The synergy between Fe and Ru in N2O decomposition over FeRu-FER catalysts: A mechanistic explanation

Article abstract of DOI:10.1016/j.jcat.2006.08.006

Fe-FER is an active catalyst for the abatement of N₂O in the tail gas of nitric acid plants. The activity of Fe-FER can be increased if Ru is added as a second active component. This is a surprising finding, because noble metals are usually strongly inhibited by NO, which is always present in tail gas. Yet the bimetallic FeRu-FER catalyst is more active than the sum of the components, Fe-FER and Ru-FER. A synergy between Fe and Ru can explain this phenomenon. This work discusses the role of Fe and Ru in the reaction mechanism as well as the interplay of these two components. In situ IR measurements show that the preferential adsorption of NO and its reaction products on Fe in the bimetallic catalyst reduces the inhibiting effect of NO on the Ru component; this effect largely contributes to the synergy between Fe and Ru. Moreover, in situ X-ray absorption data are presented, which allow for tracing the average oxidation state of the two active components Fe and Ru under reaction conditions.

Full text of DOI:10.1016/j.jcat.2006.08.006

Journal of Catalysis 243 (2006) 340–349
www.elsevier.com/locate/jcat
The synergy between Fe and Ru in N2O decomposition
over FeRu-FER catalysts: A mechanistic explanation
Gerhard D. Pirngruber a,∗,1, Lukas Frunz a, Johannis A.Z. Pieterse b
a
Institute for Chemical and Bioengineering, ETH Zurich, CH-8093 Zurich, Switzerland
b
ECN Hydrogen & Clean Fossil Fuels, PO Box 1, Westerduinweg 3, 1755 ZG, Petten, The Netherlands
Received 10 May 2006; revised 8 August 2006; accepted 12 August 2006
Abstract
Fe-FER is an active catalyst for the abatement of N O in the tail gas of nitric acid plants. The activity of Fe-FER can be increased if Ru is
2
added as a second active component. This is a surprising finding, because noble metals are usually strongly inhibited by NO, which is always
present in tail gas. Yet the bimetallic FeRu-FER catalyst is more active than the sum of the components, Fe-FER and Ru-FER. A synergy between
Fe and Ru can explain this phenomenon. This work discusses the role of Fe and Ru in the reaction mechanism as well as the interplay of these
two components. In situ IR measurements show that the preferential adsorption of NO and its reaction products on Fe in the bimetallic catalyst
reduces the inhibiting effect of NO on the Ru component; this effect largely contributes to the synergy between Fe and Ru. Moreover, in situ
X-ray absorption data are presented, which allow for tracing the average oxidation state of the two active components Fe and Ru under reaction
conditions.
© 2006 Elsevier Inc. All rights reserved.
Keywords: In situ spectroscopy; NO adsorption; XANES; IR; Ferrierite
1. Introduction
[12,13], precluding their use in the tail gas of nitric acid facto-
ries. Recently it was shown that using a bimetallic FER catalyst
containing iron and ruthenium significantly enhances N2O de-
composition if some NO is present [14,15]. This effect also
manifests in the presence of O2 and H2O. The bimetallic cat-
alyst is more active than the sum of the single components,
Fe and Ru; there is a synergy between Fe and Ru [16]. Pos-
sible explanations for this synergy have been offered [15], but
the effect remains incompletely understood. The goal of the
present contribution is to unravel the mechanism by which Fe
and Ru cooperate, with the aid of in situ spectroscopy (IR and
XANES).
Direct catalytic decomposition of N2O in the tail gases of
nitric acid plants is a safe, cost-efficient method for N2O abate-
ment. Consequently, much effort is invested in developing im-
proved catalysts that convert N2O at moderate temperatures
(below 773 K). Iron zeolites, especially Fe-ZSM-5, have been
studied extensively for use as catalysts in N2O decomposition
[1–5]. NOx, which is usually present in off-gases, promotes
N2O decomposition over iron zeolites [6–9]. Fe-FER [4,10] and
Fe-BEA [5] perform particularly well. However, a temperature
of 773 K is still necessary to achieve 75% N2O conversion with
Fe-FER in the presence of H2O and NOx [10]. Noble metal cat-
alysts have a higher intrinsic activity for N2O decomposition
compared with iron [11]. Unfortunately, N2O decomposition
over noble metal catalysts is inhibited by NO, O2, and H2O
2. Experimental
2.1. Catalyst preparation
NH4-FER was obtained from Na/K-FER (Tosoh HSZ-
720KOA, Si/Al = 9.2) by threefold ion exchange with a
NH4NO3 solution for 1 h at room temperature. The sample
was filtered and washed with a large amount of demineralized
*
Corresponding author.
E-mail address: gerhard.pirngruber@ifp.fr (G.D. Pirngruber).
Present address: Catalysis and Separation Division, Institut Français du Pét-
1
role, 69390 Vernaison, France.
0021-9517/$ – see front matter © 2006 Elsevier Inc. All rights reserved.
doi:10.1016/j.jcat.2006.08.006

G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
341
Table 1
temperature. Pseudo-first-order rate constants were calculated
using the formula
Elemental composition of the samples and intensity of Brønsted OH band
a
Sample
Fe
Ru
Fe/Al Ru/Al (Na + K) / Exchange
F
(wt%) (wt%)
Al
Brønsted OH (%)
k = −
· ln(1 − X),
mcat · p
b
Fe-FER
FeRu-FER
Ru-FER
Fe-FER-2
FeRu-FER-2 0.6
2.2
2.2
n.d.
0.6
n.d.
0.4
0.4
n.d.
0.45
0.45
0.31
–
–
17
32
28
20
68
68
0.31 0.03
0.22
0.29
–
where F is total flow, mcat mass of the catalyst, p is total pres-
sure, and X is the conversion. The equilibrium between NO and
NO2 was calculated using HSC software.
