Superhalogens: A Bridge between Complex Metal Hydrides and Li Ion

Mar 12, 2015 - Joseph A. Teprovich , Héctor Colón-Mercado , Aaron L. Washington II , Patrick A. Ward , Scott Greenway , David M. Missimer , Hope Har...
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Superhalogens: A Bridge between Complex Metal Hydrides and Li Ion Batteries Puru Jena Physics Department, Virginia Commonwealth University, Richmond, Virginia 23284-2000, United States ABSTRACT: Complex metal hydrides and Li ion batteries play an integral role in the pursuit of clean and sustainable energy. The former stores hydrogen and can provide a clean energy solution for the transportation industry, while the latter can store energy harnessed from the sun and the wind. However, considerable materials challenges remain in both cases, and research for finding solutions has traditionally followed parallel paths. In this Perspective, I show that there is a common link between these two seemingly disparate fields that can be unveiled by studying the electronic structure of the anions in complex metal hydrides and in electrolytes of Li ion batteries; they are both superhalogens. I demonstrate that considerable progress made in our understanding of superhalogens in the past decade can provide solutions to some of the materials challenges in both of these areas.

W

e live in a world of limits. Of all of the things that are limited, nothing impacts our society, economy, and lifestyle more than fossil fuels do. Not only are these fuels limited, but also they have an adverse effect on the environment and are the primary cause for global warming. With the growing world population and increasing demand on energy in the developing countries, it is not a question of if but when the world will face an energy crisis. According to the U.S. Energy Information Administration,1 world energy consumption will increase from 524 quadrillion Btu in 2010 to 630 quadrillion Btu in 2020 and 820 quadrillion Btu in 2040. While energy conservation and efficiency can address some of the issues, to maintain the current economic growth and standard of living without affecting the climate, it is imperative that clean, abundant, cost-effective, and safe alternative energy sources be developed. Meeting the energy demand will not only require multiple renewable energy sources such as solar, wind, hydro, biomass, and hydrogen but also advanced technologies to store this energy. In each category, numerous materials challenges remain and to address them will require out-of-the-box ideas.

In this Perspective, I discuss one such idea where the knowledge gained in one field can be applied in another to accelerate materials discovery. For example, consider superhalogens, complex hydrides, and Li ion batteries; at the outset, they do not appear to have anything in common. Superhalogens2 are a class of molecules that mimic the chemistry of halogen atoms but have electron affinities (EAs) that far exceed that of any halogen and are used as oxidizing agents. Complex metal hydrides such as alanates [M(AlH4)x] and borohydrides [M(BH4)x] (M is a metal atom with valence x) contain a large amount of hydrogen and are considered as promising materials for hydrogen storage. Li and Na ion batteries, due to their light weight and high energy densities, are widely used for electrochemical storage of energy. While considerable research has been done both in hydrogen storage materials3,4 and Li ion batteries,5 many material challenges remain that need to be addressed for a successful hydrogen economy as well as for safer and more efficient electrochemical storage. In the following, I discuss these challenges and show how superhalogens that form a bridge between these two apparently disparate fields can address some of these challenges. First is the need and challenges in a hydrogen economy. Because about 75% of the oil used in the United States goes to meet the need of the transportation industry, it is necessary that an alternate clean energy solution be found to meet the energy needs of this industry. Hydrogen is considered to be a good candidate because it is the third most abundant element on the Earth and is renewable and clean as water is the only byproduct when hydrogen burns. For these reasons, President Bush announced a national hydrogen initiative in 2001 with the

Meeting the energy demand will not only require multiple renewable energy sources such as solar, wind, hydro, biomass, and hydrogen but also advanced technologies to store this energy. In each category, numerous materials challenges remain and to address them will require out-ofthe-box ideas. © XXXX American Chemical Society

