Mechanistic interpretation of CO oxidation turnover rates on supported Au clusters

Article abstract of DOI:10.1016/j.jcat.2011.09.015

Kinetic and isotopic data are used to interpret the mechanistic role of gaseous H₂O molecules and of non-reducible (Al₂O ₃) and reducible (TiO₂, Fe₂O₃) supports on CO oxidation turnovers catalyzed by small Au clusters (<5 nm). H₂O acts as a co-catalyst essential for O₂ activation and for catalyst stability in CO oxidation at near-ambient temperatures, but also inhibits rates via competitive adsorption at higher H₂O pressures. The effects of CO, O₂, and H₂O pressures on CO oxidation turnover rates, the absence of ¹⁶O₂/¹⁸O ₂ and ¹⁶O₂/H₂¹⁸O exchange, and the small H₂O/D₂O kinetic isotope effects are consistent with quasi-equilibrated molecular adsorption of CO, O ₂, and H₂O on Au clusters with the kinetic relevance of H₂O-mediated O₂ activation via the formation of hydroperoxy intermediates (OOH), which account for the remarkable reactivity and H₂O effects on Au clusters. These elementary steps proceed on Au clusters without detectable requirements for support interface sites, which are no longer required when H₂O is present and mediates O₂ activation steps. Rate enhancements by H₂O were also observed for CO oxidation on Pt clusters (1.3 nm), which is also limited by O₂ activation steps, suggesting H₂O-aided O₂ activation and OOH species in oxidations involving kinetically-relevant O₂ activation. These intermediates have also been proposed to account for the ability of O₂/H₂O mixtures to act as reactants in alkene epoxidation on Au-based catalysts.

Full text of DOI:10.1016/j.jcat.2011.09.015

Journal of Catalysis 285 (2012) 92–102
Contents lists available at SciVerse ScienceDirect
Journal of Catalysis
journal homepage: www.elsevier.com/locate/jcat
Mechanistic interpretation of CO oxidation turnover rates on supported Au clusters
1
2
⇑
Manuel Ojeda , Bi-Zeng Zhan , Enrique Iglesia
Department of Chemical Engineering, University of California and E.O. Lawrence Berkeley National Laboratory, Berkeley, CA 94720, United States
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 18 July 2011
Revised 13 September 2011
Accepted 14 September 2011
Available online 19 October 2011
Kinetic and isotopic data are used to interpret the mechanistic role of gaseous H O molecules and of non-
2
reducible (Al
clusters (<5 nm). H
dation at near-ambient temperatures, but also inhibits rates via competitive adsorption at higher H
pressures. The effects of CO, O , and H O pressures on CO oxidation turnover rates, the absence of
O exchange, and the small H O/D O kinetic isotope effects are consistent with
quasi-equilibrated molecular adsorption of CO, O , and H
of H O-mediated O
for the remarkable reactivity and H
ters without detectable requirements for support interface sites, which are no longer required when H
is present and mediates O activation steps. Rate enhancements by H O were also observed for CO oxi-
activation steps, suggesting H O-aided O acti-
activation. These intermediates
2
3
O ) and reducible (TiO
2
, Fe
2 3
O ) supports on CO oxidation turnovers catalyzed by small Au
2
O acts as a co-catalyst essential for O
2
activation and for catalyst stability in CO oxi-
2
O
2
2
1
6
18
16
18
2
O /
O
2
and
2
O /H
2
2
2
Keywords:
CO oxidation
Au
H O effect
2
Mechanism
2
2
O on Au clusters with the kinetic relevance
activation via the formation of hydroperoxy intermediates ( OOH), which account
ꢀ
2
2
2
O effects on Au clusters. These elementary steps proceed on Au clus-
2
O
Hydroperoxy
2
2
dation on Pt clusters (1.3 nm), which is also limited by O
vation and OOH species in oxidations involving kinetically-relevant O
2
2
2
ꢀ
2
have also been proposed to account for the ability of O
idation on Au-based catalysts.
2 2
/H O mixtures to act as reactants in alkene epox-
Published by Elsevier Inc.
1
. Introduction
catalysts with similar cluster size, Au content, support material,
and synthetic provenance [1,9]. The absence of consensus about
these catalyst properties reflects, at least in part, the presence of
The catalytic oxidation of CO has been extensively used to probe
mechanisms and site requirements in heterogeneous catalysis and
to remove CO from combustion effluent and H -containing streams
1–3]. Au-based materials catalyze CO oxidation at sub-ambient
temperatures when present as small Au clusters (<5 nm) dispersed
on oxide supports (TiO , Fe , Al , etc.) [4–8]. The high reactiv-
adventitious H
for catalysis. Previous studies have concluded that H
[5,10,11], inhibits [12], or does not affect [13] CO oxidation cataly-
sis. The most extensive studies [10,14] included Au clusters depos-
ited on different supports (TiO
2
O at levels difficult to detect but consequential
O promotes
2
2
[
2
2
O
3
2
O
3
2 2 3 2
, Al O , and SiO ) and concluded
ity of small Au clusters in CO oxidation catalysis has been variously
attributed to the following: (i) low-coordination Au surface atoms,
isolated Au atoms, or cationic Au species; (ii) electron transfer from
the support to Au clusters; or (iii) supports that become active
when placed in atomic contact with Au clusters [1]. The impor-
tance and relative contributions of these effects are controversial,
and the specific evidence for each proposal remains equivocal
and often contradictory.
that H O impurities in CO/O reactants strongly increase CO oxida-
tion rates, even at trace concentrations (0.1–1 ppm), which are
undetectable by conventional analytical methods, and to very dif-
2
2
2
ferent extents on the various supports [14]. These H O effects have
been attributed to the promotion of O adsorption and/or dissoci-
2
ation, decomposition of unreactive carbonate deposits, and assis-
tance in reducing inactive Au cations to Au0 by CO [7,11,14–17],
but their respective contributions remain the subject of persistent
speculation and active debate. Recent reports also show that
increasing concentrations of hydroxyl groups favor CO oxidation
Large differences in measured rates (>10-fold), in the observed
effects of reactant and H
order ranges of 0.02–1.01 and 0.07–0.46 for CO and O
tively), and in deactivation rates have been reported on Au
2
O concentrations on such rates (kinetic
2
, respec-
rates, either in alkaline aqueous phase using Au/TiO
2
and Au/C cat-
alysts [18] or in the gas phase with Au/SiO
NaOH [19].
2
catalysts doped with
Every plausible sequence of CO oxidation elementary steps and
active sites has already been proposed, based on direct or indirect
evidence, without even a modicum of consensus and with the
implication of structures ranging from cationic or metallic Au spe-
cies in some cases, but not others, in chemical or physical synergy
⇑
Corresponding author. Fax: +1 510 642 4778.
E-mail address: iglesia@berkeley.edu (E. Iglesia).
Present address: Institute of Catalysis and Petrochemistry (CSIC), C Marie Curie 2,
1
2
8049 Madrid, Spain.
2
Present address: Chevron Corporation, Richmond, CA 94802, United States.
0
021-9517/$ - see front matter Published by Elsevier Inc.
doi:10.1016/j.jcat.2011.09.015

M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
93
with reducible or non-reducible supports, and with H
promoter, participant, inhibitor, or spectator species [1,5,16,20–
3]. The only complete sequence of plausible elementary steps
supported by the comparison with rigorous rate data, in our
assessment, involves parallel competitive and non-competitive
2
O acting as a
measured rates on reactant pressure in the presence and in the ab-
sence of co-fed H O; these materials were also used to evaluate CO
oxidation pathways operating with Au clusters deposited on
reducible and non-reducible metal oxide supports. The mean Au
cluster size (dAu) derived from TEM images are 3.3 ± 0.7 and
2
2
adsorption of CO and O
steps [24–26]; this mechanistic proposal did not consider what is
currently accepted as a critical and perhaps even essential role of
2
and their reactions in kinetically-relevant
3.6 ± 0.7 nm for Au/TiO
A Pt/Al catalyst (2.03% wt., Pt clusters of 1.3 nm) was also
used in reactions of CO–O and CO/O /H O mixtures to probe any
effects of H O on CO oxidation rates on Pt clusters. The synthesis
2 2 3
and Au/Fe O , respectively.
2 3
O
2
2
2
2
H O
in mediating CO oxidation catalysis at near-ambient
2
temperature.
and characterization protocols used for this Pt catalyst have been
reported elsewhere [31].
We provide here kinetic and isotopic evidence for a sequence of
elementary steps for CO oxidation on stable Au catalysts with
added H
2
O, present as a co-catalyst essential for stable turnover
2.2. Steady-state CO oxidation rate measurements
ꢀ
rates. These steps are mediated by hydroperoxy species ( OOH),
which may form on Au surfaces [27], and which account for the
CO oxidation rates were measured in a tubular packed-bed
reactor with plug-flow hydrodynamics. Typically, catalysts (25–
30 mg, 0.250–0.425 mm pellet size) were diluted with quartz gran-
co-catalytic effect of H
bility and for the ability of these materials to catalyze propene
epoxidation with H O/O mixtures [28]. These kinetic data were
obtained under conditions of strict kinetic control and without
detectable deactivation in the presence of H O as an essential com-
ponent; the data and their mechanistic interpretation aim to re-
solve long-standing controversies that reflect, at least in part, the
instability and mechanistic promiscuity imposed by uncontrolled
2
O on CO oxidation rates and on catalyst sta-
2
2
ules (ꢃ1 g; washed with 1 M HNO
3
at 298 K for 2 h and then trea-
ted in ambient air at 1023 K for 5 h). Samples were treated in
3
ꢂ1 ꢂ1
2
flowing pure H
2
(28 cm s
g , 99.999%, Praxair) at 373 K (heat-
ꢂ1
3 ꢂ1 ꢂ1
ing rate of 0.167 K s ) for 0.5 h and in H
2
O/H
2
(28 cm s
g ,
2
1 vol.% H O) at 373 K for 0.5 h, using a previously reported proce-
dure that forms stable Au metal clusters [32]. The catalyst was
brought to the reaction temperature (282–303 K) in flowing He
and undetected concentrations of adventitious H
2
O in catalytic sol-
ids and reactant streams.
