Mechanistic Evaluation of Li x O y Formation on Î´-MnO2 in

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Article x

y

A Mechanistic Evaluation of LiO Formation on #-MnO in Non-Aqueous Li-air Batteries 2

Zhixiao Liu, Luis R. De Jesus, Sarbajit Banerjee, and Partha P. Mukherjee ACS Appl. Mater. Interfaces, Just Accepted Manuscript • DOI: 10.1021/acsami.6b05988 • Publication Date (Web): 17 Aug 2016 Downloaded from http://pubs.acs.org on August 17, 2016

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ACS Applied Materials & Interfaces

A Mechanistic Evaluation of LixOy Formation on -MnO2 in Non-Aqueous Liair Batteries Zhixiao Liu,1 Luis R. De Jesus,2 Sarbajit Banerjee,2,3,* Partha P. Mukherjee1,*

1

Department of Mechanical Engineering, Texas A&M University, College Station, TX 77843, USA 2

3

Department of Chemistry, Texas A&M University, College Station, TX 77843, USA

Department of Materials Science and Engineering, Texas A&M University, College Station, TX 77843, USA

*

Correspondence: [email protected] (Partha P. Mukherjee); [email protected] (Sarbajit Banerjee)

1 ACS Paragon Plus Environment


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Abstract Transition metal oxides are usually used as catalysts in the air cathode of lithium-air (Li-air) batteries. This study elucidates the mechanistic origin of the oxygen reduction reaction (ORR) catalyzed by δ-MnO2 monolayers and maps the conditions for Li2O2 growth using a combination of first-principles calculations and mesoscale modeling. The MnO2 monolayer, in the absence of an applied potential, preferentially reacts with a Li atom instead of an O2 molecule to initiate the formation of LiO2. The oxygen reduction products (LiO2, Li2O2, and Li2O molecules) strongly interact with the MnO2 monolayer via the stabilization of Li—O chemical bonds with lattice oxygen atoms. As compared to the disproportionation reaction, direct lithiation reactions are the primary contributors to the stabilization of Li2O2 on the MnO2 monolayer. The energy profiles of (Li2O2)2 and (Li2O)2 nucleation on -MnO2 monolayer during the discharge process demonstrate that Li2O2 is the predominant discharge product, and that further reduction to Li2O is inhibited by the high overpotential of 1.21 V. Interface structures have been examined to study the interaction between the Li2O2 and MnO2 layers. This study demonstrates that a Li2O2 film can be homogeneously deposited onto -MnO2, and that the Li2O2/MnO2 interface acts as an electrical conductor. A mesoscale model, developed based on findings from the first-principles calculations, further shows that Li2O2 is the primary product of electrochemical reactions when the applied potential is smaller than 2.4 V. Key words: Lithium-air battery cathode; oxygen reduction reaction; disproportionation reaction; Li2O2 growth kinetics; first-principles calculation; mesoscale modeling.


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Introduction Vehicle electrification has attracted tremendous recent attention as a means of reducing the energy footprint of the transportation sector and mitigating environmental pollution induced by the consumption of fossil fuels. Electric vehicles (EVs) powered by conventional lithium-ion batteries (LIBs) have already been commercialized. The bottleneck for further EV development is that the specific energy density of intercalation cathode materials is yet to exceed 250 Wh/kg, which is imperative to meet range requirements 1-3. Beyond LIBs, Li-air batteries have a great potential to promote the development of EVs. Li-air batteries exhibit a high specific energy density of 11,140 Wh/kg, which approaches the specific energy density of the gasoline engine 4-6. In Li-air batteries, the active material is oxygen (O2), which is abundant in the atmosphere. For non-aqueous Li-air batteries, experimental and theoretical studies propose the following electrochemical/chemical reactions during the discharge process 6-9: O (g) + Li (sol. ) + e = LiO , (1) LiO + Li + e = Li O ,

(2)

2LiO = Li O + O (g),

(3)

The first step is the electrochemical reduction of O2 to superoxide LiO2; subsequently, LiO2 can be further reduced to Li2O2 (Eq. 2), or be disproportionated into Li2O2 and O2 (Eq. 3). Theoretically, Li2O2 can be further reduced to Li2O, but only Li2O2 is observed in Li-air batteries with the discharge voltage cut-off above 2.0 V

10-11

formation of Li2O is electrochemically irreversible

12

. McCloskey et al. have found that the

. Therefore, it is necessary to inhibit the

formation of Li2O during the discharge process. 3 ACS Paragon Plus Environment


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One key challenge for Li-air batteries is that the discharge product Li2O2 is insoluble and acts as an electrical insulator. The primary mechanisms of charge transport have been predicted to be hole tunneling with hole polaron hopping thought to assume greater importance at higher temperatures