Received: January 2, 2015 Accepted: March 12, 2015

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expectation that “the first car driven by a child born today could be powered by hydrogen and pollution-free”. Unfortunately, hydrogen does not occur freely in nature and has to be produced from organic compounds or water, and because it is lighter than air, it needs to be stored. A subsequent report6 by the U.S. Department of Energy outlined the basic research needs for a successful hydrogen economy by focusing on materials challenges in hydrogen production, storage, and use. Among the many challenges, reversible storage of hydrogen with sufficient gravimetric and volumetric densities at ambient thermodynamic conditions was considered to be the biggest hurdle. This initiative has led to a worldwide effort for developing solid-state hydrogen storage materials as the current methods for storing hydrogen under high pressure or in liquid form have inherent drawbacks and are not suitable for widescale commercial applications. The promising solid-state materials for hydrogen storage are complex light metal hydrides that are composed of a metal cation and a complex anion. Typical examples are alanates [M(AlH4)x] and borohydrides [M(BH4)x], where M is a metal atom with valence x. Although the gravimetric and volumetric densities of hydrogen in these materials are high, hydrogen is bonded covalently. Consequently, these materials suffer from slow dehydriding or rehydriding kinetics.3 In addition, during dehydrogenation, some of these materials exhibit intermediate phases that affect their reversibility. Safety can also be an issue; Al(BH4)3, for example, is highly volatile and pyrophoric. While considerable progress has been made in our understanding of the nature of hydrogen bonding in these materials, much more needs to be done before identifying the ideal material for commercial application. In the following, I discuss how the concept of superhalogens fits into this scheme and how it also provides a bridge between hydrogen storage materials and electrolytes in Li ion batteries. In Figure 1, I show, as an example, the low- and hightemperature crystal structure7 of LiBH4. It is composed of a Li+

defined as the energy difference between the ground states of the negative ion and its neutral, while the latter is defined as the energy difference between the ground state of the anion, and its neutral at the anion geometry. The EA is always smaller than the VDE, and in most cases, the difference between the two is less than 0.2 eV. However, in rare cases where the neutral and anion ground-state geometries are very different, this difference can be large, as is the case with BH4. The EA and VDE of BH4 are, respectively, 3.42 and 4.42 eV. These values are far greater than the EA of the H atom, which is only 0.794 eV. In the early 1980s, Gutsev and Boldyrev2 had shown that the EA of molecules composed of a metal atom at the center surrounded by halogen atoms is larger than that of the halogen atom, as long as their number exceeds the valence of the metal atom by one. They named such clusters as “superhalogens”. Although H is not a halogen and the EA of BH4 is not larger than that of the halogen atom Cl, one can regard this moiety as a “superhalogen” because its EA is much larger than that of its building block, H, and its VDE is larger than the EA of Cl. Numerous studies of superhalogens have been carried out over the last 3 decades, and their existence has been confirmed by experiments.8−11 In analogy with the normal salts, which are composed of metal cations and halogen anions, salts composed of metal cations and superhalogen anions can be termed as “supersalts”.12 In this sense, all complex metal hydrides are “supersalts”. In addition to superhalogens describing the electronic structure of anions in complex metal hydrides, I will show in the following that they also play a role in understanding the intermediate phases during the decomposition of metal borohydrides as well as guide a path to make them safer. Orimo and co-workers13 were the first to discover that the formation of an intermediate phase, Li2[B12H12], during the dehydrogenation of LiBH4 is one of the main reasons that hinders its reversibility. Later experiments14−24 have confirmed the presence of Na2B12H12, MgB12H12, and CaB12H12 during the decomposition of NaBH4, Mg(BH4)2, and Ca(BH4)2, respectively. Theoretical studies25 have also shown that while Li2B12H12 and MgB12H12 are respectively the main intermediates in the decomposition of the LiBH4 and Mg(BH4)2, the decomposition of Ca(BH4)2 leads to two energetically nearly degenerate reaction pathways, yielding either CaB 6 or CaB12H12. Experimental studies of magnesium and yttrium borohydrides subsequently showed26,27 significant concentration of the octahydrotriborates Mg(B3H8)2 and Y(B3H8)3 during decomposition of Mg(BH4)2 and Y(BH4)3. Recently, it has also been found by Zuttel and co-workers that the decomposition of Ca(BH4)2 depends upon temperature (unpublished results). It decomposes into CaB6 at 350 °C and into elemental boron at temperatures above 400 °C. A crucial intermediate phase, CaB2H6, forms at 350 °C, circumventing the formation of boron sinks such as CaB12H12. Consequently, Ca(BH4)2 becomes reversible without any additives. Why do different borohydrides exhibit different intermediate phases? Theoretical studies carried out using both crystals25 and cluster models28,29 have found an answer; B3H8 and B12H12 are also superhalogens like BH4, and the preferred intermediate phases depend upon the binding energies of the metal cations with these complex anions. However, unlike the original superhalogen concept where the core atom is a metal and the surrounding atoms are halogens, B3H8 and B12H12 are superhalogens for different reasons; they do not obey the usual octet electron counting rule that makes BH4 a superhalogen.

