reference.
The amplitude and phase-shift parameters were also generated
theoretically for Zn  Zn (R = 0.308 nm with an N value of 6),
Zn  Zn (R = 0.53347 nm with an N value of 6), and Zn  Zn pairs(R = 0.616 nm with an N value of 6) using a Zn57(OH)72 model(Scheme 1A-i), and for the Zn  Al pair (R = 0.308 nm with an N
value of 2) using a Zn45Al12(OH)72 model (Scheme 1A-ii) using theab initio calculation code FEFF 8.4 [43] in the self-consistent ﬁeld
mode. One-fourth of the Zn sites were periodically replaced by Al
atoms in the Zn45Al12(OH)72 model. The exchange–correlation
potential of Hedin–Lundqvist was chosen. The Debye temperatureof the LDH models was assumed to be equal to that of ZnO(370 K) [44]. The goodness of ﬁt was given as requested by the Committee on Standards and Criteria in X-ray Absorption Spectroscopy.
The Zn, Cu, and Ga K-edge X-ray absorption near-edge structure(XANES) spectra were theoretically generated using FEFF 8.4 operated in the self-consistent ﬁeld mode and the fully multiple scattering mode.
2.4. Photocatalytic reduction of CO2
Photocatalytic reactions were conducted in a closed circulating
system using a photoreaction cell with a ﬂat quartz bottom (total
volume of reaction system 186 ml, Fig. S1) [45]. One hundred milligrams of the LDH catalyst was homogenously spread in a quartz
reaction cell (bottom plate area 23.8 cm2) connected to the circulation loop and evacuated (106 Pa) at 290 K for 2 h as pretreatment
until the desorbed gas was detected by an online gas chromatograph (GC). Either 21.7 kPa of H2 (1.67 mmol) or 2.3 kPa of CO2(0.177 mmol) + 21.7 kPa of H2 (1.67 mmol) was introduced and allowed to circulate for 30 min in contact with the catalyst to reach asorption equilibrium before the photoillumination.
The catalyst was then illuminated with UV–visible light obtained from a 500-W xenon arc lamp (Ushio, Model UI-502Q)
through the ﬂat bottom of the quartz reactor (Fig. S1A) for 5 h.
The distance between the bottom of the reactor and the lamp exit
window was set to 20 mm. The light intensity at the wavelength of
555 nm was 42 mW cm2 at the center of the sample cell and
28 mW cm2 at the periphery of the bottom plate of the sample
cell. Note that the intensity was measured at 555 nm but the Xe
arc lamp irradiated in a wide spectrum between 200 and1100 nm. The temperature was within the range 305–313 K at
the catalyst position during the illumination (Fig. S2). The transmitted light intensity was also evaluated by changing the catalystamount between 0 and 0.2 g (Fig. S3). Half of the incident light wasabsorbed or scattered by 0.1 g of catalyst charged and half wastransmitted. Because the transmitted light intensity monotonicallydecreased as the catalyst amount increased, most of 0.1 g of catalyst should have been effectively illuminated. As a comparison,
one of the LDH catalysts was illuminated with UV–visible light obtained from another 500-W xenon arc lamp (Ushio, Model SXUID502XAM) through a UV hot mirror (3.3 mm thick; Edmond), awater ﬁlter (100-mm thickness), and the ﬂat bottom of the quartz
reactor (Fig. S1B). The distance between the bottom of the reactor
and the lamp exit window was set to 206 mm. The light intensity
at 555 nm was 106 mW cm2 at the maximum of the sample celland 76 mW cm2 at the periphery of the bottom plate of the sample cell.
To evaluate the dependence of the photocatalytic reaction rates
on wavelength, a sharp cutoff ﬁlter, UV-32 (>320 nm), L-37(>370 nm), L-42 (>420 nm), or Y-48 (>480 nm) (Kenko), was set
at the lamp exit window [45]. The liquid nitrogen trap was used126

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

Table 12 mH2 O (x = 0, 1.5; MIII = Al, Ga) layered double hydroxides and reference catalysts.
Physicochemical characterization of ½Zn3x Cux MIII ðOHÞ8 þ
2 ðCO3 Þ
SBET (m2 g1)

SampleEg (eV)
Extrapolated

Fit to a  hm / (hm  Eg)n, n =Interlattice distance (nm)

1/2

3/2

2

3

(0 0 3)(1 1 0)III(A) Zn3x Cux Al LDHs
2
½Zn3 AlðOHÞ8 þ mH2 O2 ðCO3 Þ

57

5.7

5.9

5.5

5.3

5.0

0.757

0.154

2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O2 ðCO3 Þ

38

4.1

4.5

3.8

3.4

2.7

0.753

0.154

2
 mH2 O
½Zn3 GaðOHÞ8 þ
2 ðCO3 Þ

70

5.6

5.9

5.4

5.2

5.0

0.751

0.155

2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

57

3.5

4.2

3.2

3.0

2.6

0.753

0.155

(B) Zn3xCuxGa(C) References
CuAZnOaGa2O3b
a
b

III

LDHs

41
2.0

The CuAZnO catalyst was prepared via the reported inverse precipitation method starting from copper and zinc nitrates (atomic ratio Cu:Zn = 54:46) with NaHCO3 [27].
Commercial sample (Kanto Chemical, 99.0%). The major crystalline phase was a-Ga2O3.

for the reaction in H2 gas only to condense the formed CO2, CH3OH,
and H2O. Products and reactants were analyzed using packed columns of Molecular Sieve 13X-S and PEG (polyethylene glycol)6000 supported on Flusin P (GL Sciences) set in the online GC
equipped with a thermal conductivity detector (Shimadzu, Model
GC-8A). Two blank tests were performed: a reaction test illuminated with UV–visible light in the absence of photocatalysts andanother test for catalyst in the absence of any light. For the latter
test, the reactor was completely wrapped with 20-lm-thick aluminum foil. To conﬁrm the photocatalytic methanol synthesis, reactivity tests on carbonates in LDHs in H2 gas and photocatalytic
tests on LDHs in CO2 gas were also performed. Furthermore, thechange in the photocatalytic performance when the gas was
switched to a mixture of CO2 + H2 was investigated.

3. Results
3.1. Nitrogen adsorption, XRD, SEM, and UV–visible spectra
The SBET value was measured on the basis of the N2 adsorption
at 77 K for LDH compounds preheated at 383 K (Table 1). The SBET
2value for ½Zn3 GaðOHÞ8 þ mH2 O (70 m2 g1) was 21%
2 ðCO3 Þþgreater than that for ½Zn3 AlðOHÞ8 2 ðCO3 Þ2  mH2 O. The SBET value2
decreased by 19% for ½Zn3 GaðOHÞ8 þ mH2 O and by 34%
2 ðCO3 Þ2
for ½Zn3 AlðOHÞ8 þ
ðCO
Þ

mH
O
by
substitutionof Cu atoms at
3
2
2
half of the Zn sites.

2
Fig. 1. XRD spectra of fresh samples of ½Zn3 AlðOHÞ8 þ mH2 O (a),2 ðCO3 Þ2
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O (b), ½Zn3 GaðOHÞ8 þ mH2 O (c), and2 ðCO3 Þ
2 ðCO3 Þ
2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O
(d)
and
the
used
samplesof
2 ðCO3 Þ2
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O (b-u) and ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ
2 ðCO3 Þ(d-u) in 2.3 kPa of CO2 + 21.7 kPa of H2 under UV–visible light for 5 h.

