Li2O2 Wetting on the (110) Surface of RuO2, TiO2, and SnO2: An

Subscriber access provided by Clark University Libraries

Article 2

2

2

2

LiO Wetting on the (110) Surface of RuO, TiO, and SnO: An Initiating Force for Polycrystalline Growth 2

Wen-Tong Geng, and Takahisa Ohno J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp508896s • Publication Date (Web): 17 Dec 2014 Downloaded from http://pubs.acs.org on December 19, 2014

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

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

Page 1 of 32

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

The Journal of Physical Chemistry

Li2O2 Wetting on (110) Surface of RuO2, TiO2, and SnO2: An Initiating Force for Polycrystalline Growth W. T. Geng1,2* and T. Ohno1† 1

GREEN, National Institute for Materials Science, Tsukuba 305-0044, Japan

2

Materials Science and Engineering, University of Science and Technology Beijing, China

Abstract We report a first-principles study on the initial deposition of Li2O2 on three rutile oxide surfaces, RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110). The intermediate discharge product in a Li-air battery, LiO2, is found to be less stable on all rutile surfaces and will be further reduced to Li2O2 through disproportionation reaction. For the first and second layers of deposited Li2O2, the adsorption energy is comparable to the cohesive energy of bulk Li2O2, suggesting Li2O2 will likely wet the oxide surfaces and grow into thin films rather than particles. Electronic structure analyses of interfaces demonstrate Li2O2 /TiO2(110) is metallic and Li2O2 /SnO2(110) is semiconducting with a bandgap of 0.2 eV, substantially smaller than in bulk Li2O2. The large lattice mismatch at these interfaces could create amorphousness of Li2O2 and grain boundaries might form abundantly thereafter, both of which can provide charge and ion transport channels needed for oxygen reduction and evolution reactions in Li-air batteries. Therefore, coating nanostructured carbon cathode with thin films of TiO2 or employing mesoporous TiO2 nanostructures as cathode, could possibly lead to the formation of low-resistance Li2O2 thin films and thereby enhance the rate capacity of Li-air batteries.

Keywords: Li-Air Battery; Lithium Peroxide; Polycrystalline; Grain Boundary; Wetting; Rutile *

[email protected]

†

[email protected]

1

ACS Paragon Plus Environment


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

Page 2 of 32

2


Page 3 of 32

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


1. INTRODUCTION Rechargeable Li-air (Li-O2) batteries, first explored by Abraham and Jiang, [1] have a theoretical energy density almost ten times higher than their Li-ion counterparts, are therefore bearing the hope in development of high-performance power sources. The dominant discharge product in Li-air batteries, lithium peroxide (Li2O2), is intrinsically a wide band-gap insulator [2] as a perfect crystal, and this low electrical conductivity results in high overpotential, low rate capacity, and poor reversibility. [3] Extensive exploration on a wide variety of catalysts and constant emerging reports of improved cathode performance notwithstanding, the real effectiveness of catalysts, and the true mechanism of catalyst-induced overpotential lowering in particular, remain elusive. [4] Since the interface between the catalysts and the discharge product Li2O2 is of solid-solid type, catalysts are bound to be buried up by Li2O2 after the very initial discharge stage. As a consequence, the effect of catalysts on the oxygen evolution reaction (OER) in charging process is presumably indirect, through their possible influence on the cathode (porous carbon in many cases) - Li2O2 interfacial conductivity and the morphology of Li2O2.

Lithium carbonate, Li2CO3, is a side product formed at the cathode, mainly from the decomposition of the electrolyte and the remainder from direct consumption of carbon cathode. [5] There is computational evidence that the more insulating Li2CO3 between carbon and Li2O2 further increases the interfacial electrical resistance and is therefore very unwelcome. [6] It appears that catalysts which are effective in reducing overpotentials, such as Au [7], Ru [8], Pd [9], and Co3O4 [10], suppress to some extent the formation of insulating Li2CO3 and noble metals add no electrical resistance to the interface. The efficacy of catalysts in retaining the interfacial conductivity is, nonetheless, very challenging to quantify due to the difficulty in characterizing the interfacial structure. To get rid of the decomposition problem associated with carbon cathode altogether, development of carbon-free cathodes for Li-O2 batteries is emerging with strong interests. [11][12] 3



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

Page 4 of 32

On the other hand, it is now well established that the morphology of Li2O2 plays a vital role in reducing the charge overpotential. Small particles were found to exhibit lower charge overpotential than large ones which have longer paths for charge transport, [13] in accordance with the poor conductance of Li2O2. Moreover, discharged Li2O2 particles made small at cathodes composed of binder-free multi-walled carbon nanotube paper, [14] hierarchical nanoporous carbon, [15] or three-dimensional ordered macroporous LaFeO3, [16] are all found to be helpful in reducing the charge overpotential. Although the catalysts employed by researchers are greatly diverse, one feature in common of the resulted morphology of Li2O2 is its low dimensionality. Radin et al.’s first-principles density functional theory (DFT) calculations demonstrated that albeit insulating in the bulk, oxygen rich Li2O2 surfaces could be metallic. [17] This surface metallicity, if preserved well during charge and discharge, helps electron transport within the forefront surface of the Li2O2 film but cannot offer transportation path across the film. It was found that lithium vacancies could induce holes in the valence band of Li2O2 and thus make it a band conductor. [18] However, generation of a substantial amount of lithium vacancies is energetically unfavorable.

In a more detailed analysis, Lu et al. proposed that Li2O2 formed on carbon black decorated with Pd nanoparticle are polycrystalline. [19] Concurrently, Geng et al. demonstrated in an independent work using DFT calculations that some grain boundaries in Li2O2 can indeed serve as charge transport channels. [20] These findings prompt us to expect that polycrystalline nanoscale thin films, which can make more efficient use of the cathode surface and bear grain boundaries as charge conduction channels, should be the ideal morphology of Li2O2 as a discharge product. We note that when using glassy carbon area as cathode, which has smooth surface and hence small specific surface, Li2O2 will deposit in thin film with thickness in nanoscale. [6] But unfortunately, formation of Li2CO3 next to carbon, which is even 4


Page 5 of 32

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


more insulating than Li2O2, further hinders electronic transport from carbon electrode into Li2O2 and hence introduces significant additional charge overpotential.