–
0.03
–
0.09
0.09 0.04
0.04
0.73
n.d.
Ru-FER-2
n.d.
–
2.3. In situ IR spectroscopy
a
Mainly K, only traces of Na.
n.d. = not determined.
b
A 6-mg catalyst sample was pressed into a self-supporting
pellet, placed in a gold sample holder, and inserted in the in situ
IR cell [17]. The sample was heated in a flow of 5% O2 in He to
673 K and kept there for 30 min, to remove all impurities. Sub-
sequently, the cell was purged for 30 min in He flow at 673 K
to facilitate autoreduction of the catalyst. The cell was cooled
to 573 K. A spectrum of the catalyst was recorded and used
as a reference for the subsequent measurements. The reaction
was started by switching from He to a mixture of ∼3000 ppm
N2O, ∼800 ppm NO, and 0–2% O2. The concentration of NO
and N2O was twice as great as in the catalytic tests described
above. The range of the mass flow controllers imposed this
choice. The gas flow was kept at 20 mlNTP/min, correspond-
ing to a GHSV of ∼100,000 h−1. After ∼30 min (once steady
state was reached), the gas inlet was switched back to He, the
catalyst was regenerated at 673 K for 30 min, and the proce-
dure was repeated at the next reaction temperatures (i.e., 623
and 673 K). The IR spectra were recorded on a Biorad FTS
3000 MX spectrometer equipped with a broadband MCT detec-
tor. The spectral resolution was 4 cm−1. A total of 16–128 scans
were co-added. 16 scans correspond to a time resolution of
∼10 s. The reactor effluent was analyzed by mass spectrom-
etry [17]. To remove impurities of NO2 in the feed, the NO/He
mixture was passed through a cold trap at 173 K.
water to remove nitrate. Fe-FER (2.2 wt% Fe) was prepared by
impregnating NH4-FER with iron nitrate.
Ru-FER was obtained by exchanging Ru(NH4)3Cl6 with
25 g Na/K-FER at 353 K for 16 h (nominal Ru loading, 0.5%;
initial pH, 8; final pH, 6.5), followed by two exchanges with
0.05 M NH4NO3 for 1 h. The purpose of the treatment with
NH4NO3 was to maximize the exchange capacity for the sub-
sequent impregnation with iron nitrate (vide infra). The final Ru
loading was determined by ICP to be 0.40 wt%.
FeRu-FER was obtained by incipient-wetness impregnation
of Ru-FER with a solution of iron nitrate, to obtain 2.2 wt% Fe.
The samples were calcined under flowing air at 793 K for 3 h
(at a ramp of 3 K/min).
A second batch of Ru-FER (coded as Ru-FER-2) was pre-
pared at a larger scale, in which 100 g of Na/K-FER was
exchanged with Ru(NH4)3Cl6 for 16 h at 353 K. The sample
was washed and filtered, followed by a single exchange with
0.05 M NH4NO3 for 1 h. The final loading was determined by
ICP to be 0.45 wt% Ru. FeRu-FER-2 was obtained by incipient-
wetness impregnation of Ru-FER-2 (0.45 wt% Ru) with a more
diluted solution of iron nitrate to obtain 0.6 wt% Fe. Similarly,
Fe-FER-2 was obtained by impregnation of NH4-FER with a
diluted solution to obtain 0.6 wt% Fe (Table 1).
The samples were characterized by UV–vis, TPR, IR, and
XANES/EXAFS spectroscopy. UV–vis spectra were recorded
on a Cary 400 UV–vis spectrometer with a Praying Mantis sam-
ple stage from Harrick. H2-TPR spectra were recorded with an
Altamira AMI-1 apparatus equipped with a thermal conductiv-
ity detector and a Balzers MS-detector, applying 30 ml/min
flow of 10% H2 in argon at a heating rate of 20 K/min. IR and
XAS are explained in more detail below. The elemental com-
position of the samples was determined by ICP or AAS (for Na
and K).
2.4. In situ XANES spectroscopy
The in situ XANES measurements were performed at beam-
line BM26 (DUBBLE) at ESRF, Grenoble. An EXAFS cell
dedicated for in situ measurements in the fluorescence mode
was used [18]. A 10-mg catalyst sample was gently pressed into
the sample compartment through which the gas flow passes. An
Al foil (99.999% purity, 15 µm thickness) was used to seal the
sample and also served as the window for the EXAFS radiation.
A double-crystal Si(111) monochromator selected the energy
of the X-rays, and a nine-channel monolithic Ge detector was
used to collect the fluorescence radiation. Measurements were
performed at both the Fe K-edge and the Ru K-edge. The scan
time was ∼8 min.
2.2. Activity measurements
The catalytic tests were conducted in a computer-controlled
six-flow setup. First, 50 mg of catalyst (sieve fraction, 0.25–
0.5 mm) was placed on a quartz grid, then the quartz reactors
(4 mm i.d.) were placed in an oven. The total gas flow was
100 ml/min, corresponding to a GHSV of ∼60,000 h−1. Quan-
titative analysis of the gas-phase components was performed
using a micro-gas chromatograph and a NOx analyzer. Data
were collected at ascending temperatures from 533 to 773 K
(ramp 5 K/min). Preconditioning was set for 20 min at each
In the standard procedure, the sample was heated in 3%
O2 in He to 673 K and kept at that temperature for 30 min.
After the cell was flushed in He for 30 min, three different re-
actions were performed: (i) 3000 ppm N2O + 800 ppm NO,
(ii) 3000 ppm N2O + 800 ppm NO + 2% O2, and (iii) 5000
ppm N2O and the balance He. Each reaction was performed
in situ for 30 min, followed by purging in He for 30 min. The

342
G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
Fig. 2. H -TPR of Fe-FER, FeRu-FER and Ru-FER.
