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Capture Lithium in αMnO2: Insights from First Principles Chen Ling* and Fuminori Mizuno Toyota Research Institute of North America, 1555 Woodridge Avenue, Ann Arbor, Michigan 48105, United States S Supporting Information *

ABSTRACT: Here we report the study of Li and Li oxides insertion in αMnO2 with firstprinciples density functional theory calculations. For Li insertion, the redox reaction is characterized with a particular order of the reduction of Mn ions. It induces asymmetric changes of lattice parameters through Jahn−Teller distortion of MnO6 units. At composition of LiMn2O4, the ratio between lattice parameter a and b extends to the highest value, 1.19. The severe structural deformation is concluded to be the major cause of the irreversible capacity when αMnO2 is used as Li-ion battery cathodes. On the other hand, the structure of αMnO2 is stable to reversibly insert and remove Li oxides. Li oxides inserted αMnO2 is half metallic instead of insulating. On the basis of our results, a mechanism is proposed to explain the role of αMnO2 in the catalyzed decomposition of Li2O2. The key step in this mechanism is the insertion of Li oxides in αMnO2. The insertion transfers electrons from Li oxides to MnO2 and partially oxidized Li oxides, making the subsequent release of O2 easier. KEYWORDS: αMnO2, Li-ion battery, Li-air battery, intercalation, capacity fading, Jahn−Teller effect, Li2O2 decomposition structure of αMnO2 was achieved with the initial capacity reaching as high as 360 mA h/g.18 The utilization of αMnO2 as Mg ion battery cathode was also reported.22 The discharge capacity of Mg inserted αMnO2 was reported to be about 280 mA h/g in the first cycle.22 αMnO2 was also studied as possible catalyst to improve the performance of oxygen electrode in Li-air batteries.19−21 Bruce and co-workers first showed that electrodes with αMnO2 nanowires provided much higher capacity and cyclability than other metal oxides when used as Li-air battery electrocatalyst.19 Later on, Giordani et al studied the decomposition of Li2O2 using various metal oxide electrocatalysts.20 They found αMnO2 is the best among the studied metal oxides.20 Although whether αMnO2 in Li-air battery electrodes catalyzes the decomposition of Li2O2 or in fact catalyzes deleterious side reactions is still under intensive studies,23−25 it is undoubtedly important to thoroughly investigate the role of αMnO2 in Li-air batteries. The great potential of αMnO2 in the electrochemical capture of Li raises the importance to improve the limited fundamental understanding about the reaction between Li and αMnO2. The lack of such knowledge challenges the efforts to improve the sustainability in the practical usage of αMnO2 as Li capture medium. For instance, as a rechargeable battery cathode, αMnO2 suffers from poor cyclability. In Li inserted αMnO2, the capacity dropped ∼10% of its second discharge capacity after 10 cycles,8 and ∼20−60% after 100 cycles.18 The capacity fading was even more serious for Mg insertion, in which it dropped ∼30% and ∼10% in the first and second cycle,

1. INTRODUCTION Polymorphs of MnO2 provide a library of plentiful attractive structures as the host for the insertion of cations. All MnO2 polymorphs vary in the network of MnO6 octahedra connected at the edge or corners with open spaces accommodating the insertion of cations. Hollandite MnO2, often regarded as αMnO2, contains a 2 × 2 edge shared MnO6 octahedral units (Figure 1a) and is stable to accommodate alkaline, alkaline

Figure 1. (a) Crystal structure of αMnO2. The pseudocubic oxygen cavity is shown with the dotted lines. (b). Possible binding sites in the pseudocubic cavity. Oxygen ions are denoted as red spheres while Mn ions are green spheres. Black, 2a; purple, 2b; yellow, 4e; blue, 8h; and orange, 8h′ sites.

earth and other cations.1−4 The capture of Li in αMnO2 is of particular importance for its potential application as supercapacitors,5−7 rechargeable Li-ion battery cathodes8−18 and Liair battery catalysts.19−21 The study of αMnO2 as rechargeable battery cathodes has been carried out for more than two decades.8,13,14 Recyclable insertion and removal of Li in the © 2012 American Chemical Society

Received: July 25, 2012 Revised: September 18, 2012 Published: September 24, 2012 3943

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respectively.22 The loss in capacity is generally believed to be associated with the instability of the host structure during electrochemical cycling. However, the detailed mechanism that causes the instability is still unclear.8,14 It leads the capacity fading to be an unsolved challenge in Li-ion battery research at current stage. In this paper, we perform a theoretical study for the insertion of Li and Li oxides into αMnO2 through first-principles Density Functional Theory (DFT) calculations. For Li insertion, we investigate the reaction between Li and αMnO 2 and demonstrate the redox reaction happens with an ordered reduction of Mn ions. The capacity fading that has been observed experimentally is explained as a result of severe structural deformation directly related to the ordered reduction of Mn. Through the study of the insertion of Li oxides, insights into the function of αMnO2 in Li-air batteries are provided with possible mechanisms to explain the enhanced performance of oxygen electrode.

also shown at low concentrations (x = 0.0625−0.25, see Figure S1 in the Supporting Information). The comparison between our predictions and the experiments is difficult due to the lack of precise XRD data about αLixMn2O4. Nonetheless, previous studies suggested that 8h or 8h′ sites are more favorable for small cations because the cavity of hollandite compounds is too big to stabilize them at the center 2a or 2b sites.36 Indeed, 8h site was speculated to be the most stable binding sites for Na+ ions.36 Because the size of Li+ is even smaller than Na+, it can be postulated that 8h (8h′) sites should be more stable than 2a (2b) sites, consistent with DFT calculated results. The stability of 8h sites over other interstitial sites can also be understood from the coordination of Li ions and Li−O in the channel. As listed in Table 1, at high symmetric 2a (2b) sites, Table 1. Coordination of Li Ion and Li−O Bond Lengths at Different Interstitial Sites after DFT Relaxation

