Selective catalytic reduction of NO2 with urea in nanocrystalline NaY zeolite

Article abstract of DOI:10.1016/j.jcat.2005.06.025

In this study, the selective catalytic reduction (SCR) of NO₂ with urea in nanocrystalline NaY zeolite was investigated with in situ transmission Fourier transform infrared (FTIR) spectroscopy and solid-state nuclear magnetic resonance spectroscopy. At T=473 K, the reaction rate for urea-SCR of NO₂ in nanocrystalline NaY zeolite was significantly greater than that in commercial NaY zeolite with a larger crystal size. In addition, a dramatic decrease in the concentration of undesirable surface species, including biuret and cyanuric acid, was observed in nanocrystalline NaY compared with commercial NaY after urea-SCR of NO₂ at T=473 K. The increased reactivity for urea-SCR of NO₂ was attributed to silanol groups and extra-framework aluminum species located on the external surface of nanocrystalline NaY. Specifically, NO_x storage as nitrate and nitrite on the internal zeolite surface was coupled to reactive deNO_x sites on the external surface. Isotopic labeling combined with IR analysis suggest that NN bond formation involved both an N-atom originating from NO₂ and an N-atom originating from urea. This is the first clear example demonstrating that the increased external surface area (up to ～40% of total surface area) of nanocrystalline zeolites can be used as a reactive surface with unique active sites for catalysis.

Full text of DOI:10.1016/j.jcat.2005.06.025

Journal of Catalysis 234 (2005) 401–413
www.elsevier.com/locate/jcat
Selective catalytic reduction of NO with urea in nanocrystalline
2
NaY zeolite
∗
∗
Gonghu Li, Conrad A. Jones, Vicki H. Grassian , Sarah C. Larsen
Department of Chemistry, The University of Iowa, Iowa City, IA 52242, USA
Received 24 March 2005; revised 23 May 2005; accepted 2 June 2005
Available online 10 August 2005
Abstract
In this study, the selective catalytic reduction (SCR) of NO with urea in nanocrystalline NaY zeolite was investigated with in situ
2
transmission Fourier transform infrared (FTIR) spectroscopy and solid-state nuclear magnetic resonance spectroscopy. At T = 473 K, the
reaction rate for urea-SCR of NO in nanocrystalline NaY zeolite was significantly greater than that in commercial NaY zeolite with a larger
2
crystal size. In addition, a dramatic decrease in the concentration of undesirable surface species, including biuret and cyanuric acid, was
observed in nanocrystalline NaY compared with commercial NaY after urea-SCR of NO at T = 473 K. The increased reactivity for urea-
2
SCR of NO was attributed to silanol groups and extra-framework aluminum species located on the external surface of nanocrystalline NaY.
2
Specifically, NOx storage as nitrate and nitrite on the internal zeolite surface was coupled to reactive deNOx sites on the external surface.
Isotopic labeling combined with IR analysis suggest that N–N bond formation involved both an N-atom originating from NO and an N-atom
2
originating from urea. This is the first clear example demonstrating that the increased external surface area (up to ∼40% of total surface area)
of nanocrystalline zeolites can be used as a reactive surface with unique active sites for catalysis.

2005 Elsevier Inc. All rights reserved.
Keywords: Selective catalytic reduction; NO ; Y Zeolite; Urea; Isotopic labeling; Nanocrystals
2
1
. Introduction
thermally decomposes in two steps (reactions (1) and (2)) to
form ammonia and carbon dioxide [12–14]:
Selective catalytic reduction (SCR) using ammonia or hy-
(
NH2)2CO → NH3 + HNCO,
(1)
(2)
drocarbons is a promising technology for post-combustion
treatment of NOx (NO and NO2) [1]. NH3-SCR has been
developed and used worldwide for the control of NOx emis-
sions in fuel combustion from stationary sources due to its
efficiency, selectivity, and economics [2]. However, NH3
is not a practical reducing agent for NOx emissions from
mobile sources, because of its toxicity and difficulties in
storage, transportation, and handling [1]. Much interest has
focused on using urea as a safer source of ammonia in au-
tomotive applications [3–7]. Currently, a solution of 32.5%
urea in water is the preferred choice among different precur-
sors for ammonia [8–11]. It is generally accepted that urea
HNCO + H2O → NH3 + CO2.
Transition metal-containing zeolites have been exten-
sively studied as SCR catalysts [15–20]. Several recent stud-
ies have also used alkali and/or alkaline earth-substituted Y
zeolites as selective oxidation catalysts [21–23] and deNOx
catalysts [24–29]. In the absence of transition metals, zeo-
lite Y potentially offers novel SCR pathways different from
those in transition metal containing catalysts. As recently
discussed, nanocrystalline Y zeolites and metal oxides with
particle sizes less than 100 nm may be particularly useful
in environmental applications because of their small crystal
sizes and large internal and external surface areas [30–33].
The influence of intracrystalline diffusion on SCR reaction
rate has been previously investigated for transition metal-
exchanged zeolites with different crystal sizes [34–38]. Most
*
Corresponding authors. Fax: +1-319-335-1270.
E-mail addresses: vicki-grassian@uiowa.edu (V.H. Grassian),
sarah-larsen@uiowa.edu (S.C. Larsen).
0021-9517/$ – see front matter  2005 Elsevier Inc. All rights reserved.
doi:10.1016/j.jcat.2005.06.025

402
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
applicable to the current study is recent work on propylene-
SCR of NO2 over nanocrystalline NaY zeolite with a crystal
size of 23 nm [32]. At T = 473 K, significantly higher NOx
conversion and higher selectivity toward N–N bond forma-
tion was observed in nanocrystalline NaY zeolite compared
with commercial NaY zeolite with a larger crystal size [32].
It was also observed that nanocrystalline NaY zeolite was
more sensitive to the presence of adsorbed water than com-
mercial NaY zeolite in propylene-SCR of NO2 at T = 473 K
NaY sample was calcined at T = 773 K under oxygen flow
for 16 h to remove organic templates.
The synthesized nanocrystalline NaY zeolite and the
commercial NaY zeolite (purchased from Aldrich) were
characterized by several different techniques, including
scanning electron microscopy (SEM), X-ray diffraction, ni-
trogen adsorption isotherms, and solid-state nuclear mag-
netic resonance (NMR) [30]. The average crystal size of
nanocrystalline NaY zeolite was determined to be 23 nm
[30]. The commercial NaY zeolite is composed of large
intergrown NaY crystals with diameters of ∼1 µm. The ex-
ternal, internal, and total surface areas of nanocrystalline
[
(
32]. The silanol groups and extra-framework aluminum
EFAL) species with Bronsted acidity located on the external
surface of nanocrystalline NaY zeolite were found to ac-
count for the observed differences in activity (see Fig. S1
in Supplementary information [32]).
In the study reported here, in situ transmission FTIR spec-
troscopy was used to investigate urea-SCR of NO2 over
nanocrystalline NaY zeolite at relatively low temperature
2
NaY zeolite were determined to be 178, 406, and 584 m /g,
respectively. The total surface area of commercial NaY zeo-
2
lite was 477 m /g, and, as discussed in previous work [30],
the external surface area of commercial NaY zeolite is esti-
2
mated from the crystal size to be ∼4 m /g. The Si/Al ratios
(
T = 473 K). This work is motivated by the fact that the
of nanocrystalline NaY and commercial NaY zeolite were
determined to be 1.8 and 2.0, respectively [30].
silanol groups and EFAL sites on the external surface of
nanocrystalline NaY may be effective surface sites for the
decomposition of urea to NH3 and the subsequent SCR of
NO2 [2,8]. The performance of nanocrystalline NaY zeo-
lite in urea-SCR reactions is compared to commercial NaY
zeolite with a larger crystal size to explore the potential ad-
vantages of the small crystal size of nanocrystalline NaY
zeolite and to determine whether the chemical and physical
properties of nanocrystalline NaY are important in facilitat-
ing urea-SCR reactions.
Aluminum isopropoxide, TEOS, and sodium hydroxide
were purchased from Alfa Aesar. Commercial NaY zeo-
lite and TMAOH solution (20 wt%) were purchased from
Aldrich. Solid urea (>99.5%) was purchased from Acros
Organics. 15N- and C-labeled urea were purchased from
Cambridge Isotopes. Research-grade purity oxygen and car-
bon dioxide were purchased from Air Products. Anhydrous
ammonia, research-grade purity nitrogen dioxide, nitric ox-
ide, and nitrous oxide were purchased from Matheson. All
gases were used without further purification.
13
2
. Experimental section
2.2. Transmission FTIR spectroscopy
2
.1. Zeolite sample preparation and materials
Nanocrystalline NaY zeolite was synthesized using clear
Four mg urea in aqueous solution (∼0.3 mol) and 7 mg
NaY zeolite were mixed and sonicated for 30 min at room
temperature. The resulting hydrosol was coated onto a 3 ×
2 cm tungsten grid held in place by nickel jaws. The nickel
jaws were attached to copper leads so that the sample could
be resistively heated. A thermocouple wire attached to the
tungsten grid was used to measure the temperature of the
sample. The tungsten grid with zeolite sample was placed
inside a stainless steel cube, which had two BaF2 windows
for FTIR measurements and was connected to a vacuum/gas
handling system [23].
solutions according to the general method reported by
Creaser et al. [39] and further modified by Song et al. [30].
Sodium hydroxide, aluminum isopropoxide, tetramethylam-
monium hydroxide (TMAOH), and distilled water were
mixed and stirred until the resulting mixture became a clear
solution. Then tetraethylorthosilicate (TEOS) was added to
the clear solution. The clear solution mixture was stirred
overnight to ensure complete hydrolysis of the aluminum
and silicon sources. The final clear synthesis gel had the fol-
lowing composition: 0.07 Na:2.4 TMAOH:1.0 Al:2.0 Si:132
H2O:3.0 i-PrOH:8.0 EtOH. The latter two alcohols were
formed from the hydrolysis of aluminum isopropoxide and
TEOS, respectively. To optimize the formation of small NaY
crystals, the sodium content was intentionally set too low rel-
ative to the aluminum content in the synthesis gel. The clear
solution was transferred into a 500-mL flask equipped with
an air-cooled condenser and was heated to T = 368 K in an
oil bath for 84 h with stirring. Nanocrystalline Y powders
were recovered from the milky, colloidal suspension of NaY
after two cycles of centrifugation, washed with deionized
water, and dried at T = 393 K in air. The nanocrystalline
The stainless steel FTIR cell was held in place by a lin-
ear translator inside the sample compartment of a Mattson
Galaxy 6000 FTIR spectrometer equipped with a narrow-
band MCT detector. The linear translator allowed each half
of the sample grid to be moved into the FTIR beam path,
permitting detection of gas phase and adsorbed species in ze-
olites under identical reaction conditions. Each absorbance
spectrum shown was obtained by referencing 64 scans at
−1
an instrument resolution of 4 cm to the appropriate back-
ground of the zeolite or the blank grid, unless noted other-
wise.
The NaY zeolite sample with adsorbed urea was dried in
air and evacuated at room temperature overnight to remove