2
The XRD spectra measured for ½Zn3 AlðOHÞ8 þ mH2 O and2 ðCO3 Þ
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ
ðCO
Þ

mH
O
are
depictedinFig.
1a and b.
3
2
2
The diffraction peaks common at 2hB = 11.7°, 23.5°, 34.6°, 37.3°,
39.2°, 43.9°, 46.7°, 52.9°, 56.3°, 60.2°, 61.5°, and 65.5° were assigned to (0 0 3), (0 0 6), (0 0 9), (1 0 4), (0 1 5), (1 0 7), (0 1 8),(1 0 1 0), (0 1 1 1), (1 1 0), (1 1 3), and (1 1 6) diffraction for the regular layered structure, respectively [46,47]. The interlayer interval
was evaluated as 0.757 and 0.753 nm on the basis of the (0 0 3) diffraction angle (Table 1A), consistent with the values 0.757 and
0.756 nm, which were evaluated based on the (0 0 6) diffraction
angle. The interlayer interval value was 0.777 nm, on the basis ofthe (0 0 9) diffraction. The small difference may be due to the overlap of the diffraction peak (0 0 9) with the (0 1 2) diffraction. The
in-plane (1 1 0) diffraction angle (0.154 nm, Table 1A) within alayer corresponds to R(ZnAO) of 0.218 nm if complete ZnO6 octahedra are assumed. Compared to the R(ZnAO) values provided by2
the Zn K-edge EXAFS analyses for ½Zn3 AlðOHÞ8 þ mH2 O2 ðCO3 Þþ2
and ½Zn1:5 Cu1:5 AlðOHÞ8 2 ðCO3 Þ  mH2 O (0.205–0.207 nm), theLDH cation sheets were found to be slightly elongated in the direction of the cationic layer plane.
The 2hB angles of plane-to-plane diffraction peaks (0 0 3),(0 0 6), and (0 0 9) perpendicular to the cationic layers changed
negligibly when Ga atoms were substituted at the trivalent Al sites(Fig. 1c and d). In contrast, other peaks signiﬁcantly shifted toward
smaller angles, reﬂecting the greater ionic radius of GaIII than of
AlIII, e.g., 2hB for (1 1 0) diffraction of 60.2° (0.154 nm) for Al-containing LDHs to 59.5° (0.155 nm) for Ga-containing LDHs (Table 1).
2
In the spectrum d for Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O, peaks at
2 ðCO3 Þ
2hB = 27.9°, 30.3°, 35.7°, 57.3°, and 62.9° were also observed, whichwere derived from the impurity.
2
A representative SEM image for ½Zn3 AlðOHÞ8 þ mH2 O is
2 ðCO3 Þ
depicted in Fig. 2. Flat ﬂakes smaller than 100 nm were observed,
which suggested the coagulation of the stacked layered structure.
As indicated in Fig. 1b-u and d-u, the LDH samples after thephotocatalytic tests in 2.3 kPa of CO2 + 21.7 kPa of H2 underUV–visible light for 5 h differed negligibly from the corresponding
data for fresh samples (Fig. 1b and d); therefore, the layered2
structures of as-synthesized Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O2 ðCO3 Þþ
2
and Zn1:5 Cu1:5 GaðOHÞ8 2 ðCO3 Þ  mH2 O samples should be wellretained.
An extrapolation of the absorption edge in the UV–visible
absorption spectra was drawn (Fig. 3A). The intercepts with the2
2
x-axis for ½Zn3 AlðOHÞ8 þ mH2 O and ½Zn3 GaðOHÞ8 þ
2 ðCO3 Þ
2 ðCO3 Þ 
mH2 O were 218 and 222 nm, respectively (Spectra a, c), corre2
sponding to bandgap values of 5.7 for ½Zn3 AlðOHÞ8 þ
2 ðCO3 Þ 
2mH2 O and 5.6 eV for ½Zn3 GaðOHÞ8 þ
ðCO
Þ

mH
O
(Table1).
32
2The bandgap values based on the ﬁt to the equation of Davis andN. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

2
Fig. 2. SEM image for ½Zn3 AlðOHÞ8 þ mH2 O.
2 ðCO3 Þ

127

mH2 O (Spectrum d). The extrapolation of the absorption edge wasdrawn and the intercepts with the x-axis for ½Zn1:5 Cu1:5 AlðOHÞ8 þ22
ðCO3 Þ2 H2 O and ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2 ðCO3 Þ H2 O were 306 and354 nm, respectively, corresponding to the bandgap values of2
4.1 for ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O and 3.5 eV for
2 ðCO3 Þþ2
½Zn1:5 Cu1:5 GaðOHÞ8 2 ðCO3 Þ  mH2 O (Table 1). The bandgap
values calculated on the basis of the ﬁt to the Davis andMott equation by applying different n values were 2.7–4.52
for ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O and 2.6–4.2 eV for
2 ðCO3 Þ
2
½Zn1:5 Cu1:5 GaðOHÞ8 þ
ðCO
Þ

mH
3
2 O (Table 1 and Fig. 3B). The band2
gap values in the case of n = 3/2 were similar to the corresponding
bandgap values based on simple extrapolation. The electronic tran2
sition for ½Zn1:5 Cu1:5 MIII ðOHÞ8 þ mH2 O might be a forbidden
2 ðCO3 Þ
direct process from oxygen 2p to metal ns or np levels (n = 4 for Cu,
Zn, and Ga and n = 3 for Al).
2
3.2. XANES spectra for fresh ½Zn3x Cux XðOHÞ8 þ mH2 O2 ðCO3 Þ

Fig. 3. (A) Diffuse reﬂectance UV–visible absorption spectra of fresh samples of22
½Zn3 AlðOHÞ8 þ mH2 O(a),
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O(b),
2 ðCO3 Þ
2 ðCO3 Þ2
2
½Zn3 GaðOHÞ8 þ mH2 O (c), and ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (d).
2 ðCO3 Þ
2 ðCO3 Þ(Inset) Expanded spectra (a) and (c). (B) Fit to the equation of Davis and Mott(a  hm / ðhm  Eg Þn ) for data of spectrum (d). (a  hm)2, (a  hm)2/3, (a  hm)1/2, or(a  hm)1/3 vs hm was plotted and the intercept with the x-axis estimates the Eg
value.

Mott by applying different n values were within the range
2
5.0–5.9 eV both for ½Zn3 AlðOHÞ8 þ mH2 O and for2 ðCO3 Þ
2
½Zn3 GaðOHÞ8 þ
ðCO
Þ

mH
O.
The
bandgap
valuesin the case of322
n = 1/2 or 3/2 were relatively similar to the corresponding bandgap
values based on simple absorption edge extrapolation (Table 1).
2
The electronic transition for ½Zn3 MIII ðOHÞ8 þ mH2 O could
2 ðCO3 Þ
be a direct process from oxygen 2p to metal ns or np levels (n = 4
for Zn and Ga and n = 3 for Al). In addition, a weak peak was ob2served centered at 291 nm for ½Zn3 GaðOHÞ8 þ mH2 O and2 ðCO3 Þ
2a weaker peak was observed for ½Zn3 AlðOHÞ8 þðCO mH2 O3Þ2
(Fig. 3A, insets c, a).
By the inclusion of Cu, the UV absorption edge shifted toward the2lower-energy side for ½Zn1:5 Cu1:5 AlðOHÞ8 þ
2 ðCO3 Þ  mH2 O (Fig. 3A-b)
2
and shifted to further lower energy for ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2 ðCO3 Þ 