2. COMPUTATIONAL METHODS DFT calculations were performed with the Vienna ab initio simulation package (VASP) using projector augmented waves (PAW) pseudopotentials and the exchange-correlation functionals parametrized by Perdew, Burke, and Ernzerhof for the generalized gradient approximation (GGA).26−28 Numerical convergence to less than 2.5 meV per MnO2 unit was ensured by using cutoff energy 550.0 eV and appropriate Gamma centered k-point mesh density of at least 0.03 Å−1. In the search for stable configurations, all ions were fully relaxed as well as the shape and the volume of the supercell. Traditional DFT-GGA calculations for strongly correlated compounds fail to characterize the localization of transition metal delectrons and give poor descriptions their electronic, magnetic and energetic properties.29,30 An efficient way to treat the static correlations is the GGA+U method by introducing a Hubbard type potential to describe the d-part of the Hamiltonian. In all our calculations we employed GGA+U approach with U−J = 3.9.31,32 Previous reports using similar U values showed good agreements with experiments for various Mn oxides and Li-inserted Mn compounds.33−35 Calculations with different U−J values were tested and we found our conclusions were not affected.

a

site

coordination

2a 2b 4e 8h 8h′

4 8 4 5 2a

Li−O bond length (Å) 2.45 2.88 2.44 2.07 2.03

× × × × ×

4 8 4 1, 2.30 × 2, 2.40 × 2 2

The next shortest Li−O bond length at 8h′ sites is 2.95 Å (× 2).

after relaxation Li is coordinated with four (eight) oxygen ions with Li−O bond lengths 2.45 (2.88 Å), too high to stabilize the small Li+ ion. At 8h sites, Li atom is coordinated in a pseudosquare pyramid environment with five Li−O bonds from 2.06 to 2.40 Å. The shortened Li−O bond length is thus helpful to stabilize the inserted Li ion at 8h site. We then investigate the reaction path for Li intercalation in αMnO2 using the convex hull method.37,38 A total of 74 symmetrically distinct configurations of αLixMn2O4 were included in our calculation with Li occupying 8h sites. Mn2O4 and Li2Mn2O4 were used as the reference state to calculate the formation energy, Ef, of LixMn2O4 as Ef = Ex − 0.5xE Li 2Mn2O4 − (1 − 0.5x)E Mn2O4

3. RESULTS 3.1. Li Insertion in αMnO2. As shown in Figure 1, the crystal structure of tetragonal of αMnO2 belongs to space group I4/m with tetragonal symmetric structure. The lattice constants for αMnO2 free of any stabilized ions are calculated as a = 9.772 and c = 2.891 Å, in good agreement with various experimental measurements.8,10,12 The 2 × 2 tunnel of αMnO2 is constructed by the stacking of oxygen squares (Figure 1a) along the c-axis. The tunnel wall consists of oxygen ions forming a connection of pseudocubic cavities (Figure 1b). The large size of the cavity provides multiple binding sites for cations. As illustrated in Figure 1b, five symmetrically distinct sites exist as denoted with Wyckoff position as 2a, 2b, 4e, 8h, and 8h′.36 The high symmetric 2a (2b) sites locate at the center of the oxygen square (cavity). 4e sites locate between 2a and 2b sites. 8h(8h′) sites deviate from 2a(2b) sites in the directions normal to the tunnel axis or c-axis. The site preference for the guest Li ion is explored in the dilute limit by calculating the energy of a single Li ion in distinct binding site of four unit cell of αMnO2 (concentration corresponding to x = 0.0625 in LixMn2O4). At the dilute limit, 8h site has the lowest binding energy, followed by the 8h′, 4e, 2a, and 2b sites (see Figure S1 in the Supporting Information). The preference on 8h site is

(1)

Here Ex, ELi2Mn2O4, and EMn2O4 is the total energy of LixMn2O4, Li2Mn2O4, and Mn2O4, respectively. Figure 2 shows the calculated formation energies for αLixMn2O4. All configurations are stable with the inserted ions relaxed at 8h sites. Up to one Li per MnO2 formula (four

Figure 2. Formation energies and the convex hull for LixMn2O4. 3944

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Li per cavity) the calculations converge to reasonable structures with negative binding energies relative to metallic Li. Thus the maximum concentration considered in this study is Li2Mn2O4, corresponding to the capacity 307.8 mA h/g. A recent report showed the initial capacity of Li intercalated αMnO2 could even reach 360 mA h/g.18 The excess capacity might come from the adsorption of Li ions on the surface.39 The convex hull of the formation energies shows the ground states of the inserted compounds. Besides the fully delithiated and lithiated phase five structures of intermedium Li concentration are included in the convex hull at x = 0.25, 0.375, 0.5, 0.75, and 1 for αLixMn2O4. The schematic configurations of the ground states are shown in Figure 3.

The enthalpy changes for conversion reaction are calculated to be −2.54 eV, smaller than the enthalpy of intercalation reaction. It confirms the voltage profile that was observed experimentally actually came from the intercalation of Li into αMnO2. In the study of rechargeable battery cathodes, the deformation of crystal structure is critical for cycling performance.42,43 Figure 4 shows the evolution of the lattice parameters

Figure 4. Average lattice constants (solid squares, a; and solid circles, b) and unit cell volume (open diamond) of LixMn2O4.

