Mechanisms of LiCoO2 Cathode Degradation by Reaction with HF

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Mechanisms of LiCoO2 Cathode Degradation by Reaction with HF and Protection by Thin Oxide Coatings Jonathon L. Tebbe,† Aaron M. Holder,†,‡ and Charles B. Musgrave*,†,‡ †

Department of Chemical and Biological Engineering, University of Colorado at Boulder, Boulder, Colorado 80309-0596, United States ‡ Department of Chemistry and Biochemistry, University of Colorado at Boulder, Boulder, Colorado 80309-0215, United States

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S Supporting Information *

ABSTRACT: Reactions of HF with uncoated and Al and Zn oxide-coated surfaces of LiCoO2 cathodes were studied using density functional theory. Cathode degradation caused by reaction of HF with the hydroxylated (101̅4) LiCoO2 surface is dominated by formation of H2O and a LiF precipitate via a barrierless reaction that is exothermic by 1.53 eV. We present a detailed mechanism where HF reacts at the alumina coating to create a partially fluorinated alumina surface rather than forming AlF3 and H2O and thus alumina films reduce cathode degradation by scavenging HF and avoiding H2O formation. In contrast, we find that HF etches monolayer zinc oxide coatings, which thus fail to prevent capacity fading. However, thicker zinc oxide films mitigate capacity loss by reacting with HF to form a partially fluorinated zinc oxide surface. Metal oxide coatings that react with HF to form hydroxyl groups over H2O, like the alumina monolayer, will significantly reduce cathode degradation. KEYWORDS: LiCoO2 cathode, capacity fading, cathode thin films, ald coatings, battery cycling

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INTRODUCTION Rechargeable lithium-ion batteries (LIBs) have become the dominant energy storage device for portable electronics because they possess relatively high gravimetric energy and power densities among commercial secondary batteries.1,2 Commercial LIBs are typically comprised of a graphite anode, a LiMxOy cathode, where M is often Co, Ni, Mn, Fe, or a combination of these metals, and an organic electrolyte, typically comprised of ethylene carbonate or other carbonates mixed with a 1 M LiPF6 conductive additive salt.3−7 LiCoO2 is the prevailing cathode material for LIBs due to its comparatively high and constant operating voltage, high energy density, adequate power density, and moderately long lifetime compared to other LIB cathodes.8−10 Current LIBs, including those based on LiCoO2 cathodes, suffer from capacity fading an exponential decay in the battery’s discharge capacity over time. Capacity fading has been attributed to a loss of active cathode material due to interactions at the interface of the cathode and electrolyte.1,11,12 Unfortunately, increasing the cathode surface area, which enables greater LIB power densities,11−14 may significantly increase cathode degradation and other unfavorable side reactions with the electrolyte.15,16 Several possible causes of degradation of the cathode surface have been suggested, including decomposition and deposition of the liquid electrolyte at the surface,17−19 reactions between the cathode surface and HF or PF6−,15,16 and fatigue cycling from cathode lattice strain upon Li+ intercalation and deintercalation. In this study, we examine whether reactions © 2015 American Chemical Society

between HF formed in the electrolyte and the cathode surface are active pathways for cathode degradation and capacity fading. Although LIB liquid electrolytes are composed of organic carbonates mixed with lithium salts and do not initially contain HF, it has been proposed that HF impurities are generated via decomposition of the LiPF6 salt that is added to the electrolyte and reacts with trace amounts of water via the following reactions:20,21 LiPF6 → LiF(s) + PF5

(1)

H 2O + PF5 → POF3 + 2HF

(2)

This mechanism suggests that decomposition of the LiPF6 additive salt results in LiF precipitating from the electrolyte and the formation of the highly reactive PF5 species. Previous studies have shown that the stability of PF5 depends on the specific organic electrolyte and that PF5 decomposes in organic electrolytes such as ethylene carbonate (EC).22−24 Common LIB manufacturing processes typically lead to water contamination of the organic electrolyte of at least 20 ppm.2 A previous computational study demonstrated the reaction of H2O with PF5 to generate HF,25 which is known to etch metal oxides.26 Although HF reactions with the cathode appear limited by the trace amounts of water present in the electrolyte, a Received: August 24, 2015 Accepted: October 12, 2015 Published: October 12, 2015 24265

DOI: 10.1021/acsami.5b07887 ACS Appl. Mater. Interfaces 2015, 7, 24265−24278

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ACS Applied Materials & Interfaces

impeding Li+ transport at the cathode−electrolyte interface.11,12 A 90% capacity retention over more than 120 charge/discharge cycles was exhibited by a cathode coated with a 1.1 Å thick ALD alumina film, or approximately a single monolayer of alumina.11,12 Although thicker ALD films resulted in even lower rates of capacity fading, they also caused larger initial reductions in capacity, apparently due to the larger impediment to Li+ transport at the interface, which led to overall lower capacities. This observed capacity reduction is consistent with other film deposition techniques that deposit thicker coatings.33 In contrast to monolayer alumina coatings, a monolayer zinc oxide ALD film exhibited no improvement in capacity fading and was etched away during battery cycling.29,33,34 However, it has been shown that thicker zinc oxide coatings on cathodes exhibit similar capacity retention to the monolayer thick ALD alumina films and have a similar initial capacity loss due to the thicker film obstructing Li+ transport.34,35 Currently, the cathode in LIBs accounts for nearly half of the cost of the battery,1,2,8,9,15 and while efforts have been made to reduce this cost by using new, less expensive cathode materials, cathodes continue to be the most expensive component of LIBs, and cathode degradation is a principle cause of their relatively short lifetimes; LIBs lose up to 60% of their capacity after only ∼200 charge/discharge cycles. Unfortunately, an incomplete understanding of the interfacial reactions between the species present in the electrolyte and the cathode surface that lead to capacity fading has hindered the development of LIBs less susceptible to capacity fading by cathode degradation. A detailed understanding of the reactions between electrolyte species and the cathode surface that lead to capacity fading can aid the development of cathode materials and coatings that resist degradation by this mechanism and that therefore may possess significantly longer lifetimes. In the case of degradation of LiCoO2 cathodes in contact with organic electrolytes containing LiPF6 additives, HF produced by reactions of H2O impurities with PF5 has been suspected to react with the cathode to reduce its capacity. Here, we employed density functional theory (DFT) calculations to study HF dissociation at the cathode−electrolyte interface comprised of both uncoated and metal oxide-coated LiCoO2 surfaces to understand the nature of HF reactions with cathode surfaces and provide insight into the protective mechanism of cathode coatings. The reactions of HF with both monolayer alumina and zinc oxide films on (101̅4) LiCoO2, as well as sixlayer thick zinc oxide films, underlie how monolayer alumina and multilayer zinc oxide coatings prevent capacity fading, while monolayer zinc oxide films do not.

neutralization reaction between HF and hydroxyls on the LiCoO2 surface produces H2O and CoF* or LiF, which subsequently reacts with PF 5 to form additional HF. Consequently, the cycle of H2O reacting with PF5 to produce HF in the electrolyte, followed by a HF neutralization reaction with the LiCoO2 surface to produce H2O, drives breakdown of the electrolyte salt and sustains degradation of LiCoO2 and thus, continued capacity fading.27 Possible cycles of cathode and electrolyte degradation and capacity loss are shown in Scheme 1. Scheme 1. Cathode Degradation Cycle via HF Attack and Water Regeneration

In this scheme, cathode degradation results from dissolution of surface Co atoms from the CoO2 lattice by reaction with HF (Route 1) or from a loss of working Li+ caused by formation of LiF precipitates as a result of HF attack (Route 2). Acidcatalyzed dissolution (Route 1) of transition metals from cathodes is of particular concern for cathodes containing Mn, although dissolution of Ni, Fe, and Co from their lithium metal oxides occurs as well, however, to a lesser extent. A recent study found that dissolved transition metals from the cathode migrate and deposit at the anode surface, leading to a significant drop in the capacity.28 This capacity loss associated with incorporating transition metals into the graphite anode surface-electrolyte interface (SEI) was determined to be more detrimental to LIB lifetimes than the capacity fading that directly results from dissolution of the cathode. Degradation via Route 2, however, occurs as a result of LiF precipitates forming on the cathode surface and may lead to capacity loss by blocking Li + intercalation channels or by forming trap states that bind Li+ and thus hinder Li+ transport. Several approaches for preventing LIB capacity fading and the loss of active cathode material have been attempted, including: the use of novel cathode materials that resist dissolution and chemical attack,1 incorporation of electrolyte additives to prevent electrolyte decomposition,15 and deposition of protective coatings on the cathode or anode to prevent HF attack of the active material.1,6,15,28,29 While coating the electrodes with protective thin films can mitigate dissolution and degradation reactions,30−32 these coatings may also impede the transport of Li+ across the cathode-electrolyte interface, and consequently limit Li+ intercalation into the cathode, reducing the power density and discharge capacity of the LIB. Ultrathin films of alumina deposited on LiCoO2 cathodes by atomic layer deposition (ALD) have been shown to significantly reduce capacity fading in LIBs, while minimally

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COMPUTATIONAL DETAILS Quantum Chemical Methods. Ab initio calculations of LiCoO2 cathode surfaces and metal oxide films were performed using the PBE+U method.36 PBE+U combines the Perdew− Burke−Ernzerhof (PBE) DFT generalized gradient approximation (GGA) exchange correlation functional37 with an effective Hubbard onsite Coulomb repulsion term to account for the self-interaction error of GGA in describing first row transition metals. All calculations were performed using a plane wave basis set and the projector augmented wave (PAW) method,38 as implemented in the Vienna Ab initio Simulation Package (VASP).39,40 PAW pseudopotentials33 were used to model the core electrons and explicitly describe the valence 2s and 2p electrons of carbon, fluorine, aluminum, and oxygen; 24266


