20 First-Principles Design of Complex Chemical

20

First-Principles Design of Complex Chemical Hydrides as Hydrogen Storage Materials S. Bhattacharya and G. P. Das

Contents 20.1 Renewable Energy......................................................................................... 415 20.2 Materials for Hydrogen Storage..................................................................... 416 20.3 Computational Approach............................................................................... 418 20.4 Complex Hydride: Case Study of Lithium Imide.......................................... 419 20.5 Chemical Hydride: Case Study of Monoammoniated Lithium Amidoborane................................................................................................. 424 20.6 Concluding Remarks..................................................................................... 428 Acknowledgments................................................................................................... 428 References............................................................................................................... 429

20.1 Renewable Energy Energy is undoubtedly the most critical issue being faced by humanity for their very survival on Earth. Although there are various sources of energy that are available, starting from the Sun, which is a very large source, humans tend to use mostly fossil fuels to satisfy their need as well as their greed. However, the fossil fuel reserve on Earth is finite, and with our increasing population and materialistic needs, the intense use of fossil hydrocarbons causes some disproportionate increase of CO2 in the atmosphere, thereby causing global warming with disastrous environmental consequences. So no matter how long our fossil fuel reserve is going to last, it is a forgone conclusion that we do need to switch over to renewable energy sources, such as solar, hydroelectric, geothermal, tidal wave, wind, and bioenergy, and thereby reduce our dependence on fossil fuels [1]. However, it is easier said than done because of the simple reason that the energy contents per kilogram or liter of fossil fuels are typically more than the other sources; in addition, it is easier to simply dig out coal or petroleum that Mother Nature has taken millions of years to generate. This makes fossil fuel the most natural choice for domestic, commercial, automobile, aviation, and other sectors. Also, there are other issues that have got to do with global politics rather than science and technology. To 415

416

Modern Theoretical Chemistry: Electronic Structure and Reactivity

exploit the solar radiation incident on Earth (~85,000 TW), which is several orders of magnitude more than the entire energy need of the world (~15 TW), one needs to have suitable material that can store the energy and make it available for usage as and when needed. Extensive efforts are on to develop suitable materials for solar photovoltaic and solar thermal systems that can maximize efficiency as well as cost effectiveness [2]. Energy production, energy storage, and energy transportation are three issues that require multidisciplinary efforts. Hydrogen, which is the lightest element in the periodic table and the most abundant element in the universe, exhibits the highest energy density per unit mass (heating value) of all chemical fuels [3,4]. Furthermore, hydrogen can be produced by electrolysis from renewable energy sources and is therefore regenerative. It produces only water when burnt and is therefore environment-friendly. Hydrogen is not a “source” of energy but a “carrier” of energy, and it does not occur naturally as a gas on Earth. The primitive phase diagram of hydrogen shows that the critical point lies at a very low temperature of 33 K, implying that free H2 molecules experience strong repulsive interaction; for hydrogen storage, one needs to somehow reduce this via interaction with another suitable material that has affinity for it. In fact, there are some metal hydrides that can store hydrogen with a density higher than that of liquid hydrogen. It is exciting as well as challenging to probe the possibility of storing hydrogen in a more compact and safer way compared with pressurized gas and cryogenic liquid [5]. Efficient hydrogen storage is a challenge faced by the materials scientists who have been trying out various kinds of materials, from bulk materials to nanomaterials, that show high volumetric density of H-atoms present in the host lattice. High storage capacity, satisfactory kinetics, and optimal thermodynamics are some of the essential criteria for a potential hydrogen storage material.

