Adsorption Kinetics and Dynamics of CO, NO, and CO2 on Reduced

Apr 18, 2008 - Brian H. Solis , Joachim Sauer , Yi Cui , Shamil Shaikhutdinov , and Hans-Joachim Freund. The Journal of Physical Chemistry C 2017 121 ...
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J. Phys. Chem. C 2008, 112, 7390-7400

Adsorption Kinetics and Dynamics of CO, NO, and CO2 on Reduced CaO(100) E. Kadossov and U. Burghaus* Department of Chemistry, Biochemistry, and Molecular Biology, North Dakota State UniVersity, Fargo, North Dakota ReceiVed: January 25, 2008; In Final Form: February 26, 2008

The adsorption kinetics and dynamics of typical combustion gases have been studied on CaO(100) by means of temperature programmed desorption and molecular beam scattering. In addition, the sample has been characterized by Auger electron spectroscopy. Whereas CO interacts rather weakly with CaO and adsorbs molecularly on pristine and oxygen vacancy sites, CO2 and NO interact strongly with CaO, leading to a variety of structures in TPD. CO2 adsorption leads to the formation of carbonates. NO adsorption results in the desorption of N2 and N2O species presumably by the formation of NO dimers stabilized at defect sites, which is consistent with prior reports. In assisting the assignment of the TPD structures that have been seen, density functional cluster calculations have been performed and are compared with results from prior studies. Trends seen for the initial adsorption probabilities, S0, are consistent with a simple mass-match model, S0CO > S0-CO2. The coverage dependence of the adsorption dynamics is dominated by precursor effects for CO2 but shows direct Langmuirian adsorption dynamics of CO.

1. Introduction CaO has been identified as a component of the plume in coal combustion plants.1,2 Therefore, it is interesting to characterize the adsorption kinetics and dynamics of typical combustion gases such as CO, NO, and CO2 on CaO surfaces. Furthermore, CaO can serve to study the acid-base properties of metal oxides3-5 as compared with, for example, MgO, with the former being in most cases catalytically more active.6,7 In simple terms, a higher activity of CaO is related with a lower Madelung potential, which leads to a more delocalized electron distribution of surface oxygen and hence to an eventually more efficient overlap with the orbitals of adsorbing molecules.7 In addition, oxygen vacancy sites are highly catalytically active,7 allowing for an efficient charge transfer to adsorbates, as is commonly assumed for metal oxides.8 A variety of applications for CaO are known, including catalytic gasification of Ca-containing coal,9,10 NO storage catalysts used for operating combustion engines at lean (oxygen rich) conditions with alkaline earth oxides as promising catalysts, the capture of SO2 (ref 9) and CO2 (ref 11) by CaO, the dimerization of, for example, methane,12 the reduction of NO by CO,13 and the decomposition of chlorinated hydrocarbons.14 For the decomposition of NO, CaO has been considered as a catalyst less expensive than systems based on noble metals.15 In the following, a brief literature survey is given respecting CO, NO, and CO2 adsorption on CaO as far as it is pertinent for the presented data. However, studies on CaO single crystals are rather rare, and the literature is dominated by catalysis studies conducted on powder samples; no molecular beam scattering data appear to be available for CaO. More extensive is the database for the related MgO (see, e.g., refs 16-19) and NiO (see, e.g., refs 20 and 21) alkaline earth oxides. CO adsorption has been studied on CaO powders by temperature programmed desorption (TPD) and Fourier trans* To whom correspondence should be [email protected]. Fax: 701.231.8831.

addressed.

E-mail:

form infrared spectroscopy (FTIR), indicating the importance of oxygen vacancy sites.22 A CO TPD peak at 400-500 K after adsorption at room temperature has been detected, but oxygen desorption has not. Isotope labeling indicates the exchange of surface oxygen with the adsorbate via a carbonate intermediate; however, CO decomposition has been ruled out. Thus, carbonate decomposition appears to involve the desorption of CO. Furthermore, the formation of adsorbed CO2-n species has been concluded in a number of FTIR studies.22-24 The CO/CaO system has also been considered theoretically by means of Hartree-Fock (HF) calculations. Accordingly, CO molecules adsorb upright (C atom down) on Ca sites.6 The literature concerning NO adsorption is more extensive even for CaO.5,15,25-29 On CaO powders, the desorption of N2, NO, and N2O after NO exposure suggests the exchange of surface oxygen.22,26,27 On reduced (defected) CaO an even larger variety of species has been identified by IR spectroscopy such as NO, NO2, NO3, and N2O2. However, according to an electron paramagnetic resonance (EPR) study on CaO powders, basic O2- surface sites alone appear to catalyze the formation of various NxOy species.5 Interestingly, molecularly adsorbed NO, as well as N2O and its decomposition in N2 and surface oxygen, were reported for iron oxide,30 SrTiO3,31 and MgO19 surfaces after NO exposure at low temperatures. Similarly, NO dimers and their decomposition in adsorbed N2O intermediates at low temperatures have been seen on a number of metal surfaces.15,27,29,32 N2O further decomposes in N2 and surface oxygen, with oxygen vacancies presumably stabilizing the N2O intermediate. Density functional theory (DFT) calculations predict NO adsorption on CaO terrace sites, with tilted NO adsorbed on oxygen sites with the N atom toward the surface.27 Possible adsorption/decomposition pathways have been studied theoretically, for example, in refs 15 and 29. CO2 adsorption leads, according to X-ray photoelectron spectroscopy (XPS) and X-ray absorption fine structure (NEXAFS) data, to the formation of carbonates (CO3ads). Accordingly CO2 chemisorbs on oxygen sites of the CaO(100) surface,

10.1021/jp800755q CCC: $40.75 © 2008 American Chemical Society Published on Web 04/18/2008