Zn K-edge XANES spectra for LDH samples in comparison to thereference spectra are shown in Fig. 4A. The entire spectral patterns2
2
for ½Zn3 AlðOHÞ8 þ mH2 O and ½Zn1:5 Cu1:5 AlðOHÞ8 þ
2 ðCO3 Þ
2 ðCO3 Þ 
mH2 O (Spectra b and c) closely resemble each other. In both
spectra, an intense peak at 9671.7 eV is accompanied by weak,broad peaks at 9688, 9715, and 9724 eV. The Cu K-edge XANES
pattern for the latter sample closely resembles those of thecorresponding Zn K-edge data (Fig. 4A-c and B-c), strongly suggesting an identical coordination environment for Cu and Zn in2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O. The energy values of postedge
2 ðCO3 Þ
peaks relative to the ﬁrst intense peak were 16, 44, and 52 eV for
the Cu K-edge, similar to the gaps of 16, 43, and 52 eV for the Zn
K-edge.
The patterns of these experimental spectra were reproduced in
the theoretically generated Zn K-edge spectrum for a complete octahedral model (Fig. 4A-f) and theoretically generated Cu K-edge spectrum for a complete octahedral model (Fig. 4B-h and Scheme 1A-v,
ii). The energy values of the Zn K postedge peaks relative to the ﬁrst
intense peak were 16, 39, and 47 eV in the theoretical spectrum(Fig. 4A-f) compared to 16, 43, and 52 eV in the corresponding
experimental spectrum (Fig. 4A-b). The energy values of the Cu K
postedge peaks relative to the ﬁrst intense peak were 14, 39, and46 eV in the theoretical spectrum (Fig. 4B-h) compared to 16, 44,
and 52 eV in the corresponding experimental spectrum (Fig. 4B-c).
2
Thus, the Zn and Cu sites in ½Zn3 AlðOHÞ8 þ mH2 O and2 ðCO3 Þ
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ
ðCO
Þ

mH
O
shouldshow
nearlycomplete322octahedral coordination.
The Zn K-edge XANES spectra for ½Zn3x Cux GaðOHÞ8 þ
2
ðCO3 Þ2  mH2 O (x = 0, 1.5) also consisted of intense peaks at
9671.7–9672.2 eV accompanied by weak, broad peaks at 9688–
9689 eV (Fig. 4A, Spectra d and e), similar to spectra b and c for2½Zn3x Cux AlðOHÞ8 þ mH2 O (x = 0, 1.5). A minor difference
2 ðCO3 Þ
was that weak, broad peaks at 9715 and 9724 eV observed for2
½Zn3x Cux AlðOHÞ8 þ mH2 O merged into a weak, broad peak
2 ðCO3 Þ
2ranging from 9718 to 9720 eV for ½Zn3x Cux GaðOHÞ8 þ
2 ðCO3 Þ 
mH2 O. The energy position of the experimental three peaks for2
½Zn3x Cux GaðOHÞ8 þ mH2 O was reproduced in the theoret2 ðCO3 Þ
ically generated Zn K-edge XANES spectrum for a complete octahedral model (Scheme 1A-i; Fig. 4A-g). Note that 30Zn, 29Cu, and 31Ga
are similar photoelectron backscatters compared to the lighter
element 13Al.
The Cu K-edge XANES spectrum for ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2
ðCO3 Þ2 H2 O closely resembled the corresponding Zn K-edge data
(Fig. 4A-e and B-e). The energy values of the postedge peaks relative to the ﬁrst intense peak were 17 and 48 eV for the Zn K-edge,
identical to the gaps for the Cu K-edge. The energy position of
the three experimental peaks in the Cu K-edge spectrum e for2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O was closely reproduced in the
2 ðCO3 Þ

128

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

2
2
Fig. 4. (A) Normalized Zn, (B) Cu, and (C) Ga K-edge XANES spectra for ZnO (A-a), CuO (B-a), ½Zn3 AlðOHÞ8 þ mH2 O (b), ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O (c),2 ðCO3 Þ
2 ðCO3 Þ2
2
½Zn3 GaðOHÞ8 þ mH2 O (d), and ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (e). Theoretical spectra generated using FEFF 8.4 for Zn45Al12(OH)72 (f), Zn57(OH)72 (g),2 ðCO3 Þ
2 ðCO3 Þ
CuZn44Al12(OH)72 (h), CuZn56(OH)72 (i), and GaZn56(OH)72 models (j). The central, X-ray absorbing atoms were Zn (f, g), Cu (h, i), and Ga (j).

theoretical Cu K-edge spectra for the complete octahedral model2
(Fig. 4B-i). Thus, as in the case of ½Zn3x Cux AlðOHÞ8 þ
2 ðCO3 Þ 
mH2 O, the Zn and Cu sites show nearly complete octahedral
coordination.
The Cu K-edge XANES pattern for CuO (Fig. 4B-a) was similar to2
those for ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (spectrum e) and the2 ðCO3 Þ
CuZn56(OH)72 model (spectrum i). The Cu sites in CuO are coordinated with four equatorial O atoms (average CuAO distance
0.19521 nm) and two farther axial O atoms (average CuAO distance 0.2817 nm). The Cu K-edge XANES pattern could be relatively
insensitive to the variation in CuAO distances [41].
2The Ga K-edge XANES spectra for ½Zn3x Cux GaðOHÞ8 þ
2 ðCO3 Þ 
mH2 O (x = 0, 1.5) are depicted in Fig. 4C-d, e. The intense peak situated just above the edge (10,380.6 eV) and the three weak, broad
postedge peaks appeared. The energy values of the postedge peaks
relative to the ﬁrst intense peak were 16–17, 40–41, and 47–48 eV,
somewhat different from the gaps of 17 and 43 eV for the theoretically generated Ga K-edge spectrum for the complete octahedralmodel (Fig. 4C-j). A possible explanation for this discrepancy is that
there is a signiﬁcant distortion of GaIII site coordination comparedto the nearly complete octahedral ZnO6 or CuO6 coordination.

