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ArticleLocal Silver Site Temperature Critically Reﬂected Partial andComplete Photooxidation of Ethanol Using Ag−TiO2 as Revealed by
Extended X‑ray Absorption Fine Structure Debye−Waller Factor
Daiki Fukuhara, Moses Tuhafeni Joseph, Tarik Loumissi, Chao Zhang, Takaomi Itoi, Hongwei Zhang,and Yasuo Izumi*
Cite This: J. Phys. Chem. C 2021, 125, 14689−14701

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sı Supporting Information*

ABSTRACT: The critical factors in choosing the reaction pathway
for photocatalysis are often unclear. In this study, for a typical Ag−
TiO2 photocatalyst, the control factors of partial and complete
oxidation of ethanol were investigated using kinetics, UV−visible
and emission spectroscopy, and silver K-edge extended X-ray
absorption ﬁne structure. Low concentrations of O2-derived
intermediate species that were not detected by curve-ﬁtting analysis
could be monitored under the photocatalytic conditions based on
the local temperature of Ag sites provided by the Debye−Waller
factor and the correlated Debye model. The site temperaturemonitoring was extended to Ag nanoparticles growing from 0.5 to
3.6 nm under photocatalytic reaction conditions. The Ag site
temperature reached 404 K under reductive conditions todehydrogenate ethanol into acetaldehyde, whereas it was 363−368
K under oxidative conditions owing to O2-derived species forming CO2, CH4, and H2O to suppress localized surface plasmon
resonance. Under UV light, partial ethanol oxidation to acetaldehyde and O2 activation for complete oxidation to CO2 and H2Oproceeded over TiO2. However, under visible light, the C−C bond cleavage to CO2 and CH4 and complete oxidation to CO2 andH2O combined with the O2-derived species transferred from the TiO2 surface proceeded over Ag0 nanoparticles.

■(LSPR) transferring to metal oxides13,14 and/or adsorbed
molecules12 have also been reported. Recently, the origins of
the photocatalytic reaction steps have been investigated:consecutive reaction mechanism (1) to reduce CO2 to CO on
ZrO2 owing to charge separation at the band gap (BG) and (2)subsequent H2 and/or H2O activation over the Ag/Au
nanoparticle surface owing to heat converted from LSPR
spilling over reactive H species to the CO2 reduction sites over
ZrO2.15,16 Using Ni nanoparticles, the subsequent reaction (2′)
consisted of further reduction steps of CO to methane (CH4)owing to heat converted from absorbed light energy.17
The photooxidation rates of ethanol were on the order of >1
mmol h−1 gcat−1 in comparison to CO2 photoreduction rates
under reductive conditions: <1 mmol h−1 gcat−1.15−17 The
oxidation states of active Ag sites and the reactive oxygen species

INTRODUCTION
Photocatalysis consists of redox reactions by separated charges,
and the photooxidation reaction often proceeds faster than thecomplementary photoreduction reaction.1−3 Recently, photoreduction has received much more attention, e.g., the reductionof CO2 and/or N2 to fuels/feedstocks.4,5 Metal oxide/hydroxide
semiconductors are used as the major components of photocatalysts because they are often inexpensive and easy to handle,
e.g., TiO2.6 Among them, it remains open for discussion how tocontrol partial7 and complete photooxidation reactions that are
key techniques for sustainable applications.1
In this paper, the most typical photocatalyst consisting of
metal oxide (TiO2) and metal nanoparticles (Ag) is chosen, andthe processes of photocatalytic dehydrogenation (partial
oxidation) and complete oxidation to CO2 and H2O were
tested. The key reaction mechanism to select each product wasinvestigated by extended X-ray absorption ﬁne structure(EXAFS), based on the Fourier transform, curve-ﬁt analysis,
and the interpretation of the Debye−Waller factor.8,9
Although the trapping eﬀect of excited electrons on metal
nanoparticles transferred from the conduction band (CB) of ametal oxide semiconductor has been widely accepted,10−12
excited electrons owing to localized surface plasmon resonance© 2021 American Chemical Society

Received: May 7, 2021
Revised: June 18, 2021
Published: July 1, 2021

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suppressed from saturated pressure (4.91 kPa @290 K) tomonitor consecutive reaction to acetaldehyde and then to CO2
and H2O (eqs 1 and 3).

were monitored simultaneously by the peak intensities of Ag−O
and Ag−Ag interatomic pairs in the Fourier transform and the
Debye−Waller factor of EXAFS. The local temperaturemonitoring via the Debye−Waller factor based on the correlated
Debye model8,9 was advantageous to monitor atomic-level local
catalytically active elements.
Among the atomic-level state-of-the-art techniques, EXAFS
directly monitors the thermal vibration of X-ray-absorbing
atoms in contrast to surface-enhanced Raman scattering (SERS)
monitoring which requires a probe molecule, e.g., p,p′dimercaptoazobenzene18,19 or adonitol,20 phonon monitoringby Raman which requires Si dots in the Au cavity,21 Raman
spectroscopy which requires locally doped ions, e.g., Yb3+ and
Er3+,22 and absorption spectroscopy for Au nanoparticles which
requires laser irradiation.23 The local temperature monitoring by
EXAFS has been reported for Ag, Au, and Ni nanoparticles overZrO2 under CO2, H2 (or H2O), and UV−visible light, but thetemperature increase exclusively originated from light energy
because of the slow reaction rates to convert thermodynamicallystable CO2.15−17 In this paper, the exothermic heats of partial(eq 1) and complete (eq 2) oxidation reactions of ethanol weremonitored as well as the heat converted from irradiated UV−
visible light.

CH3CHO + 5/2O2 → 2CO2 + 2H 2O
Δr H ° = −1104.47 kJ mol−1(1)

C2H5OH + 3O2 → 2CO2 + 3H 2O
Δr H ° = −1277.38 kJ mol−1(3)

As the partial dehydration (eq 1) and complete oxidation of
ethanol (eq 2) require 0.5 and 3 equiv of O2, we varied theC2H5OH:O2 ratio between 1:0 and 1:4 to obtain insight into thereaction mechanism.
The photocatalyst was irradiated with UV−visible light from a
500 W xenon arc lamp (Ushio, Tokyo, Japan; model UI-502Q)
from the bottom of the quartz cell for 5 h. The distance between
the light exit and the photocatalyst sample was 10 mm. The lightintensity was 70 mW cm−2 at the center of the photocatalyst
position.28 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. L39 and U330 (2.5 mm thick; Hoya, Japan)
ﬁlters were used to pass light with wavelengths of >390 nm(L39), 245 < λ < 386 nm, and >686 nm (U330).29 A control
kinetic test irradiated by UV−visible light while the Ag (2.0 wt%)−TiO2 photocatalyst and the reactor were immersed in aquartz bath ﬁlled with water controlled at 295 K was also
performed.
The gas composition during the light irradiation wasmonitored using an online gas chromatograph equipped witha thermal conductivity detector (Shimadzu, Kyoto, Japan;
Model GC-8AT). A packed column of Molecular Sieves 13X-S(3 m length, 3 mm internal diameter; GL Sciences Inc., Japan)
was used for the separation of O2, N2, CH4, and CO, and apacked column of poly(ethylene glycol)-6000/Flusin P (3 mlength, 3 mm internal diameter, GL Sciences) was also used forthe separation of CO2, CH3CHO, acetone, C2H5OH, and H2O.
Helium (0.13 MPa, purity >99.9999%) was used as the carriergas. However, argon (0.13 MPa, purity >99.998%) was used as
the carrier gas for the separation of H2 through the packed
column of Molecular Sieves 13X-S.
UV−Visible Spectroscopy. UV−visible spectra wererecorded on a double-beam model V-650 spectrophotometerusing D2 and halogen lamps below and above 340 nm,
respectively, equipped with a photomultiplier tube and an
integrated ISV-469 sphere (JASCO, Tokyo, Japan) for diﬀusereﬂectance detection. The wavelength range was 200−800 nm.
The data were transformed using the Kubelka−Munk function.
Fresh photocatalyst samples and those used in various
photocatalytic reaction conditions were quickly set in a sample
holder for UV−visible spectrum measurements. A poly(tetraﬂuoroethylene) plate was used as the reference.
Fluorescence Spectroscopy. The absorption−ﬂuorescence spectra were recorded on an FP-8600 (JASCO; Chiba
Iodine Resource Innovation Center) using a 150 W Xe arc lamp
(UXL-159, Ushio) equipped with a photomultiplier tube. The
excitation was at 300−480 nm, and the ﬂuorescence range was
350−800 nm. The incident excitation light from the Xe lamp
was monitored using a Si photodiode, and the monitored
ﬂuorescence light emitted from the sample was normalized on
the basis of 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.