2
peratures (Fig. 2) similar to the reduction for Fe2O3 [21]. The
H2/Fe consumption up to 1050 K is 0.90. Assuming that this
figure results from a mixture of large Fe2O3 clusters that re-
duce to Fe0 and of other isolated or oligonuclear iron species
that reduce only to Fe2+, we can deduce that the fraction of
Fe2O3 should be 40%.
Ru-FER has a rather weak Brønsted OH-band despite the
low Ru/Al ratio, because the sample still contains Na+ and
K+ counter ions (Table 1). The UV–vis spectrum of Ru-FER
(Fig. 1) is similar to that of RuO2; its features can be explained
by the band structure of RuO2 [22]. The strong absorption be-
low 17,000 cm−1 is due to absorptions within the conduction
band, composed of Ru 4d orbitals. The transitions at higher
energies (i.e., 28,500 and 43,500 cm−1) can be assigned to
transitions from the valence band (i.e., from O 2p orbitals) to
the conduction band. Compared with crystalline RuO2, these
are strongly red-shifted, due to the reduction of the band gap
as a result of the small cluster size [23]. The EXAFS/XANES
spectrum of Ru-FER is very similar to that of a RuO2 refer-
ence sample. In the Fourier-transformed EXAFS spectrum, the
similarity extends up to 4 Å, corresponding to two complete
coordination shells. From this, we can deduce that the average
size of the RuO2 clusters should be at least 1–2 nm.
Fig. 1. (a) UV–vis spectra of Fe-FER, FeRu-FER and Ru-FER. (b) Zoom on the
low wavenumbers. Spectra are offset for better visibility.
total gas flow was 17.5 mlNTP/min, corresponding to GHSV
∼53,000 h−1, and the reaction temperature was 673 K. The
reactor effluent was analyzed by mass spectrometry. After the
reaction sequence was completed, Fe-FER was reduced in H2
at 673 K, to generate a reference spectrum of a fully reduced
Fe2+ sample.
3. Results
The UV–vis absorption of Ru-FER is much weaker than that
of Fe-FER; thus, it is not surprising that the UV–vis spectrum
of FeRu-FER more closely resembles Fe-FER than Ru-FER.
However, closer inspection of the low-energy region of the
spectrum shows that FeRu-FER also exhibits free electron ab-
sorption within the conduction band, which is typical for RuO2
clusters (Fig. 1b). The TPR spectrum of FeRu-FER is similar
to that of Fe-FER, but with an additional peak at low tempera-
ture due to the reduction of ruthenium (Fig. 2). The subsequent
reduction of iron shifts to lower temperatures because Ru aids
the reduction.
Fe-FER-2 was prepared by incipient wetness impregnation
of a more dilute iron nitrate solution than Fe-FER. In contrast
to Fe-FER, Fe-FER-2 is colorless. Its UV–vis spectrum (not
shown) exhibits only the two LMCT transitions at 42,000 and
36,500 cm−1. Oligonuclear iron species are absent in this sam-
ple. Although the iron loading is three times lower than that in
3.1. Characterization
The Fe/Al ratio of Fe-FER is 0.31, corresponding to an ex-
change degree of 62% (Table 1). Yet the Brønsted OH band
at 3590 cm−1 loses only 17% of its intensity compared with
the parent H-FER sample. This effect is due to the impregna-
tion method used, which results in physical deposition of iron
on the surface rather than to a true ion exchange. The UV–vis
spectra of Fe-FER shows an intense band at 29,000 cm−1 in ad-
dition to the typical O → Fe LMCT transitions at 42,000 and
36,500 cm−1 (Fig. 1). The band at 29,000 cm−1 is character-
istic of oligonuclear iron species [19,20]. Additional shoulders
at 23,000 and 20,000 cm−1 demonstrate that the sample also
contains larger Fe2O3 particles, which are probably deposited
on the outer surface of the zeolite during impregnation. The
presence of Fe2O3 was confirmed by H2-TPR showing that a
considerable fraction of the iron sites reduced at very high tem-

G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
343
Table 2
Pseudo first order rate constants of N O decomposition in a feed of 1500 ppm
2
a
N O, 400 ppm NO, 0 or 2% O
2
2
T
k (mmol/(s g bar))
With O
(K)
Without O
2
2
Ru-FER Fe-FER FeRu-FER Ru-FER Fe-ER FeRu-FER
652
672
692
722
0.00
0.03
0.12
0.35
0.30
0.49
0.76
1.48
0.45
0.81
1.44
2.93
0.09
0.24
0.49
1.19
0.35
0.57
0.89
1.65
0.86
1.40
2.25
3.96
a
−1
GHSV = 60,000 h , p = 1 bar.
Fe-FER-2, the fraction of exchanged Brønsted sites is similar
(Table 1). This confirms that a large part of the additional iron
in Fe-FER is not exchanged with Brønsted sites, but is only de-
posited on the catalyst.
3.2. N2O decomposition activity
The main results of the catalytic tests with the first batch of
samples were presented earlier [15] and are only summarized
here. The N2O decomposition activity in a feed of 1500 ppm
N2O + 400 ppm NO (+2% O2) increases in the order Ru-
FER < Fe-FER < FeRu-FER. The pseudo-first-order rate con-
stants of FeRu-FER are higher than the sum of the two other
catalysts (Table 2). Adding O2 to the feed decreases the activ-
ity of all three catalysts; the inhibiting effect of O2 decreases
in the order Ru-FER > FeRu-FER ꢀ Fe-FER. The NO2 con-
centration profiles show some interesting differences between
the catalysts (Fig. 3). NO2 can be formed by one of two reac-
tions:
Fig. 3. NO concentration profile during NO decomposition with 1500 ppm
2
N O, 400 ppm NO and (a) 2 or (b) 0% O over Fe-FER (2), FeRu-FER (P),
2
2
Ru-FER (!). Solid lines represent equilibria (a) NO + (1/2) O → NO (feed
2
2
400 ppm NO + 2% O ) and (b) NO → NO + (1/2)O (feed 400 ppm NO ).