3.3. XANES spectra for partially dehydrated and CO2-adsorbed
2½Zn3x Cux XðOHÞ8 þ mH2 O2 ðCO3 Þ
2When the ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O sample was
2 ðCO3 Þ
heated at 383 K for 1 h in vacuum, the intensity of the ﬁrst post-edge peak at 9689 eV in the Zn K-edge XANES for a fresh sample(Fig. 5A-e) decreased to 65% (Fig. 5A-e-383). In fact, the postedge
peak intensity decreased to 58% when the theoretically generated
XANES spectra for Zn57(OH)7238H2O model (Scheme 1A-i and
Fig. 5A-k) lost all the structural water (Fig. 5A-g). In contrast, aclear additional peak was not found in the Fourier transform of
the theoretically generated Zn K-edge EXAFS besides those derived
from the in-plane ZnAO, Zn  Zn (R = 0.308 nm), Zn  O
(R = 0.37722 nm), Zn  Zn (R = 0.53347 nm), or Zn  Zn pair(R = 0.616 nm) on going from the spectrum for the Zn57(OH)72
model to the one for the Zn57(OH)7238H2O model. When
34.3 kPa of CO2 was introduced into the 383 K-heated LDH sample,
the postedge peak intensity increased slightly to 72% of the fresh
sample (Fig. 5A-e-CO2-290). When the LDH sample was heated to423 K in CO2, the peak intensity further increased to 83% of spectrum e (spectrum e-CO2-423, Table 2A). Note that the postedge
peak is in the ‘‘multiple scattering’’ region, and the peak intensity
increased in the presence of either water or CO2 molecules between the LDH cationic layers.
A similar in situ trend was observed in the Cu K-edge XANES for2
the LDH compound of ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O. The
2 ðCO3 Þintensity of the ﬁrst postedge peak at 9016 eV decreased to 76%
by heating in vacuum at 383 K (Fig. 5B-e ? e-383). Upon introduction of CO2, the peak intensity increased back to 87% of the fresh
sample at 290 K (spectrum e-CO2-290) and 91% of the fresh one
at 423 K (spectrum e-CO2-423, Table 2B).
It should be noted that these semireversible in situ peakintensity changes in both the Zn and Cu K-edges for2
2
Fig. 5. Normalized Zn (A, D), Cu (B), and Ga K-edge XANES (C, E) spectra for ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (A–C) and ½Zn3 GaðOHÞ8 þ mH2 O (D and E). As2 ðCO3 Þ
2 ðCO3 Þsynthesized (e), sample (e) in vacuum at 383 K (e-383), and sample (e-383) in 34.3 kPa of CO2 at 290 K (e-CO2-290) or 423 K (e-CO2-423). Theoretical spectrum using FEFF 8.4
for Zn57(OH)72 (g), CuZn56(OH)72 (i), GaZn56(OH)72 (j), Zn57(OH)7238H2O (k), CuZn56(OH)7238H2O (l), and GaZn56(OH)7238H2O (m). The central, X-ray absorbing atoms wereZn (g, k), Cu (i, l), and Ga (j, m).

129

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135
Table 2First postedge peak intensity in normalized Zn, Cu, and Ga K-edge XANES spectra.a
SampleAssynthesizedIn
vacuumIn CO2(34.3 kPa)

383 K

290 K

423 K

0.12
(64%)
0.06
(65%)

0.12
(64%)
0.06
(72%)

0.12
(63%)
0.07
(83%)

0.08

0.06
(76%)

0.07
(87%)

0.08
(91%)

2
½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

0.23

2
 mH2 O
½Zn1:5 Cu1:5 GaðOHÞ8 þ
2 ðCO3 Þ

0.29

0.16
(70%)
0.22
(76%)

0.16
(70%)
0.22
(76%)

0.16
(70%)
0.24
(84%)

(A) At Zn K-edge (9688–9689 eV)2
½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

0.19

2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

0.09(B) At Cu K-edge (9016 eV)2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ(C) At Ga K-edge (10,397–10,398 eV)aValues in parentheses: peak intensity ratio to the as-synthesized sample.

2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O were not observed for the
2 ðCO3 Þ
2ﬁrst postedge peak at 9688 eV for ½ZnGaðOHÞ8 þ mH2 O2 ðCO3 Þ(Fig. 5D). The peak intensity decreased to 64% of that for the fresh
one in vacuum at 383 K and remained unchanged in CO2 at 290–
423 K (Table 2), suggesting speciﬁc interaction of Cu sites with
CO2 molecules.
The intensity of the ﬁrst postedge peak of the Ga K-edge XANES2
of the fresh ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O sample at
2 ðCO3 Þ
10,398 eV decreased to 76% by heating at 383 K (Fig. 5C-e ? e383). The peak intensity remained the same upon introducing
CO2 at 290 K (spectrum e-CO2-290), but it increased again to 84%
of that of the fresh sample at 423 K (spectrum e-CO2-423,
2
Table 2C). In ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O, the extent of
2 ðCO3 Þ
recovery of the ﬁrst postedge peak intensity due to the interaction
with CO2 was in the following order: Cu sites > Zn sites > Ga sites.
The peak recovery was always enhanced by heating from 290 to423 K, suggesting activated sorption of CO2 between the LDH
layers.

Table 3
Photocatalytic rates of CO2 reduction with H2 using LDH and reference photocatalysts.a
I
Photocatalyst

Formation rate (nmol h

1

g1
cat )

Quantum yield (ppm)f
CH3OH

Conversion

CH3OH

CO

R

13
(±1)
42
(±3)

30
(±2)
66
(±5)

43

0.01b

110

0.02b

2.8

0.83

3.6

0.02

250

1.5

0.19

1.7

0.03

Ga2O3e

<2

80
(±6)
79
(±6
620
(±6)
370
(±4)
580
(±6)
30
(±2)
47
(±3)

130

CuAZnOd

51
(±4)
170
(±14)
39
(±3)
130
(±10)
200
(±16)
<2

CO

Selectivity to CH3OH(mol%)(%, C-base)

R(A) In H2 (21.7 kPa)2
½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ
2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

30
(±3)
39
(±4)

(B) In CO2 (2.3 kPa) + H2 (21.7 kPa)2
½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ
2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ
2
½Zn3 AlðOHÞ8 þ mH2 O2 ðCO3 Þ
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O2 ðCO3 Þ
2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 Oc
2 ðCO3 ÞIIPhotocatalyst

660

0.16

500

0.11

780

0.16

30

0.006

39
(±4)
68
(±4)
5.9
(±0.5)
26
(±2)
26
(±2)
<7

47

0.01

<4

Formation rate per speciﬁc surface area
(nmol h1 m2)

Formation rate per amount of Cu
1(nmol h1 mmolCu )

CH3OH

CO

R

CH3OH

CO

R

2
½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

0.19

0.43

0.62

–

–

–

2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

0.74

1.2

1.9

13

20

34

2
 mH2 O
½Zn3 GaðOHÞ8 þ
2 ðCO3 Þ

0.73

1.1

1.9

–

–

–

2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ

3.0

1.4

4.4

53

25

78

0.68

11

12

–

–

–

2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O2 ðCO3 Þ

3.4

9.7

13

37

100

140

2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 Oc
2 ðCO3 Þ
CuAZnOdGa2O3e

5.3

15

21

56

160

220

<0.05
<1

0.73
23

0.73
23

<0.30
–

4.5
–

4.5
–

(A) In H2 (21.7 kPa)

(B) In CO2 (2.3 kPa) + H2 (21.7 kPa)2 mH2 O
½Zn3 AlðOHÞ8 þ
2 ðCO3 Þa

The catalyst amount was 100 mg. Values in parentheses are experimental errors for evaluation.
Calculated on the basis of molar amount of interlayer carbonate ions converted to methanol and CO.
The photon ﬂux illuminated using Model SX-UID502XAM was 2.5 times greater than other photocatalytic tests using Model UI-502Q in the comparison of the whole
wavelength region.
d
The CuAZnO catalyst was prepared via the reported inverse precipitation method starting from copper and zinc nitrates (atomic ratio Cu:Zn = 54:46) with NaHCO3 [27].
eCommercial sample (Kanto Chemical, 99.0%). The major crystalline phase was a-Ga2O3.
fTransmitted light intensity (Fig. S3) was subtracted from the incident light intensity for the evaluation.
b