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ACS Applied Materials & Interfaces the 1s and 2s electrons of lithium; the 4s and 3d electrons of cobalt; and the 3d and 4p electrons of zinc. All calculations utilized a 500 eV cutoff energy and were conducted at the Γ-point. A plane-wave expansion study conducted over a range from 250 to 550 eV determined that a 500 eV cutoff energy was found to provide convergence of the total energy to within 0.8 meV/atom. Extensive Brillouin zone folding resulting from the relatively large supercells reduces the need for k-point sampling. K-point sampling confirmed that the Γ-point calculated total energies were also converged to within 0.8 meV/atom. DFT calculations on first-row transition metal oxides, such as the CoO2, found in LiCoO2, require careful treatment of onsite correlation effects, which often lead to partial occupancies of degenerate 3d states, and thus selfinteraction errors when described by various DFT functionals. To avoid these errors we employed the rotationally invariant formulation of the onsite Hubbard model36 to treat correlation of the 3d electrons at Co and Zn sites. A Hubbard term of Ueff = U − J = 3.3 eV was used to describe the Co d-electrons in LiCoO2,41,42 while a Ueff = 8.5 eV was used for Zn.43 Geometry optimizations were converged to within 1 meV per cell. The lattice dimensions of the 96-atom LiCoO2 R3m ̅ hexagonal unit cell, constructed from 24 replications of the reduced four-atom primitive cell, were computed to be a = b = 2.82 Å and c = 13.96 Å, all within 1% of their experimentally determined values.44 Transition-state searches were conducted using the nudged elastic band method.45,46 Bader charge analysis was performed, using software from the Henkelman group,47,48 to identify the changes in charge for each atom along the computed reaction pathways. In addition, calculations were simplified by approximating reaction energies by the change in internal energy at 0 K (ΔE0). Thus, because ΔH298K = ΔE0 + ΔEZPE + ΔEthermal + PΔV, this approximation relies on significant cancellation of the zero-point energies, heat capacity contributions to the enthalpy and expansion work between the reactants and products, which is usually valid. LiCoO2 Cathode Surface Model. The (101̅4) LiCoO2 surface has been demonstrated to be the dominant active cathode surface for lithium transport, especially when produced under reducing conditions.42 The other predominant LiCoO2 surface is the (0001) surface. However, the (0001) face of LiCoO2 consists of a complete layer of CoO2 octahedra that impedes lithium intercalation, making it inactive for LIB operation.42 Thus, we focused this study on reactions between HF and the (101̅4) LiCoO2 surface because it is the primary active cathode surface. LiCoO2 cathode particles are observed to exhibit a variety of surface films comprised of the degraded organic carbonate-based electrolyte, which are expected to contribute to capacity fading. In this study, we model HF reactions with the LiCoO2 surface to develop a better understanding of the role of HF in cathode degradation and to determine whether this is an active degradation pathway. In addition to limiting the scope to only the (101̅4) surface, we also modeled the cathode in a discharged state (Li0.5CoO2) because delithiation of the LiCoO2 cathode during normal LIB charging does not extend beyond Li0.5CoO2, and we expect similar results for various degrees of delithiation. The (101̅4) LiCoO2 surface was modeled using a 120 atom (101̅4) slab consisting five CoO2 layers, as shown in Figure 1 and described in detail in the Supporting Information and in Table S1 of the Supporting Information. Our results predict hydroxyl-termination of at least half of the surface Co sites to be the most energetically favorable surface composition. OH

Figure 1. Side view of the 120-atom slab model of the (101̅4) LiCoO2 surface, with a 15 Å vacuum gap between the surface and backside of its periodic image, denoted by “∼”. The surface of the (101̅4) LiCoO2 slab is comprised of six Li, six Co, and 12 O surface atoms with hydroxyl groups terminating the surface Co atoms. Hydroxylation of the surface Co sites resulting in fractional coverages of 0.5 or greater were found to produce the most favorable surface terminations. We examined HF dissociation reactions with this model surface that react with CoOH* sites, bare surface Co* sites (no OH termination), surface Li atoms, and the Co−O back-bonds. Co, O, Li, and H are represented as blue, red, green, and white spheres.

termination of the LiCoO2 cathode surface was experimentally shown to persist even under prolonged heating and ultrahigh vacuum.49 Consequently, LiCoO2 surfaces are expected to at least partially consist of CoOH* surface sites, where “*” denotes a surface site, as shown in Figure 1.

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RESULTS AND DISCUSSION Reactions of HF with the OH-Terminated (101̅4) LiCoO2 Surface. Before considering reactions of HF with the hydroxyl-terminated (101̅4) LiCoO2 surface (see Computational Details) we examined the adsorption of EC of the electrolyte at CoOH* sites and found an adsorption energy of −0.44 eV. This suggests that in order for HF to react with the cathode, it must first displace adsorbed EC. We then calculated the adsorption of HF at CoOH* sites. We found that as HF approaches a CoOH* site the minimum energy pathway (MEP) is characterized by a flat region where HF adsorbs intact with an adsorption energy of −0.88 eV and, thus, displaces adsorbed EC. However, this molecularly adsorbed state is metastable and dissociates with no barrier into H+ and F− where H+ bonds to the O atom of the CoOH* site to produce an adsorbed H2O, while F− bonds to a neighboring surface Li+ to produce LiF (see Figure 2a). HF neutralization to form H2O and a LiF precipitate occurs via dissociation from the metastable molecularly adsorbed state and is exothermic by 1.53 eV, and the overall dissociative adsorption of HF at CoOH* surface sites is exothermic by 2.41 eV. We find that the LiF formed as a result of HF dissociation at the CoOH* site interacts with a neighboring surface Li+ to form a lithium fluoride precipitate on the cathode surface, as shown in Figure 2a. Bader charge analysis confirms the full ionic character of the Li+ and F− of the lithium fluoride surface species. The two Li−F bond distances are nearly equal in the product state (1.83 and 1.84 Å), suggesting that this adsorbed LiF species is Li2F+which has been shown to also be insoluble in the organic electrolyte.20,21 However, the Li−F−Li angle is 99.6°, while the gas-phase Li−F−Li molecule is linear. In addition to the high strain in this bond angle, the surface Li+ that Coulombically interacts with the LiF product (see Figure 2a) is also bound to surface O atoms in a similar manner to other unreacted surface Li atoms. This indicates that the product of HF reaction with the LiCoO2 surface is LiF−Li+, which is LiF Coulombically bound to a surface Li+, as shown in Figure 2a. 24267


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HF molecule that reacts with the cathode liberates an additional hydrogen from a CoOH* site until all available hydrogens reside in either HF or H2O molecules. This is illustrated by the H balance described by eq 3 through the concentration of H2O in the system: [H 2O] = [H 2O]init + 1/2[CoOH*] − 1/2[HF] − [H 2O]ads + [H 2O]des

(3)

The extent of capacity loss and cathode degradation depends on the amount of LiF precipitate that deposits on the cathode to block Li+ transport. The level of precipitate formation, in turn, depends on the extent to which HF reacts to form precipitate, as described by the concentration of HF, which is the F balance in the system described by eq 4: [HF] = 2[POF3] − [LiF]surf − x[CoFx ]

(4)

HF dissociation at the cathode surface consumes HF but increases the amount of lithium fluoride precipitate and the concentration of H 2 O, which depends on the initial concentration of H2O in the electrolyte, the concentration of CoOH* surface sites, the concentration of HF, and the concentration of adsorbed H2O, as shown by the H balance described by eq 3. Thus, the amount of LiF precipitate, [LiF]surf, (eq 4) is coupled to the H2O concentration in the system (eq 3) through the concentration of HF. The watercatalyzed cycle shown in Scheme 1 illustrates how the concentrations of H2O and HF in the electrolyte are coupled. The coupling of the concentrations of HF and H2O, as shown in eqs 3 and 4 and Scheme 1, suggests that an initially low H2O contaminant concentration will continually produce HF in the electrolyte until steady-state is reached, where all hydrogen atomsincluding those initially bound to surface CoOH* groupsare contained in either H2O or HF. The prediction of a steady-state concentration of HF is consistent with experimental observations,27 which show an initially low HF concentration followed by a rapid increase in [HF] until reaching steady-state. Although the formation of LiF−Li+ sequesters F and thus lowers the concentration of HF, it deposits onto the LiCoO2 cathode and likely hinders Li+ transport, which reduces the capacity of the battery. Thus, it is vital to limit the reactions of HF with metal oxide cathodes, which either dissolve the cathode or form lithium fluoride precipitates at the Li+ intercalation planes. In the case of HF reactions with LiCoO2 or other layered LiMO2 cathodes, such as LiMn1/3Co1/3Ni1/3O2, preventing precipitate formation is a critical challenge to overcome to minimize cathode degradation and the resulting capacity loss. We examined the effect of including EC electrolyte molecules on the reactivity of the surface toward HF, which we discuss in greater detail in the Supporting Information. In addition to adsorbing directly onto the CoOH* site and thus requiring EC desorption before HF can react with the cathode, EC affects the reactivity of HF toward the CoOH* surface site. We find that HF approaches and displaces EC adsorbed at the reaction site as it adsorbs. In this case HF adsorption is exothermic by 0.36 eV, 0.08 eV less exothermic than for adsorption in the absence of EC. However, the MEP shows that HF approaches the adsorbed EC and CoOH* reaction site along a relatively flat region and then dissociates barrierlessly to form H+ and F− upon reaching the reaction site. The H+ bonds to the O of the CoOH* site to form adsorbed H2O, while F− bonds to surface Li+ to form an LiF that interacts with a

Figure 2. Reactants, products, and reaction schemes of HF dissociations on the OH-terminated LiCoO2 (101̅4) surface at (a) a CoOH* site to form a LiF precipitate and adsorbed H2O, which is barrierless and exothermic by 1.53 eV, referenced to the HF adsorbed state; (b) an exposed Co* site to form a CoF* and a protonated surface OH*, which occurs barrierlessly and is exothermic by 1.09 eV; (c) a CoF* site neighboring a LiF−F+ adsorbed precipitate to form a second LiF adsorbed precipitate, which occurs barrierlessly and is exothermic by 1.09 eV; and (d) a CoF* site to dissociate the Co−O back-bond and form an adsorbed CoF2 species, which occurs with a barrier of 3.02 eV and is endothermic by 0.52 eV, indicating that dissolution of Co as CoF2 is not kinetically active. Here, Co atoms are shown as blue, O as red, Li as green, H as white, and F as light blue spheres.