2
2
2
2
NO + (1/2)O2 → NO2
or
(1)
NO + N2O → NO2 + N2.
(2)
Reaction (1) is equilibrium-limited, but reaction (2) is not.
Its ꢀG0r at 673 K is −135 kJ/mol. When O2 is present in
the feed, Ru-FER and FeRu-FER establish the equilibrium
between NO and NO2 above 570 and 590 K, respectively
(Fig. 3). On Fe-FER, the NO2 concentration exceeds the
NO/NO2 equilibrium, because the source of NO2 is mainly
reaction (2). The NO2 formation reaches its maximum at
685 K. At higher temperatures, NO2 is rapidly consumed by
reaction with N2O, and its concentration decreases (vide in-
fra).
Fig. 4. N O conversion in a feed of 1500 ppm N O, 400 ppm NO, balance N .
GHSV = 60000 h , p = 1 bar.
2
2
2
−1
In the absence of O2, Ru-FER is not active for NO2 for-
mation. Fe-FER forms little NO2 at low temperatures, demon-
strating that NO2 formation at these temperatures is due mainly
to reaction (1). Reaction pathway (2) is hardly affected by the
presence or absence of O2. The NO2 profile of Fe-FER at
high temperatures is very similar to the one in the presence of
O2. FeRu-FER behaves like Fe-FER at low temperatures. The
NO2 concentration reaches a maximum at 595 K and then ap-
proaches the equilibrium NO2 → NO + (1/2)O2; that is, the
NO/NO2 ratio shifts towards NO with increasing temperature.
Fig. 3 does not show the NO concentration, but because the sum
of NO and NO2 is constant at 400 ppm, the profile of NO can
be deduced.
The catalysts of the second batch are more active than those
of the original preparation (Fig. 4). The higher activity of Fe-
FER-2 compared with Fe-FER confirms that a large fraction
of iron in Fe-FER does not participate in the reaction. The
higher activity of Ru-FER-2 compared with Ru-FER can be at-
tributed to slight differences in the preparation procedures. As
for the first batch, the first-order rate constants of FeRu-FER-2
are larger than the sum of Fe-FER-2 and Ru-FER-2 (see Ta-
ble 3).

344
G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
Table 3
however, the Fe–OH and the nitrate/nitro bands are stable and
decrease very slowly over time. The spectra at 623 K are very
similar, but the nitrosyl band is much weaker. At 673 K, the
intensity of the nitrosyl band is zero, and only nitrates are ob-
served.
Pseudo first order rate constants of N O decomposition in a feed of 1500 ppm
N O and 400 ppm NO
2
2
a
T
k (mmol/(s g bar)), without O
2
(K)
Ru-FER-2
Fe-FER-2
FeRu-FER-2
Fig. 6 shows the corresponding spectra with Ru-FER. Some
differences can be immediately observed. No OH-groups are
created, and the negative band of the Brønsted OH groups is
much less intense. Likewise, the NO+ band is missing. The ni-
trosyl band is at a different wavelength (i.e., 1876 cm−1), and
the nitro/nitrate band is shifted to 1634 cm−1. In contrast to Fe-
FER, the nitrosyl band is rather stable and decreases only after
several minutes of He purging. The nitro/nitrate band is also
very stable during the He purge. The spectra at higher temper-
atures are almost identical to those at 573 K. The intensity of
the nitrosyl and nitro band hardly decreases, indicating the high
stability, particularly of the nitrosyl group, on the RuO2 sur-
face. The high stability of the surface nitrosyl is probably at the
heart of the inhibiting effect of NO on decomposition activity
of Ru-FER.
The objective of the study was to compare the behavior
of the bimetallic catalyst with the single components Fe-FER
and Ru-FER. Fig. 7 shows the IR spectra of all three sam-
ples measured in steady state at 573 and 623 K. The spectra
of FeRu-FER resemble those of Fe-FER with respect to band
intensity, band position, and shape. The similarity of Fe-FER
and FeRu-FER is also observed in the time-on-stream behavior
(see Fig. 8). Both Fe-FER and FeRu-FER show a rapid initial
increase of the nitrosyl band, followed by a decrease of the
intensity between 1 and 5 min, which is ascribed to the oxi-
dation of NO species to NOx. This effect is not observed for
Ru-FER. During He purging, the nitrosyl species desorb very
652
672
692
722
0.17
0.41
0.84
2.12
0.49
0.79
1.22
2.25
1.14
2.06
3.49
6.85
a
−1
GHSV = 60,000 h , p = 1 bar.
3.3. In situ IR spectroscopy in absence of O2
Fig. 5 shows the IR spectra of Fe-FER during reaction with
NO + N2O at 573 K. An intense nitrosyl band at 1866 cm−1
appears almost immediately after switching to the reaction mix-
ture. After 1 min, its intensity gradually decreases to a steady-
state value. A band at 1625 cm−1 slowly increases in intensity
and reaches steady state after ∼15 min. It is accompanied by a
broad shoulder at 1550 cm−1. Both bands can be assigned to ni-
tro/nitrate species [24,25]. The doublet at 2238 and 2205 cm−1
is due to gas-phase N2O. The intensity of the N2O band seems
to increase over time. The increase is due to a broad band
of NO+ hidden below the N2O doublet. The NO+ band be-
comes visible when switching from the reaction mixture back
to He. In the region of the OH stretching vibrations, a band at
3646 cm−1 and a shoulder at 3605 cm−1 appear, which can
be assigned to Fe–OH stretching vibrations [26]. The nega-
tive band at 3575 cm−1 is due to the replacement of Brønsted
protons by NO+ [27]. The nitrosyl band rapidly disappears on
switching back to He. The NO+ band and the corresponding
negative Brønsted OH band disappear within a few minutes;
Fig. 5. IR spectra measured during reaction of Fe-FER with NO + N O at 573 K, after 0.5, 1, 2, 4, 8 and 18 min on stream.