20.2 Materials for Hydrogen Storage There are various ways for efficient storage of molecular hydrogen in solid state [6,7], for which a fundamental understanding of how hydrogen interacts with materials is of utmost importance [4]. Most of the metals in the periodic table, their alloys, or intermetallic compounds react with hydrogen to form metal hydrides. Hydrogen tends to go into the metal lattice as an octahedral/tetrahedral interstitial and form M + (x/2)H2 ↔ MH x (x can be integral or nonintegral for stoichiometric and offstoichiometric hydrides) by hybridizing with the metal band. The bonding between hydrogen and the metal can range from very covalent to very ionic as well as multicentered bonds and metallic bonding. Some elemental metals, such as Mg, Al, Ti, and Pd, show special affinity for hydrogen, whereas some, such as Pt, Ru, and Ni, act as good hydrogenation catalysts. A classic textbook example is palladium hydride (PdH x) that can retain a substantial quantity of hydrogen within its crystal lattice. At room temperature and atmospheric pressure, palladium can adsorb up to 900 times its own volume of hydrogen in a reversible process. However, Pd is a heavy metal and hence does not yield good gravimetric efficiency, apart from the fact that it is quite costly. Intermetallic hydrides ABxHn are formed with A, an alkaline earth (AE) or a rare earth (RE) metal acting as a H-adsorber, and B, a transition metal (TM) acting as a H-activator. Appropriate A–B combination allows tailoring of hydride properties with hydrogen landing in the interstitial space as Ho neutral (e.g., FeTiH x and

First-Principles Design of Complex Chemical Hydrides

417

LaNi5H6). Perovskite hydrides of ABH3 structure (A is a monovalent alkali metal like K, Sr, Cs, and Rb, whereas B is a divalent alkaline earth metal like Ca and Mg), particularly Mg-based compounds, receive particular attention because of their lightweight characteristic and low-cost production. There are plenty of open areas of research that are yet to be explored both experimentally and theoretically for perovskite hydrides and their possible application in hydrogen storage [8]. Now, we come to complex hydrides formed by a combination of metals or metalloids (e.g., imides, alanates, borates, borohydrides, and aminoboranes) of low-Z elements, where the basic interaction tends to have partially ionic, partially covalent character that can be tuned. The hydrogen atoms are bonded covalently to a metal or metalloid atom to form a complex anion such as (NH2)–, (BH4)–, and (A1H4)–. This anion is then bonded ionically to the M-cation present to form a complex metal hydride [9]. In general, complex metal hydrides have the formula A xMyHz, where A is an alkali metal or alkaline earth metal cation or cation complex and M is a metal or metalloid. Well-known examples feature anions of hydrogenated group 3 elements, particularly boron and aluminum. Compounds such as LiBH4 (lithium borohydride) and NaAlH4 (sodium alanate) are among the most widely studied. The variety in complex metal hydrides is very large. The possibility of forming complex metal hydrides using lightweight elements opens a promising route to achieve very high hydrogen content by weight; for example, LiBH4 contains 18 wt.% hydrogen. Accordingly, there is an increasing interest to explore complex metal hydride systems and their subsequent optimization for practical use. Combining several complex hydrides into one storage system might improve the storage characteristics, but the complexity of reaction mechanisms requires further fundamental research on such materials. Apart from complex hydrides, there are other kinds of novel materials that have been investigated, for example, carbon-based materials activated with nanocatalysts [10], clathrate hydrates [11], metal-organic complexes [12], and more recently, nanostructured cages, viz., fullerenes and nanotubes, decorated with simple or transition metals that serve to attract hydrogen in molecular form [13–17]. Nanostructure materials built from lightweight elements, such as boron, carbon, and nitrogen, have several attractive features, viz., large surface area, low density, and high structural stability, which can be exploited for efficient storage of hydrogen. The storage takes place as hydrogen molecules are adsorbed on the surface of the solids. The possibility of storing hydrogen in molecular form is advantageous over chemical storage in atomic form, which requires the dissociation of the hydrogen bond and the formation of a hydride. To understand and exploit these materials for H-storage, it is crucial to know the way hydrogen interacts with the surface or the bulk. There are mainly three ways in which hydrogen can be adsorbed on a material [4]:

1. Physisorption, where hydrogen remains in the molecular state (H2) and gets bound on the surface rather weakly (BE approximately 10−100 meV) 2. Chemisorption, where H2 dissociates into H atoms that migrate and gets strongly bound to the material (BE approximately 2–4 eV range) 3. Molecular chemisorption, where H–H bonding gets weakened but not broken (still H2 molecular state is retained) and the strength of the binding is intermediate between physisorption and chemisorption

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It is this third form of the quasimolecular bonding that is most suitable for optimal absorption and desorption of hydrogen. The basic quasimolecular interaction and bonding of hydrogen can be explained via what is known as Kubas interaction [18], that is, donation of charge from H2 molecule to the unfilled d-orbitals of the TM atoms and back-donation from the TM atom to the antibonding orbital of H2 molecule. Kubas interaction has been exploited for designing transition metal decorated nanomaterials [19,20], metal organic frameworks (MOFs) [21,22], spillover catalysts [23], and other kinds of materials for H-storage. However, it has limitations in explaining the bonding in alkaline earth metal complexes and alkali metaldoped nanostructures [24]. In the case of metal clusters containing a few atoms, the way hydrogen interacts is fundamentally different from bulk, and the reactivity and adsorption behavior change drastically with the addition and subtraction of a few metal atoms [25,26].

