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Cite This: J. Am. Chem. Soc. 2019, 141, 6292−6301

Dual Photocatalytic Roles of Light: Charge Separation at the Band
Gap and Heat via Localized Surface Plasmon Resonance To Convert
CO2 into CO over Silver−Zirconium Oxide
Hongwei Zhang,† Takaomi Itoi,‡ Takehisa Konishi,† and Yasuo Izumi*,††Department of Chemistry, Graduate School of Science, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522, Japan
Department of Mechanical Engineering, Graduate School of Engineering, Chiba University, Yayoi 1-33, Inage-ku, Chiba 263-8522,
Japan

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‡

S Supporting Information*

ABSTRACT: Conﬁrmation of 13CO2 photoconversion into a 13C-product is
crucial to produce solar fuel. However, the total reactant and charge ﬂow
during the reaction is complex; therefore, the role of light during this reactionneeds clariﬁcation. Here, we chose Ag−ZrO2 photocatalysts because
beginning from adventitious C, negligible products are formed using them.
The reactants, products, and intermediates at the surface were monitored via
gas chromatography−mass spectrometry and FTIR, whereas the temperatureof Ag was monitored via Debye−Waller factor obtained by in situ extended Xray absorption ﬁne structure. With exposure to 13CO2, H2, and UV−visible
light, 13CO selectively formed, while 8.6% of the 12CO mixed in the productdue to the formation of 12C-bicarbonate species from air that exchanged withthe 13CO2 gas-phase during a 2 h reaction. By choosing the light activation
wavelength, the CO2 photoconversion contribution ratio was charge separated
at the ZrO2 band gap (λ < 248 nm): 70%, localized at the Ag surface plasmon resonance (LSPR) (330 < λ < 580 nm): 28%, andcharacterized by a thermal energy of 295 K: 2%. LSPR at the Ag surface was converted to heat at temperatures of up to 392 K,
which provided an eﬃcient supply of activated H species to the bicarbonate species, combined with separated electrons andholes above the ZrO2, which generated CO at a rate of 0.66 μmol h−1 gcat−1 with approximately zero order kinetics.
Photoconversion of 13CO2 using moisture was also possible. Water photo-oxidation step above ZrO2 was rate-limited, and theside reactions that formed H2 above the Ag were successfully suppressed instead to produce CO via the Mg2+ addition to trap
CO2 at the surface.

1. INTRODUCTION
Numerous studies have investigated photocatalytic CO2 fuelconversion to initiate a carbon-neutral cycle, which includes
the use of solar fuel in contrast to the irreversible consumptionof fossil fuels.1,2 However, pitfalls exist in these previous
investigations that mistake methane and/or C-containing
products converted from catalyst impurities as products
converted from CO2.1,3−6 Therefore, ﬁnding an eﬀective
photocatalyst to convert CO2 (standard formation enthalpy= −393.5 kJ mol−1) using only sustainable energy while
conﬁrming labeled 13CO2 conversion, not adventitious C, isessential.1,3,6−12
Previous studies have rarely used time course monitoring forisotope labeled-13CO2 to conﬁrm CO2 conversion. This study
monitors the 13CO 2 photoreduction time course andinvestigates the 13C and 12C product origins. On the basis of
this mechanistic investigation, we observe and clarify the dual
roles that UV−visible light has for the ﬁrst time using ZrO2based photocatalysts:13−15 (i) charge separation at the bandgap (BG) and (ii) heat transformed via localized surface
plasmon resonance (LSPR) via extended X-ray absorption ﬁne© 2019 American Chemical Society

structure (EXAFS). Temperature monitoring of the Ag via
Debye−Waller factor, which was obtained with in situ EXAFS,
was direct in comparison with surface-enhanced Raman
techniques that use adsorbed probe molecules.
First, this study evaluates the rates of step 2 that follows
water photo-oxidation in step 1 below.16−18
2H 2O → O2 + 2H 2 (or 4H+ + 4e−),Δr H ° = 483.64 kJ mol−1(1)

CO2 + H 2 (or 2H+ + 2e−) → CO + H 2O,
Δr H ° = 41.16 kJ mol−1(2)

CO2 + H 2O → CO + O2 + H 2 ,Δr H ° = 524.80 kJ mol−1(3)Received: January 4, 2019
Published: March 27, 2019
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DOI: 10.1021/jacs.8b13894
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inserted in front of the photoreaction cell. Data accumulation was
between 128 and 256 scans (approximately 2 s per scan).
UV−visible spectra were recorded on a double-beam model V-650
spectrophotometer using D2 and halogen lamps below and above 340
nm equipped with a photomultiplier tube and an integrated ISV-469
sphere (JASCO) to diﬀuse reﬂectance detection within the range
from 200 to 800 nm.16,23 Data was transformed using the Kubelka−
Munk function. A formed polytetraﬂuoroethylene plate was used as areference. Absorption−ﬂuorescence spectra were recorded on model
FP-8600 (JASCO; Chiba Iodine Resource Innovation Center, Chiba
University) using 150-W Xe arc lamp equipped with a photomultipliertube within the excitation range from 200 and 300 nm andﬂuorescence range from 300 to 800 nm.
Silver K-edge EXAFS spectra were measured at 290 K intransmission mode at the Photon Factory Advanced Ring, HighEnergy Accelerator Research Organization (KEK, Tsukuba, Japan) on
the NW10A beamline.25 A Si(3 1 1) double-crystal monochromatorand a Pt-coated focusing cylindrical mirror were inserted into the pathof the X-ray beam. A Piezo transducer was used to detune the X-ray to
two-thirds of the maximum intensity to suppress higher harmonics.
The Ag K-edge absorption energy was calibrated at 25516.5 eV using
the X-ray spectrum of a Ag metal foil (40 μm thick).26 A disk (Φ = 10mm) of a Ag (3.0−5.0 wt %)−ZrO2 photocatalyst (125 mg) was set
in a Pyrex glass reactor equipped with a Kapton ﬁlm (Dupont,
Wilmington, DE, USA; 50 μm thick) for X-ray transmission and apolyethylene terephthalate (PET) ﬁlm (Teijin, Japan, G2; 50 μm
thick) for both UV−visible light and X-ray transmission ﬁlled with 2.3
kPa of CO2 and 21.7 kPa of H2. The sample was irradiated with UV−
visible light from a Xe arc lamp through the PET ﬁlm at thebeamline.18,24,27,28
The obtained Ag K-edge EXAFS data were analyzed using the
XDAP software package.29 The pre-edge background was approximated with a modiﬁed Victoreen function, i.e., C2/E2 + C1/E + C0,
where E is the photon energy. The background for postedge
oscillation, μx, was approximated with a smoothing spline function
and was calculated for a particular number of data points:

We analyzed the photocatalytic mechanism in step 2, and we
also attempted the photoconversion of 13CO2 using moisture(step 3).

2. EXPERIMENTAL METHODS
An aqueous Ag nitrate (>99.8%, Wako Pure Chemical, Japan)
solution was reduced in a liquid phase using NaBH4 (>95%, Wako
Pure Chemical) in the presence of ZrO2 (JRC-ZRO-3, Catalysis
Society of Japan; tetragonal phase, speciﬁc surface area = 94.4 m2g−1).17 We varied the sample Ag content between 0.50 and 10 weights(wt) %. The sample is denoted as Ag−ZrO2. Mg nitrate (>99.5%,
Wako Pure Chemical) was impregnated into the Ag−ZrO2 aqueous
suspension. The molar ratio of Mg:Ag was 1:1. The water was distilled
at 358 K, and the powder product was dried at 373 K for 40 h.
The photocatalyst (0.100 g) was placed in a quartz photoreactor
and evacuated at 295 K for 2 h while connected to a Pyrex glass
circulation system (206.1 mL) and both rotary and diﬀusion pumps(10−6 Pa).16 We then introduced 2.3 kPa of 13CO2 (13C 99.0%, 17O
0.1%, 18O 0.7%, chemical purity >99.9%, Cambridge Isotope
Laboratories, Inc., Tewksbury, MA, USA) and 21.7 kPa of H2(>99.99%). For comparisons, the pressure of 13CO2 and H2 wasvaried in the ranges 1.1−6.7 kPa and 21.7−66.7 kPa, respectively. The
reactor catalyst was irradiated with UV−visible light from a 500 W
xenon arc lamp (Model OPM2−502, Ushio, Japan). The distance
between the UV−visible light exit and the photocatalyst was 20 mm.
The light intensity was 90.2 mW cm−2 at the center of thephotocatalyst. The intensity distribution of the Xe arc lamp was
measured using a spectroradiometer (Model USR45DA, Ushio,
Japan) at a position 20 mm apart from the UV−visible light exit
(Figure S1, Supporting Information). A small portion of distribution
lower than λ = 248 nm exists enabling the BG excitation of ZrO2. Inproﬁle kinetic data were collected as a function of the light’s excitation
wavelength by inserting a sharp-cut ﬁlter (2.5 mm thick) at thelighthouse exit. We used the UV32 and O58 (Hoya, Japan) types topass light with λ > 320 nm and λ > 580 nm, respectively.19,20 Control
tests with exposure to 13CO2, H2, and no light were performed bycompletely wrapping the reactor with Al foil. We also performed
control tests with exposure 13CO2 gas only and UV−visible light. The
exchange reaction with 0.67 kPa of 13CO2 was also performed using asimilar procedure.
A packed column of 13X-S molecular sieves (3 m length, 3 mm
internal diameter; GL Sciences, Inc., Japan) was employed for online
gas chromatography−mass spectrometry analyses (GC−MS; Model
JMS-Q1050GC, JEOL, Japan).21−23 Helium (purity >99.9999%) wasused as the carrier gas at 0.40 MPa. The sampling loop comprised a
Pyrex glass system kept under vacuum using rotary and diﬀusion
pumps (10−6 Pa) connected to the GC−MS via 1.5 m deactivatedfused silica tubes (No. 160-2845-10, Agilent, Santa Clara, CA, USA;
internal diameter 250 μm), which were maintained at 393 K during
analysis to avoid gas adsorption.
The surface species were monitored with a single-beam Fouriertransform infrared (FTIR) instrument (JASCO, Japan; Model FT/IR4200) equipped with a mercury−cadmium−tellurium-M detector at aconstant temperature of 77.4 K.24 A 20 mm-Φ self-supporting disk of
ZrO2 or Ag (5.0 wt %)−ZrO2 disk (50 mg) was placed in a quartz
photoreaction cell equipped with NaCl windows on both sides. The
photoreaction cell was connected to the Pyrex glass circulation system
as well as the GC−MS to enable simultaneous surface species
monitoring via FTIR and isotope distribution in the gas with the
GC−MS. The photocatalyst disk was evacuated (10−6 Pa) at 295 K
for 2 h prior to FTIR and GC−MS measurements.16
In situ FTIR measurements were performed at 295 K in a range
from 4000 to 650 cm−1. The sample disk was irradiated with UV−
visible light from a 500 W Xe arc lamp using quartz ﬁber light guide(San-ei Electric Co., Japan; Model 5Φ-2B-1000L). The distance
between the ﬁber light exit and sample disk was 50 mm. The lightintensity at the center of sample was 88 mW cm−2. The
spectrometer’s energy resolution was 1 cm−1. A 10%-cut ﬁlter was

Data Points (μx − background )2i
i
exp(−0.075ki 2)
i=1∑≤ smoothing factor(4)

where k is the angular photoelectron wavenumber.
Multiple-shell curve ﬁt analyses were performed on the Fourierﬁltered k3-weighted EXAFS data in k- and R-space (R: interatomic
distance) using the empirical amplitude extracted from the EXAFS
data for the Ag metal foil (40 μm thick). The R and coordination
number (N) values for the Ag−Ag interatomic pair were set to 0.288
9 nm and 12.30 We assumed that the many-body reduction factor, S02,
is identical for both the sample and reference.
The high-resolution transmission electron microscopy (HR-TEM)
investigations were performed using a JEM-2100F (JEOL) equippedwith a ﬁeld emission gun at an acceleration voltage of 200 kV.27 The
samples were mounted on a Cu mesh (250 mesh per inch) coatedwith a copolymer ﬁlm of poly(vinyl alcohol) and formaldehyde(Formvar, Monsanto, St. Louis, MO, USA) and coated with carbon.
Chemical compositions and elemental distributions were analyzed
using energy dispersive spectra equipped with a Si(Li) detector in the
TEM.