Polycrystalline thin films have wide applications spanning from optical, magnetic, and tribological coatings, to diffusion and thermal barriers. Extensive studies of the structure dependence on deposition parameters have been carried out and it is well recognized that the nucleation stage plays a crucial role in the final morphology of the film. [21, 22] It is particularly so when the thickness of the film is in nano-scale, as is the case of Li2O2 in nano-structured carbon cathode in Li-air battery. For such thin films, the layer-by-layer (Fran- van der Merwe) growth mode, rather than the island (Volmer – Weber) mode is likely to be more favorable. Therefore, to obtain nanocrystalline thin film of Li2O2 in the absence of Li2CO3, wetting at the Li2O2-cathode interface is desired. Wetting at solid-solid interface [23] requires an interfacial Li2O2-cathode interaction stronger than or comparable to Li2O2-Li2O2 interaction. But if Li2O2 binds too strongly with the cathode, a higher charge potential will be needed for the initial layers of Li2O2. Thus, an optimal Li2O2-cathode interaction should have similar strength to the Li2O2-Li2O2 interaction. First-principles study of Li2O2 binding on pristine carbon structures [24] and defective graphene sheet [25] demonstrated that Li2O2 cannot wet bare carbon, in accordance with experimental observations that Li2O2 often grows into particles in discharge. [7-9, 13-16] Therefore, if we keep carbon as the cathode, we need to coat carbon with some other material which can be wetted by Li2O2. Such a buffer layer, which serves as the seed crystal in the common technique used to produce polycrystalline thin films, should have compatibility with carbon in structure, charge conductivity no lower than Li2O2 so as not to introduce additional electrical resistance, stability under operating voltage, and stability under nucleophilic attack by O2- and O22-. Besides, it is also expected to be cost effective and environment benign.

Inspired by the effectiveness of RuO2 supported on reduced grapheme cathode in 5



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

Page 6 of 32

reducing the charge overpotential, [8] we conjecture that TiO2 and SnO2 with similar rutile structure but being much cheaper, might be efficient in modulating the film growth of Li2O2. Both RuO2, [26] TiO2, [27] and SnO2, [28] in rutile structure have been successfully coated onto nanostructures of carbon, in studies of energy conversion and storage. When formed on the surface of TiC that serves as cathode in a Li-O2 battery, the stability of TiO2 could be responsible for the stable and reversible formation and decomposition of Li2O2. [11] Although unlike RuO2, which is metallic and will add little to the electrical resistance when inserted in between carbon and Li2O2, TiO2 (3.0 eV) and SnO2 (3.6 eV) are both wide-gap semiconductors, their bandgaps are significantly smaller than that of Li2O2 (4.9 eV), [2] and can be made even smaller by introducing defects. [29]

Here we report first-principles DFT calculations on the adsorption of 0.5, 1.0, and 2.0 atomic layers of Li2O2 on the (110) surface of three oxides with rutile structure, namely, RuO2, TiO2, and SnO2. We find that at the coverage of 0.5, the binding energies (per formula) of Li2O2 on these three surfaces are 3.28, 3.61, and 4.25 eV, respectively, similar or moderately larger than that of Li2O2 in its bulk system, 3.30 eV; at the coverage of 1.0 and 2.0, they are reduced to 3.23, 2.69, 3.21 eV, and 3.22, 2.99, 3.30 eV. We therefore expect Li2O2 to wet these surfaces in discharge and end up with the formation of thin films. Lattice mismatch between Li2O2 and the rutile surfaces will probably results in the appearance of grain boundaries which could help to release the interfacial stresses. In the meantime, the grain boundaries could also provide paths for electron conductance through the film. In addition, we find that the interface between Li2O2 and TiO2 is metallic and that between Li2O2 and SnO2 is semiconducting with a very small band-gap and thus adds little resistance to the charge transport needed for the oxygen reduction and evolution reactions.

2. MODEL The rutile (110) surfaces were modeled with periodic slabs separated by a vacuum 6


Page 7 of 32

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


layer no less than 12 Å in thickness [Fig. 1(a)]. Each slab consists of eight repeated unit of O–X2O2–O (X=Ru, Ti, and Sn) trilayers. According to the report of Kiejna et al. on a study of the surface structure of rutile TiO2 (110) surface, [30] such a slab is thick enough to minimize the surface–surface interaction. The stoichiometric XO2 (110) surface is terminated by bridge-coordinated oxygen atom and unsaturated X atoms. For RuO2, it is known [31] that exposure to gaseous O2 leads to the formation of two additional surface species: a chemisorbed molecule bridging two neighboring surface Ru atoms and a bonded O atom positioned right above the Ru atom. The atomic adsorption is more stable than the molecular one and occupies 75% of the Ru sites at saturation, [31] thus we have taken this kind of termination as the starting RuO2 (110) structure, denoted as RuO2(110)-(1×1)-O, to study the deposit of Li2O2. Between 100K and 300K, O2 does not adsorb on the rutile TiO2 (110) surface in the absence of surface oxygen vacancies. [32] Similarly, oxygen cannot be exothermically adsorbed from the gas phase on the stoichiometric surface of SnO2 (110). [33] Therefore, we have started from the stoichiometric (110) surfaces for both TiO2 and RuO2 to study the deposition of Li2O2.

Since the aim of the present work is to study the thermodynamic limitations of the electrochemical reactions at the cathode, we have simplified the elctrochemical process of the deposition of Li2O2 on the cathode to chemical adsorption reactions, ignoring altogether the influence of electrolyte involvement, current density and other factors on the morphology of Li2O2. Obviously, the Li2O2 crystal growth on either RuO2(110)-(1×1)-O, TiO2 (110), or SnO2 (110) in a chemical process begins with the adsorption of Li and gaseous O2, which, in the mean time, undergo the reaction 2Li + O2 → Li2O2.