Cathode degradation via the HF neutralization reactions described above, with CoOH* sites on the cathode acting as bases to produce the LiF salt and H2O, create a cycle of reactions that ultimately leads to capacity fading. The H2O produced from this reaction is weakly bound to the surface and requires only 0.16 eV to desorb. On the basis of this result, we propose that H2O produced at the cathode dissolves into the electrolyte, where it reacts with PF5 to generate two additional HF molecules, as shown in Scheme 1. Thus, we predict that cathode degradation by reaction with HF is autocatalytic where (1) a H2O contaminant in the electrolyte reacts with PF5 to produce two HF molecules, (2) HF then reacts with the LiCoO2 surface to produce H2O, and (3) this H2O then desorbs from the cathode to react with PF5 to generate two additional HF molecules, reinitiating the cycle and further driving capacity fading by reacting with LiCoO2. Each cycle of the HF reaction with the cathode described above not only forms H2O but also produces a LiF−Li+ precipitate that deposits onto the cathode surface. Because this LiF−Li+ precipitate is adsorbed at the Li+ intercalation planes of LiCoO2, as shown in Figure 2a, it impedes Li+ exchange between the cathode and the electrolyte and binds two Li+ cations. The combined effect reduces the working Li+ of the cathode and results in capacity loss. Furthermore, each 24268


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ACS Applied Materials & Interfaces neighboring surface Li+, analogous to the model without EC. We predict that dissociative adsorption of HF at the CoOH* site with EC is exothermic by 2.35 eV, 0.06 eV less exothermic than without the EC molecule. Thus, we conclude that inclusion of EC solvent molecules has little effect on the predicted HF dissociation barrier or reaction energy. We find that EC interacts with the H2O produced by HF dissociation. Without the EC, the O of the H2O is positioned directly above the surface Co* and donates an O lone pair to the dz2 orbital of Co to complete its octahedral coordination, as shown in Figure 2a. We predict that desorption of the H2O product is endothermic by 0.16 eV, while adsorption of EC at a Co* site is exothermic by 0.66 eV, suggesting that EC readily displaces adsorbed H2O from the cathode surface. Although EC directly interacts with the reacting HF, it does not significantly affect the kinetics and overall mechanism of HF dissociation. Thus, explicitly including electrolyte in the model of the surface reaction site does not lead to qualitatively different results, and so we did not include the electrolyte in further calculations in the interest of computational efficiency. Reactions of HF with Exposed Co* Sites on the Partially OH-Terminated (101̅4) LiCoO2 Surface. In addition to studying HF reactions with CoOH* sites of the hydroxyl-terminated (101̅4) LiCoO2 surface, we also considered reactions on partially hydroxyl-terminated surfaces containing exposed and undercoordinated Co* centers susceptible to HF attack. We were motivated to investigate reaction at bare Co* surface sites for three main reasons: (1) cathode fabrication involves prolonged heating of the cathode under ultrahigh vacuum, which has been shown to produce partially hydroxylated LiCoO2 surfaces,49 (2) reaction of HF at CoOH* sites, as shown above, produces a bare Co* site on the surface, as shown in Figure 2a above, and (3) additional reaction mechanisms at bare Co* sites may be active. On partially hydroxylated cathode surfaces with CoOH* sites neighboring bare Co* sites, HF may react at either site or across these two sites or with surface Li. We calculated that HF dissociation across neighboring CoOH* and Co* sites to form CoF* and adsorbed H2O has a reaction energy of −1.58 eV, whereas reaction of HF at neighboring CoOH* and surface Li sites to form adsorbed H2O and LiF precipitate has a reaction energy of −1.37 eV. In both cases, we predict dissociative adsorption of HF occurs without a barrier, similar to the previous case of HF dissociation at the CoOH* site. Thus, in contrast to HF reaction at CoOH* sites which led to formation of LiF−Li+, our calculations show that on partially hydroxylated cathode surfaces HF preferentially reacts at Co* sites to produce CoF* rather than at the Li centers to form the LiF− Li+ precipitate. Accounting for HF dissociation at CoOH* sites and across neighboring CoOH* and Co* sites, our results predict that HF reactions on LiCoO2 cathode surfaces result in the production of CoF* sites, LiF precipitates, and H2O. HF reaction at either CoOH* sites or across neighboring CoOH* and Co* sites liberates hydrogen from surface CoOH* sites by forming H2O, which desorbs into the electrolyte to react with F-containing salts to regenerate HF and continue the cycle of HF reaction at the cathode. Reaction of HF with unterminated surface Co* sites may also occur at sites produced by previous reactions with HF. We investigated HF dissociation at bare Co* sites across a Co−O bond on the (101̅4) surface, which neighbors LiF−Li + produced by HF dissociation, to examine whether reactions

with HF disrupt the CoO2 lattice. We found that HF reacts at bare Co* sites neighboring surface LiF−Li+ species by dissociative adsorption to produce CoF* and a protonated surface O; this reaction also modifies the structure of the neighboring LiF−Li+ species, as shown in Figure 2b. We calculated a reaction energy for this reaction of −1.09 eV and that the reaction occurs barrierlessly. We find that the CoF* product interacts strongly with the LiF−Li+ surface species that distorts significantly from the initial LiF−Li+ species neighboring the CoF* site. Similar to the LiF precipitate formed by HF reaction at CoOH* sites, this precipitate sequesters Li+ and hinders Li+ transport between the cathode and the electrolyte. We find that protonation of the surface O and fluorination of the surface Co only slightly distorts the CoO2 lattice, increasing the distance between the protonated surface O and fluorinated surface Co by 0.08 Å, suggesting that the Co−O surface bond is weakened by reaction with HF. However, our calculations predict no significant change in the Co−O subsurface backbond and that the fluorinated Co atom remains octahedrally coordinated, which suggests that the (101̅4) surface remains intact upon reaction with an HF molecule. We also investigated subsequent HF dissociation at the cathode surface, which after reaction with one HF molecule, consists of CoF* and the neighboring distorted LiF−F+ precipitate (shown in Figure 2b). We explored HF dissociation at three sites: (a) reaction at a surface Co*site to form a second CoF*, (b) at a surface Li+ to form a second LiF precipitate, and (c) at CoF* to form CoF2 and dissociate the Co−O subsurface bond. This allows us to examine whether additional reaction with HF disrupts the CoO2 lattice to eventually dissolve the cathode (Route 1) or continues the trend of LiF precipitate formation (Route 2). We investigated dissociative adsorption of HF at a bare Co* neighboring the CoF* species to produce two neighboring CoF* sites and two protonated O sites. We predict that this reaction is exothermic by 1.23 eV and occurs without barrier, similar to the previous dissociation reaction. Additionally, we found that this reaction also leads to no significant change in the Co−O back-bonds, indicating that further fluorination is necessary to dissolve Co from the (1014̅ ) LiCoO2 surface. We examined reaction of HF with a surface Li+ and CoF* site that neighbors a LiF−F+ (case (b) above, see Figure 2b) to form two neighboring LiF−Li+ species and a CoF* rather than two neighboring CoF* sites and a single LiF, as shown in Figure 2c. We predict dissociative adsorption of HF occurs barrierlessly to form two LiF−Li+ species and two protonated surface O atoms, as shown in Figure 2c, with a reaction energy of −1.79 eV, 0.56 eV more exothermic than reaction at the neighboring Co* sites, as described above. Consequently, our calculations indicate that when HF reacts at a bare Co* site, the F preferentially bonds to a neighboring surface Li+ site over bonding to the Co of the bare, under-coordinated surface Co* site. Formation of additional LiF−Li+ species further hinders Li+ transport and increases cathode degradation. While formation of a second LiF precipitate forms a second protonated surface O, we calculate little distortion of the Co−O bonds of this protonated O atom. Thus, subsequent dissociation of a second HF at the site of reaction with HF comprised of a CoF*, LiF−Li+, and protonated surface O results in no significant weakening of the subsurface Co−O bonds upon formation of either a second LiF precipitate species or a second CoF*. On the basis of the relative exothermicities of the barrierless dissociative adsorptions described above, we 24269


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ACS Applied Materials & Interfaces expect that the HF degradation product distribution of LiF, LiF−Li+, and CoF* surface species is initially kinetically controlled and determined by the availability of the reaction sites on the surface involved in each adsorption. However, over the long time-scale of hundreds of charge/discharge cycles, this distribution will shift away from the CoF* productsbecause it is formed by a reaction that is exothermic by only 1.09 eV and is thus reversible over relatively short time scalesand toward a distribution of LiF−Li+ and LiF productswhich are formed by exothermic reactions of 1.79 and 2.53 eV, respectively, and thus undergo reverse reaction at a rate that is ∼1 × 1012 to 1 × 1024 times slower than CoF* formation. Furthermore, this suggests that degradation of other layered LiMO2 cathode materials via reaction with HF may occur as a result of LiF precipitation on the surface as well as transition metal dissolution from MFx. However, HF formation in the electrolyte has been observed to be slow at typical operating conditions,27,50 and due to fast reaction of HF at the cathode surface, we expect HF-based degradation to be limited by the rate of HF formation in the electrolyte. Finally, we examined a mechanism by which HF reacts to cleave the Co−O subsurface bonds and etch Co from the cathode surface as CoF2, as shown in Figure 2d. We located a transition state for this pathway, which produces an adsorbed CoF2 species. This reaction involves a high barrier of 3.02 eV and is endothermic by 0.52 eV. Thus, reaction of HF to cleave subsurface Co−O bonds is not kinetically active by this route at the temperatures realized during battery operation. Instead, our results predict that reaction to form LiF precipitates is favored and that subsequent reactions with HF form additional LiF. We predict that precipitate formation by Route 2 is favored over dissolution of the CoO2 lattice by Route 1 and is the primary mechanism of degradation for LiCoO2 cathodes employing EC electrolytes with LiPF6 additives, which is consistent with experimental observations of LiF and Li2F+ species on the cathode surface.51,52 Reactions of HF with Alumina-Coated (101̅4) LiCoO2 Cathodes. Ultrathin ALD films of alumina deposited on LiCoO2 cathode particles using between 2 and 12 ALD cycles have been shown to dramatically reduce capacity fading relative to uncoated LiCoO2 cathode particles over 120 charge and discharge cycles.11,53 Furthermore, ∼1 Å thick coatings resulted in the optimum operating capacity by significantly reducing capacity fading while causing a relatively small initial decrease in capacity due to impeding cation transport.11 To understand the mechanism by which ALD films of alumina mitigate capacity fading, we examined HF reactions with the alumina-coated (101̅4) LiCoO2 surface modeled using a similar approach to the one we used to investigate HF reactions with the uncoated cathode surface described above. The trimethylaluminum/water (TMA/H2O) ALD process deposits amorphous films of alumina (Al2O3) with varying degrees of hydroxyl termination, depending on the ALD process conditions and subsequent annealing or exposure to humidity.54−56 XPS of the alumina-coated LiCoO2 surfaces showed that the surface is composed of Al centers bound to either three or four oxygen atoms.56 Consequently, we considered two different Al coordinations for the monolayer Al2O3 ALD films grown using the TMA/H2O ALD process: the first, as shown in Figure 3a, consists of Al bound to the two O atoms of neighboring Co−O* groups and terminated by a single OH group (AlOH*; Al bonded to three O atoms); the second, as shown in Figure 3b, consists of Al bound to the two

Figure 3. Models of OH-terminated Al sites on alumina monolayers deposited on the (101̅4) LiCoO2 surface by ALD. The two expected hydroxyl terminations are (a) an AlOH* site with Al terminated by a single hydroxyl group, and bound to three O atoms, and (b) an Al(OH)2* site where Al is terminated by two hydroxyl groups, and bound to four O atoms. The Al center is bound to the LiCoO2 surface through bridging O atoms that result from reaction of TMA with CoOH* groups during ALD. Here, Co atoms are shown as blue, O as red, Li as green, H as white, and Al as gray spheres.

