Polyselenide Anchoring Using Transition-Metal Disulfides for

Jan 27, 2018 - While selenium has recently been proposed as a lithium battery cathode as a promising alternative to a lithium–sulfur battery, dissol...
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Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

Polyselenide Anchoring Using Transition-Metal Disulfides for Enhanced Lithium−Selenium Batteries Dong Shin Choi,† Min Sun Yeom,‡ Yong-Tae Kim,§ Heejin Kim,*,⊥ and Yousung Jung*,† †

Graduate School of Energy, Environment, Water, and Sustainability, Korea Advanced Institute of Science and Technology, 291 Daehak-ro, Daejeon 34141, Republic of Korea ‡ Center for Computational Science and Engineering, Korea Institute of Science and Technology Information, 245 Daehak-ro, Daejeon 34141, Republic of Korea § School of Mechanical Engineering, Pusan National University, Busan 46241, Republic of Korea ⊥ Electron Microscopy Research Center, Korea Basic Science Institute, 169-148 Gwahak-ro, Daejeon 34133, Republic of Korea S Supporting Information *

ABSTRACT: While selenium has recently been proposed as a lithium battery cathode as a promising alternative to a lithium−sulfur battery, dissolution of intermediate species should be resolved to improve its cycle stability. Here, we report the promising results of transition-metal disulfides as an anchoring material and the underlying origin for preventing active material loss from the electrode using density functional theory calculations. Group 5 and 4 disulfides (VS2, NbS2, TaS2, TiS2, ZrS2, and HfS2) in particular show anchoring capabilities superior to those of group 6 disulfides (CrS2, MoS2, and WS2). The governing interaction controlling the latter relative anchoring strengths is shown to be charge transfer as understood by crystal-field theory. The current findings and methodologies provide novel chemical insight for the further design of inorganic anchoring materials for both lithium−selenium and lithium−sulfur batteries.

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INTRODUCTION In line with the rapidly growing demand for electrical energy, battery technology has also steadily progressed for consumer electronics, electrical vehicles, and energy storage systems. Nevertheless, commercial improvement in the energy density of batteries has been limited because development has mainly been on modification of the electrode materials and manufacturing processes but less on changes in the fundamental mechanism of conventional lithium (Li)-ion intercalation. To make a breakthrough in the energy density, therefore, extensive research has been carried out to develop novel batteries based on new chemistry, such as Li−oxygen (O2), Li−sulfur (S), anion redox, alloying, and conversion.1−3 Among them, the Li−S battery is considered to be one of the promising technologies close to commercialization.4−7 The Li−selenium (Se) battery is analogous to the Li−S battery because of the chemical similarity between S and Se. While the gravimetric capacity of the Li−Se battery is inferior to its S counterpart (678 vs 1672 mAh g−1) because of the atomic mass of Se, its volumetric capacity is comparable (3250 vs 3467 mAh cm−3). A higher electrical conductivity of elemental Se (10−3 S m−1) than of S (10−27 S m−1) can be particularly beneficial to it being used as an electrode material.8,9 Also, the Li−Se battery can employ carbonate-based electrolytes, which are typically incompatible with the Li−S system because of electrolyte decomposition.10 Because of these advantages, the utilization of Se has recently attracted a great deal of interest in either the form of pure Se or a mixture of S and Se, Se1−xSx.8,11,12 © XXXX American Chemical Society

The major bottleneck of Li−Se batteries, among others, is an undesired migration of lithium polyselenides (LiPSe) toward the anode side, a phenomenon well-known as a shuttle effect in Li−S battery systems,13 which can degrade the cycle life and reduce utilization as a cathode material. While the shuttle effect can be suppressed by using carbonate-based electrolytes, a perceptible dissolution of LiPSe is still observed in previous works.14−16 Also, a combined cathode of S and Se (Se1−xSx) would require an ether-based electrolyte (which can lead to the dissolution of polysulfides and polyselenides) because of the incompatibility of carbonate-based electrolytes with polysulfides.10,17 These observations suggest that the prevention of LiPSe dissolution is essential for improving the performances of Li−Se batteries. Indeed, the carbon interlayer between the cathode and separator, which can physically capture the dissolved LiPSe, increased the capacity retention of the Li−Se cell for repeated cycling.18−21 To further prevent the dissolution of LiPSe, nitrogen-doped carbon materials have been introduced for utilizing strong interaction between the polar surface and LiPSe, by which the discharge capacity, polarization, and cycle life were indeed improved.22,23 In fact, such a chemical adsorption through polar surfaces beyond the nitrogen-doped carbon has widely been applied in the Li−S system. For example, oxides, sulfides, and nitrides have been examined, and these inorganic particles markedly improved the Li−S battery performances;24−30 however, Received: November 30, 2017