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
403
weakly adsorbed water. Reactant gases (NO2 and O2) were
loaded into the NaY zeolite through the gas handling sys-
tem. Two absolute pressure transducers were used to mon-
itor the pressure. The extinction coefficients of individual
gases were calibrated using the characteristic FTIR absorp-
tion band and measuring the pressure using an absolute pres-
sure transducer. Typically, the NaY zeolite was equilibrated
with gases before a spectrum was recorded. In SCR reac-
tions, NO2 was first introduced into the FTIR cell at room
temperature. After 30 min to allow adsorption equilibrium,
an excess amount of oxygen was added into the FTIR cell.
Time course experiments were conducted by automatically
recording FTIR spectra of the gas phase every 1 min.
2
.3. Solid-state magic-angle spinning NMR experiments
Approximately 200 mg of zeolite were impregnated with
13
15
an aqueous solution of labeled urea ( C and N). Samples
with adsorbed urea were allowed to dry at room temper-
ature and were placed into Pyrex sample tubes that were
degassed on a vacuum rack. To prepare samples containing
labeled urea and nitric oxide and oxygen, the nitric oxide
and oxygen were introduced to the zeolite sample while still
attached to the vacuum rack. Then the Pyrex sample tube
was sealed while the sample was immersed in liquid nitro-
gen. The sealed samples were heated for 1 h at the desired
temperatures. Samples were allowed to cool to room tem-
perature for the NMR measurements.
Fig. 1. FTIR spectrum of commercial NaY with adsorbed urea at
T = 298 K. The zeolite sample with adsorbed urea was dried in air and
evacuated at T = 298 K overnight. The FT-IR spectrum of a clean commer-
cial NaY sample is also plotted (dotted line). The blank grid is used as a
reference. All spectra were recorded at T = 298 K.
dried in air and evacuated at room temperature overnight to
remove weakly adsorbed water. The FTIR spectrum of com-
mercial NaY zeolite with adsorbed urea is shown in Fig. 1.
The FTIR spectrum of a clean commercial NaY sample is
also plotted in Fig. 1 (dotted line) for comparison.
Magic-angle spinning (MAS) NMR spectra were ob-
tained using a 7-Tesla wide bore Bruker cryomagnet and
a TecMag Discovery console. The Larmour frequencies for
Adsorbed urea in NaY zeolite can be identified by several
−1
absorption bands at 1456, 1584, 1610, 1650, and 1712 cm
in the FTIR spectrum as shown in Fig. 1. The absorption
13
15
C and N were 75.470 and 30.425 MHz, respectively.
−1
band at 1456 cm is assigned to the asymmetric N–C–N
stretching mode of adsorbed urea [14]. The absorptions at
A Chemagnetics double-channel 7.5-mm pencil MAS probe
was used to spin rotors containing the sealed samples. The
spinning speed was approximately 4.0 kHz. Single-pulse di-
−
1
1584 and 1610 cm are both due to the NH2 bending mode
1
3
15
−1
rect excitation was used for C and N. The π/2 pulse
of adsorbed urea [14]. The bands at 1650 and 1712 cm
13
width was 6.0 µs for C; the π/2 pulse width was 4.5 µs
for N. Recycle delay was 2 s, and a line broadening of
0 Hz was used for both C and N. There were 4000 scans
for the C NMR and 25,000 scans for the N MAS NMR.
Solid adamantane mixed with KBr (38.5 ppm) and saturated
(both C=O stretching mode) can be assigned to urea ad-
sorbed at different sites or in different coordination envi-
ronments in NaY zeolite [14]. Some additional bands in
15
13
15
5
13
15
−1
the spectral region of 3100–3800 cm due to O–H/N–H
stretching are also seen (not shown). The general features
of the FTIR spectrum of nanocrystalline NaY zeolite with
adsorbed urea (not shown) are similar to those of commer-
cial NaY zeolite, except for increased spectral broadening
of the peaks in nanocrystalline NaY. This broadening has
been previously attributed to increased site heterogeneity for
nanocrystalline zeolite samples [32].
15
15
−
NH4 NO3 solution (0 ppm for NO3 ) were used as ex-
13
15
ternal chemical shift standards for C and N, respectively.
All of the chemical shifts for 13C and 15N are reported rel-
13
15
ative to TMS and CH3NO2, respectively. All C and
spectra were obtained at room temperature.
N
The NaY sample with adsorbed urea was then exposed
to 0.5 Torr pressure of NO2 at room temperature and equi-
librated for 30 min. The changes in the FTIR spectra of
the NaY sample upon exposure and equlibrium with NO2
are clearly seen in the difference spectra shown in Fig. 2.
Each difference spectrum was obtained by subtracting the
spectrum of NaY zeolite before adsorbing NO2 from the
corresponding spectrum after the adsorption equilibrium. As
shown in the difference FTIR spectrum of commercial NaY
3
. Results
.1. Adsorption of NO2 in NaY zeolite with adsorbed urea
3
at T = 298 K
As discussed earlier, an aqueous urea solution and the
NaY zeolite powder were mixed at room temperature. Be-
fore adsorbing NO2, the NaY zeolite sample with area was