(99.999%, Praxair). Gas reactants (10 vol.% CO in He, 25 vol.% O
in He, UHP grade, <10 ppm H O, Praxair) were metered by elec-
tronic controllers, and H O (doubly-distilled and deionized) was
introduced into heated transfer lines using a syringe pump (Cole
Parmer 74900 Series). Helium was used as balance. CO, O , and
2
2
2
2
. Experimental methods
2
2
.1. Catalyst preparation and characterization
He streams were further purified with moisture traps (Matheson
Tri-Gas). Reactants and products concentrations were measured
by gas chromatography (Agilent 6890 GC) using a Porapak Q
packed column (80–100 mesh, 1.82 m ꢄ 3.18 mm) connected to a
thermal conductivity detector (TCD). All transfer lines were heated
Al
prepared by the deposition–precipitation (DP) methods [5,20]. Tet-
rachloroauric acid hydrate (0.24 g, HAuCl O, 99.999%, Aldrich)
ꢁxH
was dissolved in doubly-distilled deionized H
-Al (5 g, Alcoa) was treated in flowing dry air (1.67 cm s
UHP grade, Praxair) at 923 K for 5 h, and then dispersed in doubly-
2 3 2 3
O -supported Au clusters (denoted here as Au/Al O ) were
4
2
3
2
O (80 cm ) at 353 K.
to 400–415 K to prevent H
were detected in the effluent when reactors contained only quartz
diluent or Al supports. CO conversions were kept well below
10% in most experiments (and below 15% in cases) by changing
residence times while keeping CO, O , and H O concentrations at
2
O condensation. No reaction products
3
ꢂ1 ꢂ1
c
2
O
3
g ,
2 3
O
3
distilled deionized H
-Al surface at 353 K and a pH of 7 (adjusted with 0.5 M NaOH,
98%, Fluka) by stirring for 1 h. The resulting solids were rinsed
and washed with doubly-distilled deionized H O (323 K) and dried
2
O (120 cm ) at 353 K. Au was deposited onto
c
>
2
O
3
2
2
constant values. These low conversions ensure the differential nat-
ure of measured rates and avoid their rigorous but more cumber-
some interpretation in terms of integral equations.
2
at ambient temperature for 24 h. Samples were stored away from
light without further treatment. The Au content was measured by
inductively-coupled plasma emission spectroscopy (0.61% wt.; Gal-
braith Laboratories, Inc.). The mean diameter of these Au clusters
Propene epoxidation rates using O
sured on Au/TiO , while HCOOH dehydrogenation and water–gas
shift reactions were performed with all Au/Al samples. Gas
reactants (C , 25 vol.% O in He, 10 vol.% CO in He, UHP grade,
Praxair) were metered by electronic controllers, and liquids
HCOOH or H O (doubly-distilled deionized) were introduced using
2 2
/H O mixtures were mea-
2
2 3
O
(
3.5 ± 1.2 nm) was determined by high-resolution transmission
H
3 6
2
electron microscopy (TEM) and reported elsewhere [29]. Mean Au
diameters (dAu) were calculated using dAu
d
of Au clusters of diameter d
sured by N
ASAP 2000 apparatus. Pore size distributions were obtained from
these adsorption data using the Barrett–Joyner–Halenda (BJH)
equation [30].
P
P
3
i
2
¼
n
i
d =
n
i
d , where
2
i
i
is the diameter measured from TEM images, and n
i
is the number
a syringe pump. Reactants and products concentrations were mea-
sured with a mass spectrometer (Inficon Transpector) and a gas
chromatograph (Hewlett–Packard 5890) equipped with a Porapak
Q packed column (80–100 mesh, 1.82 m ꢄ 3.18 mm) connected
to a thermal conductivity detector and a HP-1 capillary column
i
. The pore size distribution was mea-
2
adsorption–desorption at 77 K using a Micromeritics
(50 m ꢄ 0.32 mm ꢄ 1.05
lm) connected to a flame ionization
Three different portions of the Au/Al
2
O
3
solids were treated in
detector (FID).
3
ꢂ1 ꢂ1
O
2
/He (25 vol.%, 25 cm g
s
) by increasing the temperature
ꢂ1
from ambient to 873, 950, or 1023 K at 0.17 K s and holding at
each temperature for 2 h. These samples are denoted as treated
2.3. Isotopic exchange rates and kinetic isotope effects
Isotopic 16O /18 and 16 18O exchange rates were mea-
catalysts, Au/Al
ture (X = 873, 950, or 1023), and the Au/Al
ent temperature is named as untreated Au/Al
Two reference Au catalysts (1.56% wt. Au/TiO
Au/Fe , prepared by deposition–precipitation and co-precipita-
tion, respectively) were provided by the World Gold Council
WGC). These two samples were used to examine the effects of
support on CO oxidation rates and on the kinetic dependence of
2
O
3
-X, where X represents the treatment tempera-
2
O
2
2 2
O /H
2
O
3
solid dried at ambi-
sured on Au/Al using a glass recirculating flow reactor [33].
2 3
O
2
O
3
.
Samples were treated as described above, and the recirculating
3
2
and 4.44% wt.
and reactor volumes (550 cm total) were evacuated with mechan-
2
O
3
ical and diffusion pumps before introducing reactants as vapors
1
6
18
18
18
(
2
O , 99.999%, Praxair;
2 2
O , 99%, Isotec; H O, 99 at.% O, Sig-
(
ma–Aldrich). Chemical and isotopic compositions were measured
by direct sampling into a gas chromatograph (Hewlett–Packard

9
4
M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
5
890) equipped with a thermal conductivity detector and an elec-
4.0
2
tron-impact mass selective detector (Hewlett–Packard 5972). D O
3
3
2
2
1
1
0
0
.5
.0
.5
.0
.5
.0
.5
.0
(
99.9% D, Cambridge Isotope Laboratories, Inc.) was used to mea-
sure H O/D O kinetic isotope effects (KIE) during CO oxidation
2
2
using the tubular plug-flow reactor described in the previous
section.
3
. Results and discussion
3
.1. Effects of H O on CO oxidation rates and catalyst stability
2
Fig. 1 shows CO oxidation rates at 288 K on untreated Au/Al
with and without added H O. H O strongly influenced CO oxida-
tion turnover rates (per surface gold atom, Au ; from mean TEM
cluster sizes, 3.5 ± 1.2 nm), as also shown previously [10]. CO oxi-
2 3
O
2
2
s
0.0
0.2
0.4
0.6
0.8
1.0
1.2
ꢂ1
ꢂ1
dation rates did not vary ð2:55 ꢅ 0:10 mol s g-at Aus Þ with
time-on-stream for the duration of this experiment (ꢃ25 ks;
Fig. 1) or of similar experiments on all other catalysts reported here
H O pressure (kPa)
2
Fig. 2. Influence of H O pressure (0.03–7.15 kPa) on CO oxidation turnover rates
2
when H
2
O (0.5 kPa) was present in CO/O
2
reactant mixtures
measured with the untreated Au/Al
size: 0.125–0.250 mm (d); 0.250–0.425 mm ( ).
2 3 2
O (0.61% wt., 288 K, 5 kPa CO, 2 kPa O ). Pellet
(
5 kPa CO; 2 kPa O
2
). These conditions led to stable CO oxidation
rates at levels that rank among the highest reported on Au-based
catalysts [9]. In contrast, CO conversion rates decreased rapidly
oxidation in the absence of H
implicate, without specific experimental evidence, a role of H
in elementary steps required for O activation, possibly via the for-
mation of activated oxygen species that react rapidly with CO to
form intermediates that ultimately decompose to CO [14,35].
The possible involvement of H O in the kinetically-relevant steps
in CO oxidation catalytic sequences is examined next by comparing
CO oxidation turnover rates (288 K, 5 kPa CO, 2 kPa O ) measured
on Au/Al using H O or D O (0.5 kPa) as additives. As we show
later (Section 3.4), the small measured H O/D O isotope effects
2
O [16,34]. In contrast, other studies
ꢂ1
ꢂ1
to ꢃ 0:08 mol s g-at Aus (first-order deactivation rate constant,
2
O
ꢂ1
ꢂ1
k
d
, 0.76 ks initially and 0.21 ks after 3 ks on stream) when H
was removed from the reactant stream and the concentration of
residual H O decreased with time. The initial CO oxidation rates
were fully recovered upon reintroduction of H O (0.5 kPa) into
the reactant mixture (Fig. 1), indicating that the effects of H
2
O
2
2
2
2
2
2
O
2
were reversible and that the absence of H O did not cause sintering
2
or other irreversible structural changes.
O
2 3
2
2
The effects of H
turnover rates are shown in Fig. 2. Turnover rates initially in-
creased with H O pressure at low concentrations (<0.2 kPa) and
reached a maximum value at ꢃ0.5 kPa H O; higher H O concentra-
tions inhibited CO oxidation reactions. These positive and revers-
ible effects of H O are consistent with some previous studies
16,20,34]; the inhibition observed at higher pressures may ac-
count for persistent contradictory reports about the specific effects
of H O on Au-catalyzed CO oxidation [12]. The promotion effects of
O are not caused by parallel water–gas shift (WGS) reactions
CO + H O ? CO + H ), because CO was not formed from CO/
O mixtures (5 kPa CO; 2 kPa H O) on this Au/Al catalyst, even
at higher temperatures (ꢃ350 K).