O atoms of neighboring Co−O* groups and terminated by two OH groups (Al(OH)2*; Al bonded to four O atoms). These two models of OH* sites on LiCoO2 cathode surfaces coated with a monolayer of an Al2O3 ALD film are consistent with the experimentally determined 1.1 Å Al2O3 film thickness. First, we examined HF dissociation at the AlOH* site. We predict that HF dissociatively adsorbs through a metastable adsorbate state that subsequently dissociates without a barrier, similar to reactions of HF at the hydroxylated LiCoO2 surface. This results in an AlF* species with an OH ligand, AlF(OH)*, as well as a second OH group bound to the neighboring AlOH* site, as shown in Figure 4a. This reaction does not produce H2O dative bonded to the Al* center. This is in contrast to the case of HF dissociating at the CoOH* sites, which we predicted forms H2O dative bonded to Co*, which

Figure 4. Reactant, transient, and product structures of HF dissociation on the hydroxyl-terminated (101̅4) LiCoO2 surface with an ALD-like monolayer of alumina at: (a) an AlOH* site to form AlF(OH)*, which occurs without barrier and is exothermic by 2.46 eV, referenced to the reactant, and (b) an AlOH* site neighboring an AlF* site to form two neighboring AlF(OH)* sites, which occurs barrierlessly with a reaction energy of −3.68 eV. Both cases involve formation of a transient H2O along the MEP dative bonded to an Al center. This H2O immediately dissociates and transfers a H to a bridging O atom, creating a neighboring Al center with two hydroxyls and avoiding H2O formation. Unlike the previous reactions with HF on the LiCoO2 surface, HF reaction with the alumina coating does not produce water, which disrupts the HF generation cycle illustrated in Scheme 1. Here, Co atoms are shown as blue, O as red, Li as green, H as white, Al as gray, and F as light blue spheres. 24270


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coating due to increased interactions between neighboring AlF(OH)* sites upon transitioning to a tetrahedral coordination. As HF dissociates across the neighboring AlF*(OH) and Al(OH)2* sites the proton transfers to the OH of AlF*(OH) to form a transient H2O intermediate (Figure 4b). Examination of the MEP indicates that this H2O immediately transfers a proton to the O of the neighboring Al−O−Co bridging O atom to form AlF(OH)2* and a partially dissociated Co−O bond (Figure 4b). Figure 4b shows the proton transfer to the neighboring Al−O−Co site that leads to weakening of the Co− O bond. Dissociation of HF at neighboring AlOH* and AlF* sites leads to the formation of two neighboring AlF(OH)* sites, which sequester both H and F on the surface, and a weakened Co−O bridging bond, thus avoiding formation of both AlF2 and H2O, which prevents liberation of H from the surface, which forms additional HF to drive autocatalytic degradation of the cathode, as demonstrated in the uncoated case. Our calculations demonstrate that Al(OH)* sites on monolayer alumina coatings on LiCoO2 cathodes act as proton relays57−63 where H2O is only formed transiently as AlOH* accepts and donates a proton to transfer protons from dissociating HF molecules to form a hydroxylated surface. Relaying protons to neighboring surface sites as HF dissociates at Al−OH* rather than forming H2O suggests that alumina coatings effectively mitigate H2O formation at the cathode surface even under high HF concentrations, and thus disrupt the cycle of degradation with electrolytes possessing higher contaminant concentrations. For HF dissociation at the AlF(OH)* site, our calculations predict that formation of AlF2* proceeds through a barrier of 2.11 eV and a reaction energy of 0.73 eV, suggesting that this reaction is not kinetically active at 310 K. Thus, our results predict that HF preferentially reacts at AlOH* sites that neighbor AlF(OH)* sites to form neighboring AlF(OH)* sites over reacting at the AlF(OH)* site to produce AlF2*. Because etching to produce AlF3 necessarily proceeds through AlF2* surface intermediates, this result suggests that a partially fluorinated aluminum oxide layer is preferentially formed over etching by reactions with HF that form AlF2* or AlF3. Thus, these results predict that the alumina coating inhibits etching. We also investigated subsequent reactions of HF with two neighboring AlF(OH)* surface sites that result from HF reacting across neighboring AlF*(OH) and Al(OH)2* sites (Figure 4b). For this reaction, we examined two different possible products: (a) three neighboring Al(OH)F* sites and (b) an AlF2* site neighboring an AlF* site. We predict that dissociative adsorption of HF forming three neighboring AlF(OH)* sites has a reaction energy of −2.14 eV, while forming an AlF2* site has a calculated reaction energy of −2.28 eV, both referenced to the initial AlF(OH)* and Al(OH)2* sites and a solvated HF molecule; we calculate that both reactions are barrierless. In both cases, dissociative adsorption of HF involves proton transfer from the reacting AlOH* to the lone pair of the O atom at a neighboring Al−O−Co bridge site, avoiding H2O formation and sequestering H on the surface in a similar manner as the reaction of HF at an AlOH* site as described above (Figure 4a). These results indicate that at the low concentrations of HF observed in the electrolyte, AlF* formation is likely while AlF2* formation requires enough HF to form a sufficiently high surface concentration of AlF* so that AlF2* formation becomes competitive, which is unlikely if HF dissociation does not form H2O. Furthermore, etching requires reaction of HF with AlF2*

then dissolves into the electrolyte to reinitiate the cathode degradation cycle of Scheme 1. We calculate that HF dissociates at the AlOH* site to form AlF* with a reaction energy of −2.84 eV, referenced to the initial solvated HF, as shown in Figure 4a. Examination of the structures along the MEP, as shown in Figure 4a, indicates that that the Al center transitions from a trigonal-planar to a tetrahedral geometry. As the reaction proceeds, the proton from the OH of the reacting AlOH* site transfers to a bridging O atom of Al−O−Co of the neighboring AlOH* site. This proton transfer weakens the Co−O and Al−O bridging bonds, increasing the Co−O distance from 1.88 to 2.00 Å and the Al− O distance from 1.68 to 1.73 Å. This leaves the AlOH* monolayer film intact because not only does the Al−O bond not dissociate but the new Al−F and O−H bonds are formed only at the expense of the H−F bond and a weakening of a Co−O bond and a Al−O bridging bond. Reaction of HF with AlOH* to form AlF(OH)* avoids the formation of H2O, as shown in Figure 4a. Thus, our calculations show that HF dissociation on the Al2O3 coating getters HF impurities from the electrolyte without producing H2 O, which is central to disrupting the autocatalytic degradation cycle. These results predict a different mechanism for the reaction of HF with the alumina-coated cathode surface than for HF reaction with the uncoated, hydroxylated LiCoO2 surface. We considered dissociation of HF at the AlOH* site to form AlF(OH)* and calculated that this reaction is 0.43 eV more exothermic than HF dissociation at a CoOH* site to produce LiF precipitate and H2O. This difference in reaction energies corresponds to a ratio in the Boltzmann factors of the respective equilibrium constants at 310 K (typical of LIB operation) of ∼1 × 108 relative to precipitate formation, which suggests that even incomplete alumina films reduce cathode degradation by HF attack. Additionally, we predict LiF formation does not occur as a result of reaction with HF if an Al center is present. Our results predict that HF dissociation will preferentially form AlF* over LiF precipitates and thus that alumina films getter HF impurities without producing precipitates that block Li+ transport, mitigating degradation via Route 2. As a result, our calculations indicate that the concentration of HF in the electrolyte will be lower for alumina-coated cathodes than for uncoated LiCoO2 cathodes because the initial reactions of HF with the coated cathodes do not generate H2O and subsequent regeneration of HF by reaction of H2O with PF5. Similar to our previous results for reaction of HF with the uncoated LiCoO2 cathode, we also investigated subsequent HF reactions at the AlF(OH)* surface site produced from previous reaction with HF (see Figure 4a). In this case, we examined HF dissociation to form two different products: (a) reaction at AlOH* neighboring the AlF(OH)* site to form two neighboring AlF(OH)* groups, as shown in Figure 4b, as well as at (b) AlF(OH)* produced by previous reaction with HF to form AlF2*. For HF dissociation at neighboring AlF(OH)* and AlOH* sites, our results predict that HF dissociatively adsorbs without barrier to form two neighboring AlF(OH)* sites with a reaction energy of −3.68 eV referenced to the solvated HF. As the reaction proceeds, the Al of the reacting site transitions from a trigonal-planar geometry to a tetrahedral geometry, as shown in Figure 4a. This reaction is even more exothermic than the reaction of HF with AlOH* sites of the unreacted alumina 24271


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ACS Applied Materials & Interfaces sites to form AlF3, and AlF2* sites are expected to be present in low concentrations for the reasons just outlined. However, experimental results on aluminum fluoride coatings on LiCoO2 cathodes show similar capacity retention rates as alumina coatings,64 which suggests that AlF2* sites formed from HF dissociation are unlikely to be etched away from the cathode surface. In addition to calculating HF reaction at the AlOH* site itself and the subsequent reactions of HF with the various sites this reaction produces, we also examined HF dissociation at a surface terminated with Al(OH)2* sites (Figure 3b), which coexist with the AlOH* sites as a result of film deposition via ALD.54−56 We predict that dissociative adsorption of HF at Al(OH)2* surface sites is barrierless and produces AlF(OH)* and H2O with a reaction energy of −2.12 eV, referenced to the Al(OH)2* site and a solvated HF. HF dissociation at Al(OH)2* sites scavenges F, but produces H2O, suggesting that Al(OH)2* surface sites do not prevent the autocatalytic cycle of degradation, which we describe in detail in the Supporting Information. This indicates that an alumina film mitigates LiCoO2 cathode degradation resulting from reaction with HF predominantly by causing HF to react with AlOH* sites, which sequesters H and F as AlF* and OH*, over HF reacting with Al(OH)2* sites. We expect a similar protective effect for alumina coatings on the surface of other layered LiMO2 cathode materials. Reactions of HF with Monolayer Zinc Oxide Coated (101̅4) LiCoO2 Cathodes. In addition to alumina coatings, ultrathin, monolayer coatings of zinc oxide were also deposited on LiCoO2 cathodes in an effort to prevent capacity fading. However, monolayer zinc oxide ALD films deposited on LiCoO2 using less than 10 ALD half-cycles fail to prevent capacity fading, in contrast to monolayer alumina films on LiCoO253. This may be the result of incomplete films containing pinholes that expose the underlying cathode or a difference in the reactivity of thicker films, as cathodes with thicker zinc oxide films exhibit a reduction in capacity fading.34,35 To examine the effect of film thickness, we considered HF reactions with both monolayer and six-layer thick zinc oxide films. First, we investigated reactions of HF with a monolayer zinc oxide film on LiCoO2 to determine why monolayer zinc oxide films fail to prevent capacity fading. We modeled the zinc oxide monolayer using a similar approach to how we modeled the alumina monolayer on LiCoO2, as detailed above. ALD of zinc oxide using the Zn(CH2CH3)2 (diethylzinc) precursor and H2O as an oxidant is generally expected to result in ZnOH* and Zn−O−Zn* bridging sites, depending on the temperature of deposition, the number of deposition cycles, the post deposition thermal anneal and exposure to humidity. Deposition of metal oxide coatings by ALD using H2O as the oxygen source typically results in mixtures of OH termination as observed both experimentally and predicted computationally.65−70 We first examined HF reacting at a ZnOH* site on the zinc oxide-coated LiCoO2 surface (Figure 5a). Our calculations predict that HF dissociates at the ZnOH* site to form H2O and ZnF* without barrier and with a reaction energy of −1.13 eV, referenced to a solvated HF (Figure 5a). The lack of a barrier and the moderate exothermicity suggests that this reaction is thermodynamically and kinetically favorable. However, the ZnOH* reaction site does not transfer its proton to a neighboring ZnOH* upon HF dissociation, unlike HF reaction at AlOH* (Figure 4a). Instead, dissociation of HF at ZnOH*