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DOI: 10.1021/acs.inorgchem.7b03001 Inorg. Chem. XXXX, XXX, XXX−XXX

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Figure 1. (a) Stable solvation geometries of LiPSe with DME molecules (Li2Sen, where n = 8, 6, 4, 2, and 1). Green, orange, red, gray, and white spheres are Li, Se, O, C, and H, respectively. (b) Relative free energy of Li2Sen as a function of the number of coordinate DME molecules. used for integration.38 For slab calculations (anchored states), the slab models were generated from the bulk crystal with a sufficient vacuum layer along the c direction (∼17, ∼17, and ∼25 Å in the a, b, and c directions, respectively). To obtain the adsorption geometry, cell parameters were fixed as the bulk value, while the atomic positions were relaxed until the remaining forces were Ti > V > Zr > Hf > Cr > Mo > W in its decreasing order (Figure 3a), where all TMDs are more effective than the graphene surface for capturing LiPSe. Because these anchoring strengths are neither too strong nor too weak, they could effectively prevent dissolution of LiPSe without degradation of active sites. Overall, the group 5 TMDs (Nb, Ta, and V) have the largest negative anchoring energy, followed by the group 4 TMDs (Ti, Zr, and Hf) and the group 6 TMDs (Cr, Mo, and W). For the high-order LiPSe molecules (n = 8 and 6), all TMDs have almost zero or negative

on the graphene surface, while the short-chain Li2Sen molecules (n = 2 and 1) mostly would precipitate into the solid phase. The anchored state is not preferred in any length of LiPSe because of the weak interaction between LiPSe and the nonpolar graphene surface. This result indicates that, although graphene can aid electronic conduction, it is not able to prevent dissolution of LiPSe. Also, an enlarged NPE value of the anchored state with decreasing chain length (n) suggests that the low-order LiPSe are less probable to contact with the carbon electrode, and thus the reduction kinetics can be deteriorated for reactions involving such low-order LiPSe. Therefore, defective sites, functional groups of oxygen, or heteroatom dopants such as nitrogen are necessary in carbon materials for preventing dissolution of LiPSe, as demonstrated in previous experimental studies.49,50 We note that the stabilities of the anchored states of LiPSe in end-on and side-on modes (filled and empty triangles in Figure 2b, respectively) are almost identical because they rely mainly on the dispersion interactions. Also, we assumed that the energy values of intermediate Li2Sen (n = 3, 5, and 7) are in a trend line. As a comparative example, we first evaluated the energy states of LiPSe on the TiS2 surface (Figure 2c) because TiS2 is one of the most effective anchoring materials in analogous Li−S systems because of its polar surface.29 In contrast to the graphene surface, the soluble species (n = 8−2) prefer the anchored state by ∼1 eV on the TiS2 surface. Here, the interaction between LiPSe and the TiS2 surface is composed of Li−Ssurf electronic and Se−Ssurf dispersion interactions, where Ssurf indicates the surficial S of TiS2. Thus, the high-order Li2Sen molecules (n = 8 and 6) bind in the end-on mode (filled triangles in Figure 2c) to maximize the Se−Ssurf dispersion interaction, whereas the short-chain Li2Sen molecules (n = 2 and 1) are more stable in

Figure 3. (a) NPE of the anchoring state of LiPSe (anchoring energy) on graphene (Gr) and MS2 surfaces, where M = Ti, V, Cr, Zr, Nb, Mo, Hf, Ta, and W. (b) Anchoring energy as a function of the electronegativity difference between M and S of the MS2 substrate. (c) Correlation between the anchoring energy and the extent of charge transferred from LiPSe to the substrate. (d) Electron density isosurface (red volume) for gained charge after LiPSe adsorption. D