Positive H O effects have been previously attributed to the
2
O concentration on steady-state CO oxidation
2
2
indicate that O–H(D) bonds are not cleaved in kinetically-relevant
steps in the CO oxidation catalytic sequence.
2
2
2
Previous studies have proposed that H
>0.5 kPa) reflect competitive adsorption by H
diates [10,36]. It is also possible that H O condensation within
catalyst mesopores imposes diffusional corruptions of chemical
reaction rates by decreasing intrapellet CO or O concentrations
O inhibition effects
2
(
2
O-derived interme-
2
2
[
2
2
and thus the kinetic driving force for CO oxidation at active sites.
We find, however, that CO oxidation turnover rates did not depend
on pellet size (0.125–0.250 and 0.250–0.425 mm diameter; Fig. 2);
thus, even if condensation occurred, the presence of H O within
2
alumina pores did not impose transport barriers, which would
have caused smaller rates on the larger pellets. Moreover, the pore
H
2
(
2
2
2
2
H
2
2
2 3
O
2
decomposition of carbonates formed on Au active sites during CO
size distribution of the Al
detectable pore filling below ꢃ1.5 kPa H
tures (Supplementary Content). We conclude that the
inhibition observed above ꢃ0.5 kPa (Fig. 2) reflects competitive
adsorption of H O-derived adsorbed species with reactive interme-
diates derived from CO or O on Au clusters and arise from chem-
ical origins instead of any transport corruptions.
2
Thermal treatments of Au/Al in 25 vol.% O /He at tempera-
2
O
3
used in this study did not lead to
2
O at reaction tempera-
2
H O
H O addition
2
10
2
0.5 kPa
0 kPa
0.5 kPa
2
-1
kd = 0.76 ks
2
O
1
3
tures up to 1023 K did not lead to detectable changes in Au cluster
size (transmission electron microscopy, TEM; Supplementary Con-
tent) or in CO oxidation rates in the presence of H
CO, 2 kPa O , 0.5 kPa H O; Table 1), indicating that Au clusters on
Al resist sintering even at these high temperatures and that
the deactivation of Au-based catalysts reported during CO oxida-
tion in the absence of H O at near-ambient temperatures cannot
reflect cluster growth. These data also show that CO oxidation
occurs on the Au clusters visible in transmission electron micro-
graphs. In contrast, the rates of HCOOH dehydrogenation (353 K,
2
O (288 K, 5 kPa -
0
.1
2
2
-1
2 3
O
kd = 0.21 ks
2
0.01
0
5
10
15
20
25
30
35
40
Time-on-stream (ks)
2
2
kPa HCOOH) and water–gas shift (523 K, 5 kPa CO, 2 kPa H O)
Fig. 1. H
2
O effect in CO oxidation turnover rates measured with the untreated Au/
). Symbols: (d) 0.5 kPa H O; ( ) 0 kPa
decreased markedly with increasing treatment temperature,
apparently as a result of the disappearance of smaller Au structures
Al (0.61% wt., 288 K, 5 kPa CO, 2 kPa O
2
O
3
2
2
2
H O.

M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
95
Table 1
2 3 2
Influence of Au/Al O treatment temperature (O /He, 2 h) on Au clusters size obtained by TEM and on rates for CO oxidation, HCOOH dehydrogenation, and water–gas shift
reactions.
ꢂ1
ꢂ1
Catalyst
Treatment temp. (K)
Mean Au size (nm)
Rate (mol h g-at Au )
CO–O
2
–H
2
Oa
HCOOHb
CO–H
2
Oc
Au/Al
Au/Al
Au/Al
Au/Al
2
O
2
O
2
O
2
O
3
3
3
3
–
3.5 ± 1.2
4.0 ± 1.0
4.0 ± 1.1
4.3 ± 1.2
2589
2827
2535
2610
477
201
85
110
74
57
-873
-950
-1023
873
950
1023
40
30
a
b
c
2
3
5
88 K, 5 kPa CO, 2 kPa O
53 K, 2 kPa HCOOH.
23 K, 5 kPa CO, 2 kPa O
2
, 0.5 kPa H
.
2
O.
2
undetectable at the resolution of these transmission electron
micrographs (Table 1) [29].
CO oxidation turnover rates (5 kPa CO, 2 kPa O
2 2
, 0.5 kPa H O) on
Au/Al catalysts increased with temperature (282–303 K) in a
2 3
O
manner consistent with an apparent activation energy of 36 ±
ꢂ1
3
kJ mol (Fig. 4); these values lie within the range reported by
3
.2. Kinetic dependence of CO oxidation rates on reactants pressure
ꢂ1
other authors (22–40 kJ mol ) on Au/Al
O
2 3
(1.2–5.6 nm Au clus-
(0.25–75 kPa) at
03–453 K [14,34,39]. However, in the absence of H O, other
and temperature
2 2
ters) in the presence of H O (40 ppm) or H
3
2
The stable CO oxidation turnover rates observed in the presence
authors found lower CO oxidation activation energies (12 ±
of H
over a wide range of CO, O
the kinetic effects of CO and O
rates on Au/Al (0.5 kPa H O; Fig. 3). CO oxidation turnover rates
increased with increasing CO (0.80–8.25 kPa) and O (0.25–
.15 kPa) pressures, with empirical fractional orders of 0.24 ± 0.02
and 0.60 ± 0.02, respectively. Au/Al treated in O /He at various
higher temperatures (873–1023 K) gave similar kinetic effects of
CO (0.15–0.25 order) and O (0.55–0.60 order) (288 K, 1–6 kPa CO,
.5–6 kPa O , 0.5 kPa H O) and they also requires co-fed H O for
2
O allowed accurate and reproducible kinetic measurements
, and H O pressures. Next, we report
pressures on CO oxidation turnover
ꢂ1
1
kJ mol , 295–373 K) with Au/Al
2
O
3
(1.2 nm Au) [34]. Moreover,
2
2
erratic and large differences in activation energies, ranging from 8
to 75 kJ mol , have been reported in the literature for nominally
anhydrous CO/O
and non-reducible supports (TiO
2
ꢂ1
O
2 3
2
2
reactions on Au clusters deposited on reducible
, Fe , CeO , SiO , etc.).
2
2
2
O
3
2
2
7
2
O
3
2
2 2 3 2
3.3. CO–O reaction pathways on Au/Al O in the presence of H O
2
0
2
2
2
Scheme 1 depicts a sequence of elementary steps for low-tem-
perature CO oxidation reaction consistent with the kinetic data dis-
cussed in Section 3.2. Steps 1, 2, and 3 represent quasi-equilibrated
non-dissociative adsorption steps for all species in the reactant
stream (CO, O , and H O). O adsorption is assumed to be molecu-
high and stable CO oxidation turnover rates (Supplementary Con-
tent). These fractional reaction orders, taken at face value, cannot
be reconciled with any sequence of elementary steps. Yet, it is in
such a manner that kinetic data on Au catalysts have been invariably
reported; we do so here, only momentarily, to show that these frac-
tional orders lie within the broad range reported in previous studies
2
2
2
lar and reversible. Irreversible O2 adsorption (whether molecular
or dissociative) would lead to CO oxidation rates proportional to
using nominally anhydrous CO/O
CO and O , respectively [1,7,9,24,37,38]). The positive orders in both
CO and O indicate that either both CO and O are present below
2
mixtures (0.1–0.5 and 0.3–0.6 for
O pressures, in contradiction with our kinetic data and with most
2
2
previous reports [1]. Reversible O adsorption and subsequent
2
ꢀ
2
2
reactions with the CO species present as most abundant reaction
saturation coverages on Au surfaces during CO oxidation or that
they adsorb non-competitively on distinct active sites.
intermediates (MARI) would also lead to first-order O2 kinetics,
2
1
1
1
0
0
O pressure (kPa)
2
Au/Al O3
2
-1
E = 36 ± 3 kJ mol
0
1
2
3
4
5
6
7
8
9
1
0
1
2
3
4
a
6
5
4
3
2
1
0
3
x10
0
-
10
4
-
-
-
x10
10
10
10
Pt/Al O3
2
-1
E = 66 ± 8 kJ mol
a
2.2
2.4
2.6
3.2
3.4
3.6
3
(
10 K/T)
0
1
2
3
4
5
6
7
8
9
Fig. 4. Arrhenius plots for the CO oxidation (5 kPa CO; 2 kPa O
measured with the untreated Au/Al (d, 0.61% wt., 3.5 nm Au, 282–303 K) and Pt/
Al , 2.03% wt., 1.3 nm Pt, 373–423 K). Rates with Pt/Al were determined at
1 kPa CO and 10 kPa O
established CO and O
orders for O
2 2
; 0.5 kPa H O)
CO pressure (kPa)
2 3
O
2
O
3
(
2 3
O
Fig. 3. Kinetic dependence of CO oxidation turnover rates on CO pressure (0.80–
.25 kPa CO, 2 kPa O
.5 kPa H O) measured with the untreated Au/Al
2
and recalculated according to previously reported and well-
pressures effects on measured kinetics (+1 and ꢂ1 reaction
8
0
2
, 0.5 kPa H
2
O) and O
2
pressure (0.25–7.15 kPa O
(0.61% wt., 288 K).
2
, 5 kPa CO,
2
2
2
O
3
2
and CO pressure, respectively) [31].