(1)

As there is experimental evidence [34] of the formation of intermediates LiO2 in oxygen reduction reaction (ORR), we also need to examine the energetic of the intermediate reaction Li + O2 → LiO2

(2) 7



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

Page 8 of 32

and the disproportionation of superoxide into peroxide 2LiO2 → Li2O2 + O2

(3)

on the three surfaces. The assumption that O2 is not dissociated in both ORR and oxygen evolution reaction (OER) has strong experimental support. [5-6] The deposition of Li2O2 on rutile (110) commence with the adsorption of LiO2 .

3. COMPUTATIONAL DETAILS The first-principles density functional theory (DFT) computation was performed using Vienna Ab initio Simulation Package. [35] The electron-ion interaction was described using projector augmented wave (PAW) method. [36] The exchange correlation between electrons was treated both with generalized gradient approximation (GGA) in the Perdew-Burke-Ernzerhof (PBE) form [37] and with the screened Heyd-Scuseria-Ernzerhof (HSE) hybrid density functional. [38, 39] To make computational efforts manageable, we performed only PBE calculations for geometry optimization, and HSE treatments were started from the structure given by PBE. We used an energy cutoff of 500 eV for the plane wave basis set for all systems to ensure equal footing. The Brillouin-zone integration was performed within Monkhorst-Pack scheme using k meshes of (4×8×1) for both free and Li/O2 adsorbed. The energy relaxation for each strain step is continued until the forces on all the atoms are converged to less than 3×10-2 eV Å-1. In HSE calculations, one quarter of (α=0.25) the local DFT exchange is replaced by the unscreened and non-local Fock exchange and we adopted a parameter of µ=0.2 Å-1 for the separation of short-range and long-range electron-electron exchange interaction, assuming that at a distance of 2/µ=10 Å, the short-range interaction becomes negligible.

The overbinding of the O2 molecule by GGA calculations is a well-known effect, [37] which is the key factor in underestimating the oxidation energy of metals. [40] Here in the study of Li2O2, we argue that the correction to account for this overbinding is not very essential because upon oxidization, O2 (O=O) molecules do not dissociate 8


Page 9 of 32

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


into O atoms (or O2- ions) but are stretched into (O-O)2- ions with a O-O bondlength of 1.55 Å. Our GGA calculation [20] shows that the energy needed to stretch an isolated O2 molecule in equilibrium (O=O bondlength, 1.23Å) to an O-O separation of 1.55 Å is only 1.76 eV, much smaller than the total binding energy of an O2 molecule. As a result, we did not consider this correction in this work. Since the chemisorbed O atom on RuO2(110) is spin-polarized, we have allowed spin-polarization for all the systems based on this surface. Besides, in view of the fact that the ground state of a LiO2 molecule is spin-polarized, 0.42 eV more stable than the spin un-polarized state, we have also treated adsorption of LiO2 on all three surfaces with spin-polarization. On the other hand, spin-polarization lowers the total energy of a Li2O2 molecule by only 0.01 eV; thus we have treated the Li2O2@TiO2 and Li2O2@SnO2 systems without spin-polarization.

4. RESULTS AND DISCUSSION A. Deposition of Li2O2 First, we have calculated the lattice constants for all three rutile oxides. The first-principles results are listed in Table 1, comparing well with X-Ray-Diffraction (XRD) measurements. Clean surfaces of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(100) were then studied as the reference systems for Li2O2 deposition. The optimized atomic structure of the slab models of these surfaces, with lattice in (110) plane fixed at calculated bulk values and the interlayer distance along [110] allowed to relax, are displayed in panels (a), (b), and (c) respectively in Fig. 1. The calculated surface energy of TiO2(110), 0.53 J/m2, is in good agreement with previous first-principles calculations. [30] Table 1. The calculated lattice constants (Å) of RuO2, TiO2, and SnO2 in rutile structure. They match well with the experimental results from standard XRD data file, JCPDS Nos. 43-1027, 76-1940, and 41-1445, respectively. RuO2 Expt.

TiO2 Theory

Expt.

SnO2 Theory

Expt.

Theory 9



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

Page 10 of 32

a, b (Å)

4.49

4.53

4.59

4.62

4.74

4.77

c (Å)

3.11

3.13

2.96

2.96

3.19

3.22

Figure 1. Slab models of the RuO2(110)-(1×1)-O (a), TiO2(110) (b), and SnO2(110) (c) surfaces. Small (red) circles denote O, and large (gray, blue, and purple) ones denote Ru, Ti, and Sn.

To investigate the deposition of Li2O2 and LiO2 molecules on the rutile surfaces, we started with (1×2) supercells with respect to the ones shown in Fig.1. Different from the adsorption of single atoms, where the number of possible adsorption sites is usually very limited, deposition of Li2O2 and LiO2 molecules could have a great variation in meta-stable geometry. Obviously, an exhausted search for the most stable adsorption configuration is beyond the computational effort affordable in this study. Since, as aforementioned, the O2 molecules cannot be exothermically adsorbed on any of these three surfaces, it is the Li that anchor Li2O2 or LiO2 onto the surfaces. The preferred sites for lone Li atoms are therefore good hints in searching for the stable configurations for Li2O2 and LiO2 molecules. We have examined all the on-top, bridge, and hollow sites defined by the surface O atoms. It turns out that the lowest-energy position for Li on RuO2(110)-(1×1)-O is on-top the saturated surface O and shifted slightly toward the bridge site of the chemically adsorbed O atoms (with respect to the stoichiometric surface); and on TiO2(110) and SnO2(110) Li is most 10


Page 11 of 32

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


stable on hollow site surrounded by two unsaturated and one saturated O atoms. (See Column a in Fig. 2) The adsorption energy of Li is -2.77, -1.62, and -1.99 eV (in reference to the elemental crystal of Li) on the three surfaces. It means that at this surface concentration, Li metal will wet these surfaces. Moreover, we have calculated the solution energy of Li in the bulk of these oxides, using a X24O48 supercell (X=Ru, Ti, and Sn). It is +1.03, -0.70, and -1.85 eV respectively, always less stable than on the corresponding surface.

