Surface Structure Dependence of SO2 Interaction with Ceria

Subscriber access provided by TUFTS UNIV

Article 2

Surface Structure Dependence of SO Interaction with Ceria Nanocrystals with Well-Defined Surface Facets Uma Tumuluri, Meijun Li, Brandon Cook, Bobby G. Sumpter, Sheng Dai, and Zili Wu J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.5b07946 • Publication Date (Web): 02 Dec 2015 Downloaded from http://pubs.acs.org on December 2, 2015

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry C is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Surface Structure Dependence of SO2 Interaction with Ceria Nanocrystals with Well-defined Surface Facets Uma Tumuluri†, Meijun Li§, Brandon G. Cook‡, Bobby Sumpter‡, Sheng Dai†§, and Zili Wu†‡* †

Chemical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee, 37831. §

‡

Department of Chemistry, University of Tennessee, Knoxville, TN 37996-1600.

Center for Nanophase Material Science, Oak Ridge National Laboratory, Oak Ridge, Tennessee, 37831.

Notice: This manuscript has been authored by UT-Battelle, LLC under Contract No. DE-AC0500OR22725 with the U.S. Department of Energy. The United States Government retains and the publisher, by accepting the article for publication, acknowledges that the United States Government retains a non-exclusive, paid-up, irrevocable, world-wide license to publish or reproduce the published form of this manuscript, or allow others to do so, for United States Government purposes. The Department of Energy will provide public access to these results of federally sponsored research in accordance with the DOE Public Access Plan (http://energy.gov/downloads/doe-public-access-plan).

ACS Paragon Plus Environment

1


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 33

ABSTRACT The effect of the surface structure of ceria (CeO2) on the nature, strength and amount of species resulting from SO2 adsorption was studied using in situ IR and Raman spectroscopy coupled with mass spectrometry along with first principle calculations based on density functional theory (DFT). CeO2 nanocrystals with different morphologies: rods (representing defective structure), cubes (100) and octahedra (111), were used to represent different CeO2 surface structures. IR and Raman spectroscopic studies showed that the structure and binding strength of adsorbed species from SO2 interaction depend on the shape of the CeO2 nanocrystals. SO2 adsorbs mainly as surface sulfites and sulfates at room temperature on CeO2 rods, cubes and octahedra that were either oxidatively or reductively pretreated. The formation of sulfites is more evident on CeO2 octahedra while surface sulfates are more prominent on CeO2 rods and cubes. This is explained by the increasing reducibility of the surface oxygen following octahedra < cubes < rods. Bulk sulfites are also formed during SO2 adsorption on reduced CeO2 rods. The formation of surface sulfites and sulfates on CeO2 cubes is in good agreement with our DFT results of SO2 interaction with CeO2 (100) surface. CeO2 rods desorb SO2 at higher temperature than cubes and octahedra nanocrystals, but bulk sulfates are formed in CeO2 rods and cubes after high temperature desorption while only some surface sulfates/sulfites are leftover on octahedra. The difference is rationalized by the fact that CeO2 rods have the highest surface basicity and largest amount of defects among the three nanocrystals, so that they bind and react with SO2 strongly and are the most degraded after SO2 adsorption cycles. The fundamental understanding obtained on the effect of the surface structure and defects on SO2 interaction with CeO2 provides insights for designing more sulfur-resistant CeO2-based catalysts.


2

Page 3 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1. INTRODUCTION CeO2 is widely used as a major component in three-way catalyst in the automobile industry because of its high oxygen storage capacity.1-4 The automobile exhaust consists of 5-20 ppm SO2. Several studies reported that even at such low concentration sulfur oxide could severely reduce the catalytic performance of CeO2-based catalysts.5-12 Therefore, there has been a strong research interest devoted to understanding how SO2 deactivates the catalyst. The interaction of SO2 with CeO2 (pure) or mixed oxides has been studied extensively using IR13-16, Raman17, 18, XPS19-21, TPD22 techniques. It has been established that surface and bulk sulfates are formed during the interaction of SO2 with CeO2. The main focus of these studies was to evaluate the SOx species formed during SO2 adsorption on polycrystalline CeO2 whose surface structure is a complex mixture of various facets and structural defects. Therefore, it was difficult to correlate the type of SOx species to the specific surface structure of CeO2, thus preventing an improved understanding of the deactivation mechanism and further development and design of more sulfur-resistant CeO2 catalysts. One of the approaches to gain fundamental insights into the SO2 poisoning effect is to utilize model CeO2 systems such as nanoshaped CeO2 with well-defined crystallographic facets. To the best of our knowledge, the effect of the surface structure (facets) of CeO2 on the nature and strength of adsorbed SOx species has not been previously investigated. Recently, CeO2 nanocrystals with different morphologies including rods, cubes and octahedra have been synthesized and their surface facets have been carefully characterized.23 The crystallographic planes for CeO2 cubes and octahedra are unambiguously identified as (100) and (111), respectively while those for CeO2 rods are still in debate. While previous literature24, 25

has concluded that the rods consist of a mixture of (110) and (100) facets, more recent work 26,

27

suggested the surface termination of CeO2 rods is debatable and depends on the preparation


3


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

method and post-treatment conditions.

Page 4 of 33

Here CeO2 rods are considered representative of a

surface with the presence of large amount of defects.

Taking advantage of these CeO2

nanoshapes, the surface dependent redox and acid-base properties of CeO2 has been better understood via several probing molecules and reactions.28, 29 Along this line, in the current work we studied the adsorption and reaction of SO2 on three CeO2 nanocrystals, rods (defective structure), cubes (100) and octahedra (111).

Our hypothesis is that different degrees of

coordination unsaturation of the surface cations (Ce) and anions (O)29 on the three CeO2 nanocrystals could result in the different acid-base properties, which in turn results in the different binding strength with the acidic SO2.

As discussed below, our IR and Raman

spectroscopic studies indicate that adsorption of SO2 results in the formation of surface sulfates/sulfites and bulk sulfites on CeO2 nanocrystals, and that the structure of the adsorbed SOx species depends on the shape of the CeO2.

2. EXPERIMENTAL Material synthesis: The synthesis of CeO2 nanocrystals was discussed in detail elsewhere.30 CeO2 rods, cubes and octahedra nanocrystals were prepared by hydrothermal treatment in Teflon-lined stainless steel autoclave.24,

25

The surface structure of CeO2 was controlled by

varying the pH, temperature and duration of the reaction. Ce(NO3)3.6H2O and NaOH were used as precursors for CeO2 rods and cubes, whereas Ce(NO3)3.6H2O was used as a precursor for CeO2 octahedra. The fresh white precipitate was centrifuged after the hydrothermal synthesis, and washed with DI water and ethanol for three times. The Na impurities were removed by


4

Page 5 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


base-acid wash which was discussed in detail in our previous work.30 The base-acid washed products were dried in vacuum overnight and then calcined at 400oC in air for 4 hours. Material characterizations: The electron microscopy (TEM and SEM) images and XRD patterns of the CeO2 rods, cubes and octahedra were presented in our previous work.23,

30, 31

Representative TEM images of the three nanocrystals are shown in Figure S 1 in the Supporting Information. The mean particle sizes, determined from the Scherrer equation (the shape factor of 0.9 was used for all CeO2 samples), are 11, 43, and 83 nm for rods, cubes and octahedra, respectively. The high resolution electron micrographs of the nanoshapes23 show the cube and octahedra shapes from which (100) and (111) terminations, respectively, are inferred, in agreement with the previous studies of CeO2 nanomaterials.24,

25, 32

The high resolution

micrographs23 showed that the CeO2 rods used in our study are characterized as having a higher concentration of defected area than the cubes and octahedra, as evidenced by the presence of black dots in the TEM image of the rods. Here we consider CeO2 rods representative of a surface with the presence of large amount of defects. Brunauer-Emmett-Teller (BET) surface areas of the CeO2 nanocrystals, calcined at 400 oC, were measured via nitrogen adsorption at 196 oC using a Micrometrics Gemini 275 system. CeO2 rods, cubes and octahedra have surface areas of 72, 17 and 10m2/g, respectively. In situ IR and Raman studies: The catalyst was pretreated either in oxidative or reducing environments. Oxidation: the catalyst was heated to 400 oC at the rate of 10 oC/min and held for 1 hour in 5% O2/He flow at 25 cm3/min. Reduction: the catalyst was heated to 400 oC at the rate of 10 oC/min and held for 30 min in 5%O2/He flow at 25 cm3/min; the catalyst was then purged with He for 5 minutes before exposing to 4%H2/He gas for another 60 min at 400 oC. The in situ spectroscopic studies consist of the following steps (ii) SO2 adsorption, (ii) He purge, and (iii)


5


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 33

temperature programmed desorption (TPD). SO2 adsorption: the catalyst was exposed to 100 ppm SO2/He flow at 25 cm3/min for 60 min at 25 oC. He Purge: the residual gas in the reactor was purged by He flow at 25 cm3/min for 10 min at 25 oC.