13

. Point defects such as Li vacancies

14

, Si dopants

15

, and Co dopants

16

can

activate crystalline Li2O2, mitigate polaron formation, and yield “in-gap” conduction states, thereby enhancing electrical conductivity. However, introducing these point defects within crystalline Li2O2 is energetically unfavorable. Radin and Sigel have investigated the mechanisms of charge transport in Li2O2, and found that changes in the charge state of O2 dimers control the defect chemistry and conductivity of Li2O2; Li vacancies, as well as small hole polarons, have been identified as the dominant charge carriers 17. Viswanathan et al. have found that the Li2O2 film thickness determines its conductivity by using both experimental and computational techniques

18

. Subsequent work from Siegel and co-workers has shown that the (0001) Li2O2

surface is half-metallic and indeed it is this surface that dominates the facets of crystalline Li2O2 19

. Such a facet selectivity implies a key role for interfacial charge transport and thus suggests

that particle morphology of the discharged products is of considerable importance. In addition, Li2O2 morphology also affects the overpotential of the oxygen evolution reaction (OER) during the charging process, likely as a result of the intrinsic conductivity and defect tolerance of specific crystallographic planes.20-21 It has been found that small Li2O2 particles with shorter paths require lower charge overpotential as compared to larger particles. As mentioned above, the surface of crystalline Li2O2 is metallic; consequently, such surfaces can provide electron transport pathways. Indeed, it has been argued that generating nanostructured polycrystalline Li2O2 films that expose the appropriate crystallographic facets on the cathode substrate can improve the performance of Li-air batteries.22 For instance, Xiao et al. have found that defects in


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graphene are beneficial for the formation of the Li2O2 nanoislands and can improve the performance of Li-air batteries 11. Yilmaz et al. have dispersed RuO2 nanoparticles as a catalyst on a carbon nanotube cathode to promote the formation of Li2O2 crystalline films; the catalyst can lower the charge overpotential owing to the large contact area between Li2O2 and the cathode surface.23 Gent et al. have theoretically studied the interaction and reaction mechanisms of LixOy species on a rutile RuO2 surface, and have found that the formation of nanocrystalline Li2O2 film on the RuO2 surface is favored over that on the surface of the carbon cathode.22 Recently, firstprinciples calculations have been employed to evaluate monolayer RuO2 as a potential catalytic material for non-aqueous Li-O2 batteries.24 As compared to rutile RuO2, monolayer RuO2 can further reduce the overpotential corresponding to the formation of Li2O2 species. Recently, Yang et al. have found that CeO2 can facilitate Li2O2 film growth during discharge.25 In addition, the Li2O2 films provide the benefits of achieving higher reversibility and faster Li2O2 reduction during charge as compared to toroidal Li2O2; this is likely a result of interfacial conductivity achieved within continuous thin films. Nazar and her colleagues have also recently reviewed several classes of ORR and OER catalysts26. Recently, δ-MnO2 nanosheets have been added as a catalyst to graphene-based Li-air battery cathodes to control Li2O2 growth and to achieve superior cycling stability and low overpotential

27-28

. The catalytic mechanism of δ -MnO2 remains to be elucidated. The

fundamental motivation for this article arises from the fact that design principles for cathodes of Li-air batteries are essentially unknown in the absence of mechanistic understanding of the formation of discharge products and the interfacial structure stabilized between the discharge products and cathode material. This substantial knowledge gap is a direct reflection of the inadequacies of experimental methods in probing the properties of interfacial structures under



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operando conditions. In this study, a mesoscale simulation approach is proposed to elucidate the effects of δ-MnO2 on ORR and to investigate the properties of the Li2O2/MnO2 interface. As a first step, the stable structures for Li, O2, and LixOy molecules adsorption on δ-MnO2 have been determined by calculating the adsorption energies, and the adsorbate-substrate interaction mechanisms are elucidated via analysis of the interfacial electronic structure. Based on findings from the first step, two reaction pathways are evaluated to explore the kinetic feasibility of discharge product formation. Finally, a kinetic Monte Carlo model is developed based on the outputs from the previous steps, and allows for the primary discharge products to be mapped to specific electrochemical windows. Computational Details A first-principles approach based on density functional theory (DFT) 29-30 within the plane wave basis set

31-32

is used to study interactions between -MnO2 and discharge products in Li-Air

battery cathodes. All first-principle calculations are performed using Vienna ab-initio Simulation Package (VASP)