c

130

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

Similarly to the ﬁrst postedge in the Zn K-edge for2
½Zn3 GaðOHÞ8 þ mH2 O, the intensity of the ﬁrst postedge
2 ðCO3 Þ
peak in the Ga K-edge spectrum decreased to 70% of the fresh
one and remained the same in CO2 at 290–423 K (Fig. 5E andTable 2C, A), suggesting poor interaction of the Zn/Ga sites with
CO2 molecules.
3.4. Photocatalytic reduction rates of CO2
First, the reactivity of the interlayer carbonate ions of LDHs was2tested in 21.7 kPa of H2 for ½Zn3x Cux GaðOHÞ8 þ mH2 O2 ðCO3 Þ(x = 0, 1.5) (Table 3I-A and Fig. 6A and B). Methanol was detected
as a byproduct using both catalysts under UV–visible light in addition to the major product of CO (70–61 mol%). The conversion ofcarbonate ions into methanol + CO was 0.01–0.02%.
2
The photocatalysis of ½Zn3 GaðOHÞ8 þ mH2 O was tested
2 ðCO3 Þ
in 2.3 kPa of CO2 + 21.7 kPa of H2 under illumination by UV–visible
light (Fig. 6C). The major product was CO (selectivity 61 mol%), andmethanol was also formed. The catalyst was compared to2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (Fig. 6D). The formation rate
2 ðCO3 Þ1of CO was essentially identical (79 nmol h g1cat ), whereas that of
1 1methanol increased (170 nmol h gcat ) by a factor of 3.3 for½Zn1:5 Cu1:5 GaðOHÞ8 2 ðCO3 Þ2  mH2 O. The methanol selectivity was
improved to 68 mol% (Table 3I-B).
2
Next, the ½Zn3 AlðOHÞ8 þ mH2 O catalyst was tested
2 ðCO3 Þ
2
(Fig. 6E) in comparison to the ½Zn3 GaðOHÞ8 þ mH2 O catalyst
2 ðCO3 Þ(Fig. 6C). Total product formation rate was enhanced by a factor of2
5.1 of ½Zn3 GaðOHÞ8 þ mH2 O by the substitution of Ga for2 ðCO3 Þ
Al. The major product was CO (selectivity 94 mol%). Then the2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O catalyst (Fig. 6F) was compared2 ðCO3 Þ
2to the ½Zn3 AlðOHÞ8 þðCO mH2 O catalyst. The total product3Þ
2

formation rate decreased by 24%. Similarly to the Cu substitution2
into ½Zn3 GaðOHÞ8 þ mH2 O, the methanol formation rate
2 ðCO3 Þ
2was promoted by a factor of 3.3 of ½Zn3 AlðOHÞ8 þ mH2 O.
2 ðCO3 Þþ
2The conversion using ½Zn1:5 Cu1:5 AlðOHÞ8 2 ðCO3 Þ  mH2 O was improved by increasing the light intensity of Model UI-502Q (0.11%)
to 2.5 times using Model SX-UID502XAM (0.16%, Table 3I-B) in thecomparison of the whole wavelength region.
The formation rates per speciﬁc surface area of catalyst andthose per amount of Cu are also summarized in Table 3II. Especially, a concern that 500-W high-intensity light may lead to a local
high temperature for the catalyst can be estimated based on theformation rates per amount of Cu, although the catalyst temperature was monitored to be 305–313 K (Fig. S2). Conventional
CuAZnO catalyst was quite inactive per amount of Cu and formed
only CO. The Cu sites in LDH are CuII, which is normally considered
to be inactive for thermal methanol synthesis [27]. Further,2
LDHs that do not contain Cu (½Zn3 GaðOHÞ8 þ mH2 O,2 ðCO3 Þ
2
½Zn3 AlðOHÞ8 þ
ðCO
Þ

mH
O)
did
producemethanol(Table3I-B).
3
2
2
Thus, the catalysis listed in Table 3 should be photocatalytic below
313 K.
The formation rates of water were between 0.5 and1
2:0 lmol h g1
cat in the kinetic tests (Fig. 6C–F). These rates exceed
the amount of water that was catalytically formed as the productsof the following equations:

CO2 þ H2 ! CO þ H2 O
CO2 þ 3H2 ! CH3 OH þ H2 O:This discrepancy is due to the desorption of interlayer water molecules of LDHs [15–19].
A control reaction experiment was performed using2
½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O in the absence of light2 ðCO3 Þ(Fig. 6BL-2). Another control experiment included a reaction

Fig. 7. Time course of photoreactions in CO2 (2.6 kPa) only in ﬁrst 3 h and in CO2(2.6 kPa) + H2 (21.7 kPa) after 3 h from the start of the reactions for 100 mg of2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O. Other reaction conditions and legends are the2 ðCO3 Þsame as those for Fig. 6.

Table 42Comparison of photoconversion of gases over ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 Oa.
2 ðCO3 Þ
Condition

Formation rate
(nmol hFig. 6. Time course of photoreactions in only H2 alone (21.7 kPa; A, B) and in CO2(2.3 kPa) + H2 (21.7 kPa) (C–F, BL-1, BL-2). One hundred milligrams of the LDH
catalyst was charged except for the blank test BL-1 (empty reactor):2
2
½Zn3 GaðOHÞ8 þ mH2 O (A, C), ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (B, D),
2 ðCO3 Þ
2 ðCO3 Þ2
2
½Zn3 AlðOHÞ8 þ mH2 O (E), and ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O (F, BL-2).
2 ðCO3 Þ
2 ðCO3 Þ
The reactor was illuminated with a 500-W Xe arc lamp except for blank test BL-2
(reactor in darkness). CO2 (; diamond), H2O (N; triangle), CH3OH (j; square), andCO (d; circle).

In methanol (1.7 Pa), no light
In methanol (1.7 Pa), under lightb
In methanol (1.7 Pa) + H2 (21.7 kPa), no light
In methanol (1.7 Pa) + H2 (21.7 kPa), under lightb
In H2 (21.7 kPa), under lightbab

The catalyst amount was 100 mg.
Photocatalytic tests using Model UI-502Q.

1

g1
cat )

CH3OH

CO

–
26
–
24
42

<2
26
<2
82
66

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

131

Fig. 8. Dependence of the formation rates of methanol and CO on the cutoff wavelength for the photoreactions in CO2 (2.3 kPa) + H2 (21.7 kPa) for 100 mg of2
2
½Zn3 GaðOHÞ8 þ mH2 O (A) or ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O (B). Sharp cutoff ﬁlters UV-32, L-37, L-42, or Y-48 were used at the exit of the Xe arc lamp. Other
2 ðCO3 Þ
2 ðCO3 Þreaction conditions and legends are the same as those for Fig. 6. (Inset) Action spectrum of product formation rates vs wavelength.

without any catalyst under the illumination of UV–visible light(reactor only, Fig. 6BL-1). In both experiments, methanol and CO
were not found.
To ensure methanol formation starting from gaseous CO2 andH2, the products were monitored when ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2
ðCO3 Þ2  mH2 O was in 2.6 kPa of CO2 for the ﬁrst 3 h under UV–
visible light and then 21.7 kPa of H2 was added for the next 3 h,
keeping the UV–visible light on (Fig. 7). Methanol was not observed during the ﬁrst 3 h; however, upon the addition of H2 gas,
1
methanol formation started at a rate of 160 nmol h g1
cat over the
next 3 h. The formation rate was similar to that observed in1Fig. 6D. The methanol formation rate was only 42 nmol h g1cat
for the reaction of interlayer carbonate with gaseous H2(Table 3A). Thus, methanol was formed predominantly starting2
from gaseous CO2 + H2 using ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O.
2 ðCO3 Þ
The stability of 1.7 Pa (0.13 lmol) of methanol was checked2
over ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O in the presence/absence
2 ðCO3 Þof light and with/without H2. Methanol was stable in the absenceof light both with and without H2, and no CO desorption
was observed (Table 4). Instead, methanol adsorbed on2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O at a rate of (more than)
2 ðCO3 Þ1 1
510 nmol h gcat within 1 h.
Next, when the catalyst was under UV–visible light without H2,
most of adsorbed methanol desorbed within 15 min. The formation
1
rate of CO was 26 nmol h g1cat , starting from methanol. When the
catalyst was under UV–visible light with H2, most of the adsorbed
methanol desorbed within 15 min again. The formation rate of CO
1
was 82 nmol h g1
cat starting from both methanol and the interlayer carbonate of LDH. The methanol amount increased at
1
24 nmol h g1
cat due to the balance of formation starting from car1
bonate of LDH reacted with H2 (42 nmol h g1
cat under a similar
photoreaction condition, Table 3A) and photodecomposition to
CO (Table 4).