Figure 5. Reactants, products, and reaction schemes of HF dissociations on ALD zinc oxide films coating the hydroxyl-terminated (101̅4) LiCoO2 surface at (a) a ZnOH* site to form ZnF* and H2O, which occurs without barrier and with a reaction energy of −0.41 eV, referenced to the reactant, and (b) a ZnF* site to form ZnF2 adsorbed to the LiCoO2 surface, which occurs with a barrier of 0.70 eV and is exothermic by 0.97 eV. Similar to the uncoated LiCoO2 case, H2O forms as a result of HF attack, continuing the HF generation and cathode degradation cycle. Furthermore, HF dissociation across the ZnO back-bonds results in etching of the ZnO coating, exposing the underlying LiCoO2 surface that is susceptible to HF attack. Here, Co atoms are shown as blue, O as red, Li as green, H as white, Zn as gray, and F as light blue spheres.

produces H2O, which propagates HF generation and the cathode degradation cycle, similar to the uncoated LiCoO2 surface. Furthermore, this reaction is 1.29 eV less exothermic than reaction with CoOH* sites on the LiCoO2 surface, which suggests that any exposed LiCoO2, which may result from gaps or defects in the extremely thin zinc oxide film, will preferentially react with HF to form LiF precipitates and degrade the cathode. Thus, although the model monolayer zinc oxide coating has a similar structure to that of the aluminum oxide film examined above, our results show that its reactivity toward HF is qualitatively different. Next, we investigated subsequent reaction of HF at the ZnF* site produced from reaction of HF with ZnOH* sites (Figure 5a). Our calculations show that HF adsorbs at the exposed O of the Co−O−Zn bridging site, enabling dissociation of the Zn− O bond that anchors the Zn atom to the surface. This reaction proceeds over a barrier of 0.70 eV, with a reaction energy of −0.96 eV, and forms ZnF2 adsorbed on the surface (Figure 5b). Subsequent desorption of ZnF2 into the electrolyte exposes the underlying OH-terminated LiCoO2 cathode, as shown in Figure 5b. This relatively low barrier suggests that this reaction is active at ambient temperature and thus that monolayer zinc oxide coatings will etch from the surface and expose the underlying LiCoO2 cathode to degradation reactions. We also examined HF dissociation at a ZnOH* site neighboring the newly formed ZnF* site to form adjacent ZnF* sites and an adsorbed H2O, instead of ZnF2. We calculate that this reaction occurs with a reaction barrier of 0.66 eV and that it is exothermic by −0.25 eV. Again, this reaction is expected to be active at ambient and operating temperatures of LIBs. We found that the neighboring ZnF* site sterically hinders H2O desorption from the surface, reducing the relative exothermicity compared to the HF dissociation that produced the first ZnF* site. This suggests that HF dissociation at the 24272


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ACS Applied Materials & Interfaces ZnF* site to produce ZnF2 is thermodynamically favored over dissociation to form neighboring ZnF* sites. Thus, we expect that HF reaction with cathodes coated with a monolayer of ZnO2 to produce a mixture of both ZnF* surface sites and ZnF2, which dissolves into the electrolyte. However, because ZnOH* sites are present in significantly higher surface concentrations than ZnF* sites at the typically low initial concentrations of H2O in the electrolyte, ZnF2 dissolution will be relatively slow, and over time, zinc oxide monolayer coatings will be etched away to leave the LiCoO2 surface vulnerable to reaction with additional HF formed in the electrolyte. Not only does the ZnO2 monolayer etch upon reaction with HF but reaction with HF also produces H2O, driving the HF degradation cycle. This is in stark contrast to aluminum oxide monolayers that, as described above, react with HF to form a fluorinated layer that sequesters both H and F in the film. Our calculations demonstrate that a monolayer zinc oxide film on the (1014̅ ) LiCoO2 cathode surface exhibits cathode degradation similar to that observed for the uncoated cathode, which is consistent with experimental observation.33 The relative ease with which HF dissolves the zinc oxide monolayer results from the divalent nature of Zn and the relatively exposed bridging O bonds that bind the Zn center to the surface. Because Zn is divalent, the three Zn−O bonds of the ZnOH* site possess less covalent character, making them weaker and more susceptible to dissociation by reaction with HF. Thus, HF readily attacks and cleaves the Zn−O back-bonds to produce ZnF2 and etch the zinc oxide coating. Consequently, these results indicate that although a monolayer zinc oxide film initially scavenges HF from the electrolyte, it ultimately reacts with HF to both dissolve and produce H2O, which reacts to form more HF and drives decomposition in a manner similar to the degradation mechanism we describe for the uncoated cathode. While a monolayer of zinc oxide on LiCoO2 will etch as a result of HF neutralization reactions with ZnO2, cathodes with thicker zinc oxide films exhibit a reduction in capacity fading.34,35 Monolayer zinc oxide films dissolve via dissociation of HF at the exposed bridging Zn−O bonds. However, thicker coatings hinder access to these back-bonds and are expected to limit etching reactions. Thus, we examined reactions of HF with a two-layer thick zinc oxide film, where the Zn centers have fourfold coordination with O atoms, as shown in Figure 6a. We calculated reaction at a Zn* site to produce ZnF* and OH* (see Figure 6a). Our calculations show that molecular adsorption of HF at Zn* is exothermic by 0.40 eV. However, similar to previous cases, this molecularly adsorbed state is metastable and dissociates barrierlessly to form ZnF* and OH* with a reaction energy of −1.10 eV relative to solvated HF. The calculated product structures exhibit no significant strain in the Zn−O subsurface bonds, while the Zn−O surface bond across which HF dissociates lengthens by 0.15 Å, from 1.81 to 1.96 Å, indicating a weakening of the surface bonds of the film. Additionally, we considered HF dissociation at a surface ZnOH* site. We found that HF dissociates at ZnOH* to form ZnF* and H2O with an activation barrier of 0.34 eV and a reaction energy that is endothermic by 0.19 eV. Consequently, we expect HF to predominantly dissociate across Zn−O surface bonds, and so we examined subsequent HF dissociation at the resulting ZnF* site. We considered subsequent reaction of HF at ZnF* surface sites to determine if reaction of HF with the thicker two-layer

Figure 6. Reactant and product structures of HF dissociation on the two-layer ALD zinc oxide coating on the LiCoO2 surface (not shown) at: (a) a Zn* site and across a Zn−O bond to form ZnF* and ZnOH*, which occurs without barrier and with a reaction energy of −0.70 eV, and (b) a Zn* site neighboring ZnF* to form two neighboring ZnF* sites, rather than ZnF2, as was predicted to occur for the zinc oxide monolayer case. The dissociation in (b) occurs without barrier and with a reaction energy of −0.72 eV. While ZnF* formation in (b) dissociates a Zn−O back-bond, two additional Zn−O back-bonds remain intact, preserving the coatingunlike the monolayer ZnO coating, where the Zn−O back-bonds are exposed and thus susceptible to reaction with the dissociating HF. We find dissociation preferentially occurs at a Zn* site neighboring ZnF*, analogous to the alumina monolayer coating. Additionally, we again find that HF preferentially adsorbs at a ZnO* site over the surface ZnOH* site. Here, O atoms are shown as red, H as white, Zn as gray, and F as light blue spheres.

ZnO film forms ZnF2 and dissolves ZnO as it does to the zinc oxide monolayer. Our calculations predict that HF does not react with the ZnF* site to form ZnF2 but instead forms two neighboring ZnF* sites (see Figure 6b). We predict that HF undergoes dissociative adsorption to form two neighboring pairs of ZnF* and OH* sites with a reaction energy of −0.72 eV. This suggests that reaction of HF with zinc oxide films without exposed back-bonds preferentially form a fluorinated ZnF* surface rather than etching to expose the unprotected cathode surface. We discuss HF reactions with thicker zinc oxide films below. Reactions of HF with Thick Zinc Oxide Coatings. Thick zinc oxide coatings28,29 on LiCo1/3Ni1/3Mn1/3O2 and LiMn2O4 reduce capacity fading, similar to a monolayer film of alumina. Reduction of capacity fading of LiMn2O4 LIBs has been demonstrated for zinc oxide coatings as thick as 10 nm, with little reduction in power density as a result of film thickness due to the higher Li+ conductivity of zinc oxide relative to alumina.35 Additionally, films of at least six layers of zinc oxide deposited on LiMn2O4 cathodes by ALD34 prevent capacity fading as well. Effective protection is particularly important in these cases because metal dissolution is a significant issue for cathodes containing Mn. The dissolved Mn deposits on the graphite anode, disrupts the SEI, and leads to significant capacity loss.71 Thus, in addition to examining reactions of HF with a monolayer of zinc oxide on LiCoO2, which does not prevent capacity fading, we also considered reactions of HF with six-layer thick zinc oxide films with fourfold coordination of the subsurface Zn centers. We expect that the reaction energetics on thicker films will be similar to those of the six-layer films we investigate. In contrast to the case of a monolayer of zinc oxide on LiCoO2, we modeled the surface of the thicker zinc oxide 24273


