Toward a Realistic Description of NOx Storage in BaO: The Aspect of

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J. Phys. Chem. B 2005, 109, 9613-9621

9613

Toward a Realistic Description of NOx Storage in BaO: The Aspect of BaCO3 Peter Broqvist,† Itai Panas,‡ and Henrik Gro1 nbeck*,† Department of Applied Physics and Competence Centre for Catalysis, Chalmers UniVersity of Technology, SE-412 96 Go¨teborg, Sweden, and Department of EnVironmental Inorganic Chemistry, Chalmers UniVersity of Technology, SE-412 96 Go¨teborg, Sweden ReceiVed: January 21, 2005; In Final Form: March 17, 2005

NOx storage over hexagonal BaCO3(110) is investigated using first-principles calculations. Special focus is put on the importance of surface decarbonation. Upon decarbonation, supported BaO quasi-molecules are formed and a small drive toward (BaO)n cluster formation is predicted. Introduction of NO2 makes the decarbonation energetically relevant, while forming NO2-BaO-NO2 units, on the decarbonated surface. With this configuration, it is possible to replace all surface carbonates with nitrites and nitrates, forming a BaCO3 supported BaNO3NO2 overlayer. Thermodynamic considerations are employed to elaborate on the thermal stability of the formed NOx overlayers.

I. Introduction The need of more fuel efficient transportation has stimulated the introduction of lean burn gasoline and diesel engines. NOx storage and reduction (NSR) catalysts are used to remove harmful emissions from such engines.1,2 Given that the conventional three-way catalyst only operates efficiently at stoichiometric conditions, the NSR catalyst allows for net excess oxygen combustion conditions while still having the capacity to reduce NOx. The NSR concept is based on controlling the combustion conditions, i.e., whether oxidative or reducing. During lean (oxidizing) operation of the engine, NOx is temporarily stored in an absorbent, whereas at short fuel rich pulses, NOx is released and reduced to N2 over noble metals (Rh, Pd, Pt) by ordinary three-way catalysis technology. The platinum is not only active in the reduction of NOx, it is essential at lean conditions for the oxidation of NO to NO2 before storage.3,4 Typically, the storage component is an alkaline earth metal oxide, where barium presently is the metal of choice. Upon NO2 storage, Ba(NO3)2 is formed.5 In addition to the technological importance, NSR catalysts have offered several scientific challenges where the understanding of NOx chemistry on alkaline earth metal oxides, carbonates, and nitrates is one example. In fact, the composition of the storage material is ill-defined at lean conditions. The uncertainties arise because of the high CO2 concentration in the engine exhaust. Still, several experimental studies indicate that the storage compound is initially BaCO3.4,6-9 Estimates on relative stabilities of BaO, Ba(NO3)2 and BaCO3 have been based upon bulk thermodynamics.4,8,10 To put the present work in context, some of the considerations are repeated here. The relative stabilities of the relevant materials are calculated from

dG ) ∆H - T∆S + RT ln Q

(1)

dG is the Gibbs free energy whereas ∆H, T and ∆S are the enthalpy of formation, temperature and entropy, respectively. * Corresponding author. Electronic address: [email protected]. † Department of Applied Physics and Competence Centre for Catalysis. ‡ Department of Environmental Inorganic Chemistry.

Values for ∆H and ∆S are taken from ref 11. Q is the ratio of partial pressures for the products and reactants, and R is the molar gas constant. The reactions we concentrate on are

BaO + CO2 h BaCO3

(R1)

BaO + 3NO2 h Ba(NO3)2 + NO

(R2)

BaO + 2NO2 + 1/2O2 h Ba(NO3)2

(R3)

BaCO3 + 3NO2 h Ba(NO3)2 + NO + CO2

(R4)

BaCO3 + 2NO2 + 1/2O2 h Ba(NO3)2 + CO2

(R5)

Here, the last two reactions are the sum reactions for (R1) and (R2), and (R1) and (R3), respectively. In the upper panel of Figure 1a, dG for reactions R1-R3 is shown as a function of temperature. In the figure, stoichiometric conditions are assumed, which leads to the last term in eq 1 vanishing. For (R1), BaCO3 is formed from BaO at temperatures below ∼1200 °C. (R2) and (R3) in Figure 1a, show that Ba(NO3)2 is preferred over BaO and NO2 in the gas phase up to temperatures about 750 °C. The crossing between (R1) and (R2) or (R3) gives the transition from Ba(NO3)2 to BaCO3 at ∼610650 °C, depending on which reaction is used. The origin of the temperature induced transition from Ba(NO3)2 to BaCO3 is entropic. The transition from BaCO3 to BaO at very high temperature indicates that the storage material under a broad range of reaction conditions is BaCO3. Realistic engine conditions imply excess CO2 compared to NO2. This will have large effects on dG for the assumed reactions. The lower panel in Figure 1a shows dG for (R5) at constant gas composition (5% CO2, 8% O2, 600 ppm NO2 with inert fill-up gas to normal pressure) as a function of temperature. The assumed gas composition drives the reaction toward carbonate formation. Here, the transition from Ba(NO3)2 to BaCO3 occurs at ∼380 °C. By varying the partial pressure of CO2 in the gas, one may investigate the sensitivity of dG on changing RT ln Q for the outlined reactions. In Figure 1b, dG vs pCO2 for three different temperatures are given, illustrating

10.1021/jp050384l CCC: $30.25 © 2005 American Chemical Society Published on Web 04/09/2005

9614 J. Phys. Chem. B, Vol. 109, No. 19, 2005

Broqvist et al. The present study addresses the competition between CO2 and NO2 toward the intermediate BaO. We extend our previous studies by considering BaCO3 as initial composition of the storage material. The key processes of concern are decarbonation of the BaCO3 surface and the subsequent nitration. Density functional theory is employed to map out relevant regions on the potential energy surface and for electronic and structural characterizations. To evaluate the replacement reactions, where one CO2 molecule is replaced with two NO2 molecules, the entropy aspect, decisive to the relative thermal stability, is taken into consideration in a semiquantitative manner. II. Computational Method