3. RESULTS AND DISCUSSION
We ﬁrst compared the reactions exposed to CO2 and H2 using
several ZrO2-based photocatalysts doped with varying amountsof Ag under UV−visible light irradiation (Table 1A-a−e andFigure S2). The major product using these catalysts was always
CO. The formation rate using 5.0 wt % Ag−ZrO2 (0.57 μmol
h−1 gcat−1) was higher by a factor of 3.9 than when using ZrO2
(0.15 μmol h−1 gcat−1; Table 1A-a, d). When the photocatalyst
Ag content varied between 0.50 and 10 wt %, we were able tomaximize the total CO formation rate by using Ag (5.0 wt %)−6293

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Table 1. Kinetic Data on Photoconversion of CO2 Using the
ZrO2-Based Photocatalyst under UV−Visible Light(A) 13CO2 (2.3 kPa) and H2 (21.7 kPa)
formation rate
(μmol h−1 gcat−1)
entrycatalyst

light irradiated13

CO

Full light0.10
λ > 320 nm
0.018
Ag (0.50 wt %)−ZrO2
Full light
0.20
Ag (3.0 wt %)−ZrO2
Full light
0.38
Ag (5.0 wt %)−ZrO2
Full light0.52
λ > 320 nm
0.15
λ > 580 nm
0.0093
Ag (10 wt %)−ZrO2
Full light
0.46(B) 13CO2 (2.3 kPa) and H2O (2.7 kPa)

a
a′
b
c
d
d′
d′′
e

ZrO2

12

CO

0.045
0.0084
0.11
0.079
0.049
0.017
0.0034
0.065

formation rate (μmol h−1 gcat−1)
entry

13

catalyst

12

CO

CO

H2

a
ZrO2
0.018
0.16
0.074
b
Ag (5.0 wt %)−ZrO2
0.00310.00490.15cMg2+−Ag (5.0 wt %)−ZrO2
0.00100.092<0.002(C) Pressure Dependence of CO Formation Rate Using 13CO2, H2, and Ag(5.0 wt %)−ZrO2

Figure 1. Time course formation of photocatalytic 13CO and 12CO
during exposure to 12CO2 (2.3 kPa) and H2 (21.7 kPa) using Ag (5.0
wt %)−ZrO2(0.100 g) irradiated by (A) full UV−visible light and (B)ﬁltered light at λ > 320 nm.

formation rate of 13CO(μmol h−1 gcat−1)
formation rate of 12CO(μmol h−1 gcat−1)
13

CO2 pressure (kPa)H2 pressure (kPa)0

21.7<0.0009<0.001

0.39
0.045
0.52
0.049

1.1
2.3
4.8
6.7

0.50
0.035

43.8

66.7

0.59
0.049
0.61
0.034

0.52
0.033

0.62
0.036

ZrO2 (Table 1A-a−e and Figure S3). We conﬁrmed the
predominant photocatalytic formation of 13CO, which wasderived from 13CO2. However, minor 12CO also formed
continuously (Figures 1A and S2). Caution should be
exercised for formation rates of CO that are of the order of
μmol h−1 gcat−1; the rates of tens of mmol h−1 gcat−1 have beenreported for methane formation.2,31 The diﬀerences of reaction
conditions, products, and their associated free energy changemust be carefully considered; however, this study isotopicallyclariﬁes the mechanism from CO2 to CO activated by light inthe course of time.
The 13CO and 12CO formation rates using Ag (5.0 wt %)−ZrO2 were 0.52 μmol h−1 gcat−1 and 0.049 μmol h−1 gcat−1,
respectively (Figure 1A). The 12C ratio in reactant CO2 was
1.0%, whereas the 12C ratio in formed CO was 8.6% (Table
1A-d). To understand this inconsistency, we performed a
13
CO2 (0.67 kPa) exchange reaction with Ag (5.0 wt %)−ZrO2
under UV−visible light irradiation (Figure 2B). The exchange
reaction proceeded with 12CO2, which was adsorbed from air
and remained after pretreatment under vacuum. The exchange
reaction reached equilibrium after 2 h. We assumed that thisreaction followed the ﬁrst-order kinetics and that the rate
constants, kr and kr′, are the exchange between gas-phase13
CO2 with adsorbed 12CO2 and gas-phase 12CO2 withadsorbed 13CO2, respectively.