DOI: 10.1021/acs.inorgchem.7b03001 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry anchoring energy values, indicating that dissolution of high-order LiPSe can be prevented effectively by TMDs. On the other hand, the low-order LiPSe molecules have an opposite trend in the anchoring energy for group 5 and 6 TMDs: with decreasing chain length (n), the anchoring energy becomes gradually negative (favorable) for the group 5 TMDs, while it increases positively (unfavorable) for the group 6 TMDs. This difference between the transition-metal (TM) groups will be discussed in more detail below. In previous studies, TMDs have been considered to be efficient anchoring materials because of their surface polarity.29 The polarity of the chemical bonds can be roughly estimated from the electronegativity difference (ΔEN) between two elements, each TM and S in TMDs. When we draw an anchoring energy as a function of ΔEN (Figure 3b), the weak anchoring of MoS2 and WS2 can be explained by the small ΔEN; however, the anchoring energies of other TMDs are not described well with the ΔEN values alone. This result in Figure 3b implies that other factors beyond the surface polarity play a role in the anchoring phenomena of TMDs. To elucidate the interaction between LiPSe and TMDs, we performed charge analysis using the Bader method. The x axis of Figure 3c is the variation of the charge density in LiPSe molecules after adsorption on each TMD, where the negative value indicates that the electrons have transferred from LiPSe molecules to the TMD substrates. In a detailed picture of Figures 3d and S4, most of the transferred electrons are located between the TM and S of TMDs rather than between LiPSe and TMDs. This indicates that the electrons are almost completely transferred from LiPSe to TMD and their interaction is electrostatic, as in a typical ionic bond. Indeed, as shown in Figure 3c, as the extent of electron transfer from LiPSe to TMDs increases, the anchoring energy between them becomes more significant. Such a linear correlation between charge transfer and the binding energy can also be seen in a previous study on the analogous Li−S system (Figure S5),47 even when different types of materials such as oxides, sulfides, and chlorides were compared as anchoring substrates. We expect that the correlation between charge transfer and the anchoring energy might hold in a different level of theory calculations because it is a comparative result. Also, we note that the charge density of an explicit DME molecule remains the same as that of a free molecule, indicating that solvent molecules do not contribute to the net charge transfer between LiPSe and TMDs. Upon formation of the Li−Ssurf bond, the electrons involved are polarized toward the TMD side, in spite of the fact that Se in LiPSe also has a high electronegativity (2.55) compared to the Ssurf (2.58) or TM elements (1.3−2.36) of TMDs. This implies that the electrons are further stabilized when placed in the crystal of TMDs and thus get transferred. To estimate this effect of the crystal structure on the stability of transferred electrons, two different structures of TMDs were considered, as shown in Figure 4: one has an octahedral coordination of S ligands, as appears in the 1T structure, and the other has a trigonal-prismatic coordination of S ligands, which is observed in the 2H or 3R structure. According to the crystal structure database, group 4 TMDs are stable under the 1T structure, while group 5 and 6 TMDs (except VS2) prefer the 2H structure.51 For these experimentally resolved structures and for their respective hypothetical counterparts, we evaluated the binding energy of a single Li atom on TMD surfaces (Figure 4a) to isolate the electronic effects without the influence of dispersion interaction. As shown in Figure S6, this simplified Li-atom binding model shows a

Figure 4. (a) Binding energy of the Li atom on sulfides with octahedral and trigonal-prismatic structures. Both experimentally discovered and hypothetical phases are presented respectively by filled and empty marks. (b) Crystal-field splitting of the octahedral and trigonalprismatic coordinations. The electron configuration of the bare state (black arrow) and additional electrons due to Li binding (red arrow) are presented.

reasonable correlation with the anchoring energy of LiPSe, validating this Li binding model as well as supporting the theory that Li−Ssurf interaction is a key factor for interpreting the difference between TMDs. While the binding energy of Li overall decreases with increasing atomic number, two aspects are noticeable in Figure 4a. First, the trigonal-prismatic structure has stronger Li binding energy compared to the octahedral structure for group 4 and 5 TM elements, which have d0 and d1 electron configuration, respectively. Second, group 6 TM elements with d2 electron configuration in a trigonal-prismatic structure have exceptionally weak Li binding energy. These points consistently appear in the anchoring energy of LiPSe (Figure 3a) and can be understood by crystal-field theory. As illustrated in Figure 4b, d orbitals of the metal center split into two degenerate levels in the octahedral ligand field (wellknown antibonding eg and nonbonding t2g), whereas they split into three levels under the trigonal-prismatic environment (antibonding dxz and dyz on the top, intermediate dxy and dx2−y2 on the middle, and nonbonding dz2 on the bottom).52,53 The d0, d1, and d2 electron configurations of bare TMDs (without anchored Li) E

DOI: 10.1021/acs.inorgchem.7b03001 Inorg. Chem. XXXX, XXX, XXX−XXX

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together (e.g., 790, 618, and 5 S cm−1 for TiS2, MoS2, and VS2, respectively), the TMDs can act as a channel for supplying both electrons and Li ions, in addition to the anchoring substrate.