404
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
Fig. 2. Difference FTIR spectra of nanocrystalline NaY with adsorbed urea
and commercial NaY zeolite with adsorbed urea following adsorption of
Fig. 3. The FTIR spectra of the gas phase upon adsorption of NO in
2
15
(
a) commercial NaY zeolite with adsorbed N-urea, (b) nanocrystalline
0
.5 Torr NO at T = 298 K. The spectra use the corresponding NaY zeolite
2
15
NaY zeolite with adsorbed N-urea, (c) commercial NaY zeolite with ad-
sorbed unlabeled urea, and (d) nanocrystalline NaY zeolite with adsorbed
unlabeled urea. All spectra were recorded at T = 298 K after adsorption
with adsorbed urea prior to NO adsorption as a background. In addition,
2
gas-phase absorptions have been subtracted from each of the FTIR spectra.
−
1
equilibrium. The NO signal at ∼1875 cm has been increased by a fac-
tor of 5 in the spectrum labeled (c) and is representative of the NO signal
intensity in all of the spectra in this figure.
zeolite (Fig. 2), several absorption bands appeared at 1272,
−
1
1
372, 1514, and 1770 cm . The absorption bands at 1272
−1
and 1372 cm are attributed to the formation of nitrite and
nitrate on Na sites, respectively [24,31,32,40,41]. The ap-
+
analysis indicate that EFAL sites are located on the external
surface of nanocrystalline NaY zeolite and are not present to
any great extent in the commercial NaY sample because of
its low external surface area [30–32].
−1
pearance of two absorption peaks at 1514 cm (νCN +δNH)
−
1
and 1770 cm (νCO) indicate the formation of a surface
species with a linear imido structure [e.g., –(C=O)–NH–
Unlabeled urea and 15N-labeled urea were used in SCR
reactions over both nanocrystalline NaY and commercial
NaY zeolite. Compared with the FTIR spectrum of NaY
zeolite with unlabeled urea, several differences were ob-
served when 15N-urea solution was mixed with the zeo-
lite sample. For example, the absorption band for adsorbed
(
C=O)–] [12]. A simultaneous loss in intensity of the bands
−
1
around 1456 and 1650 cm was observed, indicating that
adsorbed urea reacted with NO2 (or surface species gener-
ated in NO2 adsorption).
The difference FTIR spectrum of nanocrystalline NaY
with adsorbed urea after being exposed to NO2 is shown
in Fig. 2 (bottom spectrum). Similar to the commercial
1
5
−1
H NCO shifted from 2276 to 2258 cm when nanocrys-
−
1
15
NaY zeolite, the formation of nitrite (1272 cm ) and dis-
talline NaY with adsorbed N-urea was exposed to NO2
−1
appearance of adsorbed urea (1456 and 1650 cm ) were
observed for nanocrystalline NaY zeolite. However, several
differences in the spectra of nanocrystalline NaY zeolite and
commercial NaY zeolite were observed. For example, for-
mation of the surface species with the linear imido structure
at room temperature (not shown). A similar isotopic shift
was also observed for gas-phase HNCO. The FTIR spectra
of the gas phase after NO2 adsorption in commercial and
15
nanocrystalline NaY zeolite with unlabeled urea and N-
−
1
urea are shown in Fig. 3. Gas-phase NO2 (1616 cm ), NO
−1
−1
−1
−1
(
1514 and 1770 cm ) occurred to a much lesser extent
(1875 cm ), HNCO (2268 cm ), and CO2 (2349 cm )
in nanocrystalline NaY (Fig. 2). Furthermore, NO2 adsorp-
tion in nanocrystalline NaY zeolite resulted in the forma-
tion of nitrate adsorbed on external (vide infra) EFAL sites
were observed when NaY zeolite with adsorbed urea was ex-
posed to NO2 at room temperature [27,31,42]. For gas-phase
−1
HNCO, an isotopic shift of 8 cm was observed after sub-
−1
+
15
14
(
1336 cm ) [31,32], as well as nitrate on internal Na
stituting N for N [27]. The complete list of isotopic shifts
for nitrogen-containing gas phase and adsorbed species is
given in Table 1 and is discussed in detail later in this paper.
In the spectra shown in Fig. 3, the absorption band for gas-
phase NO2 over nanocrystalline NaY zeolite is less intense
than that over commercial NaY, indicating that more NO2
was adsorbed/reacted with nanocrystalline NaY zeolite rela-
tive to commercial NaY. This finding is consistent with our
−1
sites (1372 cm ). Finally, the formation of isocyanic acid
HNCO) adsorbed on EFAL sites is observed in nanocrys-
talline NaY zeolite, as indicated by a broad absorption band
(
−1
around 2276 cm [27,31,32]. In NO2 adsorption at room
temperature, the formation of surface-adsorbed isocyanic
acid is not appreciable in commercial NaY zeolite (Fig. 2).
Previous studies of pyridine adsorption coupled with IR

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
405
Table 1
List of vibrational frequencies of nitrogen-containing gas-phase and adsorbed species which showed an isotopic shift upon substitution of N for N in the
adsorbed urea ((NH ) CO) precursor
15
14
2
2
1
4
15
a
14
15
Product
HNCO
Assignment
ν( N)
ν( N)
ꢀν
ν( N)/ν( N)
Gas phase
ν
2268
2224
1285
965
2260
2201
1270
960
8
23
15
5
1.004
1.010
1.012
1.005
NCO
νNN
N O
2
N O
ν
2
NO
δNH3
NH
3
Adsorbed in
nanocrystalline
NaY
HNCO
OCN
ν
ν
ν
ν
2276
2169
1527
1271
2258
2153
1522
1266
18
16
5
1.008
1.007
1.003
1.004
NCO
NC
−
(NH ) CO
2
2
CNH
CNH
(
NH ) CO
5
2
2
Adsorbed in
commercial
NaY
HNCO
OCN
ν
ν
ν
ν
2282
2175
1759
1735
2250
2159
1751
1727
32
16
8
1.014
1.007
1.005
1.005
NCO
NC
−
(NH CO) NH
2
2
CO
(
HNCO)
8
3
CO
a
ν = ν(14N) − ν(15N).
ꢀ
15
Fig. 4. FTIR spectra of the gas phase as a function of reaction time in N-urea SCR of NO over nanocrystalline NaY zeolite in the presence of oxygen. The
2
spectra shown were recorded at T = 473 K and at t (reaction time) = 0, 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, 19, 21, 23, 25, 27, 30 min,
respectively. The blank grid is used as a reference.
previous observation that nanocrystalline NaY zeolite has a
greater adsorption capacity for NO2 than commercial NaY
zeolite [30,32].
nanocrystalline NaY zeolite at T = 473 K. In the beginning
15
of the SCR reaction (t = 0), gas-phase NO , NO, H NCO,
2
CO , and a small amount of H O were observed in the
2
2
1
5
FTIR spectrum. At t = 1 min, more H2O and H NCO ap-
3
.2. Urea-SCR of NO2 in NaY zeolite at T = 473 K: Gas
peared in the gas phase, as shown by the spectra (Fig. 4).
15
phase products
Gas-phase H NCO decayed quickly after t = 2 min. Sev-
eral other changes in gas phase composition can be seen in
these time course experiments, including the rapid produc-
An excess amount of O2 was introduced into the FTIR
cell after the NaY zeolite sample was equilibrated with NO2
for 30 min at room temperature. Urea-SCR of NO2 was then
carried out at T = 473 K. A control experiment was done by
heating the commercial NaY zeolite with adsorbed urea at
T = 473 K in the presence of oxygen.
tion of 15NH3 (927, 960, 1623, and 3331 cm ) and CO2
−1
15
at t = 2 min. Gas-phase N-urea appeared and is iden-
tified by several characteristic absorptions, including three
−1
broad bands between 3100 and 3520 cm and several ab-
sorptions between 1290 and 1720 cm in the FTIR spectra
shown in Fig. 4. Furthermore, gas-phase NO2 disappeared
−1
Fig. 4 shows the FTIR spectra of gas-phase species
formed during the first 30 min of 15N-urea SCR over
15
−1
quickly relative to NO. NNO (2201 and 1270 cm ) ap-