(Fig. 8A) in accordance with the UV absorption edge, which canbe extrapolated to 222 nm (Fig. 3A-c). In addition, the CO formation rate dropped as the cutoff wavelength increased; however, a1signiﬁcant rate of formation (32 nmol h g1cat ) was observed when
the cutoff wavelength was 370 nm. This trend may originate fromthe UV absorption peak between 260 and 375 nm at the foot of the2main absorption edge for ½Zn3 GaðOHÞ8 þ mH2 O (Fig. 3A, in2 ðCO3 Þ
set c). The additional UV light absorption could be related to thesurface sites of LDH to transform CO2 to CO. When the cutoff wavelength was 420 nm, products were not detected using2
½Zn3 GaðOHÞ8 þ mH2 O.
2 ðCO3 Þ
2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O was photocatalytically ac2 ðCO3 Þtive when the cutoff wavelength was less than 370 nm, but theactivity decreased to zero when the cutoff wavelength was
420 nm (methanol formation) or 480 nm (CO formation)(Fig. 8B). This dependence is principally in accord with the UV–visible absorption edge, which can be extrapolated to 354 nm, but
slower CO formation when the cutoff wavelength was 420 nm
could be ascribed to a shoulder peak at the foot of the UV absorption edge and/or small visible light absorption between 400 and
480 nm (Fig. 3A-d). These trends are also clear in the action spectra
as insets of Fig. 8A and B. Major product formation corresponded to2
wavelengths less than 320 nm for ½Zn3 GaðOHÞ8 þ mH2 O,2 ðCO3 Þ
whereas both methanol and CO formation corresponded to wave2
lengths centered at 390 nm for ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O.
2 ðCO3 Þ

3.5. Dependence of photocatalytic reduction of CO2 on lightwavelength and catalyst durability
The dependence of rates of methanol and CO formation fromCO2 + H2 gases on light wavelength was investigated for2
2
½Zn3 GaðOHÞ8 þ mH2 O and ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2 ðCO3 Þ
2 ðCO3 Þ 
mH2 O photocatalysts using several sharp cutoff ﬁlters (Fig. 8).
2
Methanol formation using ½Zn3 GaðOHÞ8 þ mH2 O de2 ðCO3 Þ
creased sharply as the cutoff wavelength increased to 320 nm

Fig. 9. Time course of photoreactions in CO2 (2.3 kPa) + H2 (21.7 kPa) for 100 mg of2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O for 20 h. Other reaction conditions and legends
2 ðCO3 Þ
are the same as those for Fig. 6.

132

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

2
Finally, the durability of the ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þ
catalyst was tested for 20 h. The photocatalyst exhibited the best
selectivity to methanol (68 mol%) in 5 h (Table 3B and Fig. 6D).
In the 20-h test, methanol formation continued and the selectivity
was 68–57 mol% (Fig. 9). In every 5 h, the methanol formation rate
gradually decreased, starting from 79, to 57, 30, and then
1
19 nmol h g1cat , whereas the CO formation rate remained constant1(36 nmol h g1
cat Þ for 20 h. Throughout the 20-h test, no otherproducts were found except for methanol, CO, and water.

4. Discussion
4.1. Catalyst syntheses and photocatalysis
LDH compounds were synthesized with chemical compositions2 mH2 O, where M was Al or Ga and the½Zn3x Cux MIII ðOHÞ8 þ
2 ðCO3 Þx value was 0–1.5. The synthesis of the LDH structure was conﬁrmed on the basis of the XRD spectra (Fig. 1a–c). The divalentCu was successfully incorporated into the octahedral sites in the
cationic layers with a maximum Cu:Zn atomic ratio of 1:1, as conﬁrmed by Cu and Zn K-edge XANES spectra (Fig. 4A and B). The
incorporation of Cu with a maximum Cu:Zn atomic ratio of 7:3
III2
was reported for ½Zn3x Cux Al ðOHÞ8 þ mH2 O [27]. In the
2 ðCO3 Þcase where M was Ga and the x value was 1.5, the diffraction peaks
derived from the impurity phase(s) were observed in the XRD
spectrum at 2hB = 27.9°, 30.3°, 35.7°, 57.3°, and 62.9° (Fig. 1d).
2The SBET value of ½Zn3 GaIII ðOHÞ8 þ mH2 O was slightly high2 ðCO3 ÞIII2
er (70 m2 g1) than that of ½Zn3 Al ðOHÞ8 þ mH2 O2 ðCO3 Þ
2 1(57 m g ; Table 1). The SBET value decreased by 19–34% whenequimolar amounts of Cu and Zn metal were included in the LDHs
and when M was either Al or Ga. The coagulated ﬂat ﬂakes smaller
than 100 nm in the SEM images were suggested to be stacked lay2
ers of ½Zn3 AlðOHÞ8 þ mH2 O (Fig. 2).
2 ðCO3 Þ
2Synthesized LDHs formulated as ½Zn3x Cux MIII ðOHÞ8 þ
2 ðCO3 Þ 
mH2 O were applied to the photocatalytic reduction of CO2 (Table 3
and Fig. 6). First, the conversion of interlayer carbonate ions of2
½Zn3x Cux GaðOHÞ8 þ mH2 O (x = 0, 1.5) was tested in the
2 ðCO3 Þpresence of H2 gas. Methanol formation was conﬁrmed at rates of
1
13—42 nmol h g1
cat and the total conversion of carbonate to methanol + CO was 0.01–0.02% (Table 3A). The CO2 formation derived
from the carbonate ions was of the same or a slightly higher orderthan the formation of methanol and CO (Fig. 6A and B). Because
the interlayer carbonate ions balance the charge with the½Zn3x Cux GaðOHÞ8 þ cation layers, it could be difﬁcult to convert
and remove carbonate ions from the interlayer space if an extra carbonate source or alternative anion source were not provided.
Photocatalytic methanol formation tests were performed in
the presence of CO2 and H2 gases using several LDH samples incomparison to a conventional CuAZnO catalyst and Ga2O3, whichhave primarily an a-crystalline phase (Table 3B). This is the ﬁrst
report to synthesize methanol starting from CO2 and H2
using photocatalysts under only UV–visible light. When2
½Zn3x Cux AlðOHÞ8 þ mH2 O (x = 0, 1.5) catalysts were used,
2 ðCO3 ÞCO (selectivity 94–74 mol%) was a major product. In this study,
exclusive CO photoformation was reported using Rh/TiO2 [29],ZrO2 [30,31], MgO [32], and Ga2O3 [33], and was observed using
CuAZnO or Ga2O3 (Table 3B). Among the LDH samples tested in2
this study, ½Zn3x Cux AlðOHÞ8 þ mH2 O catalysts were the2 ðCO3 Þmost active. The conversions (0.16–0.11%) were an order of magnitude greater than those obtained using references CuAZnO andGa2O3 (0.01–0.006%). The formation of methanol + CO was accel1erated by increasing the light intensity (500 nmol h g1cat ) to 2.5
1 1
times (780 nmol h gcat ) for ½Zn1:5 Cu1:5 AlðOHÞ8 2 ðCO3 Þ2  mH2 O
in the comparison of the whole wavelength region. Methanol
was the major product when the ½Zn1:5 Cu1:5 GaðOHÞ8 2 ðCO3 Þ2 