It is therefore advisable to inspect the positioning of a Li2O2 molecule with both Li atoms located in low energy sites for lone Li atoms. Since the Li-Li distance in a Li2O2 molecule is 3.08 Å, the two cannot both sit in the most stable position. The other factor is that the negatively charged O2 ion is likely to sit above metallic ions, not O ions on the surface. In addition, the O2 ion has two options to set the orientation of its axis, either parallel or perpendicular to the Li-Li axis. With these hints and guidelines in mind, we have calculated four geometries for a Li2O2 molecule adsorbed on the rutile (110) surfaces, which are shown in Columns b, c, d, and e in Fig. 2. The one with the lowest energy for each surface is listed in Column b. On all surfaces, the most natural positioning of the two Li is the symmetric site of the first Li along [1-10] direction, regardless of the O2 molecule. For both TiO2(110) and SnO2(110), such an arrangement is luckily optimal, only small adjustment is needed to form the adsorbed Li2O2 structure. This is, nevertheless, not the case for the RuO2(110)-(1×1)-O surface which has quite different alignment of O atoms. Interestingly, we find that the lowest-energy geometry of a Li2O2 molecule absorbed on this surface has parallel axis for Li-Li and O-O, different from a free Li2O2 molecule. Since Li atoms bond more strongly on RuO2(110)-(1×1)-O than on TiO2(110) and SnO2(110), reduction of O2 is less significant than on the latter surfaces. This is well indicated by the O-O bond lengths, which are 1.30, 1.48, and 1.53Å, respectively. It is worth noting the O-O bond-length in a Li2O2 molecule adsorbed on SnO2(110), 1.53Å, is very close to that in bulk Li2O2 ( 1.55Å). The adsorption energy of Li2O2 is -3.28, -3.61, and -4.25 eV 11



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

Page 12 of 32

on RuO2(110), TiO2(110) and SnO2(110), respectively.

Figure 2. Top view of the adsorption structures of Li (Column a) and Li2O2 (Columns b, c, d, and e) on the (1×2) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces. Green circles are Li, red are O, grey, blue, and purple are Ru, Ti, and Sn, respectively. The most stable adsorption structure of Li2O2 on each surface is listed in Column b.

B. Deposition of LiO2 and its disproportionation into Li2O2 Since there is experimental evidence of the formation of intermediates LiO2 in ORR, [34] we have also studied the adsorption of on the rutile (110) surfaces at the very initial stage of discharge. The strategy to search low energy adsorption geometry is similar to that employed for Li2O2. One more kind of configuration considered here is the one with Li sitting above the O2 molecule (see Fig. 3). Since there are no unsaturated metal atoms on the RuO2(110)-(1×1)-O surface, the oxygen molecule can only adsorb on top of Li (top row in Fig. 3). The most stable adsorption structure of Li2O2 on each surface is listed in Column b. Different from the Li2O2 case, the positioning of Li in LiO2 on RuO2 is nearly the same as a lone Li, an indication of a weak Li-O2 bonding in this adsorbed LiO2 molecule. We note the O-O band length is 12


Page 13 of 32

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


1.23 Å, only slightly larger than that of a free molecule (1.22 Å). As for TiO2(110) and SnO2(110), we find in the most stable adsorption configuration, both Li and O2 are in positions very close to those in adsorbed Li2O2 molecule, i.e., Li in hollow site defined by surface O and O2 molecule bridging two surface Ti or Sn. (Column a in Fig. 3). The adsorption energy for LiO2 is -2.92, -2.59, and -3.18 eV on RuO2(110), TiO2(110) and SnO2(110), respectively. Our calculations show that the O-O bond is stretched to 1.34 and 1.36 Å in the latter two cases. It is worth noting that if we neglect the spin polarization of the LiO2 molecule on TiO2(110) and SnO2(110), its adsorption energy on these two surfaces will be -2.70 and -3.27 eV, and the O-O bond-length will be 1.33 and 1.35 Å, slightly shorter than in the spin-polarized case. This means that spin-polarization of LiO2 slightly weakens its adsorption on TiO2(110) and SnO2(110).

Figure 3. Top view of the adsorption structures of LiO2 on the (1×2) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces. Green circles are Li, red are O, grey, blue, and purple are Ru, Ti, and Sn respectively. Shown in Column a is the most stable adsorption structure on each surface.

We summarize the calculated adsorption energy of a lone Li, LiO2 and Li2O2 molecules on the rutile surface in Table 1. With calculated total energies of both clean 13



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

Page 14 of 32

surfaces and adsorbed systems, we are now able to evaluate the formation heat associated with the disproportionation of superoxide into peroxide 2LiO2 → Li2O2 + O2, which is defined as Edis = [ELi2O2@surf – Esurf – 2ELi – EO2] – 2[ELiO2@surf – Esurf – ELi – EO2] = ELi2O2@surf – 2ELiO2@surf + Esurf + EO2

(4)

The reference systems are chosen as the clean surface and gaseous oxygen. In O-rich conditions, the chemical potentials of O as a function of temperature and pressure were taken from Ref. [8]. Similarly, we also focus on the 300K and one atmosphere condition, under which the chemical potential of O is -10.06 eV/molecule. Our numerical results show that such a disproportionation is exothermic on all three surfaces. The energy release is 0.13, 0.32, and 0.86 eV per Li2O2 molecule, respectively. The present calculations thus strongly suggest that even in the very initial discharge stage, lithium peroxide Li2O2 is the dominant discharge product.