Temperature programmed

desorption: the catalyst was heated to 400 oC at the rate of 10 oC/min; the temperature of the catalyst was held at 400 oC for 1 min, followed by cooling the catalyst to 25 oC at the rate of 10 o

C/min. Two SO2 adsorption and desorption cycles were performed on each sample. IR spectra were collected continuously through the SO2 adsorption cycle using Thermo

Nicolet Nexus 670 spectrometer in Diffuse Reflectance mode (DRIFTS), the outlet gases from the DRIFTS reactor (Pike Technologies HC-900) were analyzed using quadrupole mass spectrometer (Omnistar GSD-301 O2, Pfeiffer Vacuum). The Raman measurements were performed with a catalytic Raman reactor (Linkam CCR1000) using in-house built multiple-wavelengths Raman system33 that includes laser excitations at λ=325 and 532 nm. Raman scattering was collected using fiber optics connected directly to the spectrograph stage of the triple Raman spectrometer (Princeton Instruments Acton Trivista 555). Samples pretreatment and SO2 adsorption were carried out similarly as in the IR studies. DFT calculations: Spin-polarized density functional theory calculations were performed using the Vienna Ab Initio Simulation Package (VASP)34-37 with core electrons represented with projector augmented wave (PAW)38, 39 potentials and PBE exchange and correlation functional.40 For Cerium we include 12 valence electrons and for Sulfur and Oxygen we used 6 valence electrons. To account for on-site Coulomb effects we use U=5 eV for the Ce 4f states.44 Plane waves with an energy cutoff of 434.4 eV were used in all calculations, at least 18 Å of vacuum is included between periodic images, and all geometries were optimized until atomic forces were


6

Page 7 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


less than 0.02 eV / Å. Brillouin zone sampling was done at the Gamma point and we use Gaussian smearing with σ=0.2 eV. For the (100) surface, a 3 x 3 slab with 15 atomic layers is used. To satisfy the CeO2 formula we use an O terminated surface with half of the surface O atoms moved to the opposite face of the slab. The surfaces are symmetric and identical with the O atoms aligned along the same crystallographic direction. Adsorbates are added to both sides of the symmetric slab to counter dipole effects and at least 15 Å of vacuum is included between periodic images.

Vibrational frequencies are calculated by computing finite difference

approximations to the dynamical matrix. Adsorption energies are calculated as ࡱࢇࢊ࢙ = −ሾࡱሺࡿࡻ૛ + ࢙࢒ࢇ࢈ሻ − ૛ࡱሺࡿࡻ૛ ሻ − ࡱሺ࢙࢒ࢇ࢈ሻሿ

(1)

where, E(SO2+slab), E(SO2), and E(slab) are the total energies of the combined system, SO2 and the clean surface. Each geometry is optimized and identical simulation cell and Brillouin zone samplings are used in each case.

A positive value of adsorption energy indicates stable

adsorption geometry. The adsorption of SO2 on the (111) surface of ceria was previously studied by Lu et al.,41 with a 2 x 2 supercell and adsorbates on one face of the slab with dipole corrections. In contrast we used a slab with more layers and adsorbates on both faces to counter dipole effects. To ensure our calculations are consistent we calculated the vibrational modes of monodentate sulfite geometry and find 1019, 992, and 736 cm-1 in good agreement with the values reported by Lu et.al., (1016, 977, 709 cm-1).41 With such agreement we focus on this work on CeO2 (100) and its interaction with SO2.

3. RESULTS AND DISCUSSION 3.1. Structure of SOx species from SO2 adsorption


7


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 8 of 33

Figure 1(a) shows the IR spectra during SO2 adsorption on oxidized CeO2 rods, cubes, and octahedra at room temperature as a function of time. Absorbance spectra during SO2 adsorption was calculated as Abs= -log(I/Io), where I is the single beam spectrum during SO2 adsorption and Io is the single beam spectrum before SO2 adsorption. The IR spectra during SO2 adsorption on CeO2 rods exhibited characteristic peaks at 1422 (S=O), 1350 (S=O), 1235 (S-O), 1200 (S-O), 1023 (S-O, str), and 897 cm-1 (S-O, str), which indicates the formation of polynuclear surface sulfate species as shown in Scheme 1(a).12, 14, 15, 42-47 The broad IR features at 1064 and 1095 cm-1 (S-O, str) can be also ascribed to polynuclear surface sulfate species,47 possibly formed at different surface sites of CeO2 rods. The IR peak at 997 cm-1 indicates the formation of sulfite species48 as shown in Scheme 1(b), during SO2 adsorption. All the IR features generally grow in intensity with SO2 adsorption time. Note that the IR spectra exhibited some negative peaks in the 1450-1150 cm-1 region, which are due to perturbed surface carbonates. The negative peaks show that the SO2 can displace the pre-adsorbed carbonates that are formed during the sample storage in ambient atmosphere that were stable enough to stand the pretreatment at 400 oC. The rise in CO2 MS intensity profile during SO2 adsorption, shown in Figure S 2 in Supporting Information, confirms the displacement of pre-adsorbed CO2 on the catalyst by SO2 at room temperature. This displacement implies a stronger interaction between SO2 and CeO2 rods than between CO2 and CeO2, due to the more acidic nature of SO2. The Raman spectrum of oxidized CeO2 rods before SO2 adsorption in Figure 1(d) exhibits peaks at 1049 and 1174 cm-1. The peak at 1049 cm-1 is due to nitrate impurities formed during CeO2 rods synthesis and the peak at 1174 cm-1 represents the second-order longitudinal optical mode (2LO) of CeO2.23

The Raman

spectrum of oxidized CeO2 rods after SO2 adsorption exhibits peaks at 1021 (S-O) and 1372 (S=O) cm-1, characteristic of sulfates17, 18.


8

1185 1137 1063 1017 995 897

1342

1433

Octahedra 0.009

1399

Cubes

1174

1012

SO2 ads fresh

0.3 SO2 ads

Octahedra

1174

0.08

Time (min) 0 0.5 (e) 1 1.5 2 5 10 15 30 45 60 (f)

Normalized Intensity (a.u.)

1111 1014 962 897

1430 1373 1338 1232

Cubes

(c)

λ=532 nm

0.3

0.02

(b)

1372

(d)

1174

Rods

1021 1049

1235 1200 1095 1064 1023 997 897

Rods

(a)

Absorbance (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1422 1350

Page 9 of 33

fresh

0.3 SO2 ads fresh

1500

1200

900 -1

Wavenumber (cm )

900

1200 1500 -1 Raman shift (cm )

Figure 1: IR spectra during different time intervals of SO2 adsorption on oxidized CeO2 (a) rods, (b) cubes, (c) octahedra, and Raman spectra of CeO2 (d) rods, (e) cubes, (f) octahedra before and after SO2 adsorption. (λ=532 nm), temperature during adsorption=25 oC.

