New Investigations on the Surface Reactivity of Layered Lithium Oxides

Article pubs.acs.org/JPCC

New Investigations on the Surface Reactivity of Layered Lithium Oxides Nathalie Andreu, Isabelle Baraille,* Hervé Martinez, Rémi Dedryvère, Michel Loudet, and Danielle Gonbeau IPREM/ECP, UMR5254, Université de Pau, Hélioparc Pau Pyrénées, 2 av. Pierre Angot, 64053 Pau cedex 9, France ABSTRACT: Density Functional Theory is applied to understand the large difference in surface reactivity of LiCoO2 due to Al/Co substitution which is experimentally observed. In this way, we explore the SO2 and CO2 adsorption modes on the (110) surface of LiCoO2 and α-LiAlO2. For SO2 adsorption, chemisorption produces sulfite species (for LiCoO2 and α-LiAlO2) and sulfate species (in the case of LiCoO2). We demonstrate that the modification of the surface reactivity when Co3+ ions are substituted by Al3+ ions is due to a change from an adsorption mode controlled by redox properties for LiCoO2 to a less energetically favorable adsorption mode controlled by acid−base properties for α-LiAlO2. For CO2 adsorption, the formation of carbonate species is observed for both compounds, illustrating the fundamental difference in the factors controlling SO2 adsorption compared to CO2 adsorption.

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INTRODUCTION LiCoO2 is one of the most widely used positive electrode material of today’s lithium-ion batteries. Its theoretical capacity is 272 mAh·g−1, but only ∼140 mAh·g−1 reversible capacity is reached in practical cells. This value corresponds to the deintercalation of half of the Li+ ions from LiCoO2 to Li0.5CoO2 upon charging the battery up to a 4.2 V cutoff voltage. To obtain higher capacities with LiCoO2, the cells have to be charged at higher voltages than 4.2 V to remove more Li+ ions from the structure. However, increasing the charge cutoff voltage can lead to structural degradation of LiCoO2 and dissolution of cobalt in the electrolyte.1 Both phenomena result in an increased capacity fading upon cycling. In recent years, it was shown that modifying the surface of LiCoO2 particles by application of a metal oxide or phosphate coating (especially by aluminum-based coating as Al2O3 or AlPO4) can significantly improve the capacity retention upon cycling at high voltages without loss of the initial reversible capacity. However, although it is admitted that these coatings act as a protection for the electrode against cobalt dissolution in the electrolyte at high potential, the exact mechanisms are not totally understood. The appearance of a LiCo1‑xAlxO2 solid solution between the material and the Al-based coating has been proposed and then evidenced by X-ray photoelectron spectroscopy (XPS).2,3 Compounds of this solid solution have been the subject of structural and electrochemical studies.4,5 Although they show lower reversible capacities due to the presence of electrochemically inactive Al3+ ions, they show higher lithium intercalation potentials than LiCoO2 and are effective to limit cobalt dissolution at 4.5 V (vs Li+/Li).4 Since dissolution phenomena are linked to interactions at the electrolyte/electrode interface, in a previous paper we have explored the surface reactivity of these materials as compared to LiCoO2.5 In this way, we carried out adsorption experiments of © 2012 American Chemical Society

gaseous probes (NH3 or SO2) followed by XPS analyses, which allow the identification of active sites (except the weakest ones) and a quantitative determination of their concentration. The results obtained after sulfur dioxide adsorption evidence that the Al substitution modifies the surface reactivity of LiCoO2. Indeed, only sulfate species are identified for LiCoO2, whereas both sulfate and sulfite species are characterized for the LiCo1‑xAlxO2 solid solutions (x = 0.5).5 In the present paper, in order to better understand the difference in surface reactivity due to Al/Co substitution, density functional theory (DFT) calculations are used to explore the thermodynamically favorable SO2 adsorption modes on LiCoO2 and α-LiAlO2. The study is extended to the adsorption of CO2, a reference gaseous adsorbate to probe the surface basicity of oxide materials. LiCoO2 and α-LiAlO2 crystallize in a layered structure with alternating Li, O, Co/Al, and O planes along the [001] direction. Among the most probable surfaces of LiCoO2,6 the (001) surface is polar (type 3 according to Tasker7), whereas the (110) surface, which cleaves the crystal perpendicular to the (001) planes, is a nonpolar surface belonging to Tasker type 17 (Figure 1). The (110) surface with equivalent layers composed of all three elements in a stoichiometric ratio was selected as a first simple model to simulate and analyze the adsorption of gaseous molecules (SO2, CO2) on these materials. This surface is actually present in significant proportion in LiCoO2 crystals as shown by SEM images (Figure 1b). To our knowledge, this is the first theoretical approach on the surface reactivity of these layered lithium oxides and an important step is the understanding of the systematics of their interactions with different adsorbates. Received: May 3, 2012 Revised: August 22, 2012 Published: August 22, 2012 20332

dx.doi.org/10.1021/jp304293u | J. Phys. Chem. C 2012, 116, 20332−20341

The Journal of Physical Chemistry C

Article

the positions of all the atoms were optimized. The translation vector normal to the surface was set to 15 Å and thus the vacuum between successive slabs is larger than 10 Å in thickness. The single unit cell (1 × 1) (as = 4.952 Å, bs = 5.716 Å) is reported in Figure 2a.