406
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
Fig. 5. FTIR spectra of the gas phase as a function of reaction time in urea-SCR of NO over commercial NaY zeolite in the presence of oxygen. The spectra
2
shown were recorded at T = 473 K and at t (reaction time) = 0, 1, 2, 3, 4, 5, 6, 7, 8, 10, 12, 14, 16, 18, 20, 25, 30, 35, 40, 50, 60, 70, 80, 90, 120 min,
respectively. The blank grid is used as a reference.
peared in these spectra as the only isomer of N2O (Table 1)
the presence of large amounts of HNCO during the thermal
decomposition of urea at T = 473 K in commercial NaY ze-
olite (data not shown).
15
15
[
43]. Other isomers of N2O, such as N NO and N2O,
−1
have different absorption frequencies (2178/1280 cm for
N NO and 2155/1265 cm for N2O) and were not de-
15
−1
15
The time course concentrations of the gases in the IR
cell can be determined from the integrated absorbance of the
corresponding IR absorptions using calibrated extinction co-
efficients, as described in previous work [31,32]. The time
course concentrations of gas-phase NH3, CO2, and H2O in
urea-SCR of NO2 are shown in Fig. 6. The time course con-
centrations for the urea thermal decomposition over com-
mercial NaY zeolite are also included in Fig. 6. It should
be noted that the concentration of gas-phase H2O rapidly in-
creased to a maximum at t = 2 min in urea-SCR of NO2 over
nanocrystalline NaY zeolite (Fig. 6c). The concentration of
gas-phase H2O then decreased sharply until a minimum was
reached at approximately t = 15 min and started increasing
again thereafter. Most likely, strongly adsorbed water that re-
mained in nanocrystalline NaY zeolite was released initially
upon heating to T = 473 K [31], followed by the consump-
tion and production of H2O in SCR reactions [12].
The time course concentrations of NO and N2O in the
gas phase are shown in Fig. 7. In urea-SCR of NO2 over
nanocrystalline NaY zeolite, the concentration of NO de-
creased monotonically and gas-phase NO disappeared com-
pletely around t = 40 min (Fig. 7a). When commercial NaY
zeolite was used as the SCR catalyst, the NO concentration
initially increased until it reached a maximum level and de-
creased slowly thereafter. The time course concentration of
N2O demonstrates that N2O was produced more quickly and
in much higher concentration in nanocrystalline NaY zeolite
than in commercial NaY zeolite (Fig. 7b).
15
tected in our experiments [43]. The formation of NNO in
the gas phase indicates that N–N bond formation at 473 K
involved a 15N-atom originating from urea and a N-atom
originating from NO2 and that the N=O bond in NO2 re-
mained intact during the formation of 15NNO. There was
no significant change in the gas phase after 30 min of SCR
reactions over nanocrystalline NaY zeolite.
14
Compared with the SCR reaction in nanocrystalline NaY
zeolite, the urea-SCR of NO2 in commercial NaY zeolite
was much slower. The FTIR spectra of the gas phase species
present during the 2 h of urea-SCR over commercial NaY
zeolite are shown in Fig. 5. In urea-SCR using commercial
NaY catalyst, NO2 disappeared quickly, but NO persisted
in the gas phase during the 2 h reaction period. The con-
centration of HNCO reached a maximum at t = 5 min, and
HNCO remained detectable in the IR cell for at least 60 min
during the SCR reaction process (Fig. 5). There was an in-
duction period (∼3 min) for the production of NH3 and
CO2 that did not appear in the gas phase with significant
concentrations until t = 4 min (Fig. 5). N2O appeared as
one of the final products in the gas-phase [31,32]. A gas
−1
phase species with an absorption around ∼2185 cm was
present in the FTIR spectra from t = 4 to 10 min. Because
−1
another absorption band at 1255 cm was also observed
(
(
not labeled), this species is assigned to formonitrile oxide
HCNO) [42,44,45]. Further supporting this assignment is
−1
the fact that HCNO (2190 and 1255 cm ) was produced in

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
407
Table 2
Rates for formation of gas-phase NH , CO and N O and loss of gas-phase
3
2
2
NO in the first 10 min during urea-SCR of NO and thermal decomposition
2
of urea at T = 473 K
Catalyst
Initial rate (µmol/(L min))
NH
CO
NO
N O
2
3
2
Nanocrystalline NaY
Commercial NaY
Commercial NaY
10.3
2.4
3.0
8.9
1.5
1.6
−0.4
−0.2
–
1.1
0.1
–
a
b
a
The initial rate of NO loss during the urea-SCR of NO over commer-
2
cial NaY zeolite was measured from t = 9 to 18 min.
b
Thermal decomposition of urea in the presence of O .
2
Table 3
Gas-phase product distribution after urea-SCR of NO and thermal reaction
2
of urea at T = 473 K for 2 h in 7 mg NaY zeolite. The loadings of NO ,
2
urea and O are 70 ± 5, 170 ± 20, and 550 ± 20 µmol/L in the IR cell with
2
a volume of 0.32 ± 0.01 L, respectively
a
Catalyst
Concentration (µmol/L)
NH
CO
NO
N O
2
3
2
Nanocrystalline NaY
Commercial NaY
Commercial NaY
56
17
53
138
74
69
0
7
–
12
4
–
b
a
The concentration of gas-phase NH was measured at T = 473 K,
Fig. 6. Evolution of gas-phase (a) NH , (b) CO and (c) H O during
3
3
2
2
the concentrations of gas-phase CO , NO and N O were measured at
urea-SCR of NO in nanocrystalline NaY zeolite ("), urea-SCR of NO
2
2
2
2
T = 298 K.
in commercial NaY zeolite (!) and thermal decomposition of urea in com-
mercial NaY zeolite (×) at T = 473 K in the presence of oxygen.
b
Thermal reaction of urea in the absence of NO .
2
of CO2 formation. In the thermal reaction of urea over com-
mercial NaY zeolite, the ratio of NH3 formation rate to CO2
formation rate is close to 2:1, which matches the fact that
1
1
mol of urea thermally decomposes into 2 mol of NH3 and
mol of CO2 (reactions (1) and (2)). Furthermore, the ini-
tial rates for NH3 and CO2 formation are much higher in
nanocrystalline NaY than in commercial NaY. Table 2 also
includes the initial rates of N2O formation and NO loss. It
can be concluded from the data listed in Table 2 that urea
SCR of NO2 occurred more quickly in nanocrystalline NaY
than in commercial NaY zeolite.
Table 3 summarizes the product distribution in the gas
phase after urea-SCR of NO2 and thermal reaction of urea.
The concentrations of NH3 and CO2 after SCR reaction are
much higher in nanocrystalline NaY zeolite than in com-
mercial NaY zeolite. In addition, urea-SCR of NO2 over
nanocrystalline NaY zeolite resulted in 100% NOx (NO2
and NO) conversion and higher selectivity toward N2O for-
mation relative to commercial NaY zeolite. Based on mass
balance for nitrogen, more than 80% of NO2 was converted
into N2 in urea-SCR over nanocrystalline NaY zeolite at
T = 473 K (Table 3).
Fig. 7. Evolution of gas-phase (a) NO and (b) N O during urea-SCR of NO
2
2
in nanocrystalline NaY zeolite (") and in commercial NaY zeolite (!) at
T = 473 K in the presence of oxygen.
The initial rates for formation of gas-phase NH3, CO2,
and N2O and loss of gas-phase NO can be obtained by linear
fitting of the concentrations of individual gas-phase species
in the beginning of the reactions, as listed in Table 2. In urea-
SCR of NO2 over nanocrystalline NaY zeolite, the initial
rate of NH3 formation is slightly higher than the initial rate
3.3. Urea-SCR of NO2 in NaY zeolite at T = 473 K:
adsorbed products
Besides kinetics studies, FTIR analysis of surface ad-
sorbed intermediates or products can improve understanding
of the mechanism behind urea-SCR of NO2 [31,32]. Differ-