IR spectra during SO2 adsorption on oxidized CeO2 cubes at room temperature, shown in Figure 1(b) exhibit peaks at 1430 (S=O, str), 1373 (S=O, str), 1338 (S=O, str), 1232 (S-O, str), 1111 (S-O, str), 1014 (S-O, str), and 897 cm-1 (S-O, str), indicating the formation of polynuclear surface sulfates 36-41 as shown in Scheme 1(a). The IR spectra also exhibit peaks at 1014 and 962 cm-1 which are due to the formation of sulfite species49 as shown in Scheme 1(b). The intensity of all the peaks in the range of 1430 - 897 cm-1, increases gradually with SO2 adsorption time except for the peaks at 1014 and 962 cm-1 which instead exhibit a subtle decrease in the intensity. This may suggest the conversion of some of the sulfite species into sulfates on the cubes surface. The Raman spectrum after SO2 adsorption on CeO2 cubes shown in Figure 1(e) exhibits a peak at 1399 (S=O) and 1012 cm-1 (S-O) which confirms the formation of sulfates.17, 18


9


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 10 of 33

Scheme 1: Proposed structure of (a) polynuclear surface sulfate, (b) surface sulfite, (c) bulk sulfite,17, 46 and (d) isolated sulfate.

The IR spectra during SO2 adsorption on oxidized CeO2 octahedra shown in Figure 1(c) exhibit peaks at 1017 (S-O, str) and 995 cm-1 (S-O, str), due to sulfite species as shown in

Scheme 1(b). The IR spectra also exhibits the peaks at 1435 (S=O, str), 1342 (S=O, str), 1185(S-O, str), 1137 (S-O, str), 1063(S-O, str), and 897 cm-1 (S-O, str), due to polynuclear surface sulfates as shown in Scheme 1(a). The peaks at 1017 and 995 cm-1 are more intense than those at 1435, 1342, 1185, 1137, 1063, and 897 cm-1 at the early stage of SO2 adsorption, indicating that the sulfite species are formed predominantly during initial SO2 adsorption on oxidized CeO2 octahedra. These IR features increase in intensity and then stabilize with SO2 adsorption time. The Raman spectrum after SO2 adsorption on octahedra shown in Figure 1(f) is essentially the same as that of the fresh sample. The absence of characteristic Raman peaks for the sulfite and sulfate species on octahedra may indicate that the amount of SOx species is limited on octahedra. Yet another contributing factor is that Raman is a bulk technique, so it is less sensitive than IR to the surface species on CeO2 especially for octahedra which have the largest particle size (83 nm) and lowest surface area (10 m2/g) among the three CeO2 nanocrystals.

Figure 2(a) shows the IR spectra during SO2 adsorption on reduced CeO2 rods. The IR peaks at 1412 and 1235 cm-1 indicate the formation of polynuclear surface sulfates. In the range


10

Page 11 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


below 1200 cm-1, the IR spectra were dominated by a peak at 995 cm-1 during the first ten minutes of SO2 adsorption, ascribed to surface sulfite species. The peak at 995 cm-1 was overtaken by a higher wavenumber one at 1039 cm-1 with extended SO2 adsorption time. The broad IR peak at 1039 cm-1 which is at lower wavenumber than the peaks (a broad peak centering 1145 cm-1 with a shoulders at 1240, 1060, and 990 cm-1) reported for bulk sulfate in the literature43 , suggests the formation of bulk sulfites on the reduced CeO2 rods as shown in

Scheme 1(c). The Raman spectrum during SO2 adsorption on reduced CeO2 rods shown in Figure 2(d), exhibits a peak at 1021 cm-1 similar to that observed on oxidized CeO2 rods, indicating the formation of sulfates/sulfites on reduced CeO2 rods upon SO2 adsorption. The peak at 1374 cm-1 indicates the formation of sulfates. The IR spectra during SO2 adsorption on reduced CeO2 cubes (Figure 2(b)) exhibit similar features as those on oxidized CeO2 cubes (Figure 1(b)). The peaks at 1435, 1378, 1343, 1232, 1111, and 897 cm-1 imply the formation of polynuclear surface sulfates, while those at 1014 and 962 cm-1 suggest the formation of sulfites. The Raman spectrum after SO2 adsorption shown in

Figure 2(e), exhibits the characteristic peak at 1389 cm-1 for sulfates. The presence of a weak peak at 991 cm-1 indicates the formation of sulfites on reduced CeO2 cubes upon SO2 adsorption. The IR spectra during SO2 adsorption on reduced CeO2 octahedra shown in Figure 2(c) exhibit peaks due to polynuclear surface sulfates at 1452, 1380, 1339, 1185, 1063, and 897 cm-1. The peaks at 1017 and 995 cm-1 indicates the formation of sulfite species. The Raman spectrum after SO2 adsorption on reduced CeO2 octahedra shown in Figure 2(f) is the same as that of the fresh sample. This again implies that Raman is not as sensitive as IR to surface SOx species on the large octahedra particles. The assignment for the IR and Raman bands from SO2 interaction on differently pretreated CeO2 nanocrystals is summarized in Table 1.


11

0.004

λ=532 nm

1374

1174

Page 12 of 33

1021

Rods

1039 995

1235

Rods

0.009

1500

1389

1174

Cubes

991

SO2 ads fresh

0.3 SO2 ads fresh

Octahedra

1174

1063 1017 995 897

1452 1380 1339

Octahedra

1185

0.08

Time (min) 0 0.5 1 1.5 2 5 10 15 30 45 60


1111 1014 962 897

Cubes

1232

0.4

1435 1378 1343

Absorbance (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

1412


SO2 ads fresh

0.15

1200 900 -1 Wavenumber (cm )

900

1200

1500 -1

Raman shift (cm )

Figure 2: IR spectra during different time intervals of SO2 adsorption on reduced CeO2 (a) rods, (b) cubes, and (c) octahedra, and Raman spectra of reduced CeO2 (d) rods, (e) cubes, and (f) octahedra before and after SO2 adsorption. (λ=532 nm), temperature during adsorption=25 oC.


12

Page 13 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 1: Summary of band assignments of adsorbed SOx on CeO2 rods, cubes and octahedra for oxidation and reduction pretreatment cases. R: Wavenumbers for Raman spectra Wavenumbers (cm-1) 1422, 1350,1372R 1235, 1200, 1023, 1064,

CeO2 sample

CeO2 rods Oxidized

897

Assignment S=O; surface sulfates S-O; surface sulfates

997 1021 R,

S-O; surface sulfite S-O; sulfate

CeO2 rods reduced

1412, 1374 R 1235 1021R 1039

S=O; surface sulfates S-O; surface sulfates S-O; sulfates S-O; bulk sulfites

CeO2 cubes oxidized

1430, 1373, 1338,1399R 1232, 1111, 1014,, 897 1012R 1014, 962

S=O; surface sulfates S-O; surface sulfates S-O; sulfates S-O; surface sulfites

CeO2 cubes reduced

1435, 1378,1343, 1389R 1232, 111, 897, 991R

S=O; surface sulfates S-O; surface sulfates S-O; sulfites

CeO2 octahedra oxidized

1433, 1342, 1185, 1137, 1063, 897 1017, 995,

S=O; surface sulfates S-O; surface sulfates S-O; surface sulfites

CeO2 octahedra reduced

1452, 1380, 1339 1185, 1063, 897 1017, 995

S=O; surface sulfates S-O; surface sulfates S-O; surface sulfites

In perspective, SO2 adsorption was previously investigated on CeO2 powders and thin films. Lavalley48 and Waqif et.al.,43 reported that adsorption of SO2 on the polycrystalline CeO2 results in the formation of sulfite species at room temperature and sulfate species at higher temperature (300-800 oC). Overbury et.al.,19 and Smirnov et.al.,21 also reported that the SO2 adsorbs on CeO2 (111) thin films in the form of sulfites at low temperature (≤200 oC) and as sulfates at high temperatures (≥ 300 oC). According to the theoretical calculations reported by Lu et.al.,41 SO32like structure forms on both (110) and (111) surface of both stoichiometric and partially reduced


13


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 33

CeO2 and SO42- like structure forms on (110) surface. The IR and Raman spectroscopic results in this study showed that the adsorption of SO2 on CeO2 octahedra leads to mostly sulfite species. The main IR features observed at 995 and 1017 cm-1 for the surface sulfite species on octahedra are in nice agreement with the DFT prediction by Lu et al. for SO2 interaction with CeO2 (111) surface.41 The observation of some surface sulfate on octahedra is likely associated with the presence of scare oxygen-vacancies and the unsaturated sites on the edges and corners of the octahedra where the surface oxygen is more reactive than the oxygen on the CeO2 (111) basal plane. Since the O-O distance (3.9 Å) on an intrinsic (111) is not possible for sulfate species to form,41 the defect sites and coordination unsaturated sites on CeO2 octahedra can result in relaxation of the surface lattice oxygen so that the formation of sulfates is possible on the (111) surface of octahedra. The formation of similar SOx species in both oxidatively and reductively pretreated CeO2 octahedra nanocrystals is likely due to the fact that octahedra have the most stable (111) surface and are more difficult to reduce at the reduction temperature of 400 oC as compared to the CeO2 rods and cubes.23 SO2 adsorption on CeO2 (100) surface was not previously investigated via either experiments or DFT. Since our CeO2 (111) model well replicates the DFT result from Lu et al.,41 on the same surface, we used CeO2 (100) surface with a similarly sized slab model. Our DFT calculations considered various structures of the adsorbed SOx on CeO2 (100) surface (Figure 3). The adsorption energies (Eads), vibrational frequencies (S-O), and S-O and Ce-O bond lengths of the optimized structures are summarized in Table S1 in Supporting Information.