Figure 2. Time course exchange reaction of 13CO2 (0.67 kPa)
irradiated by UV−visible light using (A) ZrO2 and (B) Ag (5.0 wt%)−ZrO2. Photocatalyst used was 0.100 g.

dP 13CO2dt= −k rP 13CO2 + k r′P 12CO2(5)

P 13CO2 + P 12CO2 = P 13CO2(initial)(6)∴ P 12CO2 = P 12CO2(equilibrium){1 − e−(kr + kr ′)t }(7)

On the basis of the ﬁt (Figure 2B), the sum of the rate
constants (kr + kr′) required to attain an exchange equilibrium
was 2.0 h−1. A quicker reaction of 13.5 μmol 13CO2 along thefree sites of the Ag (5.0 wt %)−ZrO2 surface had a rate
constant of 9.0 h−1 in comparison to the 13CO2/12CO2
exchange (3.5 μmol). The converged 12CO2 partial pressure
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in total CO2 was 9.3%, which is in good agreement with the12
C ratio that formed in the CO (8.6%). We also performed
the exchange test when exposed to dark conditions (Figure
S4B). The total 13CO2 uptake changed negligibly. However,
the amount of exchangeable 13CO2/12CO2 was greater by afactor of 1.46 when exposed to light (Figures S4B and 2B),
which suggests that the CO2-derived species is light activated.
13
CO began to form after an induction period of 0.5 h due toan exchange delay with the 12C surface species (Figure S2,Left), as well as the Ag activation via heating by light (seebelow).
Using ZrO2 (Figure 2A), the converged partial pressure of
12
CO2 in the total CO2 was 7.3%, whereas the 12C ratio in the
total CO was as high as 27% (Table 1A-a and Figure S2A),
which suggests the existence of an irreversible CO2-derived
species leading to CO formation associated with O vacancy in
ZrO2 (2.3 μmol, Figure 2A and Scheme 1c).32,33 The uptake of
Scheme 1. Proposed Intermediate Species Starting from
CO2 and H2 to CO during CO2 Exchange andPhotocatalytic CO2 Reduction

Figure 3. FTIR spectra of ZrO2 (50 mg; A, B) and Ag (5.0 wt %)−ZrO2 (50 mg; C, D). (A, C) Under CO2 for 2 h (1) and undervacuum for 30 s (2) and 7 min (3). (B, D) Under UV−visible light,
13
CO2 (2.3 kPa), and H2 (21.7 kPa) for 2 h (1) and under vacuum for
90 s (2) and 24 h (3).

O atoms at the O vacancy site in ZrO2 correlated well with theintensity ratio between the Zr 3d and O 1s peaks in the X-ray
photoelectron spectra.34,35 The O vacancy was more reactivein Ag−ZrO2 when exposed to UV−visible light due to H
activation by Ag (3.5 μmol, Figure 2B), whereas we observed
no Ag eﬀects when exposed to dark conditions (Figure S4A,B).
The number of total CO2-derived sites on the ZrO2 (0.100 g)
surface was 1.3 CO2 molecules per nm2, which decreased to
1.1 CO2 molecules per nm2 of Ag−ZrO2 (Figure 2). This
demonstrates that CO2-derived species were above the ZrO2
surface and not above Ag.
The rate dependence on reactant pressure approachingatmospheric pressure was also studied using Ag (5.0 wt %)−ZrO2 catalysts (Table 1C). Extremely weak dependence on
both the pressure of CO2 and H2 was found with a maximum
CO formation rate of 0.66 μmol h−1 gcat−1 at 6.7 kPa CO2 and66.7 kPa H2, demonstrating approximately zero order kinetics
due to the strong adsorption of CO2 and H.
FTIR spectra were measured for the ZrO2 pretreated undervacuum at 295 K for 2 h. With the adsorption of CO2 at 295 K
for 2 h, peaks appeared at 1624, 1421, and 1220 cm−1, whichwere assigned to bicarbonate species (monodentate or
bridging; Figure 3A1, Left).36 Simultaneously, the ZrO2hydroxy peak at 3700 cm−1 decreased due to a reaction with
CO2 while the bicarbonate hydroxy peak increased at 3615
cm−1 (Figure 3A1, Right). Much weaker shoulder peaks, due
to carbonate species, also appeared at 1553 and 1334 cm−1.37
When exposed to vacuum for 30 s and 7 min, the peaks, due tothe bicarbonate species, decrease by 1/2−2/3 (Figure 3A2, 3,Left). Since carbonate species were present in the pretreatedsample under vacuum (data not shown), the changes underboth CO2 and vacuum were minimal and we, therefore, regard
carbonate as inert.
We then monitored FTIR changes associated with UV−
visible light irradiation (Figure 3B). At 2.3 kPa of 13CO2 and21.7 kPa of H2 irradiated by UV−visible light for 2 h, thebackground level, e.g., the level in the wavenumber region of3650−3500 cm−1, increased as a result of the IR excitation ofUV−visible-excited electrons trapped beneath the conduction
band of ZrO2 (Figure 3B2,3),27,38 while the vibrational peaksof surface species showed negligible change before and after
irradiation by UV−visible light. A similar trapping state in
CdSe during IR excitation and further hot electron injections
from LSPR Au nanoparticles were reported.39 The antisymmetric and symmetric stretching vibration peaks at 1624 and1421 cm−1 for bicarbonate and at 1553 and 1334 cm−1 forcarbonate in natural CO2 (Figure 3A, Left) shifted to 1588 and
1389 cm−1 and 1518 and 1304 cm−1 in 13CO2 and H2 (Figure
3B, Left), respectively, based on the following equation for
harmonic oscillation:24

ν̃ =
ν13
̃ CO
ν12
̃ CO
6295

1
2πc

=

k
μ
1
1
+ 16
13
1
1
+ 16
12

(8)

= 0.97778(9)
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where ν̃ is the wavenumber, c is the speed of light, k is the force
constant, and μ is the reduced mass. We chose a peak at 1389
cm−1 to evaluate the amount of change in the bicarbonate
species. Under vacuum for 90 s, the bicarbonate peaks
decreased to one-third of their intensity under 13CO2 and H2(Figure 3B1,2), which corresponds to the amount of exchange
between 13CO2 and 12CO2 (2.3 μmol per 0.100 gcat; Figure2A). The chemisorbed bicarbonate species, which were
exchangeable in gaseous CO2, became bridged while thebicarbonate species desorbed under vacuum since CO2
monodentately coordinates to the Zr site (Scheme 1b,c).36
Similarly, isotope labeled bicarbonate exchanged with CO2,
which was monitored using attenuated total reﬂectance IR.
Previous studies have suggested these as an intermediate to
CO formation above Au.40
We also conﬁrmed that the amount of exchange of CO2 andrate constants (simple adsorption, exchange) negligibly varied
between the powder (Figure 2) and disk samples (Figure 3).
When exposed to vacuum and UV−visible light for 24 h, thebicarbonate peaks further decreased to 85% of their intensity
under vacuum for 90 s (Figure 3B2,3). The decreased amountwas
2.3 μmol ×