20.3 Computational Approach First-principles computational approach plays a crucial role in predicting the H2 adsorption and desorption processes in complex hydrides and also their decomposition pathways. Calculations are performed within the so-called density functional theory (DFT) [27,28], where the exchange correlation potential is treated via some mean field approximation, and the problem of solving an inhomogeneous many-electron system is reduced to that of solving an effective one-electron Schrödinger equation with an effective potential. Such an effective single-particle approach has been embraced by materials scientists mainly because it provides a reliable computational tool yielding material-specific quantitative results with desirable accuracy for the ground state (cohesive, electronic, magnetic, etc.) properties of a large variety of systems. Firstprinciples DFT calculations, based on local density approximation (LDA) and its improved variants such as generalized gradient approximation (GGA), have reached an unprecedented level of accuracy and reliability such that one not only can explain but also can predict material properties and phenomena [29]. The most widely used first-principles electronic structure method for materials with fixed geometry is based on either plane wave–based methods or linear methods or localized basis set methods. From the ground state total energy, one can estimate the force acting on the atoms, which is essential to do ab initio molecular dynamics simulations à la recipe proposed by Car and Parrinello within the DFT framework. Such dynamical simulation enables one to determine the so-called “energy landscape,” that is, how the energy of a system evolves with the position of the atoms, to monitor the making and breaking of chemical bonds, for example, desorption of hydrogen molecule from a nanohost material as a function of temperature. For the present studies, we have used state-of-the-art DFT-based methods with plane wave basis set, viz., VASP [30] with PAW potentials [31] for extended systems, and with localized atomic orbital or Gaussian basis set, viz., DMol3 [32] or GAUSSIAN03 [33] for molecular or cluster systems. In all our calculations, the ions are steadily relaxed toward equilibrium until the Hellmann–Feynman forces are converged to less than 10 –3 eV/Å. Available experimental structural data have been used as input for some of the hydrides whenever they are available.

419


The calculations that we have carried out for studying the above-mentioned materials and phenomena can be broadly classified into the following three categories: (1) ground state geometry, electronic structure, and activation barrier estimation using the so-called nudge elastic band method [34] for different possible configurations; (2) transition state calculations and reaction pathways; and (3) ab initio molecular dynamics with Nose thermostat for estimating the desorption kinetics. There are a number of detailed investigations reported in the literature [35–37] where different kinds of complex hydrides have been treated using DFT at various levels of sophistication.

20.4 Complex Hydride: Case Study of Lithium Imide Complex hydrides involving light metals show impressive gravimetric efficiencies, but the desorption temperature of H2 is rather high. For example, amides and imides of low-Z alkali metals such as Li and Na are prospective candidates for hydrogen storage with ∼7% to 10% gravimetric efficiency. Here, the dehydriding reaction takes place in one or more steps at varying desorption temperatures depending on the kinetic reaction barrier. It was demonstrated by Chen et al. [38] how lithium amide (LiNH2) reacts with lithium hydride (LiH) to yield lithium imide (Li2NH) or lithium nitride (Li3N) and molecular hydrogen. The forward reaction results in desorption, whereas the reverse reaction results in absorption. The total reaction is a two-step reaction process (Figure 20.1) as follows: Step 1: Li3N + H2 ↔ Li2NH + LiH Step 2: Li2NH + H2 ↔ LiNH2 + LiH Total reaction: Li3N + 2H2 ↔ Li2NH + LiH + H2 ↔ LiNH2 + 2LiH.

(20.1) (20.2) (20.3)

The reaction is exothermic with ∆H ∼ −96 kJ/mol H2, whereas the gravimetric efficiency turns out to be ∼10 wt.%. This dehydrogenation reaction leading to 100 10 1 0.1 100 10 1 0.1 100 10 1 0.1

10

wt. % H2

Pressure (bar)

Absorption

8 6 4

Desorption

2 0 50

350 150 250 Temperature (ºC)

TG results of a Li3N sample

450

Ab

Ab

255°C Des.