Figure 2. Top view of the unit cells/adsorption modes on (110) surface of LiMO2(a) (1 × 1) unit cell, (b) (2 × 2) supercell (low coverage), and (c) 2 × 1 supercell (intermediate coverage).

Adsorption of SO2 and CO2 was simulated on (110) surface of LiCoO2 and α-LiAlO2 using a four-layer slab. Different approaches were considered with two coverages of the surface with a (2 × 2) supercell and a (2 × 1) supercell (Figure 2b,c). In the first case (low coverage−LC) the adsorbed molecules are ∼9 Å (as direction) and ∼11 Å (bs direction) apart from each other, a value large enough to validate the approximation of the “isolated molecule limit”. In the second case (intermediate coverage−IC), the adsorbed molecules are ∼9 Å (as direction) and ∼6 Å (bs direction) apart from each other. Adsorption energies are calculated as follows:

Figure 1. (a) Schematic representation of the bulk, (001) and (110) surfaces of LiMO2 (b) SEM (Scanning Electron Microscopy) images of LiCoO26.

We start by a comparison of the ideal bulk optimized for LiCoO2 and α-LiAlO2. Several ab initio quantum chemical studies have been carried out on LiCoO26,8−10 but, to our knowledge, such an approach does not exist for α-LiAlO2. Then, we briefly present the results on LiCoO2 and α-LiAlO2 (110) surfaces and we focus on the adsorption of SO2 and CO2 on these layered lithium oxides. We explore the preferential adsorption sites, discuss the nature of adsorption interactions, and compare the factors that control SO2 and CO2 adsorption on LiCoO2 and α-LiAlO2.

Eads = Esurface + SO2 /CO2 − Esurface − ESO2 /CO2

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where Esurface+SO2/CO2 represents the energy of the super system (surface + adsorbate), Esurface the energy of the (110) surface, and ESO2/CO2 the energy of the isolated molecule. The ESO2/CO2 energy was calculated with the VASP program considering SO2 (or CO2) in a 10 × 10 × 10 Å3 box. The volume of the cell was kept fixed and the positions of all of the atoms were optimized. For the super system, the translation vector normal to the surface was 20 Å and thus the vacuum distance between successive slabs was larger than 10 Å. By this convention, a negative value of Eads indicates an exothermic adsorption. It is also noteworthy that when the geometry of the adsorbed species was close to a sulfate, we have modified, as proposed in a previous work,16 the spin polarization of the whole system in order to account for the transfer of two electrons from the molecule to the surface; this improves the SO2 adsorption energy in the case of LiCoO2 at low and intermediate coverages. The charge distribution on the atoms was investigated using Bader’s topological analysis.17,18 In this approach, atomic charges are calculated using the decomposition of electronic charge density into atomic contributions by dividing the space into atomic regions with surfaces at a minimum in the charge density.

COMPUTATIONAL DETAILS DFT calculations were performed with the periodic supercell plane-wave basis approach implemented in the Vienna Ab initio Simulation Package (VASP).10,11 The interactions between the valence and the frozen core electrons were simulated with the projector augmented wave (PAW) method,12 and electron exchange and correlation were treated within the functional proposed by Perdew and Wang, 13,14 using (nonlocal) Generalized Gradient Approximation (GGA). Plane waves were included to a cutoff of 400 eV. A grid of 38 k-points yielded well-converged bulk results. Spin-polarized calculations were performed, and the rotationally invariant formulation of the on-site Hubbard-U model given by Dudarev et al.15 was used for LiCoO2 (the on-site correction U−J was set to be 4 eV for Co 3d electrons in order to obtain a reasonable value for the band gap). The optimization of the ideal bulk, with hexagonal structure without defects was performed for both materials. For the (110) surface, an eight-layer slab was built having equal surfaces on both sides and containing a stoichiometric ratio of constituents. The volume of the cell was kept fixed and 20333



