Introduction
In this study, the time course of m/z = 17 at the retention timeof methane was monitored for 13CH4, which was converted from
13
CO2. Moreover, the time course of m/z = 16 at the same
retention time was monitored for 12CH4, which was converted
from mostly 12CO2 by subtracting the 13CH3+ fragment derived
from 13CH4 (Figure S1).

Figure S1. (A) Mass spectrum at 10’17” and (B) mass
13
chromatogram at m/z = 17 sampled for 6.0 h under CO2 (2.3 kPa),
H2 (21.7 kPa), and UV–visible light using the Ni (10 wt%)–ZrO2Reduced photocatalyst.

Experimental
ZrO2 [0.500 g; JRC-ZRO-3, Catalysis Society of Japan; major
monoclinic and minor tetragonal phase, specific surface area(SA) 94.4 m2 g−1] powder and NiII nitrate hexahydrate (>99.9%,
Wako Pure Chemical, Japan; 0.129–0.995 g) were added towater (80 mL, <0.055 µS cm−1) supplied by an RFU424TA
system (Advantec, Japan). The mixture was agitated using
ultrasound (430 W, 38 kHz) for 20 min and magnetically stirred
at 900 rpm for 1 h. Sodium borohydride (>95%, Wako Pure
Chemical; 0.129–1.04 g) dissolved in water (20 mL) was added
and magnetically stirred at 900 rpm. Then, the suspension was
filtered using a polytetrafluoroethene-based membrane filter(Omnipore JVWP04700, Millipore, Burlington, MA, USA; poresize 0.1 µm) and washed with water (250 mL). The resultant
powder was dried overnight at 373 K, and is denoted as Ni–ZrO2.
The loading of Ni was between 5.0 and 30 wt%. The 5.0 wt% Ni
sample was light green, and the 10–30 wt% Ni samples were
grayish-black.
ZrO2 (0.500 g) and Ni(NO3)2·6H2O (0.260 g) were mixed inwater, and the water was distilled at 353 K. The resultant
powder was dried at 373 K and calcined in air at 723 K for 2 h.
The obtained powder is denoted as NiO–ZrO2, with a Ni loadingof 10 wt% and gray color.
For comparison, SiO2 (mesostructured MCM-41, Sigma–
Aldrich, St. Louis, MO, USA; specific SA 1000 m2 g−1) was used
as a support for Ni using a preparation procedure similar to thatused for Ni–ZrO2. The obtained gray powder is denoted as Ni–
SiO2, with a Ni loading of 10 wt%.
The ZrO2, Ni–ZrO2, NiO–ZrO2, or Ni–SiO2 sample (each 20mg) was placed in a quartz U-tube (internal volume 46.0 mL)
connected to a Pyrex glass circulation system (volume 206.1

S1mL)[1] and treated in vacuum for 1 h using both rotary anddiffusion pumps (10−6 Pa). Then, 20 kPa of H2 (purity >99.99%)
was introduced into the sample, increased to 723 K at a ramping
rate of 15 K min−1, and maintained at the temperature for 10 min.
Thus-treated samples are denoted as ZrO2-Reduced, Ni–ZrO2Reduced, NiO–ZrO2-Reduced, or Ni–SiO2-Reduced withbrownish-black color, although ZrO2-Reduced was white.
The photocatalytic uptake/exchange/conversion tests of 13CO2
were conducted using 20 mg of the catalyst samples. For 13CO2
photo-uptake/exchange tests, 13CO2 (0.67 kPa; 13C 99.0%, 17O
0.1%, 18O 0.7%, purity >99.9%, Cambridge Isotope Laboratories,
Inc., Tewksbury, MA, USA) was used. For 13CO2 photoreduction
tests, 13CO2 (2.3 kPa) and H2 (21.7 kPa; purity >99.99%) wereused. For comparisons, photocatalytic tests under 13CO2 (4.6–
9.2 kPa) and H2 (43.4–86.8 kPa) were also performed.
Separately, 13CO2 (2.3 kPa) and H2O (2.3 kPa) were introducedfor 13CO2 photoconversion tests. Furthermore, 13CO2 (2.3 kPa),
H2O/D2O (2.3 kPa, D 99.9%, chemical purity >99.5%,
Cambridge Isotope Lab), and H2 (21.7 kPa) were introduced for13
CO2 photoconversion tests.
For these kinetic tests, the quartz reactor containing thecatalyst was irradiated with UV–visible light from both top andbottom using a 500 W Xe arc lamp (SX-UID502XAM, Ushio,
Japan) through a Y-shaped quartz fiber light guide (Optel, Tokyo,
Japan; Model 1.2S15-1000F-1Q7-SP-RX; 40-cm-long fiber and
80-cm-long branches). The distance between the fiber light exit
(Φ = 5 mm) and the photocatalyst was 20 mm. The light intensity
at the photocatalyst center was 142 mW cm−2. For comparison,
13
CO2 photoreduction tests were performed under light of 186
mW cm−2 in the pressure dependence tests (Table 1c’, c”, c’’’), in
the presence/absence of external heat (393 K; Figure 7), andirradiated by full light but in water bath at 295 K (Table 1c””””).
The light intensity distribution of the Xe arc lamp on wavelength
was measured using a spectroradiometer (Model USR45DA,
Ushio, Japan) placed 20 mm from the UV–visible light source(Figure S2).