The DFT

calculations show sulfite (structure III), chemisorbed SO2 (structures IV-VI), sulfate (structure VIII) are the favorable species that are formed upon SO2 interaction with the CeO2 (100) surface. Compared to gas phase SO2 (S-O 1.46 Å), there is 1 longer S-O bond (1.47 Å ) and 2 shorter S-O


14

Page 15 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


bonds in sulfite species (1.43, 1.43 Å). For sulfate species, there is one shorter S-O bond (1.43 Å) and three longer S-O bonds (1.53, 1.53, 1.54 Å) than the gas phase SO2. In both cases, the longer S-O bond lengths are associated with surface oxygen. On the pristine CeO2 (100), the S-O bond length of sulfite species closest to the surface to the surface oxygen (S-Osurf) is slightly higher than the sulfate species. The IR results show that the characteristic peaks for sulfate species are observed at higher wavenumbers (1232, 1111 cm-1) than the sulfite species (962 cm1

) which are consistent with the DFT calculations. The large O-O distance(3.71 Å)50 in pristine

(100) is not favorable for the formation of sulfate species, however the low vacancy formation energy (2.27 eV compared to 2.60 eV for (111) surface)51 of (100) surface allows the formation of the sulfate species. Our experimental IR results showed the formation of sulfites and sulfates when SO2 adsorbs on CeO2 cubes, in agreement with the theoretical calculations. Yet the characteristic peaks for the chemisorbed SO2 were not conspicuous in the IR spectra. Furthermore, the DFT calculations show that the sulfites (structure X and XI) with S-Osurf bond length of 1.58 and 1.62 Å are energetically more favorable when an oxygen vacancy is introduced into the CeO2 (100) surface. This is also reflected in the vibrational frequency changes of the sulfite species formed on an intact and O-vacancy inducedCeO2 (100) surfaces. According to the DFT calculations, the IR features of the sulfite species on reduced CeO2 (100) surface should be seen at lower wavenumber than the sulfites on oxidized CeO2 (100) surface. However, adsorption of SO2 on both oxidized and reduced CeO2 cubes results in similar IR features of sulfites (see Figures 1 and 2).

This discrepancy between DFT and the IR

observations is probably related to the surface characteristic of CeO2 cubes as revealed in our recent microscopy work.52. It was found that the CeO2 (100) facet on cubes is terminated with a mixture of O and Ce, in contrast to the theoretical full O or Ce termination due to the polar


15


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

character.28,

53

Page 16 of 33

The real (100) surface of CeO2 cubes is composed of patches of O and Ce

clusters, rendering the surface highly defective even under an oxidation state. So the type of sulfite species is similar upon SO2 adsorption irrespective of the oxidative or reduction pretreatment of CeO2 cubes. This surface property also makes the formation of surface sulfate possible as SO2 interacts with the O patches on the CeO2 cube surface. It is known that comparison of the absolute vibrational frequencies of adsorbates is challenging between experimental results and DFT calculations due to the systematic errors inherent in DFT. For example, the calculated harmonic symmetric and antisymmetric stretching frequencies (Table S 1) for gas phase SO2 from our DFT are underestimated (about 80 cm-1) when compared with the measured ones for SO2.54 So it is more advantageous to compare the trend, i.e., frequency shift, instead of absolute frequencies. The introduction of one O-vacancy into the CeO2 (100) slab substantially lowers the frequencies of surface sulfite species, i.e., ~ 170 cm-1 for the ω1 mode: ~ 1370 → ~ 1200 cm-1. So an even larger frequency shift is expected for surface sulfite species when more O-vacancies are introduced into the (100) slab model, as is the case of true CeO2 cube surfaces. Thus the observation of the main sulfite feature in the IR spectra at around 1000 cm-1 is reasonable. The surface of CeO2 rods is highly defective and reducible, so the formation of both sulfite and sulfate upon SO2 adsorption is reasonable. If we adapt the previously proposed (110) + (100) facets for CeO2 rods, both Lu’s DFT41and our calculations predict the formation of sulfite and sulfate species. This can also explain our IR and Raman observations on the CeO2 rods. The utilization of CeO2 with well-defined crystallographic facets is helpful for understanding the nature of adsorbed SOx species and how they depend on the surface structure of CeO2. Although the type of surface species (sulfites and sulfates) formed from SO2 adsorption does not vary with


16

Page 17 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


the surface structure of CeO2 nanocrystals, the relative amount of sulfites vs. sulfate depends on the shape of the CeO2 nanocrystals. According to the relative IR peak intensity (Figures 1 and

2), the extent of sulfate formation on the three CeO2 surfaces follows rods ≥ cubes > octahedra while the formation of sulfites follows the reverse order. Since the formation of sulfates from SO2 requires the reduction of CeO2, the trend can be explained by the reactivity or vacancy formation energy of the surface oxygen of the CeO2 nanocrystals that follows the trend: rods > cubes > octahedra.23, 29 We also note that the frequency of the S=O of the surface sulfate species exhibited subtle variation (1422 and 1350 cm-1 for rods, 1430 and 1373 cm-1 for cubes, 1433 and 1342 cm-1 for octahedra) with the surface structure of CeO2 nanocrystals. This variation could be due to the different atomic arrangements on the surface of CeO2 nanocrystals, i.e., surface OO distance and surface O coordination number are different on the three CeO2 surfaces29, 41 so that the formed SOx species can undergo different stress and thus show variation in vibrational frequencies.


17


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 33

Figure 3: Optimized SOx structures on CeO2 (100). (I) sulfite – 1, (II) sulfite – 2, (III) sulfite-3, (IV) chemisorbed SO2 – 1, (V) chemisorbed SO2 – 2, (VI) chemisorbed SO2 – 3, (VII) physisorbed SO2 – 1, (VIII) sulfate – 1, (IX) Chemisorbed SO2 adjacent to O vacancy, (X) sulfite adjacent to O vacancy, (XI) sulfite adjacent to O vacancy-2. Color coding: red – oxygen, yellow – sulfur, light blue – Ce4+, dark blue – Ce3+. The adsorption energy of sulfate species (VIII) was not presented the due to the additional chemistry required, e.g. reduction of ceria surface. We identify the nonmagnetic Ce4+ (4f0) and magnetic Ce3+ (4f1) by the spin charge associated with each Ce atom. Bader volumes are assigned using the all-electron charge density and spincharges are calculated from the magnetization density.55


18

Page 19 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


3.2. Strength of SOx species from SO2 adsorption Figure 4(a) shows the IR spectra collected during TPD after room temperature SO2 adsorption on the oxidized CeO2 rods.