Figure 1. Crystal structures of the (a) low-temperature phase and (b) high-temperature phase of LiBH4 [reprinted with permission from ref 46, copyright 2014].

cation and a perfectly tetrahedral BH4− complex anion. Note that because B is trivalent, to accommodate the fourth hydrogen, an extra electron is needed. This extra electron is provided by the metal atom through charge transfer, and therefore, complex metal hydrides can be considered as salts because of the ionic interaction between the metal cation and the complex anion. The enhanced stability of the complex anion is due to the fact that the extra electron is distributed over all four H atoms, thus reducing electron−electron repulsion. The stability of an anion is measured by the EA and the vertical detachment energy (VDE). The former is 1120

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Table 1. Binding energies, ΔELi+ of Li+ Ions to Complex Boron−Hydrogen Anions [ΔELi+ = E(Li+ ) + E(complex anion) − E(salt)]a

Instead, B3H8 and B12H12 are superhalogens because their anions obey the aromaticity30,31 and Wade−Mingos rules,32,33 respectively. To understand this behavior, we first compare the geometries of BH4−, B3H8−, and B12H122− in Figure 2. BH4−, as

salt

complex anions

binding energies (eV)

LiBH4 LiB3H8 Li2B12H12

BH4− B3H8− B12H122−

6.62 6.25 7.47

a Note that for Li2B12H12, energies needed to remove the first and second Li+ ions are 5.94 and 9.00 eV, respectively. In this table, the average removal energy per Li+ is listed.

electron is added, just as a halogen needs an extra electron to be stable. Because BH4 mimics the chemistry of a halogen, it can be easily seen that Al(BH4)4 would also mimic a halogen and can be stabilized only if an extra electron is added. Furthermore, the EA of Al(BH4)4, namely, 5.56 eV, is higher than that of BH4 superhalogen. One can, thus, regard Al(BH4)4 as a hyperhalogen38,39 because it is composed of BH4 superhalogens as building blocks. When countered with a metal cation such as K+, Al(BH4)4− can then form a salt, KAl(BH4)4, which can be termed as a hypersalt. The question is whether KAl(BH4)4 would be as volatile and pyrophoric as Al(BH4)3 or if will it be a safer material. The answer has come from a recent experiment40 where KAl(BH4)4 was experimentally synthesized by mixing KBH4 with Al(BH4)3. These compounds are compared in Figure 3. Note that KAl(BH4)4 is

Figure 2. Equilibrium geometries of BH4−, B3H8−, and B12H12− [reprinted from ref 29].

discussed earlier, is a perfect tetrahedron, and its structure is nearly identical to that in crystalline metal borohydrides (see Figure 1). The geometry of B3H829,34−36 is characterized by a pair of H atoms radially bonded to each of the three B atoms and two bridge-bonded H atoms. The geometry of B12H122− is a perfect icosahedron37 where 12 B atoms occupy the vertices of the icosahedron and all 12 H atoms are radially bonded to the B atoms.