CO, NO, and CO2 on Reduced CaO(100) yielding carbonate-like species involving lattice oxygen.33 Similar conclusions have been drawn from ultraviolet photoelectron spectroscopy (UPS) and metastable impact electron spectroscopy (MIES) experiments on CaO thin films.34 CO2 adsorption on CaO powders leads to CO2 desorption temperatures as high as ∼1100 K, as the result of carbonate decomposition.9,35 In an early IR study the formation of bidentate and monodentate carbonates was also proposed on powder samples. Interestingly, isotope scrambling indicated that surface oxygen is involved in the CO3ads formation.36 The formation of carbonates has also been theoretically confirmed,4,7 although bidentate species have been objected.7 In this study, we present a thorough characterization of the adsorption kinetics (by TPD) and adsorption dynamics (by molecular beam scattering) of typical combustion gases (CO and CO2) on a CaO(100) single crystal. In addition, NO TPD data are included, as well as supplemental AES (Auger electron spectroscopy) scans and ab initio calculations based on the DFT. 2. Experimental Procedures The measurements have been collected using a home-built, triply differentially pumped molecular beam system.37 The supersonic beam is attached to a scattering chamber that contains a shielded mass spectrometer for TPD38 and adsorption probability measurements, as well as a combined AES/low-energy electron diffraction (LEED) system. The impact energy, Ei, of the molecules has been varied by seeding 3% of CO and CO2 in He combined with a variation of the nozzle temperature within 300-700 K. The measuring error in the initial adsorption probabilities, S0, amounts to (0.04. The coverage, Θ, dependent adsorption probability curves, S(Θ), have been smoothed while conserving the shape of the transients. Measurements have been taken for normal impact angles. The saturation coverage of the probe molecules has been set to 1 ML at low temperatures (as common) indicating saturation of the surface. Data collected at greater temperatures have been scaled accordingly by integrating the transients. NO adsorption has solely been studied by TPD. The exposures are given in Langmuir (1 L ) 1 × 10-6 mbar s). In the case of CO2, for some of the TPD experiments the gas has also been dosed on the sample with the molecular beam system, which maintains the background pressure while dosing CO2 below 2 × 10-8 mbar and entirely eliminates contributions from the sample holder. An exponential background has been subtracted from the TPD curve; a heating rate of 2 K/s has been used for all TPD measurements. The sample temperature could be reduced to ∼90 K for the less conducting metal oxides by bubbling He gas through a lN2 dewar. (For details of the TPD setup, see ref 38.) The reading of the thermocouple has been calibrated in situ ((5 K) by measuring the condensation temperature of alkanes; a nonlinear correction was required.39 We would like to note that thermally activated bond breaking has been seen for n/iso-butane; the details will be reported elsewhere. 3. Computational Procedure To assist the assignment of the TPD features (rather than as a quantitative modeling), DFT calculations have been performed to model CO and CO2 adsorption on CaO clusters using the Gaussian03 package.40 NO adsorption has already been studied extensively by a variety of computational techniques (see, e.g., refs 15, 27, and 28 and Table 1). In the simulations, stoichiometric CaxOx clusters (x ) 9, 12, or 16) were used considering various adsorption sites: fivecoordinated sites on a terrace of the regular (100) surface, four-

J. Phys. Chem. C, Vol. 112, No. 19, 2008 7391 coordinated sites on cluster edges, and three-coordinated corner sites (Figure 1). Ab initio simulations were performed using hybrid Barone’s modified Perdew-Wang 1991 exchange method41 with the Lee, Yang, and Parr correlation function including both local and nonlocal terms MPW3LYP.42 The 6-31G+(d) basis set was utilized for all carbon and oxygen atoms, while the 6-31G(d) basis set was employed for calcium atoms. In all geometry optimizations only the adsorbates (CO, CO2) were allowed to relax freely, while the position of cluster atoms were kept fixed. The value of the lattice constant of the CaxOx clusters was obtained by performing cluster optimization in a “breathing mode” in the absence of the adsorbate and ranged from 4.671 to 4.693 Å depending on the cluster size. Allowing the edge and corner atoms to relax would increase the binding energy of the adsorbates, which is, however, already larger for those defect sites than for the adsorption on terrace sites, as described below. Therefore, including a relaxation of the clusters would not change the results qualitatively but would significantly increase the computational time. The main object of the DFT calculations was to assist in assigning the TDS features. Adsorption energies for each configuration were calculated by

Eads ) E(CaxOx - Y) - E(CaxOx) - E(Y)

(1)

where Y denotes the adsorbate (CO, CO2). A positive value of Eads corresponds to exothermic adsorption. The basis set superposition error (BSSE) was calculated for each structure using the standard counterpoise procedure implemented in Gaussian03 with subsequent correction of adsorption energy values. 4. Possible Adsorption Configurations According to DFT Some of the most stable adsorption configurations on Ca9O9 clusters, as determined by DFT, are shown in Figure 1. Table 1 summarizes binding energies obtained in this study by DFT as compared with other theoretical results from the literature. Eads has been calculated by eq 1; Eads (BSSE) is the adsorption energy including the basis set superposition error, and Rads denotes the adsorbate-cluster binding distance (see Table 1). 4.1. CO Adsorption. Adsorption of CO on CaO(100) clusters via the C or O atom and on Ca or O lattice sites has been considered. While no stable adsorption configurations for CO were found on O sites or Ca corner sites, CO exothermically adsorbs on Ca terrace and edge sites of Ca9O9 clusters through the C atom (see parts I and II of Figure 1 and Table 1). The adsorption energy of CO on edge sites is greater than on the terrace sites, as expected. However, the absolute values of the binding energies are quite small (0.028-0.126 eV) and indicate weak adsorption. CO can also adsorb on the edge site via its oxygen atom with a binding energy of 0.030 eV. The relatively large bond length between the surface and CO of Rads ) 2.742.99 Å is also consistent with a weak adsorption. To investigate the effect of the cluster size on adsorption energies, simulations on larger clusters such as Ca12O12 and Ca16O16 have been conducted. The absolute value of binding energy increases with cluster size; for Ca12O12 it increases by 31% as compared with Ca9O9. However, the difference in CO adsorption energy for Ca16O16 and Ca12O12 clusters amounts to only 13%. In addition, the adsorption bond length is not affected by the cluster size. Therefore, we are confident that trends can correctly be extracted by the DFT conducted here. The most stable adsorption configurations for CO are Ca edge sites as well as Ca sites on terraces, where CO binds via the C atom. Although CO

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Kadossov and Burghaus

TABLE 1: Results of Quantum Chemical Calculations (See Section 4) as Compared with the Literaturea cluster CO2