Figure 3. Schematic of the occupancy of 8h sites in the ground state structures of LixMn2O4. Each cavity is illustrated with the square corresponding to a Mn4O8 formula. The unoccupied and occupied 8h sites are shown with the open and solid circles, respectively.

and unit cell volumes of ground state LixMn2O4. The unit-cell volume gradually increases with Li concentration. From MnO2 to LiMn2O4, the volume expands by 7.4%, comparable to the discharge from FePO4 to LiFePO4. The volumetric expansion is achieved mostly by the change of lattice parameter a and b with nearly unaltered c during the lithiation (see Figure S2 in the Supporting Information). Interestingly, the changes on lattice parameter a and b are not symmetric. In tetragonal symmetric αMnO2, a and b axis are equivalent. The tetragonal symmetry is broken with the insertion of guest Li ions as evidenced by the difference between a and b values. In the stage I of the lithiation both a and b slightly increase by less than 1%. The crystal structure remains at tetrahedral or near-tetrahedral symmetry. In stage II, a increases rapidly and b gradually decreases. It lowers the symmetry of the structure to be orthorhombic (see Figure S3 in the Supporting Information). In stage III, the major change is the expansion along b-axis with nearly unaltered a. The fully lithiated Li2Mn2O4 recovers to a neartetrahedral symmetry with the difference between a and b less than 0.04 Å (0.36%). Figure 5 plots the ratio between lattice parameter a and b as a function of the average oxidation state of Mn ions. While in stage I, a/b increases by only 0.5%; stage II and III experience a dramatic increase and decrease in a/b, respectively. It reaches the maximum value, 1.19, at LiMn2O4. This value is comparable to the highest c/a ratio in the cubic-to-tetragonal deformation of spinel LixMnO2,15 which comes from the cooperative Jahn− Teller elongation of Mn−O bonds along the electronic zdirection.15 The highest c/a ratio in the structural deformation of spinel LixMnO2 was reported to be 1.16 at the composition LiMnO2.15 For Li insertion in MnO2, Jahn−Teller effect is directly related to the reduction of Mn ions from 4+ to 3+ because only Mn3+ (t2g3eg1) ion is the Jahn−Teller active species. To investigate the influence of Jahn−Teller distortion to the structural deformation of αMnO2, we first study the redox reaction between guest Li and host MnO2 by analyzing the

Depending on the number of Li ions inserted in the cavity, the intercalation is divided into three stages. In stage I from pristine MnO2 to Li0.5Mn2O4, 0−1 Li ion is inserted in each cavity. At composition Li0.5Mn2O4, all cavities are occupied by one Li ion. In stage II from Li0.5Mn2O4 to LiMn2O4, each cavity is occupied by 1−2 Li ions. To achieve the smallest electrostatic repulsion between Li ions, the second Li preferably occupies the 8h site the farthest from the first Li in the same cavity. The two occupied 8h sites align each other near parallel to one of the lattice direction, and perpendicular to the direction of the channel (c-axis). For convenience, we define the lattice direction that is parallel to two occupied (unoccupied) 8h sites to be a(b)-axis. The maximum Li concentration in this stage is LiMn2O4 with all cavities occupied by two Li ions. Further insertion in stage III occupies the rest of 8h sites. Because each cavity contains four 8h sites, at the end of Li intercalation, all these sites are occupied corresponding to the composition Li2Mn2O4. The voltage (vs Li/Li+) to electrochemically insert Li along the convex hull is calculated as40,41 v = −(Ex 2 − Ex1 − (x 2 − x1)E Li)/(x 2 − x1)e

(2)

Here Ex1, Ex2, and ELi are the total energy of Lix1Mn2O4, Lix2Mn2O4, and metallic Li, respectively. As shown in Figure 2, the voltage drops from 3.51 to 2.31 V during the lithiation. Experimentally, the voltage profile of αMnO2 showed a distint slope curve from approximately 3.7 to 2.5 V,13,14 in agreement with our predicted values. It is also important to evaluate the phase stability of intercalated compounds with other possible competing phases. Thus we also consider the conversion reaction as Li + 2MnO2 → 0.5Li 2O2 + Mn2O3

(3) 3945

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Mna3+ ions indicate the predominance of an elongated tetragonal Jahn−Teller distortion where dz2 orbital lowers its energy vs dx2−y2 orbital. Subsequently, lithiation from LiMn2O4 to Li2Mn2O4 reduces Mnb from 4+ to 3+, leading to an evolution of lattice parameter along the b-axis in the same manner. These results reveal a particular order of reduction of Mn ions, which explains the tetragonal-to-orthorhombic deformation as shown in Figures 4 and 5. In stage I and II, up to composition LiMn2O4, only Mna ions are reduced from Mna4+ to Mna3+. In stage I, the concentration of Jahn−Teller active Mna3+ is less than 50%, not high enough to induce strong lattice deformation. Therefore in stage I the deformation of the crystal lattice is not obvious. Only slight increase of a/b ratio is observed. Severe structural deformation appears after the concentration of Mna3+ is higher than 50% in stage II. The cooperative Jahn−Teller distortion elongates the structure along a-axis and shortens b-axis to compensate the elongation. The largest a/b reaches its highest value at LiMn2O4 where all Mna ions are reduced to 3+. Subsequent insertion of Li reduces Mnb ions in stage III and induces the elongation along b-axis. The whole process can be described with the following reactions stage I/II:

Figure 5. Change of the average ratio between lattice constant a and b as a function of the average oxidation state of Mn ions in the ground states of αLixMn2O4.

electronic states of Mn ions during the lithiation. As illustrated in Figure 6a, in αMnO2, each cavity is surrounded by MnO6

e− + Li+ + Mna 4 +Mnb 4 +O4 → Li 2Mna 3 +Mnb 4 +O4

(4)

stage III: e− + Li+ + lim na 3 + Mnb 4 +O4 → Li 2Mna 3 +Mnb 3 +O4

(5)