mH2 O photocatalyst was used (selectivity 68 mol%; Table 3I-B).
Photocatalytic conversion of CO2 into methanol in aqueous solution was reported using Co-phthalocyanine/TiO2 [11] and Nd3+/
TiO2 [12]. The maximum formation rates of methanol were 9.3
1
and 23 lmol h g1cat , respectively, in addition to a higher forma1tion rate of formic acid using CoPc/TiO2 (150 lmol h g1cat ). The
disadvantage of LDH catalysts for slower photoconversion of
CO2 using hydrogen in this study may be due to slow gas diffusion in the interlayer space compared to the cases in thesereferences.
The dependence of the photocatalytic methanol formation rateon the light wavelength (Fig. 8) was compared to the UV–visible
absorption spectra for ½Zn3 GaðOHÞ8 2 ðCO3 Þ2  mH2 O (Fig. 3A-c
and A-d). The higher limit of the wavelength used to form metha2
nol was between 370 and 420 nm for ½Zn1:5 Cu1:5 GaðOHÞ8 þ
2 ðCO3 Þ 
mH2 O (Fig. 8B), in accordance with the UV absorption edge extrapolated to 354 nm or the foot (400 nm) of a smaller shoulder on theabsorption edge (Fig. 3A-d). The higher limit of wavelengths at
which methanol formed was less than 320 nm for2
½Zn3 GaðOHÞ8 þ mH2 O (Fig. 8A), in accordance with the
2 ðCO3 ÞUV absorption edge extrapolated to 222 nm (Fig. 3A-c). The COformation rates also decreased as the cutoff wavelength increased;
however, the trend was less steep than for methanol, and thus,
the CO selectivity was 100 mol% when the cutoff wavelength2
was 370 nm(½Zn3 GaðOHÞ8 þ mH2 O) and 420 nm
2 ðCO3 Þ
2
(½Zn1:5 Cu1:5 GaðOHÞ8 þ
ðCO
Þ

mH
O)
(Fig.
8). These minor differ32
2
ences probably occur because the UV–visible absorption is theelectronic transition to the energy level slightly below the bottomof the conduction band and this level was derived from the surface
sites selective for transforming CO2 into CO.
The action spectra in the insets of Fig. 8A and B were comparedto the light intensity distribution for Model UI-502Q obtainedusing a USR40 spectroradiometer (Ushio). Compared to the lightintensity peak centered at 380 nm (not shown), the quantum yield2
for methanol formation using ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O2 ðCO3 Þwas evaluated at 1.5 ppm (Table 3I-B). The quantum yield for all
products was 1.7 ppm. On the other hand, the action spectrum2
for ½Zn3 GaðOHÞ8 þ mH2 O (Fig. 8A, inset) was compared2 ðCO3 Þto the light intensity peak centered at 315 nm. The quantum yields
for methanol and CO were 2.8 and 0.83 ppm, respectively(Table 3I-B). As the bandgap of LDHs that do not contain Cu was
wider (5.6–5.7 eV), they would utilize UV light only between 200
and 220 nm (Fig. 3A).
The methanol formation and minor CO formation using the2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O photocatalyst was conﬁrmed
2 ðCO3 Þ
for as long as 20 h (Fig. 9), suggesting the suitability of the photocatalysts to capturing CO2 and converting it to fuel under continuous natural light. The gradual decrease of the methanol formation
rate may be affected by the readsorption of product methanol incontrast to the constant formation rate for CO, because product2
CO readsorption on ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O catalyst
2 ðCO3 Þshould be weaker.
2
For ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O under UV–visible light,
2 ðCO3 Þ
1the methanol formation rates increased from 42 nmol h g1
cat in1
H2 to 170 nmol h g1
in
CO
+
H
(Table3).
The
presenceof car2
2
catbonate is possibly the key to methanol photosynthesis; however,gaseous CO2 promoted the formation rates by a factor of 4.0. On
the other hand, the catalyst under UV–visible light produced no
product in CO2 and the photocatalysis required the addition of
H2 to the system (Fig. 7). No methanol was detected after the catalyst was suspended in deionized water under UV–visible light for
5 h (see the Supporting materials). Furthermore, no impurity that
contained carbon was included in the reagents used throughout
the LDH syntheses, except for sodium carbonate. Therefore, thesource of methanol was gaseous CO2 and interlayer carbonates.
The photoconversion of carbonates to methanol was favorable inN. Ahmed et al. / Journal of Catalysis 279 (2011) 123–135

the presence of gaseous CO2 and it most likely compensated for theinterlayer anions.
4.2. Spectroscopic insight and reaction mechanism of photocatalysis
The speciﬁc photocatalysis thus found to form methanol from
CO2 were correlated with the desorption trend of interlayer molecules (water and/or carbonates) and the adsorption of the substrate. The LDH structure is ultimately simple from the viewpointof heterogeneous catalysis because all the surface sites are hydroxygroups bound to [MO6] octahedra and the type of metal elements(Zn, Cu, Ga, or Al) is the only difference (Scheme 1A). The presenceor absence of interlayer molecules was difﬁcult to monitor using
EXAFS. However, the ﬁrst postedge peak (9688–9689 eV at theZn K-edge, 9016 eV at the Cu K-edge, and 10,397–10,398 eV at
the Ga K-edge) was found to be a critical indicator of the presenceor absence of these molecules (Fig. 5).
2
For the LDH compound ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O,2 ðCO3 Þwhich was the most methanol-selective, the intensity of the ﬁrst
postedge peak at the Zn K-edge decreased to 65% of that for the
fresh one in vacuum at 383 K due to the loss of structural watermolecules and/or carbonates. The Zn K postedge peak intensity de2
creased similarly for Cu-free ½Zn3 GaðOHÞ8 þ mH2 O to 64%
2 ðCO3 Þof that for the fresh one. The peak intensity increased to 83% in
CO2 at 423 K in the former Cu-containing LDH, whereas the peakremained unchanged for the latter Cu-free LDH in CO2 at 290–
423 K.
The desorption–adsorption trend monitored at the Zn K-edge2
for ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O was also evident in both2 ðCO3 Þ
the Cu and the Ga K-edge spectra. The recovery extent (=peak
increase due to CO2 sorption/peak decrease caused by the loss ofwater and carbonates) of the ﬁrst postedge peak intensity in thenormalized XANES spectra was of the following order: Cu sites(63%) > Zn sites (51%) > Ga sites (11%). In contrast, the ﬁrst Ga K
postedge peak remained unchanged for Cu-free ½Zn3 GaðOHÞ8 þ
2
ðCO3 Þ2  mH2 O in CO2 gas at 290–423 K.
More importantly, the methanol formation rate was enhanced2
by a factor of 3.3 by adding Cu to ½Zn3 GaðOHÞ8 þ mH2 O2 ðCO3 Þ(Table 3B). The Cu site should play the photocatalytic role by binding CO2 and coupling it with protons and photogenerated electrons, presumably utilizing a CuI and CuII redox couple. Thisbinding between CO2 and Cu in turn affects the multiple scattering
region of Zn and Ga sites as well, because interlayer hydroxygroups shared by Cu and Zn (Scheme 2a, M = Zn) and by Cu and
Ga (Scheme 2a, M = Ga) inserted CO2 to form hydrogen carbonate(Schemes 2b and 1B). The steps b ? c ? d in Scheme 2 should be