Ca9O9

Ca12O12

Ca16O16 Ca19O19 Ca24O24 Ca30O30

TPD peak R10 R10 R9 R8 R10 R10 R9 R8 R9 R8 R10 R10 R9 R8

Ca5O5b [OCa5]8+

site C/O-edge C/O-corner C/O-terrace (0°) C/O-terrace (45°) C/O-edge C/O-corner C/O-terrace (0°) C/O-terrace (45°) C/O-terrace (0°) C/O-terrace (45°) C/O-edge C/O-corner C/O-terrace (0°) C/O-terrace (45°) C/O-terrace (45°)

Eads (eV)

Eads (BSSE) (eV)

Eads (BSSE) (kJ/mol)

2.26 1.75 0.30 0.23 1.91 1.84 0.51 0.47 0.66 0.53

2.14 1.63 0.14 0.043 1.79 1.72 0.35 0.32 0.51 0.38 2.27 2.30 1.04 0.62

206 157 13.5 4.1 173 166 33.8 30.9 49.2 36.7 219 222 100 60 87

1.426 1.396 1.545 1.502 1.434 1.391 1.522 1.485 1.517 1.479 1.39 1.36 1.43 1.40 1.49 1.23 1.26 1.384

2.7 12.2 -3.9 2.9 3.9 4.4

2.862 2.821 2.991 2.740 2.862 2.872

1.00 0.88

periodic CO

Ca9O9 Ca9O9 Ca12O12 Ca16O16 periodic

NO

R1 R2 R1 R1 R1

Ca5O5b Ca9O9b Ca13O13b Ca13O13b Ca10O10b Ca8O8b Ca11O11b Ca13O12b Ca13O12b

R4 R4 R4 R3

C/Ca-terrace C/Ca-edge O/Ca-terrace O/Ca-edge C/Ca-terrace C/Ca-terrace C/Ca O/Ca

0.098 0.172 0.006 0.063 0.112 0.120 0.55 -0.004

N/O-terrace (0°) N/O-terrace (45°) N/O-terrace (0°) N/O-terrace (45°) N/O-terrace (0°) N/O-terrace (45°) N/1O-terrace (dimer) N/2O-terrace (dimer) O/2Ca-terrace (dimer) N/O-edge O/F-center O/F+-center

0.028 0.126 -0.040 0.030 0.040 0.046

Rads (Å)

ref this study

7

75 HF4 MP2c 76 this study

6 0.54 0.57 0.48 0.49 0.58 0.62 0.73 0.39 0.0 1.37 3.67 1.67

52.1 55.0 46.3 47.2 56.0 59.8 70.4 37.6 0.0 132 354 161

1.586 1.572 1.609 1.617 1.517 1.497 2.105 1.697 2.487 2.646 1.446

27

a O/Ca terrace: bonding via O atom of the adsorbate (either CO or CO ) on Ca terrace sites of the cluster; C/O corner: bonding via C atom of 2 the adsorbate on the O corner sites of the cluster (1 eV ) 96.485 kJ/mol). b Embedded.

adsorption has not been studied extensively before by theoretical methods, these findings are in agreement with intuition and prior results, as summarized in Table 1. 4.2. CO2 Adsorption. The adsorption of CO2 has already been considered in more detail theoretically (see Table 1). However, to complement this project, DFT calculations are included. In contrast to CO, the most stable adsorption configuration of CO2 is via the C atom adsorbed on O sites of CaO (see parts III-VI of Figure 1). CO2 does not adsorb either via the O atom or on Ca sites. This result is fully consistent with a number of other theoretical studies (see Table 1) indicating carbonate formation. There are four possible adsorption sites on the Ca9O9 cluster used here: two on terrace sites, one on an edge site, and one on a corner site. The two terrace site adsorption configurations result from the fact that the CO2 plane can either be parallel to the Ca-O-Ca plane (part III of Figure 1) of the cluster or at an angle of 45° (part IV of Figure 1). In all cases, the bond length between the C atom in CO2 and O sites is 1.40-1.55 Å, which corresponds to a strong adsorption. The O-C-O angle of adsorbed CO2 amounts to ∼130°, indicating the formation of carbonates; this is in agreement with experimental data on powder samples (see section 1). Similar to CO, for Ca9O9 clusters the strongest CO2 adsorption was found on the edge site; on the terrace site, it was weaker. This decrease in adsorption energy is consistent with an elongation of the adsorbate-cluster bond length (see Table 1).

The cluster size effect was also studied for CO2 adsorption. The geometric and energetic parameters do not significantly change for adsorption of CO2 on edge and corner sites of Ca12O12 clusters. However, the effect of cluster size is larger concerning terrace sites. The binding energy increases from 0.14 to 0.35 eV for a “0°” configuration and from 0.043 to 0.3 eV for a “45°” configuration. Simulations employing Ca16O16 clusters produced 0.51 and 0.38 eV correspondingly. It is evident that for CO2 adsorption on terrace sites the absolute value of the binding energy increases with cluster size. Prior DFT calculations7 considering, for example, Ca30O30 clusters, revealed the same trend (Table 1). Although this limits a quantitative modeling, it is evident that two different adsorption sites for CO2 on terraces exist with distinctly different binding energies. This result assists in the assignment of TPD structures seen experimentally. We are not aware of detailed theoretical studies considering the effect of oxygen vacancy sites for CaO surfaces. 5. Data Presentation and Discussion 5.1. Sample Cleaning and Characterization. The CaO(100) surface (from Goodfellow) was prepared by cleaving in air and polishing, as described in prior studies.43 The sample (10 × 10 × 2 mm) was immersed in petroleum for storage (CaO is hygroscopic) and was rinsed in acetone and ethanol prior to mounting in the vacuum chamber.33 Part I of Figure 2(curve a)

CO, NO, and CO2 on Reduced CaO(100)

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Figure 1. Different adsorption configurations of CO and CO2 on CaxOx clusters: (I) CO on the terrace site of Ca9O9. (II) CO on the edge site of Ca9O9. (III) CO2 on the terrace site of Ca9O9 (0° mode). (IV) CO2 on the terrace site of Ca9O9 (45° mode). (V) CO2 on the edge site of Ca9O9. (VI) CO2 on the corner site of Ca12O12.