To further confirm the sequential order of reduction, we analyzed the charge transfer between Li and Mn/O using topological analysis of the total electron density within the atoms in molecules approach to get the Bader charge around individual nucleus.46,47 From MnO2 to LiMn2O4, the Bader charge of Mna decreases from +2.14e to +1.89e, whereas for Mnb ions, it decreases only to +2.10e. This change suggests the reduction of Mna ions with Mnb remaining its oxidation state.48 The Bader charges of Mnb ions decrease when LiMn2O4 is discharged to Li2Mn2O4, consistent with the speculated reduction of Mnb ions as in eq 5. Finally, we evaluate the mobility of Li ions in the channel of αMnO2. Unlike spinel-LiMn2O4, in which the interstitials to accommodate Li forms a 3D-connected network, the cavities of αMnO2 extend only in the direction of c-axis. For simplicity we only consider the 1D migration of Li along the channel at dilute concentration (Li0.0625Mn2O4). In this case, the migration barrier is calculated to be 0.47 eV, comparable to the barrier in its spinel counterpart. A detailed analysis of Li migration at nondilute concentration is still necessary to fully explore the kinetic limit of αMnO2 in Li-ion battery. For instance, it is possible that at high concentration the preoccupied cations would block the diffusion of Li ions through the channel. Nonetheless, we speculate that the kinetics of Li migration in αMnO2 should be comparable to that in the spinel phase at least at dilute concentration. 3.2. Li Oxide Insertion in αMnO2. The open channel of αMnO2 provides the possibility to the insertion of small molecules as well as cations.9 Experimental studies on the crystal structures of Li2O inserted αMnO2, xLi2O·MnO2, showed that Li ions preferred to occupy 8h′ sites with oxygen staying at or slightly deviating from 2b sites.8,10 As discussed earlier, in pristine αMnO2 Li prefers to occupy 8h sites instead

Figure 6. (a) Schematics of Mna (green) and Mnb (orange) in αMnO2. The dz2 electron cloud is illustrated with the blue color. (b− d) Partial density of states of Mna (red) and Mnb (blue) at the ground states of LixMn2O4 with x = 0, 1, and 2. The insertions show the distance between oxygen and Mna/Mnb ions (unit Å) and the Bader charges of Mn ions.

units connected with edge and corner shared oxygen ions. Mn ions can be categorized into two types as denoted as Mna and Mnb, of which the electronic dz2 orbital are nearly parallel to aand b-axis, respectively.44,45 Figure 6b−d shows the partial Density of States (pDOS) of Mn ions in the ground states of MnO2, LiMn2O4 and Li2Mn2O4. In MnO2 the electronic configurations of Mna and Mnb ions are clearly identical as a result of the tetragonal symmetry. With the insertion of Li, the tetragonal symmetry is broken and asymmetric characteristics are observed for Mna and Mnb ions. From MnO2 to LiMn2O4 the additional electrons are first added preferentially to the dz2 orbital of Mna sites and reduce the oxidation states of Mna from 4+ to 3+. The reduction induces cooperative Jahn−Teller distortion of MnaO6 units, as evidenced by the characteristic Mn−O distances. The longer Mn−O bonds along dz2 orbital of 3946

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of 8h′ sites. We first checked the preferable sites in xLi 2 O·MnO 2 for configurations with Li(O) occupying 8h′(2b) or 8h(2a) sites with x values from 0.03125 to 0.25. Consistent with the experimental results, for Li2O insertion Li(O) prefers to occupy 8h′(2b) sites with lower binding energies (see Figure S4 in the Supporting Information). We consider the electrochemical insertion of Li oxides with the formula LixOy (x ≤ 4y) into αMnO2 with the following reaction x e− + x Li+ + 0.5yO2 + MnO2 → LixOy · MnO2

(6)

The voltage of the intercalation can be calculated as V = −(E LixOy − xE Li − 0.5yEO2 − E MnO2)/x e

(7) Figure 7. Lattice parameters (Å) and unit-cell volume (v, Å3) of xLi2O·MnO2. Experimental results came from ref 15.

Here ELixOy, ELi, EO2, and EMnO2 is the energy for LixOy·MnO2, metallic Li, O2 molecule, and pristine MnO2, respectively. The reference energy for O2 molecule is calibrated using the voltage for Li2O2 decomposition (2.96 V vs Li/Li+). Table 2 lists the

αMnO2 remains nearly unchanged with only slight expansion along a- and b-axis (less than 3%). More interestingly, although the initial configuration used in DFT calculations was asymmetric in a and b direction with the insertion of Li2O, after relaxation the differences between lattice parameter a and b are most likely negligible. The largest difference between a and b is still less than 0.1 Å (0.8%) in LiMn2O4.5. This trend is different from the data discussed earlier for Li insertion in αMnO2. To compare, for LiMn2O4, a is about 0.6 Å (6%) larger than b. We then investigate the deformation of 0.25LixO·MnO2 with x varied from 1 to 4. The ratio between lattice parameter a and b is used to qualitatively describe the tetragonal to orthorhombic deformation as listed in Table 1. The maximum a/b ratio is calculated to be 1.06 as shown in Table 1. It is significantly lower than the values discussed earlier for the insertion of Li. The distortion of MnO6 is also partially suppressed when Li is cointercalated with oxygen. For instance, in Li2Mn2O4.5 the lengths of Mn−O bond, 2.25(2.11) Å, are smaller than that in LiMn2O4 or Li2Mn2O4 (Figure 6). These results demonstrate that αMnO2 is structurally stable to accommodate the insertion of Li oxides. Combining with the calculated voltage, we conclude that the insertion of Li oxides can happen reversibly during the operation of Li-air batteries. Interestingly, we find that the insertion of Li oxides in αMnO2 results in compounds with half-metallic characteristics. Pristine αMnO2, Li2O, and Li2O2 are all insulators with the band gap to be 0.7 eV (DFT calculations, Figure 8), approximately 8 eV,49 and around 2 eV,50 respectively. Figure 8 shows the total DOS for αMnO2 inserted with different level of Li oxides. The spindown states of the Li oxides inserted αMnO2 are insulating, whereas the spin-up states are conducting, suggesting the compound is actually half metallic. The majority of the occupied states near Fermi energy are composed of Mn 3d and framework O 2p orbitals, with minor contribution from the inserted O 2p orbitals (see Figure S5 in the Supporting Information). To understand the insulator-to-half-metallic transition in Li oxides inserted αMnO2, we examine the electron transfer between Li oxides and MnO2 with the Bader charge analysis. In pristine αMnO2, the Bader charges on Mn and oxygen ions are calculated to be +2.14e and −1.07e. When Li oxides is inserted, a small but non-negligible amount of electrons are transferred from Li oxides to framework ions. For example, in 0.25Li2O·MnO2, Bader charges of Mna and framework oxygen