29, 33

. The Perdew-Burke-Ernzerhof (PBE) functional

augmented wave (PAW) method

35-36

34

and the projector

are employed to describe the electron—electron exchange

correlations and electron-ion interactions, respectively. The cut-off energy of the plane wave basis set is set to 400 eV in this study. The Hellman-Feynman forces are less than 0.02 eV/Å for optimization of atom positions. Since the 3d shell of Mn atom is not fully occupied, the GGA+U method developed by Dudarev et al. is used to describe the strong electron correlations value of Ueff is set to 3.9 eV according to previous theoretical studies

38-40

37

. The

. Spin polarization is

also considered in the present study given the ferromagnetism intrinsic to -MnO2 40. The adsorption of Li, O2, or LixOy (x, y = 1, 2) on -MnO2 is modeled by placing the adsorbate onto a (2 × 2) δ-MnO2 monolayer. -MnO2 slabs are separated by 16 Å vacuum along 6 ACS Paragon Plus Environment

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the z direction to avoid the interaction between two consecutive images. A 6 × 6 × 1 k-mesh is generated in the Brillouin zone (BZ) by the Monkhorst-Pack (MP) technique 41. The effect of van der Waals’ interactions is also considered by applying DFT-D2 dispersion correction 42, which is known to allow for better optimization of layered structures 43. Results and Discussion It is important to note that the calculations in this work are not entirely conjectural but instead relate directly to experimental findings. The atomistic structure of -MnO2 in this work is constructed based on previous experimental measurements. The space group of -MnO2 is P63/mmc 44, and the experimentally measured lattice parameter ranges from 2.84—2.88 Å 4445

. In the present study, the calculated lattice parameter is =2.89 Å, which agrees well with the

experimental results. The present simulation shows that the MnO2 monolayer is ferromagnetic with a 3.411 magnetic moment on each Mn atom and a magnetization of -0.196 on each O atom, which agrees well with previous theoretical studies of this structure.40 To evaluate the interaction between the adsorbate and the substrate, the adsorption energy is calculated as = Here,

!" #$ @%&#'

!" #$ @%&#'

*

− %&#' − ) ! − #' . (4)

is the total energy of the system, %&#' is the energy of a clean MnO2

monolayer, ! is the energy per Li atom for metallic Li crystallized in the body-centered cubic (bcc) structure, and #' is the energy of an isolated O2 molecule. A negative indicates an exothermic reaction and attractive interaction between the adsorbate and substrate.



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The adsorption energies of Li and O2 on -MnO2 are listed in Table 1 and the corresponding atomistic structures are shown in Figure 1. There are three typical Li adsorption sites on the MnO2 monolayer: (1) an O-top site, which implies that the Li atom is directly placed atop an oxygen atom in the upper sublayer; (ii) a Mn-hollow site, which implies that the Li atom is placed at a hollow site above a Mn atom of the middle sublayer; and (iii) an O-hollow site, which means the Li atom is placed at a hollow site situated directly above an O atom of the bottom sublayer. Our simulations demonstrate that O-hollow site is the most favored for Li adsorption and the adsorption energy of Li on MnO2 is -2.76 eV, which indicates a strong attractive force. When a Li atom is adsorbed at the O-hollow site, it can coordinate with three oxygen atoms to form Li—O bonds. In addition, the Li atom suffers weaker Columbic repulsions because the O-hollow site is situated relatively further away from the positively charged Mn atoms. A previous study suggests that the adsorption energy of Li on pristine graphene is 0.10 eV, which implies that there is no strong chemical interaction between Li and graphene. Hence, dopants need to be introduced to graphene to capture Li for further electrochemical reduction reactions during discharge 46. Geng et al. have calculated the adsorption energies for binding of Li atoms to the surface of several transition metal oxides (TMO2 with TM = Ti, Ru, Sn) with the rutile structure

22

. These authors also found strong attractions between Li and TMO2. The

adsorption energy of Li on -MnO2 monolayer is comparable to the adsorption energy (-2.77 eV) calculated for the RuO2 (110) surface 22. The optimized atomistic structure of Li on the -MnO2 monolayer shows that the bond length between Li and O atoms in the substrate varies from 1.86 Å to 1.94 Å, which is very close to the Li-O bond length, 1.94 Å, of crystalline Li2O2. Goodenough and his colleagues studied the crystal structure of Li2MnO2, and found that the Li-O 8 ACS Paragon Plus Environment