IR spectra during TPD was calculated as Abs= -

log(ITPD,/IP), where ITPD is the single beam spectrum during TPD, and IP is the single beam spectrum collected during cooling at the same temperature after the pretreatment. The intensity of the IR peaks at 1422, 1235, 1350, 1192, 1023, and 897 cm-1 decreases with the increase of desorption temperature from 40 – 200 oC, indicating the desorption of some of the weakly bound surface sulfates. At the temperature above 200 oC, the intensity decrease of the peak at 997 cm-1 is accompanied by intensity increase of the peak at 1363 cm-1, indicating the conversion of the sulfite species to isolated sulfate species (Scheme 1(d)). An intensity increase is also evident for the feature at around 1200 cm-1 at temperature above 200 oC, suggesting the formation of some bulk sulfate species. The IR spectra also exhibits decrease in the area of the broad peak between 1095 and 1023 cm-1 that represents the surface sulfates. The decrease in the peak area in the range 1150 – 1000 cm-1 is plotted as a function of desorption temperature and shown in Figure

4(b). The decrease in peak area suggests the decomposition/desorption of surface sulfates during TPD. The temperature of maximum desorption shown in Figure 4(c) is a qualitative indicator of the strength of the adsorbed species. The MS intensity profile shows two SO2 desorption peaks at 90 and 235 oC, assignable to the desorption of surface sulfate and sulfite species. The large temperature difference between the two desorption peaks suggests the presence of more than one type of surface sulfite and sulfate due to the heterogeneous nature of CeO2 rods surface.


19

360 280 Heating 200 120 40 1200 900 -1 Wavenumber (cm )

Peak area (a.u.)

400

1500

Page 20 of 33

1.7

Absorbance (a.u.)

0.03

Normalized MS intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

1422 1363 1350 1235 1200 1192 1095 1064 1023 997 897


1.6 1.5 1.4 1.3

0

100

200

300

400

o

Temperature ( C)

o

235 C

4.50E-012 3.60E-012 2.70E-012 o

90 C

1.80E-012 9.00E-013 0.00E+000

40

140

240

340 400 o

Temperature ( C)

Figure 4: (a) IR absorbance spectra and (c) normalized SO2 MS intensity (m/e=64) profile during TPD on oxidized CeO2 rods. The MS intensity was normalized to the surface area of the CeO2 nanocrystals. (b) Intensity change of the IR features in the range of 1150 – 1000 cm-1 vs. desorption temperature.

Figure 5(a) - (c) show the room temperature IR spectra of CeO2 rods, cubes, and octahedra nanocrystals from SO2 adsorption and then after He purge and after TPD. Absorbance was calculated as Abs= -log(I/Io), where I is the single beam spectrum after He purge and after TPD of 1st cycle and Io is the single beam spectrum of fresh sample before SO2 adsorption. The IR spectrum of CeO2 rods, Figure 5(a) after He purge exhibits similar peaks to those in the presence of SO2 (Figure 1(a)), which indicates that the surface sulfates/sulfites were not


20

Page 21 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


removed during He purge at room temperature. The slight decrease in the intensities of the peaks at 1422, 1235, 1064, 997, and 897 cm-1 after TPD indicates that some of the polynuclear surface sulfates and sulfites species have desorbed from the catalyst upon thermal desorption. The Raman spectrum of CeO2 rods, shown in Figure 5(d) exhibits peaks at 995, 1021 and 1372 cm-1 that confirms the presence of surface sulfite species and bulk sulfates and after TPD. The IR spectra of CeO2 cubes shown in Figure 5(b) exhibit decrease in intensity for peaks at 1430, 1373, and 1232 cm-1, indicating the desorption of some portion of sulfates after TPD. The IR and Raman spectra (Figure 5(e)) of CeO2 cubes after TPD still present peaks due to surface sulfates and sulfites as well as some bulk sulfates (~1200 cm-1). The IR spectrum of CeO2 octahedra after TPD shown in Figure 5(c) exhibits the peaks at 1185, 1137, 1063, 1017 and 995 cm-1. In comparison to the spectrum before TPD, the peak at 1433 and 897 cm-1 disappear and the peaks between 1185 - 897 cm-1 show reduction in intensities, implying that the most of the adsorbed species are removed during TPD. The IR and Raman spectra for reduced CeO2 nanoshapes before and after TPD are compared in Figure S 3 in Supporting Information. Briefly, the spectra show that the SOx species on reduced CeO2 rods and cubes do not desorb completely as in the oxidized cases. The reduced CeO2 octahedra nanocrystals retain little of the adsorbed SOx species after TPD, similar to the case on the oxidized counterpart.


21

λ=532 nm

1372

995 1021 1045

(d)

0.02

1174

Page 22 of 33

Rods

897

(a)

1235 1200 1095 1064 1023 997

Rods

0.3

after TPD

after TPD fresh

after TPD He purge

(c)

1185 1137 1063 1017 995 897

Octahedra

0.009

1500

(f)

after TPD He purge 1200

900

1399

1174

972 1012

Cubes 0.3

after TPD fresh

Octahedron

1174

0.08

(e)


1111 1014 962 897

1232

Cubes

1433

(b)

1430 1373

He purge

Absorbance (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

1422 1350


0.3 after TPD 900

-1

Wavenumber (cm )

fresh 1500

1200 -1

Raman shift (cm )

Figure 5: IR spectra of oxidized CeO2 (a) rods, (b) cubes, (c) octahedra obtained after He purge (T=25 oC) and after TPD to 400 oC (T=25 oC). Corresponding Raman spectra of fresh and after TPD for oxidized CeO2 (d) rods, (e) cubes and (f) octahedra.

Figure S4(a) in Supporting Information compares the IR spectra during 2 cycles of SO2 adsorption on oxidized CeO2 rods. In comparison to the 1st adsorption cycle, the 2nd cycle shows a decrease in intensities of the characteristic bands for adsorbed SOx species on CeO2 rods. This can be attributed to the degradation of the catalyst due to the strongly bound SOx species from 1st adsorption cycle. The IR spectra during 1st cycle of SO2 adsorption on CeO2 rods exhibited characteristic peaks of polynuclear surface sulfate species at 1422, 1350, 1235, 1023, 897 cm-1 whereas, the absorbance spectrum during 2nd cycle SO2 adsorption exhibit peaks at 1367, 1328, 1235, 1043 cm-1 which indicate the SO2 adsorbed in the form of isolated sulfate species shown in

Scheme 1(d) on degraded catalyst. The peak at 975 cm-1 is more evident on the 2nd cycle sample than on the 1st cycle, suggesting that more SO2 adsorbs in the form of sulfites on the degraded


22

Page 23 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


catalyst than on the fresh CeO2 rods. IR spectrum during 2nd cycle SO2 adsorption on reduced CeO2 rods, shown in Figure S4(d), does not exhibit any characteristic peaks of the adsorbed SOx species. The reason for the absence of IR peaks is not yet clear since the corresponding MS results in Figure 6(b) show that appreciable amount of SO2 was desorbed during 2nd cycle from CeO2 rods. The IR spectrum during 2nd SO2 adsorption cycle on oxidized (Figure S4(b)) and reduced CeO2 cubes (Figure S4(e)) exhibits similar features as the IR spectrum during 1st SO2 cycle but with decreased band intensity. The decrease in the intensity of the IR peaks during 2nd cycle indicates the degradation of the catalyst due to the accumulation of the adsorbed SOx. The IR spectrum of adsorbed SOx during 2nd cycle on oxidized (Figure S4(c)) and reduced (Figure

S4(f)) CeO2 octahedra exhibits only slight decrease in IR intensities compared to that after 1st cycle, indicating that the catalyst was only slightly degraded. This is consistent with the fact that CeO2 octahedra bind SO2 more weakly than CeO2 rods and cubes.