Figure 1. (a) Upper panel: dG vs T for (R1)-(R3) at stoichiometric conditions. Lower panel: dG vs T for 5% CO2, 8% O2 and 600 ppm NO2 for R5. (b) dG vs pCO2 in % at three different temperatures for R5. Same O2 and NO2 concentrations as in (a), lower panel.

that very low CO2 concentration is enough to form bulk BaCO3 at temperatures above 400 °C. Thus, at typical working conditions (∼350 °C), there is a delicate balance between the bulk formation of Ba(NO3)2 and BaCO3, even at excess CO2. Experimentally, decarbonation of alumina supported BaCO3 has been observed at temperatures above 250 °C using temperature programmed desorption measurements.9 The effect of high CO2 concentrations in the gas phase on the NOx storage performance is not completely clear. On one hand, Cant and Patterson8 concludes that there is no effect of high CO2 concentrations on the NO2 uptake in a BaO/Al2O3 catalyst. On the other hand, refs 4, 6, 12, and 13, observe a clear effect of the NOx storage performance in the presence of CO2. Having said this, it is not clear to what degree the bulk thermodynamical limit is applicable for a working storage material. Erosive BaCO3 to Ba(NO3)2 interconversion at short time scales may be taken to exclude the relevance of the bulk limit. Previous theoretical work on NOx storage for NSR catalysts assume MeO (Me being alkaline earth metal) as the active storage material.14-22 This is reasonable as interconversion between BaCO3 and Ba(NO3)2 requires a BaO intermediate. The nature of this intermediate is, however, not clear, and clusters, embedded clusters and slab models have been employed in the above studies to model the active storage material. A central observation is the importance of pairwise NO2 adsorption.16 Thus, an additional energy gain was predicted upon adsorption of NO2 in a pairwise configuration utilizing both a surface barium and oxygen site owing to odd-even electron counting effects.

The calculations performed in this study utilize the density functional theory (DFT)23,24 in the implementation with planewaves and pseudopotentials (PWPP).25 The cutoff energy was 340 eV and the reciprocal space was sampled using a Monkhorst-Pack grid26,27 with a k-point spacing of 0.05 Å-1. Ultrasoft pseudopotentials as proposed by Vanderbilt have been used to describe the interaction between the core and the valence.28,29 We have used the gradient corrected approximation to the exchange-correlation potential as proposed by Perdew et al.30 (PBE). The calculations employ periodic boundary conditions. For the surface calculations, a four layer slab with two atomic layers kept frozen at theoretical bulk distances for representing the hexagonal bulk BaCO3(110) surface was used. The repeated slabs were separated by a 12 Å vacuum width to ensure that interslab interactions do not affect the energetic an structural results. The geometries were optimized and considered converged when the forces on the atoms were less than 3 × 10-2 eV/Å and the change in total energy was below 1 × 10-5 eV. To check the accuracy of our calculations, we calculated the lattice parameters, bond lengths, cohesive and binding energies for bulk BaO, BaCO3, Ba(NO3)2 and the molecules CO2, NO2 and O2. Generally, good agreements between experimental data and present results were obtained.31 III. Results and Discussion A. Bulk Properties. Thermal decomposition of bulk BaCO3 has been studied in the literature.32,33 During decarbonation, two phase transitions occur, orthorombic (whiterite) to hexagonal, and hexagonal to cubic.32 In this study, the hexagonal structure (resembles the structure of calcite) and the whiterite structure (orthorombic) for bulk BaCO3 has been considered. The calcination energies for BaCO3 f CO2 + BaO were calculated to 2.1 and 2.7 eV for the hexagonal and whiterite structures, respectively. The finding of whiterite being more stable is consistent with experiments, which has this phase as dominating for bulk BaCO3.32,33 The structural geometry optimized lattice parameters for the hexagonal phase is 5.4 × 5.4 × 18.8 Å3, whereas for the whiterite phase, 5.4 × 9.0 × 6.6 Å3 was calculated. Comparing the electronic density of states (DOS) for BaCO3 with that of BaO,34 one can see that the electronic states belonging to O(2p) are dominating the valence electronic band for both phases. B. BaCO3 Surface. CO2 desorption has been measured over a wide temperature range.9 This has been taken to reflect noncrystalline (peak at 250-600 °C) and crystalline (peak at 600-730 °C) BaCO3.9 In this study, the hexagonal structure is chosen as the model for supported BaCO3 aggregates owing to its intermediate stability and relevance in the decarbonation process of BaCO3.32 The surface investigated is the BaCO3(110) surface. The (110) facet is chosen, as it is charge neutral. The geometry optimized surface is displayed in Figure 2. Figure

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Figure 3. SDOS: (a) the BaCO3(110) surface (solid line) and PDOS for top surface carbonates (filled); (b) partially decarbonated BaCO3 surface (solid line) and PDOS for the supported BaO quasi-molecule (filled); (c) fully decarbonated BaCO3 surface (solid line) and PDOS for the two supported BaO quasi-molecules (filled); (d) (BaO)2 cluster formation on the BaCO3 surface (solid line) and PDOS for the supported Ba2O2 entity (filled).

Figure 2. Side and top view of the geometry optimized BaCO3(110) surface as cleaved from the hexagonal BaCO3 bulk structure with the used unit cell in 2 × 2 representation. Color code: white is carbon, green is barium, and red is oxygen (color on the web).