9
6
M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
the terms in the denominator of Eq. (1) containing these constants
ꢀ
ꢀ
ꢀ
(
relative surface coverage of O , OH, and OOH) to be much smaller
ꢀ
ꢀ
ꢀ
2
than the others (vacant sites, CO , O , and H O ; Supplementary
2
Content) at the conditions of our experiments. Therefore, Eq. (1)
becomes
2
=3
a
ðPCO
P
O2
P
P
H2
O
Þ
r ¼
ð6Þ
2
½
1 þ K
1
P
CO þ K
2
O2 þ K
3
P
H2
O
ꢆ
ꢀ
ꢀ
ꢀ
when assuming that OOH, OH, and O are minority surface species
⁄
ꢀ
present in much lower concentrations than vacant sites (ꢀ), CO , O ,
2
ꢀ
and H
2
O .
The parity plot in Fig. 5 shows that Eq. (6) accurately describes
the measured effects of reactants pressures (0.80–8.25 kPa CO;
Scheme 1. CO oxidation pathways on Au/Al
2
O
3
at low temperatures in the
ꢀ
ꢀ
presence of co-fed H
species).
2
O ( denotes a vacant active site and X the various adsorbed
0
.25–7.15 kPa O
rates on Au/Al
abundant surface species or about quasi-equilibrated steps consis-
tent with these measured rates. Alternate CO/O /H O reaction
2
; 0.03–1.15 kPa H
O . We did not find other assumptions about most
2 3
2
O) on CO oxidation turnover
suggesting that alternate steps must be involved in the activation
ꢀ
2
2
of O species. Step 4 represents the formation of hydroperoxy spe-
2
ꢀ
ꢀ
pathways were also considered (Supplementary Content), but rig-
orous kinetic and isotopic (shown later) analyses demonstrated
that none of them was able to successfully predict measured rate
data. The kinetic parameters derived from the nonlinear regression
cies ( OOH) via proton transfer from water to O , in a step assumed
2
to be fast and quasi-equilibrated based on the small measured
2 2
H O/D O kinetic isotope effects (Section 3.4). In this sequence,
O–O bond activation is assisted by H
2
O, in a step that avoids the
ꢀ
analysis of measured CO oxidation rates with Au/Al
2
O
3
are shown
difficult dissociation of O
species that then react with CO to form CO
2
on Au surfaces; this step forms OOH
ꢀ
in Table 2. The H
2
O adsorption equilibrium constant (3.1 ±
2
(step 5). Catalytic
ꢂ1
ꢀ
0.6 kPa ) is larger than for CO and O
2
(0.18 ± 0.05 and 0.06 ±
turnovers are completed via the recombination of OH species
ꢂ1
ꢀ
0.02 kPa , respectively); CO is indeed expected to bind more
(
step 6) to re-form the H
2
O molecule used to activate O2 (which
strongly than O
2
, based on first-principle density functional theory
therefore acts as a co-catalyst instead of a stoichiometric reactant)
ꢀ
(DFT) with generalized gradient approximation (GGA) calculations
and an O atom; the latter is ultimately removed via the reaction
ꢀ
on Au/Al
2
O
3
(100) and Au
8
–Au30 model systems, which gave CO
with CO to form another CO
2
molecule (step 7), a step that exhib-
ꢂ1
ꢂ1
adsorption energies (90–100 kJ mol ) much greater than for O
2
its a low activation barrier (ꢃ24 kJ mol ) according to previous
ꢂ1
(
5–70 kJ mol ) [41,45–47]. The measured H
2
O adsorption con-
theoretical estimates on Au(211) and Au(221) model surfaces
ꢀ
ꢀ
ꢀ
ꢀ
ꢀ
stant (K
3
) is, however, somewhat larger than the expected from
[
40]. OH recombination to HOOH ( OH + OH ? HOOH + ꢀ) is
not considered as a plausible alternate path to the proposed
H O + O reaction (E = 17–25 kJ mol [41,42]) based on DFT esti-
2 a
ꢀ
ꢀ
ꢂ1
6
5
4
3
2
1
0
mates on Au(111) showing a much higher activation energy bar-
rier (157 kJ mol ) [43]).
The assumption of pseudo-steady-state [44] for all adsorbed
species and of quasi-equilibrium for steps 1–4 in Scheme 1 leads
to the following CO oxidation rate equation, the full algebraic der-
ivation of which is included in Appendix A:
ꢂ1
0
1
2
3
4
5
6
ð1Þ
-1
-1
Measured Au-time yield (mol s g-at Au )
s
Each term in the denominator of this equation represents the con-
centration of the indicated adsorbed intermediate relative to the
Fig. 5. Parity plot for measured (0.61% wt. Au/Al
2 3
O ; 288 K; 0.80–8.25 kPa CO;
0
.25–7.15 kPa O ; 0.03–1.15 kPa H O) and calculated (from Eq. (6)) CO oxidation
2
2
concentration of vacant sites on Au cluster surfaces. In Eq. (1), K
, and K are the equilibrium constants for molecular CO, O , and
O adsorption, respectively, and , b, , and d are given by the
1
,
turnover rates.
K
2
3
2
H
2
a
c
combinations of rate and equilibrium constants for elementary
steps
Table 2
Kinetic parameters for the CO oxidation model in the presence of H
obtained from nonlinear fitting to Eq. (6) of kinetic data measured with Au/Al
TiO , and Au/Fe catalysts.
2
O (Scheme 1)
O
3
, Au/
2
=3
1=3
2
a
¼ 3=2ðK
1
K
2
K
3
K
4
K
5
Þ
Þ
ð2K
6
Þ
5
ð2Þ
ð3Þ
ð4Þ
ð5Þ
2
2 3
O
2
=3
1=3
ꢂ1=3
b ¼ ðK
c
2
K
3
K
4
2
Þ
K
ð2K
6
ðK
1
K
Þ
Kinetic parameter
2
Au/Al O
3
Au/TiO
2
2
Au/Fe O
3
ꢂ1=3
2=3
1=3
ð2Þꢂ2=3ðKꢂ7 1
ꢂ1
ꢂ2
¼ K1 ðK
d ¼ ðK
The values of the kinetic parameters b,
nonlinear regression analysis of measured CO oxidation rates make
3
K
4
K
5
Þ
ðK
6
Þ
Þ
a
K
K
K
(s kPa
)
11.0 ± 3.4
0.18 ± 0.05
0.06 ± 0.02
3.1 ± 0.6
12.3 ± 3.4
0.24 ± 0.09
0.05 ± 0.02
3.4 ± 0.9
1.3 ± 0.4
0.18 ± 0.07
0.16 ± 0.06
1.1 ± 0.5
ꢂ1
(kPa
(kPa
(kPa
)
)
)
1=3
1=3
1
2
3
ꢂ1
ꢂ1
2
K
3
K
4
K
5
Þ
ð2K
6
Þ
c
, and d derived from the
2=3
1=3.
6
a
¼ 3=2ðK
1
K
2
K
3
K
4
K
5
Þ
ð2K
Þ

M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
97
equimolecular 16O
of CO and H
mial isotopologue distributions (50%
adsorption steps would not form mixed O18O isotopes. The rate
/18
O
2
mixtures (5 kPa) at 300 K in the absence
4
3
2
1
0
2
2
O. Quasi-equilibrated dissociation would form bino-
1
6 18
O
O), while molecular
1
6
16 18
16
18
KIE = 1.21
of
smaller than CO oxidation rates (288 K, 5 kPa CO, 2 kPa O
O) even at higher temperatures and O pressures than those
used in CO oxidation catalysis (Fig. 7a); these findings are consis-
O
O formation (300 K, 5 kPa
O
2
, 5 kPa
O
2
) was ꢃ50 times
2
, 0.5 kPa
0
.5 kPa H O
0.5 kPa H O
2
H
2
2
2
0
.5 kPa D O
16 18
2
tent with previous studies that reported no detectable
O
O for-
1
6
18
mation during
Au/TiO (353 K) and Au/Fe
sorbs molecularly ðO
O / O isotopic exchange at similar pressures on
2 2
2
2
O
3
(298 K) [13,49] and show that O
2
ad-
18O
ꢀ
2
þ ꢀ $ O Þ on Au clusters.
2
1
6
The rate of isotopic exchange between
2 2
O (5 kPa) and H
(
2 3
0.5 kPa) was also measured on Au/Al O at 300 K in the absence
1
6
18
18
16
of CO. The formation of
O
O,
2 2
O , or H O molecules was not
0
2
4
6
8
10 12 14 16
detected (Fig. 7b). The absence of isotopic exchange between O
2
Time-on-stream (ks)
and H
2
O reflects the non-dissociative nature of O
2
adsorption on
exchange. We
1
6
18
Au, as also inferred from the absence of
O
2
/
O
2
Fig. 6. H
2
O/D
2
O isotope effects on CO oxidation turnover rates measured with Au/
; 0.5 kPa H O (j) or D O ( )).
note that steps 2–4 in the catalytic sequence do not lead to the
cleavage of O–O bonds, which instead occurs in latter steps via
reactions with CO⁄ in CO oxidation and with propene in propene
epoxidation on Au catalysts [28]; as a result, the equilibrated nat-
ure of these O
oxygen exchange.
Al (0.61% wt., 288 K; 5 kPa CO; 2 kPa O
2
O
3
2
2
2
the Au–H
cal treatments on model Au
TiO
2
O interaction energies reported from previous theoreti-
2
adsorption and reaction steps cannot cause isotopic
8
–Au30, Au/Al
2
O
3
(100), and Au/
ꢂ1
2
(110) surfaces (30–100 kJ mol ) [42,48].