0.050
× 0.15 = 0.17 μmol‐13CO3H
0.100

decreased to 89% of their intensity observed under vacuum for
90 s (Figure 3D2,3). The decreased amount was
0.050
3.5 μmol ×× 0.11 = 0.19 μmol‐13CO3H
(11)
0.100
versus 13CO detected by GC−MS (0.029 μmol). Fifteen
percent of the 13CO2 at the Ag−ZrO2 surface photoconverted
to 13CO. Clearly, Ag promoted intermediate species reduction
to CO associated with H supplied from Ag and O vacancy sites
in comparison with the ratio (1.1%) using ZrO2 associatedwith only O vacancy sites.
In the UV−visible spectra for the Ag−ZrO2 samples, a peak
centered at 428 nm was dominant between an Ag compositionof 0.5 and 3.0 wt %, whereas shoulder features centered at 368
and/or 470 nm increased at elevated Ag contents, i.e., 5.0−10
wt % (Figure 4). These peaks are due to LSPR induced by(10)

versus 13CO detected by GC−MS (0.001 9 μmol). Only 1.1%of the 13C-bicarbonate species along the ZrO2 surface
photoconverted to 13CO because of the simultaneous
occurrence of a reverse reaction to 13CO2. Due to thelimitations of the 13X-S molecular sieve column, desorbed
13
CO2 could not be quantitatively evaluated using GC−MS
and the total mass balance starting from bicarbonate was not
conﬁrmed. The amount of reduction of 13CO2 to 13CO should
be related to that of surface O vacancies in ZrO2. The eﬀects ofapparatus-based IR light on photocatalysis should be negligible
because no reduction in CO2 was observed, even in thepresence of H2 (see the following wavelength dependence
section, Table 1A-d′′ nor was signiﬁcant amount of heat
induced in comparison to LSPR-induced heat as a result ofUV−visible irradiation (see the following EXAFS section,
Table S2-e).
We also monitored the behavior between CO2 and the Ag(5.0 wt %)−ZrO2 photocatalyst using FTIR. The peak positionand change in intensity during CO2 adsorption at 295 K andsubsequent evacuation (Figure 3C) were similar to the changes
observed when using ZrO2 (Figure 3A), which demonstrates
that both the bicarbonate and carbonate species were above
the ZrO2 surface, not the Ag surface. A relatively greater
portion of bicarbonate species remained (3/5−1/3) undervacuum for 30 s and 7 min (Figure 3C) in comparison with theratio using the ZrO2 (Figure 3A) in consistent with theamounts of exchanged (chemisorbed) CO2, i.e., 2.3 and 3.5μmol using ZrO2 and Ag (5.0 wt %)−ZrO2, respectively(Figure 2).
At 2.3 kPa of 13CO2 and 21.7 kPa of H2 with irradiation
using UV−visible light for 2 h (Figure 3D), we observed aperfect isotope shift similar to the shift observed using ZrO2
(Figure 3B and eq 9). However, the intensity of thebicarbonate peak decreased by approximately 50% when
exposed to 13CO2 and H2 in comparison with CO2, whichsuggests the presence of H2 activation on Ag and spillover onto
the ZrO2 surface (Scheme 1d). When exposed to vacuum andUV−visible light for 24 h, the bicarbonate peaks further

Figure 4. Diﬀuse reﬂectance UV−visible spectra for the Ag−ZrO2
samples. The Ag contents were 0, 0.50, 3.0, 5.0, and 10 wt %.