Table 2. Adsorption energy Eads (eV/atom) of a Li atom, LiO2 and a Li2O2 molecule on the RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces. For Li2O2 adsorption, we studied higher coverage, mono-layer (1×1)-1L and double-layer (1×1)-2L. Also listed are the O-O bond-lengths (Å), dO-O. RuO2(110)-(1×1)-O

TiO2(110)

SnO2(110)

Adsorbate

Coverage

Eads

dO-O

Eads

dO-O

Eads

dO-O

Li

(1×2)

-2.77

/

-1.62

/

-1.99

/

Li O2

(1×2)

-2.92

1.23

-2.59

1.34

-3.18

1.36

(1×2)

-3.28

1.30

-3.61

1.48

-4.25

1.53

(1×1)-1L

-3.23

1.32

-2.69

1.46

-3.21

1.48

(1×1)-2L

-3.22

1.52/1.39

-2.99

1.51/1.49

-3.30

1.52/1.54

Li2O2

C. Growth of Li2O2 into mono- and double-layers As discussed above, the binding strength of Li2O2 molecules on the oxide surface largely determines whether it will wet the surface or not. The calculated binding energy of a Li2O2 molecule on the (1×2) RuO2(110)-(1×1)-O surface is -3.28 eV, and 14


Page 15 of 32

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


it increases to -3.61 and -4.25 eV on TiO2(110) and SnO2(110). By comparison, the cohesive energy of bulk Li2O2 is -3.30 eV per molecule. Clearly, binding of Li2O2 on these three surfaces is comparable or moderately stronger than in its bulk form. It can therefore be inferred that if thermodynamics plays a decisive role in the morphology of Li2O2, it will likely grow in a layer by layer mode at the very initial discharge process.

Previous theoretical study suggested that the (0001) surface of Li2O2 has the lowest formation energy, followed by (1-100) and (11-20). [17] However, because of remarkable lattice mismatch between Li2O2 (0001) (hexagonal, x=y=3.16Å; orthorhombic, x=5.47, y=3.16Å) and rutile (110) (RuO2: x=6.41, y=3.13 Å; TiO2: x=6.53, y=2.96 Å; SnO2: x=6.74, y=3.22 Å), perfect Li2O2 (0001) layered structure cannot be expected at the initial stage of discharge. A full investigation of the dynamics and kinetics of the Li2O2 growth is beyond the affordability of the present work. We would like to gain some knowledge of the interfacial structure through a simplified model. We have increased the coverage of Li2O2 on the oxide surfaces to full mono-layer by reducing the size of the computation cells from (1×2) to (1×1), with regard to the (110) rutile surfaces. Doubling of the coverage can be done by putting another Li2O2 molecule in parallel to the original one, that is, along [1-10] on RuO2, and along [001] on TiO2 and SnO2.

The optimized atomic structures of mono-layer adsorption of Li2O2 on the (1×1) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces are displayed in upper (top view) and middle (side view) panels in Fig. 4. A comparison with Fig. 2 shows that from half-layer to one-layer coverage on RuO2(110)-(1×1)-O, the O2 ion rotates 90 degree and its axis is now along [1-10] direction, enabling it bonding with four Li. By comparison, it tilts in the (1-10) plane on TiO2(110), and SnO2(110) surfaces, to reduce the repulsion between nearest-neighbors along [001]. It is worth noting that this tilting, which leads the axis of the O2 ion leaning toward the [110] 15



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

Page 16 of 32

direction, makes the Li2O2 monolayer more like a (0001) layer in its bulk form. As for the adsorption energy, we find that for a mono-layer Li2O2 coverage [(1×1)-1L] it is slightly lower than the half-layer [(1×2)] on RuO2(110), and remarkably lower than that TiO2(110) and SnO2(110) surfaces (see Table 1). It indicates a repulsion between Li2O2 molecules on these surfaces.

Figure 4. The calculated atomic structure of monolayer (top and middle) and double layer (bottom) adsorption of Li2O2 on the (1×1) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(100) surfaces. Top view is displayed only for monolayer adsorption.

To determine the double-layer deposition structure, we make an analogy between the mono-layer adsorbed systems and the clean surfaces, taking O2 ions as the protruding O ions in the clean surfaces. Again, we adjust the position of the second Li2O2 molecule to let at least one of Li sit on top of the O2 ion, or the bridge site defined by the O2 ions of the first layer. Meanwhile, we have considered positioning of the second Li2O2 with Li-Li axis along both [001] and [1-10]. The optimized most stable geometry of double-layer adsorption of Li2O2 on the (1×1) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces are displayed in the bottom panels in Fig. 4. Top view is not shown here because atoms in different layers cannot 16


Page 17 of 32

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


be easily distinguished, which makes the picture a mess. It has to be pointed out that although the double-layer of Li2O2 deposited on rutile surfaces, especially on TiO2(110), and SnO2(110), has an atomic structure more or less similar to that of a double-layer (0001) thin film truncated from Li2O2 bulk, it is not that close. This, we argue, is mainly due to the large lattice mismatch between Li2O2 and rutile oxides.

If the lattice mismatch between two structures is small, as is often the case for secondary phase precipitation in alloys, we can estimate the contribution of interfacial strain to the interfacial binding by considering two extreme cases: imposing lattice constants of phase A to phase B and those of B to A. The real interfacial binding strength is somewhere in between these two cases. [41] But here we are encountering a much larger lattice mismatch at the Li2O2 (0001)/rutile (110) interfaces, where lattice coherence is absent. As a result, we would try to elucidate the trend, not the exact magnitude, of the mismatch effect on the structure of Li2O2. We have evaluated the changes in adsorption energy when assuming the lattice constants of RuO2 (110) (x=6.41, y=3.13Å) which are closer to those of to Li2O2 (0001) (x=5.47, y=3.16Å), for both TiO2 (110) (x=6.53, y=2.96 Å) and SnO2 (110) (x=6.74, y=3.22 Å). Our first-principles calculations demonstrate that the adsorption energy of the first layer of Li2O2 increased from -2.69 to -3.17 eV on TiO2 (110), and the second layer increases from -2.99 to -3.29 eV; on SnO2 (110), it increases from -3.21 to -3.98 eV for the first layer, but decrease from -3.30 to -2.96 eV for the second layer. Taking the average of the first and second layers, its adsorption gets stronger on both surfaces, from -2.84 to -3.23 eV on TiO2 and from -3.26 to -3.47 eV on SnO2. It is clear that the large lattice mismatch between Li2O2 and these rutile oxides weakens the interfacial binding. This weakening, however, is very welcome indeed. If there were a much better lattice match at the interface, we can expect a significantly strengthened interfacial binding, resulting in an increased charge potential during the final stage of the charging process in a Li-air battery.