408
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
groups are present mainly on the external surface of the
nanocrystalline NaY zeolite [30,32].
The difference FTIR spectrum of nanocrystalline NaY
seen after 15N-urea SCR of NO2 for 30 min (Fig. 8a) is
similar to that seen after urea-SCR for 2 h (Fig. 8b). The
1
5
−1
15
−
formation of adsorbed H NCO (2258 cm ) and OC N
−1
15
(
2153 cm ), as well as the loss of adsorbed N-urea (1522
−1
and 1266 cm ), are apparent in this spectrum. The iso-
topic shifts for these nitrogen-containing species are appar-
ent in the spectrum and are listed in Table 1. In the spectrum
shown in Fig. 8a, a dotted baseline is drawn in the range
−1
1
200–1800 cm , revealing that a small amount of nitrate re-
mained in nanocrystalline NaY zeolite after 30 min of SCR
reaction at T = 473 K.
When commercial NaY zeolite was used as the SCR cat-
alyst, a loss of adsorbed urea was observed, as indicated by
−
1
the negative absorption bands at 1286 cm (νCNH), 1651
and 1712 cm
−1
−1
(both νC=O), and 3436 and 3555 cm
15
(both νNH) in the spectrum of commercial NaY zeolite
after urea-SCR reactions (Fig. 8c). The other CNH vibra-
Fig. 8. Difference FTIR spectra of zeolite samples following (a) N-urea-
SCR of NO in nanocrystalline NaY for 30 min, (b) urea SCR of NO
2
2
−1
in nanocrystalline NaY for 2 h, (c) urea-SCR of NO in commercial NaY
tion that should fall between 1400 and 1600 cm may
be less pronounced in the spectrum, because of the for-
mation of two absorption bands at 1455 and 1495 cm
Fig. 8c). The two absorption bands at 1455 and 1495 cm
2
for 2 h and (d) thermal decomposition of urea in commercial NaY for 2 h
at T = 473 (from top to bottom) in the presence of oxygen. The spectra
were recorded at T = 298 K. All spectra shown use the corresponding NaY
−
1
1
−
(
zeolite with adsorbed urea prior to NO adsorption as a background. In
2
addition, gas-phase absorptions have been subtracted from each of the FTIR
spectra shown.
are assigned to the bending mode of adsorbed ammo-
nium ion and stretching mode of surface species containing
imine groups (C=NH), respectively [12]. In the spectrum
shown in Fig. 8c, the corrected baseline (dotted line) in
ence FTIR spectra of nanocrystalline NaY and commercial
NaY zeolite after urea-SCR of NO2 are shown in Fig. 8. This
figure also shows the difference FTIR spectra of nanocrys-
talline NaY zeolite after 15N-urea SCR of NO2 and commer-
cial NaY zeolite after thermal reaction of urea. The zeolite
samples were cooled back to room temperature before the
spectra were recorded. It was observed that most gas phase
NH3, H2O, and urea disappeared from the gas phase after
the zeolites were cooled back to room temperature, because
of adsorption in the zeolite samples.
−1
the range 1200–1800 cm clearly demonstrates the forma-
−1
tion of adsorbed biuret (1759 cm ) [12] and cyanuric acid
−1
(
1735 cm ) [48] in commercial NaY zeolite after SCR re-
−
−1
actions. Furthermore, surface OCN (2175 cm ) [46] and
−1
nitrate (1376 cm ) were observed after urea-SCR of NO2
for 2 h in commercial NaY zeolite. In the OH/NH stretching
region (Fig. 8c), the absorptions in the spectrum of com-
mercial NaY zeolite indicate the formation of OH group on
+
−1
−1
Na (3695 cm ), surface-adsorbed H2O (3623 cm ), and
−1
NH-containing species (3220, 3380, 3489 cm ), as well as
the loss of adsorbed urea (3436 and 3555 cm ). The FTIR
In urea-SCR of NO2 over nanocrystalline NaY zeolite,
−1
−
the formation of adsorbed HNCO and OCN can be iden-
spectrum of commercial NaY zeolite after thermal reaction
of urea is almost the same as that after urea-SCR of NO2,
except that nitrate, biuret, and cyanuric acid formation is not
observed in the thermal decomposition of urea (Fig. 8d).
The comparison between the difference FTIR spectra of
−1
tified by the absorption bands at 2276 cm [27,31,32] and
169 cm [27,46], respectively, in the FTIR spectra shown
−1
2
in Fig. 8b. The loss of adsorbed urea, characterized by the
−
1
negative absorptions at 1271, 1527, and 1712 cm (νCO)
−1
and 3424 and 3514 cm (both νNH), can also be seen in
the zeolite samples after urea-SCR of NO for 2 h indicates
2
−1
the spectra. The absorption band at 1527 cm is assigned
to the CNH bend-stretch mode, in which the nitrogen and
hydrogen move in opposite directions relative to the carbon,
whereas the absorption band at 1271 cm is assigned to the
CNH open-stretch mode, in which N and H atoms move in
the same direction relative to the carbon [47]. In addition,
that greater nitrate conversion was achieved in nanocrys-
talline NaY zeolite than in commercial NaY zeolite. This
finding is consistent with the fact that urea-SCR of NO2
is much slower in commercial NaY than in nanocrystalline
NaY, as mentioned earlier (see Table 2). Furthermore, un-
desired surface products (biuret and cyanuric acid) were ob-
served in commercial NaY zeolite, but not in nanocrystalline
NaY zeolite, after urea-SCR of NO2.
−1
−1
the recovery of silanol groups (3742 cm ) and OH groups
+
−1
attached to Na (3695 cm ) are evident in the spectrum
of nanocrystalline NaY zeolite shown in Fig. 8b [30–32].
This indicates that adsorption and SCR reactions occurred to
a significant extent on the external surface, because silanol
The assignments of absorption bands in FTIR spectra
shown in Fig. 8 are further supported by isotopic studies
−
(Table 1). The absorption frequencies for HNCO and OCN

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
409
(13
C- and 15N-labeled), NO, and O2 in nanocrystalline NaY.
15
Before heating, the N peak at −301 ppm is present and is
assigned to the adsorbed urea. After heating to T = 423 K,
ammonia is produced, as indicated by the peak at −356 ppm.
After heating to T = 523 K, a peak at −74 ppm is present
and is assigned to N2. After heating to T = 623 K, a sub-
stantial increase in the N2 peak and loss of intensity for the
ammonia peak are seen.
The solid-state NMR spectra were acquired under condi-
tions of thermal equilibrium because of the long time periods
needed for signal averaging. In the presence of oxygen, the
adsorption of NO in zeolite Y occurs in a similar manner as
the adsorption of NO2 [24], because of the oxidation of NO
into NO2 [40]. Thus these NMR results can be readily com-
pared with the FTIR data in this study. The NMR results are
also consistent with the FTIR results, in that urea thermally
decomposes to form NH and CO at elevated temperatures
1
3
15
Fig. 9. C single pulse (a) and N single pulse (b) MAS NMR of la-
3
2
13
15
beled urea ( C and N), NO and O adsorbed in nanocrystalline NaY at
2
in the presence of NO and O . The most important aspect of
2
T = 298, 423, 523 and 623 K (from top to bottom). Spectra were acquired
the NMR results is the observation of N2 formation, which
cannot be observed directly by FTIR spectroscopy.
at room temperature after heating to the desired temperatures. Numbers of
13
15
scans taken for C and N spectra were 20000 and 10000, respectively.
Line broadening = 50 Hz.
4
. Discussion
in nanocrystalline NaY zeolite are different from those in
commercial NaY zeolite. In nanocrystalline NaY, HNCO
−1
−
−1
4.1. Reactions between gas-phase NO and adsorbed urea
in NaY zeolite at T = 298 K
(
2276 cm ) and OCN (2169 cm ) are most likely ad-
2
sorbed on EFAL sites [31,32]. In commercial NaY zeolite,
−
OCN is characterized by the relatively intense absorption
−1
band at 2175 cm in the FTIR spectrum shown in Fig. 8c
and is most likely attached to Na sites [27]. In nanocrys-
The adsorption of NO2 in zeolite Y has been described
in detail in the literature [24,27,31,32,40,41]. In this study,
NO2 reacted with adsorbed water by reaction (3) or (4) in
NaY zeolite, because a considerable amount of water re-
mained in the zeolite after the zeolite samples were evac-
uated at room temperature [32,41]:
+
talline NaY zeolite, the isotopic shift for the CNH vibration
−1
mode of adsorbed urea (5 cm ) is exactly the same as that
for the NH3 deformation mode of gas-phase NH3. This sug-
gests that urea was adsorbed in nanocrystalline NaY zeolite
in such a way that at least one NH2 group remained almost
unperturbed [14]. In commercial NaY zeolite, the isotopic
shift for the C=O stretching mode of adsorbed biuret and
2
NO2 + H2O → HNO3 + HONO,
(3)
(4)
3NO + H O → 2HNO + NO.
2
2
3
−1
cyanuric acid (8 cm ) is the same as that for the NCO
stretching mode of gas-phase HNCO. This may be due to
the configuration exchange
The nitric acid and nitrous acid further deprotonate and ad-
sorb on surface cationic sites to form surface nitrate and
nitrite. This explains the formation of surface nitrate, ni-
trite (Fig. 2), and gas-phase NO after NO2 adsorption in
NaY zeolite (Fig. 3). Nitrate formation on EFAL sites in
nanocrystalline NaY zeolite was also observed (Fig. 2). In
nanocrystalline NaY zeolite, NO2 adsorption occurs in part
–(C=O)–NH– ↔ –C(OH)=N–
in adsorbed biuret and cyanuric acid in commercial NaY ze-
olite.
n+
−
.4. Complementary solid-state NMR ( C and 15N)
1
3
on cationic sites of EFAL species (M –O ), according to
reaction (5) [24,31,32]:
3
spectra of labeled urea, NO, and O2 in nanocrystalline NaY
n+
−
n+
−
M
–O + NO2 → M –NO3 .
(5)
The 13C MAS NMR spectra for urea ( C- and 15N-
labeled), NO, and O2 in nanocrystalline NaY, shown in
Fig. 9a, has a 13C peak at 165 ppm before heating. This peak
is assigned to the carbonyl group of adsorbed urea [16]. Af-
ter heating to T = 423 K, a narrow intense peak is present
at 125 ppm that can be assigned to CO2. After heating to
T = 623 K, there is a substantial increase in peak intensity
for CO2 at 125 ppm, whereas the peak at 165 ppm is no
13
The formation of NO, HNCO, and a small amount of CO2
and H2O was observed when NO2 was adsorbed in NaY ze-
olite with adsorbed urea at room temperature (Figs. 3–5).
Isotopic labeling shows that the N-atom in HNCO originates
from urea (Fig. 4). The loss of adsorbed urea was observed
as well upon the adsorption of NO2 (Fig. 3). Because the de-
composition of urea into HNCO and NH3 is not significant
at room temperature [12], HNCO most likely formed in the
15
longer present. Fig. 9b shows the N MAS NMR of urea