2

G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
345
Fig. 6. IR spectra measured during reaction of Ru-FER with NO + N O at 573 K, after 0.5, 1, 2, 4, 8 and 18 min on stream. The bands of gas phase CO are an
2
2
artifact that originates from the air in the spectrometer compartment.
−1
Fig. 8. Intensity of the NO (1866 or 1877 cm , full symbols) and of the ni-
−1
tro/nitrate bands (1625 or 1635 cm , open symbols) as a function of time on
stream during the reaction of the catalysts with NO + N O at 573 K.
2
rapidly from Fe-FER and FeRu-FER, whereas the intensity of
the NO band decreases only gradually in Ru-FER (see Fig. 9),
due to the higher stability of the nitrosyl species adsorbed on
Ru. In conclusion, we can state that the bimetallic catalyst be-
haves like Fe-FER in the in situ IR experiments.
3.4. In situ IR spectroscopy in the presence of O2
Fig. 10 compares the IR spectra of Fe-FER in the presence
and absence of O2 in the feed. In the presence of O2, the in-
tensity of the surface nitrosyl band on Fe-FER decreases, and
the nitro groups increase in intensity. Among the nitro bands,
a band at 1575 cm−1 becomes particularly more intense. This
band was observed earlier during the reaction of NO with O2
on iron zeolites [28,29]. Moreover, the intensity of the Fe–
Fig. 7. IR spectra of Fe-FER (—), Ru-FER (—) and FeRu-FER (-·-·-) at
(a) 573 K and (b) 623 K after 30 min reaction with NO + N O.
2

346
G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
−1
Fig. 9. Intensity of the NO band (1866 or 1877 cm ) as function of time during
Fig. 11. IR spectra of Ru-FER after 30 min reaction with NO + N O at 573 K
2
the He purge, after reaction of the catalysts with NO + N O at 573 K.
2
in the presence and absence of O .
2
Fig. 12. IR spectra of Fe-FER, Ru-FER and FeRu-FER at 623 K after 30 min
reaction with NO + N O + O .
Fig. 10. IR spectra of Fe-FER after 30 min reaction with NO + N O at 573 K,
2
2
2
in the presence and absence of O .
2
toward the NO2-derived species, but the equilibrium remains in
favor of NO.
OH groups increases slightly compared with the reaction with-
out O2. Because the Fe–OH groups arise from the oxidation of
Fe2+ to Fe3+, it is not surprising that their intensity increases
in the presence of O2.
For FeRu-FER, shape and position of the NO and NOx bands
in the presence of O2 are similar to those of Fe-FER. At 573 K,
the nitrosyl band has similar intensity on all three catalysts.
The absorbance of the nitro/nitrate bands decreases in the or-
der FeRu-FER > Fe-FER > Ru-FER (not shown). The spectra
of the three samples at 623 K are shown in Fig. 12. As in the
absence of O2, the Ru-FER catalyst maintains a rather intense
nitrosyl band at 623 K, whereas the nitrosyl band of Fe-FER
and FeRu-FER is weak. The intensity of the nitro/nitrato bands
is highest on FeRu-FER, probably due to the higher fraction of
Fe3+ sites in FeRu-FER (see Section 3.5). We can postulate that
the NO adsorption on FeRu-FER is similar to that on Fe-FER,
also when O2 is present in the feed.
The spectra of the three samples of the second batch are
shown in Fig. 13. The region of nitrate/nitro bands is better re-
solved, and a number of additional bands appear. We do not
discuss or assign these bands in detail here; for our purposes,
the important information is that the spectra of FeRu-FER-2
do not resemble those of Fe-FER-2. The intensity and posi-
tion of the nitrosyl band in the bimetallic sample are closer to
those of Ru-FER-2. The band is very stable and desorbs only
The situation is similar for FeRu-FER. Compared with the
reaction without O2, the nitrosyl band decreases and the nitro
bands (particularly a band at 1575 cm−1) increase in inten-
sity. For Ru-FER, the nitro bands also increase (Fig. 11), and
a second NO band at 1890 cm−1 becomes visible. The higher
frequency of the NO stretching vibration indicates less back-
donation from Ru to the NO 2π∗ antibonding orbital. In the
presence of O2, the RuO2 surface is more oxidized (i.e., the
electron density at the Fermi level is reduced) and thus less
willing to donate electrons to NO [30]. Surprisingly, the total
intensity of the nitrosyl band decreases by only 15% compared
with that in the reaction without O2. We attribute this to the
fact that NO- and NO2-derived species are in chemical equilib-
rium on the RuO2 surface. The ratio between surface nitrosyl
and surface nitro species shifts in favor of NO with increasing
temperature, as does the equilibrium between NO and NO2 in
the gas phase. The presence of O2 increases the total concentra-
tion of the surface NO/NO2 species and shifts the equilibrium

G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
347
Fig. 13. IR spectra of Fe-FER-2, Ru-FER-2 and FeRu-FER-2 at 623 K after
30 min reaction with NO + N O + O .
2
2
very slowly in He. We can postulate that in FeRu-FER-2, NO
adsorbs on the ruthenium component rather than on the iron
component.