Complex metal hydrides and Li ion batteries play an integral role in the pursuit of clean and sustainable energy. For a molecule to be aromatic, it must have 4n + 2 (n = 0, 1, 2, ...) electrons. A classic example of an aromatic molecule is C6H6, which has six π electrons; thus, n = 1. To see why B3H8− is aromatic,36 we note that neutral B3H8 contains a total of 17 valence electrons, nine of which are contributed by three B atoms and eight are contributed by eight H atoms. Because in B3H8 six H atoms are bonded radially to the three B atoms while the other two are bridge-bonded, 12 of the 17 electrons are occupied in covalent B−H bonds. This leaves B3H8 with five electrons. According to the aromaticity rule, an extra electron is needed to stabilize B3H8, thus making B3H8 a superhalogen. For the closo-boranes such as B12H122−, the Wade−Mingos rule requires (n + 1) pair of electrons, where n is the number of vertices in the boron polyhedron. In B12H12, there are a total of 48 valence electrons, of which 24 electrons are involved in 12 covalent B−H bonds, thus leaving 24 electrons for icosahedral cage bonding. Thus, two extra electrons are needed to satisfy the Wade−Mingos rule [(12 + 1) × 2 = 26], rendering B12H122− its unusual stability. Indeed, the first EA of B12H12 is 4.61 eV, which gains an additional 0.9 eV of energy when the second electron is attached,28,37 and B12H122− is the most stable gas-phase dianion known in chemistry. It is this enhanced stability of the superhalogen anions that gives rise to the large binding energy with a metal cation, which in turn is responsible for the stability of the intermediate phase compounds. In Table 1, these binding energies are listed. Next, I illustrate how the superhalogen concept can be used to convert highly volatile and pyrophoric Al(BH4)3 into a safer salt. As pointed out before, BH4 is stable only when an extra

Figure 3. Al(BH4)3 in its natural state (left panel) and as a stabilized solid in the form of KAl(BH4)4 (right panel) [reprinted from ref 40].

a solid at ambient conditions and is neither volatile nor pyrophoric. Thus, changing the composition of the anion can indeed transform an unsafe ionic liquid into a safer solid! It has also been shown29 that the energy necessary to remove a H atom from KAl(BH4)4 is smaller than that from Al(BH4)3. This provides an additional advantage for using hypersalts over supersalts for hydrogen storage. In the following, I show how the superhalogen concept can also be used to design and synthesize electrolytes for a new generation of safer Li ion batteries. The three major components of these batteries are the anode, cathode, and electrolyte. The anode is usually made of graphite, and the cathode is composed of metal oxides. The electrolytes, composed of Li salts, provide the Li+ ions, which move from the cathode to the anode during charging and reverse when discharging. The most commonly used commercial electrolytes in Li ion batteries are LiAsF6, LiBF4, LiPF6, LiFePO4, LiClO4, LiN(SO2F)2, and LiN(SO2CF3)2. The electronic structure of the anionic components of these salts has been recently studied 1121

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theoretically.41 They were found to possess a common property; they are all superhalogens. This can be easily understood by counting the number of valence electrons of AsF6, BF4, PF6, FePO4, and ClO4 moieties. Note that the maximal valencies of As, B, and P are 5, 3, and 5, respectively. Fe, on the other hand, can exist in oxidation states of +2 or +3. With a maximal valencey of Cl being 7, ClO4 needs an extra electron to be stable and hence serves as a negative ion. Similarly, with Fe and P having valencies of 2 and 5, respectively, FePO4 also serves as a negative ion. The calculated geometries of these negative ions are shown in Figure 4. The

electron to satisfy its electronic shell closure. With the charge state of SO2CF3 being +1, N has to be in the −3 charge state so that N(SO2CF3)2 can be stabilized as an anion. The geometries of N(SO2F)2− and N(SO2CF3)2− are also shown in Figure 4, and their VDEs are given in Table 2. Again note that all of these moieties are suprhalogens. The question is, Why are the electrolytes composed of complex anions instead of simple halogens? The answer lies in the binding of the Li+ ions in these complex salts as well as the size of the anions. We define these binding energies as ΔE Li+ = (Eanion + E Li+) − ESalt

(1)

The results are given in Table 2. Due to the large size of the anions, the binding energy should be small, but it is compensated for by the large VDE of the complex anions. Note that the binding energy of Li+ in all of the electrolytes, with the exception of LiFePO4, is either comparable or smaller than that of LiF, which is 5.98 eV. The large size of the anions, however, allows Li+ ions greater mobility. The disadvantage of the currently used electrolytes is that they contain halogens that are toxic and not environmentally friendly. Consequently, there is a strong need to develop halogen-free electrolytes. Realizing that all of the negative ions of current electrolytes are superhalogens and that superhalogens can be created42,43 without the benefit of halogens, one naturally wonders if some of these halogen-free superhalogens can form the building blocks of a new generation of halogen-free electrolytes in Li ion batteries. Important properties that the electrolytes must satisfy are the low binding energy and greater mobility of the Li+ ion and the insensitivity of the electrolytes to water. The later quantity can be studied by calculating the binding energy of water to the electrolytes, namely ΔE H2O = (ESalt + E H2O) − ESalt + H2O