shows AES data of the as-received sample. Carbon contaminations are evident (see the AES line at 265 eV). After the annealing cycles (see part I of Figure 2; curve b, 5 min at 1000 K; curve c, 120 min at 1000 K) in vacuum, the carbon AES line intensity dropped below the detection limit, and only Ca (at 262, 280, 300, and 324 eV) and oxygen (at 519 eV) AES peaks remained. Note that the secondary electron background overlapped with the carbon AES region for curve b (part I of Figure 2). However, clearly carbon-containing impurities have finally been cleaned off, see curve c in part I of Figure 2. A shift of the AES peaks (by ∼20 eV) due to charging of the sample as compared with reference data44 was present. While annealing the samples initially a large increase in the H2O background pressure was detected (cf., ref 45). Interestingly, no adsorption of CO2 was evident by adsorption transients (CO2 pressure in the scattering chamber vs exposure time with the molecular beam system46) at 97 K for the asreceived sample (see part II of Figure 2, curve a). Note that the area above the transient and below the saturation level (indicated by the dashed lines in part II of Figure 2) corresponds to the number of adsorbed gas-phase species. After annealing the surface in vacuum, an uptake of CO2 was clearly seen (part II of Figure 2, curve c). We may note that the adsorption transient (curve c in part II of Figure 2) decreases slightly with exposure time (0-40 s range). This is a well know effect referred to as adsorbate-assisted adsorption. However, in the case of CaO, the adsorbate-assisted adsorption is so small that integrating the curves basically leads to a standard Kisliuk-like shape. It appears plausible that while reducing the water content (solved in the bulk of the crystal) by annealing the sample, a passivating Ca(OH)2 or CaCO3 layer formed on the surface (see also, e.g., ref 33). The catalytic activity regained toward, for example, CO2

Figure 2. Sample characterization by (I) AES (Auger electron spectroscopy, 1000 eV, Vpp ) 10 V) and (II) molecular beam scattering. Curve (a) as-received sample, (b) after 5 min annealing in vacuum at 1000 K, (c) after 120 min annealing at 1000 K.

adsorption after annealing shows that the passivating surface layer has been decomposed. No clear LEED pattern could be obtained. Although charging of the crystal affected the LEED/AES experiments, this result indicates a rather large density of defects (see section 5.2.1) present on the cleaved but ultrahigh vacuum (UHV) annealed CaO surface. Theoretical studies predict oxygen vacancy formation energies decreasing along the alkaline earth series.47 However, partially reduced metal oxides are often characterized by a larger catalytic activity48-51 and may be considered a good model system as compared with catalysis studies on powders. 5.2. Adsorption Kinetics. TPD data as a function of exposure, χ (as indicated), are shown in Figures 3-5. Although CO adsorption/desorption appears to obey a simple mechanism, the results obtained for CO2 and NO are quite complex. Therefore, for CO2 the effect of adsorption temperature has been characterized; in addition, the decomposition products were analyzed by multimass TPD for CO2 and NO. The experimental kinetic parameters are summarized in Table 2 and are compared with results obtained for related systems. We are not aware of experimental binding energies of CO, NO, and CO2 determined for CaO single crystals. 5.2.1. CO Adsorption. The kinetics of CO adsorption (Figure 3) was dominated by weakly adsorbed species desorbing at 124 K (R1 peak, at small exposures) and 175 K (R2 peak), which is consistent with the DFT calculations (Table 1) and suggests the adsorption of CO on pristine (R1 peak) and defect (R2 peak) sites. Similar CO TPD features have been seen for other metal oxides such as

7394 J. Phys. Chem. C, Vol. 112, No. 19, 2008

Figure 3. TPD (thermal programmed desorption) data for CO adsorption on CaO(100), as a function of exposure as indicated.

TiO2(110).52 Integrating the TPD curves led to an estimated defect density of 40%, assuming that each defect site anchors one CO molecule. However, it has been proposed that oxygen vacancy sites can stabilize up to three (oxygen) molecules.8 Thus, the defect density obtained from the TPD data is an upper limit. To the best of our knowledge, no scanning tunneling microscopy data are available for CaO surfaces, and the tools available for this project in our group do not allow for an atomistic characterization of defect sites. However, it is safe to assume that the defects are oxygen vacancy sites. No strong indications for a decomposition of CO were present. First, the TPD data were well reproducible. In contrast, decomposition would lead to site-blocking effects due to residual carbon, which were not observed. Second, AES scans collected after the adsorption/desorption cycles of CO did not indicate any change in the surface composition. In addition, no indications for carbonate formation was evident on the CaO(100) single crystal (by TPD) studied at UHV conditions after adsorption of CO, as expected theoretically (Table 1). The CO TPD peaks shifted by ∼30 K to lower desorption temperatures with increasing exposures. In contrast to NO (cf., part I of Figure 4) and as commonly observed for CO adsorption,52-56 repulsive lateral interactions was present. Note that CO condensation does not start until adsorption temperatures as low as ∼35 K were reached (see Table 2). A Redhead analysis yielded a heat of adsorption of 24-31 kJ/mol and 3845 kJ/mol for the R1 and R2 peaks, respectively, and assuming a first-order process (ν ) 1 × 1013/s). The experimental binding energies of CO for CaO were larger than those determined for MgO (see Table 2). This trend may be unexpected considering the mostly electrostatic interaction of those molecules with the surface, as pointed out initially in some theoretical studies. However, this trend is consistent with recent quantum chemical calculations (see Table 1) and may be caused by a more covalent character of the binding. Importantly, the proposal that surface defects may obscure the experimentally obtained binding energies has been objected in studies comparing MgO and NiO thin film data and results obtained on the corresponding cleaved surfaces.57 Similarly, defect structures for CaO in CO TPD appear well separated from those which can be assigned to the pristine surface sites (Figure 3). 5.2.2. NO Adsorption. Adsorption of NO at 100 K (part I of Figure 4) initially led to the detection of a TPD peak at 211 K (R4 peak), which shifted