Figure S2. The intensity distribution of light from Xe arc lamp used
for kinetic tests in this study as a function of wavelength.

In-profile kinetic data were collected as a function of the light’s
excitation wavelength by inserting a sharp-cut filter at each fiber
light exit. UV32, Y52, and WR715 [2.0, 2.0, and 2.5 mm thick;Hoya, Japan] filters were used to pass light with wavelengths of>320 nm, >520 nm, and >715 nm, respectively. For comparison,a reaction test under dark conditions was conducted bycompletely wrapping the reactor in Al foil at 298, 373, 393, and
413 K. Furthermore, comparison kinetic test irradiated by fulllight but the quartz reactor was in quartz water bath controlled at
295 K was also performed. The kinetic tests irradiated underlight at 393 K using 13CO2, H2, (and H2O) were performed using
a quartz fiber light guide and the reactor set in cylindrical heater.
A packed column of polyethene glycol-6000/Flusin P support
column (3 m length, 3 mm internal diameter; GL Sciences, Inc.,
Japan) for 13CO2 photouptake/exchange tests and that of 13X-S
molecular sieves (3 m length, 3 mm internal diameter; GL
Sciences, Inc.) for 13CO2 photoreduction/conversion tests wereused for online gas chromatography-mass spectrometry (GCMS) analyses (Model JMS-Q1050GC, JEOL, Tokyo, Japan).[1–3]Helium (0.40 MPa, purity >99.9999%) was used as the carriergas. A sampling loop (4.6 mL) composed of a Pyrex glass
system was maintained under vacuum using rotary and diffusion
pumps (10−6 Pa) and connected to GC-MS through deactivatedfused silica tubes (No. 160-2845-10, Agilent, Santa Clara, CA,
USA; 1.5 m, internal diameter 250 µm), which were maintained
at 393 K during the analysis to avoid gas adsorption.
The surface species were monitored with a single-beam
Fourier transform infrared (FTIR) instrument (JASCO, Tokyo,
Japan; Model FT/IR-4200) equipped with a mercury–cadmium–tellurium-M detector at a constant temperature of 77.4 K.[1,4,5] A
self-supporting disk (Φ = 20 mm) of Ni (10 wt%)–ZrO2-Reduced(10 mg) diluted with ZrO2 (90 mg) was prepared in a Pyrex tube
and transferred into a quartz photoreaction IR cell equipped with
NaCl windows on both sides using a glove box (model MINI400RS; Unico, Tsukuba, Japan) filled with argon (purity>99.998%). The cell was connected to the Pyrex glass
circulation system, as well as to the GC-MS, to enable
simultaneous monitoring of the surface species using FTIR andisotope distribution in products using GC-MS.[1,2] The
photocatalyst disk was evacuated (10−6 Pa) at 295 K for 2 h prior
to the measurements.
The in-situ FTIR measurements were performed at 295 K in arange of 4000–650 cm−1. The sample disk was irradiated with
UV–visible light by a 500 W Xe arc lamp using a quartz fiber
light guide. The distance between the fiber light exit and sample
disk was 46 mm. The light intensity at the sample center was 90
mW cm−2. The spectrometer’s energy resolution was 1 cm−1, andthe data accumulation included 256 scans (~2 s per scan). Fromeach spectrum, we subtracted the spectrum of Ni (10 wt%)–ZrO2
just reduced in H2 at 723 K.
UV−visible spectra were recorded on a double-beam model V650 spectrophotometer using D2 and halogen lamps below andabove 340 nm, equipped with a photomultiplier tube and an
integrated ISV-469 sphere (JASCO, Tokyo, Japan) for diffusereflectance detection within the wavelength range of 200–800
nm.[1,2] The data were transformed using the Kubelka–Munkfunction. A polytetrafluoroethene plate was used as thereference.
Transmission electron microscopy (TEM) was performed
using a JEM-2100F (JEOL) equipped with a field-emission gun(acceleration voltage of 200 kV) at the Center for Analytical
Instrumentation. The samples were mounted on a Cu mesh (250meshes per inch) coated with carbon and a copolymer film ofpoly(vinyl alcohol) and formaldehyde (Formvar, Monsanto, St.
Louis, MO, USA). High-angle annular dark-field (HAADF)
scanning TEM (STEM) and high-resolution (HR) TEM images
were observed using the JEM-2100F model.[1] The chemicalcompositions and elemental distributions were analyzed using
energy-dispersive spectra with a Si (Li) detector equipped in the
TEM.[4]