410
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
Scheme 1. Thermal decomposition of urea on silanol (SiOH) and EFAL (AlOx) sites in nanocrystalline NaY zeolite.
reactions between adsorbed urea and HNO3 (reaction (6))
14]. The reaction between the adsorbed urea and HONO
can lead to the formation of CO2 and H2O at room tempera-
ture:
t = 2 min (Figs. 6a and 6b). The initial formation rates of
gas-phase NH3 and CO2 shown in Table 2 correlate well
with the stoichiometry shown by reaction (2), in which 1
mol of HNCO hydrolyzes to form 1 mol of NH3 and 1 mol
of CO2. Most likely, NH3 formed in hydrolysis of HNCO in
the zeolite enters the gas phase immediately at T = 473 K.
Moreover, both the loss of adsorbed urea and regeneration
of surface silanol groups were observed in nanocrystalline
NaY zeolite after urea-SCR of NO2 (Fig. 8). The absorption
frequencies for the urea adsorbed in nanocrystalline NaY
are different from those in commercial NaY, which con-
tains few silanol groups (Table 1) [30,32]. We can conclude
that the adsorption and thermal decomposition of urea in
nanocrystalline NaY zeolite involved the silanol groups on
the external surface. Finally, HNCO produced in thermal de-
composition of urea adsorbed and most likely hydrolyzed on
EFAL sites in nanocrystalline NaY zeolite. This finding is
supported by the fact that HNCO adsorbed on EFAL sites
was observed before (Fig. 2) and after urea-SCR of NO2
[
(
(
NH2)2CO + HNO3 → HNCO + NH4NO3,
NH2)2CO + 2HONO → 2N2 + CO2 + 3H2O.
(6)
(7)
NH4NO3 can thermally decompose into N2O + 2H2O, pro-
viding a minor route for N–N bond formation in urea SCR
of NO2 in NaY zeolite [26,49]. Reaction (7) is used in de-
stroying excess nitrous acid in diazotization [50].
When commercial NaY zeolite was exposed to NO2 at
room temperature, surface species with a linear imido struc-
ture, such as biuret, appeared. This surface species is most
likely generated in the reactions between HNCO and ad-
sorbed urea (Fig. 2). Amide carbonyls of biuret and tri-
−1
uret have characteristic absorption bands ∼1750–1770 cm
[
12]. Biuret can be produced via reaction (8) [50],
(
Fig. 8) in nanocrystalline NaY zeolite.
Scheme 1 describes the adsorption and thermal decom-
NH2–CO–NH2 + HNCO → NH2–CO–NH–CO–NH2. (8)
The above analysis suggests that EFAL species in nano-
crystalline NaY zeolite provide additional surface sites for
NO2 adsorption and its further reactions at room tempera-
ture. Previous studies have shown that EFAL sites on the ex-
ternal surface of nanocrystalline NaY zeolite are important
for the formation of reactive nitrate and for the subsequent
SCR reactions [32]. In this study, EFAL sites in nanocrys-
talline NaY zeolite also provide important reactive surface
sites for HNCO and inhibit the formation of undesired prod-
ucts, such as biuret or triuret.
position of urea on the external surface of nanocrystalline
NaY zeolite before N–N bond formation. Initially, urea ad-
sorbed in nanocrystalline NaY zeolite by hydrogen bonding
to silanol groups on the external surface through one N atom,
leaving the other NH2 almost unperturbed [14]. This is con-
sistent with the isotopic studies showing that the isotopic
shift for the CNH vibration mode is the same as the iso-
topic shift for the NH3 deformation mode of gas-phase NH3
(
Table 1). At elevated temperature (T = 473 K), urea de-
composed into HNCO and NH3. While NH3 remained on
the silanol groups of nanocrystalline NaY zeolite via hydro-
gen bonding, HNCO adsorbed on EFAL sites and continued
to hydrolyze and produce CO2 and additional NH3. The loss
of silanol groups was observed when NH3 was adsorbed
in clean nanocrystalline NaY at room temperature and at
T = 473 K.
4
.2. Insights into urea-SCR of NO2 in nanocrystalline NaY
zeolite at T = 473 K
The FTIR spectra of the gas phase shown in Fig. 4
demonstrate that thermal decomposition of urea occurred
at T = 473 K in nanocrystalline NaY zeolite. As described
previously, gas-phase HNCO was produced initially and de-
cayed rapidly after t = 2 min (Fig. 4). The initial change in
H2O concentration (Fig. 6c) was similar to that of HNCO.
At t = 2 min, the concentration of NH3 in gas phase began
to increase sharply (Fig. 6a). According to reaction (1), the
first step in thermal decomposition of urea produced equal
amounts of HNCO and NH3. This suggests that most NH3
produced in thermal decomposition of urea remained in the
nanocrystalline NaY zeolite. In addition, the concentrations
of NH3 and CO2 in the gas phase increased sharply after
In urea-SCR of NO over nanocrystalline NaY zeolite,
2
isotopic studies indicate that the N–N bond formation in-
volved both an N atom originating from the adsorbed urea
and an N atom originating from gas-phase NO (Fig. 4).
2
NH produced in thermal decomposition of urea is the actual
3
reducing reagent in NO reduction [2,12–14]. It has been
x
suggested that OH groups and surface acid sites on the cat-
alyst surface can coordinate with and further activate NH3
[2,8].
Although the concentration of gas-phase NO decreased
monotonically in urea-SCR of NO2 over nanocrystalline