C2H5OH + 1/2O2 → CH3CHO + H 2O
Δr H ° = −172.91 kJ mol−1

Article(2)■

EXPERIMENTAL SECTION
Catalyst Preparation. AgNO3 (0.160 g; purity >99.8%,
Wako Pure Chemical, Japan) and TiO2 (3.30 g; P25 Evonik,
Essen, Germany, speciﬁc surface area 60 m2 g−1 andanatase:rutile phase = 8:2) were mixed in pure water (50 mL,<0.055 μS cm−1), agitated by ultrasound (430 W, 38 kHz) for 10
min, and stirred at 353 K and 850 rotations/min for 1 h.24,25 The
obtained white precipitate was ﬁltered using a poly(tetraﬂuoroethylene)-based membrane ﬁlter (Omnipore
JVWP04700, Millipore, Burlington, MA; pore size 0.1 μm)
and washed using pure water (each 50 mL) ﬁve times. Then, theproduct was dried at 373 K for 24 h, followed by heating to 673
K at a ramping rate of 9.9 K min−1, and maintained at 673 K for 2
h. The obtained light yellow powder became gray overnight
because of the oxidation of Ag0 formed by heating to AgI2O,
which is a stable state at 293 K. The Ag content in the sample wasdetermined to be 2.0 wt % based on X-ray absorption edge jump(see the X-ray Absorption Fine Structure (XAFS) Measurements section below). The ﬁnal product was denoted as Ag (2.0
wt %)−TiO2.
Photocatalytic Oxidation Tests. The photooxidation testsof ethanol were performed in a closed circulation system (total
inner volume 202.7 mL including the reactor and two sampling
sections) connected to a vacuum using rotary and diﬀusion
pumps (10−6 Pa). The photocatalyst samples (0.100 g) were
homogeneously dispersed in a ﬂat-cylinder quartz reactor(bottom area 23.8 cm2; inner volume 48.8 mL),26,27 connectedto the closed circulation system and evacuated for 2 h. Then,
ethanol (1.33 kPa) and/or O2 (2.67 or 5.32 kPa) was introduced
into the quartz reactor. The molar amount of ethanol was 6.0
times that of Ag atoms in the sample. The amount was14690

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reduction factor, S02, is identical for both the sample and thereference.

Transmission Electron Microscopy (TEM) Measurements. TEM was performed using a JEM-2100F (JEOL)
equipped with a ﬁeld-emission gun (acceleration voltage of 200kV) at the Center for Analytical Instrumentation. The samples
were mounted on a Cu mesh (250 meshes per inch) coated withcarbon and a copolymer ﬁlm of poly(vinyl alcohol) andformaldehyde (Formvar, Monsanto, St. Louis, MO).
X-ray Absorption Fine Structure (XAFS) Measurements. Silver K-edge XAFS spectra were measured at 290 K
in transmission mode at the Photon Factory Advanced Ring,
High Energy Accelerator Research Organization (KEK,
Tsukuba, Japan), on the NW10A beamline. A Si (311)double-crystal monochromator and a Pt-coated focusing
cylindrical mirror were inserted in the path of the X-ray beam.
A piezotransducer was used to detune the X-rays to two-thirds ofthe maximum intensity to suppress higher harmonics. The Ag Kedge absorption energy was calibrated at 25 516.5 eV using the
X-ray spectrum of a Ag metal foil (40 μm thick).30
A disk (Φ = 10 mm) of a Ag (2.0 wt %)−TiO2 photocatalyst(140 mg) was set in a Pyrex glass reactor equipped with a Kapton
ﬁlm (Dupont, Wilmington, DE; 50 μm thick) for X-ray
transmission and a poly(ethylene terephthalate) (PET) ﬁlm(Toyobo Film Solutions, Japan, G2; 38 μm thick) for both UV−
visible light and X-ray transmission ﬁlled with 1.33 kPa ofethanol and 2.67 or 5.32 kPa of O2. The sample was irradiatedwith UV−visible light from a Xe arc lamp through the PET ﬁlm
at the beamline (Figure S1, Supporting Information).15−17 As
control tests, Ag K-edge XAFS spectra were measured for the Ag(2.0 wt %)−TiO2 photocatalyst (140 mg) under 1.33 kPa ofethanol and 2.67 kPa of O2 irradiated with light from a Xe arc
lamp ﬁltered by L39 and U330 ﬁlters to conﬁrm the lightwavelength dependence and under 1.33 kPa of ethanol at 301−
403 K using an electric furnace at the beamline to conﬁrm thetemperature evaluated on the basis of the correlated Debye
model.8,9
The obtained Ag K-edge XAFS data were analyzed using the
XDAP software package.31 The pre-edge background wasapproximated with a modiﬁed Victoreen function, i.e., C2/E2 +C1/E + C0, where E is the photon energy. The background forpostedge oscillation, μx, was approximated with a smoothing
spline function and was calculated for a particular number ofdata points:
data points (μx − background )2
i
i
2
exp(
−
0.075
k
)
i
i=1

∑

≤ smoothing factor

Article■RESULTS AND DISCUSSION
Photocatalytic Oxidation Tests. The photocatalyticoxidation tests of ethanol using TiO2 and Ag (2.0 wt %)−TiO2 catalysts are summarized in Figures 1 and 2. CO was not

Figure 1. Time course changes in the photocatalytic tests duringexposure to (i) C2H5OH (1.33 kPa), (ii) C2H5OH (1.33 kPa) and O2(2.67 kPa), and (iii) C2H5OH (1.33 kPa) and O2 (5.32 kPa) using (a)TiO2 and (b) Ag (2.0 wt %)−P25 irradiated by UV−visible light. The
amount of photocatalyst was 0.100 g.