Des.

Li2NH Re-Li3N Li3N

Ab

Des

0.0 0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 H/Li-N-H P-C-T curves of Li3N and Li2NH

FIGURE 20.1 Absorption and desorption of Li amide and Li imide undergoing reversible reaction to produce H2. (Reproduced from Chen, P. et al., Nature, 420, 302, 2002. With permission.)

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release of hydrogen is reversible, which is an additional attractive feature, although this reversibility is not always guaranteed, as will be discussed in this article. Lithium atoms are ionized as Li+ cations, whereas [NH2]– forms a complex anion, and it is the strength of the interaction between Li+ and [NH2]– that dictates the enthalpy of reactions and hence the desorption kinetics of H2. One way to do this is to alloy the binary hydride with some divalent alkaline earth metal such as Ca or Mg (Figure 20.2a) [9], and these ternary complex hydrides have been synthesized in recent years [39]. For pure lithium imide, the orthorhombic Ima2 structure (Figure 20.2) is found to have lower ground state energy (by ∼0.6 eV) than that of the cubic struc ture (F-43m). The Ca-doped system, viz., Li2Ca(NH)2, has a trigonal structure space group P-3m1, with H occupying different possible 2d or 6i positions (Table 20.1) and three possible orientations of the N–H bond [40] The Li atoms in the 2d site occupy the tetrahedral hole created by the N-lattice (Figure 20.2b), whereas Ca atoms in the 1b site occupy the trigonal prismatic hole created by the N-lattice (Figure 20.2c). Li–Ca separation (≈3.077 Å) is ≈25% elongated as compared with the Li–Li bond length, which is very similar to the asymmetric Li–Li bond lengths in Li3N. The most crucial, however, is the N–H bond where hydrogen can partially or randomly occupy the three possible positions with one-third probability of occupancy at any C

Li

Li H

H

Li

N

N Ca

Ca

Ca

H

Ca

A

(a)

H

N

N

Li

Li

N

Li

N

N Ca

Ca N

N

Ca

A

Ca

N Li

Ca

Ca

Ca N H

H N

N

(c)

H

Li

Ca Ca

Ca

N Ca

Ca

Ca

Li

0

H

N

B

Ca

(b)

N Li

N

H

Li

Ca

Ca N H

Li

Ca

B

(d) Hydrogen

Lithium

Nitrogen

Calcium

FIGURE 20.2 (a) Li2Ca(NH)2 structure with hydrogen occupying any of the three equivalent positions above and below the respective nitrogen atoms. (b) Tetrahedral hole created by the N-lattice that is occupied by Li → 2d site (⅓, ⅔, 0.8841). d(Li–N) = 2.223 Å ×3, d(Li–Li) = 2.479 Å. (c) Octahedral hole created by the N-lattice that is occupied by Ca → 1b site (0, 0, ½). d(Ca–N) = 2.518 Å ×6, d(Ca–Ca) = 3.566 Å. (d) Supercell constructed by repeating the unit cell along the c-axis with different possible N–H bond orientations.

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TABLE 20.1 Estimated Bond Lengths, Reaction Enthalpies, and Hydrogen Removal Energies for Undoped and Doped Li Imides

System LiNH2 Li2NH Li2Ca(NH)2 Li2Mg(NH)2

Structure Space Group (Formula Unit) Tetragonal l-4 (4) Orthorhombic Ima2 (8) Trigonal P-3m1 (3) Orthorhombic lba2 (16)

Chemical Reaction

Reaction Enthalpy, ∆H (kJ/ mol H2)

∆EH (kJ/ mol H2)

Average N-H Bond Length (Å)

LiNH2 + LiH ⇆ Li2NH + H2

68.9

268

1.03

Li2NH + LiH ⇆ Li3N + H2

108.8

192

1.04

3Li2Ca(NH)2 + 2LiH ⇆ 4Li2NH + Ca3N2 + 2H2 3Li2Mg(NH) 2 + 2LiH ⇆ 4Li2NH + Mg3N2 + 2H2