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
411
NaY zeolite (Fig. 7a), further analysis indicates that both
the formation and consumption of NO occurred. NO is the
major gas-phase product in thermal reactions of nitrate and
nitrite in nanocrystalline NaY zeolite at 473 K [31,32]. In
the present work, the initial rate of N2O formation is much
greater than the initial rate of NO loss in urea-SCR of NO2
over nanocrystalline NaY zeolite (Table 2). The actual initial
rate for N–N bond formation should be even greater, because
talline NaY zeolite than in commercial NaY zeolite, because
of the differences in external surface area [30,32]. (The ex-
ternal surface areas of nanocrystalline NaY and commer-
2
cial NaY zeolite are 178 and ∼4 m /g, respectively [30].)
Silanol groups and EFAL species on the external surface are
responsible for the improved performance of nanocrystalline
NaY zeolite in propylene-SCR of NO2 [32]. In the present
study, urea-SCR of NO2 using nanocrystalline NaY zeolite
as the catalyst led to a significantly greater SCR reaction rate
and less formation of undesired products than seen in urea-
SCR in commercial NaY.
>
80% of the NO2 was converted to N2 in urea-SCR of NO2
over nanocrystalline NaY zeolite. In urea-SCR of NO2 over
commercial NaY zeolite, the concentration of NO initially
increased, then decreased after 9 min (Fig. 7a). A similar
time course concentration of NO was found in propylene-
SCR of NO2 in nanocrystalline NaY [31].
It has been proposed that in NH3-SCR reactions using
transition metal zeolite catalysts, NH3 adsorbed on surface
acid sites reacts directly with gas-phase NO to form N2
and H2O [2,20]. Recent studies by Yeom et al. [25,26] have
demonstrated that NO participates in a number of reactions
in the NOx reduction network over BaNaY zeolite. Am-
monium nitrate is an SCR intermediate that can thermally
decompose to form N2O or can react with NO to produce
ammonium nitrite [25]. In micrometer-sized BaNaY, thermal
decomposition of ammonium nitrate yielded N2O at 585 K
The SCR reactions in nanocrystalline NaY zeolite oc-
curred more quickly than in commercial NaY zeolite (Ta-
ble 2). Consequently, higher NOx conversion (Table 3) and
−
NOx conversion (Fig. 8a) was achieved in nanocrystalline
NaY than in commercial NaY zeolite after urea-SCR for 2 h.
The better SCR reaction rate in urea-SCR over nanocrys-
talline NaY zeolite can be explained by the greater concen-
trations of both silanol groups and EFAL sites than exist in
commercial NaY. First, the silanol groups and EFAL species
have diffusional advantages, because they both are located
on the external surface of the nanocrystalline NaY zeolite.
Furthermore, they are considered chemically distinct surface
sites. A major conclusion of this study and previous studies
is that nitrate adsorbed on EFAL sites on the external surface
of nanocrystalline NaY zeolite is more reactive than nitrate
[
26]. A significant amount of N2O was produced in urea-
SCR of NO2 over nanocrystalline NaY at 473 K (Table 3;
Fig. 7b), suggesting that ammonium nitrate decomposition
may occur at a lower temperature on nanocrystalline NaY.
But ammonium nitrate absorption around 3270, 1460, and
+
in the internal pores adsorbed to Na [32]. The activation
of NH3 by silanol groups is also important for N–N bond
formation [2].
−1
1
350 cm was not evident in the FTIR spectra presented
As shown in Scheme 1, EFAL species on nanocrystalline
NaY zeolite also provide important sites for the adsorption
and fast hydrolysis of HNCO, which is essential for avoid-
ing isocyanic acid emission [15] and inhibiting the formation
of undesired products, including biuret and cyanuric acid.
The kinetic data listed in Table 2 indicate that the hydrolysis
of HNCO occurred at much more rapidly in nanocrystalline
NaY zeolite than in commercial NaY zeolite. This finding is
consistent with the observation that gas-phase HNCO per-
sists for a much longer period (>60 min; Fig. 5) in the
presence of commercial NaY zeolite than in the presence
of nanocrystalline NaY zeolite (<20 min; Fig. 4). The time
course measurements for H2O also indicate that hydrolysis
occurs much faster in nanocrystalline NaY zeolite than in
commercial NaY zeolite (Fig. 6c).
In conclusion, silanol groups and EFAL species are phys-
ically and chemically distinct surface sites present on the
external surface of nanocrystalline NaY zeolite. These sites
are responsible for the improved performance of nanocrys-
talline NaY relative to commercial NaY, because most urea-
SCR reactions occur on the external surface. The role of the
internal surface is to store NOx before SCR reactions on
the external surface. A pictorial representation of this multi-
functional capability of the nanocrystalline zeolite catalyst is
given in Scheme 2. Thus nanocrystalline alkali zeolites can
be classified as new materials for SCR catalysts with poten-
tially significantly better performance.
in this study [51]. Although this finding suggests that no
ammonium nitrate is present, another possibility is that the
ammonium bands were obscured in the FTIR spectrum be-
cause of spectral congestion in these regions due to broad
absorption bands from other adsorbed species.
Ammonium nitrite is an important SCR intermediate that
decomposes efficiently near 373 K to N2 and H2O on Y
zeolites [25–27]. The formation of N2 and loss of NO ob-
served in the present study are consistent with the formation
of ammonium nitrite and its subsequent decomposition to
N2. However, ammonium nitrite bands are not apparent in
the FTIR spectra presented in this study. Therefore, because
ammonium nitrate and ammonium nitrite were not directly
observed in our experiments, a definite mechanistic path-
way cannot be elucidated at this time. Further studies are
needed to investigate ammonium nitrate (nitrite) reactivity
on nanocrystalline NaY.
4
.3. Origin of size effect in urea-SCR of NO2 over NaY
zeolite
Reducing the zeolite crystal size can increase the external
surface area and dramatically alter the diffusional properties
of the zeolite materials [35–37]. Spectroscopic studies have
revealed that silanol groups and EFAL species on the exter-
nal surface are present in greater concentration in nanocrys-

412
G. Li et al. / Journal of Catalysis 234 (2005) 401–413
Scheme 2. (a) SEM image of nanocrystalline NaY zeolite (scale bar = 100 nm), (b) a nanocrystalline zeolite particle with a crystal size of 23 nm, and (c) zeolite
Y supercage. The internal surface of nanocrystalline zeolite provides sites for NOx storage (as nitarte and nitrite) and the minority of the SCR reactions, the
external surface provides sites for additional NOx storage and the majority of the SCR reactions.
4
.4. Urea-SCR versus propylene-SCR of NO2 in
tive to ∼20% selectivity to N2 formation in propylene-SCR
of NO2 over nanocrystalline NaY zeolite [32]. As mentioned
earlier, NOx emissions emerge at high velocity under typi-
cal lean burn conditions [52]. At low temperature, urea-SCR
over nanocrystalline NaY zeolite appears to be an even more
promising deNOx technology than propylene SCR.
nanocrystalline NaY zeolite
An earlier study evaluated nanocrystalline NaY zeolite
in deNOx reactions at T = 473 K using propylene as the
reducing agent and found that nanocrystalline NaY zeolite
was very sensitive to preadsorbed water in propylene-SCR
of NO2 [32]. Compared with urea-SCR, propylene-SCR ap-
pears to be a less realistic deNOx technology under typical
conditions of humid, oxygen-rich exhaust streams [8,52].
Although we did not study the effect of adsorbed wa-
ter on urea-SCR in the present work, water should not be
detrimental to deNOx with urea over nanocrystalline NaY
zeolite. Water is necessary for the hydrolysis of HNCO in
urea-SCR reactions (reaction (2)); moreover, as the above
discussion suggests, at elevated temperatures water does not
inhibit the adsorption of NH3 in nanocrystalline NaY zeolite.
In urea-SCR of NO2 over nanocrystalline NaY zeolite, the
concentration of H2O in the gas phase kept increasing after
the urea-SCR reactions were completed around t = 40 min
5. Conclusions
Nanocrystalline NaY zeolite exhibits enhanced deNO at
x
low temperature (T = 473 K) compared with commercial
NaY zeolite, as demonstrated by the present study on the se-
lective catalytic reduction of NO with urea. Silanol groups
2
and EFAL species on the external surface of nanocrystalline
NaY were found to be responsible for the higher SCR re-
action rate and decreased formation of undesired products,
including biuret and cyanuric acid, relative to commercial
NaY zeolite. Isotopic labeling coupled with FTIR analysis
indicated that N–N bond formation involves both an N atom
(Fig. 6c). Fig. 6a indicates that the NH3 concentration in the
originating from NO and an N atom originating from urea.
2
gas phase further decreased after t = 40 min. Because the
oxidation of NH3 by O2 was not significant under the exper-
imental conditions described in this work [17], NH3 most
likely adsorbed on available silanol groups and/or EFAL
sites in nanocrystalline NaY zeolite in the presence of wa-
ter [30,32].
In nanocrystalline NaY zeolite, kinetic studies suggests that
NH formed in thermal decomposition of urea was activated
3
for SCR reactions by hydrogen bonding with silanol groups.
Nanocrystalline alkali zeolites can be considered as new
catalytic materials that have NO storage capacity in the
x
internal pores and high reactivity on the external surface.
Compared with urea-SCR, a significant amount of nitrate
and carbonaceous deposit remained in nanocrystalline NaY
zeolite after propylene-SCR of NO2 at T = 473 K for 6 h
These features provide multifunctional capabilities beyond
those found for zeolites with larger crystal size and smaller
external surface area. Although for shape selective catalysis,
the reactivity of the external surface may be undesirable, in
[
32]. The carbonaceous deposit in nanocrystalline NaY ze-
olite was found to hinder the adsorption of propylene and
NO2 and to further SCR reactions. In urea-SCR of NO2
over nanocrystalline NaY zeolite, surface nitrate was not
observed after 2 h of SCR reactions. Surface HNCO and
OCN were the only undesired products formed in urea-
SCR over nanocrystalline NaY (Fig. 8). Furthermore, more
than 80% of NO2 was converted into N2 in urea-SCR rela-
reactions such as deNO or hydrocarbon cracking, nanocrys-
talline zeolite catalysts may have greatly enhanced reactivity
x
because of reactions on the external surface. This is the first
example in the literature demonstrating that the increased
external surface area (up to ∼40% of the total surface area)
of nanocrystalline zeolites can be used as a reactive surface
with unique active sites for catalysis.
−