3.5. Support effects on CO oxidation rates with Au-based catalysts
2 2
3.4. H O/D O isotope effects and isotopic exchange measurements
The identity of the support has been often implicated as an
H
2
O/D
2
O isotope effects were used to probe the potential
O and of O–H bond activation in kinetically-rel-
essential second function in the reactivity of Au clusters in CO oxi-
dation [1,5,6,24,50], either via its participation in O activation or
because of its ability to stabilize small Au clusters. Reducible oxi-
des (e.g., TiO , Fe ) have been reported to catalyze CO oxidation
at higher rates than refractory oxides (e.g., Al , SiO ), a finding
attributed to the requirement for O activation on periphery sites
located at Au-support interfaces [6,51]. The data used to reach
these conclusions were obtained under nominally anhydrous con-
ditions, but possibly in the presence of catalytically-consequential
involvement of H
evant steps in the context of the CO oxidation catalytic sequence
depicted in Scheme 1. CO oxidation turnover rates (288 K, 5 kPa -
CO, 2 kPa O
O or D O (0.5 kPa) to CO/O
measured H O/D O isotope effects (rH
consistent with previous studies that report H
fects near unity on Au/Al (295 K, 1 kPa CO, 0.5 kPa O
O(D O)) [16]. These data lead us to conclude that O–H (or O–
2
2
2
2 3
O
2
) were measured on untreated Au/Al
reactant streams (Fig. 6). The small
O/rD O = 1.21; Fig. 6) are
O/D O isotope ef-
, 1.5 kPa
2
O
3
by adding
2
O
3
2
H
2
2
2
2
2
2
2
2
2
2
2
O
3
2
H
2
2
levels of H
2
O, derived, at least in part, from H
2
O adsorbed on sup-
O is de-
D) bonds are not cleaved in kinetically-relevant steps, which would
have shown large normal isotope effects at these low temperatures
ports and at concentrations that decrease with time as H
2
pleted by desorption from these supports. These processes may
have led, in turn, to the decrease in rates with time often inter-
preted as catalyst deactivation during CO oxidation on Au-based
catalysts.
Fig. 8 and Table 3 compare turnover rates on Au clusters of sim-
2 2 3
ilar size dispersed on reducible (TiO and Fe O , 3.3 ± 0.7 and
(
288 K). These small normal H
probable thermodynamic origins stemming from quasi-equili-
brated H O dissociation steps that occur before kinetically-relevant
steps in CO oxidation catalytic sequences.
Next, we probe chemisorption steps on Au/Al
measuring the rate of formation of 6O18O isotopologues from
2 2
O/D O isotope effects reflect their
2
O
2
2
O
3
by
1
2 3
3.6 ± 0.7 nm Au clusters, respectively) and non-reducible (Al O ,
1
0
0
0
0
0
.0
.8
.6
.4
.2
.0
(a)
(b)
1
0
0
0
0
.0
.8
.6
.4
.2
16
O2
18
O2
16
O2
1
6
18
16 18
18
O2
O O
O O
0.0
0
1
2
3
4
0.0
0.5
1.0
1.5
Contact time (ks)
Contact time (ks)
Fig. 7. Molar fraction of 16O
0.61% wt.; untreated).
(j), 16O18O ( ), and 18 ) during contact at 300 K of equimolar 16 /18 (a, 5 kPa) or 16 18O (b, 5 kPa O
O ( O O O /H , 0.5 kPa H O) on Au/Al O
2 2 2 2 2 2 2 2 3
2
(

9
8
M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
10
10
H O addition
H O addition
2
(a)
2
(b)
0.5 kPa
0 kPa
0.5 kPa
1
0.5 kPa
0 kPa
0.5 kPa
1
0
.1
0.1
0.01
0
2
4
6
8
10
12
14
0
2
4
6
8
10
12
14
Time-on-stream (ks)
Time-on-stream (ks)
2 2 2 2 3 2 2
Fig. 8. H O effect on CO oxidation turnover rates (288 K, 5 kPa CO, 2 kPa O ) measured with Au/TiO (a) and Au/Fe O (b). Symbols: (d) 0.5 kPa H O; ( ) 0 kPa H O.
Table 3
Effect of H
values appear to reflect the encapsulation of some TEM-visible Au
clusters within Fe as a consequence of the co-precipitation
methods used, which in contrast with the deposition–precipitation
method used for TiO and Al supports do not ensure the pres-
2
O addition (0.5 kPa) on CO oxidation turnover rates
) for Au clusters deposited onto Al
(1.56% wt. Au), and Fe (4.44% wt. Au).
(
(
288 K, 5 kPa CO, 2 kPa O
0.61% wt. Au), TiO
2
2
O
3
2 3
O
2
2 3
O
2
2 3
O
CO oxidation turnover rate
ðmol sꢂ1 g-at Ausꢂ1
ence of Au clusters at accessible support surfaces.
We conclude from these data that supports influence CO oxida-
tion turnover rates with anhydrous reactants, perhaps because
Þ
Added H
Au/Al
Au/TiO
Au/Fe
2
O (kPa)
0
0.5
2
O
3
0.08
0.54
0.07
2.55
2.70
0.59
they can contribute adsorbed H
sites that become essential for O
2
O as co-catalyst or as periphery
activation only under anhydrous
2
2
2
O
3
conditions. These support effects and specifically the reducible nat-
ure of the support become much less consequential when H
mediates the required O activation steps. The observed H O ef-
fects on CO oxidation rates with Au/TiO and Au/Fe also suggest
that the mechanistic pathways in Scheme 1 are also involved in CO
oxidation reactions on reducible supports, as long as H O is present
as a co-catalyst. These elementary steps occur on Au cluster sur-
faces without any significant requirement for a second function
provided by either interfacial or support sites. This is also consis-
2
O
3
.5 ± 1.2 nm Au clusters) oxides for CO/O
tants. As in the case of Au/Al catalysts, the addition of H
a co-catalyst was essential for stable turnover rates on Au/TiO
and Au/Fe . H O (0.5 kPa) increased CO oxidation rates on all
catalysts, but to different extents, possibly because of the presence
of support-dependent H O concentrations even for nominally
anhydrous CO–O reactants. Turnover rates with anhydrous CO–
reactants (based on TEM mean cluster sizes) on Au/TiO were
significantly higher than on Au/Al
(ꢃ7 times), in agreement
with literature claims for the benefits of reducible oxide supports.
CO/O /H O mixtures, however, gave similar rates on Au/TiO and
Au/Al (2.70 and 2:55 mol s g-at Aus , respectively) and some-
what smaller rates on Au/Fe
2 2 2
and CO/O /H
O reac-
2
2
O
3
2
O as
2
2 3
O
2
2
2
O
3
2
2
2
2
tent with previous reports [14] showing similar CO oxidation acti-
O
2
2
ꢂ1
vation energies on Au/TiO
2
and Au/Al
2
O
3
(22 and 25 kJ mol
,
2 3
O
respectively), which suggests a common reaction mechanism over
these two catalysts.
2
2
2
ꢂ1
ꢂ1
The effects of CO and O
rates on Au/TiO and Au/Fe
2
pressure on CO oxidation turnover
(Fig. 9; 288 K, 0.5 kPa H O) were
2 3
O
ꢂ1
ꢂ1
2
2
O
3
2
2
O
3
ð0:59 mol s g-at Aus Þ. The latter
O pressure (kPa)
O pressure (kPa)
2
2
(a)6
0
1
2
3
4
5
6
7
(b)1.2
0
1
2
3
4
5
6
7
5
4
3
2
1
0
1.0
0.8
0.6
0.4
0.2
0.0
0
1
2
3
4
5
6
7
0
1
2
3
4
5
6
7
CO pressure (kPa)
CO pressure (kPa)
Fig. 9. Kinetic dependence of CO oxidation turnover rates on CO (d, 1–5 kPa CO, 2 kPa O
88 K with Au/TiO (a) and Au/Fe (b).
2 2 2 2 2
, 0.5 kPa H O) and O ( , 0.5–6 kPa O , 5 kPa CO, 0.5 kPa H O) pressures measured at
2
2
2 3
O

M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
99
Table 4
Effect of H
epoxidation rates and selectivity (350 K, 4 kPa
1.56% wt. Au, WGC). Table adapted from reference [28].
2
O addition on initial (extrapolated to zero time-on-stream) propene
5
C
H
3 6
,
4 kPa with Au/TiO
O
2
)
2
(
4
ꢂ1
H
2
O pressure
Metal-time yield (mol h g-
Selectivity (%, carbon
based)
ꢂ1
2
2
1
1
0
0
.5
(kPa)
at Au
)
Propene
oxide
Acetone
.0
.5
.0
.5
.0
1
2
6
0.35
0.51
0.52
0.57
85.2
75.0
39.7
24.1
14.8
25.0
60.3
75.9
12
near-ambient temperatures on Au/TiO
finding later confirmed by propylene oxide synthesis from CO/
/propene mixtures in aqueous Au/TiO and TS-1 suspensions
2
(Table 4, 350 K) [28], a
0
.0 0.5 1.0 1.5 2.0 2.5
4
5
-1
-1
O
2
2
Measured Au-time yield (mol s g-at Au )
s
[
59]. These epoxidation reactions provide independent evidence
ꢀ
Fig. 10. Parity plot for measured (288 K; 1–6 kPa CO; 0.5–6 kPa O
.53 kPa H O) and calculated (from Eq. (6)) CO oxidation turnover rates with Au/
TiO ) and Au/Fe ).