2 shows also the unit cell used unless something else is stated. Although, in the crystal, the planar carbonate ions have their normal direction parallel to the (110) plane, the geometry optimized surface structure has this normal rotated toward the (110) surface normal direction. The introduction of an interface induces perturbation on the second atomic layer to a lesser extent. The projected electronic density of states (PDOS) on the atomic pseudo-orbitals, when summed up for the two top layers, called here surface DOS (SDOS), plotted against the highest occupied Kohn-Sham eigenvalue is reported in Figure 3a. Also shown is the contribution to the SDOS from the top layer carbonates. These bands are dominated by O(2p) states. The top band belongs to carbonates in the upper layer, whereas the lower states belong to second layer carbonates. This is as expected upon cleavage of a bulk compound, where surface electronic states are higher in energy compared to bulk electronic states. We will discuss the electronic properties of the BaCO3 surface in more detail when considering decarbonation and NO2 adsorption. C. Decarbonation of BaCO3(110). We study transient decarbonation by removing CO2 molecules from the BaCO3-

(110) surface. Of course, the decarbonation of BaCO3 is energetically unfavored but will provide information on the composite BaO/BaCO3 substrate, which is a relevant intermediate in the NOx storage process. Consider a partially decarbonated surface obtained by the removal of one CO2 from the unit cell (cf. Figure 2). This step is associated with an energy penalty of 3.5 eV. The result is the formation of a quasi-molecular BaO entity supported on the BaCO3 surface; cf. Figure 4a. For the supported entity, the Ba-O bond length (ds) and the adiabatic BaO molecular desorption energy (BE) are 2.3 Å and 3.7 eV, respectively. For the gas-phase BaO molecule, dg and the dissociation energy (Deg) are computed to 1.9 Å and 5.9 eV, respectively. dg and Deg compare well with the experimental values of 1.94 Å and 5.79 eV, respectively.35 The cause for forming a perturbed BaO molecule when supported on the surface (ds ) 2.3 Å) as compared to the gas phase (dg ) 1.9 Å) is the competition between internal binding energy in the BaO entity and the adhesion to the support. Next, we consider a fully decarbonated surface. In this case, several possible arrangements of the remaining BaO phase upon CO2 abstraction are possible. One configuration comprises two BaO molecular units. The cost of removing the second CO2 is less than removing the first, 3.2 eV as compared to 3.5 eV, and the resulting structure is displayed in Figure 4b. The molecular desorption energy of the BaO molecules has decreased from 3.7 to 2.9 eV. The Ba-O bond distance has decreased from 2.3 to 2.0 Å, which is close to the value of the gas-phase BaO molecule. The results infer that the binding in the BaO quasimolecules is stronger than the adhesion to the support. It is conceivable that upon decarbonation, lattice mismatch between the BaO crystal and BaCO3 will either result in BaO quasimolecule formation, or, in segregation, forming (BaO)n aggregates. To test the segregation scenario, we first consider cluster formation in the gas phase. Here, the BaO molecule, dimer and tetramer are considered. The cohesive energies (Deg) in these clusters are given in Table 1. The (BaO)2 cluster is a fairly stable compound in gas-phase with a Deg of 7.3 eV compared to the

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Broqvist et al.

Figure 4. Geometry optimized decarbonated BaCO3(110) surfaces. (a) Partly decarbonated surface. To the left, top view showing the periodicity using a 2 × 2 unit cell. To the right, top and side view of the BaO quasi-molecule. (b) 2 × 2 unit cells for the fully decarbonated surface. To the right and left, side and top views of the formed BaO quasi-molecules, respectively. In the figures, the two bottom layers have been removed for clarity. Same color code as in Figure 2.

TABLE 1: Calculated Cohesion in the (BaO)n Clusters per BaO Unit in the Gas Phase (Deg) and in the Configuration When Supported (Des) on the Fully Decarbonated BaCO3 Surfacea 4‚BaO 2‚(BaO)2 1‚(BaO)4 BaO bulk

Deg [eV]

Desb [eV]

Bec [eV]

∆E

5.9 7.3 8.2 9.6

5.6 6.9 8.3d

2.9 3.1 2.9

0 0.0 -0.2

a Be is the adiabatic desorption energy. ∆E is the difference in total energy relative to the 4‚BaO system. The cohesive energy for BaO bulk is reported for reference. b Des is calculated in the gas phase using the constrained supported structure. c Be is calculated with respect to the fully relaxed gas-phase cluster. d The higher Des is due to the formation of a [(BaO)4]∞ chain. See Figure 5.

5.9 eV for the BaO molecule. The (BaO)4 cluster is even more stable (Deg ) 8.2 eV). These values should be compared to the theoretical bulk cohesive energy of 9.6 eV. Consider next the supported (BaO)n clusters, and the drive to cluster formation on the BaCO3 surface. The supported (BaO)2 cluster results in the atomic configuration displayed in Figure 5a. This configuration is nearly energetically degenerated with the case of having the two BaO molecules on the surface (∆E differ less than 0.02 eV); cf. Figure 4b. This is not accidental as the nearest neighbor ionic interactions are similar in the two cases. When supported, the cluster is distorted compared to its gas-phase structure. This indicates that the potential energy surface for the BaO cluster is rather flat, and thus, the cluster is very flexible toward deformations. Taking the (BaO)2 structure obtained when supported and comparing the energy of this with the optimized gas phase species yields only 0.4 eV gain (cf. Table 1). The adiabatic desorption energy for the (BaO)2 cluster is comparable to the case with two BaO molecules.

Figure 5. BaO cluster formation on the fully decarbonated BaCO3(110) surfaces: (a) formation of a rhombic (BaO)2 cluster; (b) formation of a polymer chain consisting of (BaO)4 clusters. In the upper panels, the periodicity. In the lower panels, the supported BaO clusters. In the figures, the two bottom layers have been removed for clarity. Same color code as in Figure 2.

When the supported (BaO)4 cluster is studied, the unit cell is doubled in one direction (creating a 1 × 2 unit cell). The surface morphology makes it difficult to form the cluster on the surface where the only configuration we find is to have one BaO molecule as “anchor” owing to the large mismatch between the BaO cluster and the BaCO3 surface. The studied configuration is depicted in Figure 5b. The clusters are connected across unit cells, thus forming a chain of (BaO)4 units. The calculations for this configuration shows that the cluster growth starts to be favored compared to having BaO molecules spread out over the surface. However, the energy gain is relatively small (∆E ) -0.2 eV) and seems to originate from the formation of a (BaO)4 chain. This is seen in Table 1 where Des > Deg for (BaO)4 compared to having a (BaO)4 chain in gas phase. So, it is not clear if isolated (BaO)4 clusters without the formation of chains are favorable. Having calculated the energetics for decarbonation of the BaCO3 surface layer, we address the electronic aspects upon BaO molecule and cluster formation. The SDOS for the two uppermost layers of atoms in the BaCO3 surface as well as for the different decarbonated surfaces are shown in Figure 3a-d. The occupied bands consist mainly of O(2p) states. Moreover, in Figure 3a, contributions from the first layer carbonates is explicitly shown whereas, in Figure 3b-d, the contributions from the (BaO)n entities, resulting upon CO2 abstraction, are reported (shaded). Upon decarbonation, oxygen states belonging to the BaO entity become the highest occupied states (compare Figure 3a,b). Interestingly, three oxygen 2p states are close to degenerated, indicating a quasi-spherical ion. Considering the fully decarbonated surface with two BaO quasi-molecules, two distinguished peaks are seen close to the highest occupied state (cf. Figure 3c). These peaks originate from the two BaO entities, here with bond lengths close to the gas phase BaO molecule, leading to a symmetry breaking in the O(2p) states compared to the near spherically symmetric oxygen ion in the case of partially decarbonated surface. For the supported (BaO)2 unit, an additional splitting is observed, compare parts c and d of Figure 3. This is owing to the inequivalence of the two BaO