G. Li et al. / Journal of Catalysis 234 (2005) 401–413
413
Acknowledgments
[24] C. Sedlmair, B. Gil, K. Seshan, A. Jentys, J.A. Lercher, Phys. Chem.
Chem. Phys. 5 (2003) 1897.
[
25] Y. Yeom, J. Henao, M. Li, W.M.H. Sachtler, E. Weitz, J. Catal. 231
2005) 181.
26] M. Li, J. Henao, Y. Yeom, E. Weitz, W.M.H. Sachtler, Catal. Lett. 98
2004) 5–9.
[27] Y.H. Yeom, B. Wen, W.M.H. Sachtler, E. Weitz, J. Phys. Chem. B 108
2004) 5386.
The authors thank Dr. Weiguo Song for providing the
nanocrystalline NaY zeolite sample and the SEM image.
Although the research described in this article was funded
wholly or in part by the Environmental Protection Agency
(
[
(
(
(
grant EPA R829600 to S.C.L. and V.H.G.), it has not been
[
[
28] S.J. Schmieg, B.K. Cho, S. Oh, Appl. Catal. B 49 (2004) 113.
29] M. Richter, A. Trunschke, U. Bentrup, K.W. Brzezinka, E. Schreier,
M. Schneider, M.M. Pohl, R. Fricke, J. Catal. 206 (2002) 98.
subjected to the EPA’s required peer and policy review. Thus
this work does not necessarily reflect the views of the EPA,
and no official endorsement should be inferred.
[30] W. Song, G. Li, V.H. Grassian, S.C. Larsen, Environ. Sci. Technol. 39
2005) 1214.
(
[
[
[
31] G. Li, S.C. Larsen, V.H. Grassian, J. Mol. Catal. A 227 (2005) 25.
32] G. Li, S.C. Larsen, V.H. Grassian, Catal. Lett., in press.
33] E. Lukas, S. Decker, A. Khaleel, A. Seitz, S. Fultz, A. Ponce, W. Li,
C. Carnes, K. Klabunde, Chem. Eur. J. 7 (2001) 2505.
References
[
[
[
1] S. Bhattacharyya, R.K. Das, Int. J. Energy Res. 23 (1999) 351.
2] P. Forzatti, Appl. Catal. A 222 (2001) 221.
3] B.K. Gullett, P.W. Groff, M. Lin, J.M. Chen, Air Waste 44 (1994)
[34] M. Ogura, T. Ohsaki, E. Kikuchi, Micropor. Mesopor. Mater. 21
(
1998) 533.
35] A. Shichi, K. Katagi, A. Satsuma, T. Hattori, Appl. Catal. B 24 (2000)
7.
[
1188.
9
[
[
4] M. Koebel, M. Elsener, M. Kleemann, Catal. Today 59 (2000) 335.
5] E. Seker, N. Yasyerli, E. Gulari, C. Lambert, R.H. Hammerle, Appl.
Catal. B 37 (2002) 27.
[36] A. Shichi, A. Satsuma, T. Hattori, Appl. Catal. B 30 (2001) 25.
[37] T. Tabata, H. Ohtsuka, Catal. Lett. 48 (1997) 203.
[38] P. Praserthdam, N. Mongkolsiri, P. Kanchanawanichkun, Catal. Com-
mun. 3 (2002) 191.
[
[
6] M. Koebel, E. Olav Strutz, Ind. Eng. Chem. Res. 42 (2003) 2093.
7] R. van Helden, R. Verbeek, F. Willems, R. van der Welle, Soc. Auto-
mot. Eng. SP-1861 (2004) 25.
[
[
39] Q. Li, D. Creaser, J. Sterte, Chem. Mater. 14 (2002) 1319.
40] J. Szanyi, J. Kwak, R.A. Moline, C.H.F. Peden, Phys. Chem. Chem.
Phys. 5 (2003) 4045.
[8] P.L.T. Gabrielsson, Top. Catal. 28 (2004) 177.
[9] T. Morimune, H. Yamaguchi, Y. Yasukawa, Exp. Therm. Fluid Sci. 18
[41] J. Szanyi, J. Kwak, C.H.F. Peden, J. Phys. Chem. B 108 (2004) 3746.
[42] U. Eickhoff, F. Temps, Phys. Chem. Chem. Phys. 1 (1999) 243.
[43] A. Lapinski, J. Spanget-Larsen, J. Waluk, J. Radziszewski, J. Chem.
Phys. 115 (2001) 1757.
(
1998) 220.
10] D.H.E. Seher, M. Reichelt, S. Wickert, Soc. Automot. Eng. SP-1811
2003) 21.
11] P. Tennison, C. Lambert, M. Levin, Soc. Automot. Eng. SP-1860
2004) 349.
[
[
(
[44] R.E. Baren, M.A. Erickson, J.F. Hershberger, Int. J. Chem. Kinet. 34
2002) 12.
(
(
[
[
12] H.L. Fang, H.F.M. DaCosta, Appl. Catal. B 46 (2003) 17.
13] S. Yim, S. Kim, J. Baik, I. Nam, Y. Mok, J. Lee, B. Cho, S. Oh, Ind.
Eng. Chem. Res. 43 (2004) 4856.
[45] P.W. Beck, K. Feldl, Angew. Chem. Int. Ed. 5 (1966) 722.
[46] M.S. Lowenthal, R.K. Khanna, M.H. Moore, Spectrochim. Acta Part
A 58 (2002) 73.
[
[
[
[
14] M. Larrubia, G. Ramis, G. Busca, Appl. Catal. B 27 (2000) L145.
15] G. Delahay, B. Coq, Catal. Sci. Ser. 3 (2002) 345.
16] C.A. Jones, D. Stec, S.C. Larsen, J. Mol. Catal. A 212 (2004) 329.
17] J. Baik, S. Yim, I. Nam, Y. Mok, J. Lee, B. Cho, S. Oh, Top. Catal. 30
[47] N.B. Colthup, L.H. Daly, S.E. Wiberley, Introduction to Infrared and
Raman Spectroscopy, Chapmann & Hall, London, 1990.
[48] N.W. Cant, D.C. Chambers, A.D. Cowan, I.O.Y. Liu, A. Satsuma, Top.
Catal. 10 (2000) 13.
(
2004) 37.
[49] D.T. Burns, A. Townshend, A.H. Carter, Reactions of the Elements
and Their Compounds (Part A), Inorganic Reaction Chemistry, vol. 2,
Horwood, Chichester, 1981.
[
[
[
[
[
[
18] R.Q. Long, R.T. Yang, J. Am. Chem. Soc. 121 (1999) 5595.
19] R.Q. Long, R.T. Yang, J. Catal. 188 (1999) 332.
20] R.Q. Long, R.T. Yang, J. Catal. 194 (2000) 80.
21] F. Blatter, H. Sun, H. Frei, Catal. Lett. 35 (1995) 1.
22] F. Blatter, H. Sun, S. Vasenkov, H. Frei, Catal. Today 41 (1998) 297.
23] G. Li, M. Xu, S.C. Larsen, V.H. Grassian, J. Mol. Catal. A 194 (2003)
[50] P.A.S. Smith, The Organic Chemistry of Open-Chain Nitrogen Com-
pounds, vol. I, 1965.
[51] T.G. Koch, N.S. Holmes, T.B. Roddis, J.R. Sodeau, J. Chem. Soc.,
Faraday Trans. 92 (1996) 4787.
169.
[52] J.N. Armor, Catal. Today 26 (1995) 99.