detected above the detection limit (0.2 nmol) while themaximum H2 evolved was 0.79 μmol in 5 h of reaction, andthe amount was negligible throughout this study. Furthermore,
the authors could not distinguish between the ethanol
adsorption (Scheme 1A-b and B-b) and carbonaceous
deposition.
The time course of ethanol photooxidation under various
conditions using TiO2 and Ag (2.0 wt %)−TiO2 is summarized
in Figure 1. Using TiO2 (0.100 g), C2H5OH (1.33 kPa), andUV−visible light, any reactions proceeded negligibly (Figure 1ai). Of the introduced ethanol (112 μmol), 29.3 μmol remainedin the gas phase at 5 h of reaction. Therefore, 73.8% adsorbed as
ethoxy and hydroxy species over the TiO2 surface (Scheme 1A-aand b) before 5 h of reaction, similar to a recent report.34
Using TiO2 (0.100 g), C2H5OH (1.33 kPa), O2 (2.67 kPa),
and UV−visible light (Figure 1a-ii), the initial formation rates of
CH3CHO and H2O were 816 and 860 μmol h−1 gcat−1,
respectively, following eq 1 for 30 min of reaction. Then, CO2
was consecutively formed from CH3CHO at the rate of 512μmol h−1 gcat−1 via eq 3 between 30 min and 2 h of reaction.
Similar consecutive photooxidation of C2H5OH to CH3CHO
and then CO2 was reported using Ag−TiO2−montmorillonite
under ethanol and moisture.35 Thus, the TiO2 surface activatedO2 (Scheme 1B-c) as reported in the literature.7 The molar(4)

where 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:interatomic distance) based on the plane-wave approximation
for amplitude Ai(k), coordination number Ni, backscattering
amplitude f i(k), Debye−Waller factor σi, and mean freephotoelectron path λ for shell i:ÄÅÉÅÅ i 2 2 R i yÑÑÑNijz
ÅA i (k ) =|f (k)| expÅÅ−2jjσi k + zzÑÑÑ, i = Ag, OÅÅÇ kλ {ÑÑÖ
kR i2 i(5)
using the empirical amplitude extracted from the EXAFS data
for the Ag metal foil (40 μm thick) and Ag2O powder. The R
values for the Ag−Ag and Ag−O interatomic pairs were set to0.2889 nm with the N value of 1232 and 0.2044 nm with the Nvalue of 2, respectively.33 We assumed that the many-body14691

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ArticleFigure 2. Time course changes in the photocatalytic tests during exposure to C2H5OH (1.33 kPa) and O2 (2.67 kPa) using (a) TiO2 irradiated underfull light and (b) Ag (2.0 wt %)−TiO2 irradiated under (i) full light, (ii) full light and at 297 K, (iii) visible light of λ > 390 nm, and (iv) UV and IR lightof 245 < λ < 386 nm and λ > 686 nm. The amount of photocatalyst was 0.100 g.

Scheme 1. Major Reaction Steps that Occurred Using (A) Ag−TiO2 Photocatalyst, Ethanol, and UV−Visible Light Irradiationand (B) Ag−TiO2 Photocatalyst, Ethanol, O2, and UV−Visible Light Irradiation

amount of formed H2O during the period was equal to that of
CO2, exactly following the stoichiometry of eq 3.
When the initial molar ratio of ethanol:O2 changed from 1:2
to 1:4 (Figure 1a-iii), complete oxidation reaction via eq 2
became predominant, and the amount of CO2 produced (115μmol) was 1.92 times greater than that under ethanol:O2 of 1:2(Figure 1a-ii). However, the amount was still 51.4% of that
expected if eq 2 completely proceeded. Furthermore, H2O
formation (61.8 μmol) stopped after more than 1 h of reaction(Figure 1a-iii), implying incomplete oxidation reaction including TiO2 surface species, e.g.,

Next, the Ag (2.0 wt %)−TiO2 photocatalyst was tested. UnderC2H5OH gas only, CH3CHO was formed slowly: 18.4 μmol
over Ag nanoparticles until 4 h of reaction (Figure 1b-i and
Scheme 1A-d and e). The molar amount was half of the Ag
atoms (27.8 μmol) contained in the photocatalyst (0.100 g),
indicating the following noncatalytic reaction reducing AgI sites.
C2H5OH + Ag 2O → CH3CHO + H 2O + 2Ag
Δr H ° = −141.86 kJ mol−1(7)Essentially equivalent ethanol photodecomposition rates werereported using Ag−TiO2, Ag2O−TiO2, and AgO−TiO2 photocatalysts irradiated with a low-pressure Hg lamp,36 suggestingthat the Ag sites in all these catalysts were reduced to the Ag0
state under ethanol and UV light. This noncatalytic reaction
proceeded signiﬁcantly more slowly (18.4 μmol h−1 gcat−1) than

C2H5OH + 7/4O2→ CO2 + 1/2H 2O + CH3(surface) + 2HO(surface)
(6)
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the photocatalytic reaction using O2 over TiO2 (816 μmol h−1
gcat−1, Figure 1a-ii). The H2O formation was 74% greater thanthe CH3CHO formation contradicting the stoichiometry of eq
7. The extra water was due to the dehydration of ethanol. If thecontribution of the noncatalytic reaction (eq 7) is excluded, thephotocatalytic activity of ethanol decomposition using TiO2 andAg−TiO2 was essentially zero in the absence of O2. There wasonly 20% increase in the ethanol photodecomposition rate using
TiO2 and Ag−TiO2 in the presence of moisture irradiated bylight-emitting diode light (λ = 405 nm).37
When C2H5OH (1.33 kPa) and O2 (2.67 kPa) wereintroduced, CH3CHO, H2O, and CO2 were formed at a molar
ratio of 1:1.82:0.57 until 30 min of reaction (Figure 1b-ii), which
was diﬀerent from the ratio of 1:1.05:0.26 using TiO2 (Figure
1a-ii), suggesting that both partial oxidation to CH3CHO via eq
1 over TiO2 and complete oxidation to CO2 via eq 2 over Ag
nanoparticles occurred. The products between 30 min and 2 h ofreaction were CO2 and CH4, whereas CH3CHO was completely
consumed until 2 h of reaction. Thus, complete oxidation via eqs
1 and 3 and dissociative redox reaction via eq 8 proceeded over
the Ag nanoparticles (Scheme 1B-d−f) in contrast to the slower
consecutive reactions of eqs 1 (Scheme 1B-a−c) and 3 over
TiO2.