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ACS Applied Materials & Interfaces coatings as a six-layer slab of (211̅ 0̅ ) zinc oxide, as shown in Figure 7a. This model was chosen because surface chemistry is

cathode from capacity fading in a similar manner to the protection afforded by an aluminum oxide monolayer coating. Because we found that H2O does not form by HF dissociation at a ZnOH* site, we also examined reaction of HF with an exposed Zn* site on the (211̅ 0̅ ) zinc oxide surface. We found that dissociative adsorption of HF occurs to produce a surface ZnF* and a protonated surface O atom, as shown in Figure 7a. This reaction occurs barrierlessly and is exothermic by 1.31 eV, 0.44 eV more exothermic than reaction at a ZnOH* site. The Zn−O bond across which HF dissociates lengthens by 0.26 Å (from 1.85 to 2.11 Å) or 14%, as shown in Figure 7a. The Zn−O back-bond, however, increases by 0.11 Å (from 1.90 to 2.01 Å) or 6%, indicating that this bond weakens upon reaction with HF, although it remains intact, whereas the Zn− O bond dissociates in the monolayer zinc oxide case. Formation of a Zn−F bond increases the coordination of the exposed surface Zn atom to four, the same coordination as Zn atoms in bulk zinc oxide. These results suggest that HF preferentially reacts at exposed Zn* sites over ZnOH* sites. Next, we investigated HF dissociation at the ZnF* surface site and found that HF reacts to form an additional pair of neighboring ZnOH* and ZnF* sites, (Figure 7b); HF does not react with ZnF* to form ZnF2. Similar to HF dissociation at the ZnOH* site, HF dissociation at ZnF* occurs barrierlessly and with a reaction energy of −1.22 eV, referenced to a solvated HF (Figure 7b). Thus, dissociation of HF at a Zn center neighboring a ZnF* site is 0.09 eV less exothermic than reaction of HF at an exposed Zn* on the zinc oxide surface. This is a result of repulsion between the two negatively charged F atoms of the two neighboring ZnF* sites, as well as strain in the Zn−O back-bonds, which lengthen from 1.88 to 1.99 Å to accommodate the repulsion between the F atoms. However, while the Zn−O back-bonds lengthen, we do not find that they dissociate. Thus, reaction of HF at ZnF* forms a second ZnF* and a second OH*, rather than H2O and ZnF2, in contrast to reaction of HF with a zinc oxide monolayer. We next considered HF dissociation at a Zn* site neighboring the ZnF* and ZnOH* sites produced by dissociation of HF across a Zn−O surface bond. We calculate a reaction energy of 0.87 eV, 0.44 eV less exothermic than dissociation at an isolated Zn−O bond. The lower exothermicity of HF dissociation upon forming neighboring ZnF* sites suggests that HF dissociation reactions on thick zinc oxide-coated surfaces will preferentially lead to dispersed fluorination over formation of neighboring ZnF* sites. As a result, we expect that significant time is required to establish a sufficiently high concentration of neighboring ZnF* sites to lead to ZnF2 formation and etching of the zinc oxide film. Furthermore, because H2O is not produced as a result of dissociation, it is likely that the supply of H will become exhausted before the film dissolves. We have shown that exposed Zn−O back-bonds of the zinc oxide monolayer are susceptible to reaction with HF, which etches the coating. We also examined HF reactions with the (21̅1̅0) zinc oxide surface with a subsurface oxygen vacancy directly below the reaction site to determine whether these back-bonds affect the reactivity of the surface. In the absence of the Zn−O back-bond below the ZnF* site, we found that HF preferentially reacts at the ZnF* site to form ZnF2, as shown in Figure 7c, analogous to the zinc oxide monolayer case. In this case, ZnF2 formation occurs barrierlessly and is exothermic by 0.99 eV. Formation of two neighboring ZnF* sites above the O vacancy, however, also occurs barrierlessly, but is only

Figure 7. Reactant and product structures of HF dissociation on the six-layer (21̅1̅0) ZnO surface model of thick ZnO coatings at: (a) a Zn* site and across a Zn−O bond to form ZnF* and OH* on the surface, which occurs without barrier and with a reaction energy of −1.31 eV, (b) a Zn* site neighboring ZnF* to form two neighboring ZnF* sites, which occurs without barrier and is exothermic by 0.87 eV, and (c) a ZnF* site above a vacancy to model the case where the subsurface Zn−O bond is not present to anchor the surface Zn and prevent formation of ZnF2 on the surface, which occurs barrierlessly and is exothermic by 0.99 eV. Without the subsurface Zn−O backbond ZnF2 preferentially forms over formation of two neighboring ZnF* sites. Here, O atoms are shown as red, H as white, Zn as gray, and F as light blue spheres.

a relatively localized phenomenon and, consequently, a six-layer slab of zinc oxide will effectively model the surface chemistry of thicker films as we expect little influence from the underlying cathode substrate on the surface reactivity of films greater than 10 nm thick. We chose the (21̅1̅0) surface because this is the most stable zinc oxide surface and should capture the relevant chemistry of thick zinc oxide coatings.72 This is an accurate approximation as long as the zinc oxide surface sites are consistent with those of relatively thick coatings of zinc oxide on the cathode particles. In our investigation of the protective role of thicker zinc oxide films, we expect a combination of surface terminations. Consequently, we investigated HF dissociation at both ZnOH* and exposed Zn* sites. Additionally, we studied whether H2O formation occurs as a result of HF reaction at ZnOH* sites on the thick zinc oxide surface, as on the zinc oxide monolayer, or whether formation of two neighboring surface OH* sites occurs, as occurs on the alumina monolayer. We first examined HF reacting at ZnOH* sites. We found that HF does not dissociate at the ZnOH* site to form H2O and a ZnF*, but instead the reacting ZnOH* site transfers its proton to a neighboring Zn−O−Zn bridging O, ultimately forming a ZnF* site and two neighboring ZnOH* sites. We predict that this reaction is barrierless and exothermic by 0.87 eV. This suggests that the zinc oxide film disrupts the degradation cycle by limiting H2O formation, protecting the 24274


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Figure 8. Schematic potential energy surface (PES) of HF dissociation reactions on the hydroxyl-terminated and Al2O3 and ZnO coated (101̅4) LiCoO2 cathode surfaces comparing reactions that lead to degradation via Scheme 1 (blue) and those that prevent capacity fading (green). Protective metal oxide coatings preferentially react with HF to form a surface comprised of MF* and neighboring OH* over H2O and LiF. The reactions shown are (upper to lower): dissolution of Co from the CoO2 lattice as CoF2, dissolution of the ZnO monolayer from the cathode surface as ZnF2, dissociation of HF on thicker ZnO coatings that do not lead to degradation, LiF precipitate and H2O formation on the uncoated cathode surface, and HF dissociation at the Al2O3 monolayer coating. Each PES is referenced to an intact, solvated HF and the cathode surface.

exothermic by 0.84 eV. Thus, we find that a zinc oxide film will be etched away from the cathode if HF reacts at the Zn−O back-bonds (similar the monolayer of zinc oxide) or if a Zn center lacks Zn−O subsurface bonds. Thus, our results suggest that the Zn−O back-bonds play a key role in preventing cathode degradation for zinc oxide-coated cathodes. For thick zinc oxide coatings, HF dissociation does not lead to cleaving of the Zn−O subsurface bond. In the monolayer zinc oxide case the Zn−O bridging bonds between the zinc oxide film and the LiCoO2 substrate lengthened by 0.17 Å upon formation of ZnF*, and ultimately react with HF to form a hydroxyl-terminated cathode surface with an adsorbed ZnF2 species. However, HF dissociation on thick zinc oxide coatings with few subsurface O vacancies leads to a qualitatively different result than HF reaction with a monolayer of zinc oxide; HF reaction on the zinc oxide slab results in formation of a partially fluorinated surface that is resistant to etching and that does not form ZnF2 or H2O. This is of particular concern for amorphous ZnO coatings, which may have high defect and vacancy concentrations, and thus are susceptible to etching. For thin amorphous films, this etching will expose the underlying LiCoO2 surface to degradation by reaction with HF or to the decomposing electrolyte molecules. Accordingly, if zinc oxide coatings are used to protect LIB cathodes, our results suggest depositing coatings at least several monolayers thick to provide ample Zn−O subsurface back-bonds as well as a high surface density of neighboring Zn−O bonds, which provides multiple Zn* sites where HF can react to preferentially form neighboring ZnF* and ZnOH* groups over ZnF2 and H2O. Whereas thicker Al2O3 films were unnecessary to protect the underlying LiMO2 cathode surface and result in an initial reduction in capacity due to the barrier to Li+ transport they

impose, in the case of ZnO2 coatings, thicker films are necessary to protect the underlying cathode and result in a smaller capacity penalty compared to alumina because of the relatively high Li+ conductivity of zinc oxide.73

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CONCLUSIONS We have performed ab initio calculations of the HF dissociation reactions on the (101̅4) LiCoO2 surface and ALD aluminum oxide and zinc oxide-coated LiCoO2 to determine the mechanism of cathode degradation for LiCoO2 cathodes and of the protection afforded by metal oxide coatings, which are summarized below in Figure 8. Reactions of HF at MOH*, MF*, and M* sites were calculatedwhere M are the Co, Al, or Zn metal centers of the LiCoO2 cathode or alumina or zinc oxide coatings. We find that the hydroxyl-terminated LiCoO2 surface is highly reactive toward HF and reacts to form LiF precipitates and H2O with a reaction energy of −1.53 eV. These LiF species are insoluble in the organic electrolyte and form LiF crystallites at the Li+ intercalation channels on the LiCoO2 surface, which inhibit Li+ migration at the cathode−electrolyte interface and lead to capacity fading of the LIB. Furthermore, the H2O produced by this reaction desorbs from the surface and reacts with PF5 in the electrolyte to generate additional HF molecules, which creates an autocatalytic degradation cycle. This degradation only requires the presence of very low initial concentrations of H2O to initiate the cycle but ultimately leads to the autocatalytic and accelerated degradation of the cathode. We find a similar result for HF reacting at monolayer zinc oxide-coated LiCoO2 surfaces; HF reacts at ZnOH* sites to produce ZnF* and H2O with a reaction energy of −1.13 eV and no barrier to reaction. Furthermore, subsequent reaction of HF at ZnF* sites preferentially produces ZnF2 and H2O, which 24275



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etches the coating to expose the underlying LiCoO2 surface to further HF attack and continues the HF generation/reaction cycle. This suggests that monolayer zinc oxide coatings do not protect the underlying cathode from capacity fading, as observed experimentally. We find that the Al centers of monolayer alumina coatings are more reactive toward HF (ΔHrxn = −2.84 eV for dissociative adsorption of HF at AlOH*) than the uncoated and zinc oxide-coated surfaces. Additionally, the alumina coating is not etched by reaction with HF because HF preferentially reacts at AlOH* sites over AlF* sites, leading to a fluorinated alumina surface rather than AlF3 and etching of the alumina coating. Moreover, HF reaction with the alumina coating does not produce H2O upon reaction with HF, thus avoiding the accelerated degradation cycle observed in uncoated LiCoO2 batteries. We find that zinc oxide coatings of multiple layers with fourfold coordinated Zn centers exhibit similar results to the monolayer alumina films. These thicker zinc oxide films, with protected subsurface Zn−O bonds, preferentially form a dispersed fluorinated layer over etching the coating by forming ZnF2, unless a subsurface O vacancy is present. Furthermore, our results predict that the six-layer thick zinc oxide surface does not form H2O as a result of reaction with HF but instead forms an additional OH* site; this disrupts the HF generation/ reaction degradation cycle in a similar manner to how the monolayer alumina coating does. LiCoO2 cathode degradation resulting from HF attack is dominated by LiF precipitate formation, and thus, we recommend metal oxide cathode coatings that favor HF reaction at the metal site in the coating to form OH* and MF* over reacting at surface Co* or Li+ sites to form CoF and LiF. Moreover, we recommend metal oxide coatings for layered LiMO2 cathodes that preferentially form surface hydroxyl groups over H2O as a result of HF dissociation.