Table 2. Voltages (v, vs Li/Li+) for the Electrochemical Insertion of Li Oxides (LixOy) into αMnO2 x

y

composition

V (V)

a/b

0.0625 0.125 0.125 0.25 0.375 0.25 0.5 0.625 0.75 0.875 1

0.03125 0.0625 0.125 0.125 0.1875 0.25 0.25 0.25 0.25 0.25 0.25

0.03125Li2O·MnO2 0.0625Li2O·MnO2 0.125LiO·MnO2 0.125Li2O·MnO2 0.1875Li2O·MnO2 0.25LiO·MnO2 0.25Li2O·MnO2 0.125Li5O2·MnO2 0.25Li3O·MnO2 0.125Li7O2·MnO2 0.25Li4O·MnO2

2.79 2.71 2.73 2.74 2.69 2.79 2.86 2.90 2.95 3.05 3.02

1.004 1.002 1.002 1.004 1.002 1.005 1.009 1.029 1.060 1.055 1.013

calculated voltages for the insertion of Li oxides considered in our study. We note that the voltages for the insertion of either stoichiometric Li2O or Li2O2 is slightly lower than the voltage to form bulk Li2O (2.91 V) or Li2O2 (2.96 V). It suggests that the insertion is a slightly thermodynamically unfavorable route. However, considering all the values lie within ±0.25 V to the discharge voltage of Li2O2, comparable to the typical error in DFT+U predicted the cathode voltages,48 it is very likely that when Li-air battery electrode is discharged/charged, the insertion/removal of Li oxides in αMnO2 also happens with approximately the same voltages. Besides, the voltage increases with the insertion of excess Li. Thus we postulate the possible insertion of Li oxides in αMnO2 happens through the formation of Li2O2 on the surface of αMnO2 followed by the coinsertion of excess Li and Li2O2, which further lowers the energy of the system. To check whether the lattice of αMnO2 is stable to accommodate the insertion of Li oxides, we first check the structure change caused by the insertion of Li2O and compare it with the experimental results. Figure 7 shows lattice constants and volumes for xLi2O·MnO2. Compared with experimental measurements DFT+U slightly overestimates the unit cell volume for xLi2O·MnO2. When x increases from 0.125 to 0.25, DFT calculations predict a continuous increase of the unit-cell volumes, whereas experimentally, the unit-cell volume remains nearly constant. Nonetheless the overall agreements are acceptable. When Li2O is inserted, the structural geometry of 3947

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tolerated in the cycling charge and discharge process.15 From LiMn2O4 to LiMnO2, a/b decreases to approximately 1. The extremely long Mn−O distance in LiMnO2 even indicates possible structure collapse that begins from corner shared MnO6 units (Figure 6). All these results indicate the structure of αMnO2 is not stable enough to accommodate the recycled insertion and removal of Li ions, which eventually leads to an irreversible capacity fading during the cycling of αMnO2. Another possible mechanism that might leads to the structure instability is the displacement of framework ions to the channel, which could either destruct the framework of αMnO2 or block the diffusion of Li ions in the 1D channel.51 To study this effect, we calculate the formation energy the antisite defect that is formed either by the exchange of one Mn ion and a vacant 8h sites or one Li ion. Without Li ions, it costs 5.62 eV to form a VaMn−Mnva defect from the defect free configuration (VaVa−MnMn). With Li ions the energy to form MnLi−LiMn defect is lowered to 4.37 eV. However, this high value still suggests the formation of antisite Mn ions is not favored in αMnO2. For comparison, the formation energy of exchange one Fe and Li ions is only 0.55 eV in LiFePO4.52 Combing this analysis and the results presented above, we conclude that the structural deformation caused the anisotropic Jahn−Teller distortion of Mn3+O6 octahedrals is responsible to the reported irreversible capacity fading during the cycling of αMnO2. The irreversible capacity caused by structure deformation has also been observed in the cycling of spinel LiMn2O4. In fact, when spinel LiMn2O4 is discharged to LiMnO2, a cubic-totetragonal distortion expands the lattice parameter a/c by 16%. This expansion has been regarded as one of the major causes that lead to capacity fading for overdischarged spinel LiMn2O4.14 Numerous studies have been reported to suppress this distortion by decreasing the Jahn−Teller effect with doped cathode materials or with oxygen deficient compounds. Similar efforts may also work in αMnO2 to enhance its electrochemical performance as Li-ion battery cathode. In fact, our results have already shown that if Li is inserted with oxygen, or the insertion of Li oxides, does not damage the structure as much as in the case of Li insertion. We should note here that besides the structural deformation other factors including the dissolution of Mn ions in the electrolyte and the fracture of the particle surfaces at high rate discharge may also contribute to the capacity fading in αMnO2, as they already did in spinel LiMnO2 system.53,54 Nonetheless the intrinsic structure failure should still be treated as one of the main challenges to use αMnO2 in Li-ion batteries. We then discuss the insights from our results to the functions of αMnO2 in Li-air batteries. In nonaqueous Li-air batteries, the electrochemical reaction on the cathode involves the formation and decomposition of Li2O2 during discharge and charge process, respectively, as shown in the following

Figure 8. Electronic density of states shows αMnO2 is an insulator with band gap 0.7 eV, whereas Li oxide inserted into MnO2 is half metallic.

ions are reduced to +2.07e and −1.17e, respectively. Consequently oxygen ions in Li2O unit are partially oxidized. The electrons striped from oxygen in Li2O unit fill the unoccupied orbitals in the conduction band of αMnO2 and lower the band energies. As a result, in LixOy·MnO2 the lowered conduction band overlaps with valence band and the compound becomes half metallic.