X-ray diffraction (XRD) patterns were observed using a D8
ADVANCE diffractometer (Bruker, Billerica, MA, USA) at theCenter for Analytical Instrumentation at a Bragg angle (θB) of 2θB= 20°–60° with a scan step of 0.02° and scan rate of 1 s per step.
The measurements were performed at 40 kV and 40 mA using
Cu Kα emission (wavelength λ = 0.15419 nm) and a Ni filter. The
crystal sizes (t) were estimated using the following Scherrer
equation:

t=

0.9 λ
Peak width × cosθ B(S1)
The absorption–fluorescence spectra were recorded on FP8600 (JASCO; Chiba Iodine Resource Innovation Center) using
a 150 W Xe arc lamp (UXL-159, Ushio) equipped with aphotomultiplier tube for excitation at 200–300 nm within afluorescence range of 300–800 nm. The incident excitation lightfrom Xe lamp was monitored by Si photodiode, and monitored
fluorescence light emitted from sample was normalized basedon the incident light intensity at each wavelength. The
photocatalyst powder (2.0 mg) was mixed with purified water(3.0 mL) and ultrasonicated (430 W, 38 kHz) for 30 min. All
spectra were recorded for the suspensions in a quartz cell at
295 K.
The Ni K-edge extended X-ray absorption fine structure(EXAFS) spectra were measured at 295 K in the transmission
mode at the Photon Factory, High Energy Accelerator Research
Organization (KEK, Tsukuba, Japan), on 9C and 12C beamlines.
A Si(1 1 1) double-crystal monochromator and a Rh-coatedfocusing bent cylindrical mirror were inserted into the X-ray
beam path. A piezotransducer was used to detune the X-ray to
two-thirds of the maximum intensity to suppress the higher
harmonics. The Ni K-edge absorption energy was calibrated at
8331.65 eV[6] using the spectrum of Ni metal (5.0 µm thick).
The Ni, Ag, and Mo K-edge EXAFS spectra were measured
for Ni, Ag, and Mo foils at 300–400 K on 9C beamline and at the
Photon Factory Advanced Ring, KEK, on NW10A beamline
equipped with a Si(3 1 1) double-crystal monochromator, a Ptcoated focusing bent cylindrical mirror, and a piezotransducer.[1]Ni (10 wt%)−ZrO2-Reduced powder (120 mg) was prepared ina Pyrex glass U-tube, and transferred into a Pyrex glass XAFS
cell (Φ = 20 mm, t = 2.0 mm) equipped with polyethene
terephthalate (PET) film windows (Toyobo Film Solutions, Japan,
G2; 38 µm thick) for both UV–visible light and X-ray transmission.
The cell was directly connected to the U-tube, and the sample
did not contact air throughout the procedure. Then, the XAFS
cell was filled with the reactant gas. The sample was thenirradiated with UV–visible light from the Xe arc lamp through aquartz fiber light guide and the PET film window at the beamline.
X-rays were perpendicularly transmitted from the disk, whereas
the incident angles of UV–visible light were 45° and − 135°
relative to the X-rays.[1–3,7] The distance between the light exit of
the quartz fiber light guide and the sample was 50 mm.
The obtained Ni K-edge EXAFS data were analyzed using
XDAP,[8] and the pre-edge background was approximated with amodified Victoreen function:

C2 C1+ + C0E2 E(S2)where E is the photon energy. The background for post-edge
oscillation, µx, was approximated with a smoothing spline
function and calculated for a particular number of datapoints:

S2

Data Points∑
i=12B). The equilibrated 12CO2 ratio during the photocatalytic
reduction tests conducted using 2.3 kPa of 13CO2 (193 µmol, 13C
99.0%) and Ni (10 wt%)–ZrO2-Reduced (20 mg) was(1.93+ 0.64) /193 = 1.3 mol% .
In contrast, based on ref. 1, the equilibrated 12CO2 ratio was
7.3 mol% for the total CO2 gas using 0.67 kPa of 13CO2 and 100
mg of ZrO2, not H2-reduced. The equilibrated 13CO2 and 12CO2
on the surface were 19.8 µmol (Figure 2B) and 19.8× 0.073 =1.4µmol over ZrO2. Thus, the amount of 12CO2 on surface before
the photouptake/exchange test conducted on ZrO2 (100 mg) was
1.4 + 2.3 = 3.7 µmol. This corresponds to 0.74 µmol per 20 mg
ZrO2, essentially equivalent to 0.79 µmol for ZrO2-Reduced.
Thus, the H2-pretreatment of ZrO2 at 723 K affected exclusively
the amount of chemisorbed CO2 over ZrO2-Reduced (0.66 µmol;
Figure 2A) versus ZrO2 (0.45 µmol per 20 mg ZrO2).[1] Thisdifference is the difference of O vacancy sites over ZrO2
(Scheme 1A-c), and accounts for the increase of CO formation
rate by 1.16 times (Tables 1a and S1a).
In the 13CO2 photoreduction tests, the 13C ratio in the formed
C-containing products was 2.1–4.3 mol% when no filter or a filter
to transmit λ > 320 nm was used (Table 1c, c””), and gradually
increased to 5.7, 29, and 34 mol% when a filter to transmit λ >520 and 715 nm was used and under dark conditions,
respectively (Table 1c’’’’’, c”””, c’’’’’’’). The exchange of free
13
CO2 with 12CO2 chemisorbed from air became substantiallyslower under light of longer wavelength, and under dark
conditions, and the contribution of 12CO2 originating from the air(Scheme 1A-c) became not negligible.
Using NiO (10 wt% Ni)–ZrO2-Reduced, the formation rate of
the total C-containing compounds increased by a factor of 1200
after H2 pretreatment (Table S1c and Figure S5B), which was
66% of that using Ni (10 wt%)–ZrO2-Reduced (Table 1c). The
difference could be owing to the difference in particle size of Ni0
formed either by NaBH4 or by heating in H2 at 723 K.

(µ xi − background i )2
≤ smoothing factor
exp (−0.075ki2 )(S3)in which k is the angular photoelectron wavenumber.
Multiple-shell curve-fit analyses were performed on the
Fourier-filtered k3-weighted EXAFS data in k– and R-space (R isinteratomic distance) based on the plane-wave approximation(eq. 7 in main text) using an XDAP code,[8] in which the empirical
amplitude was extracted from the EXAFS data for the Ni metal
foil and NiO powder. The R values for the Ni–Ni interatomic pair
were set to 0.24917 nm with an N value of 12 for Ni metal (a =0.35238 nm),[9] and those for the Ni–O and Ni–Ni interatomic
pair were set to 0.2088 nm with an N value of 6 and to 0.2953
nm with an N value of 12 for NiO (a = 0.4176 nm).[10] We
assumed that the many-body reduction factor, S02, was identicalfor both sample and reference.

Results and Discussion
1. 13CO2-EXCHANGE REACTION AND PHOTOREDUCTION
ZrO2 not pretreated with H2 only formed CO (Table S1a);
however, the CO formation rate increased by 1.16 times after H2treatment (Figure 1A and Table 1a).
Moreover, the Ni (10 wt%)–ZrO2 photocatalyst was tested forCO2 photoreduction (Table S1b). The total formation rate of Ccontaining compounds, mostly methane, was only 0.97% of thatusing Ni (10 wt%)–ZrO2-Reduced (Figure 1C). However, theinitial rate increased by 3.1 times at 42–47 h. This is because
the oxidized NiII sites of Ni (10 wt%)–ZrO2 gradually transformed
into Ni0 under H2 and UV–visible light.
Based on the uptake/exchange tests conducted for 13CO2(0.67 kPa; Figure 2A), the equilibrated 12CO2 ratio was 2.4 mol%in total CO2 gas, using 20 mg of ZrO2-Reduced. Therefore, theequilibrated 13CO2 and 12CO2 on the surface were 9.8 µmol
(Figure 2A) and 9.8 × 0.024 = 0.23 µmol for ZrO2-Reduced. Thus,
the amount of 12CO2 that existed on the surface before thephoto-uptake/exchange test conducted on ZrO2-Reduced (20mg) was 0.23 − 0.098 + 0.66 = 0.79 µmol (Figure 2A). The
equilibrated 12CO2 ratio during the photocatalytic reduction tests
conducted using 2.3 kPa of 13CO2 (193 µmol, 13C 99.0%) andZrO2-Reduced (20 mg) was (1.93+ 0.79 ) /193 = 1.4 mol% .
Similarly, based on Figure 2B, the equilibrated 12CO2 ratio was
2.1 mol% in total CO2 gas using 0.67 kPa of 13CO2 and 20 mg ofNi (10 wt%)–ZrO2-Reduced. The equilibrated 13CO2 and 12CO2
on the surface were 5.9 µmol (Figure 2B) and5.9 × 0.021 = 0.12 µmol, respectively, over Ni (10 wt%)–ZrO2Reduced. Thus, the amount of 12CO2 on the surface before thephotouptake/exchange test conducted on Ni (10 wt%)–ZrO2Reduced (20 mg) was 0.12 – 0.02 + 0.54 = 0.64 µmol (Figure