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bond length is also around 1.90 Å 47. The adsorption of O2 on the MnO2 monolayer only releases 0.06 eV per molecule, which indicates that O2 interacts with -MnO2 monolayer via weak van der Waals’ interactions rather than covalent hybridization. Bader charge analysis 48 demonstrates that there is no electron migration between the adsorbed O2 and the -MnO2 substrate. The O-O bond length in Figure 1(b) is 1.23 Å, which is only 0.02 Å longer than the 1.21 Å bond-length of O2 in the gas phase 49. The slight variation of the O-O bond length also indicates the absence of a strong interaction between O2 and the -MnO2 monolayer. Table 1. Energetic and geometric parameters of reactants (Li and O2) and products (LixOy with x, y = 1, 2) adsorbed on the -MnO2 monolayer. represents the adsorption energy; + !# represents the length of the bond between Li atoms and O atoms in the MnO2 monolayer; , !# represents the Li-O bond length in the adsorbate; and ,## represents the O-O bond length in the adsorbate. Only the lowest energy configuration are listed for each adsorbate. adsorbate

(eV)

+

Li

-2.76

1.86~1.94

--

--

O2

-0.06

--

--

1.23

LiO2

-3.01

1.93~2.00

2.14

1.25

Li2O2

-5.60

1.81~2.10

2.06

1.28

Li2O

-4.24

2.10~2.15

1.84~1.85

--

!#

(Å)

,

!# (Å)

,## (Å)

The adsorption of single LixOy (x, y = 1, 2) molecules has also been examined. The most stable configurations are demonstrated in Figure 1, and the corresponding adsorption energies are listed in Table 1. Based on their negative adsorption energies, it can be inferred that LixOy molecules interact with the MnO2 monolayer via strong chemical bonds. Li2O2 adsorption energy is -5.60 eV, followed by Li2O adsorption (-4.24 eV) and LiO2 adsorption (-3.01 eV). The formation of Li2O2 is the most energetically favored, which is confirmed by the experimental observation that Li2O2 is the predominant discharge product. The atomistic structures of LixOy on 9 ACS Paragon Plus Environment


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-MnO2 are demonstrated in Figure 1(c)—(e). In the most energetically stable relaxed structure, the Li atoms bridge LixOy adsorbates and substrates via Li-O bonds. The bond length (+

!# )

is

around 2 Å as shown in Table 1. A previous theoretical study reported that the adsorption energies of LixOy on pristine graphene are around -0.30 eV 50, which suggests weak interactions between the adsorbate and the substrate. It is well recognized that wetting of the Li2O2 film on the substrate requires a strong attractive interaction at the Li2O2/cathode interface;22,

51

consequently, obtaining a strongly

wetted Li2O2 thin film on graphene remains a formidable challenge. In the present study, we find that the adsorption energy of Li2O2 on the MnO2 monolayer is -5.60 eV. The results presented here suggest the strong attraction between Li2O2 and -MnO2 will facilitate effective wetting of the -MnO2 substrate by the Li2O2 film. Difference charge densities (Fig. 2) have been analyzed to understand the interaction mechanism between LixOy and the MnO2 monolayer. It is clear that electron accumulation regions appear between Li atoms in adsorbates and O atoms in substrates, and these regions represent strong chemical bonds. In this study, the thermodynamics of the conversion of LiO2 to Li2O2 is first calculated without considering an applied potential. There are two reaction paths for the formation of Li2O2. One reaction path is that the adsorbed LiO2 is directly lithiated to Li2O2 (Eq. 1). A viable alternative is the disproportionation reaction which extracts one O2 molecule from (LiO2)2 cluster to yield Li2O2 as per Equation 2. The energy of the lithiation reaction is calculated as Δ .!/0 !' #' =

!' #' @%&#'

−

!#' @%&#'

and the disproportionation reaction energy is calculated as 10 ACS Paragon Plus Environment

− ! , (5)

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Δ !2 !' #' =

!' #' @%&#'

+ #' − (

!#' )' @%&#' .

(6)

According to the present simulations, it is found that the lithiation reaction energy Δ .!/0 !' #' is 2.59 eV per Li2O2 molecule, and the disproportionation reaction energy Δ !2 !' #' is 0.33 eV per Li2O2 molecule (Figure 3). Hence, from a thermodynamic perspective, the formation of Li2O2 on -MnO2 monolayer is expected to be dominated by the exothermic lithiation reaction rather than the endothermic disproportionation reaction. The reaction energy of Li2O2 to Li2O has also been examined. Two reaction pathways are proposed: in the first reaction pathway, a Li2O2 cluster is directly lithiated to a (Li2O)2 cluster; the second pathway involves a disproportionation reaction wherein an O2 molecule is extracted from a (Li2O2)2 cluster. Analogous to the formation of Li2O2, the lithiation reaction energy per molecule is calculated as 3

Δ .!/0 !' # = ((

!' #)' @%&#'

−

!' #' @%&#'

− 2 ! ), (7)

and the disproportionation reaction energy per molecule is calculated as Δ

!2 !' #

3

= ((

!' #)' @%&#'

+ # − (

!' #' )' ).