Kadossov and Burghaus to 274 K with increasing exposure. The small exposures required to detect this feature strongly suggest assigning it to molecular NO adsorption/desorption in the monolayer range on terrace sites (Tables 1 and 2). The observed peak shift is consistent with attractive lateral interactions, as expected since NO cluster formation was evident in a number of prior studies (section 1). A Redhead analysis of the peak position yielded binding energies in the range of 54-71 kJ/mol (see Table 2) assuming a first-order pre-exponential factor of ν ) 1 × 1013/s. At slightly larger exposures (onset 4 L) but before saturation of the R4 peak, a structure at 112 K was detected (R3 peak), which shifted by less than 2 K to greater temperatures with increasing exposures. The low-temperature edges of these R3 peaks appeared to line up as would be expected for zerothorder desorption of the multilayer; however, condensation of NO has been observed at much lower temperatures of ∼55 K.58 Therefore, it appears plausible to assign the R3 structure to the decomposition of NO clusters such as NO dimers anchored by two oxygen sites on CaO terrace sites, as suggested by prior theoretical work (see Table 1). Dimer formation in the monolayer coverage range has also been seen experimentally on MgO and SrTiO3 by TPD and FTIR.19,31 The fact that the R3 peak was not detectable for the smallest exposures is consistent with NO cluster formation. Part II of Figure 4 depicts TPD of N2 (m/e ) 28) and N2O (m/e ) 44) for an exposure of 4 L NO as compared with a typical NO TPD curve. Note that N2O also has a fragment at m/e ) 30 (co-incident with the NO parent mass) but with a much smaller intensity. In addition, the N2O intensities detected are small as compared with the NO signal. Thus, the fragmentation of nitrous oxide in the mass spectrometer will not significantly affect the NO TPD curves. No O2 (m/e ) 32) or NO2 (m/ e ) 46) signals were detected. Desorption of these molecules uniquely revealed the decomposition of NO as an adsorption pathway competing with molecular NO adsorption/ desorption (R3/R4 peak). Note that the intensity ratios of the TPD peaks are not consistent with the fragmentation pattern of gas-phase NO (by the mass spectrometer) and not all desorption temperatures of the detected features coincide. Furthermore it is quite unlikely that the TPD curve detected for m/e ) 28 is obscured by background adsorption of CO in the UHV chamber. The presence of small amount of desorbing N2O (R7 peak at 200 K) as also seen in prior studies (see section 1), suggests that NO clusters form N2O as an intermediate,15 which, in turn, results in N2 desorption (R5 peak) when the intermediate decomposes. Desorption of N2O stabilized by oxygen vacancy sites could explain the detection of the R7 peak, i.e., the N2O intermediate partially decomposes in N2 but also desorbs intact. Interestingly, the position of the R3 (dimer decomposition) and R5 TPD (N2O decomposition) peaks were correlated, which is consistent with the proposed mechanism. (NO)2 dimers, as intermediates for the formation of N2O, have also been proposed by Goodman et al.19 for MgO, as well as desorption of intact N2O as a competing reaction pathway (see also ref 31). The R6 structure in N2 TPD detected at much greater temperatures can be assigned to the direct decomposition of NO in N2 and adsorbed oxygen, Olattice, which may be catalyzed by charge transfer from defect sites. The direct NO decomposition as a reaction pathway competing with N2O bond activation also has been concluded for metal surfaces32 as well as other metal oxides.31 The N2 product can result from either NO or N2O intermediates. The N-O bond order in NO is larger than in N2O (2.5 vs 1.5). Therefore, NO would decompose at higher temperature then N2O. Since the position of the R5 peak matches

CO, NO, and CO2 on Reduced CaO(100)

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TABLE 2: Experimental Binding Energiesa TPD peak

Tpeak smallflarge (K)

R3 R4 R5 R6 R7

55 75, 84, 100 45, 145, 225 114f112 211f274 113f112 238f212 205f191

NO

CO2

Θ0: 21.2 Θ0: 55.0 28.7f28.2 54.3f71.1

80

R8 R9 R10 R11 R12 CO

118f103 203f185 270f500 105f95 289f318 27 35, 29-30 33 60 80-100 45, 145 150f100

R1 R2

39.4f29.7 34.4f31.7 43.6 36 48 29.8f25.9 52.1f47.4 26.4f23.8 75.1f82.9

17 ( 2 20-40 Θ0: 28.9

124f95 175f148

26.9;17.0 32.2 f 41.4 30-36 31f24 45f38

K (5

kJ/mol (1.5

225

units error

Ed smallflarge (kJ/mol)

assignment sublimation temperature monolayer defects defects monolayer dimer monomer NO decomposition (see eq 2)

multilayer monolayer pristine sites defect sites pristine defects adsorption on C/O-terrace (45°) adsorption on C/O-terrace (0°) adsorption on C/O-edge (corner) CO3 decomposition sublimation temperature multiplayer bilayer terrace sites (1015/s) defects close to defects monolayer structure physisorption pristine sites defects sites adsorption on C/Ca-terrace adsorption on C/Ca-edge

surface

ref

MgO(100) NiO (100) CaO (100)

58 57 56 this study

MgO (100) ZnO(0001)

77 59 38, 51

TiO2(110) CaO (100)

38, 51 this study

58, 56, 78, 56 MgO (100)

67

NiO (100)

56

ZnO TiO2(110)

55, 79, 80 52 (1014/ s)

CaO (100)

this study

this study

a Heating rate β ) 2 K/s; pre-exponential factor ν ) 1 × 1013/s (first order) unless otherwise noted; Θ data obtained in the zero coverage limit; 0 “defects” refers to oxygen vacancy sites; smallflarge denotes parameters obtained at small/large exposure; the following refers to the DFT (cf., Table 1): O/Ca terrace, bonding via O atom of the adsorbate (either CO or CO2) on Ca terrace sites of the cluster; C/O corner: bonding via C atom of the adsorbate on O corner sites of the cluster).

the position of the R3 peak (corresponding to the decomposition of NO dimers), and all N2O intermediates desorb at lower temperature (peak R7), R6 has been assigned to direct NO decomposition. Therefore, another pathway for N2O formation is likely, namely, via Nads and NOads interaction. Obviously, a number of competing adsorption/reaction pathways must exist. The following scheme, eq 2, summarizes the proposed mechanism, which is consistent with theoretical predictions and experimental results obtained on other metal oxides.

The arrows are labeled with the corresponding TPD features. It is well established from studies on other metal oxides such as TiO2 that oxygen vacancy sites form while annealing the metal oxide and “heal out” while exposing oxygen.8,51 Thus, an equilibrium such as Olattice T Ovac is present, with Ovac denoting oxygen vacancy sites. Therefore, a thermal desorption of oxygen and/or a poisoning of the catalyst by oxygen atoms may not be expected, consistent with studies on powder samples.