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Figure 7. SDOS and PDOS in case of NOx adsorption: (a) NO2 adsorption on BaCO3 (solid line) and PDOS for NO2 (filled); (b) nitrate adsorption on the BaCO3 surface (solid line) and PDOS for NO3(filled); (c) NO2 adsorption on the partially decarbonated surface forming a nitrite (solid line) and PDOS for NO2- (filled); (d) NO32formation on the partially decarbonated BaCO3 (solid line) and PDOS for NO32- (filled).

Figure 6. NO2 adsorption on the (a) BaCO3(110) surface and (b) the partly decarbonated BaCO3 surface. In the upper panels, the top view of the surface with periodicity. In the lower panels, the side view. In the figures, the two bottom layers have been removed for clarity. (c) and (d) show spin density isosurfaces for the adsorption configurations given in (a) and (b), respectively. The isosurface is plotted in transparent yellow at 0.03 e-/Å. Color code: white is carbon, green is barium, red is oxygen, and blue is nitrogen (color on the web).

units forming the rhombic cluster; cf. Figure 5a. Note that upon decarbonation and (BaO)n formation, the band-gap is decreasing, which supports the interpretation of the BaO entity as quasimolecular. D. NO2 Adsorption over BaCO3(110). In this section, we address the interactions between NO2 and the BaCO3(110) surface at different degrees of decarbonation. Moreover, as adsorption of CO2 is also possible, this adsorption to a Ba2+ site at the stoichiometric BaCO3 surface will be discussed. 1. Stoichiometric Surface. On the stoichiometric BaCO3 surface, the stable adsorption site for NO2 is in a bridged configuration over two barium ions. The adsorption configuration is displayed in Figure 6a. The adsorption energy is calculated to 0.7 eV. The molecule is adsorbed in a bridging configuration where the two oxygen atoms are coordinated to the barium cations. On the basis of the structure (rN-O ) 1.26 Å and V(O-N-O) ) 122°), the adsorbed NO2 is charged; however, no true nitrite is formed. The same picture emerges from analyzing the spin density (Fv-FV). The electronic spin is delocalized over oxygen atoms in the surface carbonates as well as on the NO2 molecule, cf. Figure 6c, indicating partial charge transfer. Also, integration of the charge density in the direction normal to the surface reveals that ∼0.5 e- has been subtracted from the surface. The SDOS for NO2 adsorption is displayed in Figure 7a. The shaded area corresponds to the PDOS of the NO2 molecule. The highest occupied molecular orbital (HOMO) for the NO2 molecule is located at the same energy as the O(2p) states of the surface. This is consistent with the partial electron

transfer of approximately 0.5 e- to the NO2 molecule. Interestingly, the lowest unoccupied molecular orbital of π* character belonging to the NO2 molecule appears in the BaCO3 band gap. As the initial capture of NO2 and the following nitrite formation is important for the storage mechanism,16 we have investigated the coverage dependence on the NO2 adsorption energy. The adsorption configuration discussed implies a coverage of 0.5 with respect to available surface barium ions. Two other coverages were studied, namely 0.25 and 1. The 0.25 coverage calculations were performed in the larger unit cell. The adsorption energies calculated for the different coverages are 1.0, 0.7, and 0.1 eV for the quarter, half and full coverage, respectively. As expected, the NO2 adsorption energy is strongly dependent on coverage. The coverage dependent adsorption energy is due to both static repulsion between adsorbed nitrites and the energy cost for electron subtraction from the surface. The magnitudes of the computed adsorption energies are similar to what has been reported for NO2 adsorption over BaO(100)16,20-22 at similar coverages. The NO2 adsorption configuration is the same for the half and the quarter coverages. At full coverage, NO2 is forced out from the optimal bridging configuration as it only utilizes one barium ion for adsorption. In the high coverage limit, a disproportionation reaction between the adsorbed NO2 molecules, forming nitrate and NO in the form of either a nitrosyl ion or NO(g) is possible.

BaCO3 + 2NO2 f BaCO3NO3-NO+

(R6)

BaCO3 + 2NO2 f BaCO3NO3 + NO(g)

(R7)

The exothermicity of (R6) is 1.3 eV/2NO2. This energy should be compared to the adsorption energy calculated for NO2 at the full coverage limit (0.1 eV) as well as at the low coverage limit (1.0 eV). This means that at full coverage, the disproportionation reaction according to (R6) is strongly favored. However, the coverage dependence of (R6) is probably minor, due to the sufficiency of the local ion pair adsorption configuration. Therefore this product should become relevant at coverages above ∼0.5. Experimentally, NO(g) is observed upon