17



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

Page 18 of 32

Given the very large lattice mismatch, we conjecture that the atomic structure of Li2O2 at the immediate cathode/Li2O2 interface might be somewhat amorphous, which is a very common phenomenon at the grain boundaries or other interfaces. Recent first-principles calculations have suggested that both charge and ion transport can be enhanced in amorphous Li2O2. [42] This amorphousness, on the other hand, is very likely to fade away as crystalline periodicity builds up with continuing Li2O2 deposition. After the initial discharge, polycrystalline lithium peroxide could be expected, as was observed in experiments. [19]

D. Electrical conductance at the interfaces We now look at the electronic structure of the rutile/Li2O2 interfaces to assess their electrical conductivity, a property that plays a vital role in determining capacity limitations in non-aqueous Li-O2 batteries. [43] The calculated density of states (DOS) projected on the interfacial O atoms, one in rutile surface and on in Li2O2, for both TiO2 and SnO2 systems (bottom middle and bottom right in Fig. 3) are plotted in Fig. 4. The DOS of interfacial Li and Ti atoms are not shown because the core levels of both Ti and Li, which lose valence electrons, are far below the Fermi level. We show both PBE and HSE results for comparison. It is seen that for both Li2O2@TiO2(110) and Li2O2@SnO2(110), the semi-local exchange-correlation functional predicts a metallic interface, with the Fermi level located just on bottom of the conduction band in Li2O2@TiO2(110) and just on top of the valence band in Li2O2@SnO2(110); whereas the screened hybrid density functional yields a bandgap of around 0.2 eV for the latter case, and much smaller than in the bulk Li2O2. As a consequence, the presence of Li2O2@TiO2(110) or Li2O2@SnO2(110) interface will probably not bring about additional electrical resistance to the cathode.

Figure 5. Projected density of states on the interfacial O atom, one in TiO2 (left) or SnO2 (right) and one in Li2O2. The interfaces are those shown in the bottom middle and bottom right panels in Fig. 4. For comparison, we show both PBE 18


Page 19 of 32

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


(top) and HSE (bottom) results. The Fermi energy is set to zero.

5. CONCLUSION AND IMPLICATIONS To summarize, we have carried out first-principles density functional theory calculations to study the initial deposition of Li2O2 on three oxide surfaces, RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110). Even under oxygen-rich conditions, the intermediate discharge product in Li-air battery, superoxide LiO2 is found to be less stable on the oxide surfaces and will be further reduced to lithium peroxide Li2O2 through disproportionation reaction, particularly so on the latter two surfaces. For the first and second layers of Li2O2 adsorbed on these surfaces, the adsorption energy is comparable to the cohesive energy in its bulk form, suggesting that Li2O2 will probably wet the oxide surfaces and grow into thin films rather than particles. Electronic structure analyses demonstrate Li2O2 /TiO2(110) is metallic and Li2O2 /SnO2(110) are semiconducting with a bandgap of 0.2 eV, much smaller than in the bulk Li2O2.

Figure 6. Schematic drawing of the predicted polycrystalline Li2O2 on TiO2 thin 19



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

Page 20 of 32

film, which coats the carbon cathode.

The remarkable lattice mismatch at these interfaces could render amorphousness of Li2O2 on the immediate cathode surface and grain boundaries will form abundantly thereafter. Since both amorphousness and grain boundaries could provide charge and ion transport channels needed for oxygen reduction and evolution reactions in Li-air batteries and hence reduce the overpotentials, we speculate that coating the nanostructured carbon cathode with thin TiO2 films, as illustrated in Fig. 6, could help modulating the morphology of discharge product Li2O2 into thin films and thereby enhance the rate capacity of charge-discharge circles. As TiO2 is cheap, stable, and environment-friendly, it can at least serve as good substitute for RuO2 in Li-air batteries. Moreover, even mesoporous nanostructured TiO2 alone, with enhanced charge conductivity by doping or co-doping, could also be harnessed as the cathode for Li-O2 batteries. Very recent experimental works [44, 45] involving TiO2 in cathodes have indeed shown improved stability compared with carbon containing batteries.

ACKNOWLEDGMENTS We are grateful to the support of MEXT Program for Development of Environment Technology using Nanotechnology. References [1] Abraham, K. M.; Jiang, E. J. Electrochem. Soc. 1996, 143, 1-5. [2] Hummlshøj, J. S.; Blomqvist, B.; Datta, S.; Vegge, T.; Rossmeisi, J.; Thygesen, K. S.; Luntz, A. C.; Jacobsen, K. W.; Norskøv, J. K. Communications: Elementary oxygen electrode reactions in the aprotic Li-air battery. J. Chem. Phys. 2010, 132, 071101-1-4. [3] Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J. M. Li-O2 and Li-S 20


Page 21 of 32

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


batteries with high energy storage. Nat. Mater. 2012, 11, 19-29. [4] McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C. On the Efficacy of Electrocatalysis in Nonaqueous Li-O2 Batteries J. Am. Chem. Soc. 2011, 133, 18038– 18041. [5] McCloskey, B. D.; Bethune, D. S.; Shelby, R. M. Girishkumar, G.; Luntz, A. C. Solvents’ Critical Role in Nonaqueous Lithium–Oxygen Battery Electrochemistry. J. Phys. Chem. Lett. 2011, 2, 1161-1166. [6] McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshøj, J. S.; Nørskov, J. K.; Luntz, A. C. Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li-O2 Batteries J. Phys. Chem. Lett. 2012, 3, 997– 1001 [7] Peng, Z.; Freunberger, S. A.; Chen, Y.; Bruce, P. G. A reversible and higher-rate Li-O2 battery Science 2012, 337, 563-566. [8] Jung, H. G.; Jeong, Y. S.; Park, J. B.; Sun, Y. K.; Scrosati, B.; Lee, Y. J.Ruthenium-Based Electrocatalysts Supported on Reduced Graphene Oxide for Lithium–Air Batteries ACS Nano 2013, 7, 3532– 3539. [9] Xu, J. J.; Wang, Z. L.; Xu, D.; Zhang, L. L.; Zhang, X. B.Tailoring Deposition and Morphology