(8)

For the formation of Li2O, the lithiation reaction energy Δ .!/0 !' # is -1.75 eV per Li2O molecule, which indicates that the lithiation reaction for Li2O formation is thermodynamically favored. The disproportionation reaction energy for Li2O formation Δ !2 !' # is 1.45 eV. The results indicate that formation of Li2O is strongly disfavored in terms of thermodynamics and that any Li2O that is formed should spontaneously be converted to Li2O2.



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In order to further understand the initial nucleation of Li2O2 on -MnO2 during the discharge process, the energy profile is then calculated with consideration of the applied potential based on Δ =

!"4 #$4 @%&#'

−

!"456 #$456 @%&#'

3

− (7& − 7&3 )#' − (7& − 7&3 )( ! − 89). (9)

Here, xn and yn are the number of Li and O atoms in the nth reaction step. 9 is the applied potential and e is the elementary charge. The energy profiles for (Li2O2)2 and (Li2O)2 formation on the -MnO2 monolayer are shown in Figure 4. The top panel in Figure 4 shows the energy profile with the applied potential 9 = 2.47 V, which is the highest potential that still makes the free energy diagram of (Li2O2)2 formation downhill. In this case, the energy decreases till (Li2O2)2 appears on the substrate; the formation of (Li2O)2 in the next reaction step corresponds to a significant energy increase. In experiments, electrochemical performance curves also demonstrate that the discharge potential is around 2.5 V with a discharge current of 0.333 mA/cm-3 for a δ-MnO2/graphene composite cathode

27

. Because a further reduction to (Li2O)2

significantly increases the total energy of the system, the formation of Li2O is thermodynamically unviable with 9 = 2.47 V. The equilibrium potential (9 ) of 2Li + 2e + O = Li O is 2.96 V

52

. Hence the overpotential for the Li2O2 formation on -MnO2 can be

estimated by = = 9 − 9. (10) The overpotential for initial (Li2O2)2 nucleation on the -MnO2 monolayer is 0.49 eV, which is much less than the overpotential of Li2O2 nucleation on pristine graphene and nitrogen-doped graphene substrates

46

, and approaches the theoretical overpotential of Li2O2 nucleation on pure

crystalline Li2O2 surfaces.9 Therefore, -MnO2 strongly abets Li2O2 film formation during 12 ACS Paragon Plus Environment

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discharge. The previous experimental study also reported that the graphene/-MnO2 cathode can provide a higher discharge voltage (lower overpotential) than the graphene cathode.27 The bottom panel in Figure 4 shows the energy profile with applied potential 9 = 1.75 V, which is the highest potential to make the free energy diagram of (Li2O)2 formation downhill. In this case, the reduction of Li2O2 to (Li2O)2 is not an endothermic reaction. However, the formation of Li2O requires a high overpotential 1.21 eV, which indicates that δ-MnO2 can inhibit the formation of Li2O during the discharge process. This study has benchmarked theoretical calculations to experimentally measured values, proposed mechanistic pathways, and finally provided other experimentally measurable parameters that can be used for validation. Indeed, the overpotential value can be provided as a measurable boundary condition for future experimentation. Wetting of δ-MnO2 by the Li2O2 layer requires a strong attractive force between these two layers. Previous theoretical and experimental studies have demonstrated that crystalline Li2O2 prefers to anisotropically grow along the (0001) surface

8, 53

. The atom layer sequence of

Li2O2 (0001) surface structure is …-OLiO-Li-OLiO-Li-OLiO-… as shown in Figure 5(a). Our simulations demonstrate that the lattice mismatch between the -MnO2 monolayer and Li2O2 (0001) surface is about 6%. Based on these findings, two interface structures have been evaluated in this study: (i) an OLiO layer is placed on the MnO2 layer as shown in Figure 5(b), and (ii) an OLiO-Li layer is placed on the MnO2 layer as shown in Figure 5(c). The interfacial energy between two layers is calculated as 3

!&/? = @ ( where A is the interface area,

%

%

− − % ), (11)

is the total energy of the interface structure. is the energy of

the OLiO (or OLiO-Li) layer, and % is the energy of the -MnO2 layer. A negative value of 13 ACS Paragon Plus Environment