5.2.3. CO2 Adsorption. (a) CoVerage-Dependent CO2 Adsorption Kinetics. In addition, the CO2 adsorption kinetics (part I of Figure 5) was more complex than that observed for CO adsorption on the same sample. Three distinct features have been detected when collecting TPD data as a function of exposure: a broad structure within the temperature range of 270-500 K (R10 peak), a peak shifting from 203 to 185 K with increasing exposure (R9 peak), and a low-temperature structure related with CO2 desorbing at about 110 K (R8 peak). Again, the R8/R9 peak positions shifting to lower temperatures with increasing exposure would be consistent with repulsive lateral interactions as seen before for CO2 adsorption on other metal oxides (e.g., ZnO,38 TiO2,51 MgO59). The position of the R10 peak was approximately independent of exposure. In addition, the CO2 TPD data shown in part II of Figure 5 were obtained while saturating the surface with the molecular beam system with CO2 at different adsorption temperatures, TS, as indicated. The same features seen before (cf. part I of Figure 5) were present at low adsorption temperatures (TS ) 105 K), i.e., the peaks R8, R9, and R10 were reproduced. Since CO2 was directly dosed on the surface with the molecular beam system, which guarantees a very low background pressure, we rule out that an artifact from the sample holder obscured the TPD data. Increasing the adsorption temperature (part II of Figure 5) simply quenched the adsorption/desorption pathways in the order of their binding energies (desorption temperatures) as expected.

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Kadossov and Burghaus

Figure 4. (I) TPD data for NO adsorption on CaO(100) as a function of exposure as indicated. (II) TPD of m/e ) 30 (NO), 28 (N2), and 44 (N2O) at 4 L of NO exposure.

Thus, at the greatest adsorption temperature of 220 K, only the high-temperature structure (R10 peak) remained. In addition, the position of the R8 and R9 peaks in part III of Figure 5 shift to higher temperatures with increasing adsorption temperature (decreasing coverage) consistent with the data shown in part I of Figure 5. However, the peak shifts are rather small. The R8 and R9 TPD peaks originated from weakly bonded CO2. Therefore, we assign the R8 and R9 peaks to CO2 adsorption on terrace sites. DFT indicates two different adsorption configurations (parts III and IV of Figure 1), which differ by their alignment with respect to the cluster, as described in section 4.2. Thus, the DFT calculations were important in assisting in the TDS peak assignment. The R10 peak is assigned to adsorption on defect sites, which is also qualitatively consistent with the DFT calculations (Table 2). Estimating the area of the R10 peak leads to a defect density of 40%, which is consistent with the value obtained from the CO TPD data. Distinct defect induced TPD structures have also been seen for CO2 adsorption on TiO2 (ref 49, 51, and 60) and ZnO (see ref 38). All CO2 TPD features are assigned to carbonate-like species (Figure 1) involving lattice oxygen, as proposed before in experimental and theoretical studies (section 1). To characterize the kinetics of the carbonate formation/ decomposition further, multimass TPD were collected. (b) Carbonate Decomposition. Part III of Figure 5 shows the result of multimass TPD as a function of exposure. The main panel shows desorption of a CO fragment after exposure of CO2 at low temperatures. A CO TPD structure at ∼105 K (R11 peak) and a broad feature at about ∼320 K (R12 peak) are evident. The inset depicts a comparison of CO2 TPD (thin line) and CO TPD (thick line) after exposure of 2 L of CO2. Again, the R8, R9, and R10 CO2 TPD peaks are

Figure 5. (I) TPD data for CO2 adsorption on CaO(100), as a function of exposure as indicated. (II) CO2 TPD as a function of adsorption temperature. The surface has been saturated with CO2 by means of the molecular beam system. (III) CO TPD as a function of CO2 exposure. The inset directly compares CO2 TPD with CO TPD data.

seen in addition to the CO TPD structures. Clearly, the shape of the CO TPD curves (part III of Figure 5) differs from those seen for CO2 (see, e.g., inset of part IIII of Figure 5). Therefore, the detected CO did not originate from the cracking of CO2 in the ionizer of the mass spectrometer. Note also that the m/e ) 28 to 44 intensity ratio detected while backfilling the scattering chamber with CO2 amounted to ∼10, while it equaled roughly one (see inset of part III of Figure 5) in the TPD runs. Furthermore, the areas of the CO and CO2 TPD peaks were approximately equal. According to information from the mass spectrometer vendor, the sensitivity factors for CO and CO2 amount to 20 × 10-5 and 13 × 10-5 A/mbar, respectively. Thus, the detected CO indicates either a decomposition of CO2 or carbonates while heating up a layer of CO2 adsorbed on CaO at low temperatures. Importantly, the desorption of oxygen was below the detection limit of the mass spectrometer and no carbon residuals were present in AES scans detected after CO2 adsorption/desorption

CO, NO, and CO2 on Reduced CaO(100)

J. Phys. Chem. C, Vol. 112, No. 19, 2008 7397

Figure 7. Initial adsorption probability, S0, vs impact energy, Ei, at constant adsorption temperature, Ts, for CO and CO2, as indicated. The curves have been parametrized as follows CO2 (Ts ) 95 K), S0 ) 0.46 exp(-Ei/1.17) + 0.03; CO2 (Ts ) 115 K), S0 ) 0.46 exp(-Ei/0.57) + 0.18; CO (Ts ) 95 K), S0 ) 0.70 exp(-Ei/0.79) + 0.20 with Ei in eV.