of

Discharge

Products

Towards

High-Rate

and

Long-Life

Lithium-Oxygen Batteries Nat. Commun. 2013, 4, 2438-1-10 [10] Black, R. W.; Lee, J.-H.; Adams, B. D.; Mlims, C. A; Nazar, L. F.The Role of Catalysts and Peroxide Oxidation in Lithium–Oxygen Batteries Angew. Chem., Int. Ed. 2013, 125, 410-414. [11] Ottakam Thotiyl, M. M.; Freunberger, S. A.; Peng, Z.; Chen, Y.; Liu, Z.; Bruce, P. G. A stable cathode for the aprotic Li–O2 battery Nat. Mater. 2013, 12, 1050– 1056. [12] Li, F.; Tang, D.-M.; Chen, Y.; Golberg, D.; Kitaura, H.; Zhang, T.; Yamada, A.; Zhou, H. Ru/ITO: A Carbon-Free Cathode for Nonaqueous Li–O2 Battery Nano Lett. 2013, 13, 4702– 4707. [13] Gallant, B. M.; Kwabi, D. G.; Mitchell, R. R.; Zhou, J.; Thompson, C. V.; Shao-Horn, Y. Influence of Li2O2 Morphology on Oxygen Reduction and Evolution 21



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

Page 22 of 32

Kinetics in Li–O2 Batteries Energy Environ. Sci. 2013, 6, 2518– 2528. [14] Li, F.; Chen, Y.; Tang, D. M.; Jian, Z.; Liu, C.; Golberg, D.; Yamada, A.; Zhou, H. Performance-improved Li–O 2 battery with Ru nanoparticles supported on binder-free multi-walled carbon nanotube paper as cathode Enengy Environ. Sci. 2014, 7, 1648-1652. [15] Lim, H.‐D.; Song, H.; Kim, J.; Gwon, H.; Bae, Y.; Park, K.‐Y.; Hong, J.; Kim, H.; Kim, T.; Kim, Y. H.; Lepró, X.; Ovalle‐Robles, R.; Baughman, R. H.; Kang, K. Superior Rechargeability and Efficiency of Lithium–Oxygen Batteries: Hierarchical Air Electrode Architecture Combined with a Soluble Catalyst Angew. Chem. Int. Ed. 2014, 53, 3926-3932. [16] Xu, J. J.; Wang, Z. L.; Xu, D.; Meng, F. Z.; Zhang, X. B. 3D ordered macroporous LaFeO3 as efficient electrocatalyst for Li–O2 batteries with enhanced rate capability and cyclic performance Energy Environ. Sci. 2014 DOI: 10.1039/C3EE42934B [17] Radin, M. D.; Rodriguez, J. F.; Tian, F.; Siegel, D. J. Lithium Peroxide Surfaces Are Metallic, While Lithium Oxide surfaces Are Not. J. Am. Chem. Soc. 2012, 134, 1093-1103. [18] Hummlshøj, J. S.; Blomqvist, B.; Datta, S.; Vegge, T.; Rossmeisi, J.; Thygesen, K. S.; Luntz, A. C.; Jacobsen, K. W.; Norskøv, J. K. Communications: Elementary Oxygen Electrode Reactions in the Aprotic Li-air Battery. J. Chem. Phys. 2010, 132, 071101-1-4. [19] Lu, J.; et al.

A nanostructured cathode architecture for low charge over

potential in lithium-oxygen batteries Nat. Commun. 2013, 4, 2383-1-9. [20] Geng, W. T.; He, B. L.; Ohno, T. Grain Boundary Induced Conductivity in J. Phys. Chem. C 2013, 117, 25222-25228. [21] Ratsch, C.; Vanables, J. A. Nucleation theory and the early stage of thin film growth J. Vac. Sci. Technol. A 2003, 21, S96-S109. [22] Petrov, I.; Barna, P. B.; Hultman, L.; Greene, J. E. Microstructural evolution during film growth J. Vac. Sci. Technol. A 2003, 21, S117-S128. 22


Page 23 of 32

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


[23] Haber, J.; Machej, T.; Czeppe, T. The phenomenon of wetting at solid/solid interface Surf. Sci. 1985, 151, 301-310. [24] Xu, Y.; Shelton, W. A. Oxygen Reduction by Lithium on Model Carbon and Oxidized Carbon Structures J. Electrochem. Soc. 2011, 158, A1177-A1184. [25] Xiao, J.; et al. Hierarchically Porous Graphene as a Lithium –Air Battery Elctrode Nano Lett. 2011, 11, 5071-5078. [26] Ye, J. S.; Cui, H. F.; Liu, X.; Lim, T. M.; Zhang, W. D.; Sheu, F. S. Preparation and

Characterization

of

Aligned

Carbon

Nanotube–Ruthenium

Oxide

Nanocomposites for Supercapacitors small 2005, 1, 560 –565. [27] George, P. P.; Pol, V. G.; Gedanken, A.; Gabashivili, A.; Cai, M.; Mance, A. M.; Feng, L.; Ruthkosky, M. S. Selective Coating of Anatase and Rutile TiO2 on Carbon via Ultrasound Irradiation: Mitigating Fuel Cell Catalyst Degradation J. Fuel Cell Sci. Technol 2008, 5, 041012-1-9 [28] Lou, X. W.; Li, C. M.; Archer, L. A. Designed Synthesis of Coaxial SnO2@carbon Hollow Nanospheres for Highly Lithium Storage Adv. Mater. 2009, 21, 2536-2539. [29] Yim, C. M.; Pang, C. L.; Thornton, G. Oxygen Vacancy Origin of the Surface Band-Gap State of TiO2 (110) Phys. Rev. Lett. 2010, 104, 036806-1-4. [30] Kiejna, A.; Pabisiak, T.; Gao, S. W. The energetic and structure of rutile TiO2 (110) J. Phys.: Condens. Matter 2006, 18, 4207-4217. [31 ] Kim, Y. D.; Seitsonen, A. P.; Wendt, S.; Wang, J.; Fan, C; Jakob, K.; Over, H; Ertl, G. Characterization of Various Oxygen Species on an Oxide Surface: RuO2 (110) J. Phys. Chem. B 2001, 105, 3752-3758. [32] Henderson, M. A.; Epling, W. S.; Perkins, C. L.; Peden, C. H. F.; Diebold, U. Interaction of Molecular Oxygen with the Vacuum-Annealed TiO2(110) Surface: Molecular and Dissociative Channels J. Phys. Chem. B 1999, 103, 5328-5337. [33] Oviedo, J.; Gillan, M. J. First-principles study of the interaction of oxygen with the SnO2(110) surface. Surf. Sci. 2001, 490, 221 - 236.