are drawn with black arrows in Figure 4b. Because dz2 in the trigonal-prismatic field lies lower than t2g of the octahedral field,52 the d1 and d2 electron configurations are stable under the trigonalprismatic coordination (except the V case, where the crystalfield splitting is less prominent than the d4 or d5 elements). When the Li atom (or LiPSe) is anchored on the surface of TMDs with d0 or d1 elements, transferred electrons (denoted with red arrows in Figure 4b) occupy t2g and dz2 orbitals in the octahedral and trigonal-prismatic structures, respectively. This leads to the stronger Li binding on the trigonal-prismatic structure than on the octahedral structure for d0 and d1 elements, as shown in Figure 4a, because of the lower-lying dz2 orbital. By contrast, in the group 6 TMDs with d2 elements, the stable dz2 orbital of the trigonal-prismatic structure is fully occupied in the bare state, and thus the additional transferred electrons from the anchored Li are allowed to occupy the next stable dxy (dx2−y2) orbital that is quite a bit more unstable than the t2g orbital of the octahedral structure or the dz2 orbital of the trigonal-prismatic structure. The latter indicates that electron transfer from Li (or LiPSe) to the group 6 TMDs is energetically unfavorable, leading to a small amount of electron transfer, as shown in Figure 3c, and, consequently, resulting in the weak anchoring energy for the group 6 TMDs, as shown in Figure 3a. In the electron microscopy experiment, a spontaneous transformation of MoS2 from trigonal-prismatic to octahedral structure was observed after Li adsorption, supporting our explanation on the structure-dependent capability of Li binding.54 The stronger anchoring energy of the group 5 TMDs is also attributed to this structural reason: they have a trigonal-prismatic structure in the bare state with an empty dz2 orbital, which can optimally stabilize the additional electrons from LiPSe. Because the present interpretation is based on the interaction between Li and Ssurf, we expect that the discussions on TMDs described above hold for the existing Li−S battery systems as well. Indeed, TMDs with strong anchoring energy (NbS2, TiS2, VS2, and ZrS2) have been demonstrated as effective additive materials for the Li−S battery.29,55 In the case of weak anchoring energy materials (MoS2 and WS2), however, the sulfur deficiency,56 edge sites,57−59 additional inorganic materials,60 and carbon interlayers61 were necessary to acquire notable improvements in the discharge capacity of the Li−S battery. These results demonstrate that the surface of group 6 TMDs by itself is insufficient to capture LiPS, as calculated in Figure 3a. Moreover, the results in Figure 4 suggest that one can improve the anchoring capability of TMDs by controlling the crystal structure as trigonal-prismatic or octahedral. Last, we note that strong anchoring TMDs such as NbS2 and TiS2 bind the Li atom by 1.8−2 eV, as shown in Figure 4a, when the coverage of Li is 0.03 monolayer (ML). Because this energy value overlaps with the discharge potential range of LiPSe (∼2 V), the electrochemical lithiation of the TMD surface (xLi + MS2 → LixMS2) could happen competitively. These surficial binding energies of Li, however, decrease constantly to 1.13 and 0.98 eV for NbS2 and TiS2, respectively, with increasing coverage to 1 ML. The latter suggests that only a part of the TMD surfaces would be lithiated during the discharge process and the rest of the TMD surfaces are still active for the anchoring process. Also, we expect that such surficial Li atoms can be utilized for the discharge process of LiPSe if they can be migrated. The surficial migration barriers of Li were evaluated to be ∼0.2 eV for all TMDs, as shown in Figure S7, suggesting that the bound Li can readily be supplied to the discharge reaction of LiPSe. When the high electrical conductivities of TMDs are considered

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CONCLUSIONS In summary, we examined nine different TMDs as additive materials for the Li−Se batteries to prevent dissolution of LiPSe. We found that the anchoring capabilities are NbS2 > TaS2 > TiS2 > VS2 > ZrS2 > HfS2 > CrS2 > MoS2 > WS2 in its decreasing order, and all of the TMDs considered here can capture LiPSe more efficiently than the graphene surface. This trend in the anchoring energy was revealed to be attributed to the extent of charge transfer from LiPSe to TMDs during the anchoring process. When the anchoring materials are grouped together, a strong anchoring capability of the group 5 TMDs (NbS2, TaS2, and VS2) and a contrarily weak anchoring capability of the group 6 TMDs (CrS2, MoS2, and WS2) are significant. This difference between groups of TMDs in the anchoring energy is satisfactorily explained in the framework of crystal-field theory: The most stable dz2 orbital of the trigonal-prismatic structure can accommodate the transferred electrons in the group 5 TMDs, which promotes electron transfer from LiPSe to TMDs and thus leads to strong electrostatic binding. The present findings and understandings can also explain the existing Li−S chemistry with TMD additives, providing novel insight for further development of the additive materials for both Li−Se and Li−S batteries.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.7b03001. Charge analysis for Li2Sen, electron gain isosurfaces, binding energy versus charge transfer for Li2S4, and anchoring energy as a function of the lithium binding energy (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. ORCID

Yong-Tae Kim: 0000-0001-9232-6558 Heejin Kim: 0000-0003-3027-6983 Yousung Jung: 0000-0003-2615-8394 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We acknowledge financial support from the Korea Institute of Science and Technology Information (Grant C17004), the Samsung Future Technology Foundation (Grant SRFC-TA 1403-04), and the R&D Convergence Program of the National Research Council of Science & Technology (Grant CAP-15-02KBSI).

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DOI: 10.1021/acs.inorgchem.7b03001 Inorg. Chem. XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.inorgchem.7b03001 Inorg. Chem. XXXX, XXX, XXX−XXX