4. DISCUSSIONS We first discuss the message from our study to use αMnO2 as rechargeable Li ion battery cathodes. Thackeray and his group first reported αMnO2 as a high capacity cathode material but with poor cyclability.13 Slight changes in unit-cell volume were found during the electrochemical cycling of αMnO2.13 The similar result was later reported in Kijima et al’s studies.10 Thus the volumetric expansion should not be responsible for the irreversible capacity. Our theoretical investigations find a similar trend in the first and second stage of the lithiation, where the unit-cell volume increases by only 7.4%. The small change of the unit cell size is contributed to the tetragonal structure that accommodates the deformation in two dimensions. The expansion of lattice a axis is compensated by the simultaneous decrease of lattice b. However, in the subsequent lithiation after LiMn2O4, lattice a nearly remains as a constant and the crystal deformation can only be accommodated in one dimension. As a result, a larger increase of unit-cell volume is observed. The unit-cell volume of Li2Mn2O4 is 13.4% larger than that of LiMn2O4. An asymmetric change of lattice parameter a and b during lithiation is observed, which is induced by Jahn−Teller distortion of MnO6 units with a particular order of reduction of Mn ions. As a result, a severe tetragonal-to-orthorhombic deformation of the crystal lattice with the maximum a/b ratio reaching 1.19 at composition LiMn2O4 is predicted. This value is comparable to that in the cubic-to-tetragonal deformation of spinel LixMnO2. Such a large deformation is unlikely to be

2Li+ + 2e− + O2 ↔ Li 2O2 (v = 2.96 V vs Li/Li + )

(8)

Here we focus on αMnO2-catalyzed Li2O2 decomposition, which is of great importance in the charge process and is believed to be more catalytically sensitive than Li2O2 formation reaction.21 The direct decomposition of Li2O2 is usually required to overcome a large kinetic barrier that results in the large overpotential of the charge voltage. With the presence of αMnO2 such overpotential is significantly reduced,20 suggesting the role of αMnO2 is to enhance the kinetics of Li2O2 decomposition. It is thus necessary to understand how 3948

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αMnO2 is involved in the electrode reactions and how it enhances the kinetics of Li2O2 decomposition. Our results answers the first question by showing that during the operation of Li-air battery the electrode process is very likely to involve reversible Li oxides extraction or reinsertion in the channel of αMnO2. It is supported by two facts: the intercalation voltage that is very close to the decomposition voltage for either Li2O2 or Li2O, and the structural stability of αMnO2 for the cycled insertion and removal of Li oxides. We note that after this manuscript was submitted, a similar study also supported our conclusion that Li oxides are reversibly inserted and removed in the channel of αMnO2 during the operation of Li-air batteries.55 Like the intercalation of Li, the insertion of Li oxides also involves the charge transfer between the intercalated species and the host. As evidenced by the Bader charge analysis, when Li oxides are inserted, a small but non-negligible amount of electrons is transferred from Li oxides to MnO2 (Figure 8, and Figures S5 and S6 in the Supporting Information). Consequently, αMnO2 here acts as an oxidant to oxidize O in the inserted Li oxides to higher oxidation states. Therefore, with the presence of αMnO2 the direct oxidation of Li2O2 to O2 is split into partially oxidation of O22− followed by the subsequent release of O2. As a result, subsequently releasing oxygen from an oxidation state higher than the stable O22− becomes easier than directly decomposing Li2O2. The above mechanism is illustrated in Figure 9. Instead of directly decomposing into Li and O2, with the presence of

consistent with our analysis. As revealed by the PDOS analysis (Figure 8 and Figure S5 in the Supporting Information), the electrons striped from oxygen ions in Li oxidizes preferably occupies Mn 3d orbitals and reduces Mn4+ to lower oxidation states. Thus our results also prove that Mn in Li oxide inserted αMnO2 stays in a mix-valence environment and, according to Trahey et al’s statement, is beneficial to the O2 evolution kinetics.55 The key step to sustainably support the mechanism illustrated in Figure 9 is the continuous insertion and partially oxidation of Li oxides in αMnO2. Compared to other polymorphs of MnO2, αMnO2 is unique because its large channel is stable enough for the cycled insertion and removal of Li oxides. As shown by Bruce and co-workers, the electrode performance was the best when used with αMnO2 nanowires.19 It suggests that the insertion of Li oxides may only happen at some special surfaces. Our study does not include the investigation of the surface properties of αMnO2. It is thus essential to investigate the incorporation and migration of Li oxides species on different αMnO2 surfaces in future studies. Similar to αMnO2, some of other transition metal oxides can also accommodate the cycled insertion/removal of Li oxides without damaging the host structure. For example, Li2O can be removed from Li5FeO4 or LiMnO3·LiFeO2 either electrochemically or chemically by acid-treatment.56 These materials, after activated by the removal of Li2O, improved the electrode performance of Li-air batteries.56 It is reasonable to hypothesize that the same mechanism governs the decomposition of Li2O2 when catalyzed with these compounds. Another unique property of αMnO2 is the half metallic behavior when Li oxides are inserted in the channel (Figure 8). If the insertion of Li oxides in αMnO2 happens in Li-air batteries, the half metallic product can improve the electron conductivity through the electrode. It was previously reported that the surface of Li2O2 is also half-metallic, despite its bulk is insulating. Our results further support the hypothesis in ref 50 that the electronic conductivity of the discharge product should not significantly affect the charge process of Li-air battery. The large overpotential in the charge process should be an intrinsic requirement for the decomposition of Li2O2. This overpotential could essentially be decreased with electrocatalysts such as αMnO2. Overall our results suggest the great potential of αMnO2 in the application of Li-air batteries.

Figure 9. Schematic of αMnO2 catalyzed decomposition of Li2O2 in the charge of Li-air battery. The decomposition begins with (1) insertion of Li oxides into αMnO2 with electron transferring from Li oxides to MnO2. The release of O2 may happen through (2) decomposition from LixOy·MnO2 to LixMnO2 and (3) decomposition of LixMnO2; or directly through (4) decomposition of LixOy·MnO2 into Li, αMnO2, and O2.