133

promoted utilizing the CuI and CuII redox couple. Further reduction
steps e ? f, f ? g, g ? h, and h ? a should be also coupled with the
CuI and CuII redox to ﬁnally produce methanol. If the species e isdehydrated, another catalytic cycle can be proposed to regenerate
species a. A clear explanation cannot be given at the moment for
why methanol and CO were exclusively obtained in the catalytic
cycles of Scheme 2. Formic acid and formaldehyde may be unstable
and/or insensitive by the gas-phase sampling of heterogeneous
catalysis using GC.
The role of the CuI sites supported on TiO2 in the CO2 photoreduction in alkaline aqueous solution under the UV light was suggested [48]. The formation of carbonate species over a MgO
catalyst was proposed, but the previous studies suggested it to
be inactive or to only lead to CO formation in H2 gas [32]. The discrepancy between these previous studies and this study should be
due to the presence or absence of Cu sites.
2
Photocatalytic activity of ½Zn3x Cux AlðOHÞ8 þ mH2 O2 ðCO3 Þ
2was greater than that of ½Zn3x Cux GaðOHÞ8 þ
ðCO
Þ

mH
3
2 O. The
2
exact photocatalytic role of MIII sites is not known at present.
One of the possibilities is the 50-nm difference of the UV lightabsorption edge between MIII = Ga and Al (Fig. 3A-d and A-b).
The peaks of product formation rates in the action spectrum for2
½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O centered at 390 nm (Fig. 8B,
2 ðCO3 Þ
inset) were fairly close to the light intensity peak from the Xe arc
lamp centered at 380 nm. The wavelength shift of 50 nm toward2
the higher energy side for ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O2 ðCO3 Þ
may be serious because intense light centered at 380 nm did notcontribute to the charge separation for the catalyst. A decrease inbasicity from substituting an Al site for Ga in the layered structure
was reported [49]. The afﬁnity of Ga sites in LDHs with CO2 wassuggested [26].
The photocatalytic energetic is depicted in Scheme 3. The band2
gap values (5.7–4.1 eV) estimated for ½Zn3 AlðOHÞ8 þ mHþ2O
2 ðCO3 Þ
and ½Zn1:5 Cu1:5 AlðOHÞ8 2 ðCO3 Þ2  mH2 O were greater than that for
ZnO, enabling CO2 photoreduction at relatively higher potentials
to form methanol at 0.32 V and to form CO at 0.11 V by supplying
photoexcited electrons to CO2 in the conduction band. On the other
hand, photogenerated holes in the valence band were neutralizedwith electrons formed during H2 oxidation (0 V). This trend and elec2
tron ﬂow are also evident for ½Zn3x Cux GaðOHÞ8 þ mH2 O2 ðCO3 Þ(bandgap values estimated as 5.6–3.5 eV; Table 1B). The selectivity
to produce methanol vs CO formation increased by adding Cu to2
½Zn3 MIII ðOHÞ8 þ mH2 O (M = Al, Ga; Table 3). The bandgap
2 ðCO3 Þ
became narrower with the addition of Cu; however, the potential
for the methanol formation step (CO2 + 6H+ + 6e ? CH3OH + H2O)
[2] was 0.21 V higher than that for the CO formation step (Scheme 3).

Scheme 2. Proposed photocatalytic catalytic cycle of CO2 reduction to methanol or CO over ZnACuAGa LDH catalysts.

134

N. Ahmed et al. / Journal of Catalysis 279 (2011) 123–1352
2
Scheme 3. Energy levels at the bandgap for ZnO, ½Zn1:5 Cu1:5 AlðOHÞ8 þ mH2 O, and ½Zn3 AlðOHÞ8 þ mH2 O vs the reaction potentials for each step of CO2
2 ðCO3 Þ
2 ðCO3 Þ
reduction to methanol/CO and H2 oxidation to protons.

Thus, the selectivity of CO2 photoreduction was not controlled by theenergetics of reaction steps in the viewpoint of electron ﬂow[36,45,50–52], but the adsorption of CO2 and the stability of the reduced species, as probed by Zn, Cu, and Ga K-edge XANES spectra(Fig. 5), should be critical.

5. Conclusions2
LDH compounds formulated as ½Zn3x Cux MIII ðOHÞ8 þ
2 ðCO3 Þ 
þIII
mH2 O, where x was within the range 0–1.5 and M was Al or GaIII,
were synthesized and an ordered layer structure consisting of octahedral Zn, Cu, Al, and/or Ga sites was conﬁrmed. Under UV–visible
light illumination, these LDH compounds synthesized methanol
2
starting from CO2 + H2 for the ﬁrst time. ½Zn3 AlðOHÞ8 þ
2 ðCO3 Þ 
mH2 O was the most active, producing CO with 94 mol% selectivity2(0.16% conversion). ½Zn1:5 Cu1:5 GaðOHÞ8 þ mH2 O was the
2 ðCO3 Þmost selective for producing methanol (68 mol%) at 0.03% conversion. The photocatalytic rates can be enhanced 45% by increasing
the light intensity by 2.5 times in the comparison of the whole
wavelength region. Catalytic methanol formation was conﬁrmed
on the basis of blank tests performed in the absence of a catalyst
under UV–visible light, in the absence of any light with a catalyst,
in H2 or CO2 gas with a catalyst under UV–visible light, dependence
on cutoff wavelength, and suspended catalyst in deionized water
under UV–visible light. The inclusion of Cu sites in the LDH layers
improved the methanol selectivity. The bandgap (3.5–5.7 eV) was
wide enough to proceed with the reaction steps CO2 þ 6Hþ þ
6e ! CH3 OH þ H2 O and CO2 þ 2Hþ þ 2e ! COþH2 O at relatively negative reaction potentials. The key factor for determining
the selectivity to methanol vs CO was suggested to be the bindingof CO2 at the Cu sites as hydrogen carbonate species
CuAO(AZn)AC(OH)@O and CuAO(AGa)AC(OH)@O, as demonstrated in the X-ray absorption spectra.
Supporting text to verify photocatalytic methanol formation derived from CO2: Fig. S1 depicts the photocatalytic reactor used inthis study, Fig. S2 monitors the temperature of the reaction cell

during the photocatalysis, and Fig. S3 evaluates the transmitted
light intensity through catalyst layer in the reactor.
Acknowledgments
The authors are thankful for ﬁnancial support from a Grant-inAid for Scientiﬁc Research C (2255 0117) from MEXT (2010) andfrom the Promotion Section of Technology Innovation, Daikin
Industries (2008–2009). The authors appreciate the setup of photocatalytic measurements by M. Morikawa and Y. Yoshida. The X-ray
absorption experiments were performed with the approval of the
Photon Factory Proposal Review Committee (2009G552,
2007G576). The authors thank Professor Yoshitake for the Xe arc
lamp and Professors Kaneko and Kanoh for the XRD apparatus.
Appendix A. Supplementary material
Supplementary data associated with this article can be found, inthe online version, at doi:10.1016/j.jcat.2011.01.004.
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