(2)

41

Recently, Giri et al. have studied this problem. They calculated the binding energy of the Li+ ion, ΔELi+ and that of water, ΔEH2O in a number of halogen-free superhalogens such as NO3, BH4, B3H8, and CB11H12. In Table 3, I list these

Figure 4. Optimized geometries of different anions in currently used electrolytes in Li ion batteries [reprinted from ref 41, copyright 2014].

stability of these moieties is determined by how much energy is gained when an electron is attached. This can be seen by calculating either the EA or the VDE. In Table 2, we give the

Table 3. Calculated VDE (eV), Li+ Binding Energy (ΔELi+, eV), and Binding Energy of H2O (ΔEH2O, eV) in Four Potential New Electrolytes for the Li Ion Battery

Table 2. Calculated VDE (eV), Li+ Binding Energy (ΔELi+, eV), and Binding Energy of H2O (ΔEH2O, eV) in Currently Used Electrolytes in a Li Ion Battery anions

VDE

ΔELi+

ΔEH2O

FePO4− ClO4− N(SO2F)2− N(SO2CF3)2− BF4− PF6− AsF6−

4.32 5.83 6.89 7.01 7.66 8.55 8.91

7.38 5.96 5.82 6.01 6.08 5.73 5.65

1.04 1.02 1.02 0.99 1.41 1.07 1.09

anions

VDE

ΔELi+

ΔEH2O

NO3− BH4− B3H8− CB11H12−

4.22 4.50 4.72 5.99

6.53 6.62 6.25 5.08

0.96 0.92 0.93 1.08

energies in LiNO3, LiBH4, LiB3H8, and LiCB11H12 salts and compare them with those of halogen-containing electrolytes given in Table 2. Note that these values are comparable, and in the case of LiCB11H12, the Li+ binding energy is significantly smaller than those in the current electrolytes. Of all of the halogen-free electrolytes, the most promising one, therefore, is LiCB11H12. To increase the Li content of the electrolyte, Giri et al.41 also studied the potential of Li2B12H12 as a halogen-free electrolyte. Here, the advantage is that there are two Li ions per each B12H122− moiety. To study its potential, they calculated the energy needed to remove the first Li+ ion to be 5.94 eV and the second to be 9.00 eV. This is to be compared with 6.62 eV

VDE values of the above complex anions. Note that they are all greater than the EA of Cl, namely, 3.6 eV. The valence electron counting in N(SO2F)2 and N(SO2CF3)2 is more subtle as the oxidation state of N can vary from −3 to +5 and that of S can vary from −2 to +6. In N(SO2F)2, the oxidation states of N and S are, respectively, +3 and +4. With the oxidation state of O and F at −2 and −1, respectively, N(SO2F)2 needs an extra 1122