Figure 6 compares the normalized SO2 MS intensity profiles during TPD after two cycles of SO2 adsorption on oxidized and reduced CeO2 rods, cubes, octahedra nanocrystals. The SO2 MS intensities were normalized to the surface area of the CeO2 nanoshapes. As shown in Figure

6(a), oxidized CeO2 rods give SO2 desorption maxima at 90 and 235 oC for first cycle and at 130 and 210 oC for second cycle. Whereas, SO2 desorption occurs at 90 oC for both cycle 1 and cycle 2 on oxidized CeO2 cubes (Figure 6(b)) and octahedra (Figure 6(c)) nanocrystals. On reduced CeO2 rods (Figure 6(d)), two SO2 desorption maxima are present at 90 and 210 oC for first cycle and at 90 and 280oC. SO2 on reduced CeO2 cubes, shown in Figure 6(e) desorbs at 102 oC for the first cycle and 96 oC for the second cycle, whereas SO2 on reduced CeO2 octahedra (Figure 6(f)) desorbs at 93 oC for the first cycle and 90 oC for the second cycle. As the first desorption peak for CeO2 rods occurs at similar temperature to the major desorption on


23


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 33

both cubes and octahedra where surface sulfites and sulfates are the most abundant species, the desorption maxima at 90 oC could be assigned to the decomposition of weakly bound surface sulfite and sulfate species. The high temperature desorption peak for CeO2 rods can be due to the decomposition of the strongly bound sulfate species. Clearly, CeO2 rods exhibited the highest SO2 desorption temperature among the three nanocrystals. Higher temperature of desorption implies that SO2 binding strength is stronger on CeO2 rods than on cubes and octahedra. The interaction between SO2 and CeO2 is determined by three factors: the coordination status of surface oxygen, the surface O-O distance, and the surface O-vacancy formation energy (for sulfate formation). Low coordination status of surface oxygen favors high surface basicity of ceria which is also promoted by the presence of large amount of O-vacancies (i.e., low O-vacancy formation energy). O-vacancies are also favorable sites for SO2 adsorption. Therefore, the interaction strength of SO2 with ceria is expected to follow the same trend of basicity strength and defects amount of the three ceria nanoshapes as studied previously, rods > cubes > octahedra.29 CeO2 rods, which are more basic (due to the low coordination status of oxygen) and have large amount of defects, can thus, adsorb SO2 more strongly than cubes and octahedra nanocrystals, giving rise to a higher temperature of the maximum desorption of SO2. Even though the basicity strength and defects amount of ceria cubes and octahedra are quite different according to our recent work,29 the SO2 binding strength does not vary much on the two ceria surfaces. This is likely due to fact that the CeO2 (111) and (100) surfaces have very similar closest surface O-O distance (about 3.7 angstrom), which is not favorable for forming stable sulfate species from SO2 interaction.

Therefore, the surface

basicity, amount of defects and the surface structure (O-O distance) of ceria work together in determining the interaction strength of SO2. The general decrease in the desorption temperature


24

Page 25 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


for CeO2 rods, cubes and octahedra during 2nd cycle could be due to the degradation of the catalysts caused by the accumulated SOx species from 1st cycle.

The degraded catalysts

generally bind SO2 weakly resulting in the decrease in the temperature of maximum desorption. This finding is similar to the one observed by Chakravatula et.al.,56 where the adsorbed CO2 binds weakly on the degraded amine sorbents than on the fresh amine sorbents. One exception is the increase in desorption maxima during 2nd SO2 adsorption cycle on reduced CeO2 rods, likely due to the difference in the decomposition temperatures of sulfite and sulfate species. The huge rise in SO2 desorption amount during 2nd cycle TPD on CeO2 rods is due to the fact that the degraded rods (after first cycle) can no longer strongly hold sulfate/sulfite species so that majority of the SOx species on CeO2 rods decomposes/desorbs during TPD. In contrast, the temperature and amount of SO2 desorption during 2nd cycle for oxidized CeO2 cubes and octahedra nanocrystals did not exhibit evident change, suggesting that the degradation of these catalysts was too slight to affect the binding strength of adsorbed SOx. Reduction pretreatment of CeO2 rods, cubes and octahedra nanocrystals generally leads to a slightly higher SO2 desorption temperature compared to the oxidized counterparts. This is due to the fact that reduction of CeO2 nanoparticles results in the increase of oxygen vacancies, which in turn results in the increase of SO2 binding strength.29 Reduced CeO2 rods exhibit a more drastic increase in the temperature of desorption than the reduced cubes and octahedral nanocrystals. The difference in the degree of shift in the temperature of desorption is due to the difference in the reducibility of the three surfaces. Our previous studies showed that the CeO2 rods are readily reducible, followed by CeO2 cubes while CeO2 octahedra exhibited slight reduction after H2 treatment at 400 oC.23 So the presence of larger amount of oxygen vacancies


25


and higher surface basicity of reduced CeO2 rods results in the strong SO2 interaction upon hydrogen pretreatment.

Oxidized

Reduced Rods

(a)

3.00E-011

o

130 C

(b)

235 C

o

2.50E-012 1.25E-012 90oC

o

1.00E-011

Cycle 1 Cycle 2 o 280 C

(d)

210 C

o

2.00E-011 Normalized MS intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 26 of 33

90 C 0.00E+000 40 140

240

0.00E+000 340 400 40 Cubes

o

210 C

140

240

340 400

(e)

4.00E-011

4.50E-011

o

102 C

o

90 C

3.00E-011

2.00E-011 0.00E+000

o

96 C

1.50E-011 40

140

240

0.00E+000 40 340 400 Octahedra

(c)

140

240

340 400

(f)

6.00E-011

o

90 C

4.00E-011 2.00E-011 0.00E+000

7.50E-011

o

93 C

5.00E-011 o

90 C

2.50E-011 40

140 240 340 400 o Temperature ( C)

0.00E+000

40

140 240 340 400 o Temperature ( C)

Figure 6: Normalized SO2 MS intensity during TPD for oxidized CeO2 (a) rods, (b) cubes, and (c) octahedra and reduced CeO2 (d) rods, (e) cubes, and (f) octahedra nanocrystals. The MS intensity was normalized to the surface area of the CeO2 nanocrystals.

The amount of desorbed SO2 during the two TPD cycles on the three CeO2 nanoshapes is quantified and compared in Table 2 for the oxidation and reduction pretreated cases. During the 1st cycle, the oxidized CeO2 octahedra desorb slightly more SO2 (2.42 µ mol/m2) than cubes


26

Page 27 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(2.07 µ mol/m2) but significantly more than rods (0.45 µ mol/m2). This is consistent with the fact that SO2 binding strength on CeO2 rods and cubes is higher than octahedra and not all the adsorbed SOx species completely desorb during TPD, especially on rods. During 2nd cycle, oxidized CeO2 rods desorb significantly larger amount of SO2 compared to 1st cycle, which suggests that the decrease in the SO2 binding strength due to the degradation of CeO2 rods caused by the accumulated SOx species. Again, the trend is similar for SO2 desorption amount on the three reduced nanocrystals during the two cycles i.e., reduced CeO2 rods give the smallest SO2 desorption for the 1st cycle and then SO2 desorption increases greatly on the regenerated rods. On the reduced CeO2 nanoshapes, the amount of SO2 desorption is generally larger than on the oxidized counterparts, except for the reduced CeO2 rods. This is due to the increase in basicity and creation of vacancies, both of which facilitates SO2 adsorption. The amount of SO2 desorbed during 2nd cycle-TPD on reduced CeO2 rods is less than the oxidized counterpart. This could be due to the fact that SO2 binds strongly on the reduced CeO2 nanoshapes and not all the SOx was desorbed from the reduced CeO2 rods even after the 2nd cycle.

Table 2: SO2 desorption amounts during TPD on oxidized and reduced CeO2 rods, cubes and octahedra nanocrystals.