UV−visible light where the Ag particle size distribution
became wider with increased Ag content. The absorption
edge was always at 248 nm (Figures 4 and S1), which indicatesthat the ZrO2 BG is 5.0 eV.
As the response to UV light absorption, ZrO2 samples
showed broad and sharp ﬂuorescence peaks at 370 and 396
nm, respectively, with excitation at 200 nm (Figure S5A-a),whereas the peaks were almost extinguished with excitation at
240 nm (Figure S5A-b). The excitation spectra for the two
peaks (Figure S5B) and the wavelength dependence of the Xe
arc lamp (Figure S1) suggested contribution of light at 200 < λ< 248 nm to charge separation in ZrO2 (Scheme 1). A
ﬂuorescence peak was reported for mean 4 nm-ZrO2 withexcitation at 300 nm and an extension of the excitation
wavelength was ascribed to the transition involving extrinsic
states.41 Such an eﬀect involving O vacancy states is also
plausible under the conditions using H2 and UV−visible lightshown in this study (Scheme 1c).
Formation rates of CO using ZrO2 and Ag (5.0 wt %)−ZrO2
irradiated under full light (0.15 and 0.57 μmol h−1 gcat−1; Table
1A-a,d) decreased to 0.027 and 0.17 μmol h−1 gcat−1,
respectively, under light at λ > 320 nm (Table 1A-a′,d′ andFigure 1B). The rate using Ag (5.0 wt %)−ZrO2 irradiated bylight at λ > 580 nm further decreased to 0.013 μmol h−1 gcat−1(Table 1A-d′′). Therefore, for ZrO2, the CO formation rate
decreased by 82% with the ﬁltration of the excitation light at λ
> 320 nm in comparison to a photocatalytic test irradiatedunder full UV−visible light (Table 1A-a,a′), i.e., UV light at λ <248 nm to separate the charges at the ZrO2 BG (Figure 4)
played a signiﬁcant photocatalytic role in the reduction of CO2
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Journal of the American Chemical Society(82%). For the Ag (5.0 wt %)−ZrO2, the CO formation rate
decreased by 70% and 98% with the ﬁltration of the excitation
light at a λ > 320 nm and λ > 580 nm, respectively, incomparison with tests under full UV−visible light (Table 1Ad,d′,d′′). UV light at λ < 248 nm separated the charges at theZrO2 BG and UV−visible light at 330 < λ < 580 nm induced
Ag LSPR (Figures 4 and S1), which also played a signiﬁcantrole in the photocatalytic reduction of CO2 (i.e., 70% and 28%,
respectively).
Total CO formation rates in the blank tests when exposed to13
CO2, H2, and no light were 0.0062 and 0.013 μmol h−1 gcat−1
using ZrO2 and Ag (5.0 wt %)−ZrO2, respectively (Table S1Aa,b). These rates proceeded thermally (295 K,
1
RT = 1.2 kJ mol−1) but were only 4.2% and 2.2%, respec2
tively, of the corresponding rates under UV−visible light(Table 1A-a,d). The other blank test under 13CO2 and UV−
visible light using Ag (5.0 wt %)−ZrO2 formed no products
above GC−MS detection limits (Table S1B-a).
Ag K-edge EXAFS spectra were monitored for Ag (3.0 or 5.0
wt %)−ZrO2 under CO2 (2.3 kPa) and H2 (21.7 kPa)
irradiated using UV−visible light at beamline. In the Fouriertransform (FT; Figure 5C), a peak due to a Ag−Ag interatomic
scattering of photoelectrons dominated at 0.28 nm (phase shiftuncorrected), which demonstrates that the Ag nanoparticles
were exclusively metallic.25 In fact, in the HR-TEM images, weobserved the metallic Ag(1 1 0) lattice fringes (i.e., the interval
0.287 nm versus the theoretical 0.289 nm30) to be in contact
with the tetragonal ZrO2(1 0 0) lattice fringe (i.e., the interval
0.364 nm versus the theoretical 0.364 nm; Figure 6).42 The
EXAFS FT peak intensity was signiﬁcantly suppressed, i.e., byas much as 31%, during a 100 min period of light irradiation(Figure 5C). On the contrary, peak intensity quickly recovered
when the light was turned oﬀ.
Such light-induced changes in the FT were quantitatively
evaluated using a curve ﬁt analysis based on the plane-wave
approximation for amplitude Ai(k), backscattering amplitude f i,
Debye−Waller factor σi, and mean free photoelectron path λfor shell i43 using a XDAP 2.2.7 code:29ÄÅÉÅÅ i 2 2 R i yÑÑÑjz
ÅA i (k ) =|f (k)|expÅÅ−2jjσi k + zzÑÑÑ, i = AgÅÅÇ k
λ {ÑÑÖ
kR i 2 i
Ni(12)

Figure 5. Time course changes of (A) coordination number N values,(B) Debye−Waller factor σ values, and (C) FT obtained from angular
photoelectron wavenumber k3-weighted Ag K-edge EXAFS χ-function
for Ag (3.0 wt %)−ZrO2 under CO2 (2.3 kPa) and H2 (21.7 kPa)
irradiated by UV−visible light for 100 min followed by dark
conditions for 20 min. (D) The change in the FT of the Ag K-edge
EXAFS for fresh Ag−ZrO2: 3.0, 5.0, and 10 wt % Ag. (E) The
correlation between the σ value and temperature for bulk sites (circle,○) and surface sites (vertical motion; square, □) in/on the Ag metal
generated by the correlated Debye model using a FEFF8 code.

The N value was 9.0 ± 0.7 before light irradiation and did notchange signiﬁcantly during both light irradiation (8.3−10.0)
and after the light was oﬀ (10.1 ± 0.7; Figure 5A). N values of
9−10 correspond to a particle size (d) of 2.5−3.7 nm assuming
the spherical face-centered cubic (fcc) nanoparticle model andthat surface dispersion (D) is 0.54−0.36 (mean value =0.45).44,45
The σ value was calculated to be 0.00995 nm for Ag metal at
290 K using the correlated Debye model46,47 with the ab initio
multiple-scattering calculation code, FEFF848 and the Debye
Ag temperature [θD(Bulk) 225 K].49 The XDAP code provides
an experimental diﬀerence for the σ2 value from that of the Ag
metal (model) based on eq 12. Initial σ values of 0.0101 nm
for Ag (3.0 wt %)−ZrO2 before light irradiation quickly
increased to 0.0107 nm (10 min irradiation) and progressively
increased to 0.0117 nm (90 min light irradiation, shown in
Figure 5B). Then, the value quickly decreased to 0.0101 nm
after the light was turned oﬀ at 112 min.
Furthermore, we evaluated the temperature at the Ag site
based on the σ values. The σ value temperature dependence

derives from the FEFF8 combined with the correlated Debye
model46,47 for both bulk and surface Ag sites using the bulk
and surface Debye temperature (Figure 5E-a and b,
respectively). We assumed that the thermodynamically stablefcc(1 1 1) face had preferable exposure for latter value[θD(Surf,⊥) 155 K].50 We also approximated the mean Ag
nanoparticle temperature as the arithmetic mean temperature
based on the θD(Surf,⊥) weighted by 1/2·1/3D [D: dispersion of
nanoparticles (0.45), for an eﬀective vertical degree of freedom
at a free hemisphere surface] and that based on the θD(Bulk)
weighted by (1 − D) + 1/2D + 1/2·2/3D (bulk site, nonfree
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Journal of the American Chemical Society