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RESULTS AND DISCUSSION

LiCoO2 − α-LiAlO2Bulk Properties. LiCoO2 (high temperature form) crystallizes in the trigonal system (R3m̅ space group, no. 166) with the α-NaFeO2-type structure. This structure, named O3, can be represented as an ordered rocksalttype with an ABCABC stacking of oxygen planes, with the Li+ and Co3+ ions ordered in alternate layers of octahedral sites of the (111) planes.19 α-LiAlO2 has the same crystal structure as LiCoO2.20 The conventional hexagonal cell used for space group R3m̅ (no. 166) was adopted in this study, the (111) cubic plane of the ordered rocksalt becoming (001). In this configuration, the structure of LiMO2 is characterized by the lattice constants ahex and chex, representing, respectively, the M− M intra and interlayer (chex/3) distances, and by zox, representing the position of oxygen atom along chex axis. Metal, lithium, and oxygen atoms occupy the Wyckoff crystallographic positions 3a (0,0,0), 3b (0,0,0.5), and 6c (0,0,zox), respectively. As several theoretical studies have been carried out on LiCoO2 but none in the case of α-LiAlO2, we started by the optimization of the ideal bulk for these two materials. The lattice parameters of the relaxed bulk LiCoO2 and α-LiAlO2 are reported in Table 1. They agree with the experimental trends when going from LiCoO2 to α-LiAlO2 (evolution of c parameter, M−O and Li−O bond lengths). Table 1. Optimized Lattice Parameters of LiMO2 Compared to Experimental Values α-LiAlO2

LiCoO2 ahex (Å) chex (Å) zox M−O Li−O

exp.19

optimized bulk

exp.20

optimized bulk

2.815 14.050 0.260 1.912 2.102

2.863 14.073 0.256 1.978 2.078

2.810 14.152 0.263 1.903 2.119

2.826 14.346 0.262 1.928 2.126

Figure 3. Total density of states (full line) and projections onto M, O, and Li states, respectively, represented as blue, gray, and dark areas (a) for LiCoO2 and (b) for α-LiAlO2.

energy range of the upper valence band of LiCoO2 (a narrow A band and a broad band B, C, D) and α-LiAlO2 (a broad band with two maxima A and B) are well reproduced. To complete this study, we have examined the charge distribution (Bader charges) in the relaxed structures. The results reveal significant differences in the iono-covalence of M−O bonds between LiCoO2 and α-LiAlO2, reflected by very different charges on atoms, (LiCoO2: qO= −1.06 e, qCo= 1.26 e; α-LiAlO2: qO= −1.67 e, qAl= 2.48 e) . Bare Surfaces of LiCoO2 and α-LiAlO2. Our previous work on LiCoO26 suggests that (001) and (110) surfaces should be present in LiCoO2 crystallites. The (001) surface is a polar surface, classified as type 3 according to Tasker, alternating anion and cation layers leading to a finite dipole moment per repetitive unit. Different mechanisms including reconstruction of the surface, faceting...have been proposed for the stabilization of such polar surfaces as reviewed by Noguera.28 Unfortunately, to our knowledge, no experimental information on such surface reconstruction is available for crystalline LiCoO 2 and the phase space of possible reconstructions is too huge to be screened by theoretical calculations alone. On the contrary, the (110) surface is a nonpolar surface, classified as type 1 according to Tasker. Each layer is composed of all of the elements in a stoichiometric ratio leading to zero

LiCoO2 can be described as a wide gap semiconductor, the 2.0 eV computed band gap agrees with the experimental values reported in the literature (2.1 eV21 and 2.7 ± 0.3 eV22). In contrast, α-LiAlO2 can be described as an insulator, as revealed by the 6 eV computed band gap but, to our knowledge, no experimental data exist. For LiCoO2, we observe a convergence to a non-magnetic state. This is consistent with previous experiments and calculations, which show that Co3+ in LiCoO2 is in a lowspin state S = 0.22−25 The computed DOS (density of electronic states) for LiCoO2 and α-LiAlO2 (Figure 3) reveal significant differences. The valence band of LiCoO2 is dominated by Co and O states, and the conduction band is dominated by Co states. The gap is between t2 g and eg Co states with some mixing with O2p states in the t2 g part. These results are consistent with previous calculations on LiCoO2.6,8,26 For α-LiAlO2, the valence and the conduction band correspond to a mixing of Al3p/Al3s and O2p states with more important contribution of oxygen states in the valence part and of Al states in the conduction part. The XPS valence spectrum we have obtained for LiCoO2 and α-LiAlO2 are compared with the total density of states modulated by the photoionization cross sections (tabulated by Scofield27) (Figure 4). A very good agreement is observed and the main characteristics of the experimental spectra in the 20334



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In the model of the four-layer slab selected, the relaxation of the surface was taken into account by optimizing all atomic positions with constant as and bs unit cell parameters (LiCoO2: as = 4.95 Å and bs = 5.72 Å; α-LiAlO2: as = 5.05 Å and bs = 5.79 Å). After full relaxation, modifications are observed (by reference with the initial atomic positions in the calculated bulk), with the same trends for LiCoO2 and α-LiAlO2 but with more important changes for α-LiAlO2. The first layer of atoms L1 is moved toward the second layer L2, and this contraction of the first two layers distance is compensated by an increase of the distance between the second and third layers (d [L2 −L3]) for both compounds (Figure 5). It is noteworthy that, with

Figure 5. Schematic view of the (110) surface in ball and stick representation (red, oxygen atoms; blue, metal atoms; and gray, lithium atoms)geometric parameters of relaxed slabs of LiCoO2 and α-LiAlO2 (δ represents the evolution in Å by reference to initial bulk values).