2
;
0.12–
for the formation of OOH species from O
2
–H
activation during CO oxidation reactions.
activation pathways occur on Au clusters surfaces with-
out detectable involvement of the support and remove the need
for support-mediated O activation at the periphery of Au clusters
as a requirement for CO oxidation catalysis. Theoretical calcula-
tions on Au and Au30 clusters indicate that coadsorbed H O and
form O O complexes via partial proton sharing or transfer
2
O reactants and for
0
2
their involvement in O
These O
2
2
(
2 3
O (
2
2
2 3
similar to those measured on Au/Al O (Fig. 3). Turnover rates
increased with CO pressure (1–6 kPa, 0.29 ± 0.04, and 0.19 ± 0.02
8
2
orders on Au/TiO
0.5–6 kPa, 0.68 ± 0.01, and 0.49 ± 0.02 orders on Au/TiO
Fe
therefore appears to be generally applicable to CO oxidation reac-
tions with CO–O –H O reactants on Au clusters dispersed on all
three supports. The rate equation derived from the steps in Scheme
(Eq. (6)) accurately described all rate data as a function of CO, O
and H O pressures on Au/TiO and Au/Fe . Fig. 10 shows the
2
and Au/Fe
2
O
3
, respectively) and O
2
pressure
and Au/
O
2
2
ꢁH
2
(
2
ꢂ1
ꢀ
with low activation barriers (17–25 kJ mol ) to form OOH species
[
2 3
O , respectively). The sequence of elementary steps in Scheme
ꢀ
41,42]. OOH species may be responsible for the well-established
1
promotional effect of H
1,34,39,60,61], as proposed previously for model Ru–Pt core shell
nanoparticles from first-principles calculations consistent with
2 2
in CO/O reaction on Au catalysts
2
2
[
1
2
,
ꢀ
ꢀ
ꢀ
atomic H addition to O to form OOH as the relevant pathway
2
2
2
2 3
O
2
in the preferential oxidation of CO in H -rich streams [62].
respective parity plots of measured and predicted turnover rates,
and Table 1 lists the regressed kinetic parameters. As found on
Au/Al
for CO and O
, K , and K are similar on all supports, consistent with the similar
2
O
3
, H
2
O adsorption equilibrium constants are larger than
2
on Au/TiO and Au/Fe (Table 2). The values of
2
2
O
3
3.7. H O effects on CO oxidation turnovers catalyzed
by supported Pt clusters
2
K
1
2
3
size of Au clusters and their exclusive involvement in mediating
the elementary steps in Scheme 1.
A catalytic effect of H O was also observed during CO oxidation
2
on Pt clusters, the origins of which have not, to our knowledge,
3
(
.6. Evidence for the presence and reactivity of hydroperoxy
OOH) species on Au
been examined in the extensive previous literature dealing with
Pt-catalyzed CO oxidation catalysis [63,64]. Fig. 11a shows CO oxi-
⁄
2 3
dation turnover rates on Pt/Al O at 423 K (2.03% wt., 1.3 nm Pt
The elementary steps for CO–O
O (Scheme 1) and the assumptions made about their reversibil-
2
reactions in the presence of
clusters) with and without added H
2
O. CO oxidation rates de-
ꢂ1
ꢂ1
H
2
creased from 0.28 to 0:08 mol s g-at Pts in ꢃ1 ks when H
2
O
ity lead to a rate equation (Eq. (6)) consistent with measured rates
on Au clusters dispersed on both reducible (TiO , Fe ) and non-
reducible (Al ) supports. These steps and assumptions are also
consistent with H O/D O isotope effects on CO oxidation rates
and with the absence of isotopic exchange in
(0.5 kPa) was removed from the reactant mixture (1 kPa CO,
10 kPa O ). Initial CO oxidation rates were fully recovered upon
reintroduction of H O (0.5 kPa, Fig. 11a). The parallel formation
of CO via water–gas shift occurs at negligible rates
2
O
2 3
2
2
O
3
2
2
2
2
ꢂ3
16
18
18
ꢂ1
ꢂ1
O
2
/
O
2
and
O
O
2
/
ð< 10 mol s g-at Pts , Pt clusters of 0.9–1.7 nm) at these tem-
peratures (423 K, 3 kPa CO, 10 kPa H O) [63,65]. The data in
Fig. 11a suggest that H O also provides a more facile route for O
1
6
ꢀ
ꢀ
H
2
O
2
mixtures. O
2
activation involves reactions of O and H
2
2
2
ꢀ
to form hydroperoxy species ( OOH) with intact O–O bonds, but
2
2
ꢀ
weaker relative to those in O
2
molecules. These OOH intermedi-
activation during CO oxidation on CO-covered Pt clusters. A recent
ates also account for the reactivity of small Au clusters in the for-
study combining theory and experiment [31] showed that anhy-
ꢀ
mation of H
2
O
2
from H
2
–O
2
mixtures [52,53], which proceeds via
drous CO oxidation on Pt occurs via CO -assisted O
steps on surfaces nearly saturated by chemisorbed CO intermedi-
2
dissociation
ꢀ
ꢀ
OOH intermediates, as shown from theoretical estimates [27],
electronic and spin resonance spectra [54], and inelastic neutron
scattering [55]. OOH species also account for propylene epoxida-
ates, which hinder O
2
adsorption and subsequent activation. The
ꢀ
ꢀ
2
presence of H O may act to decrease local CO coverages and weak-
ꢀ
tion with H
2
/O
2
mixtures on Au-based catalysts [56,57]. These
en inhibition effects by CO [64], but may also provide OOH species
ꢀ
OOH species have been also proposed to account for the promo-
, without specific direct evidence
has been detected by UV–visible
as alternate O
concomitantly higher CO oxidation turnover rates than the CO-as-
sisted O activation steps that limit CO oxidation with anhydrous
CO/O
2
activation routes with lower activation barriers and
tion of CO oxidation by H
35,39,58]. The formation of H
2
[
2
O
2
2
spectroscopy during CO oxidation on Au/TiO
2
and Au/C in aqueous
2
mixtures. Indeed, apparent activation energies decrease
ꢂ1
alkaline media [18].
from 84 ± 6 kJ mol
(403–473 K, 1 kPa CO, 10 kPa O
2
) [31] to
O,
ꢀ
ꢂ1
The formation of OOH species from O
2
2
/H O mixtures is also
66 ± 8 kJ mol
(373–423 K, 1 kPa CO, 10 kPa O
2
,
0.5 kPa H
2
evident from the formation of propylene oxide from propene at
Fig. 11b) when H
2
O was present in CO/O reactants.
2

1
00
M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
a) 1 H O addition
(b)
10
(
2
0.5 kPa
0 kPa
0.5 kPa
-1
1
E = 66 ± 8 kJ mol
a
0.1
0.1
0.01
0.01
1E-3
0
1
2
3
4
5
6
7
8
2.2
2.3
2.4
2.5
3
(10 K/T)
2.6
2.7
2.8
Time-on-stream (ks)
Fig. 11. H
Symbols: (d) 0.5 kPa H
2
O effect on CO oxidation turnover rates (a, 423 K) and Arrhenius plot (b, 373–423 K) measured with Pt/Al
O; ( ) 0 kPa H O.
2 3 2
O (2.03% wt., 1.3 nm Pt clusters, 1 kPa CO, 10 kPa O ).
2
2
Fig. 4 shows CO oxidation turnover rates (5 kPa CO, 2 kPa O
2
,
responsible for the higher measured CO oxidation rates compared
to those with anhydrous CO/O mixtures.
and 0.5 kPa H O) on Au/Al (3.5 nm Au) and Pt/Al (1.2 nm Pt)
2
2
O
3
O
2 3
2
as a function of temperature. CO oxidation turnover rates were
much higher on Au clusters than on Pt clusters. CO oxidation acti-
4
. Conclusions
Moisture levels in ‘‘anhydrous’’ CO/O2 streams well below the
vation energies were also much smaller on Au/Al
2
O
3
than Pt/Al
2
O
3
ꢂ1
(
36 ± 3 versus 66 ± 8 kJ mol ). The lower CO oxidation rates on Pt
ꢀ
clusters reflect CO -saturated surfaces, a phenomenon that de-
detection limits of typical speciation techniques are consequential
for CO oxidation catalysis; this has led to our suggestion that the
creases O
2
adsorption rates on Pt [31]. In contrast, Au surfaces
ꢀ
are not saturated with CO intermediates during CO oxidation
catalysis (as suggested by the small CO adsorption equilibrium,
presence of H
for their reproducible measurements, and for fundamental studies
of the mechanism and site requirements of these reactions. H O is
an efficient co-catalyst necessary for O activation steps and for
catalyst stability in CO/O reactions at near-ambient temperatures
with Au clusters (<5 nm) deposited on reducible (TiO , Fe ) and
non-reducible (Al ) supports. Rigorous kinetic and isotopic data
measured on stable Au catalysts have led us to propose a mecha-
nistic picture of CO oxidation in the presence of H O that involves
activation steps that form hydroperoxy species
2
O is essential both for practical CO oxidation rates,
1 2 2
K ). Consequently, the adsorption of O and H O and the subse-
quent formation of the oxidant species (hydroperoxy intermedi-
ates) is more facile on Au-based catalysts.