23



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

Page 24 of 32

[34] Hassoun, J.; Croce, F.; Armand, M.; Scrosati, B. Investigation of the O2 Electrochemistry in a Polymer Electrolyte Solid-State Cell Angew. Chem., Int. Ed. 2011, 50, 2999– 3002. [35] Kresse, G.; Furthmuller, J. Efficient iterative schemes for ab initio total-energy calculations using a plane-wave basis set. Phys. Rev. B, 1996, 54, 11169. [36] Blochl, P. E. Projector augmented-wave method. Phys. Rev. B, 1994, 50, 17953. [37] Perdew, J. P.; Burke, K.; Ernzerhof, M. Generalized gradient approximation made simple. Phys. Rev. Lett. 1996, 77, 3865-3869. [38] Heyd, S.; Scuseria, G. E.; Ernzerhof, M. Hybrid functionals based on a screened Coulomb potential. J. Chem. Phys. 2003, 118, 8207-8215. [39] Paier, J.; Marsman, M.; Hummer, K.; Kresse, G.; Gerber, I.C.; Angyan, J.G. Screened hybrid density functionals applied to solids. J. Chem. Phys. 2006, 124, 154709-1-13. [40] Wang, L.; Maxisch, T.; Ceder, G. Oxidation energies of transition metal oxides with the GGA+U framework. Phys. Rev. B 2006, 73, 195107-1-6 [41] Geng, W. T.; Ping, D. H.; Gu, Y. F.; Cui, C. Y.; Harada, H. Stability of nanoscale co-precipitates in a superalloy: A combined first-principles and atom probe tomography study Phys. Rev. B 2007, 76, 224102-1-10 [42] Tian, F.; Radin, M. D.; Siegel, D. J. Enhanced Charge Transport in Amorphous Li2O2 Chem. Mater. 2014, 26, 2952-2959. [43] Viswanathan, V.; Thygesen, K. S.; Hummelshøj, J. S.; Nørskov, J. K.; Girishkumar, G.; McCloskey, B. D.; Luntz, A. C. Electrical conductivity in Li2O2 and its role in determining capacity limitations in non-aqueous Li-O2 batteries J. Chem. Phys. 2011, 135, 214704-1-10 [44] Zhao, G. Y.; Mo, R. W.; Wang, B. Y.; Zhang, L.; Sun, K. N. Enhanced Cyclability of Li–O2 Batteries Based on TiO2 Supported Cathodes with No Carbon or Binder Chem. Mater. 2014, 26, 2551-2556. [45] Adams, B. D.; Balck, R.; Radtke, C.; Williams, Z.; Mehdi, B. L.; Browning, N. D.; Nazar, L. F. The importance of Nanometric Passivating Films on Cathodes for 24


Page 25 of 32

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


Li-Air Batteries ACS Nano DOI: 10.1021/nn505337p

Table of Contents Graphic

25



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

Figure 1. Slab models of the RuO2(110)-(1×1)-O (a), TiO2(110) (b), and SnO2(110) (c) surfaces. Small (red) circles denote O, and large (gray, blue, and purple) ones denote Ru, Ti, and Sn.

(a)

(b)

(c)

[110] [001] [110]


Page 26 of 32

Page 27 of 32

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


Figure 2. Top view of the adsorption structures of Li (Column a) and Li2O2 (Columns b, c, d, and e) on the (1×2) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces. Green circles are Li, red are O, grey, blue, and purple are Ru, Ti, and Sn, respectively. The most stable adsorption structure of Li2O2 on each surface is listed in Column b.

a

c

b

@RuO2

@TiO2

@SnO2 [001]

[110]


d

e


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

Page 28 of 32

Figure 3. Top view of the adsorption structures of LiO2 on the (1×2) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(110) surfaces. Green circles are Li, red are O, grey, blue, and purple are Ru, Ti, and Sn respectively. Shown in Column a is the most stable adsorption structure on each surface.

a

c

b

@RuO2

@TiO2

@SnO2


d

Page 29 of 32

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


Figure 4. The calculated atomic structure of monolayer (top and middle) and double layer (bottom) adsorption of Li2O2 on the (1×1) supercell of RuO2(110)-(1×1)-O, TiO2(110), and SnO2(100) surfaces. Top view is displayed only for monolayer adsorption.

@RuO2

@TiO2

[001]

[110]

[110]

[110]


@SnO2


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

Figure 5. Projected density of states on the interfacial O atom, one in TiO2 (left) or SnO2 (right) and one in Li2O2. The interfaces are those shown in the bottom middle and bottom right panels in Fig. 4. For comparison, we show both PBE (top) and HSE (bottom) results. The Fermi energy is set to zero.


Page 30 of 32

Page 31 of 32

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


Figure 6. Schematic drawing of the predicted polycrystalline Li2O2 on TiO2 thin film, which coats the carbon cathode.

Li2O2 Li2O2 TiO2 Carbon



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

Table of Contents Graphic

Li2O2 Li2O2 TiO2 Carbon


Page 32 of 32

Li2O2 Wetting on the (110) Surface of RuO2, TiO2, and SnO2: An

Recommend Documents