Catalyst

Amount of SO2 desorbed during TPD (µmol/m2)

Oxidized Cycle 1 Rods 0.45 Cubes 2.07 Octahedra 2.42

Cycle 2 3.99 2.24 2.18

Reduced Cycle 1 0.05 3.35 4.32

Cycle 2 0.33 2.49 4.08


27


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 33

4. CONCLUSIONS IR and Raman spectroscopic results reveal that the nature, strength and amount of SOx species formed from SO2 adsorption are dependent on the surface structure of CeO2 nanocrystals. Surface sulfates and sulfites are formed from SO2 adsorption on CeO2 rods, cubes and octahedra in both pre-oxidized and pre-reduced states. Surface sulfites are more prominent on CeO2 octahedra than the other two surfaces while surface sulfate species are the most favored on CeO2 rods. Bulk sulfite species are also formed on pre-reduced CeO2 rods in addition to the surface sulfites and sulfates. DFT calculations reveal that the sulfites, chemisorbed SO2 and sulfates are favored on the stoichiometric CeO2 (100) surface and sulfites are predominant on the reduced CeO2 (100) surface, which is in agreement with the experimental results on CeO2 cubes. The difference in SO2 interaction with CeO2 nanoshapes is due to the difference in the reactivity of the surface oxygen and the amount of defect sites. CeO2 rods, which have high basicity and large amount of defects, bind SO2 more strongly than CeO2 cubes and octahedra nanocrystals. The effect of surface basicity (depending on the coordination of surface O) and the presence of defect sites on SO2 binding strength on CeO2 is evidenced by the higher desorption temperature of SO2 (a) on rods than on cubes and octahedra, and (b) on the reduced CeO2 nanoshapes than the oxidized counterparts. Strong interaction of SO2 with the most reactive and defective CeO2 rods result in stable bulk sulfate species after high temperature desorption and thus being the most degraded sample among the three CeO2 nanoshapes where octahedra is the least degraded. The work here suggests that a more sulfur-resistant CeO2 catalyst maybe designed by employing the stable (111) facet-terminated CeO2. Yet it remains challenging to fine tune the balance between stability (related to low energy surface) and reactivity (related to high energy surface and defects) of CeO2-based catalysts.


28

Page 29 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


AUTHOR INFORMATION Corresponding Author *Email: [email protected]. Telephone: 865-576-1080

Funding Sources U.S. Department of Energy, Office of Science, Basic Energy Sciences.

ACKNOWLEDGEMENTS This work is supported by the Center for Understanding and Control of Acid Gas-Induced Evolution of Materials for Energy (UNCAGE-ME), an Energy Frontier Research Center funded by U.S. Department of Energy, Office of Science, Basic Energy Sciences. Part of the work including the IR and Raman was conducted at the Center for Nanophase Materials Sciences, which is a DOE Office of Science User Facility. This research used computational resources of the National Energy Research Scientific Computing Center, which is supported by the U.S. DOE Office of Science under Contract No. DE-AC02-05CH11231.

ASSOCIATED CONTENT TEM images of CeO2 nanocrystals, adsorption energies and vibrational frequencies of the optimized structures on CeO2 (100), IR and Raman spectra of reduced CeO2 nanocrystal after He purge and TPD and comparison of IR Spectra of adsorbed SOx on oxidized and reduced CeO2 nanocrystals.


29


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 30 of 33

REFERENCES 1. Kašpar, J.; Fornasiero, P.; Graziani, M., Use of CeO2-Based Oxides in the Three-Way Catalysis. Catal. Today 1999, 50, 285-298. 2. Farrauto, R. J.; Heck, R. M., Catalytic Converters: State of the Art and Perspectives. Catal. Today 1999, 51, 351-360. 3. Trovarelli, A.; de Leitenburg, C.; Boaro, M.; Dolcetti, G., The Utilization of Ceria in Industrial Catalysis. Catal. Today 1999, 50, 353-367. 4. Kašpar, J.; Fornasiero, P.; Hickey, N., Automotive Catalytic Converters: Current Status and Some Perspectives. Catal. Today 2003, 77, 419-449. 5. Hilaire, S.; Sharma, S.; Gorte, R. J.; Vohs, J. M.; Jen, H. W., Effect of SO2 on the Oxygen Storage Capacity of Ceria-Based Catalysts. Catal. Lett. 2000, 70, 131-135. 6. Beck, D. D.; Sommers, J. W., Impact of Sulfur on the Performance of Vehicle-Aged Palladium Monoliths. Appl. Catal. B - Environ. 1995, 6, 185-200. 7. Beck, D. D.; Sommers, J. W.; DiMaggio, C. L., Axial Characterization of Oxygen Storage Capacity in Close-Coupled Lightoff and Underfloor Catalytic Converters and Impact of Sulfur. Appl. Catal. B - Environ. 1997, 11, 273-290. 8. Beck, D. D.; Sommers, J. W.; DiMaggio, C. L., Impact of Sulfur on Model Palladiumonly Catalysts under Simulated Three-Way Operation. Appl. Catal. B - Environ. 1994, 3, 205227. 9. Kolli, T.; Huuhtanen, M.; Hallikainen, A.; Kallinen, K.; Keiski, R. L., The Effect of Sulphur on the Activity of Pd/Al2O3, Pd/CeO2 and Pd/ZrO2 Diesel Exhaust Gas Catalysts. Catal. Lett. 2009, 127, 49-54. 10. Wang, Y.; Li, X.; Zhan, L.; Li, C.; Qiao, W.; Ling, L., Effect of SO2 on Activated Carbon Honeycomb Supported CeO2–MnOx Catalyst for NO Removal at Low Temperature. Ind. Eng. Chem. Res. 2015, 54, 2274-2278. 11. Park, E.; Kim, M.; Jung, H.; Chin, S.; Jurng, J., Effect of Sulfur on Mn/Ti Catalysts Prepared using Chemical Vapor Condensation (CVC) for Low-Temperature NO Reduction. ACS Catal. 2013, 3, 1518-1525. 12. Liu, J.; Li, X.; Zhao, Q.; Hao, C.; Wang, S.; Tadé, M., Combined Spectroscopic and Theoretical Approach to Sulfur-Poisoning on Cu-Supported Ti-Zr Mixed Oxide Catalyst in the Selective Catalytic Reduction of NOx. ACS Catal. 2014, 4, 2426-2436. 13. Dunn, J. P.; Jehng, J.-M.; Kim, D. S.; Briand, L. E.; Stenger, H. G.; Wachs, I. E., Interactions between Surface Vanadate and Surface Sulfate Species on Metal Oxide Catalysts. J. Phys. Chem B 1998, 102, 6212-6218. 14. Luo, T.; Gorte, R. J., Characterization of SO2-Poisoned Ceria-Zirconia Mixed Oxides. Appl. Catal. B - Environ. 2004, 53, 77-85. 15. Luo, T.; Vohs, J. M.; Gorte, R. J., An Examination of Sulfur Poisoning on Pd/Ceria Catalysts. J. Catal. 2002, 210, 397-404. 16. Bazin, P.; Saur, O.; Lavalley, J. C.; Blanchard, G.; Visciglio, V.; Touret, O., Influence of Platinum on Ceria Sulfation. Appl. Catal. B - Environ. 1997, 13, 265-274. 17. Flouty, R.; Abi Aad, E.; Siffert, S.; Aboukais, A.; Aboukaïs, A., Formation of Cereous Sulphate Phase upon Interaction of SO2 with Ceria at Room Temperature - Thermal Analysis, Raman and EPR Study. J. Therm. Anal. Calorim. 2003, 73, 727-734. 18. Twu, J.; Chuang, C. J.; Chang, K. I.; Yang, C. H.; Chen, K. H., Raman Spectroscopic Studies on the Sulfation of Cerium Oxide. Appl. Catal. B - Environ. 1997, 12, 309-324.