ArticleThe excited O species on TiO2 irradiated by UV light resulting
in charge separation at the BG was found to be indispensable forthe formation of CO2 and CH4 (Scheme 1B-c′ and f).
Conversely, using a ﬁlter to pass light of 245 < λ < 386 nm andλ > 686 nm to pass UV and IR light, the time course (Figure 2biv) resembled that irradiated under UV−visible light at 297 K(Figure 2b-ii). Thus, heated Ag nanoparticles converted fromlight energy proceeded along the dissociative redox reaction via
eq 8, and the origin of the energy should be from visible light.
The reaction of eq 3 proceeded much more slowly over Ag
nanoparticles at 297 K than that at 363−368 K owing to lightirradiation; however, the reaction of eq 3 proceeded well after 2
h of reaction (Figure 2b-iv) due to a self-acceleration mechanism
by exothermic heat (see the EXAFS section).
UV−Visible Spectra. The absorption edge of the spectrum
for fresh TiO2 was extrapolated to 369.6 nm (Figure 3a),C2H5OH + O2 → CO2 + CH4 + H 2O
Δr H ° = −475.04 kJ mol−1(8)When the initial molar ratio of ethanol:O2 was changed from 1:2
to 1:4 (Figure 1b-iii), complete oxidation reaction over Ag
nanoparticles became predominant because suﬃcient O2 was
supplied for the reaction of eq 2 (ethanol:O2 = 1:3). Total Ccontaining products (185 μmol of CO2, 20.5 μmol of CH4; intotal 206 μmol) accounted for the complete conversion of
ethanol into products by the eﬀects of Ag, while a part of thehydroxy/water remained on the photocatalyst surface based on
the formed water (114 μmol, 1.38 kPa). The formed water was
still less than the water-saturated pressure (2.6 kPa).
A control experiment was performed to evaluate the eﬀect of
catalyst heating converted from irradiated light energy.15−17
When the catalyst was irradiated using UV−visible light while
keeping the reactor at 297 K using a water bath, CH4 was
negligibly formed (Figure 2b-ii). In contrast, consecutive
oxidation of C2H5OH to CH3CHO (eq 1) and then to CO2(eq 3) proceeded similar to the test under UV−visible light(Figure 2b-i). The reaction of eq 8 negligibly proceeded while
C2H5OH was consecutively oxidized to CO2 and H2O overTiO2 and Ag nanoparticles (Scheme 1B-a−e) at 297 K. In fact,
the formed CO2 was equal to that in Figure 2b-i if the amount of
CH4 was subtracted on the basis of the stoichiometry of eq 8.
The amount of water formed was half that in Figure 2a, implying
water adsorption over Ag nanoparticles at 297 K.
Next, the fast major complete oxidation reaction using
C2H5OH (1.33 kPa), O2 (2.67 kPa), and Ag (2.0 wt %)−TiO2
irradiated under UV−visible light was compared with that undervisible light, λ > 390 nm (Figure 2b-iii). Complete oxidation
and/or dissociative redox reactions were suppressed almost
completely, whereas 25.7 μmol of CH3CHO and 40.5 μmol of
H2O were formed until 4 h of reaction similar to the time course
irradiated under UV−visible light but in the absence of O2(Figure 1b-i). Thus, dehydrogenation of ethanol (eqs 1 and 7;
Scheme 1B-d and e) eﬀectively proceeded over the Ag
nanoparticles owing to the heat (see the EXAFS section
below) transformed from LSPR by the visible light irradiation.

Figure 3. UV−visible spectra for (a) TiO2 and (b−f) Ag (2.0 wt %)−TiO2 photocatalysts as (a, b) fresh and after photocatalytic tests under(c) C2H5OH (1.33 kPa), (d, f) C2H5OH (1.33 kPa) and O2 (2.67 kPa),and (e) C2H5OH (1.33 kPa) and O2 (5.32 kPa) irradiated by (c−e)
UV−visible light and (f) visible light (λ > 390 nm).

corresponding to the BG of 3.3 eV for the anatase phase. In the
spectrum for fresh Ag (2.0 wt %)−TiO2, the absorption edge
appeared to shift upward because of the very broad absorption in
the whole wavelength range by Ag2O nanoparticles.
All the spectra for Ag (2.0 wt %)−TiO2 photocatalysts used
for ethanol photooxidation reactions irradiated under UV and/or visible light exhibited a typical intense peak at 475−505 nm(Table 1), demonstrating that the Ag sites were reduced underethanol to the Ag0 state and exhibited LSPR when irradiated byvisible light. The peak maximum shifted toward a lower
wavelength as the Ag nanoparticle size became smaller becauseof the quantum size eﬀect except for one irradiated under visible
light38 (Table 1) as discussed in the EXAFS section.
Fluorescence Spectra. When the excitation was at 300 nm
for the TiO2 sample, corresponding to greater energy than thatof the BG excitation, major peaks due to the band-edge emission
appeared at 399 and 417 nm (Figure 4). By the addition of Ag
species [Ag2O for fresh Ag (2.0 wt %)−TiO2 sample], the peakintensities of major peaks as well as minor peaks at 452, 470, and573 nm related to interband levels39 were eﬀectively suppressed
owing to charge separation between TiO2 and Ag2O particles. A
new peak appearing at 376 nm for the Ag (2.0 wt %)−TiO2
photocatalyst used under C2H5OH and O2 (ratio 1:2) and UV−
visible light would be associated with the excitation at the TiO2
BG combined with LSPR excitation40 and/or the band-edge
emission due to more unsaturated TiO2.39
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ArticleTable 1. Coordination Number (N) of the Ag−Ag Interatomic Pair, Evaluated Ag Nanoparticle Size (d̅) and Dispersion (D)Based on EXAFS for Ag (2.0 wt %)−TiO2, and Absorption Peak Maximum in UV−Visible Spectroscopy
before light irradiation

10 or 20 min of lightirradiation

light

Na

d̅ (nm)b

Dc

Na

d̅
(nm)b

Dc

Na

d̅ (nm)bDc

λ > 390 nm
245 nm < λ < 386 nm, λ > 686 nm

3.2
3.2
2.5
3.2
4.0

0.65
0.65
0.50
0.65
0.80

1.0
1.0
1.0
1.0
1.0

8.3
5.9
5.5
6.6
5.4

1.9
1.1
1.1
1.2
1.0

0.64
0.89
0.93
0.82
0.94

10.1
9.6
8.6
8.7
8.1

3.6
3.1
2.1
2.2
1.8

0.38
0.45
0.60
0.59
0.66

gas (kPa)
entry

ethanol

O2

i
ii
iii
ii′
ii″

1.33

0
2.67
5.32
2.67

full light

converged @ >160 min light and light oﬀ
abs. peak max
(nm)
485
491
475
505

a

Obtained based on EXAFS. bd̅ values estimated based on the N value and ref 46. cD values estimated based on the N value and ref 46.

3d) on which O2-derived species adsorbed (see the following
EXAFS section).
TEM Images. The changes of the Ag (2.0 wt %)−TiO2
sample were observed by TEM. For the sample calcined at 673 K
for 2 h, nanocrystals of TiO2 (10−50 nm) were observed (Figure5a). In the HR-TEM image (Figure 5a, bottom), lattice fringesof anatase TiO2(101) were observed with the intervals of 0.349
nm (theoretical 0.3517 nm41), and those of rutile TiO2 (200)were also observed with the intervals of 0.230 and 0.223 nm(theoretical 0.22969 nm41). When the sample was underC2H5OH (1.33 kPa), O2 (2.67 kPa), and UV−visible light for 20
min, nanocrystals of TiO2 and the lattice fringes with theintervals of 0.349, 0.341, and 0.346 nm were observed due to theanatase TiO2(101) lattice (Figure 5b). In contrast, no images for
the Ag cluster/nanoparticle were observed.
For the sample under C2H5OH (1.33 kPa), O2 (2.67 kPa),
and UV−visible light for 5 h, the Ag(111) lattice was observedwith the intervals of 0.229 nm (theoretical 0.2359 nm32) as wellas the anatase TiO2(101) lattice with the intervals of 0.349 nm(Figure 5c). The Ag nanoparticles were detected in the TEM
image as indicated by arrows (Figure 5c, top).
EXAFS Monitoring. The Fourier transforms of the EXAFS
spectra monitoring for Ag (2.0 wt %)−TiO2 under various
photocatalytic reaction conditions (Figures 1b and 2b) aresummarized in Figure 6. Its associated k3-weighted EXAFS χ-

Figure 4. Fluorescence spectra for TiO2 and Ag (2.0 wt %)−TiO2
photocatalysts as fresh and after photocatalytic tests under (i) C2H5OH(1.33 kPa) and (ii) C2H5OH (1.33 kPa) and O2 (2.67 kPa). The
excitation wavelength was 300 nm.