irradiation. This fact guarantees that the heat originatingfrom probe synchrotron X-ray was negligible in this study.
Ag nanoparticles were always metallic for the Ag−ZrO2
samples (Figure 5D), and N values for the fresh Ag−ZrO2
samples gradually increased from 9.0 ± 0.7, 9.6 ± 0.6, and 10.7
± 0.8 for 3.0, 5.0, and 10 wt % Ag, respectively, whichcorresponds to (d, D) sets of (2.5 nm, 0.54), (3.1 nm, 0.43),
and (4.2 nm, 0.28).44,45 Thus, Ag surface atoms increase from
0.15 to 0.20 and ﬁnally to 0.26 mmol gcat−1, respectively,
whereas the CO formation rates reached maximum values at
5.0 wt % Ag (Table 1A and Figure S3), which indicates that
the ZrO2 sites were primarily responsible for CO2 reduction
and were assisted by heated neighboring Ag sites via LSPR that
activated and supplied H to move forward eq 2 (Scheme 1d).
The close correlation between Au LSPR and H2 activation wasreported using the Au−SiO2 photocatalyst due to the charge
excitation to the antibonding H2 orbital from LSPR.54 H2
activation via the Ni−H species above the Ni/SiO2·Al2O3
photocatalyst31 relates to the mechanisms discussed in thisstudy, i.e., endothermic CO formation, versus exothermic
methane formation.31
In the Ag−ZrO2 UV−visible absorption spectrum, peaks
appear centered at 428 and 368 nm (2.90−3.37 eV) due to a
LSPR at Ag (Figure 4). The Ag work function is 4.52−4.74
eV,55 i.e., a Fermi level (EF) of 0.08−0.30 V for a standardhydrogen electrode versus the reduction potential from CO2 to
CO (−0.11 V)1 or ZrO2 conduction band minimum (−1.0V).13 For nanoparticles to exhibit LSPR behavior, it waspreviously reported that the following mechanisms need tooccur: (i) charge excitation to an unoccupied adsorbate
state,54,56 (ii) hot electron injection that originates from LSPR
for support,19,57−60 (iii) electron trapping from the support to
the Schottky barrier,59 and (iv) plasmonic resonant energy
transfer (PRET).59−61 However, for cases i−iii, it is often
diﬃcult to explain how holes remained react after electron
activation/reaction because the O moiety that derives from
CO2 must transfer to the LSPR nanoparticles for oxidation. In
this study, H2 oxidizes based on its association with the
reduction of bicarbonate species above ZrO2. Therefore, holes
react with H2 to form H+ and must further react with O thatderives from CO2 to form water above the Ag surface. Thus,
Ag thermally contributes for CO generation when heated to325 K−392 K (Figure 5B), i.e., converted from LSPR (H2easily dissociates on the heated Ag surface and spills over the
ZrO2 surface) (Scheme 1d,e).
The conversion of LSPR over Ag and Au owing to heatpreceded phenol decomposition62 and thiosulfate oxidation,63
whereas previous studies considered that short-lived heat
converted from LSPR over Ag promotes reactant diﬀusion in
the aqueous solution.64 This study is the ﬁrst to report theconversion of CO2 assisted by heat that was converted via
LSPR. Peaks due to formate species were not detected in our
FTIR spectra, and we are unable to ﬁnd evidence that theformate was an intermediate to CO.14,65 The H andbicarbonate species coupled with the hole and electron,
respectively, and combine to form CO and water, which isenhanced by a factor of 3.9 due to Ag (Table 1A-a,d). The risein the Ag temperature in the ∼0.5 h experiment via LSPR
(Figure 5B) and H2 activation eﬀects (Scheme 1) agree withthe ∼0.5 h induction period for CO formation using Agcontaining photocatalysts (Figure S2, Left).
Finally, we tested the photocatalytic conversion of CO2
using moisture (eq 3). Using ZrO2, 13CO2 (2.3 kPa), moisture

Figure 6. HR-TEM image of Ag (5.0 wt %)−ZrO2.

hemisphere in contact with ZrO2, and two lateral degrees of
freedom at a free hemisphere surface) (see Supporting
Information for the detail).
As a result, the initial temperature of 286 K before
irradiation rose to 325 K after 10 min of irradiation andprogressively increased to 392 K after 90 min of lightirradiation (Figure 5B and Table S2a). The temperature
quickly dropped to 290 K after the light was turned oﬀ. Suchrise/drop in temperature is possibly due to LSPR heattransformation.51 Similarly to the result using Ag (3.0 wt%)−ZrO2 above, the initial temperature increased to 363 K
after a 90 min UV−visible light irradiation and quickly
dropped to 297 K once the light was oﬀ using the Ag (5.0 wt%)−ZrO2 sample (Table S2c). Related to the EXAFS method,
the nanoparticle temperature was also monitored based on
surface-enhanced Raman scattering for the NC stretchingvibration of adsorbed probe molecules.52 The method basedon the σ value used in this study via EXAFS is common andcan directly probe for all nanoparticles to expose LSPR, e.g.,
Au and Ag. Three-dimensional mapping of Ag nanoparticles
was performed via electron energy loss spectroscopy as afunction of energy loss.53
Furthermore, we also observed increases or decreases intemperature under argon and UV−visible light (367 K at 50
min of light, 302 K 10 min after light had been turned oﬀ;
Figure S6B and Table S2b), demonstrating that the ambient
gas was not a major factor for temperature up/down. The σvalue/temperature negligibly changed until 329 K when theincident light wavelength was ﬁltered at more than 580 nm(Table S2e), whereas an essentially identical rise/drop during
full light (365 K at 90 min, 298 K when light was oﬀ) for the σvalue/temperature was observed when λ was more than 320
nm (Table S2d), demonstrating the heating eﬀect by infrared
portion of light were marginal in comparison to that by light
whose wavelength was in the range 320 < λ < 580 nm via
LSPR of Ag (Figures 4 and S1). Reaction heat was not a factor
in this study because the reaction that reduces CO2 to CO (eq2) is endothermic. The heating and reaction promotion wasreported for the exothermic formation of methane from CO2
using Ni-based photocatalysts.31 The temperature dropped tothe same level (i.e., 290−302 K) essentially in 10−35 min after
UV and/or visible light turned oﬀ (Table S2a−e) incomparison to 286 K before UV and/or visible light
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Journal of the American Chemical Society(2.7 kPa), and UV−visible light, the CO generation rateincreased by a factor of 1.2 compared with using 13CO2 and H2(Table 1A-a, B-a). However, the molar ratio of the newly
formed CO:H2 was 1:0.41 rather than 1:1 based on eq 3. Since
90% of the newly formed CO was 12CO, the bidentate
bicarbonate species c was more stable above ZrO2 (Scheme 1),
forming CO gas and a hydroxy group. Conversely, using the Ag(5.0 wt %)−ZrO2 photocatalyst, the CO formation rate wassuppressed by 96% compared with the formation rate whenusing ZrO2 (Table 1B-a,b) and by 99% when using CO2 andH2 (Table 1A-d). Instead, H2 preferably formed at 0.15 μmol
h−1 gcat−1. Based on the eﬀects of Ag, H2 selectivity in the
reduced products increased from 29% to 95% with eq 1 (water
splitting) rather than eq 3 due to the preferential adsorption ofwater versus CO2 on ZrO2 and H2 desorption above Ag(Scheme S1). Water that blocks the O vacancy site is also
plausible. Once reaction intermediates, i.e., H and bicarbonate
form under light, CO and H2 are slowly formed via thermalenergy under dark at 295 K (Table S1C-a′). In contrast that
12
CO was preferably formed among total CO under light(61%; Table 1B-b and Figure S7, Left column), 13CO becamea major product (68%, Table S1C-a′ and Figure S7, Rightcolumn) under dark due to the gradual consumption ofpreadsorbed bicarbonate (12C 98.9%) that is in the equilibriumwith gas-phase 13CO2 (Figure 2B). The addition of Mg2+ onto
the Ag (5.0 wt %)−ZrO2 to attract CO2 to the surface
mitigated the water blocking problem. Therefore, the COformation rate increased from 0.0080 to 0.093 μmol h−1 gcat−1,
i.e., a factor of 11.7, via the addition of Mg2+ similar to the
MgO−Pt−TiO2 photocatalysts (Table 1B-c).66