30

Page 31 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


19. Overbury, S. H.; Mullins, D. R.; Huntley, D. R.; Kundakovic, L., Chemisorption and Reaction of Sulfur Dioxide with Oxidized and Reduced Ceria Surfaces. J. Phys. Chem B 1999, 103, 11308-11317. 20. Smirnov, M. Y.; Kalinkin, A. V.; Pashis, A. V.; Prosvirin, I. P.; Bukhtiyarov, V. I., Interaction of SO2 with Pt Model Supported Catalysts Studied by XPS. J. Phys. Chem C 2014, 118, 22120-22135. 21. Smirnov, M. Y.; Kalinkin, A. V.; Pashis, A. V.; Sorokin, A. M.; Noskov, A. S.; Kharas, K. C.; Bukhtiyarov, V. I., Interaction of Al2O3 and CeO2 Surfaces with SO2 and SO2 + O2 Studied by X-Ray Photoelectron Spectroscopy. J. Phys. Chem B 2005, 109, 11712-11719. 22. Rodriguez, J.; Rodriguez, T.; Jirsak, A.; Freitag, J.; Hanson, J.; Larese, S.; Chaturvedi, Interaction of SO2 with CeO2 and Cu/CeO2 Catalysts: Photoemission, XANES and TPD Studies. Catal. Lett. 1999, 62, 113-119. 23. Wu, Z.; Li, M.; Howe, J.; Meyer, H. M.; Overbury, S. H., Probing Defect Sites on CeO2 Nanocrystals with Well-Defined Surface Planes by Raman Spectroscopy and O2 Adsorption. Langmuir 2010, 26, 16595-16606. 24. Yan, L.; Yu, R.; Chen, J.; Xing, X., Template-Free Hydrothermal Synthesis of CeO2 Nano-Octahedrons and Nanorods: Investigation of the Morphology Evolution. Cryst. Growth Des. 2008, 8, 1474-1477. 25. Mai, H.-X.; Sun, L.-D.; Zhang, Y.-W.; Si, R.; Feng, W.; Zhang, H.-P.; Liu, H.-C.; Yan, C.-H., Shape-Selective Synthesis and Oxygen Storage Behavior of Ceria Nanopolyhedra, Nanorods, and Nanocubes. J. Phys. Chem B 2005, 109, 24380-24385. 26. Agarwal, S.; Lefferts, L.; Mojet, B. L.; Ligthart, D. A. J. M.; Hensen, E. J. M.; Mitchell, D. R. G.; Erasmus, W. J.; Anderson, B. G.; Olivier, E. J.; Neethling, J. H. et al., Exposed Surfaces on Shape-Controlled Ceria Nanoparticles Revealed through AC-TEM and Water–Gas Shift Reactivity. Chemsuschem 2013, 6, 1898-1906. 27. Florea, I.; Feral-Martin, C.; Majimel, J.; Ihiawakrim, D.; Hirlimann, C.; Ersen, O., ThreeDimensional Tomographic Analyses of CeO2 Nanoparticles. Cryst. Growth Des. 2013, 13, 11101121. 28. Wu, Z.; Li, M.; Mullins, D. R.; Overbury, S. H., Probing the Surface Sites of CeO2 Nanocrystals with Well-Defined Surface Planes Via Methanol Adsorption and Desorption. ACS Catal. 2012, 2, 2224-2234. 29. Wu, Z.; Mann, A. K. P.; Li, M.; Overbury, S. H., Spectroscopic Investigation of SurfaceDependent Acid–Base Property of Ceria Nanoshapes. J. Phys. Chem C 2015, 119, 7340-7350. 30. Mann, A. K. P.; Wu, Z.; Calaza, F. C.; Overbury, S. H., Adsorption and Reaction of Acetaldehyde on Shape-Controlled Ceo2 Nanocrystals: Elucidation of Structure–Function Relationships. ACS Catal. 2014, 4, 2437-2448. 31. Wu, Z.; Li, M.; Overbury, S. H., On the Structure Dependence of CO Oxidation over CeO2 Nanocrystals with Well-Defined Surface Planes. J. Catal. 2012, 285, 61-73. 32. Zhou, K.; Wang, X.; Sun, X.; Peng, Q.; Li, Y., Enhanced Catalytic Activity of Ceria Nanorods from Well-Defined Reactive Crystal Planes. J. Catal. 2005, 229, 206-212. 33. Wu, Z.; Dai, S.; Overbury, S. H., Multiwavelength Raman Spectroscopic Study of SilicaSupported Vanadium Oxide Catalysts. J. Phys. Chem C 2010, 114, 412-422. 34. Kresse, G.; Hafner, J., Ab Initio Molecular-Dynamics Simulation of the Liquid-MetalAmorphous-Semiconductor Transition in Germanium. Phys. Rev. B 1994, 49, 14251-14269. 35. Kresse, G.; Hafner, J., Ab Initio Molecular Dynamics for Liquid Metals. Phys. Rev. B 1993, 47, 558-561.


31


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 32 of 33

36. Kresse, G.; Furthmüller, J., Efficiency of Ab-Initio Total Energy Calculations for Metals and Semiconductors Using a Plane-Wave Basis Set. Comp. Mater. Sci. 1996, 6, 15-50. 37. Kresse, G.; Furthmüller, J., Efficient Iterative Schemes for Ab Initio Total-Energy Calculations Using a Plane-Wave Basis Set. Phys. Rev. B 1996, 54, 11169-11186. 38. Kresse, G.; Joubert, D., From Ultrasoft Pseudopotentials to the Projector AugmentedWave Method. Phys. Rev. B 1999, 59, 1758-1775. 39. Blöchl, P. E., Projector Augmented-Wave Method. Phys. Rev. B 1994, 50, 17953-17979. 40. Perdew, J. P.; Burke, K.; Ernzerhof, M., Generalized Gradient Approximation Made Simple. Phys. Rev. Lett. 1996, 77, 3865-3868. 41. Lu, Z.; Müller, C.; Yang, Z.; Hermansson, K.; Kullgren, J., SOx on Ceria from Adsorbed SO2. J. Chem. Phys. 2011, 134, 184703. 42. Cotton, F. A., Advanced Inorganic Chemistry. Wiley: 1999. 43. Waqif, M.; Bazin, P.; Saur, O.; Lavalley, J. C.; Blanchard, G.; Touret, O., Study of Ceria Sulfation. Appl. Catal. B-Environ. 1997, 11, 193-205. 44. Morrow, B. A.; McFarlane, R. A.; Lion, M.; Lavalley, J. C., An Infrared Study of Sulfated Silica. J. Catal. 1987, 107, 232-239. 45. Saur, O.; Bensitel, M.; Saad, A. B. M.; Lavalley, J. C.; Tripp, C. P.; Morrow, B. A., The Structure and Stability of Sulfated Alumina and Titania. J. Catal. 1986, 99, 104-110. 46. Morterra, C.; Cerrato, G.; Bolis, V., Lewis and Brønsted Acidity at the Surface of Sulfate-Doped ZrO2 Catalysts. Catal. Today 1993, 17, 505-515. 47. Morterra, C.; Cerrato, G.; Emanuel, C.; Bolis, V., On the Surface Acidity of Some Sulfate-Doped ZrO2 Catalysts. J. Catal. 1993, 142, 349-367. 48. Lavalley, J. C., Infrared Spectrometric Studies of the Surface Basicity of Metal Oxides and Zeolites Using Adsorbed Probe Molecules. Catal. Today 1996, 27, 377-401. 49. Colthup, N. B.; Daly, L. H.; Wiberley, S. E., Introduction to Infrared and Raman Spectroscopy. Elsevier Science: 1990. 50. Nolan, M.; Watson, G. W., The Surface Dependence of CO Adsorption on Ceria. J. Phys. Chem B 2006, 110, 16600-16606. 51. Nolan, M.; Parker, S. C.; Watson, G. W., The Electronic Structure of Oxygen Vacancy Defects at the Low Index Surfaces of Ceria. Surf. Sci. 2005, 595, 223-232. 52. Lin, Y.; Wu, Z.; Wen, J.; Poeppelmeier, K. R.; Marks, L. D., Imaging the Atomic Surface Structures of CeO2 Nanoparticles. Nano Lett. 2014, 14, 191-196. 53. Ganduglia-Pirovano, M. V. n.; Hofmann, A.; Sauer, J., Oxygen Vacancies in Transition Metal and Rare Earth Oxides: Current State of Understanding and Remaining Challenges. Surf. Sci. Rep. 2007, 62, 219-270. 54. Kuchitsu, K.; Morino, Y., Estimation of Anharmonic Potential Constants. Ii. Bent Xy2 Molecules. Bull. Chem. Soc. Japan 1965, 38, 814-824. 55. Tang, W.; Sanville, E.; Henkelman, G., A Grid-Based Bader Analysis Algorithm without Lattice Bias. J. Phys.: Condens. Matter 2009, 21, 084204. 56. Srikanth, C. S.; Chuang, S. S. C., Infrared Study of Strongly and Weakly Adsorbed CO2 on Fresh and Oxidatively Degraded Amine Sorbents. J. Phys. Chem C 2013, 117, 9196-9205.


32

Page 33 of 33

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


TOC: Temperature of maximum SO2 desorption during TPD on CeO2 rods, cubes, and octahedra.


33

Surface Structure Dependence of SO2 Interaction with Ceria

Recommend Documents