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!&/? represents an attractive force. When calculating and % , atom positions are the same as in the interface structure. The interfacial energy of OLiO/MnO2 is found to be only −0.01 eV/Å2, whereas that of OLiO-Li/MnO2 is -0.27 eV/ Å2. It is thus clear that the latter OLiO-Li/MnO2 interface is more energetically stable. For the OLiO/MnO2 interface, atomistic structure analysis shows that the gap between the OLiO layer and MnO2 layer is 2.56 Å. For the OLiO-Li/MnO2 interface, the distance between the OLiO sublayer and MnO2 is reduced to 2.45 Å, and Li atoms serve as the bridge between the Li2O2 and MnO2 components of the heterostructure. The bond length between bottom Li atoms and O atoms from MnO2 is around 2.18 Å, which is slightly longer than the Li-O bond in bulk Li2O2. The current study demonstrates that the formation of a Li2O2 film with a (0001) surface index on a MnO2 monolayer is energetically favored. The interface between the two layers is stabilized by the strong attraction between the MnO2 and Li2O2 layers. As noted above, it is thought that the formation of a strongly wetting Li2O2 film can lower the charge voltage as a result of the large area of contact between the MnO2 substrate and the discharge product. Indeed, a recent experimental study also reported that the graphene/-MnO2 cathode requires a lower charge voltage as compared to a graphene cathode. 27 The electronic structure of the OLiO-Li/MnO2 interface (Figure 6) has been analyzed to further understand the interfacial interaction. The difference charge density shows that electron accumulation regions (yellow isosurface in Figure 6(a)) appear between the Li2O2 layer and MnO2 layer, which further indicate that strong Li-O chemical bonds form between these two layers. According to the density of states (DOS) plot shown in Figure 6(a), it is found that Li 2s/O 2p hybridization makes the main contribution to Li-O bond formation and the hybridization states distribute in the range of -7.5 to -5.0 eV below the Fermi level. The OLiO-Li/MnO2 14 ACS Paragon Plus Environment

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interface is further an electrical conductor. It is interesting that the conduction mechanism in the MnO2 layer is quite distinct from that of the Li2O2 layer. The Li2O2 layer is a p-type conductor since the Fermi level goes across the valence band of the Li2O2 layer. A previous theoretical study reported that the pristine MnO2 monolayer is a semiconductor

40

. The total DOS of the

pristine MnO2 monolayer calculated in our work (Figure 6b) clearly shows that Fermi level crosses the gap between the valence and conduction bands. It is interesting that the -MnO2 monolayer is metallic in the Li2O2/MnO2 interface structure. Figure 6(a) clearly shows that Fermi level goes across the conduction band of the -MnO2 layer. The present study reports two reaction mechanisms leading to Li2O2 formation on the MnO2 monolayer. The first mechanism involves an electrochemically induced reduction reaction with a reaction energy that is dependent on the discharge voltage. The second mechanism is the disproportionation reaction with a reaction barrier that reflects intrinsic thermodynamics of the materials. Although -MnO2 has been used as the cathode of Li-O2 batteries, a clear understanding of the role of these mechanisms remains to be elucidated. To this end, a mesoscopic kinetic Monte Carlo (KMC) model has been developed to investigate the competition between the ORR and disproportionation models for Li2O2 formation on δ-MnO2. The rates of each of the electrochemical reaction steps (Eq. 1 and Eq. 2) can be estimated using the Butler—Volmer equation: A = A max E1, 8)G E−

H(IJ KL) MN

O O. (12)

Here k0 is the pre-exponential coefficient which is conventionally assumed to be the same for each electrochemical reaction when modeling Li2O2 formation


18, 54

. P is Boltzmann constant,


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and T is the temperature. In this model, Q = 1 is used considering only thermodynamic limitations 18. The rate of the disproportionation step (Eq. 3) can thus be estimated as: A = A exp E−