The ﬂuorescence level became the lowest for the Ag (2.0 wt%)−TiO2 photocatalyst used under ethanol and UV−visible
light. The eﬀective charge transfer from TiO2 CB to metallic Ag
nanoparticles should be the reason as shown in the most intense
LSPR absorption peak at 485 nm under these conditions (Figure3c) and O-free metallic Ag surface in comparison to Ag
nanoparticles under C2H5OH, O2, and UV−visible light (Figure

Figure 5. TEM (top column) and high-resolution TEM (bottom column) images observed for Ag (2.0 wt %)−TiO2 samples (a) calcined at 673 K for
2 h and (b, c) sample (a) under C2H5OH (1.33 kPa), O2 (2.67 kPa), and UV−visible light for 20 min (b) and 5 h (c). The arrows in the TEM image
indicate Ag nanoparticles.
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ArticleFigure 6. Time course changes of the Fourier transform obtained from the k3-weighted Ag K-edge EXAFS χ-function for Ag (2.0 wt %)−TiO2 (a) freshin air and (b) under C2H5OH (1.33 kPa) irradiated by UV−visible light (i), C2H5OH (1.33 kPa) and O2 (2.67 kPa) by UV−visible light (ii), C2H5OH(1.33 kPa) and O2 (2.67 kPa) by visible light (λ > 390 nm) (ii′), C2H5OH (1.33 kPa) and O2 (2.67 kPa) by UV and IR light (245 nm < λ < 386 nm, λ >686 nm) (ii″), and C2H5OH (1.33 kPa) and O2 (5.32 kPa) by UV−visible light (iii) followed by dark conditions ((i), (ii), (ii′), (iii), and (c): reference
spectra for Ag metal foil and Ag2O powder).

Figure 7. Time course of the coordination number (N, top) and the Debye−Waller factor (σ, bottom) of the Ag−Ag interatomic pair in the EXAFS
measured for Ag (2.0 wt %)−TiO2 using (i) C2H5OH and UV−visible light, (ii) C2H5OH and O2 with the molar ratio of 1:2 and UV−visible light, (iii)C2H5OH and O2 with the molar ratio of 1:4 and UV−visible light, (ii′) C2H5OH and O2 with the molar ratio of 1:2 and visible light (λ > 390 nm), and(ii″) C2H5OH and O2 with the molar ratio of 1:2 and UV and IR light (245 nm < λ < 386 nm, λ > 686 nm).

demonstrating the presence of the Ag0 state, whereas the color
turned to gray under air overnight, indicating oxidation tofunctions are also listed in Figure S2. The Ag (2.0 wt %)−TiO2
sample prepared just after heating at 673 K was light yellow,
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AgI2O.25 The Ag2O formation was conﬁrmed in the Fouriertransform for the sample in air overnight (Figure 6a) comparedwith the data for the standard Ag2O sample (Figure 6c).
Under 1.33 kPa of ethanol during sample preparation andtransportation to the KEK beamline (∼12 h), a major part of
Ag2O was reduced to Ag0 before UV−visible light irradiation
based on the signiﬁcantly reduced peak intensity for Ag−O
interatomic pairs (Figure 6b-i, light on, 0 min). The weak peakbecame negligible within 10 min of UV−visible light irradiation,suggesting that most of the O sites (maximum 13.9 μmol) in
Ag2O were removed from Ag, whereas only 5.5 μmol of water
was formed at 30 min of reaction, and the amount gradually
increased until 4 h of reaction (32.1 μmol, Figure 1b-i). Thus,
Ag2O quickly reacted with C2H5OH, whereas CH3CHO andH2O desorption from the photocatalyst surface was slow. As the
Ag−O peak disappeared, the peak due to the Ag−Ag
interatomic pair gradually increased (Figure 6b-i), indicatingthat the Ag nanoparticle grew starting from the N value of 3.2
quickly to 8.3 and then converged to 10.1 (Figure 7i and Table
1i).
After the UV−visible light was turned oﬀ, the Ag−Ag peakbecame more intense because of the decrease of σ values within
10 min (Figures 6b-i and 7i, bottom), whereas the Ag−O peak
recovered to the level before the UV−visible light irradiation.
These facts suggested that the temperature decrease of Ag
nanoparticles and the partial recovery of Ag2O sites shifted fromthe dynamic charge equilibrium of the excited electron supply
from the TiO2 CB to Ag nanoparticles under UV−visible lightand the oxidation eﬀect probably by the O atoms at the interface
between Ag and TiO2.
Under 1.33 kPa of C2H5OH and 2.67 kPa of O2, Ag sites in theinitial Ag (2.0 wt %)−TiO2 sample were more reduced before
the UV−visible light irradiation based on the relatively weaker
Ag−O peak and stronger Ag−Ag peak intensities compared withthe data under ethanol only (Figure 6b-i and ii). However, underUV−visible light irradiation, the behavior of the Ag sites wasvery similar to that under the conditions of 1.33 kPa ethanol only
(Figure 6b-i). Ethanol is an eﬀective reducing agent for Ag sites;
however, it may adsorb more strongly on the surface, whereas
the hydrogen would combine with Ag2O sites more quickly inthe presence of an appropriate amount of O2 by forming H2O assuggested by the quicker H2O formation under C2H5OH and O2
versus that under C2H5OH (Figure 1b-i and ii). The N value for
the Ag−Ag interatomic pair increased from 3.2 to 5.9 and thenconverged to 9.6 (Figure 7ii and Table 1ii) in a similar N value
range to that under ethanol only (Figure 7i), whereas the σvalues were, in general, smaller than those under ethanol only
(Figure 7ii, bottom).
In contrast, under 1.33 kPa of C2H5OH and 5.32 kPa of O2(Figure 6b-iii), a relatively larger portion of Ag2O sites versus
Ag0 sites remained before and after the UV−visible lightirradiation (Figure 6b-iii) by the eﬀect of the extra O2 gas. The N
value for the Ag−Ag interatomic pair increased from 2.5 to 5.5
and then converged to 8.6 (Figure 7iii and Table 1iii),
signiﬁcantly smaller than values under C2H5OH or underC2H5OH and O2 (1:2 ratio) (Table 1i and ii. Thus, in thepresence of 5.32 kPa of O2 and UV−visible light irradiation, Ag0nanoparticles were smaller as a core and wrapped with a verythin Ag2O shell.
Under the conditions of 1.33 kPa of ethanol, 2.67 kPa of O2,
and visible light irradiation (λ > 390 nm) as a control test, theAg−O interatomic peak was weak/negligible, similar to themonitoring under UV−visible light (Figure 6b-ii and ii′);