5. CONCLUSIONS To conclude, we employed first-principles calculations to study the insertion of Li and Li oxides in αMnO2. The insertion of Li is identified with a particular order of the reduction of Mn ions. It induces asymmetric change of lattice parameter a and b and lowers the symmetry of the structure to be orthorhombic. The ratio between a/b reaches as high as 1.19 at composition LiMn2O4. It is unlikely for the crystal lattice to tolerate such a large deformation, which is concluded to be responsible for the irreversible capacity fading during cycling. It has been demonstrated that the channel of αMnO2 is stable to accommodate the insertion of Li oxidizes. Li oxides inserted αMnO2 is predicted to be half metallic instead of being insulators. Based on our results a possible mechanism is proposed for the catalytic decomposition of Li2O2 in the charge of Li-air batteries with the presence of αMnO2. The key step in this mechanism is the insertion and partially oxidation of Li oxides into αMnO2. It lowers the energy barrier for the subsequent release of O2. The capability of cycling the insertion

αMnO2, the charge of Li-air battery first inserts and partially oxidizes Li2O2 in the channel of αMnO2 (Step 1). The subsequent release of O2 can be achieved either through the direct decomposition of LixOy·MnO2 (step 4), or through the release of O2 (step 2) followed by the decomposition of LixMnO2 (step 3). Current evidence is not strong enough to support which route is more reasonable. Because a kinetic barrier exists in both routes, the direct comparison of the theoretical predicted voltages for each route and the experimental measurements may not be sound enough either. However, as discussed earlier, the insertion and removal of Li in αMnO2 is accompanied by severe structural deformation that may destroy αMnO2 after a few cycles. It is thus likely that LixOy·MnO2 decomposes directly into Li and O2 at voltages appropriately to support the reverse reaction shown with eq 6. A similar mechanism proposed by Trahey et al suggested that the O2 redox reaction kinetics may be assisted if Mn ions stay in a mix-valence environment.55 In fact, we find this assumption is 3949

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(21) Shao, Y.; Park, S.; Xiao, J.; Zhang, J.; Wang, Y.; Liu, J. ACS Catal. 2012, 2, 844−857. (22) Zhang, R.; Yu, X.; Nam, K.; Ling, C.; Arthur, T. S.; Song, W.; Knapp, A.; Ehrlich, S. N.; Yang, X.; Matsui, M. Electrochem. Commun. 2012, 22, 110−113. (23) Girishkumar, G.; McCloskey, B.; Luntz, A. C.; Swanson, S.; Wilcke, W. J. Phys. Chem. Lett. 2010, 1, 2193−2203. (24) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C. J. Am. Chem. Soc. 2011, 133, 18038−18041. (25) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C. J. Am. Chem. Soc. 2011, 133, 18038−18041. (26) Kresse, G.; Furthmuller, J. Phys. Rev. B 1996, 54, 11169−11186. (27) Kresse, G.; Hafner, J. Phys. Rev. B 1994, 49, 14251−14269. (28) Kresse, G.; Joubert, D. Phys. Rev. B 1999, 59, 1758−1775. (29) Zhou, F.; Cococcioni, M.; Marianetti, C. A.; Morgan, D.; Ceder, G. Phys. Rev. B 2004, 70, 235121. (30) Zhou, F.; Kang, K.; Maxisch, T.; Ceder, G.; Morgan, D. Solid State Commun. 2004, 132, 181−186. (31) Wang, L.; Maxisch, T.; Ceder, G. Phys. Rev. B 2006, 73, 195107. (32) Jain, A.; Hautier, G.; Moore, C. J.; Ong, S. P.; Fischer, C. C.; Mueller, T.; Persson, K. A.; Ceder, G. Comput. Mater. Sci. 2011, 50, 2295−2310. (33) Franchini, C.; Podloucky, R.; Paier, J.; Marsman, M.; Kresse, G. Phys. Rev. B 2007, 75, 195128. (34) Ong, S. P.; Chevrier, V. L.; Hautier, G.; Jain, A.; Moore, C.; Kim, S.; Ma, X.; Ceder, G. Energy Environ. Sci. 2011, 4, 3680−3688. (35) Mueller, T.; Hautier, G.; Jain, A.; Ceder, G. Chem. Mater. 2011, 23, 3854−3862. (36) Michiue, Y.; Watanabe, M. J. Mater. Chem. 2001, 11, 3103− 3109. (37) Dalton, A. S.; Belak, A. A.; Van der Ven, A. Chem. Mater. 2012, 24, 1568−1574. (38) Kim, H.; Kim, D. J.; Seo, D.; Yeom, M. S.; Kang, K.; Kim, D. K.; Jung, Y. Chem. Mater. 2012, 24, 1205−1211. (39) Zhai, D.; Li, B.; Xu, C.; Du, H.; He, Y.; Wei, C.; Kang, F. J. Power Sources 2011, 196, 7860−7867. (40) Zhou, H.; Upreti, S.; Chernova, N. A.; Hautier, G.; Ceder, G.; Whittingham, M. S. Chem. Mater. 2011, 23, 293−300. (41) Eames, C.; Armstrong, A. R.; Bruce, P. G.; Islam, M. S. Chem. Mater. 2012, 24, 2155−2161. (42) Eames, C.; Armstrong, A. R.; Bruce, P. G.; Islam, M. S. Chem. Mater. 2012, 24, 2155−2161. (43) Ven, A. V. d.; Garikipati, K.; Kim, S.; Wagemaker, M. J. Electrochem. Soc. 2009, 156, A949−A957. (44) Komarek, A. C.; Isobe, M.; Hemberger, J.; Meier, D.; Lorenz, T.; Trots, D.; Cervellino, A.; Fernández-Díaz, M. T.; Ueda, Y.; Braden, M. Phys. Rev. Lett. 2011, 107, 027201. (45) Shimizu, Y.; Okai, K.; Itoh, M.; Isobe, M.; Yamaura, J.; Yamauchi, T.; Ueda, Y. Phys. Rev. B 2011, 83, 155111. (46) Bader, R. F. W. Chem. Rev. 1991, 91, 893−928. (47) Ling, C.; Banerjee, D.; Song, W.; Zhang, M.; Matsui, M. J. Mater. Chem. 2012, 22, 13517−13523. (48) Chevrier, V. L.; Ong, S. P.; Armiento, R.; Chan, M. K. Y.; Ceder, G. Phys. Rev. B 2010, 82, 075122. (49) Ishii, Y.; Murakami, J.; Itoh, M. J. Phys. Soc. Jpn. 1999, 68, 696− 697. (50) Radin, M. D.; Rodriguez, J. F.; Tian, F.; Siegel, D. J. J. Am. Chem. Soc. 2012, 134, 1093−1103. (51) Malik, R.; Burch, D.; Bazant, M.; Ceder, G. Nano Lett. 2010, 10, 4123−4127. (52) Dathar, G. K. P.; Sheppard, D.; Stevenson, K. J.; Henkelman, G. Chem. Mater. 2011, 23, 4032−4037. (53) Thackeray, M. M.; Shao-Horn, Y.; Kahaian, A. J.; Kepler, K. D.; Skinner, E.; Vaughney, J. T.; Hackney, S. A. Electrochem. Solid-State Lett. 1998, 1, 7−9. (54) Wang, L. F.; Qu, C. C.; Striebel, K. A.; Chen, J. S. J. Electrochem. Soc. 2003, 150, A905−A911. (55) Trahey, L.; Karan, N. K.; Chan, M. K. Y.; Lu, J.; Ren, Y.; Greeley, J.; Balasubramanian, M.; Burrell, A. K.; Curtiss, L. A.;