SL MN

O. (13)

Based on Eqs. (12) and (13), the two parameters that predominantly affect Li2O2 formation are the applied potential 9 and temperature T. Figure 7 shows the fraction of ORR produced Li2O2 on -MnO2. It can be seen that 9 is the predominant factor that affects the reactions occurring on the surface. Figure 7 clearly shows that Li2O2 is only produced by the ORR reaction when 9 < 2.4 V. In this potential window, the alternative electrochemical reactions (Eqs (1) and (2)) are always exothermic ( 8Φ + Δ ≪ 0 ). In other words, the disproportionation reaction is suppressed under electrochemical conditions. The reaction kinetics are sensitive to the value of Φ in the window Φ ∈ (2.4,2.6) V. In this region, the electrochemical reaction of Li2O2 formation (Eq. (2)) is gradually converted from the exothermic to the endothermic with an increase of the applied potential Φ. According to Eq. 12, it is clear that the rate of Li2O2 formation monotonically decreases and approaches k0. It should be noted that the rate of the disproportionation reaction is a constant that is independent of the applied potential. Therefore, the LiO2 pathway begins to make a greater contribution to the disproportionation reaction as the applied potential rises to 2.6 V. When Φ > 2.6 V, the formation of Li2O2 is endothermic and the reaction rate is a constant according to Eq. 12. Hence the ORR coverage does not change when the applied potential is larger than 2.6 V. It should be noted that the pre-coefficients of all reactions (electrochemical reactions and the disproportionation reaction) are simply assumed to be k0 = 1010 s-1 site-1 for the purposes of this discussion. Under a realistic set of conditions, the pre-coefficient values will be determined by the properties of the catalyst, the chemical nature of


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the electrolyte, and other physicochemical interplays. All of these factors can determine the shape of the coverage curve (e.g., the limiting coverage when Φ > 2.6 V). Conclusions Theoretical calculations in this study provide a mechanistic view of the interface between the MnO2 monolayer and the discharge products and suggest reasons for how the former improves the performance of Li-air batteries. The self-consistent set of ideas developed in this study as a potential mechanistic pathway would have been impossible to gauge entirely by experiments given the absence of in situ probes of sufficient chemical and interfacial resolution. Indeed, this ability to develop a detailed picture of the potential energy landscape and to extrapolate rigorously to kinetic barrier is what makes theory such a powerful tool for investigating interfacial phenomena. In this work, a first-principles approach is used to study the interactions between discharge products and δ-MnO2 monolayers that serve as putative Li-air battery cathodes. The adsorption energies of Li, O2, and LixOy (x, y = 1, 2) on MnO2 are calculated. Based on an evaluation of the adsorption energies, it is found that the MnO2 monolayer prefers to react with Li atoms rather than O2 molecules to initiate LiO2 formation. The incipient LiO2 is expected to be strongly chemisorbed onto the MnO2 layer. LiO2 can be subsequently converted to Li2O2 and finally Li2O via a sequence of lithiation or disproportionation reactions. According to the adsorption energy calculated in the present work, it can be inferred that the MnO2 monolayer has a strong affinity for LiO2 and Li2O2 species, which facilitates the formation of a nicely wetted Li2O2 film on the cathode surface. The calculated reaction energies show that the lithiation reactions are exothermic and thus thermodynamically favored over the disproportionation reactions that are endothermic. Although the formation of Li2O is energetically favored in the 17 ACS Paragon Plus Environment


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absence of an applied potential, Li2O formation on the MnO2 monolayer is inhibited by 1.21 V overpotential during the discharge process. Li2O2 formation on the -MnO2 monolayer requires only a 0.49 V overpotential. Consequently, Li2O2 is expected to be the predominant species in discharge products. The low calculated overpotential suggests that the MnO2 monolayer helps the battery to achieve a higher discharge voltage. Evaluation of Li2O2/MnO2 interfacial interactions indicates that strong Li-O bonds bridge these two layers together. Electronic structure studies further indicate that the interfacial structure is a metal and thus can facilitate charge transport. Based on the DFT derived parameters, a mesoscale KMC model is developed to study the effects of applied potential and temperature on the reaction kinetics occurring at -MnO2 surfaces. Our results demonstrate that Li2O2 formation is sensitive to applied potential. The electrochemical reaction makes the predominant contribution to Li2O2 formation when the applied potential is lower than 2.4 V. Acknowledgements Financial support from the Texas A&M Energy Institute is gratefully acknowledged. Supercomputing resources from Texas A&M University High Performance Computer Center are also acknowledged. L.R.D.J. acknowledges support from a National Science Foundation Graduate Research Fellowship under Grant 1252521.