Articlehowever, the Ag−Ag interatomic peak intensity did not increasemuch (Figure 6b-ii′) in comparison to the monitoring underUV−visible light (Figure 6b-ii). Conversely, after the UV−
visible light was turned oﬀ, a quick increase of the peak was
clearly observed within 20 min similar to all the other results
depicted in Figure 6b-ii′. When the irradiation was visible light
only (λ > 390 nm), the light negligibly activated TiO2, and theelectron ﬂow from the TiO2 CB to the Ag species was very
limited and the growth of Ag0 nanoparticles was signiﬁcantly
slower than that irradiated under UV−visible light.
Under these conditions, the N value increased starting from
3.2 to 6.6 and then converged to 8.7 (Figure 7ii′, top), smaller
than the corresponding value of 9.6 under C2H5OH (1.33 kPa)
and O2 (5.32 kPa) under UV−visible light (Figure 7iii, top andTable 1ii and ii′). The reason the σ values were larger after thelight was turned oﬀ than that before the irradiation (Figure 7ii′,
bottom) was due to the structural disorder factor as discussed
below.
Finally, on the basis of the curve-ﬁt analysis, under theconditions of 1.33 kPa of ethanol, 2.67 kPa of O2, and UV and IR
light irradiation (245 nm < λ < 386 nm, λ > 686 nm), the Ag−Ag
peak intensity gradually increased, greater than that under visible
light but smaller than that under UV−visible light (Figure 6b-ii,ii′, and ii″). This fact suggested that the electron transfer fromthe CB of TiO2 to Ag has a greater eﬀect than LSPR of Ag for theAg nanoparticle growth. In clear contrast to the otherexperiments using Ag (2.0 wt %)−TiO2 (Figure 6b-i, ii, ii′,
and iii), the Ag−Ag peak intensity negligibly changed when theUV−visible light was turned oﬀ (Figure 6b-ii′), suggesting that
the UV and IR light did not heat Ag sites.
The σ value was calculated to be 0.009 95 nm for Ag metal at
290 K using the correlated Debye model8,9,15 with the ab initio
multiple-scattering calculation code, FEFF8.4,42 and the Debye
temperature of Ag bulk (θD(bulk) 225 K).43 By the curve-ﬁtanalysis based on eq 5, the XDAP code31 provides anexperimental diﬀerence for the σ2 value from that of the Ag
metal (model). The initial σ value of 0.011 1 nm for Ag (2.0 wt%)−TiO2 under C2H5OH (1.33 kPa) before UV−visible lightirradiation quickly increased to 0.0123 nm (10 min of
irradiation) and in turn slightly decreased to 0.011 9 nm (150
min of irradiation) as shown in Figure 7i. Then, the value quickly
decreased to 0.0108 nm after the light was turned oﬀ at 170 min.
Furthermore, we evaluated the temperature at the Ag site
based on the σ values. The temperature dependence of the σvalue derived from FEFF8.4 combined with the correlated
Debye model8,9 for both bulk and surface Ag sites using the bulk
and surface Debye temperature was previously reported.15 We
assumed that the thermodynamically stable face-centered cubic(fcc) (111) face had preferable exposure for the latter value(θD(surf,⊥) 155 K).44,45 The Debye temperature for the lateraldegree of freedom for Ag atoms at the fcc Ag(111) surface(θD(surf, ∥)) was 226 K,45 very similar to the θD(bulk) (225 K).
Therefore, we only considered the vertical degree of freedom ata free hemispherical surface of the Ag sphere corresponding toθD(surf,⊥), whereas two lateral degrees of freedom wereconsidered to correspond to θD(bulk) if the Ag nanocrystals
were suﬃciently larger than a few nanometers. We also
approximated the mean Ag nanoparticle temperature as thearithmetic mean temperature based on the θD(surf,⊥) weighted by
1/2·1/3D [D: dispersion of nanoparticle (Table 1), for aneﬀective vertical degree of freedom at a free hemisphericalsurface] and that based on the θD(bulk) weighted by (1 − D) + 1/
2D + 1/2·2/3D for the bulk site, nonfree hemisphere in contact14696

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ArticleTable 2. Coordination Number (N) and Debye−Waller Factor (σ) of the Ag−Ag Interatomic Pair Obtained for EXAFS of Ag (2.0
wt %)−TiO2 (140 mg) (A) under C2H5OH (2.3 kPa) at the Controlled Temperatures at the Beamline versus (B) under Various
Photocatalytic Conditions for 160−200 min and then in the Dark for 40−100 min (Figure 7)
particle size and dispersion
entry

Tobs (K)aNbd̅ (nm)c

Dd

i
ii
iii
iv
v
vi
vii

301
301
343
373
403
343
301

4.9
5.7
7.1
8.4
10.4
10.3
11.0

0.97
1.1
1.4
2.0
3.9
3.8
4.5

0.99
0.91
0.77
0.63
0.33
0.34
0.24

i
ii
iii
ii′
ii″

300
300
300
300
300

10.1
9.6
8.6
8.7
8.1

3.6
3.1
2.1
2.2
1.8

0.38
0.45
0.60
0.59
0.66

correlated Debye modelσ (nm)b

σcorrelated_Debye (nm)e(A) Thermally Grown
0.011 30.010 80.011 20.011 20.011 50.012 40.011 80.011 10.010 90.010 60.010 2(B) Light-Induced Grown
0.010 70.010 30.010 60.010 30.010 70.010 40.011 10.010 40.011 00.010 4TD(bulk) (K)eTD(surf,⊥) (K)enf

σdisorder (nm)g

377.7
343.2
372.5
395.2
415.0
353.8
307.6

176.7
159.8
174.1
185.0
194.2
165.2
142.2

2.31
1.52
1.16
1
1
1
1

0
0
0
0
0.003 60.001 80.002 6

310.6
312.6
317.1
316.7
318.9

143.6
144.6
146.6
146.5
147.5

1
1
1
1
1

0.002 90.002 50.002 50.004 00.003 6aSample temperature (Tobs) during EXAFS measurements. bObtained based on EXAFS. cAg nanoparticle size (d̅) estimated based on the N value
and ref 46. dDispersion (D) estimated based on the N value and ref 46. eσ value, bulk temperature (TD(bulk)), and surface temperature(perpendicular degree of freedom, TD(surf,⊥)) corresponding to Tobs based on the correlated Debye model. fDegree of freedom eﬀectively similar tovertical motion among vertical and lateral motions for Ag atoms at Ag nanoparticle size. gStructural disorder factor (σdisorder) calculated based on eq
10.

with TiO2 and two lateral degrees of freedom at a freehemispherical surface.15
However, for small Ag nanoparticles of size less than 2 nm
based on the N values,46 the lateral motions cannot be regarded
as similar to that in the bulk but rather similar to vertical motion
toward the Ag particle surface. We introduced the value n as thedegrees of freedom for the surface Ag atom. Thus, in general, we
approximated the temperature by the arithmetic mean temperature based on the θD(surf,⊥) weighted by (1/2·n/3)D for degrees
of freedom eﬀectively similar to vertical motion among verticalsurfaces and lateral motions at a free hemispherical surface (1 ≤n < 3) and that based on the θD(bulk) weighted by (1 − D) + 1/2
D + (1/2·( 3− n)/3)D for the bulk site, nonfree hemisphere incontact with TiO2 and the rest of the degrees of freedom at a freehemispherical surface.
Teval = nD/6TD(surf, ⊥) + (1 − nD /6)TD(bulk)

Figure 8. Adjustment of degrees of freedom for surface Ag atoms tocorrespond to the Debye temperature θD(surf,⊥) rather than θD(bulk) toevaluate the Ag site temperature based on eq 9 based on EXAFS
measurements for Ag (2.0 wt %)−TiO2 under C2H5OH at 301−373 K(Table 2).