2
2
2
2
2 3
O
3.8. Role of H
2
O in CO oxidation pathways catalyzed by supported
2 3
O
noble metal (Au,Pt) clusters
2
2
The experimental evidence reported here for H O effects on CO
H
2 2
O-mediated O
ꢀ
oxidation turnover rates resolves long-standing conflicts among
measured rates on Au catalysts with similar cluster size, Au con-
( OOH), which precisely account for the remarkable co-catalytic ef-
2
fect of H O on measured CO oxidation rates. These steps occur
tents, supports, and synthetic provenance. The presence of H
as a co-catalyst in CO/O reactants, even at trace concentrations, af-
fects significantly catalyst stability and measured CO oxidation
rates. H O addition is essential for stable turnover rates, and its
2
O
exclusively on the Au clusters; hence, the influence of support
identity on CO oxidation catalysis, and specifically its reducible
2
nature, becomes much less important in the presence of H
These H O effects in CO/O reactions are also found with Pt cata-
lysts, which suggest that common elementary steps during CO/
/H O reaction occurs on supported noble metal clusters (Au,Pt).
2
O.
2
2
2
presence leads to a sequence of elementary steps that are mediated
ꢀ
by hydroperoxy species ( OOH), which account for the co-catalytic
O
2
2
effect of H
as for the ability of Au clusters to catalyze propene epoxidation
reactions with O /H O mixtures [28,59]. The elementary steps that
form the hydroperoxy species from O and H O occur exclusively
2
O on CO oxidation rates and on catalyst stability, as well
The formation of hydroperoxy species explains the significantly
higher CO oxidation rates compared to those measured with anhy-
drous CO/O mixtures.
2
2
2
2
2
on Au cluster surfaces without any significant involvement of sup-
port sites; hence, the reducible nature of the support become much
less consequential for CO oxidation catalysis in the presence of the
H O molecules that mediate required O activation steps. In con-
2 2
trast, supports influence CO oxidation turnover rates with nomi-
nally anhydrous reactants, perhaps because they contribute
Acknowledgments
This work was supported by the Director, Office of Basic Energy
Sciences, Chemical Sciences Division of the US Department of En-
ergy under Contract DE-AC02-05CH11231. M. Ojeda acknowledges
the financial support from the European Union (Marie Curie Ac-
tions, MOIF-CT2005-007651) and Spanish Ministry of Science and
Innovation (RYC-2010-06067). H. Kung and M. Kung (Northwest-
either adsorbed H
periphery sites that become essential for O
are required only in the absence of H O.
The observed H O effects in CO oxidation are not exclusive of
reactions catalyzed by Au clusters. We find that H O also influ-
ences significantly measured CO oxidation rates with Pt clusters.
This strongly suggests that CO oxidation with H O as a co-catalyst
2
O that acts as co-catalyst in CO oxidation or
2
activation, but which
2
2 3
ern University) kindly provided the initial Au/Al O catalysts and
2
synthetic protocols. We thank Dr. David Flaherty for proofreading
and helpful discussions. We are also grateful for the electron
microscopy data provided by Dr. M. Avalos, L. Rendon and F. Ruiz
from the Centro de Ciencias de la Materia Condensada, UNAM, Mex-
ico. The World Gold Council (WGC) is also acknowledged for sup-
plying the reference Au samples.
2
2
occurs on supported noble metal clusters via common elementary
steps involving the formation of hydroperoxy species that are

M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
101
Appendix A
Since r
6
7
= r ,
ꢀ
2
2=3 1=3
2=3
k
k
6
½ OHꢆ
ðK
2
K
3
1=3
K
4
k
5
Þ
k6 ðPO2
P
H2
O
1=3
Þ
ꢀ
The rate of CO oxidation via the sequence of elementary steps
shown in Scheme 1 is given by
½
O ꢆ ¼
¼
½ꢀꢆ
ꢀ
2=3
7
½CO ꢆ
K
ð2Þ
k
P
1
7
CO
2=3
r
CO2 ¼ r
5
þ r
7
ðA1Þ
ðP
O
2
P
H
2
O
Þ
¼
c
½ꢀꢆ
ðA17Þ
ðA18Þ
1
=3
P
CO
By applying the pseudo-steady-state approximation (PSSA) to
ꢀ
ꢀ
O and OH species, we find that r
Consequently, Eq. (A1) becomes
6
= r
7
and r
5
/2 = r
6
, respectively.
where
2
=3 1=3
ðK
2
K
3
1=3
K
4
K
5
Þ
k6
c
¼
3
2
3
2
ꢀ
ꢀ
2=3
r
CO2
¼
r
5
¼
k
5
½CO ꢆ½ OOHꢆ=½Lꢆ
ðA2Þ
K1 ð2Þ
k
7
The site balance can be also expressed as the following:
where [L] is the number of active sites.
2=3
The quasi-equilibrium assumption for steps 1–4 gives
ðPO2
P
H2
1=3
O
Þ
½
Lꢆ ¼ ½ꢀꢆ þ K
1
P
CO½ꢀꢆ þ K
2
P
O2 ½ꢀꢆ þ K
3
P
H2
O
½ꢀꢆ þ b
ꢀ
½
½
½
½ꢀ
CO ꢆ ¼ K
1
P
CO½ꢀꢆ
ðA3Þ
ðA4Þ
ðA5Þ
P
CO
ꢀ
2
=3
O ꢆ ¼ K
O ꢆ ¼ K
2
P
O2 ½ꢀꢆ
3
2
1=3
ðPO2
P
H2
1=3
O
Þ
ꢀ
ꢄ ½ꢀꢆ þ dðPCO
P
O2
P
H2
O
Þ
½ꢀꢆ þ
c
½ꢀꢆ
ðA19Þ
H
2
P
H2
O
½ꢀꢆ
P
CO
ꢀ
Oꢀꢆ
K
4
½O ꢆ½H
2
2
OOH ¼
ðA6Þ
ðA7Þ
ðA8Þ
ðA9Þ
Therefore,
ꢀ
½
OHꢆ
½
Lꢆ
ꢀ
The PSSA to OH gives
½ꢀꢆ ¼
2=3
2=3
ðP
O
P
H
O
Þ
1=3
ðP
O
P
2
H O
Þ
1=3
2
2
2
1
þ K
1
PCO þ K
2
PO2 þ K
3
P
H2
O
þ b
1=3
þ dðPCO
P
O2
P
H2
O
Þ
þ
c
P
P
CO
CO
2
OOH ¼ 2
k
6
½OHꢆ
½ꢀ
ðA20Þ
ꢀ
k
5
½CO ꢆ
Substituting Eq. (A20) in Eq. (A10)
Combining Eqs. (A3)–(A7)
2
=3
a
ðPCO
P
O2
P
H2
2=3
O
Þ
ꢀK K
ꢁ
r
CO
2
¼
2=3
1
=3
ðPO2
P
Þ
ðPO2
P
P
Þ
H2O
1=3
CO
1=3
H2O
1=3
CO
2
1
2
K
2k
6
3
K
4
K
5
1=3
½1 þ K P þ K PO2 þ K PH2O þ b
1
CO
2
3
þ dðPCOPO2 PH2OÞ
þ
c
ꢆ
½ꢀ
OHꢆ ¼
ðPCO
P
P
Þ
½ꢀꢆ
½ꢀꢆ
P
O2
H2
O
ðA21Þ
Substituting Eq. (A8) in Eq. (A6)
2
=3
1=3
2=3
Appendix B. Supplementary data
ðK
2
K
3
K
4
3
Þ
ð2k
6
Þ
ðP 2
O
P
P
H2
1=3
CO
O
Þ
½ꢀ
OOHꢆ ¼
1
=3
ðK
K
5
Þ
Discussion on the CO oxidation kinetics corruptions imposed by
2
O condensation and concomitant mass transfer limitations in-
H
Hence, the rate of CO oxidation (Eq. (A2)) becomes
3
side the
c
-Al
2
O
3
pores; TEM images and Au cluster size distribution
samples; kinetic
for the untreated and thermally-treated Au/Al
2 3
O
2
=3
1=3
2=3
2
r
CO2
¼
¼
ðK
1
K
2
K
3
K
4
k
5
Þ
ð2k
6
Þ
ðPCO
P
O2
P
H2
O
Þ
½ꢀꢆ =½Lꢆ
2
a
parameters obtained from nonlinear fitting to Eq. (1) of rate data
measured with Au/Al O ; kinetic dependence of CO oxidation turn-
2
=3
2
2
3
ðPCO
P
O2
P
H2
O
Þ
½ꢀꢆ =½Lꢆ
ðA10Þ
ðA11Þ
over rates on CO pressure and O
mally treated Au/Al catalysts; and examples of alternate CO
oxidation pathways, the derived kinetic expressions and the corre-
sponding parity plots for the kinetic data measured with Au/Al
288 K; 0.80–8.25 kPa CO; 0.25–7.15 kPa O ; 0.03–1.15 kPa H
2
pressure measured with the ther-
2 3
O
where
3
2
=3
1=3
2 3
O
2
O).
a
¼
ðK
1
K
2
K
3
K
4
K
5
Þ
ð2k Þ
6
2
(
2
ꢀ
ꢀ
Supplementary data associated with this article can be found, in
the online version, at doi:10.1016/j.jcat.2011.09.015.
[
OOH] and [ OH] can be written as follows:
2
=3
ðP 2
O
P
P
H2
O
Þ
½ꢀ
OOHꢆ ¼ b
½ꢀꢆ
ðA12Þ
1
=3
References
CO
[
[
1] G.C. Bond, C. Louis, D.T. Thompson, Catalysis by Gold, Imperial College Press,
London, 2006.
2] B.K. Min, C.M. Friend, Chem. Rev. 107 (2007) 2709.
and
½ꢀ
OHꢆ ¼ dðPCO
P
P
Þ
1=3
½ꢀꢆ
ðA13Þ
ðA14Þ
[3] T. Ishida, M. Haruta, Angew. Chem. Int. Ed. 46 (2007) 7154.
O2
H2
O
[
[
[
4] M. Haruta, T. Kobayashi, H. Sano, N. Yamada, Chem. Lett. 16 (1987) 405.