respect to bulk terminated surfaces, the oxygen atoms move slightly outside the plane in both cases and the cations slightly inward (Li more than Co in LiCoO2 and Al more than Li in αLiAlO2). The Bader charge analysis shows some differences for charges at the (110) surfaces compared to the bulk materials, but the same evolutions are observed with much more negative oxygen atoms for α-LiAlO2 (qO = −1.63 e) than for LiCoO2 (qO = −1.09 e). In addition, the results obtained for LiCoO2 evidence an important spin polarization (number of unpaired electrons, nα − nβ = 2.9) for the undercoordinated cobalt atoms at the surface classically associated to a high spin (HS) polarization; cobalt atoms inside the slab keep no spin polarization. This highlights the coexistence for this compound of cobalt atoms in a low spin configuration inside the surface as in the bulk and in a high spin configuration at the top layer of the (110) surface. LiCoO2−α-LiAlO2 −SO2, CO2 Adsorption. Our goal is to improve the fundamental knowledge of the surface reactivity of LiCoO2 and α-LiAlO2 through the simulated adsorption of SO2 and CO2. The challenge is to understand the factors controlling the adsorption and the reactivity of the moieties on the surface. Lewis acids such as SO2 and CO2 can be expected to bond to oxide surfaces at the anion sites and thus they are often considered as probe molecules testing the basic sites of the surfaces. SO2 can also be oxidized or reduced and CO2 can also be reduced. Only a few theoretical studies have been devoted to the SO2 adsorption on metal oxides. Density functional theory (DFT) calculations on alkaline earth oxides find SO2 to adsorb as a simple Lewis acid, adsorption energies increasing uniformly down the series.29 Another recent DFT +U study30 reports that

Figure 4. XPS valence band and modulated DOS (a) LiCoO2 and (b) α-LiAlO2.

dipole moment and net charge. As the termination for this surface is unambiguously defined because all layers are equivalent, the (110) surface was selected as a simple model to simulate and analyze the adsorption of gaseous molecules. 20335



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Figure 6. SO2 adsorption on cationic sites of (110) surface of LiCoO2 and α-LiAlO2 (a) one oxygen of SO2 (O2) on top of a metal cation, (b) one oxygen of SO2 on top of a lithium cation, (c) one oxygen between two cations (metal and lithium), and (d) two oxygens of SO2 (O1 and O2), on top of a metal cation and a lithium cation.

the SO2 retention on the Cr2O3 surface is a chemisorption process with sulfite formation on the surface, but now a stable sulfate-like species was found. CO2 adsorption has been the subject of different investigations and some trends emerge. On metal oxides whose cations are in the highest oxidation state, such as MgO and TiO2, CO2 does not bind through its C atom to surface oxygen;31−34 it does bind weakly through its O atoms. Interaction involving the C atom of CO2 occurs only when the metal oxides have cations that are not in a high oxidation state.35 For these surfaces, either the adsorption involves a redox mechanism with (CO2)2− formation or the surface O2− ions are basic enough to bind to CO2 forming carbonates.29 These theoretical studies on metal oxides have revealed that the picture for adsorption is complex, cumulating several acid− base interactions with electron transfer (redox process) occurring in some cases. In order to check acid−base and redox properties, we have undertaken a systematic study of SO2 and CO2 adsorption on (110) surface of LiCoO2 and α-LiAlO2. SO2 Adsorption. Two kinds of approaches were selected with the SO2 plane either perpendicular or parallel to the (110) surface. SO2 is a bent molecule with calculated O−S bond lengths of 1.45 Å and an O−S−O angle of 119.4°, in agreement with experimental structure (O−S = 1.43 Å and O−S−O = 119.3°). The HOMO corresponds to the S (sp) lone pair and the LUMO is of π*SO type. Adsorption on Cationic Sites. Although SO2 is classically used to test the basicity of metal oxides, we have considered SO2 adsorption through O atoms on cationic sites of the surface in order to compare the so-called “acidic properties” of the surfaces of both compounds. Different stable adsorption modes were identified for both compounds and similar results were obtained in the case of low and intermediate coverages. They are characterized by different interactions of oxygen