and H2. Instead, we found that using CO2 and moisturedisadvantageously yielded reversely formed H2 over Ag.
However, H2 formation was redirected toward CO formation
via the addition of Mg2+ as the CO2-anchor sites. The
combination of alkaline [earth] metal ion−LSPR metal (Ag,Au)−ZrO2 is required for artiﬁcial photosynthesis free from Cimpurity problems and using the dual roles of light: charge
separation and heat via LSPR.

4. CONCLUSIONSIn this study, the Ag (5.0 wt %)−ZrO2 photocatalyst was themost active during 13CO2 reduction using H2 and UV−visible
light at a rate of 0.66 μmol h−1 gcat−1. The product included
5.5−8.6% 12CO. However, this CO derives from preadsorbed
CO2 from the air that exchanged with gas-phase 13CO2 at a
rate constant of 2.0 h−1. FTIR spectra demonstrated bridging
and monodentate bicarbonate species via ZrO2 surface hydroxygroup consumption under photoreaction CO2 and/or H2
conditions. The bridging bicarbonate species were exchangeable with the gas-phase CO2 and O vacancy on the ZrO2
surface should have participated in its formation. Under H2and UV−visible light with the expense of bridging bicarbonate,
15% was directed to 13CO formation, whereas the remaining
proportion reversely desorbed 13CO2.
The charge separation contribution at the ZrO2 BG and the
Ag contribution characterized by LSPR were evaluated basedon in-proﬁle kinetic data measurements using sharp-cut ﬁlters:
70 and 28%, respectively. We further investigated thecontribution of the mean 2.5−3.7 nm Ag nanoparticles using
in situ EXAFS. A rise in temperature from 286 to 392 K when
exposed to UV−visible light irradiation and a rapid drop to
290 K under dark conditions were directly monitored based on
the Debye−Waller factor change for a Ag−Ag interatomic pairinterference. H2 should oxidize based on its association withbicarbonate species reduction above ZrO2, but no evidence foravailable O on Ag was found in this study. Thus, the heated Ag
surface activated H2, and the H atoms spill over to thebicarbonate species above ZrO2 rather than hot electron
injection or PRET to/with ZrO2.
Such dual roles for light during charge separation at the BG
and heat via LSPR were remarkably eﬀective when using CO2

Notes

■

ASSOCIATED CONTENT* Supporting InformationS

The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/jacs.8b13894.
Intensity distribution of Xe arc lamp, detailed data fromthe photocatalytic tests and data from control experiments, absorption−ﬂuorescence spectra, theoreticaldetails and synchrotron X-ray experiments under Ar, asummary of the Ag temperature evaluation based on the
Debye−Waller factor, and time course as well as the
proposed reaction mechanisms under CO2 and moisture
(PDF)

■

AUTHOR INFORMATION

Corresponding Author*yizumi@faculty.chiba-u.jp
ORCID

Yasuo Izumi: 0000-0001-8366-1864
Author Contributions

YI, HZ, TI, and TK contributed equally.
The authors declare no competing ﬁnancial interest.

■

ACKNOWLEDGMENTS
The authors are grateful for ﬁnancial support from the Grantin-Aid for Scientiﬁc Research C (17K05961) from the Japan
Society for the Promotion of Science, Grant for Academic
Research from the Japan Gas Association (2018), and Leading
Research Promotion Program (2015−2019) from the Institutefor Global Prominent Research, Chiba University. X-ray
absorption experiments were performed with the approval of
the Photon Factory Proposal Review Committee (2016G577,
2018G649). The authors thank Dr. K. Niki, Department ofChemistry, Chiba University, for helpful discussion on LSPR.
The authors would like to thank Enago (www.enago.jp) for the
English language review.

■

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