9618 J. Phys. Chem. B, Vol. 109, No. 19, 2005 NO2 storage.3,16 It could be speculated that the nitrosyl formed in (R6) can be one source of NO(g) formation according to (R7). However, the exothermicity of (R7) is 0.2 eV/2NO2 at full coverage. This energy should be compared to the adsorption energy of NO2 at half coverage (0.7 eV). Consequently, this reaction is not the channel for the NO(g) formation observed in experiments. The different degree of electron abstraction from the surface between NO2 and NO3 is seen in the SDOS and PDOS reported in Figure 7a,b. The two features close to the HOMO level for NO2 in Figure 7a are merged in the case of NO3 (cf. Figure 7b). For NO2, this spin polarization splitting was associated with partial charge transfer, forming a NO2δ-. The stronger electronegativity of NO3 results in close to full electron subtraction from the substrate, and thus, the spin polarization splitting is decreased. Integrating the electron density normal to the surface reveals a charge transfer of ∼0.9 e-. 2. Decarbonated Surface. As discussed above, surface BaO quasi-molecules or clusters are formed upon decarbonation of the BaCO3 surface. As no significant energy gains were found upon cluster formation, we consider NO2 adsorption on supported BaO quasi-molecules. Adsorption of NO2 on BaO clusters has been studied previously.17 An enhancement of the pairwise adsorption was observed as compared to the BaO(100) surface. The larger energy gain is due to the intrinsic instability of small oxide clusters. Two configurations of the adsorbed NO2 molecule on the partially decarbonated surface were investigated: over the oxygen ion or over a barium ion associated with the BaO quasimolecule. The former is the stable adsorption configuration with an adsorption energy of 2.6 eV. The adsorption configuration is displayed in Figure 6b. Upon the adsorption, a slightly strained NO32- species is formed. This assignment is based on the N-O bond lengths together with a tilted angle between O2N-O; cf. Figure 6b. The bond length (N-O) for a gas-phase NO32- is 1.37 Å.36 Furthermore, upon NO32- formation, the Ba2+ ion moves back to the original position in the lattice. The spin density is reported in Figure 6d. NO2 has one unpaired electron and upon formation of the NO32- species, the spin is localized over the entire NO32- complex with no electronic spin found in the surface carbonates. Possible formation of strained NO32species on BaO has been discussed in the literature.16,17,21 The second (metastable) adsorption configuration is similar to the NO2 adsorption on the BaCO3(110) surface. However, in this case, a true nitrite is formed with a bond angle of 116°. The adsorption energy for this species is 1.1 eV. Thus, the adsorption energy is higher than that on BaCO3 prior to decarbonation at the same coverage. The reason for the higher binding energy is that the electron in this case is abstracted from the more electropositive subsystem of BaO quasi-molecules. SDOS of the different adsorption configurations of NO2 on the partially decarbonated surface are shown in Figure 7c,d. Figure 7c,d corresponds to nitrite and NO32- formation on the decarbonated surface, respectively. The PDOS for the true nitrite formation over the decarbonated surface show differences compared to NO2 adsorption over the stoichiometric BaCO3 surface (Figure 7a). In particular, the spin polarized splitting has vanished. In contrast to the nitrite and nitrate formation, for the NO32- formation, the highest occupied state in the DOS is located solely on the NO32- species (cf. Figure 7d). Having effectively replaced one carbonate with an NO32species, a second NO2 molecule is introduced in the unit cell to investigate the effect of pairwise adsorption on the decarbonated BaCO3 surface. In this case, both a barium and an

Broqvist et al.

Figure 8. Pairwise NO2 adsorption on (a) the partially decarbonated surface and (b) the fully decarbonated surface. To the right and left, side and top views, respectively. In (a), lines are inserted for 3D visualization. In (b), the periodicity is shown to illustrate the formation of a complete NO2-BaO-NO2 overlayer. Same color code as in Figure 6.

oxygen site are utilized in the adsorption. Adsorption of the second NO2 molecule on adjacent barium sites to the NO32species leads to a large energy gain. The adsorption configuration is given in Figure 8a. The adsorption energy for two NO2 molecules on the partly decarbonated surface is 5.1 eV. This is considerably higher than previously reported for pair formation on BaO(100)16,21 (by ∼2-3 eV). It is also higher than pair formation on a (BaO)9 cluster, by ∼1 eV. Thus, BaO quasimolecule formation in the carbonate matrix will lead to an enhanced stability of Ba(NO3)2 complexes formed upon NO2 storage. As discussed above, the decarbonation of the BaCO3 surface is associated with a high energy cost (∼3.5 eV), which is much more than what is obtained for single NO2 adsorption on the BaCO3 surface. Here, the formation of a NO2-BaONO2 pair is 1.6 eV more stable than having the partially decarbonated surface. To check if it is possible to complete the replacement reaction and terminate the BaCO3 slab with nitrites and nitrates, we also calculated the binding energy of four NO2 molecules on the fully decarbonated surface. In this calculation, the relaxed structure resembles the structure of a 2D-BaNO3NO2 film, where nitrites also penetrate into the surface (cf. Figure 8b). Interestingly, this configuration where the nitrates has replaced the carbonates, matches the structure of BaCO3; i.e., the carbonatenitrate replacement is close to epitaxial. Of course, the adsorption energy per NO2 pair is less than for pairwise NO2 adsorption on the partly decarbonated surface, but still ∼2.2 eV more stable than having the surface terminated with carbonates. The formation of a NO2-BaO-NO2 monolayer may in fact prevent the formation of bulk Ba(NO3)2. E. CO2 Adsorption on the Clean BaCO3 Surface. As discussed above, there is a competition between NO2 and CO2 on the decarbonated surface, where two NO2 molecules are needed to replace one CO2. However, as the adsorption energy for NO2 on the stoichiometric BaCO3 surface is low, at least at high coverage, there could also be a competition between NO2 and CO2 adsorption at Ba2+ sites. For CO2, only the 0.5 coverage was investigated. The adsorption energy is calculated to 0.3 eV. The reason for the lower adsorption energy is associated with the lower electron negativity of CO2 compared to NO2 and the

NOx Storage in BaO

Figure 9. Proposed surface decarbonation scenario in the presence of NO2: (1) stoichiometric BaCO3 surface; (2) adsorption of NO2 over barium sites; (3) formation of a NO32- species replacing a carbonate in the lattice; (4) pairwise adsorption of NO2, forming a BaCO3 supported NO2-BaO-NO2 complex; (5) formation of a complete NO2-BaO-NO2 overlayer. The energy difference (dE) is with respect to the stoichiometric BaCO3 surface (1) and NO2 in gas phase.