Table S1. Kinetic data on photoconversion ofvisible light

13

13

a
b
c

13

CO2 (2.3 kPa) using H2 (21.7 kPa) and the ZrO2-based photocatalysts (0.020 g) under UV–Formation rate (µmol-C h

Entry

12

Figure S3. Formation rate dependences of CH4, CH4, CO, and12
CO on the Ni content in ZrO2 in the photocatalytic test exposed to13
CO2 (2.3 kPa) and H2 (21.7 kPa). The catalysts were previously
reduced in H2 at 723 K.

Catalyst
ZrO2
Ni (10 wt%)–ZrO2
NiO (10 wt% Ni)–ZrO2-Reduced

13

CO

1.1
0.11
0.020

12

CO

0.64
<0.002<0.002

S3

13

CH4

<0.002
2.6
220

12

−1

−1

gcat )

CH4<0.0020.60
1.6

Σ C-containing
compounds
1.7
3.3
220

12

C-product
ratio (mol%)
37
18
0.73

2. UV–VISIBLE ABSORPTION SPECTRAThe UV–visible spectrum for ZrO2 contained edge absorption
at 248 nm (Figure S6). The spectrum for Ni (5.0 wt%)–ZrO2 wassimilar to that for ZrO2. However, a weak shoulder (250–290 nm)
to the absorption edge and a tiny peak centered at 404 nm
appeared. The spectrum suggested that the Ni nanoparticles
formed through a liquid-phase reduction, partially re-oxidized.
The Ni (5.0 wt%)–ZrO2-Reduced sample showed extensive
absorption in the entire wavenumber region of 200–800 nm(Figure S6), suggesting the presence of Ni0.
Unlike Ni (5.0 wt%)–ZrO2, Ni (10 wt%)–ZrO2 exhibited
substantial absorption over the whole wavelength region in
Figure S6. The difference suggested that the Ni0 nanoparticles
formed through a liquid-phase reduction were relatively stable inair probably owing to the difference in pH during the liquid-phasereduction and the resultant particle size of Ni.

13

Figure S6. Diffuse-reflectance UV–visible absorption spectra for
ZrO2, Ni (5.0 wt%)–ZrO2, Ni (5.0 wt%)–ZrO2-Reduced, Ni (10 wt%)–
ZrO2, and Ni (10 wt%)–ZrO2-Reduced samples.

12

Figure S4. Time course of CH4 and CH4 formation during (A–C)
13
photocatalytic and (D) catalytic tests exposed to CO2 (2.3 kPa) andH2 (21.7 kPa) using Ni (10 wt%)–ZrO2-Reduced. (A–C) Catalyst wasirradiated under Xe arc lamp filtered at wavelengths (A) λ > 320 nm,(B) λ > 520 nm, and (C) λ > 715 nm and (D) catalyst under dark
conditions at 298 K. The amount of catalyst used was 0.020 g.

13

12

13

12

Figure S5. Time course of CO, CO, CH4, and CH4 formation
13
during the photocatalytic test exposed to CO2 (2.3 kPa) and H2(21.7 kPa) using (A) Ni (10 wt%)–SiO2-Reduced and (B) NiO (10wt% Ni)–ZrO2-Reduced. The amount of catalyst used was 0.020 g.

S4

3. FTIR MONITORING UNDER 13CO2 AND H2

13

subtracted from each data. (A) Under CO2 (2.3 kPa) and H2 (21.7
13kPa) for 1 h. (B) Under CO2, H2, and UV–visible light for 2 h. (C)13
Under CO2, H2, and darkness for 2 h. (D) Under vacuum for 1 min,and then UV–visible light for 24 h.

4. TEM OBSERVATION

Figure S9. (A) TEM and (B) HAADF-STEM images observed for Ni(10 wt%)–ZrO2-Reduced sample. The arrows indicate Ni
nanocrystals dispersed on ZrO2 crystals.