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List of Figures Figure 1. Atomistic structures of (a) Li, (b) O2, (c) LiO2, (d) Li2O2 and (e) Li2O adsorbed on MnO2 monolayers. The first row shows the top view, whereas the second row indicates the side view. Dark blue spheres represent Mn atoms; violet spheres represent Li atoms; red spheres represent O atoms in the substrate; and cyan spheres represent O atoms in the adsorbate. Figure 2. Difference charge densities of (a) Li, (b) LiO2, (c) Li2O2, and (d) Li2O adsorbed on a MnO2 monolayer. The yellow isosurface (0.005 |e|/Å3) indicates the electron accumulation region, whereas the cyan isosurface (0.005 |e|/Å3) indicates the electron depletion region. Blue spheres, violet spheres, and red spheres represent Mn atoms, Li atoms, and O atoms, respectively. Figure 3. Reaction energies for the formation of Li2O2 and Li2O by different reaction mechanisms. The positive reaction energy indicates an endothermic reaction, whereas the negative reaction energy indicates an exothermic reaction. Figure 4. Energy profiles for the nucleation of (Li2O2)2 and (Li2O)2 with a discharge potential of Φ = 2.47 V and Φ = 1.75 V. The solid black line represents the electrochemical reaction path culminating in the formation of (Li2O2)2, whereas the dashed red line represents the electrochemical reaction culminating in the formation of (Li2O)2. Φ = 2.47 V is the highest potential to make the free energy diagram of (Li2O2)2 formation downhill. Φ = 1.75 V is the highest potential that makes the free energy diagram of (Li2O)2 formation downhill. Figure 5. Atomistic structures of (a) the layer sequence of the Li2O2 (0001) surface, (b) the OLiO/MnO2 interface structure, and (c) the OLiO-Li/MnO2 interface structure. Figure 6. (a) Electronic structure of the OLiO-Li/MnO2 interface. The left plot demonstrates the projected density of states for atoms in different sublayers; red curves represent spin-down states and black curves represent spin-up states. The right graphic demonstrates the difference charge density at the interface. The yellow isosurface (0.005 |e|/Å3) represents the electron accumulation region, whereas the cyan isosurface (0.005 |e|/Å3) represents the electron depletion region. (b) Total density of states of the pristine MnO2 monolayer. The Fermi energy level crosses the gap between the valence band and the conduction band, which indicates that pristine MnO2 monolayer is not a conductor. Figure 7. The population of ORR-produced Li2O2 as a function of applied potential.



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Figure 1. Atomistic structures of (a) Li, (b) O2, (c) LiO2, (d) Li2O2 and (e) Li2O adsorbed on MnO2 monolayers. The first row shows the top view, whereas the second row indicates the side view. Dark blue spheres represent Mn atoms; violet spheres represent Li atoms; red spheres represent O atoms in the substrate; and cyan spheres represent O atoms in the adsorbate.


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Figure 2. Difference charge densities of (a) Li, (b) LiO2, (c) Li2O2, and (d) Li2O adsorbed on a MnO2 monolayer. The yellow isosurface (0.005 |e|/Å3) indicates the electron accumulation region, whereas the cyan isosurface (0.005 |e|/Å3) indicates the electron depletion region. Blue spheres, violet spheres, and red spheres represent Mn atoms, Li atoms, and O atoms, respectively.



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Figure 3. Reaction energies for the formation of Li2O2 and Li2O by different reaction mechanisms. The positive reaction energy indicates an endothermic reaction, whereas the negative reaction energy indicates an exothermic reaction.


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Figure 4. Energy profiles for the nucleation of (Li2O2)2 and (Li2O)2 with a discharge potential of Φ = 2.47 V and Φ = 1.75 V. The solid black line represents the electrochemical reaction path culminating in the formation of (Li2O2)2, whereas the dashed red line represents the electrochemical reaction culminating in the formation of (Li2O)2. Φ = 2.47 V is the highest potential to make the free energy diagram of (Li2O2)2 formation downhill. Φ = 1.75 V is the highest potential that makes the free energy diagram of (Li2O)2 formation downhill.



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Figure 5. Atomistic structures of (a) the layer sequence of the Li2O2 (0001) surface, (b) the OLiO/MnO2 interface structure, and (c) the OLiO-Li/MnO2 interface structure.


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Figure 6. (a) Electronic structure of the OLiO-Li/MnO2 interface. The left plot demonstrates the projected density of states for atoms in different sublayers; red curves represent spin-down states and black curves represent spin-up states. The right graphic demonstrates the difference charge density at the interface. The yellow isosurface (0.005 |e|/Å3) represents the electron accumulation region, whereas the cyan isosurface (0.005 |e|/Å3) represents the electron depletion region. (b) Total density of states of the pristine MnO2 monolayer. The Fermi energy level crosses the gap between the valence band and the conduction band, which indicates that pristine MnO2 monolayer is not a conductor. 31 ACS Paragon Plus Environment


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Figure 7. The population of ORR-produced Li2O2 as a function of applied potential.


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TOC graphic:



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Schematic illustration of a catalyzed interface for Li2O2 formation.


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Mechanistic Evaluation of Li x O y Formation on Î´-MnO2 in

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