Figure 6. Typical examples of adsorption transients for (I) CO and (II) CO2.

cycles. Therefore, we rule out a decomposition of CO2 or CO3 into carbon and oxygen. This suggests the decomposition of carbonates into CO and lattice oxygen, as also seen in studies of CaO powders. The absence of desorbing oxygen could be related with the equilibrium of thermally formed oxygen vacancy sites and the healing of these vacancies by oxygen formed while the carbonates decomposed, as described above.48,61 Isotopelabeling experiments on CaO powder samples (see section 1) support our conclusion of the involvement of lattice oxygen in the carbonate formation.22 Tables 1 and 2 summarize the assignment of all TPD features detected in comparison with experimental and theoretical studies on related systems. 5.3. Adsorption Dynamics. The adsorption dynamics, i.e., the gas-surface energy transfer processes, were mapped by measuring adsorption probabilities as a function of impact energy and adsorption temperatures by means of the molecular beam scattering technique. The results are summarized in Figures 6-10. 5.3.1. Typical Examples. Typical examples of adsorption transients (for two different temperatures) are depicted in Figure 6, i.e., background pressure vs exposure time, t, curves are shown. Normalizing the background level to zero and the saturation level to one results in a plot of 1-S vs exposure time, with S denoting the adsorption probability. (The start of the transient has been set to t ) 0.) The ratio of the initial pressure rise and the saturation pressure equals the initial adsorption probability, S0. Integrating the transient leads to the coverage dependence of the adsorption probability, S(Θ). The area above the transient and below the saturation level equals the saturation coverage at the given measuring conditions. Interestingly, the shapes of the transients for CO (part I of Figure 6) and CO2 (part II of Figure 6) are very different. An almost perfectly linear

Figure 8. Initial adsorption probability as a function of surface temperature and parametric in impact energy (I) for CO and (II) for CO2. The CO data have been parametrized according to CO: S0 ) 0.90 exp(-Ei/313) + 0.10 with Ei in eV.

increase in the CO intensity (decrease in S) up to saturation of the surface is seen for CO, while in the case of CO2, S remains initially constant. Thus, the adsorption dynamics (gas-surface energy transfer processes) for CO and CO2 are quite different, as discussed in the following by considering the effect of surface temperature, Ts, and impact energy, Ei, on S0 and S(Θ). 5.3.2. Initial Adsorption Probabilities. (a) Effect of Impact Energy. At low temperatures, S0 of CO is systematically larger than for CO2 (see Figure 7). This result is consistent with a simple mass-match model.62,63 The better the mass match, the more

7398 J. Phys. Chem. C, Vol. 112, No. 19, 2008

Figure 9. Coverage dependence of the adsorption probability at low adsorption temperatures and parametric in impact energy for (I) CO and (II) CO2.

efficient the gas-surface energy transfer processes should be and the larger S0. By assumption of a simple rigid molecule model where the entire molecule interacts either with O sites or with Ca sites of the CaO surface, the CO2-to-O or CO2-to-Ca massmatch amounts to 2.75 and 1.1, respectively, while it equals 1.75 and 0.7 for CO. Quantum chemical calculations predict CO bonded with the C atom to Ca sites and CO2 bonded via the C atom to O sites as the energetically most stable equilibrium adsorption configurations (section 4). All other impact scenarios, such as CO molecules impinging first with the O atom on Ca sites, are sterically/energetically hindered. Therefore, we ran rule out the CO2-to-Ca combination and CO has the more favorable mass-match. Thus, S0-CO > S0-CO2 is predicted and indeed observed here (see Figure 7). Simple mass-match considerations have also been applied successfully for CO adsorption on the oxygen and Zn-terminated surface of ZnO.63,64 This simple model certainly does not always lead to correct predictions65,66 since the surface corrugation, defects, and effect of internal degrees of freedom are ignored. It appears that these effects do not significantly alter the adsorption dynamics of CO and CO2 on CaO(100). However, regarding a larger data set by including a number of different systems does not lead to a fully consistent picture. For example, a few other metal oxides have been studied by the molecular beam scattering technique: CO-MgO(100), S0 ) 0.47 (pure room temperature beam);67 CO2-TiO2(110), S0 ) 0.50 (Ei ) 0.11 eV); CO-TiO2(110), S0 ) 0.84 (Ei ) 0.047 eV); CO-ZnO(0001)-Zn, S0 ) 0.85 (Ei ) 0.047 eV); CO2-ZnO(0001), S0 ) 0.72 (Ei ) 0.11 eV) (see, e.g., Figure 3 in ref 51). A decrease in S0 with increasing Ei (Figure 7) is commonly observed for molecular adsorption,68,69 which is consistent with the conclusions drawn from the TPD data (section 5.2.2).

Kadossov and Burghaus

Figure 10. Coverage dependence of the adsorption probability at constant impact energy and parametric in surface temperature for (I) CO and (II) CO2.

Dissociative adsorption should lead to an increase in S0 above a given threshold energy (see, e.g., ref 70). The decrease in S0 simply results from less efficient energy transfer processes at large impact energies (shorter interaction time) and the requirement to dissipate more kinetic energy of the gas-phase species to the surface at large Ei. (b) Effect of Surface Temperature. The effect of surface temperature, Ts, on S0 is summarized in Figure 8. In the case of CO, S0 decreases roughly exponentially with Ts (part I of Figure 8). According to the CO TPD data (Figure 3), the majority of the CO molecules start to desorb at about 120 K. Thus, the decrease in the initial adsorption probability with Ts is the result of a temperature dependent adsorption/ desorption equilibrium and quantifies simply a kinetic (or net) adsorption probability, Skin, according to Skin ) FS0 - kdΘCO, rather than the dynamic adsorption probability, S0 (kd denotes the temperature-dependent desorption rate coefficient, F the CO flux, and ΘCO the CO coverage). The temperature dependence of the adsorption dynamics is more complicated for CO2 than for CO, as expected since carbonates are formed according to the TPD results shown in Figure 5 and many prior studies. S0 increases initially (section s1 in part II of Figure 8) and decreases afterward (section s2 in part II of Figure 8). The onset of the decrease in S0 at ∼140 K is too low to be entirely explained by an adsorption/desorption equilibrium of CO2. Therefore, we assign the decrease in S0 (section s2, part II of Figure 8) to the decomposition of carbonates (related with the R10 and R12 TPD peaks, Figure 5) rather than the thermal desorption of CO2. Thus, again, an effective or kinetic adsorption probability is detected. A decrease in Skin is expected in the