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REFERENCES

(1) Etacheri, V.; Marom, R.; Elazari, R.; Salitra, G.; Aurbach, D. Challenges in the Development of Advanced Li-Ion Batteries: A Review. Energy Environ. Sci. 2011, 4, 3243−3262. (2) Goodenough, J. B.; Kim, Y. Challenges for Rechargeable Li Batteries. Chem. Mater. 2010, 22, 587−603. (3) Borgel, V.; Markevich, E.; Aurbach, D.; Semrau, G.; Schmidt, M. On the Application of Ionic Liquids for Rechargeable Li Batteries: High Voltage Systems. J. Power Sources 2009, 189, 331−336. (4) Chen, Z.; Qin, Y.; Amine, K.; Sun, Y.-K. Role of Surface Coating on Cathode Materials for Lithium-Ion Batteries. J. Mater. Chem. 2010, 20, 7606. (5) Chen, Z.; Zhang, L. Z.; West, R.; Amine, K. Gel Electrolyte for Lithium-Ion Batteries. Electrochim. Acta 2008, 53, 3262−3266. (6) Zhang, S. A Review on Electrolyte Additives for Lithium-Ion Batteries. J. Power Sources 2006, 162, 1379−1394. (7) Tasaki, K.; Kanda, K.; Nakamura, S.; Ue, M. Decomposition of LiPF6 and Stability of PF5 in Li-Ion Battery Electrolytes. J. Electrochem. Soc. 2003, 150, A1628. (8) Aricò, A. S.; Bruce, P.; Scrosati, B.; Tarascon, J.-M.; van Schalkwijk, W. Nanostructured Materials for Advanced Energy Conversion and Storage Devices. Nat. Mater. 2005, 4, 366−377. (9) Souquet, J. L.; Duclot, M. Thin Film Lithium Batteries. Solid State Ionics 2002, 148, 375−379. (10) Zhao, Y.; Truhlar, D. G. The M06 Suite of Density Functionals for Main Group Thermochemistry, Thermochemical Kinetics, Noncovalent Interactions, Excited States, and Transition Elements: Two New Functionals and Systematic Testing of Four M06-Class Functionals and 12 Other Function. Theor. Chem. Acc. 2008, 120, 215−241. (11) Jung, Y. S.; Cavanagh, A. S.; Dillon, A. C.; Groner, M. D.; George, S. M.; Lee, S.-H. Enhanced Stability of LiCoO2 Cathodes in Lithium-Ion Batteries Using Surface Modification by Atomic Layer Deposition. J. Electrochem. Soc. 2010, 157, A75. (12) Scott, I. D.; Jung, Y. S.; Cavanagh, A. S.; Yan, Y.; Dillon, A. C.; George, S. M.; Lee, S.-H. Ultrathin Coatings on Nano-LiCoO2 for LiIon Vehicular Applications. Nano Lett. 2011, 11, 414−418. (13) Xia, H.; Lu, L.; Ceder, G. Li Diffusion in LiCoO2 Thin Films Prepared by Pulsed Laser Deposition. J. Power Sources 2006, 159, 1422−1427. (14) Zhu, C.; Yang, C.; Yang, W. D.; Hsieh, C. Y.; Ysai, H. M.; Chen, Y. S. High Performances of Ultrafine and Layered LiCoO2 Powders for Lithium Batteries by a Novel Sol-Gel Process. J. Alloys Compd. 2010, 496, 703−709. (15) Aurbach, D.; Markovsky, B.; Salitra, G.; Markevich, E.; Talyossef, Y.; Koltypin, M.; Nazar, L.; Ellis, B.; Kovacheva, D. Review on Electrode−electrolyte Solution Interactions, Related to Cathode Materials for Li-Ion Batteries. J. Power Sources 2007, 165, 491−499. (16) Dedryvère, R.; Martinez, H.; Leroy, S.; Lemordant, D.; Bonhomme, F.; Biensan, P.; Gonbeau, D. Surface Film Formation on Electrodes in a LiCoO2/graphite Cell: A Step by Step XPS Study. J. Power Sources 2007, 174, 462−468. (17) Park, Y.; Shin, S. H.; Lee, S. M.; Kim, S. P.; Choi, H. C.; Jung, Y. M. 2D Raman Correlation Analysis of Formation Mechanism of Passivating Film on Overcharged LiCoO2 Electrode with Additive System. J. Mol. Struct. 2014, 1069, 183−187. (18) Edstrom, K.; Gustafsson, T.; Thomas, J. The Cathode− electrolyte Interface in the Li-Ion Battery. Electrochim. Acta 2004, 50, 397−403. (19) Aurbach, D.; Markovsky, B.; Rodkin, A.; Cojocaru, M.; Levi, E.; Kim, H. J. An Analysis of Rechargeable Lithium-Ion Batteries after Prolonged Cycling. Electrochim. Acta 2002, 47, 1899−1911. (20) Aurbach, D.; Markovsky, B.; Rodkin, A.; Levi, E.; Cohen, Y. S.; Kim, H.-J; Schmidt, M. On the Capacity Fading of LiCoO 2 Intercalation Electrodes: The Effect of Cycling, Storage, Temperature, and Surface Film Forming Additives. Electrochim. Acta 2002, 47, 4291−4306.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsami.5b07887.

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Research Article

LiCoO2 surface model; inclusion of explicit ethylene carbonate electrolyte molecule at the CoOH* reaction site; HF dissociation at Al(OH)2 surface sites. (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported in part by National Science Foundation (NSF) Grant No. CHE-1214131 (C.B.M.). We also gratefully acknowledge the use of XSEDE supercomputing resources, which are supported by NSF ACI-1053575, and the Janus supercomputer, which is supported by NSF Grant No. CNS-0821794 and the Univ. of Colorado Boulder. We also thank Dr. C. Muhich for useful discussions. 24276


Research Article

ACS Applied Materials & Interfaces (21) Li, C.; Zhang, H. P.; Fu, L. J.; Liu, H.; Wu, Y. P.; Rahm, E.; Holze, R.; Wu, H. Q. Cathode Materials Modified by Surface Coating for Lithium Ion Batteries. Electrochim. Acta 2006, 51, 3872−3883. (22) Kawamura, T.; Okada, S.; Yamaki, J. ichi. Decomposition Reaction of LiPF6-Based Electrolytes for Lithium Ion Cells. J. Power Sources 2006, 156, 547−554. (23) Yang, H.; Zhuang, G. V.; Ross, P. N. Thermal Stability of LiPF6 Salt and Li-Ion Battery Electrolytes Containing LiPF6. J. Power Sources 2006, 161, 573−579. (24) Tasaki, K.; Kanda, K.; Kobayashi, T.; Nakamura, S.; Ue, M. Theoretical Studies on the Reductive Decompositions of Solvents and Additives for Lithium-Ion Batteries near Lithium Anodes. J. Electrochem. Soc. 2006, 153, A2192. (25) Okamoto, Y. Ab Initio Calculations of Thermal Decomposition Mechanism of LiPF6-Based Electrolytes for Lithium-Ion Batteries. J. Electrochem. Soc. 2013, 160, A404−A409. (26) Kang, J. K.; Musgrave, C. B. A Quantum Chemical Study of the Self-Directed Growth Mechanism of Styrene and Propylene Molecular Nanowires on the Silicon (100) 2 × 1 Surface. J. Chem. Phys. 2002, 116, 9907−9913. (27) Lux, S. F.; Lucas, I. T.; Pollak, E.; Passerini, S.; Winter, M.; Kostecki, R. The Mechanism of HF Formation in LiPF6 Based Organic Carbonate Electrolytes. Electrochem. Commun. 2012, 14, 47−50. (28) Aurbach, D. Review of Selected Electrode−solution Interactions Which Determine the Performance of Li and Li Ion Batteries. J. Power Sources 2000, 89, 206−218. (29) Martha, S. K.; Markevich, E.; Burgel, V.; Salitra, G.; Zinigrad, E.; Markovsky, B.; Sclar, H.; Pramovich, Z.; Heik, O.; Aurbach, D.; Exnar, I.; Buqa, H.; Drezen, T.; Semrau, G.; Schmidt, M.; Kovacheva, D.; Saliyski, N. A Short Review on Surface Chemical Aspects of Li Batteries: A Key for a Good Performance. J. Power Sources 2009, 189, 288−296. (30) Kozen, A. C.; Lin, C.-F.; Pearse, A. J.; Schroeder, M. A.; Han, X.; Hu, L.; Lee, S.-B.; Rubloff, G. W.; Noked, M. Next-Generation Lithium Metal Anode Engineering via Atomic Layer Deposition. ACS Nano 2015, 9, 5884−5892. (31) Zhang, J.; Lu, Q.; Fang, J.; Wang, J.; Yang, J.; NuLi, Y. Polyimide Encapsulated Lithium-Rich Cathode Material for High Voltage Lithium-Ion Battery. ACS Appl. Mater. Interfaces 2014, 6, 17965− 17973. (32) Wu, F.; Zhang, X.; Zhao, T.; Li, L.; Xie, M.; Chen, R. Multifunctional AlPO4 Coating for Improving Electrochemical Properties of Low-Cost Li[Li0.2Fe0.1Ni0.15Mn0.55]O2 Cathode Materials for Lithium-Ion Batteries. ACS Appl. Mater. Interfaces 2015, 7, 3773− 3781. (33) Liu, J.; Li, W.; Manthiram, A. Dense Core-Shell Structured SnO2/C Composites as High Performance Anodes for Lithium Ion Batteries. Chem. Commun. (Cambridge, U. K.) 2010, 46, 1437−1439. (34) Zhao, J.; Wang, Y. Surface Modifications of Li-Ion Battery Electrodes with Various Ultrathin Amphoteric Oxide Coatings for Enhanced Cycleability. J. Solid State Electrochem. 2013, 17, 1049− 1058. (35) Arrebola, J. C.; Caballero, A.; Hernán, L.; Morales, J. ReExamining the Effect of ZnO on Nanosized 5 V LiNi0.5Mn1.5O4 Spinel: An Effective Procedure for Enhancing Its Rate Capability at Room and High Temperatures. J. Power Sources 2010, 195, 4278−4284. (36) Dudarev, S.; Botton, G.; Savrasov, S.; Humphreys, C.; Sutton, A. Electron-Energy-Loss Spectra and the Structural Stability of Nickel Oxide: An LSDA+ U Study. Phys. Rev. B: Condens. Matter Mater. Phys. 1998, 57, 1505−1509. (37) Perdew, J.; Burke, K.; Ernzerhof, M. Generalized Gradient Approximation Made Simple. Phys. Rev. Lett. 1996, 77, 3865−3868. (38) Blöchl, P. E. Projector Augmented-Wave Method. Phys. Rev. B: Condens. Matter Mater. Phys. 1994, 50, 17953−17979. (39) Kresse, G.; Furthmüller, J. Efficient Iterative Schemes for Ab Initio Total-Energy Calculations Using a Plane-Wave Basis Set. Phys. Rev. B: Condens. Matter Mater. Phys. 1996, 54, 11169−11186.