5. XRDIn the XRD pattern for the incipient Ni–ZrO2-Reduced samples(Figure S10), the diffraction peaks at 24.3°, 28.1°, 31.4°, 34.3°,
35.3°, 38.7°, 41.0°, 45.1°, 50.3°, 54.2°, and 55.7°, assignable to(1 1 0), ( 1 1 1 ), (1 1 1), (0 2 0), (2 0 0), (1 2 0), (1 0 2), ( 2 0 2 ),( 1 2 2 ), (3 0 0), and (1 3 0) planes of monoclinic crystal of ZrO2,
respectively, were observed. When the Ni content was >10 wt%,
two new peaks appeared at 44.4° and 51.8°, assigned to (1 1 1)
and (2 0 0) planes of Ni metal, respectively (Figure S10).

Figure S7. FTIR spectra of Ni (10 wt%)–ZrO2-Reduced (10 mg) in−1−1
wavenumber regions of 3800–3500 cm (left) and 1800–1100 cm
13(right). (A) Under CO2 (2.3 kPa) and H2 (21.7 kPa) for 1 h. (B)13
13
Under CO2, H2, and UV–visible light for 20 h. (C) Under CO2, H2,
and darkness for 2 h. (D) Under vacuum for 1 min, and then UV–
visible light for 20 h. The spectrum for Ni (10 wt%)–ZrO2 justreduced in H2 at 723 K was subtracted from each data.

Figure S10. XRD patterns of Ni–ZrO2-Reduced. The Ni content was0, 5.0, 10, 15, and 30 wt%.

6. FLUORESCENCE SPECTRA

Ni–ZrO2-Reduced samples were all brownish-black, where Ni0
absorbed most part of UV–visible light while ZrO2 mostly
absorbed UV light of λ < 248 nm (Figures S2 and S6). The
intensity at 370 nm decreased by 72% and 90% with theaddition of 5.0 and 30 wt% of Ni, respectively (Figure 5A), as theNi species shielded ZrO2 from the incident light and also

−1

Figure S8. FTIR spectra of ZrO2 (50 mg) for 3000–2700 cm (left)−1
and 2200–1900 cm (right). The spectrum for fresh ZrO2 was

S5accepted and isolated the photoelectrons from ZrO2 (chargetransfer), to suppress the competing fluorescence emission step.
Figure 5B shows the weaker fluorescence above 248 nm
caused by the low concentration of O vacancy level between the
ZrO2 bandgaps. The fluorescence with the excitation at 248 nm
was clearly confirmed for ZrO2 (Figure S11B), while it was quite
suppressed due to the charge separation with Ni nanocrystals
for Ni (10 wt%)–ZrO2-Reduced (Figure S11B). Using Xe arc
lamp and quartz cell both for kinetic tests (Figure 1) andfluorescence spectroscopy (Figures 5 and S11), ZrO2 could be
excited via band-gap excitation near the foot of wavelength
distribution for Xe lamp (λ < 248 nm; Figure S2) and also theelectronic transition related to the O vacancy of ZrO2 (λ > 248
nm; Figures 5B and S11B).

using Ni (5.0 wt%)–ZrO2-Reduced (Table 1b, c) because of (i)the increase of active Ni sites, (ii) Ni0 site temperature differenceby >10 K, (iii) relatively stable Ni0 state @10 wt% in contrast to
the oxidized NiII sites in air @5.0 wt% (Figure S6), and (iv) thequantum size effects of the Ni particles (mean 1.7 nm@10 wt%
versus 1.1 nm @5.0 wt%). Conversely, when the Ni amountbecame more than 10 wt%, the Ni nanocrystals predominantly
covered ZrO2 surface (Figures 5A and S11). By the effects, thefirst step from CO2 to CO over ZrO2 surface was significantly
suppressed and methane formation rate drastically dropped(Figure S3).
The temperature dependence of experimental σ values formetal foils as examples agreed well with that by theory.[11] We
confirmed the deviation between experiment and theory was <8
K for metal foils in the range of 300–400 K (Figure S13), andapplied to the local temperature evaluation of mean 1.7 nm Ni
nanocrystals based on correlated Debye model.[11,12] Taking
systematic error to experimentally determine σ value led to theerror <20 K (Figure 6D) into account, total evaluation errors asroot mean square would be <22 K (Scheme 1C). Anyway, tomonitor the local temperature of 1.1 or 1.7 nm Ni nanocrystals
dispersed on ZrO2 by thermocouple is impossible.
When the Ni (10 wt%)–ZrO2-Reduced was irradiated underlight of λ > 320 nm (Figure S14, left), the trend of changes in N
and σ values was similar to that of the data irradiated under fulllight (Figure 6C, D). The temperature increased to 378–387 K(Figure S14D, left). The mean temperature was lower by <10 K
than that irradiated by full light (Table S2b, c). In contrast, thetemperature reached during irradiation of light of λ > 715 nm was
322–334 K (Figure S14D, right, and Table S2d).
Figure S15 shows the monitoring of Ni sites in the Ni (10wt%)–ZrO2-Reduced photocatalysts under CO2 and moisture
where a reversible temperature change is observed. Table S2elists the conditions and temperature obtained.
The increase in temperature for the Ni, Ag, and Au sites,combined with ZrO2 irradiated by light, is summarized in Table
S2. The reversible temperature rise was primarily controlled by
the incident light wavelengths for Ni (b–d), Ag (h–j), and Au (k, l).
The increase in Ni and Ag ratio in the composites slightly
increased the temperature. The temperature rise was in theorder Ni > Ag > Au (Table S2b, f, k), and did not correlate to themolar heat capacity Cpo value of metals. Brownish–black Ni
nanoparticles of 1–2 nm absorbed visible light more effectively
than the localized surface plasmon resonance[13] of Ag and Au
nanoparticles (mean 3–4 nm), and enabled Ni to have atemperature of 394 K (Figure 6D).