atoms (SO2) with cobalt/aluminum and lithium atoms at the surface. For LiCoO2, the adsorption energy varies from −0.14 to −0.29 eV and from −0.20 to −0.47 eV for α-LiAlO2, corresponding to physisorption or weak chemisorption when the interaction involves two cationic sites at the surface (Figure 6). This is reflected in the geometry of the adsorbate, which is only slightly perturbed compared to the gas phase structure (increase of S−O = 0.02−0.04 Å, decrease of OSO = 2− 4°) and in the long M−O (SO2)/Li−O (SO2) distances (from 0.3 to 0.4 Å longer than the bond lengths inside the bulk structures). Small charge transfers from the surface to SO2 are observed (from ∼0.08−0.11 e to ∼0.14−0.18 e when one or two bonds are formed). The more important stabilizations observed for α-LiAlO2 compared to LiCoO2 are due to stronger interactions as evidenced by greater electronic transfers from the surface to sulfur dioxide. Adsorption on O Sites. We also investigated the adsorption of SO2 (oxygen atoms noted O1 and O2) with sulfur atom on top of an oxygen site of the surface (O3) in order to probe the so-called “basicity” of the surface. Notations are referred to in Figure 7. The stable adsorption geometries can be divided into two sets, corresponding to the formation of sulfite-like (SO32−) or sulfate like (SO42−) surface species as illustrated in Figure 7. In the first case corresponding to the sulfite species, the situation is more favorable in energy for α-LiAlO2 (Eads (LC) = −1.82 eV, Eads (IC) = −1.67 eV) than for LiCoO2 (Eads (LC) = −1.34 eV, Eads (IC) = −1.16 eV). These sulfite species appear strongly attached to the surface through the formation of an S− O3 bond, with the coexistence of several interactions. Besides the expected sulfur−oxygen interaction corresponding to similar S−O3 distance of ∼1.6 Å for LiCoO2 and α-LiAlO2, we also observed secondary interactions (Table 2, Figure7). These pronounced interactions occur between the oxygen 20336



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the surface. This interaction is evidenced in the difference electron charge density map (Figure 8). The other adsorption mode leading to the formation of sulfate like species is only observed for LiCoO2. Upon adsorption, only one oxygen atom of the SO2 molecule (O1) is oriented toward the nearest surface cations (Co and Li) and interacts with them; the other SO bond is not perturbed by adsorption (Figure 9). The most important interaction occurs between the S atom (especially sp lone pair) and oxygen atoms at the surface, resulting in S−O (S−O3 and S−O4) distances in the range 1.52−1.55 Å (Table 3). This adsorption mode corresponds to a strong binding (chemisorption, bonds created are shown in the difference electron charge density in Figure 10) and is associated with a significant electronic transfer (∼1.1 e) from SO2 (and more precisely from sulfur atom, qS (SO2) = +2.26 e and qS (sulfate LC, IC) = +3.67 e) to the surface (Figure 10). The Bader charge analysis shows the delocalized nature of this electronic transfer at the surface. It affects the oxygen atoms in direct interaction with sulfur atom, the cobalt atom Co1 at the surface but also cobalt atoms (Co2 and Co3) of the second layer. The sensitivity of the sub surface to the presence of ad-atoms has been evidenced in a recent theoretical work.36 These results reflect the variation of the oxidation state of sulfur from IV to VI and in parallel the reduction process of cobalt cations at the surface/subsurface of LiCoO2. The clearest evidence for the electron transfer from SO2 to the cobalt cations of the sub surface (Co2 and Co3) is the variation of the number of unpaired electron nα−nβ [from 0 (Low Spin) to ∼3 (High Spin)]. This sulfate adsorption mode corresponds to a strong binding (chemisorption) (Eads (LC) = 2.64 eV, Eads (IC) = −2.52 eV) more favorable in energy that the “sulfite” mode (Eads (LC) = −1.34 eV, Eads (IC) = −1.16 eV) also existing for LiCoO2. This result points out that the SO2 adsorption mode for LiCoO2 is not controlled by acid−base properties but by a redox mechanism. CO2 Adsorption. To improve our knowledge on the surface reactivity of both materials, we have also examined CO2 adsorption on the (110) surface of LiCoO2 and α-LiAlO2. In view of the previous results, we only considered the intermediate coverage. CO2 is a relatively weak Lewis acid and some differences are expected compared to SO2. As before, we analyzed adsorption on cationic sites of the surface, through O atoms of CO2 (O1 and O2) and adsorption with a C atom on top of the oxygen sites of the surface. CO2 is a linear molecule with calculated C−O bond lengths of 1.18 Å. The HOMO and the LUMO are respectively associated to the πg symmetry type bonding combination of oxygen lone pairs and to the πu antibonding orbital (π*C−O). Adsorption on Cationic Sites. For adsorption on cationic sites of the surface, different initial configurations were investigated, but only one stable configuration was identified for both compounds at low and intermediate coverage. For this adsorption mode, the CO2 geometry is unchanged relative to the isolated molecule (CO2 linear, C−O ≈ 1.17 Å) with one oxygen interacting with a lithium atom at the surface (Lis−O distances in the range 2.2−2.3 Å). The adsorption values are weak (∼ −0.1 eV for LiCoO2 and ∼ −0.17 eV for α-LiAlO2) and correspond to a physisorption process (Figure 11). These results are in line with a very small charge transfer from the surface to CO2 (0.01−0.02 e). Adsorption on Oxygen Sites. We also investigated adsorption with C atom on top of oxygen sites (O3) at the

Figure 7. SO2 adsorption on oxygen sites of the (110) surface of LiCoO2 and α-LiAlO2Formation of sulfite species(a) LiCoO2 and (b) α-LiAlO2.