fact that the stoichiometric BaCO3 surface is only capable to interact strongly via ionic interactions. F. Emerging Scenario. This section concerns the aspect of BaCO3 on NO2 storage. The proposed NO2 induced surface decarbonation following the minimum energy path is shown schematically in Figure 9. Decarbonation of BaCO3 before nitration (not included the figure) is associated with a 3.5 eV energy cost. Introducing NO2 reduces the energy penalty to 0.9 eV by the simultaneous formation of a NO32- species. Adsorption of a second NO2, producing an NO2-BaO-NO2 pair, is more stable than the surface carbonate. Further stabilization is obtained by completing the NO2-BaO-NO2 monolayer, whereby full surface decarbonation is achieved. Previous studies have shown the pairwise storage configurations to be critical for BaO.16 Here, we demonstrate that the pair mechanism may drive the decarbonation of BaCO3, which is a necessary prerequisite for NSR catalysts working in excess CO2 conditions. IV. Thermodynamics and Kinetics of NOx Storage over BaCO3 Surfaces In the previous section, the adsorptions of NO2 and CO2 over a BaCO3 surface have been described. The reaction energetics for CO2 replacement with NO2, forming nitrites and nitrates on the surface were also studied. In the Introduction, we elaborated on the well-known role of entropy and gaseous environment for thermodynamics of bulk reactions. Here, similar considerations are applied for the surface reactions. First, we investigate the temperature dependence on the initial adsorption of NO2 and CO2 on the BaCO3(110) surface. The standard procedure for studying characteristics of adsorption is to use a mean-field approach.37 Here, we study the temperature

J. Phys. Chem. B, Vol. 109, No. 19, 2005 9619

Figure 10. (a) NO2 and CO2 equilibrium coverages vs temperature when only allowing adsorption and desorption. (b) Gibbs free energy for full (FC) and partial (PC) replacement of surface carbonates with nitrite-nitrate pairs. Two cases are investigated: stoichiometric conditions and the case with 5% CO2 and 600 ppm NO2. Full line corresponds to the lowest value of Gibbs free energy according to eq 1 and 2, respectively. Dashed lines correspond to the continuation of the line in case of partially and fully decarbonated surfaces.

dependence of the adsorption-desorption equilibrium for NO2 and CO2 in a monolith catalyst. The catalyst is assumed to consist of BaCO3(110) as active phase for the adsorption. A gas flow of 5% CO2 and 600 ppm NO2 together with an inert add-up gas at normal pressure is assumed. Kinetic gas theory is used to calculate preexponential factors for the applied adsorption.38 Transition state theory is used for the desorption processes, assuming that the adsorbed state and the transition state for desorption have the same entropy. The adsorption is treated unactivated whereas for the desorption process we use the calculated adsorption energies. For CO2, the adsorption energy is fairly small and is therefore taken to be independent of coverage. For NO2, however, a linear coverage dependence (E(Θ) ) E0 + E1(Θ - 1)) was used, as supported by the results obtained in the previous section. This means that the adsorption energy for NO2 will be in the range of 1.2-0.1 eV for the complete range of coverages. The adsorption-desorption reactions considered are the following:

BaCO3(110) + CO2 h BaCO3(110) - CO2

(R8)

BaCO3(110) + NO2 h BaCO3(110) - NO2

(R9)

This model does not include surface reactions between adsorbed NO2 molecules. Rather, it reveals wether the calculated adsorption energies for NO2 are high enough to allow accumulation of nitrites on the surface at elevated temperatures, and high CO2 pressures. The result for the equilibrium coverages of NO2 and CO2 vs temperature are reported in Figure 10a. Significant coverages of NO2 exist at elevated temperatures. This is important for the replacement reaction to occur.

9620 J. Phys. Chem. B, Vol. 109, No. 19, 2005

Broqvist et al.

However, it is doubtful if the disproportionation reaction will occur, due to the low coverages at higher temperature in conjunction with an expected substantial activation barrier associated with this reaction. To increase the storage capability at higher temperatures, the very stable pairwise adsorption configurations have to be considered. The residence time and the coverage of adsorbed NO2 molecules in the form of nitrites on the BaCO3 surface will be crucial for the formation of surface nitrates. The replacement reaction between one CO2 and two NO2 will most likely lead to a large change in entropy (∆S). Consequently, the replacement reaction will display a large temperature dependence. Values of ∆S for adsorption on surfaces are uncertain. The value for bulk formation is known and corresponds to the maximum ∆S value, removing all translational and rotational degrees of freedom. This change in entropy is denoted ∆Sloc. Surface species, however, display higher mobility compared to bulk species. The lower limit of ∆S corresponds to having the adsorbates as a 2-dimensional gas (∆S2D). Consequently, the upper and lower limit of ∆S can be estimated if the change in entropy for the bulk material is considered to be constant (cf. Appendix A for details). The replacement reactions considered are

S + 2NO2 h S - NO2BaONO3 + CO2

(R10)

S + 4NO2 h S - (NO2BaONO3)2 + 2CO2

(R11)

and

S is the BaCO3 slab and (R10) and (R11) corresponds to nitritenitrate pair formation on the partially and fully decarbonated surfaces, respectively. For NO2, two different adsorption configurations exist. One is the nitrate, which is treated as a localized species. The nitrite, however, is treated as mobile (∆S2D). For the backward reactions of (R10) and (R11), we use ∆Sloc for the CO2 adsorption, thus assuming the formation of localized carbonate species in the same manner as the nitrate formation. We have used the same approximation for both the partially and the fully decarbonated surface. The entropy change (∆Sr), replacing one carbonate with two NO2 molecules becomes NO2 NO2 2 ∆Sr ) ∆SCO loc - ∆Sloc - ∆S2D

(2)