(9)

To evaluate the dependence of n values on the Ag nanoparticle
size, several Ag K-edge EXAFS spectra were measured for Ag(2.0 wt %)−TiO2 under C2H5OH (1.33 kPa) by heating from
301 to 403 K and then back to 301 K in the dark (Table 2A-i−vii), and for a mean Ag particle size of 0.97−2.0 nm, the n values’
dependence was drawn (Figure 8). When the mean Ag particle
size was greater than 2.0 nm [N(Ag−Ag) > 8.4], the n value was1.
The temperature-controlled EXAFS data for Ag (2.0 wt %)−TiO2 under C2H5OH (1.33 kPa) required consideration of thestructural disorder factor (σdisorder)47 once the sample was heatedat 403 K as in eq 10 owing to the increase of structural disorderfor thermally grown Ag nanocrystals of mean size of 3.9−4.5 nm:
σ 2 = σcorrelated_Debye 2 + σdisorder 2suggested a relatively unsaturated metallic Ag surface undercatalytic conditions.48
As a result, using Ag (2.0 wt %)−TiO2 (0.100 g), C2H5OH(1.33 kPa), and UV−visible light, the initial temperature of 290
K before irradiation rose to 423 K after 10 min of irradiation andthen reached an equilibrium temperature of ∼404 K after 50 minof light irradiation (Figure 7i, bottom). The initial CH3CHO
formation rate via eq 7 was 15.3 μmol h−1, corresponding to an
exothermic reaction at the rate of 2.17 J h−1 (Table 3i). On thebasis of the molar heat capacity of Ag and TiO2 (Cp,m° = 25.350
and 55.100 J K−1 mol−1, respectively, 0.100 g of photocatalystused),49,50 the heat of reaction for CH3CHO formationaccounted for sample heating at a rate of 16 K per 30 min(Table 3i). This temperature accounts for the diﬀerence
between the initial temperature peak at 10 min (423 K) andthe equilibrium temperature (404 K) at 50−150 min of(10)

Assuming σdisorder values of 0.0018−0.0036 nm for such
samples (Table 2A-v−vii), σcorrelated_Debye values corresponded
well to the experimental temperature based on the approximation of eq 9 (Table 2A-iv−vii). The increase of structural
disorder for Ag nanoparticles under C2H5OH (1.33 kPa)
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ArticleTable 3. Heat of Reaction (Q) in the Initial 30 min of Reaction and Heat De-excited via LSPR under Photocatalytic Conditions ofEthanol Oxidation Using Ag (2.0 wt %)−TiO2 (0.100 g)
gas (kPa)
ent.

ethanol

O2

i
ii

1.33

0
2.67

iii

5.32

ii′
ii″

2.67

light

product

reactionrate(μmol h−1)

T increase rate by Q
(K min−1)

T increase rate by LSPR
(K min−1)

eq 7
eq 1
eq 2
eq 1
eq 2
eq 7
eq 2

15.3
64.2
51.3
71.0
64.0
18.6
9.4

0.54
11

49

λ > 390 nm
245 nm < λ < 386 nm, λ > 686 nm

CH3CHOCH3CHO
CO2
CH3CHO
CO2
CH3CHO
CO2

full light

13
0.65
1.5

0

common eﬀect of surface O2-derived species on Ag reducing
LSPR (Figure 3d and e).
It is implausible that surface O2-derived species transferredover metallic Ag nanoparticles of 1.1−3.1 nm eﬀectively block
the irradiation of visible light of λ = 400−600 nm (Table 1ii and
iii). Thus, surface O2-derived species eﬀectively suppressed the
LSPR; the collective translational movement of surface Ag
atoms and the light energy dissipated to TiO2 and then the glass
reactor resulting in lower equilibrium temperatures of 363−368
K in comparison to 404 K under ethanol only (Figure 7,
bottom). Such a comparison was enabled by the local site
temperature monitoring based on the correlated Debye modelfor the σ value of the local site in contrast to SERS monitoring,which requires a probe molecule;18−20 phonon monitoring by
Raman, which requires Si dots;21 Raman spectroscopy, which
requires locally doped ions;22 and absorption spectroscopy,
which requires laser irradiation.23
The Ag site temperature change under C2H5OH (1.33 kPa),
O2 (2.67 kPa), and visible light (λ > 390 nm) was interesting tocompare with the other changes in view of the photocatalytic
activation mechanism. Under the reaction conditions, O2
activation over the TiO2 surface was suppressed in the absenceof UV light. The partial oxidation to CH3CHO via eq 7 (Figure
1b-ii, right) accounted for the temperature increase of 0.65 K
min−1, whereas visible light led to a temperature increase of 49 K
min−1 if all the light was absorbed (Table 3ii′). Thus, thetemperature increase to 411 K after 20 min of reaction was due
to major light energy via Ag LSPR. The equilibrated temperaturewas ∼406 K, similar to ∼404 K under ethanol and UV−visible
light (Figure 7i and ii′, bottom). Thus, irradiated under visible
light, O2 was not activated over TiO2, and surface O-free Ag
nanoparticles were more eﬀectively heated owing to LSPR. The
broader Ag LSPR peak under the conditions suggested thediﬀerence of light absorption compared to that irradiated underUV−visible light (Figure 3d and f).
In clear contrast to all the other EXAFS monitoring (Figure 7i,
ii, iii, and ii′), the Ag site temperature progressively increased
under C2H5OH (1.33 kPa), O2 (2.67 kPa), and UV and IR light(245 < λ < 386 nm and λ > 686 nm) (Figure 7ii″) by a selfacceleration mechanism due to the heat of reaction corresponding to a temperature increase of 1.5 K min−1 whereas there wasno heat eﬀect by light because LSPR did not occur in the absenceof visible light irradiation (Table 3ii″). The catalyst heated byexothermic heat progressively increased the reaction rate.
In view of the accuracy of the site temperature evaluation
based on the Debye−Waller factor, the method has been
conﬁrmed by tests on Ni, Mo, and Ag metal foils. The deviation
between experiment and theory was already proved to be lessthan <8 K for the metal foils in the temperature range of 300−

photoreaction, the major energy source for heating Ag being
light.
The Ag nanoparticles absorb light in the wavelength range
400−600 nm17,28 (Figure 3) due to LSPR that de-excited toresult in heat at a rate of 49 K min−1 (Table 3i). Thus, Ag
nanoparticles were heated to the equilibrium temperature 404 K
due to light energy and the heat dissipation to TiO2 and then theglass reactor. The heat energy for CH3CHO formation (eq 7)
accounted for the initial extra temperature rise by 19 K
compared with the equilibrium temperature of 404 K.
Under C2H5OH and O2 with the molar ratio of 1:2, CH3CHO
and CO2 were initially formed via eqs 1 and 2 at the rates of 64.2
and 51.3 μmol h−1, respectively (Figure 1b-ii). The total heat of
the reaction accounted for a temperature rise of 11 K min−1(Table 3ii). The heat converted from Ag LSPR also heated theAg nanoparticles at a rate of 49 K min−1 if all the light ofwavelength 400−600 nm was absorbed. The temperature of theAg nanoparticles soon reached equilibrium at ∼363 K within 10
min of reaction, signiﬁcantly lower than those under ethanol
only by 41 K (Figure 7i and ii, bottom). The formation of
CH3CHO and CO2 stopped after 2 h of reaction because thereactants were totally consumed in the exothermic reactions of
eqs 1, 2, and 8 (Figure 1b-ii).
In contrast, the equilibrium temperature remained constant at∼363 K irrespective of the heat of reaction. Thus, the Ag site
temperature rise and drop under light on and light oﬀ was mostly
due to heat converted from Ag LSPR. A part of the heat derived
from light energy and exothermic reaction was also used for Ag
nanoparticle growth. The Ag sites were mostly the Ag0 state
during the EXAFS monitoring under UV−visible light (Figure
2b-ii); however, the LSPR peak intensity was weaker underC2H5OH and O2 (1:2) compared with that under C2H5OH only
(Figure 3c and d). Thus, surface O2-derived species on Ag
suppressed the light absorption by LSPR and lowered theequilibrated temperature due to the balance of mostly heat deexcited from LSPR and heat dissipation to ∼363 K (Figure 7ii,
bottom).
Under C2H5OH and O2 with the ratio of 1:4, initial CH3CHO
and CO2 formation rates were 71.0 and 64.0 μmol h−1,
respectively, which accounted for a temperature rise of 13 K
min−1 (Table 3iii). The temperature reached a maximum of 386
K at 20 min of reaction and equilibrated at ∼368 K. Similar to
the monitoring under C2H5OH only and C2H5OH and O2(1:2), the temperature rise was mostly due to heat converted
from Ag LSPR, but the initial maximum temperature was higher
than equilibrium by 18 K due to the extra heat of reaction fromfaster CH3CHO and CO2 formation under C2H5OH and O2
with the ratio of 1:4 (Figure 1b-iii). The similar equilibrium
temperature, 368−363 K, under C2H5OH:O2 = 1:2 was by the14698