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needed to remove the Li+ ion from LiBH4. Thus, Li2B12H12 may not offer significant advantage over LiBH4 for use in Li ion batteries. Use of LiBH4 as a solid-state electrolyte in Li ion batteries has already been studied by the Orimo group.44 The authors demonstrated that LiBH4 has the potential to be utilized as a fast lithium ionic conductor. LiBH4 undergoes a structural phase change7 (occurring at ∼390 K) from orthorhombic to hexagonal upon heating (see Figure 1). The high-temperature phase (hexagonal) has a conductivity of 10−3 S cm−1, while the low-temperature phase (orthorhombic) has between 10−8 and 10−6 S cm−1. The Li ion conductivity in various metal hydrides has also been recently reviewed by Orimo and co-workers.45,46 The mobility of the Li ion in the solid state can be further enhanced by the introduction of inorganic salts, anions, and metal dopants. Stable cycling in an all-solid-state lithium ion battery utilizing pure LiBH4 as a solid-state electrolyte with LiCoO2 and Li metal as the cathode and anode, respectively, has recently been demonstrated.47 These works have clearly shown that metal hydrides originally thought of as hydrogen storage materials can also be used as components of Li ion batteries. The above analogy has prompted some of the ideas developed in hydrogen storage materials to be applied directly to Li ion batteries. One such example is the use of C60 as a dopant to facilitate hydrogen release. It was shown48 earlier that C60 serves as a catalyst and is as effective as TiCl3 in the dehydrogenation of NaAlH4. This was possible because C60, which is electronegative with an EA of 2.6 eV, helps to weaken the ionic bond between Na+ and AlH4−. Teprovich et al.49 have recently explored if the addition of C60 to LiBH4 would similarly enhance the conductivity of the Li+ ion in LiBH4 electrolyte. This was indeed found to be the case, and the result is similar to the effect observed when LiBH4 is ball milled with lithium halides.44 In Figure 5, the result of Teprovich et al.49 on the ionic and electronic conductivity of the LiBH4/C60 nanocomposite utilizing a symmetrical cell setup Li/SSE/Li (SSE = LiBH4-based solid-state electrolyte) is reproduced. The activation energy for lithium ion conduction in the LiBH4/C60 nanocomposite calculated from these data was found to be

lower than the activation energy of lithium ion conduction in pure LiBH4. The demonstration that the LiBH4/C60 nanocomposite as an electrolyte is compatible with a pure lithium metal electrode during extended cycling is a significant departure from the commonly utilized inorganic additives45 currently used to enhance Li ion mobility in LiBH4-based materials. This methodology could potentially be utilized to enhance alkali ion mobility in other complex hydrides.

Superhalogens serve as a bridge between complex metal hydrides and electrolytes in Li ion batteries, and the knowledge gained in superhalogen research can be useful in designing safer complex metal hydrides and electrolytes. In summary, I have tried to show that knowledge gained from research in one field can be helpful in making progress in another and that there often is a common link whose fundamental understanding may guide focused discovery of materials. The examples that I have chosen are complex metal hydrides and electrolytes in Li ion batteries, the common link between them being the superhalogen building blocks. Complex light metal hydrides are salts composed of complex anions and metal cations. The electrolytes in Li ion batteries are also salts with similar composition, namely, Li+ cation and complex anions. Complex anions in both cases are superhalogens. Recent work has shown that LiBH4, which, with 18 wt % hydrogen, is attractive as a hydrogen storage material,3 also can be used as an electrolyte in Li ion batteries.44 C60, which serves as a catalyst for dehydrogenation of complex metal hydrides,48 can also improve Li ion conductivity in Li ion batteries.49 Realization that the complex anions in currently used electrolytes are superhalogens has been helpful in designing halogen-free electrolytes.41 In addition, superhalogens can also account for the underlying stability of intermediate phases29 during dehydrogenation of complex hydrides as well as provide a path to make them safer.40 Thus, superhalogens serve as a bridge between complex metal hydrides and electrolytes in Li ion batteries, and the knowledge gained in superhalogen research can be useful in designing safer complex metal hydrides and electrolytes. Finally, a word on how this understanding can further help in the future design and synthesis of improved hydrogen storage and Li ion battery materials is provided. As pointed out before, bimetallic borohydride, KAl(BH4)4, is not only safer but also has better thermodynamic properties than its building blocks,40 KBH4 and Al(BH4)3. These studies can be extended to a large class of bimetallic borohydrides, and recent experiments50−54 on NaZn2(BH4)5, NaZn(BH4)3, KCd(BH4)3, K2Cd(BH4)4, LiSc(BH4)4, and NaSc(BH4)4 are already showing promising results. In a recent article, Cerny and co-workers55 have demonstrated that merging the concepts of molecular chemistry with ceramic host lattices can lead to the design and synthesis of an unusual class of complex hydride perovskite materials. Note that the traditional perovskites are metal oxides with composition ABO3, where A and B are metal cations. The lightest element, hydrogen, is rarely encountered in oxide perovskites. Yet, these authors have been able to synthesize 30 new complex hydride perovskite-type materials by using the