5] M.C. Kung, R.J. Davis, H.H. Kung, J. Phys. Chem. C 111 (2007) 11767.
6] M.M. Schubert, S. Hackenberg, A.C. van Veen, M. Muhler, V. Plzak, R.J. Behm, J.
Catal. 197 (2001) 113.
where
2
=3
2=3
ðK
2
K
3
K
4
Þ
1
ð2k Þ
6
[
7] G.C. Bond, D.T. Thompson, Gold Bull. 33 (2000) 41.
b ¼
1
=3
ðK
k
5
Þ
[8] C.J. Jia, Y. Liu, H. Bongard, F. Schüth, J. Am. Chem. Soc. 132 (2010) 1520.
9] V. Aguilar-Guerrero, B.C. Gates, Catal. Lett. 130 (2009) 108.
[
[
[
[
10] M. Daté, M. Haruta, J. Catal. 201 (2001) 221.
and
11] F. Gao, T.E. Wood, D.W. Goodman, Catal. Lett. 134 (2010) 9.
12] P. Landon, J. Fergusson, B.E. Solsona, T. Garcia, S. Al-Sayari, A.F. Carley, A.A.
Herzing, C.J. Kiely, M. Makkee, J.A. Moulijn, A. Overweg, S.E. Golunski, G.J.
Hutchings, J. Mater. Chem. 16 (2006) 199.
ꢀK
ꢁ
1
=3
1
2 3 4 5
K K K K
2k
6
d ¼
ðA15Þ
ðA16Þ
[
[
13] H. Liu, A.I. Kozlov, A.P. Kozlova, T. Shido, Y. Iwasawa, Phys. Chem. Chem. Phys.
1
(1999) 2851.
A site balance gives
14] M. Daté, M. Okumura, S. Tsubota, M. Haruta, Angew. Chem. Int. Ed. 43 (2004)
2129.
ꢀ
ꢀ
ꢀ
ꢀ
ꢀ
ꢀ
½
Lꢆ ¼ ½ꢀꢆ þ ½CO ꢆ þ ½O ꢆ þ ½H
2
O ꢆ þ ½ OOHꢆ þ ½ OHꢆ þ ½O ꢆ
2
[15] F. Boccuzzi, A. Chiorino, J. Phys. Chem. B 104 (2000) 5414.

1
02
M. Ojeda et al. / Journal of Catalysis 285 (2012) 92–102
[
16] C.K. Costello, J.H. Yang, H.Y. Law, Y. Wang, J.N. Lin, L.D. Marks, M.C. Kung, H.H.
Kung, Appl. Catal. A: Gen. 243 (2003) 15.
17] Z. Wu, S. Zhou, H. Zhu, S. Dai, S.H. Overbury, J. Phys. Chem. C 113 (2009) 3726.
18] W.C. Ketchie, M. Murayama, R.J. Davis, Topics Catal. 44 (2007) 307.
19] K. Qian, W. Zhang, H. Sun, J. Fang, B. He, Y. Ma, Z. Jiang, S. Wei, J. Yang, W. Huang,
J. Catal. 277 (2011) 95.
20] H.H. Kung, M.C. Kung, C.K. Costello, J. Catal. 216 (2003) 425.
21] V. Aguilar-Guerrero, B.C. Gates, J. Catal. 260 (2008) 351.
22] C.K. Costello, M.C. Kung, H.S. Oh, Y. Wang, H.H. Kung, Appl. Catal. A: Gen. 232
[43] M. Neurock, A. Wieckowski, Adv. Phys. Chem., in press.
[44] M. Boudart, G. Djéga-Mariadassou, Kinetics of Heterogeneous Catalytic
Reactions, Princeton University Press, Princeton, 1984.
[45] C. Shang, Z.P. Liu, J. Phys. Chem. C 114 (2010) 16989.
[46] N. Lopez, J.K. Norskov, J. Am. Chem. Soc. 124 (2002) 11262.
[47] M. Mavrikakis, P. Stoltze, J.K. Norskov, Catal. Lett. 64 (2000) 101.
[48] T. Yan, J. Gong, D.W. Flaherty, C.B. Mullins, J. Phys. Chem. C 115 (2011)
2057.
[
[
[
[
[
[
[49] B. Schumacher, Y. Denkwitz, V. Plzak, M. Kinne, R.J. Behm, J. Catal. 224 (2004)
449.
(
2002) 159.
[
[
[
[
23] S.T. Daniells, A. Overweg, M. Makkee, J.A. Moulijn, J. Catal. 230 (2005) 52.
24] S.D. Lin, M. Bollinger, M.A. Vannice, Catal. Lett. 17 (1993) 245.
25] M.A. Bollinger, M.A. Vannice, Appl. Catal. B: Environ. 8 (1996) 417.
26] C.G. Long, J.D. Gilbertson, G. Vijayaraghavan, K.J. Stevenson, C.J. Pursell,
B.D. Chandler, J. Am. Chem. Soc. 130 (2008) 10103.
27] D.G. Barton, S.G. Podkolzin, J. Phys. Chem. B 109 (2005) 2262.
28] M. Ojeda, E. Iglesia, Chem. Commun. (2009) 352.
29] M. Ojeda, E. Iglesia, Angew. Chem. Int. Ed. 48 (2009) 4800.
30] E.P. Barrett, L.G. Joyner, P.P. Halenda, J. Am. Chem. Soc. 73 (1951) 373.
31] A.D. Allian, K. Takanabe, K.L. Fujdala, X. Hao, T.J. Truex, J. Cai, C. Buda, M.
Neurock, E. Iglesia, J. Am. Chem. Soc. 113 (2011) 4498.
[50] D. Widmann, Y. Liu, F. Schüth, R.J. Behm, J. Catal. 276 (2010) 292.
[51] M. Haruta, S. Tsubota, T. Kobayashi, H. Kageyama, M.J. Genet, B. Delmon, J.
Catal. 144 (1993) 175.
[52] G. Li, J. Edwards, A.F. Carley, G.J. Hutchings, Catal. Today 114 (2006) 369.
[53] M. Okumura, Y. Kitagawa, K. Yamaguchi, T. Akita, S. Tsubota, M. Haruta, Chem.
Lett. 32 (2003) 822.
[54] B. Chowdhury, J.J. Bravo-Suarez, N. Mimura, J. Lu, K.K. Bando, S. Tsubota, M.
Haruta, J. Phys. Chem. Lett. 110 (2006) 22995.
[55] C. Sivadinarayana, T.V. Choudhary, L.L. Daemen, J. Eckert, D.W. Goodman, J.
Am. Chem. Soc. 126 (2003) 38.
[
[
[
[
[
[56] B. Taylor, J. Lauterbach, G.E. Blau, W.N. Delgass, J. Catal. 242 (2006) 142.
[57] P. Landon, P.J. Collier, A.J. Papworth, C.J. Kiely, G.J. Hutchings, Chem. Commun.
18 (2002) 2058.
[
32] C.K. Costello, J. Guzman, J.H. Yang, Y.M. Wang, M.C. Kung, B.C. Gates, H.H.
Kung, J. Phys. Chem. B 108 (2004) 12529.
[
[
[
33] J.A. Biscardi, E. Iglesia, J. Catal. 182 (1999) 117.
34] J.T. Calla, R.J. Davis, Ind. Eng. Chem. Res. 44 (2005) 5403.
35] E. Quinet, L. Piccolo, F. Morfin, P. Avenier, F. Diehl, V. Caps, J.L. Rousset, J. Catal.
[58] C. Rossignol, S. Arrii, F. Morfin, L. Piccolo, V. Caps, J.L. Rousset, J. Catal. 230
(2005) 476.
[59] J. Jiang, S.M. Oxford, B. Fu, M.C. Kung, H.H. Kung, J. Ma, Chem. Commun. 46
(2010) 3791.
2
68 (2009) 384.
[
[
[
36] M. Olea, M. Tada, Y. Iwasawa, J. Catal. 248 (2007) 60.
37] J.T. Calla, M.T. Bore, A.K. Datye, R.J. Davis, J. Catal. 238 (2006) 458.
38] T.V. Choudhary, C. Sivadinarayana, C.C. Chusuei, A.K. Datye, J.P. Fackler, D.W.
Goodman, J. Catal. 207 (2002) 247.
[60] L. Piccolo, H. Daly, A. Valcarcel, F.C. Meunier, Appl. Catal. B: Environ. 86 (2009)
190.
[61] M. Shou, H. Takekawa, D.Y. Ju, T. Hagiwara, D.L. Lu, K.I. Tanaka, Catal. Lett. 108
(2006) 119.
[
39] E. Quinet, F. Morfin, F. Diehl, P. Avenier, V. Caps, J.L. Rousset, Appl. Catal. B:
Environ. 80 (2008) 195.
40] Z.P. Liu, P. Hu, A. Alavi, J. Am. Chem. Soc. 124 (2002) 14770.
41] A. Bongiorno, U. Landman, Phys. Rev. Lett. 95 (2005) 106102.
42] C.R. Chang, Y.G. Wang, J. Li, Nano Res. 4 (2011) 131.
[62] S. Alayoglu, A.U. Nilekar, M. Mavrikakis, B. Eichhorn, Nat. Mater. 7 (2008) 333.
[63] R.H. Nibbelke, M.A.J. Campman, J.H.B.J. Hoebink, G.B. Marin, J. Catal. 171
(1997) 358.
[64] J. Bergeld, B. Kasemo, D.V. Chakarov, Surf. Sci. 495 (2001) L815.
[65] P. Panagiotopoulou, D.I. Kondarides, Catal. Today 112 (2006) 49.
[
[
[