With this approximation for the entropy contribution in eq 1, we compute the thermal stability of the nitrite-nitrate pairs at different coverages according to (R10) or (R11). Two cases have been investigated, either reaction at stoichiometric conditions, or, assuming 5% CO2 and 600 ppm NO2 in gas phase. The results are presented in Figure 10b, where Gibbs free energy is plotted as function of temperature. At lower temperatures (T < 150 °C), the full coverage (100%) of nitrite-nitrate pairs is obtained as the most stable phase. Increasing the temperature, the partially (50% coverage) replaced phase becomes the preferred. At temperatures corresponding to dG > 0, the backward reactions of (R10) and (R11) is preferred. The introduction of CO2 has a striking effect on the transition temperature, which is shifted by ∼450 °C. These trends in stability ordering with temperature should be regarded as qualitative, given the large simplifications made estimating ∆Sr as well as the various approximations built into the model. However, the effect of CO2 to reduce the transition temperature is unambiguous. V. Conclusions This study has employed first principles methods to address critical aspects of the NO2 storage property of a NOx storage

and reduction catalyst. It has been particularly concerned with (a) NO2 adsorption on the stoichiometric BaCO3(110) surface, (b) the nature of the decarbonated surface, (c) how adsorption of NO2 affects the decarbonation, and (d) how NO2 is incorporated in this NOx storage material. The NO2 interaction with the stoichiometric BaCO3(110) surface results in the formation of adsorbed NO2δ- species. The adsorption energy is found to correlate linearly with the surface coverage. At the high coverage limit, disproportionation of two NO2 molecules forming nitrate and nitrosyl ions becomes energetically favored. This is, however, not believed to be the main source of NO(g) observed experimentally upon NO2 storage.3 CO2 adsorption competes with NO2 for the Ba2+ sites at the stoichiometric BaCO3 surface but is found to bind moderately. To check whether the calculated adsorption energies are sufficient for accumulation of NO2 on the surface, a mean field model was constructed that allowed solely for adsorption of NO2 and CO2. This model predicts a significant NO2 coverage even at temperatures above 400 °C. However, this coverage is not sufficient for describing the storage capacity of NSR catalysts. For this, incorporation of NO2 in the BaCO3 surface, as a result of CO2 desorption, is needed. Decarbonation in inert atmosphere implies the formation of supported BaO quasi-molecules at the BaCO3 surface. The drive for (BaO)n cluster formation was tested by allowing for supported (BaO)2 and (BaO)4 clusters to form. The supported (BaO)2 cluster displays the same stability as having isolated BaO quasi-molecules at the surface. For the (BaO)4 cluster, intercluster interactions due to (BaO)4 chain formation give a small energy gain compared to having the isolated BaO quasimolecules on the surface. This ambiguity is one of the main results of the present study. It demonstrates the structural complexity of the partly decarbonated BaCO3 surface, and this particular issue must be taken into account for approaching a complete understanding of an NSR catalyst under realistic working conditions. The presence of NO2 is predicted to promote BaCO3 decarbonation. The possibility of forming an NO32- species significantly reduces the energy cost for CO2 desorption. To obtain an exothermic decarbonation energy, pairwise adsorption of NO2 molecules is needed, forming NO2-BaO-NO2 complexes. Similar structures have shown to be important in the case of NO2 storage over BaO.16 This pairwise adsorption allows for complete surface decarbonation, forming a BaNO3NO2 overlayer, which in fact may prevent bulk nitrate formation. The thermal stability of the NO2 storage related compounds were investigated at either stoichiometric conditions or in excess CO2. Excess CO2 conditions unambiguously reduces the thermal stability of BaNO3NO2 overlayers. The storage of NO2 over a BaO(100) surface16 and on (BaO)9 clusters has been studied by us previously.17 Here, one step further was taken toward a mechanistic understanding of NOx storage under combustion conditions. This was done by considering the corresponding carbonated substrate, i.e., BaCO3. Although additional complications regarding decarbonation and the nature of surface BaO have been unraveled and shown to be essential for the performance of an NSR catalyst, it is gratifying to note how the storage mechanism proposed in refs 16 and 17 applies to the NO2 storage on BaCO3. Acknowledgment. We thank Erik Fridell, Magnus Skoglundh and Sture Nordholm for valuable discussions.The Competence Centre for Catalysis is hosted by Chalmers University of Technology and financially supported by the

NOx Storage in BaO

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Swedish Energy Agency and the member companies AB Volvo, SAAB Automobile Powertrain AB, Johnson Matthey CSD, Perstorp AB, Akzo Nobel Catalyst, AVL-MTC AB and the Swedish Space Corporation. A. Appendix: Calculation of ∆S In section 4, calculations of ∆S were performed. The calculations were performed by assuming the molecule goes from the gas phase to either a complete local, ∆Sloc, or mobile in two dimensions (2D), ∆S2D, adsorbed species. The entropy change in the two cases were calculated according to the following equations, neglecting the entropy contribution from vibration:

∆Sloc ) Str,3D + Srot,3D

(A1)

∆S2D ) Str,3D + Srot,3D - (Str,2D + Srot,2D)

(A2)

In the calculations performed here, we have assumed Srot,3D ) Srot,2D. The entropy S, is calculated from

S ) R ln(Q) + RT

d(ln(Q)) dT

(A3)

where Q is the partition function, in this case

Q ) qtr,iqrot,i

(A4)

The partition functions qtr,i are calculated from

qtr,3D )

(2πmkbT)3/2 V h3

(A5)

qtr,2D )

(2πmkbT) Asite h2

(A6)

The partition functions for rotation qrot,i are calculated from

qrot,3D )

8π2IxkbT σrh2

(linear)

(A7)

and

qrot,3D )

8π2x8π2IxIyIz(kbT)3/2 σrh3

(nonlinear)

(A8)

where σr is the rotational number for the molecule in question (2 for NO2 and CO2 displaying C2V and D∞h symmetry, respectively). Ii is the moment of inertia obtained from ref 43. References and Notes (1) Takahashi, N.; Shinjoh, H.; Iijima, T.; Suzuki, T.; Yamazaki, K.; Yokota, K.; Suzuki, N.; Miyoshi, N.; Matsumoto, S.; Tanizawa, T.; Tanaka, T.; Kasahara, K. Catal. Today 1996, 27, 63. (2) Bo¨gner, W.; Kra¨mer, M.; Krutzcsh, B.; Pischinger, S.; Voigtla¨nder, D.; Wenninger, G.; Wirbeleit, F.; Brogan, M.; Brisley, R.; Webster, D. E. Appl. Catal. B 1995, 7, 153. (3) Fridell, E.; Persson, H.; Westerberg, B.; Olsson, L.; Skoglundh, M. Catal. Lett. 2000, 66, 71.