Figure 5. Ionic (●) and electrical (■) conductivity measurements of the LiBH4/C60 nanocomposites utilizing a symmetrical cell set up Li/ SSE/Li (SSE = LiBH4-based solid-state electrolyte). The activation energy for each material was also calculated from this data. Blue, LiBH4/C60 (70:30) annealed at 300 °C; green, LiBH4/C60 (50:50) annealed at 300 °C; red, LiBH4/C60 (70:30) as prepared; gray, LiBH4/ C60 (50:50) as prepared [reprinted from ref 49]. 1123

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BH4− anion instead of oxygen. Because perovskites possess a host of functionalities such as ferroelectricity and have potential for solar cells,56 the ability to replace oxygen anions with superhalogen moieties such as BH4− will expand the scope of nontraditional perovskites and provide opportunities for focused discovery of new materials. Similarly, use of temperature to control decomposition pathways of complex hydrides may lead to unexpected results, as seen by Zuttel and coworkers (unpublished results) in the case of Ca(BH4)2. A fundamental understanding of the electronic structure of the anionic components of cathode materials of Li ion batteries may also lead to new cathodes with improved performance.

(4) Jena, P. Materials for Hydrogen Storage: Past, Present, and Future. J. Phys. Chem. Lett. 2011, 2, 206−211. (5) Goodenough, J. B.; Kim, Y. Challenges for Rechargeable Li Batteries. Chem. Mater. 2010, 22, 587−603. (6) Dresselhaus, M.; Crabtree, G.; Buchanan, M.; Mallouk, T.; Mets, L.; Taylor, K.; Jena, P.; DiSalvo, F.; Zawodzinski, T. Basic Energy Needs for the Hydrogen Economy: Report of the Basic Energy Sciences Workshop on Hydrogen Production, Storage, and Use, May 13−15, 2003. http:// www.sc.doe.gov/bes/hydrogen.pdf (2003). (7) Soulie, J.-P.; Genaudin, G.; Cerny, R.; Yvon, K. Lithium Borohydride LiBH4: I. Crystal Structure. J. Alloys Compd. 2002, 346, 200− 205. (8) Wang, X.-B.; Ding, C.-F.; Wang, L.-S.; Boldyrev, A. I.; Simons, J. First Experimental Photoelectron Spectra of Superhalogens and Their Theoretical Interpretations. J. Chem. Phys. 1999, 110, 4763−4771. (9) Gutsev, G. L.; Rao, B. K.; Jena, P.; Li, X.; Wang, L.-S. On the Origin of the Unusual Stability of MnO4−. Chem. Phys. Lett. 1999, 312, 598−605. (10) Alexandrova, A. N.; Boldyrev, A. I.; Fu, Y.-J.; Yang, X.; Wang, X.B.; Wang, L.-S. Structure of the NaxClx+1 (x=1−4) Clusters via Ab Initio Algorithm and Photoelectron Spectroscopy. J. Chem. Phys. 2004, 121, 5709−5719. (11) Elliott, B. M.; Koyle, E.; Boldyrev, A. I.; Wang, X.-B.; Wang, L.S. MX3− Superhalogens (M = Be, Mg, Ca; X = Cl, Br): A Photoelectron Spectroscopic and Ab Initio Theoretical Study. J. Phys. Chem. A 2005, 109, 11560−11567. (12) Giri, S.; Behera, S.; Jena, P. Superalkalis and Superhalogens as Building Blocks of Supersalts. J. Phys. Chem. A 2014, 118, 638−45. (13) Orimo, S.; Nakamori, Y.; Ohba, N.; Miwa, K.; Aoki, M.; Towata, S.; Züttel, A. Experimental Studies on Intermediate Compound of LiBH4. Appl. Phys. Lett. 2006, 89, 021920. (14) Ohba, N.; Miwa, k.; Aoki, M.; Noritake, T.; Towata, S.; Nakamori, Y.; Orimo, S.; Zuttle, A. First-Principles Study on the Stability of Intermediate Compounds of LiBH4. Phys. Rev. B 2006, 74, 075110. (15) Hwang, S. J.; Bowman, R. C.; Reiter, J. W., Jr.; Rijssenbeek, J.; Soloveichik, G. L.; Zhao, J. C.; Kabbour, H.; Ahn, C. C. NMR Confirmation for Formation of [B12H12]2− Complexes during Hydrogen Desorption from Metal Borohydrides. J. Phys. Chem. C 2008, 112, 3164−3169. (16) Her, J. H.; Yousufuddin, M.; Zhou, W.; Jallsatgl, S. S.; Kulleck, J. G.; Zan, J. A.; Hwang, S. J.; Bowman, R. C., Jr.; Uovic, T. J. Crystal Structure of Li2B12H12: A Possible Intermediate Species in the Decomposition of LiBH4. Inorg. Chem. 2008, 47, 9757−9759. (17) Purewal, J.; Hwang, S. J.; Bowman, R. C., Jr.; Rönnebro, E.; Fultz, B.; Ahn, C. C. Hydrogen Sorption Behavior of the ScH2−LiBH4 System: Experimental Assessment of Chemical Destabilization Effects. J. Phys. Chem. C 2008, 112, 8481−8485. (18) Her, J. H.; Zhou, W.; Stavila, V.; Brown, C. M.; Udovic, T. J. Role of Cation Size on the Structural Behavior of the Alkali-Metal Dodecahydro-closo-Dodecaborates. J. Phys. Chem. C 2009, 113, 11187−11189. (19) Kim, C.; Hwang, S. J.; Bowman, R. C., Jr.; Reiter, J. W.; Zan, J. A.; Kulleck, J. G.; Kabbour, H.; Majzoub, E. H.; Ozolins, V. LiSc(BH4)4 as a Hydrogen Storage Material: Multinuclear HighResolution Solid-State NMR and First-Principles Density Functional Theory Studies. J. Phys. Chem. C 2009, 113, 9956−9968. (20) Friedrichs, O.; Remhof, A.; Hwang, S. J.; Züttel, A. Role of Li2B12H12 for the Formation and Decomposition of LiBH4. Chem. Mater. 2010, 22, 3265−3268. (21) Li, H. W.; Kikuchi, K.; Sato, T.; Nakamori, Y.; Ohba, N.; Aoki, M.; Miwa, K.; Towata, S.; Orimo, S. Synthesis and Hydrogen Storage Properties of a Single-Phase Magnesium Borohydride Mg(BH4)2. Mater. Trans. 2008, 49, 2224−2228. (22) Li, H. W.; Kikuchi, K.; Nakamori, Y.; Ohba, N.; Miwa, K.; Towata, S.; Orimo, S. Dehydriding and Rehydriding Processes of WellCrystallized Mg(BH4)2 Accompanying with Formation of Intermediate Compounds. Acta Mater. 2008, 56, 1342−1347.