3.5. In situ XANES spectroscopy
The in situ XANES measurements were carried out to follow
the changes in the oxidation state of the catalyst during pre-
treatment, reaction, and regeneration. Fe-FER contains a large
amount of large Fe2O3 clusters, which dominate the X-ray ab-
sorption. These clusters do not reduce easily and do not par-
ticipate in the reaction. Hence, only very small shifts of the
Fe K-edge are seen in the in situ experiments. A quantitative
comparison of the redox behavior of Fe-FER and FeRu-FER
is, therefore, not possible. To circumvent the dominating con-
tribution of the large Fe2O3 clusters on the X-ray absorption
spectra, a second batch of Fe-FER (i.e., Fe-FER-2) was pre-
pared that does not contain oligonuclear iron clusters or Fe2O3
particles. Fe-FER-2 exhibits a high redox activity during the in
situ XANES measurements. Fig. 14 shows the in situ XANES
spectra of Fe-FER-2 in the three reaction mixtures N2O + NO,
N2O + NO + O2, and N2O, as well as after reduction in H2.
The spectrum in N2O corresponds to the fully oxidized state
(Fe3+); the spectrum in H2, to the fully reduced state (Fe2+).
The fully reduced form has a characteristic peak in the edge at
7118.5 eV. The spectra measured during reaction with N2O +
NO (+ O2) correspond to mixed oxidation states. The fraction
of oxidized iron sites was determined by linear combination
of the reference spectra in N2O and H2. During reaction with
NO + N2O, approximately 30% of the iron sites are reduced
to Fe2+ (Table 4). The fraction of reduced sites during reaction
hardly depends on the presence or absence of O2. Note that
O2 does not inhibit N2O decomposition over Fe-FER. Autore-
duction of Fe3+ to Fe2+ occurs when purging with He before
switching to another reaction mixture. The fraction of reduced
sites in He is about 40%.
Fig. 14. In situ XANES spectra of (a) Fe-FER-2 and (b) FeRu-FER-2 during
reaction with N O + NO (- - -), N O + NO + O (-·-·-), N O only (—) and
2
2
2
2
after reduction in H (– – –), at 673 K.
2
Table 4
2+
Fraction of Fe sites during pretreatment in O , autoreduction in He and re-
2
action with NO + N O (+ O ) at 673 K
2
2
Fe-FER-2
FeRu-FER-2
O
He
∼50
19
31
27
30
18
29
2
a
55
NO + N O
33
42
31
40
2
b
He
NO + N O + O
2
2
c
He
a
He purge after treatment with O .
2
b
c
He purge after reaction with NO + N O.
2
He purge after reaction with NO + N O + O .
2
2
creases. Catalytic results indeed show a small inhibiting effect
of O2 on the bimetallic catalyst (Table 2), which may be as-
cribed to the lower concentration of Fe2+ sites (vide infra).
The comparison of Fe-FER-2 and FeRu-FER-2 reveals that the
bimetallic catalyst reduces less readily than the iron sample.
4. Discussion
4.1. Catalytic synergy
The same analysis was performed for FeRu-FER-2 (Ta-
ble 4). During reaction with NO + N2O, FeRu-FER-2 has a
similar fraction of Fe2+ sites as seen in Fe-FER-2. When O2
is added to the reaction mixture, the concentration of Fe2+ de-
To set the basis for discussing the spectroscopic data, we
briefly recall what is already known about the catalytic prop-
erties of iron and ruthenium ferrierite. The mechanism of N2O

348
G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
decomposition on iron zeolites in the presence of NO can be
described by the following set of equations:
This mechanism is in accordance with the kinetic data and
could also operate in the presence of O2 in the feed, albeit with
lower efficiency.
N2O + NO → NO2 + N2,
N2O + NO2 → NO + N2 + O2,
Sum: 2N2O → 2N2 + O2.
(2)
(3)
4.2. Adsorptive synergy
The in situ IR data show that a second mechanism may con-
tribute to the synergy between Fe and Ru. The presence of high
concentrations of iron influences the sorption properties of NO
on the RuO2 particles. In the bimetallic FeRu-FER catalyst, NO
adsorbs mainly on the Fe sites (Figs. 7 and 12). Consequently,
the coverage of the Ru sites with NO is significantly reduced,
especially at high temperatures. NO acts as a strong inhibitor
for N2O decomposition on RuO2. The function of Fe in the
bimetallic catalyst is to trap NO and thereby leave the RuO2
surface free to exert its high intrinsic activity for N2O decom-
position. This effect is not observed on the FeRu-FER-2 catalyst
with low iron loading. FeRu-FER-2 preferentially adsorbs NO
on the Ru component because of the stronger interaction of
NO with Ru compared with Fe. Iron can work effectively as an
NO trap only if present in high concentrations, that is in large
excess and in intimate contact with Ru.
The question may be raised as to whether our IR data are rep-
resentative, because for experimental reasons they were mea-
sured at higher concentrations than the catalytic data. Compar-
ative measurements with similar NO and N2O concentrations as
used for catalysis showed that the spectra do not change much.
Between 400 and 800 ppm NO, the adsorption isotherm of NO
on iron is rather flat. Under conditions farther away from satura-
tion with NO (i.e., at even lower concentrations or temperatures
above 673 K), we can expect NO adsorption to be preferred on
Ru rather than on Fe, and the adsorption synergy loses impor-
tance.
NO functions as intermediate oxygen storage [6–8,31]. It takes
up oxygen from N2O and is converted to NO2. NO2 releases
the oxygen atom again by reaction with a second molecule of
N2O, and O2 is formed. This mechanism greatly enhances the
rate of O2 formation, which otherwise is a very slow step in
N2O decomposition (in the absence of NO) [6,17]. NO thereby
functions as a catalyst for N2O decomposition on iron zeolites.
More detailed discussions of the reaction mechanism and the
role of the surface oxygen species from N2O are provided else-
where [6,8,31].
In Ru-FER, NO is a strong inhibitor for N2O decomposi-
tion; however, the catalyst is very active for the oxidation of
NO by O2. Reaction (1) reaches equilibrium already at 560 K.