(4) Rodrigues, F.; Juste, L.; Potvin, C.; Tempe´re, J.; Blanchard, G.; Dje´ga-Mariadassou, G. Catal. Lett. 2001, 72 (1-2), 59. (5) Westerberg, B.; Fridell, E. J. Mol. Catal. A: Chem. 2001, 165 (12), 249. (6) Balcon, S.; Potvin, C.; Salin, L.; Tempe´re, J.; Dje´ga-Mariadassou, G. Catal. Lett. 1999, 60, 39. (7) Lietti, L.; Forzatti, P.; Nova, I.; Tronconi, E. J. Catal. 2001, 204, 175. (8) Cant, N.; Patterson, M. J. Catal. Today 2002, 73, 271. (9) Centi, G.; Arena, G. E.; Perathoner, S. J. Catal. 2003, 216, 443. (10) Amberntsson, A.; Persson, H.; Engstro¨m, P.; Kasemo, B. Appl. Catal. B: EnViron. 2001, 31, 27. (11) Lide, D. R., Ed. Handbook of Chemistry and Physics, 71st ed.; CRC Press, Inc.: Boca Raton, FL, 1990-1991. (12) Epling, W. S.; Campbell, G. C.; Parks, J. E. Catal. Lett. 2003, 90 (1-2), 45. (13) Epling, W. S.; Parks, J. E.; Campbell, G. C.; Yezerets, A.; Currier, N. W.; Campbell, L. E. Catal. Today 2004, 96, 21. (14) Schneider, W.; Li, J.; Hass, K. J. Phys. Chem. B 2001, 105, 6972. (15) Schneider, W.; Hass, K.; Miletic, M.; Gland, J. J. Phys. Chem. B 2002, 106, 7405. (16) Broqvist, P.; Panas, I.; Fridell, E.; Persson, H. J. Phys. Chem. B 2002, 106, 137. (17) Broqvist, P.; Gro¨nbeck, H.; Fridell, E.; Panas, I. J. Phys. Chem. B 2004, 108, 3523. (18) Broqvist, P.; Gro¨nbeck, H.; Fridell, E.; Panas, I. Catal. Today 2004, 96, 71. (19) Valentin, C. D.; Figini, A.; Pacchioni, G. Surf. Sci. 2004, 556 (23), 145. (20) Branda, M.; Valentin, C.; Pacchioni, G. J. Phys. Chem. B 2004, 108, 4752. (21) Schneider, W. F. J. Phys. Chem. B 2004, 108, 273. (22) Karlsen, E.; Nygren, M.; Petterson, L. G. M. J. Phys. Chem. B 2004, 107, 7795. (23) Hohenberg, P.; Kohn, W. Phys. ReV. 1964, 136, 864. (24) Kohn, W.; Sham, L. J. Phys. ReV. 1965, 140, A1133. (25) Payne, M. C.; Teter, M. P.; Allan, D. C.; Arias, T. A.; Joannopoulus, J. D. ReV. Mod. Phys. 1992, 64 (4), 1045. (26) Monkhorst, H.; Pack, J. Phys. ReV. B 1976, 13, 5188. (27) Pack, J.; Monkhorst, H. Phys. ReV. B 1977, 16, 1748. (28) Vanderbilt, D. Phys. ReV. B 1990, 41, 7892. (29) We have used the ultrasoft pseudopotentials as distributed together with the CASTEP code from Accelrys Inc. For the diferent atoms, the number of electrons treated explicitly in the valence shell are Ba(10), O(6), C(4) and N(5). (30) Perdew, J.; Burke, K.; Ernzerhof, M. Phys. ReV. Lett. 1996, 77, 3865. (31) The cohesive energy and lattice parameter for bulk BaO are calculated to be 9.76 eV (10.10 eV) and 5.59 Å (5.50 Å), respectively (experimental data in parentheses39). The reaction enthalpies for BaO + CO2 f BaCO3 and BaO + 2NO2 + 1/2O2 f Ba(NO3)2 are calculated to -2.7 eV (-2.7 eV) and -5.2 eV (-5.2 eV), respectively (tabular value11). The bond angle and bond length in NO2 are calculated to 134° (134°) and 1.21 Å (1.20 Å), respectively (experimental data from ref 40). The bond length in CO2 is calculated to be 1.17 Å (experimental 1.16 Å40). (32) Lander, J. J. J. Am. Chem. Soc.1951, 73, 5794. (33) Arvanitidis, I.; Sichen, D.; Seetharaman, S. Metallurgical Mater. Trans. B 1996, 27 (3), 409. (34) Broqvist, P.; Gro¨nbeck, H.; Panas, I. Surf. Sci. 2004, 554, 262. (35) Huber, K. P.; Herzberg, G. Molecular spectra and molecular structure: IV. Constants of diatomic molecules; Van Nostrand Reinhold Co.: New York, 1979. (36) The calculations were made employing the GGA-PBE and a double numerical basis set with polarization (DNP) as implemented in Dmol3.41,42 (37) Dumesic, J.; Rudel, D.; Aparico, L.; Rekoske, J.; Trevin˜o, A. The Micro-kinetics of Heterogeneous Catalysis; ACS Professional reference book; American Chemical Society: Washington, DC, 1993. (38) The preexponential factors are calculated using kinetic gas theory by calculating the collision coefficient at normal pressure. The area per site is taken as one-fourth of the unit cell (1.47 × 10-19 m2). The number of sites are assumed to be 0.18 mole/kgcat and the sticking coefficient is assumed to be 1 for both CO2 and NO2. (39) Ko¨nigstein, M.; Catlow, C. J. Solid State Chem. 1998, 140, 103. (40) Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements; Pergamon Press: Oxford, U.K., 1984. (41) Delley, B. J. Chem. Phys. 1990, 92, 508. (42) Delley, B. J. Chem. Phys. 2000, 113, 7756. (43) Knox, D.; Dadyburjor, D. Chem. Eng. Comun. 1981, 11, 99.