Table 2. SO2 Adsorption As Sulfite Speciesa on Oxygen Sites of (1 1 0) Surfaces of LiCoO2 and α-LiAlO2 α-LiAlO2

LiCoO2 sulfite O1−S/O2−S O3−S O1−M1/ O2−M2 O1−Li1/ O2−Li2 O1SO2 isolated systems S−O (SO2) OSO (SO2) O−M (bulk) O−Li (bulk)

low coverage

intermediate coverage

low coverage


1.53/1.59 1.59 2.10/1.96

1.54/1.55 1.63 2.13/2.06

1.48/1.66 1.62 1.80/>3

1.46/1.63 1.69 1.78/>3

2.55/2.49

2.41/2.41

2.12/>3

2.15/>3

117°

116°

110°

107°

1.45 119.4° 1.98 2.08

1.93 2.13

a

The bond distances are given in Å and angles in degrees. See Figure 7 for the orientation of SO2 on the surface and for notations of atoms.

atoms of SO2 and the cations of the surface (M1,2 and Li1,2) neighbors to O3. In the case of LiCoO2, both SOi fragments interact and lead to O−Co1,2/Li1,2 distances slightly longer than the bond lengths in the bulk material (O−Co1,2: ∼2 Å, O− Co(bulk): 1.98 Å − O−Li1,2: ∼2.5 Å, O−Li(bulk): 2.08 Å) . For αLiAlO2, only one oxygen atom of SO2 interacts with Al1/Li1 of the surface, but in this case, we observe O−Al1/Li1 distances slightly shorter or equal to the bond lengths in bulk oxides (O− Al1: ∼1.8 Å, O−Al(bulk): 1.93 Å − O−Li1: ∼2.12 Å, O−Li(bulk): 2.13 Å). Similar observations have been reported previously for the adsorption of SO2 on Ca−O23a. These interactions lead to a total electronic transfer from the surface to SO2, evaluated by Bader charge analysis higher for α-LiAlO2 [0.35 e (LC) − 0.37 e (IC)] than for LiCoO2 [0.08 e (LC) − 0.13 e (IC)] (Figure 8). This is mainly due to a more important O3 → S transfer in relation with the most negative O3 atoms of LiAlO2 (qO3 = −1.63 e) compared to LiCoO2 (qO3 = −1.09 e). This O3 interaction of “π” type appears to be the driving force for this sulfite adsorption mode controlled by the basic properties of 20337



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Figure 8. Formation of sulfite species on LiCoO2 and α-LiAlO2Bader charges on isolated systems (a) and )b) supersystem at low coverage (c) and (d) and difference electron charge density maps between the total and the atomic density (the isodensity values varie from −0.05 to 0.05) (e) and (f).

Table 3. SO2 Adsorption As Sulfate Speciesa on Oxygen Sites of (110) Surface of LiCoO2 LiCoO2

Figure 9. SO2 adsorption on oxygen sites of the (110) surface of LiCoO2formation of sulfate species (a) low coverage, (b) intermediate coverage.

sulfate

low coverage


O1−S/O2−S O3−S/O4−S O1−M1 O1−Li1 O1SO2 isolated systems S−O (SO2) OSO (SO2) O−M (bulk) O−Li (bulk)

1.52/1.44 1.55/1.52 2.07 2.16 115°

1.51/1.44 1.55/1.53 2.03 2.24 115° 1.45 119.4° 1.98 2.08

a

The bond distances are given in Å and angles in degrees. See Figure 9 for the orientation of SO2 on the surface and for notations of atoms.

surface. The results obtained for an intermediate coverage of the surface evidence a strong chemisorption for LiCoO2 as for α-LiAlO2 with the formation of quasi-planar “carbonate”-like species strongly attached to the surface. Indeed, the CO2 geometry no longer remains that of the molecule and is very similar in both systems. The C−O distances are in the range 1.24−1.30 Å and the molecule is bent (O−C−O = 128°)

(Table 4). For LiCoO2, as for α-LiAlO2, the C−O3 distance is short ∼1.35 Å, close to the other C−O distances and strong interaction occur between the oxygen atoms of CO2 and Co/Al and Li in the neighborhood of O3 at the surface. The M−O 20338



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distances are slightly shorter than the M−O distances inside the bulk structures (Co−O (CO2): 1.91 Å, Co−O(bulk): 1.98 Å; Al−O (CO2): 1.81 Å, Al−O(bulk): 1.93 Å) and the Li−O distances ∼2.1 Å are the same as in the bulk structures (2.08− 2.13 Å). It is also to be noted that the formation of carbonates results from some extraction phenomena of the surface lithium participating in the tridentate especially significant in the case of α-LiAlO2 (Figure 13). The formation of this kind of species is evidenced in the difference electron charge density map (Figure 13). This lead to rather strong adsorption with the largest energy obtained for α-LiAlO2 (−1.62 eV) compared to LiCoO2 (−1.16 eV) (Figure 12). These different interactions, involving the π*C−O of the adsorbate, lead to small electron transfers, without any modification of spin state, from the surface to CO2. These charge transfers are less important in the case of LiCoO2 (0.13 e) than for α-LiAlO2 (0.35 e), in relation with a control by the basicity of oxygen anions at the surfaces.