Fe-FER is less active for the oxidation of NO by O2. Thus, the
synergistic effect of Fe and Ru can be explained by the reaction
sequence
2NO + O2 → 2NO2 (over Ru),
N2O + NO2 → N2 + O2 + NO (over Fe),
that is, rapid oxidation of NO to NO2 occurs over Ru, and
the NO2 enters the N2O decomposition cycle on Fe. Several
arguments do not support this proposal, however. The NO2
concentration on Ru-FER and Fe/Ru-FER is limited by equi-
librium (1), which is unfavorable for NO2 at higher temper-
atures. On Fe-FER, NO2 is formed by reaction (2), which is
not equilibrium-limited. As a result, the gas-phase NO2 con-
centration at higher temperatures is greater on Fe-FER than on
Ru-FER and on the bimetallic catalyst; that is, the bimetallic
catalyst cannot accelerate the N2O decomposition cycle over
iron by feeding more NO2 to step (2). Moreover, the reaction
sequence proposed above depends on the presence of O2 in the
feed. However, it has been shown that synergy between Ru and
Fe also exists in an O2-free feed mixture [15].
In the absence of O2 in the feed, Ru-FER does not produce
any NO2; that is, it is not active for reaction (1). The bimetallic
catalyst produces NO2 at low temperature. At higher tempera-
tures, NO2 decomposes to NO and O2 [the reverse reaction of
Eq. (1)], and thermodynamic equilibrium among NO2, NO, and
O2 is approached. The catalytic behavior of Fe-FER is hardly
affected by the presence or absence of O2. This allows us to sug-
gest a different mechanism by which Fe and Ru can cooperate:
Fe-FER is very active for the oxidation of NO to NO2 by N2O
[reaction (2)], but O2 formation by reaction of N2O with NO2
[reaction (2)] is rate-limiting. Ru offers an alternative route for
O2 formation by decomposition of NO2 to NO and O2, which
is fast (close to equilibrium). The whole reaction sequence can
be written as
4.3. Redox synergy
N2O decomposition on iron zeolites proceeds via a redox cy-
cle between Fe2+ and Fe3+ [32]. Therefore, it is not surprising
that the concentration of Fe2+ under reaction conditions largely
determines the catalytic activity of Fe-ZSM-5 in N2O decom-
position [33]. The concentration of Fe2+ becomes even more
important in the presence of NO [17], because Fe2+ centers act
as adsorption sites for NO, which promotes N2O decomposi-
tion over iron zeolites (vide supra). The in situ XANES data
show that the fraction of Fe2+ sites in a working iron ferrierite
catalyst is high, ∼30% if Fe2O3 particles are excluded. Adding
Ru to the catalyst does not increase the concentration of Fe2+
;
in the presence of O2, it even decreases (Table 4). Ru does not
have a positive influence on the redox properties of Fe. There is
no redox synergy between Fe and Ru.
4.4. Comparison of Fe-FER and Fe-ZSM-5
We recently studied NO-assisted N2O decomposition on a
series of Fe-ZSM-5 catalysts [17] and found that the behav-
ior of Fe-FER was markedly different from Fe-ZSM-5. On Fe-
ZSM-5, weakly bound surface NO2 species were identified as
intermediates in the NO-assisted N2O decomposition cycle.
N2O + NO → NO2 + N2 (over Fe),
NO2 → NO + (1/2)O2 (over Ru).

G.D. Pirngruber et al. / Journal of Catalysis 243 (2006) 340–349
349
Stable surface nitrates did not seem to be involved in the reac-
tion. On Fe-FER, weakly bound surface NO2 species were not
identified. Stable, probably charged nitro/nitrate species dom-
inated the IR spectra in the region of 1650–1500 cm−1. This
indicates that the reaction intermediates on Fe-FER differ from
those on Fe-ZSM-5.
Kukula (ETH Zurich) for participating in the XAS experi-
ments. We also thank Martin Kuba for the Na/K determina-
tions, Dr. Ruud van den Brink (ECN) for valuable discussions,
and the reviewers for their constructive comments and sugges-
tions.
References
A second, marked difference between Fe-FER and Fe-ZSM-
5 is in terms of redox activity. The Fe sites in Fe-FER read-
ily reduce to Fe2+ (if Fe2O3 clusters are excluded), whereas
XANES measurements on Fe-ZSM-5 show that only a negli-
gible fraction of the iron sites are in the oxidation state Fe2+
during reaction with NO + N2O. A more detailed discussion of
these effects is beyond the scope of this paper, but we believe
that the comparison of Fe-ZSM-5 and Fe-FER merits a more
in-depth investigation [34,35].
[1] M. Kogel, B.M. Abu-Zied, M. Schwefer, T. Turek, Catal. Commun. 2
(2001) 273.
[2] J. Perez-Ramirez, F. Kapteijn, G. Mul, J.A. Moulijn, Chem. Commun.
(2001) 693.
[3] J. Perez-Ramirez, F. Kapteijn, G. Mul, J.A. Moulijn, Appl. Catal. B 35
(2002) 227.
[4] A. Guzman-Vargas, G. Delahay, B. Coq, Appl. Catal. B 42 (2003) 369.
[5] J.A.Z. Pieterse, S. Booneveld, R.W. van den Brink, Appl. Catal. B 51
(2004) 215.
[6] J. Perez-Ramirez, F. Kapteijn, G. Mul, J.A. Moulijn, J. Catal. 208 (2002)
211.
5. Conclusion
[7] C. Sang, C.R.F. Lund, Catal. Lett. 73 (2001) 73.
[8] C. Sang, B.H. Kim, C.R.F. Lund, J. Phys. Chem. B 109 (2005) 2295.