of Li oxides and the half metallic properties of Li oxides inserted αMnO2 provides instructive insights to the functions of αMnO2 when used in Li-air batteries.

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ASSOCIATED CONTENT

S Supporting Information *

The preference of interstitial sites for Li and Li oxides insertion, additional information about the evolution of the structure, partial density of state for αMnO2 with and without the insertion of Li2O and schematic about the band structure change, and the decomposed DOS of LiMn2O4. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The discussions with Dr. Masaki Matsui, Dr. Ruigang Zhang, Dr. Tim S. Arthur, Dr. Nikhilendra Singh, and Dr. Yoshinari Makimura are greatly appreciated. Images of crystal structures were generated with the VESTA program.57

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REFERENCES

(1) Brock, S. L.; Duan, N.; Tian, Z. R.; Giraldo, O.; Zhou, H.; Suib, S. L. Chem. Mater. 1998, 10, 2619−2628. (2) Tanaka, Y.; Tsuji, M.; Tamaura, Y. Phys. Chem. Chem. Phys. 2000, 2, 1473−1479. (3) Thanh, D. N.; Singh, M.; Ulbrich, P.; Stepanek, F.; Strnadova, N. Mater. Res. Bull. 2012, 47, 42−50. (4) Cheng, F. Y.; Su, Y.; Liang, J.; Tao, Z. L.; Chen, J. Chem. Mater. 2010, 22, 898−905. (5) Wu, Q.; Tao, H.; Xie, K.; Liu, N.; Yu, L.; Hu, Z. J. Mater. Chem. 2011, 21, 17904−17908. (6) Wang, H.; Lu, Z.; Qian, D.; Li, Y.; Zhang, W. Nanotechnology 2007, 18, 115616. (7) Zhang, H.; Cao, G.; Wang, Z.; Yang, Y.; Shi, Z.; Gu, Z. Nano Lett. 2008, 8, 2664−2668. (8) Johnson, C. J.; Dees, D. W.; Mansuetto, M. F.; Thackeray, M. M.; Vissers, D. R.; Argyriou, D.; Loong, C.-K.; Christensen, L. J. Power Sources 1997, 68, 570−577. (9) Johnson, C. S.; Thackeray, M. M. J. Power Sources 2001, 97−98, 437−442. (10) Kijima, N.; Ikeda, T.; Oikawa, K.; Izumi, F.; Yoshimura, Y. J. Solid State Chem. 2004, 177, 1258−1267. (11) Kijima, N.; Sakata, Y.; Takahashi, Y.; Akimoto, J.; Kumagai, T.; Igarashi, K.; Shimizu, T. Solid State Ionics 2009, 180, 616−620. (12) Kijima, N.; Takahashi, Y.; Akimoto, J.; Awaka, J. J. Solid State Chem. 2005, 178, 2741−2750. (13) Rossouw, M. H.; Liles, D. C.; Thackeray, M. M.; David, W. I. F.; Hull, S. Mater. Res. Bull. 1992, 27, 221−230. (14) Thackeray, M. M. Prog. Solid State Chem. 1997, 25, 1−71. (15) Thackeray, M. M.; David, W. I. F.; Bruce, P. G.; Goodenough, J. B. Mater. Res. Bull. 1983, 18, 461−472. (16) Johnson, C. S. J. Power Sources 2007, 165, 559−565. (17) Komaba, S.; Sasaki, T.; Kumagai, N. Electrochim. Acta 2005, 50, 2297−2305. (18) Zhang, C.; Feng, C.; Zhang, P.; Guo, Z.; Chen, Z.; Li, S.; Liu, H. RSC Adv. 2012, 2, 1643−1649. (19) Débart, A.; Paterson, A. J.; Bao, J.; Bruce, P. G. Angew. Chem., Int. Ed. 2008, 47, 4521−4524. (20) Giordani, V.; Freunberger, S. A.; Bruce, P. G.; Tarascon, J. M.; Larcher, D. Electrochem. Solid-State Lett. 2010, 13, A180−A183. 3950

dx.doi.org/10.1021/cm302347j | Chem. Mater. 2012, 24, 3943−3951

Chemistry of Materials

Article

Thackeray, M. M. Adv. Energy Mater. 2012, DOI: 10.1002/ aenm.201200037. (56) Trahey, L.; Johnson, C. S.; Vaughey, J. T.; Kang, S. H.; Hardwick, L. J.; Freunberger, S. A.; Bruce, P. G.; Thackeray, M. M. Electrochem. Solid-State Lett 2011, 14, A64−A66. (57) Momma, K.; Izumi, F. J. Appl. Crystallogr. 2008, 41, 653−658.

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