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CONCLUSIONS This study highlights clear differences in the (110) surface reactivity of LiCoO2 and α-LiAlO2 and provides new insight into the factors that control the SO2 and CO2 adsorption. The distinction between acid−base and oxidation/reduction (redox) reactions, fundamental from a chemical point of view, is illustrated. In all cases, adsorption on cationic sites appears less favorable than on oxygen sites, indicating that the (110) surfaces of LiCoO2 and α-LiAlO2 are basic. Considering the SO2 adsorption, we evidence the formation of sulfite and sulfate species strongly attached to the surfaces with the coexistence of several acid−base interactions. For sulfites species existing for α-LiAlO2 and LiCoO2, no formal charge transfer occurs. The largest adsorption energies obtained for α-LiAlO2 are indicative of a stronger basic character of the surface compared to LiCoO2. When chemisorption produces sulfate species, the relevance of redox chemistry is apparent and leads to a strong adsorption, the most favorable for low as for intermediate coverage in the case of LiCoO2. An electron transfer process occurs from sulfur to the surface that accommodates this extra electron by electron delocalization; this is clearly evidenced by a variation of the spin state of near surface cobalt atoms. For CO2 adsorption, the formation of carbonate species strongly attached to the surface is observed for LiCoO2 as for α-LiAlO2. The strongest adsorption determined for α-LiAlO2 is in line with the high basicity of (110) surface in this case and the absence of a redox mechanism for LiCoO2. The picture that emerges from these results is that the surface reactivity of α-LiAlO2 is dominated by the Lewis basicity of surface oxide anions (O2−). The situation is more complex for LiCoO2 with an adsorption mode controlled either by electron transfer/redox (SO2 chemisorption producing sulfate species) or by acid−base interactions (CO2 chemisorption producing carbonate species with no electron transfer occurring). These results allow for understanding the significant evolutions experimentally observed upon SO2 adsorption when going from LiCoO 2 to α-LiAlO 2 . Indeed, the modification of the surface reactivity when Co3+ ions are substituted by Al3+ ions is due to a change from an adsorption mode controlled by redox properties for LiCoO2 to a less energetically favorable adsorption mode controlled by acid base

Figure 10. Formation of sulfite species on LiCoO2Bader charges on isolated systems (a) supersystem at low coverage (b) and difference electron charge density maps between the total and the atomic density (the isodensity values vary from −0.05 to 0.05) (c).

Figure 11. CO2 adsorption on cationic sites of the (110) surface of LiCoO2 and α-LiAlO2.

Table 4. CO2 Adsorptiona on Oxygen Sites of (110) Surfaces Of LiCoO2 and α-LiAlO2 carbonate

LiCoO2

α-LiAlO2

O1−C/O2−C O3−C O1−Li1 O2−M2 O1CO2 isolated systems C−O (CO2) O−M (bulk) O−Li (bulk)

1.23/1.30 1.37 2.12 1.91 128°

1.24/1.31 1.35 2.05 1.81 128°

1.17 1.98 2.08

1.93 2.13

a

The bond distances are given in Å and angles in degrees. See Figure 12 for the orientation of SO2 on the surface.

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Figure 13. Formation of carbonate species on LiCoO2 and α-LiAlO2Bader charges on isolated systems (a) and (b) supersystem at intermediate coverage (c) and (d) and difference electron charge density maps between the total and the atomic density (the isodensity values varie from −0.05 to 0.05) (e) and (f).

species from sulfates for LiCoO2 to sulfites for α-LiAlO2 are in line with this evolution. This theoretical work also demonstrates, for the first time to our knowledge, fundamental differences in the factors controlling SO2 adsorption compared to CO2 adsorption that may be exploitable in materials containing reducible cations.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Professor C. Minot (LCT, UPMC, Paris, France) for stimulating discussions.

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REFERENCES

(1) Amatucci, G. G.; Tarascon, J-M; Klein, L. C. Solid State Ionics 1996, 83, 167−173. (2) Verdier, S.; Ouatani, L.; El; Dedryvère, R.; Bonhomme, F.; Biensan, P.; Gonbeau, D. J. Electrochem. Soc. 2007, 154, A1088. (3) Appapillai, A. T.; Mansour, A. N.; Cho, J.; Shao-Horn, Y. Chem. Mater. 2007, 19, 5748−5757. (4) Myung, S.-T.; Kumagai, N.; Komaba, S.; Chung, H.-T. Solid State Ionics 2001, 139, 47−56. (5) Dahéron, L.; Dedryvère, R.; Martinez, H.; Flahaut, D.; Ménétrier, M.; Delmas, C.; Gonbeau, D. Chem. Mater. 2009, 21, 5607−5616.