Figure S11. Fluorescence spectra for (A) Ni (0–30 wt%)–ZrO2 withthe excitation at 200 nm and (B) ZrO2 and Ni (10 wt%)–ZrO2Reduced with the excitation at 248 nm; 2 mg of the sample powder
was suspended in 3 mL of water.

7. Ni K-EDGE EXAFS
The intensity of the Ni–Ni shell at 0.21 nm (phase shiftuncorrected) for Ni (5.0 wt%)–ZrO2-Reduced was similar to thatfor Ni metal, and decreased by 44% when exposed to air for 1
week, suggesting that Ni0 core remained inside the NiIIO shell.
Moreover, the change in the Ni K-edge EXAFS for Ni (5.0wt%)–ZrO2-Reduced photocatalyst irradiated by UV–visible light
was monitored (Figure S12) in comparison to that for Ni (10wt%)–ZrO2-Reduced photocatalyst (Figure 6). The σdisorder value
was evaluated as 0.00292 nm. Then, the temperature at the Nisite based on the σ values was evaluated. Based on the meanN(Ni–Ni) value (7.1) for Ni (5.0 wt%)–ZrO2-Reduced, the mean
particle size of 1.1 nm and dispersion of 0.77 were evaluated.
The temperature rise (372–388 K) using Ni (5.0 wt%)–ZrO2Reduced decreased by >10 K compared to 389–394 K using Ni(10 wt%)–ZrO2-Reduced (Table S2a, b).
The reason why methane formation drew a curve peaked at
Ni 10 wt% in photocatalytic tests exposed to 13CO2 and H2(Figure S3) is considered. The rate using Ni (10 wt%)–ZrO2Reduced was higher by a factor of 16 times compared to one

S6

Figure S13. The temperature dependence of Debye–Waller factor
for Ag, Mo, and Ni metal foils (a) monitored by thermocouple in anelectric furnace at constant temperature (square, ◻; cross, ×; star, ⭐︎)and (b) generated by the correlated Debye model using a FEFF8.4
code (circle, ○; diamond, ♢; triangle, △).

Figure S12. Ni K-edge EXAFS for Ni (5.0 wt%)–ZrO2-Reduced
under CO2 (2.3 kPa) and H2 (21.7 kPa). (A, B) Time-course change
3of Fourier transform of k χ irradiated by light from Xe arc lamp (A)and under dark conditions (B). (C, D) N (C) and σ values (D)obtained by curve-fit analyses for the Ni–Ni shell at 0.21 nm inpanels A and B.
Figure S14. Ni K-edge EXAFS for Ni (10 wt%)–ZrO2-Reduced under
CO2 (2.3 kPa) and H2 (21.7 kPa). (A, B) Time-course change of3
Fourier transform of k χ irradiated by light from Xe arc lamp filteredat λ > 320 nm (Left panel) and λ > 715 nm (Right panel) (A) andunder dark (B). (C, D) N (C) and σ values (D) obtained by curve-fit
analyses for the Ni–Ni shell at 0.21 nm in panels A and B.