[9] D. Kaucky, Z. Sobalik, M. Schwarze, A. Vondrova, B. Wichterlova,
J. Catal. 238 (2006) 293.
[10] B. Neveu, C. Hamon, K. Malefant, French Patent WO 99/34901, 1999.
[11] Y. Li, J.N. Armor, Appl. Catal. B Environ. 1 (1992) L21.
[12] G. Centi, A. Galli, B. Montanari, S. Perathoner, A. Vaccari, Catal. To-
day 35 (1997) 113.
[13] G. Centi, S. Perathoner, F. Vazzana, M. Marella, M. Tomaselli, M. Man-
tegazza, Adv. Environ. Res. 4 (2000) 325.
[14] J.A.Z. Pieterse, G. Mul, I. Melian-Cabrera, R.W. van den Brink, Catal.
Lett. 99 (2005) 41.
[15] J.A.Z. Pieterse, S. Booneveld, G. Mul, R.W. van den Brink, Stud. Surf.
Sci. Catal. 157 (2005) 1915.
[16] V. Boissel, S. Tahir, C.A. Koh, Appl. Catal. B 64 (2006) 234.
[17] G.D. Pirngruber, J.A.Z. Pieterse, J. Catal. 237 (2006) 237.
[18] N. Weiher, E. Bus, B. Gorzolnik, M. Moller, R. Prins, J.A. van Bokhoven,
J. Synch. Rad. 12 (2005) 675.
The observed synergy between Fe and Ru in bimetallic
FeRu-FER catalysts for N2O decomposition in the presence of
NO results from two effects. The first of these effects is a cat-
alytic cooperation between Fe and Ru. The Fe component is
very active for the oxidation of NO by N2O, that is, N2O +
NO → NO2 + N2. The subsequent O2 formation via N2O +
NO2 → O2 + N2 + NO is rate-limiting. The Ru component
offers a second channel for O2 formation, which is rapid: The
reaction NO2 ꢀ NO +(1/2)O2 is close to equilibrium over
Ru. In the forward direction, which is thermodynamically fa-
vored at high temperature, it closes the catalytic cycle of N2O
decomposition. In theory, this mechanism should also work for
other transition or noble metals that are highly active in the de-
composition of NO2 to NO and O2.
[19] M.S. Kumar, M. Schwidder, W. Grünert, A. Brückner, J. Catal. 227 (2004)
384.
[20] J. Perez-Ramirez, J.C. Groen, A. Brückner, M.S. Kumar, U. Bentrup, M.N.
Debbagh, L.A. Villaescusa, J. Catal. 232 (2005) 318.
[21] R.W. van den Brink, S. Booneveld, M.J.F.M. Verhaak, Catal. Today 75
(2002) 227.
[22] A.K. Goel, G. Skorinko, F.H. Pollak, Phys. Rev. B 24 (1981) 7342.
[23] A.P. Alivisatos, Science 271 (1996) 933.
[24] X. Wang, H.Y. Chen, W.M.H. Sachtler, J. Catal. 197 (2001) 281.
[25] K. Hadjiivanov, B. Tsyntsarski, T. Nikolova, Phys. Chem. Chem. Phys. 1
(1999) 4521.
[26] B.R. Wood, J.A. Reimer, A.T. Bell, M.T. Janicke, K.C. Ott, J. Catal. 225
(2004) 300.
[27] K. Hadjiivanov, J. Saussey, J.L. Freysz, J.C. Lavalley, Catal. Lett. 52
(1998) 103.
[28] L.J. Lobree, I.C. Hwang, J.A. Reimer, A.T. Bell, Catal. Lett. 63 (1999)
233.
A second effect that may contribute to the synergy between
Fe and Ru is based on the adsorption properties of the bimetal-
lic catalyst. In a bimetallic sample with high iron loading, NO
and NO2-derived surface species adsorb preferentially on the
Fe component, thereby significantly reducing the inhibiting ef-
fect of NO on the Ru component. The effect is not observed
in samples with low iron loadings. Unfortunately, our data do
not allow us to quantify the contribution of the adsorption ef-
fect to the synergy between Fe and Ru; this would require a
more detailed kinetic study. The in situ XANES data show that
the presence of Ru does not increase the fraction of Fe2+ under
reaction conditions; there is no positive influence of Ru on the
redox properties of Fe.
[29] H.Y. Chen, T. Voskoboinikov, W.M.H. Sachtler, J. Catal. 180 (1998) 171.
[30] Y. Wang, K. Jacobi, G. Ertl, J. Phys. Chem. B 107 (2003) 13918.
[31] G. Mul, J. Perez-Ramirez, F. Kapteijn, J.A. Moulijn, Catal. Lett. 77 (2001)
7.
[32] G.D. Pirngruber, P.K. Roy, Catal. Today 110 (2005) 199.
[33] G.D. Pirngruber, M. Luechinger, P.K. Roy, A. Cecchetto, P. Smirniotis,
J. Catal. 224 (2004) 429.
[34] R. Brosius, D. Habermacher, J.A. Martens, L. Vradman, M. Herskowitz,
L. Capek, Z. Sobalik, J. Dedecek, B. Wichterlova, V. Tokarova, O. Gon-
siorova, Top. Catal. 30–31 (2004) 333.
Note added in proof
While this manuscript was being processed, a synergy be-
tween Fe and Pt in N2O decomposition was reported by
D. Kaucky, K. Jisa, A. Vondrova, J. Novakova, Z. Sobalik,
J. Catal. 242 (2006) 2.
Acknowledgments
[35] L. Capek, V. Kreibich, J. Dedecek, T. Grygar, B. Wichterlova, Z. Soba-
lik, J.A. Martens, R. Brosius, V. Tokarova, Microporous Mesoporous
Mater. 80 (2005) 279.
The authors thank Can Pinarci (ETH Zurich) for helping
with the IR measurements and Martin Kuba and Dr. Pavel