(40) Kresse, G.; Joubert, D. From Ultrasoft Pseudopotentials to the Projector Augmented-Wave Method. Phys. Rev. B: Condens. Matter Mater. Phys. 1999, 59, 1758−1775. (41) Dahéron, L.; Dedryvère, R.; Martinez, H.; Ménétrier, M.; Denage, C.; Delmas, C.; Gonbeau, D. Electron Transfer Mechanisms upon Lithium Deintercalation from LiCoO2 to CoO2 Investigated by XPS. Chem. Mater. 2008, 20, 583−590. (42) Kramer, D.; Ceder, G. Tailoring the Morphology of LiCoO2: A First Principles Study. Chem. Mater. 2009, 21, 3799−3809. (43) Palacios, P.; Aguilera, I.; Wahnón, P. Electronic Structure and Optical Properties in ZnO:M(Co, Cd): Effect of Band-Gap Variation. Thin Solid Films 2010, 518, 4568−4571. (44) Akimoto, J.; Gotoh, Y.; Oosawa, Y. Synthesis and Structure Refinement of LiCoO2 Single Crystals. J. Solid State Chem. 1998, 141, 298−302. (45) Henkelman, G.; Uberuaga, B. P.; Jónsson, H. Climbing Image Nudged Elastic Band Method for Finding Saddle Points and Minimum Energy Paths. J. Chem. Phys. 2000, 113, 9901−9904. (46) Peters, B.; Heyden, A.; Bell, A. T.; Chakraborty, A. A Growing String Method for Determining Transition States: Comparison to the Nudged Elastic Band and String Methods. J. Chem. Phys. 2004, 120, 7877−7886. (47) Henkelman, G.; Arnaldsson, A.; Jónsson, H. A Fast and Robust Algorithm for Bader Decomposition of Charge Density. Comput. Mater. Sci. 2006, 36, 354−360. (48) Sanville, E.; Kenny, S. D.; Smith, R.; Henkelman, G. Improved Grid-Based Algorithm for Bader Charge Allocation. J. Comput. Chem. 2007, 28, 899−908. (49) Moses, A. W.; Flores, H. G. G.; Kim, J. G.; Langell, M. A. Surface Properties of LiCoO2, LiNiO2 and LiNi1-xCoxO2. Appl. Surf. Sci. 2007, 253, 4782−4791. (50) Plakhotnyk, A. V.; Ernst, L.; Schmutzler, R. Hydrolysis in the System LiPF6propylene Carbonatedimethyl carbonateH2O. J. Fluorine Chem. 2005, 126, 27−31. (51) Lee, D. J.; Lee, K. S.; Myung, S. T.; Yashiro, H.; Sun, Y. K. Improvement of Electrochemical Properties of Li1.1Al0.05Mn1.85O4 Achieved by an AlF3 Coating. J. Power Sources 2011, 196, 1353−1357. (52) Myung, S.-T.; Amine, K.; Sun, Y.-K. Surface Modification of Cathode Materials from Nano- to Microscale for Rechargeable Lithium-Ion Batteries. J. Mater. Chem. 2010, 20, 7074. (53) Jung, Y. S.; Cavanagh, A. S.; Riley, L. A.; Kang, S. H.; Dillon, A. C.; Groner, M. D.; George, S. M.; Lee, S. H. Ultrathin Direct Atomic Layer Deposition on Composite Electrodes for Highly Durable and Safe Li-Ion Batteries. Adv. Mater. 2010, 22, 2172−2176. (54) Xu, Y.; Musgrave, C. B. A DFT Study of the Al2O3 Atomic Layer Deposition on SAMs: Effect of SAM Termination. Chem. Mater. 2004, 16, 646−653. (55) Heyman, A.; Musgrave, C. B. A Quantum Chemical Study of the Atomic Layer Deposition of Al2O3 Using AlCl3 and H2O as Precursors. J. Phys. Chem. B 2004, 108, 5718−5725. (56) George, S. M. Atomic Layer Deposition: An Overview. Chem. Rev. 2010, 110, 111−131. (57) Kang, J. K.; Musgrave, C. B. The Mechanism of HF/H2O Chemical Etching of SiO2. J. Chem. Phys. 2002, 116, 275−280. (58) Somani, S.; Mukhopadhyay, A.; Musgrave, C. Atomic Layer Deposition of Tantalum Nitride Using a Novel Precursor. J. Phys. Chem. C 2011, 115, 11507−11513. (59) Lim, C. H.; Holder, A. M.; Musgrave, C. B. Mechanism of Homogeneous Reduction of CO2 by Pyridine: Proton Relay in Aqueous Solvent and Aromatic Stabilization. J. Am. Chem. Soc. 2013, 135, 142−154. (60) Lim, C.-H.; Holder, A. M.; Hynes, J. T.; Musgrave, C. B. Reduction of CO2 to Methanol Catalyzed by a Biomimetic OrganoHydride Produced from Pyridine. J. Am. Chem. Soc. 2014, 136, 16081− 16095. (61) Bianco, R.; Hay, P. J.; Hynes, J. T. Proton Relay and Electron Flow in the O–O Single Bond Formation in Water Oxidation by the Ruthenium Blue Dimer. Energy Environ. Sci. 2012, 5, 7741−7746. 24277


Research Article

ACS Applied Materials & Interfaces (62) Fetter, J. R.; Qian, J.; Shapleigh, J.; Thomas, J. W.; GarcíaHorsman, A.; Schmidt, E.; Hosler, J.; Babcock, G. T.; Gennis, R. B.; Ferguson-Miller, S. Possible Proton Relay Pathways in Cytochrome c Oxidase. Proc. Natl. Acad. Sci. U. S. A. 1995, 92, 1604−1608. (63) Pines, E.; Pines, D. Proton Dissociation and Solute-Solvent Interactions Following Electronic Excitation of Photoacids. In Ultrafast Hydrogen Bonding Dynamics and Proton Transfer Prosesses in the Condensed Phase; Elsässer, T., Van den Akker, H., Eds.; Springer: Netherlands, 2002; Vol. 23, pp 155−184. (64) Sun, Y. K.; Han, J. M.; Myung, S. T.; Lee, S. W.; Amine, K. Significant Improvement of High Voltage Cycling Behavior AlF3Coated LiCoO2 Cathode. Electrochem. Commun. 2006, 8, 821−826. (65) Widjaja, Y.; Musgrave, C. B. Quantum Chemical Study of the Mechanism of Aluminum Oxide Atomic Layer Deposition. Appl. Phys. Lett. 2002, 80, 3304−3306. (66) Mukhopadhyay, A. B.; Musgrave, C. B.; Sanz, J. F. Atomic Layer Deposition of Hafnium Oxide from Hafnium Chloride and Water. J. Am. Chem. Soc. 2008, 130, 11996−12006. (67) Widjaja, Y.; Musgrave, C. B. Quantum Chemical Study of the Mechanism of Aluminum Oxide Atomic Layer Deposition. Appl. Phys. Lett. 2002, 80, 3304−3306. (68) Guziewicz, E.; Godlewski, M.; Wachnicki, L.; Krajewski, T. A.; Luka, G.; Gieraltowska, S.; Jakiela, R.; Stonert, A.; Lisowski, W.; Krawczyk, M.; Sobczak, J. W.; Jablonski, A. ALD Grown Zinc Oxide with Controllable Electrical Properties. Semicond. Sci. Technol. 2012, 27, 074011. (69) Guziewicz, E.; Godlewski, M.; Krajewski, T.; Wachnicki; Szczepanik, A.; Kopalko, K.; Wójcik-Glodowska, A.; Przeździecka, E.; Paszkowicz, W.; Łusakowska, E.; Kruszewski, P.; Huby, N.; Tallarida, G.; Ferrari, S. ZnO Grown by Atomic Layer Deposition: A Material for Transparent Electronics and Organic Heterojunctions. J. Appl. Phys. 2009, 105, 105. (70) Kim, S. K.; Hwang, S.; Park, S. K.; Yun, S. Comparison between ZnO Films Grown by Atomic Layer Deposition Using H2O or O3 as Oxidant. Thin Solid Films 2005, 478, 103−108. (71) Park, J. S.; Meng, X.; Elam, J. W.; Hao, S.; Wolverton, C.; Kim, C.; Cabana, J. Ultrathin Lithium-Ion Conducting Coatings for Increased Interfacial Stability in High Voltage Lithium-Ion Batteries. Chem. Mater. 2014, 26, 3128−3134. (72) Breedon, M.; Spencer, M. J. S.; Yarovsky, I. Adsorption of Atomic Nitrogen and Oxygen on ZnO (21̅1̅0) Surface: A Density Functional Theory Study. J. Phys.: Condens. Matter 2009, 21, 144208. (73) Shenouda, A. Y.; Liu, H. K. Studies on Electrochemical Behaviour of Zinc-Doped LiFePO4 for Lithium Battery Positive Electrode. J. Alloys Compd. 2009, 477, 498−503.

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Mechanisms of LiCoO2 Cathode Degradation by Reaction with HF

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