S7

Table S2. Temperature of metal (Ni, Ag, and Au)–ZrO2 photocatalyst samples monitored by EXAFS and related factors
Metal
SampleoSampleLight
TheatedCp (J
Entry
Purged gas
Content
θDebye amount−1
color
irradiated(K)
Element
K
(mg)
(wt%)
(K)
−1
g )
a
5.0
Full light372–388b
Full light389–394
CO2 (2.3 kPa),
H2 (21.7 kPa)c
Brownishλ > 320 nm
378–387
Ni
0.444
450
120
black
10
d
λ > 715 nm
322–334
CO2 (2.3 kPa),e378–384
H2O (2.3 kPa)
CO2 (2.3 kPa),f
325–392
Full lightH2 (21.7 kPa)
3.0
g
Argon
348–367
Dark
Ag
0.235
225
125
h
yellow
351–363i
5.0
λ > 320 nm
331–365
CO2 (2.3 kPa),j
λ > 580 nm
307–329
H2 (21.7 kPa)kFull light
314–321
l
Au
5.0
0.129
165
80
Wine redλ > 715 nm
299–300m
Argon
Full light
320–324

Ref

This
work

1

2

Figure S15. Ni K-edge EXAFS for Ni (10 wt%)–ZrO2-Reduced
under CO2 (2.3 kPa) and H2O (2.3 kPa). (A, B) Time-course change
3in Fourier transform of k χ irradiated by light from Xe arc lamp (A)and under dark (B). (C, D) Coordination number N (C) and Debye–
Waller factor σ (D) obtained by curve-fit analyses for the Ni–Ni shellat 0.21 nm in panels A and B.

8. 13CO2 PHOTOCONVERSION IN H2 AND H2O/D2O

13

12

Figure S16. Time course of CH4, CH4, and H2 formation during
13
the photocatalytic test with (A) CO2 (2.3 kPa) and H2O (2.3 kPa),
13(B) CO2 (2.3 kPa), H2 (21.7 kPa), and H2O (2.3 kPa), and (C)13
CO2 (2.3 kPa), H2 (21.7 kPa), and D2O (2.3 kPa) using Ni (10wt%)–ZrO2-Reduced (0.020 g).

S8

Finally, we tested CO2 photoconversion using water and Ni(10 wt%)–ZrO2-Reduced catalyst (eq. 3 of main text). H2 wasquickly formed in the initial 1 h by the stoichiometric oxidation of
Ni0 surface by H2O (Figure S16A, left), but the subsequent,steady photocatalytic reduction rate of water was 3.8 µmol h−1
gcat−1 (Table 2a). In contrast, CH4 was formed constantly fromthe beginning, but the preferable 12CH4 formation switched toboth 13CH4 and 12CH4 formation at ~10 h of the reaction. The
total formation rate of the C-containing products, mostly
methane, was only 0.23% of that obtained using H2 (Tables 1cand 2a).

Based on Table S3, the following correlations were obtained
for the kinetic tests in 13CO2, H2, and D2O.

e=Ed≃Dc ≃ C − ⎛⎜ 1 D + E ⎞⎟ f⎝4⎠⎛
b ≃ B − ⎜ 1 C + 3 D ⎞⎟ f + ⎛⎜ 1 D + 1 E ⎞⎟ f 2⎝2
4 ⎠ ⎝8
2 ⎠
⎛
⎞
⎛
3
1
3
a ≃ A − ⎜ B + C ⎟ f + ⎜ C + 11 D + 1 E ⎞⎟ f 2 − ⎛⎜ 3 D + 3 E ⎞⎟ f 3
⎝4
⎝8
⎝ 32
2 ⎠
16
2 ⎠
8 ⎠

Table S3. Mass number analysis for the product methane in the photocatalytic tests under
13
13
13
*1,*2
CH2D2, CHD3, and CD4 with the molar amount of a, b, c , d, and e, respectively.
m/z
15
16
17
18
12

CH4

1
fa
100

13

CH4

1
fb
400

13

13

CH3D

CH2D2

13

fa +

3
fb
400

1
1
fb +
fc
4
200

a+

19

20

13

CH4,

13

CH3D,

21

1
b
100

3
1
fb +
fc
4
200

1
3
fc +
fd
2
400

b+

1
c
100

1
1
fc +
fd
2
400
3
1
fd +
fe
4
100

CD4

Σ

CO2, H2, and D2O to produce

1
a
100

CHD3

13

13

c+

1
d
100

1
fd
4

d+

1
e
100

fe

=A

=B

=C

e

=D

=E

*1

The f value is the fragment ratio for natural methane, i.e. intensity ratio of peak at m/z = 15 to that at m/z = 16. The f value was assumed
13
13
to be equal for CH4. The observed f value for CH4: 0.691.
*2A, B, C, D, and E values are the intensity in mass chromatogram at m/z = 17, 18, 19, 20, and 21, respectively.
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