The aim of this Perspective is to show how seemingly disparate fields have a common link and how a fundamental understanding of this link may be useful for advancing materials design and synthesis. The aim of this Perspective is to show how seemingly disparate fields have a common link and how a fundamental understanding of this link may be useful for advancing materials design and synthesis.

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AUTHOR INFORMATION

Notes

The authors declare no competing financial interest. Biography Puru Jena is Distinguished Professor of Physics at Virginia Commonwealth University (VCU) and Jefferson Science Fellow alumnus at the U.S. Department of State. He is a Fellow of the American Physical Society and recipient of Virginia’s Outstanding Scientist and Outstanding Faculty Award as well as VCU’s Presidential Medallion, Award of Excellence, and Distinguished Scholar Award. His research interests include clusters, cluster-assembled materials, surfaces, interfaces, and hydrogen storage. http://physics.vcu.edu/ jenasgroup/

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ACKNOWLEDGMENTS This work is partially supported by the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Materials Sciences and Engineering under Award # DE-FG0296ER45579. I am thankful to Professors Qiang Sun, Qian Wang, and Anil Kandalam for a critical reading of the manuscript. I acknowledge the computational resources of the National Energy Research Scientific Computing Center, which is supported by the Office of Science of the U.S. Department of Energy under Contract No. DE-AC02-05CH11231.

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REFERENCES

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