CO, NO, and CO2 on Reduced CaO(100) case of a decomposition pathway competing with the adsorption process. Similar effects have commonly been seen for the bond activation of alkanes on metal surfaces, see, e.g., ref 71. Thus, the adsorption dynamics (part II of Figure 8) and kinetics data (Figure 5) are consistent; both sets of experiments indicate the thermal decomposition of carbonates. The initial increase in the adsorption probability by ∼20% (section s1 in part II of Figure 8) could indicate the thermally activated formation of the more weakly bonded carbonates (related with the R8/R9 TPD peaks). 5.3.3. CoVerage-Dependent Adsorption Probabilities. Figure 9 summarizes the effect of impact energy, Ei, on the coverage, Θ, dependence of the adsorption probability, S(Θ), at low temperature; Figure 10 depicts the effect of adsorption temperature, Ts, on S(Θ) for constant Ei. (a) Effect of Impact Energy. Unexpectedly, S(Θ) of CO on CaO(100) follows a simple Langmuirian adsorption dynamics even at low impact energies (see parts I of Figures 9 and 6). Thus, S(Θ) ) S0(1 - Θ) is obeyed, indicating a negligible influence of the precursor state on the adsorption dynamics. A turnover from precursor mediated to direct adsorption dynamics has often been observed at large impact energies (see, e.g., ref 72). However, in the case considered here, even at thermal impact energies (i.e., with a pure room temperature beam of CO) a perfect linear dependence of S(Θ) on coverage is seen. To the best of our knowledge, for all other metal oxides studied so far by molecular beam scattering techniques, precursor-mediated adsorption was present. In contrast to CO adsorption, S(Θ) for CO2 on the same CaO(100) surface is dominated by the effect of precursor states (see part II of Figure 9) since S(Θ) remains approximately constant up to saturation and over the entire impact energy range. This type of curve shape (part II of Figure 9) is predicted by Kisliuktype73,74 precursor models. Although the lifetime of CO in a precursor state is somewhat limited by the lowest temperature accessible in this study, the most obvious difference between CO and CO2 adsorption on CaO(100) is the bonding type of these adsorbates. According to the TPD data and quantum chemical calculations, CO is weakly bonded (physisorbed), whereas CO2 interacts much more strongly with CaO(100). (b) Effect of Surface Temperature. The temperature dependence of S(Θ), see Figure 10, resembles the same effect as described above: strong precursor effects for CO2 adsorption (S(Θ) approximately independent of Θ, part II of Figure 10) and negligible precursor effects for CO adsorption (S decreases linearly with the number of free adsorption sites, part I of Figure 10). With increasing adsorption temperatures the saturation coverage naturally decreases and the shape of the S(Θ) curves scale accordingly. 6. Summary The following information has been collected for CO, NO, and CO2 adsorption on CaO(100) by means of kinetics experiments (temperature programmed desorption TPD) and molecular beam scattering, which maps the gas-surface energy transfer processes (adsorption dynamics). Assisting the assignment of TPD peaks, DFT (density functional theory) cluster calculations have been performed (see Table 1). • According to DFT, CO does not adsorb on O lattice sites. CO mainly adsorbs via the C atom on Ca sites. In contrast, the most stable adsorption configuration of CO2 is via the C atom adsorbed on O sites of CaO. CO2 does not adsorb either via the O atom or on Ca sites (see Figure 1, Table 1). • The as-received sample is chemically passivated, most likely by a Ca(OH)2 layer formed by water dissolved in the bulk which

J. Phys. Chem. C, Vol. 112, No. 19, 2008 7399 diffuses to the surface region when annealing the crystal. This layer can be thermally decomposed, which activates the catalyst (see part II of Figure 2). The sample cleaning induced a maximum density of defects of 40%, as estimated from the TPD data. • Whereas CO interacts weakly with CaO(100), NO and CO2 are strongly bound to the surface (see Figures 4 and 5). • No indications for CO dissociation have been obtained on the CaO(100) single crystal studied at UHV conditions. • Molecular desorption of NO (part I of Figure 4) as well as competing reactive adsorption pathways have been seen. • NO adsorption leads to desorption mainly of N2 and smaller amounts of N2O (part II of Figure 4). A mechanism including NO dimer formation has been proposed, which is consistent with related studies on other systems and DFT calculations. A number of competing adsorption/reaction pathways must exist. However, spectroscopic data on single crystals will be required to further verify the proposed mechanism. • Carbonates are formed readily when adsorbing CO2 on CaO(100), as is consistent with prior reports applying spectroscopic techniques mostly to powder samples (see Figure 5). • The carbonates decompose into CO (part III of Figure 5) as evident from multimass TPD. • Whereas CO and CO2 show TPD peak shifts consistent with repulsive lateral interactions, the NO-NO interaction is attractive (Figure 4). • The adsorption dynamics is rather uncommon as well. Whereas CO2 adsorption is dominated by precursor-mediated adsorption dynamics (as expected), CO adsorption follows simple and direct Langmuirian adsorption dynamics even at small impact energies (see Figure 9). This experimental fact may be related with the different binding types (weak vs strong adsorption) of these molecules on CaO(100). • The trends seen for the initial adsorption probabilities, S0, are consistent with predictions by a simple mass-match model. Accordingly, S0 of CO is systematically larger than S0 of CO2 (see Figure 7). • The structures seen in the S0(Ts) curves of CO2 could be related with the TPD results and carbonate formation. • Kinetics parameters have been determined and are summarized in Table 2 as well as DFT results, which are given in Table 1. • CaO is catalytically more active than, for example, the extensively studied MgO system, e.g., larger binding energies of small molecules are evident (Table 2). Acknowledgment. Assistance by S. Funk in the initial stages of the project is acknowledged. Financial support by the Department of Energy (DE-FG02-06ER46292; ND state grant, “coal combustion emission mechanism”) is acknowledged as well as the use of the NDSU computer facility to conduct the Gaussian03 calculations. References and Notes (1) Chen, Y.; Shah, N.; Huggins, F. E.; Huffman, G. P.; Dozier, A. J. Microsc. 2005, 217, 225-234. (2) Seames, W. S. Fuel Process.Technol. 2003, 81, 109-125. (3) Cox, P. A. Clarendon Press, Oxford, 1995. (4) Pacchioni, G.; Ricart, J. M.; Illas, F. J. Am. Chem. Soc. 1994, 116, 10152. (5) Paganini, M. C.; Chiesa, M.; Martino, P.; Giamello, E. J. Phys. Chem. B 2002, 106, 12531-12536. (6) Halim, W. S. A. Appl. Surf. Sci. 2007, 253, 8974-8980. (7) Jensen, M. B.; Pettersson, L. G.M.; Swang, O.; Olsbye, U. J. Phys. Chem. B 2005, 109, 16774-16781. (8) Henderson, M. A.; Epling, W. S.; Perkins, C. L.; Peden, C. H.; Diebold, U. J. Phys. Chem. B 1999, 103, 5328.

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