Figure 12. CO2 adsorption on oxygen sites of the (110) surface of LiCoO2 and α-LiAlO2formation of carbonate species.

properties for α-LiAlO2. In addition, the experimental results have evidenced that Al substitution modifies the surface chemical reactivity of LiCoO2. Indeed, only sulfate species were identified for LiCoO2, but both sulfate and sulfite species were characterized for LiCo1‑xAlxO2 compounds. 5 The theoretical results with a change in preferential adsorbed 20340



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(6) Dahéron, L.; Martinez, H.; Dedryvère, R.; Baraille, I.; Ménétrier, M.; Denage, C.; Delmas, C.; Gonbeau, D. J. Phys. Chem. C 2009, 113, 5843−5852. (7) Tasker, P. W. J. Phys. C Solid State Phys. 1979, 12, 4977−4984. (8) Czyzyk, R.; Potze, R.; Sawatzky, S. Phys. Rev. B 1992, 46, 3729− 3735. (9) Kramer, D.; Ceder, G. Chem. Mater. 2009, 21, 3799−3809. (10) Kresse, G.; Hafner, J. Phys. Rev. B 1993, 47, 558−561. (11) Kresse, G.; Furthmuller, J. Comput. Mater. Sci. 1996, 6, 15−50. (12) Kresse, G.; Joubert, D. Phys. Rev. B 1999, 59, 1758−1775. (13) Perdew, J. P.; Chevary, J. A.; Vosko, S. H.; Jackson, K. A.; Penderson, M. R.; Singh, D. J.; Fiolhais, C. Phys. Rev. B 1992, 46, 6671−6687. (14) Perdew, J. P.; Wang, Y. Phys. Rev. B 1992, 45, 13244−13249. (15) Dudarev, S. L.; Botton, G. A.; Savrasov, S. Y.; Humphreys, C. J.; Sutton, A. P. Phys. Rev. B 1998, 57, 1505−1509. (16) Hammami, R.; Dhouib, A.; Fernandez, S.; Minot, C. Catal. Today 2008, 139, 227−233. (17) Bader, R. F. W. Atoms in Molecules. A Quantum Theory; Oxford University Press: New York, 1990. (18) http://theory.cm.utexas.edu/vtsttools/bader/ (19) Orman, H. J.; Wiseman, P. J. Acta Crystallogr., Sect. C 1984, 40, 12−14. (20) Marezio, M.; Remeika, J. P. J. Chem. Phys. 1966, 44, 3143−3144. (21) Kushida, K.; Kuriyama, K. Solid State Commun. 2002, 123, 349− 352. (22) van Elp, J.; Wieland, J. L.; Eskes, H.; Kuiper, P.; Sawatzky, G. A.; de Groot, F. M. F.; Turner, T. S. Phys. Rev. B 1991, 44, 6090−6103. (23) Ménétrier, M.; Rougier, A; Delmas, C. Solid State Commun. 1994, 90, 439−442. (24) Marichal, C.; Hirschinger, J.; Granger, P.; Ménétrier, M.; Rougier, A.; Delmas, C. Inorg. Chem. 1995, 34, 1773−1778. (25) Saadoune, I.; Delmas, C. J. Mater. Chem. 1996, 6, 193−199. (26) Ensling, D.; Thissen, A.; Laubach, S.; Schmidt, P. C.; Jaegermann, W. Phs. Rev. B 2010, 82, 195431. (27) Scofield, J. H. J. Electron Spectrosc. Relat. Phenom. 1976, 8, 129− 137. (28) (a) Noguera, C. J. Phys.: Condens. Matter 2000, 12, R367. (b) Goniakowski, J.; Finocchi, F.; Noguera, C. Rep. Prog. Phys. 2008, 71, 016501. (29) (a) Schneider, W. F. J. Phys. Chem. B 2004, 108, 273−282. (b) Karlsen, E. J.; Nygren, M. A.; Pettersson, L. G. M. J. Phys. Chem. B 2003, 107, 7795−7802. (30) Ranea, V. A.; Hernandez, S. N.; Medina, S.; Irurzun, I. M.; Coria, I. D.; Mola, E. E. Surf. Sci. 2011, 605, 489−493. (31) Calatayud, M.; Markovits, A.; Menetrey, M.; Mguig, N.; Minot, C. Catal. Today 2003, 85, 125−143. (32) Calatayud, M.; Markovits, A.; Minot, C. Catal. Today 2004, 89, 269−278. (33) Markovits, A.; Fahmi, A.; Minot, C. J. Mol. Struct. Theochem. 1996, 371, 219−235. (34) Pacchioni, G.; Ricart, J. M.; Illas, F. J. Am. Chem. Soc. 1994, 116, 10152. (35) Hammami, R.; Batis, H.; Minot, C. Surf. Sci. 2009, 603, 3057− 3067. (36) Asaduzzaman, A. M.; Krüger, P. J. Phys. Chem. C 2008, 112, 19616−19619.

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New Investigations on the Surface Reactivity of Layered Lithium Oxides

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