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400 K.17 Our tentative study suggested that the Debye−Waller
factor for Zr of ZrO2 was negligibly changed by the UV−visible
light irradiation15−17 as well as that for Ti of TiO2 in this study.
As the Ag volume percent in Ag (2.0 wt %)−TiO2 is only 0.80%,
the heat exclusively transformed from LSPR at the Ag
nanoparticle surface quickly dispersed to supporting TiO2, andonly a part of the light energy contributed to the dynamic
temperature equilibrium at Ag (404−363 K).
This study extended the application to Ag nanoparticles
growing from 0.50 to 3.6 nm in size under photocatalytic
conditions (Table 1). By the introduction of an empirical n value
for the degrees of freedom of surface Ag sites in eq 9 dependent
on the Ag nanoparticle dispersion, the diﬀerence of theequilibrated temperature under light was demonstrated: 404−
406 K under reductive conditions allowed partial oxidation in
the absence of O2 or the absence of O2 activation (no UV light)(Figure 7i and ii′) versus 363−368 K under oxidative conditions
to allow complete oxidation in the presence of both O2 and UVlight (Figure 7ii and iii).
The remaining diﬃculty in determining the temperature forgrowing Ag nanoparticles under photocatalytic conditions is the
structural disorder factor (eq 10). The factor σdisorder was
evaluated for a thermally heated Ag (2.0 wt %)−TiO2 sample
under C2H5OH at 403 K as reference tests: 0.0018−0.0036 nm(Table 2A-v−vii). The sample temperature after the UV and/or
visible light was turned oﬀ in Figure 7 should converge soon to
300 K (ambient temperature inside the beamline after
photocatalytic tests), whereas the evaluated temperature basedon EXAFS was 326, 319, 317, 344, and 336 K (Figure 7). This isdue to the increase of σdisorder, and the value was evaluated to
0.0025−0.0040 nm due to the light-driven growth of Ag
nanocrystals (Table 2B) in a similar range to 0.0018−0.0036 nm
due to furnace heat-driven growth of Ag nanocrystals (Table
2A). This factor leads to a maximum 20 K overestimation for theAg site temperature during light irradiation. Still, the eﬀect of
surface O species on Ag LSPR and accordingly heatedtemperature is certain in the comparison of the equilibrated
temperature between 404−406 and 363−368 K. This method is
expected to apply to real-time monitoring of catalysis thatcritically depends on the nanoparticle size, e.g., Au.51

Articleequilibrated temperature was 363−368 K under O2 and UVlight, lower by suppressed Ag LSPR by O2-derived species
activated on TiO2 and transferred to Ag to allow complete
oxidation to CO2 and the other product due to C−C bondcleavage: CH4. In oxidative conditions, O2-derived species were
not detected by EXAFS curve-ﬁt analysis but could bemonitored as the temperature decreased by 40 K to suppress
Ag LSPR in accord with the LSPR peak intensity change underphotocatalytic conditions.

■

ASSOCIATED CONTENT

sı Supporting Information*

The Supporting Information is available free of charge athttps://pubs.acs.org/doi/10.1021/acs.jpcc.1c04076.
Experimental details for photocatalytic tests and synchrotron X-ray measurements and the summary of EXAFS χfunction oscillations (PDF)■

AUTHOR INFORMATION

Corresponding Author

Yasuo Izumi − Department of Chemistry, Graduate School of
Science, Chiba University, Chiba 263-8522, Japan;
orcid.org/0000-0001-8366-1864; Phone: +81-43-2903696; Email: yizumi@faculty.chiba-u.jp; Fax: +81-43-2902783

Authors

Daiki Fukuhara − Department of Chemistry, Graduate Schoolof Science, Chiba University, Chiba 263-8522, Japan
Moses Tuhafeni Joseph − Department of Chemistry, Graduate
School of Science, Chiba University, Chiba 263-8522, Japan;
orcid.org/0000-0003-4663-1628
Tarik Loumissi − Department of Chemistry, Graduate School of
Science, Chiba University, Chiba 263-8522, Japan;
orcid.org/0000-0002-4840-8625
Chao Zhang − Department of Chemistry, Graduate School of
Science, Chiba University, Chiba 263-8522, Japan
Takaomi Itoi − Department of Mechanical Engineering,
Graduate School of Engineering, Chiba University, Chiba 2638522, Japan; orcid.org/0000-0002-2864-6650
Hongwei Zhang − Department of Chemistry, Graduate Schoolof Science, Chiba University, Chiba 263-8522, Japan;
orcid.org/0000-0002-6773-8769
Complete contact information is available at:https://pubs.acs.org/10.1021/acs.jpcc.1c04076■

CONCLUSIONSIn this paper, the reactivity and selectivity in C2H5OH
photooxidation using Ag−TiO2 photocatalysts were correlated
to the temperature monitoring of Ag nanoparticles during thephotocatalytic tests. On the basis of kinetic tests by varying thereactant ratio (C2H5OH:O2) between 1:0 and 1:4, a major roleof TiO2 was the O2 activation by UV light, whereas the Ag
nanoparticle surface eﬀectively allowed partial oxidation of
C2H5OH to CH3CHO and subsequent C−C bond cleavage to
CO2 and CH4 by the combination of O species activated overTiO2.
On the basis of the correlated Debye model for surface andbulk Ag sites in Ag nanoparticles, the Ag site temperature during
the C2H5OH photooxidation reaction was monitored. In thispaper, the temperature for Ag nanoparticles growing from 0.5 to
3.6 nm was successfully monitored on the basis of the evaluationof the eﬀective degrees of freedom for translational motion of
surface Ag sites as well as the correlated Debye model. The
equilibrated temperature during photocatalysis was as high as
404−406 K to allow partial oxidation to CH3CHO over the Ag0
surface heated by the converted energy from Ag LSPR due tovisible light when O2 or UV light was absent. In contrast, theAuthor Contributions

D.F., M.T.J., T.L., T.I., C.Z., and H.Z. wrote part of themanuscript. Y.I. supervised the research and wrote most of themanuscript.
Notes

The authors declare no competing ﬁnancial interest.

■

ACKNOWLEDGMENTS
The authors are grateful for the ﬁnancial support obtained fromGrants-in-Aid for Scientiﬁc Research B (20H02834) and C(17K05961) from the Japan Society for the Promotion of
Science. X-ray absorption experiments were performed with theapproval of the Photon Factory Proposal Review Committee(2020G676, 2019G141, and 2018G649). The authors thankEnago (www.enago.jp) for the English language review.
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