102.6

181

1.05

82.8

183

1.05

Source: Bhattacharya, S. et al., J. Phys. Chem. B, 112, 11381, 2008. With permission.

instant of time. To mimic this quantum delocalization effect, we have prescribed a supercell by repeating the unit cell three times along the c-axis and considered all these three possible H-positions built in the first, second, and third one-third of the super cell (Figure 20.2d). We have estimated the total energy for all the three different configurations designated as 6i (set-1), 6i (set-2), and 6i (set-3) and found that configuration 6i (set-2) is energetically most favorable. The electronic structures of each of the constituents of Equations 20.1 through 20.3 and their Ca-doped counterparts as well as the heats of reactions (exothermic) have been estimated from first-principles density functional calculations [40]. The resulting total as well as partial densities of states (Figure 20.3) clearly show a semiconducting behavior with a GGA band gap of ∼2.3 eV and a two-humped structure of the occupied part of electronic density of states (DOS) (Figure 20.3). N-2p bands predominantly contribute to the occupied DOS, as expected, whereas H-s character prevails in the lowest occupied band. Li behaves like a cation transferring its electron to [NH]2–. The parent compound Li2CaN2 of the ternary hydride Li2Ca(NH)2 turns out to be metallic with the Fermi level lying near the antibonding peak. However, on introduction of hydrogen in this anti-La2O3 structure, this peak is pushed down below the Fermi level, thereby opening a band gap of ∼2.3 eV, which is very similar to that of the pure imide. The Ca-4s band has a dominant contribution to the DOS and also affects the bonding between the Li+ cation and the [NH]2– anion. The lower bonding peak arises because of strong hybridization between H-s and N-p orbitals. The upper bonding peaks arise out of Li–N interactions. The average N–H bond lengths, hydrogen removal energies, and the enthalpy of formation of Li2Ca(NH)2 and Li2Mg(NH)2 have been estimated, and the results have been compared with the same quantities estimated for the pure Li imides and amides. The enthalpy of formation is the most fundamental and important quantity for hydrogen storage materials, which can be estimated from the difference between

422

Modern Theoretical Chemistry: Electronic Structure and Reactivity Li2NH

10

N–DOS

2s 2p

0.75

0.4

2s 2p

0.2

Li–DOS

Li–DOS

0.4

0.0 –8 –6 –4 –2 0 2 4 6 8 2

0.00 –8 –6 –4 –2 0 2 4 6 8

0.2

0.0 –8 –6 –4 –2 0 2 4 6 8 1.50

2s 2p

1

2s 2p

2s 2p

N–DOS

1.50

0 –8 –6 –4 –2 0 2 4 6 8

0 –8 –6 –4 –2 0 2 4 6 8

N–DOS

Li–DOS

0.0 –8 –6 –4 –2 0 2 4 6 8

0.75

0.00 –8 –6 –4 –2 0 2 4 6 8

0 –8 –6 –4 –2 0 2 4 6 8

0.4

1s 2s

H–DOS

H–DOS

Density of states (states/eV f.u)

2s 2p

0.2

0.2

20

10

0 –8 –6 –4 –2 0 2 4 6 8

Li2Ca(NH)2

Total DOS

Total DOS

10

0.4

30

20

20

0.4

Li2CaN2

30

Total DOS

30

1s 2s

0.2

0.0 –8 –6 –4 –2 0 2 4 6 8

0.0 –8 –6 –4 –2 0 2 4 6 8

1

Ca–DOS

Ca–DOS

2

3p 3d 4s

2

1

3p 3d 4s

0 –8 –6 –4 –2 0 2 4 6 8

0 –8 –6 –4 –2 0 2 4 6 8

Energy in eV

FIGURE 20.3 Total and partial electronic densities of states calculated for Li2NH, Li2CaN2, and Li2Ca(NH)2.

the energies before and after hydriding reaction (Equation 20.4). The enthalpy change in a reaction at 0 K was calculated using

∆H =

∑E

products

−

∑E

reactants

(20.4)

where E is the total energy of one of the bulk structures of interest as calculated by DFT. We investigate the thermodynamics of hydrogen release from the mixture of Li2Ca(NH)2 and LiH, which allows us to draw comparisons with the thermodynamics of hydrogen release from the other Li–N–based compounds, viz., parent imides and amides, along with Li2Ca(NH)2 with LiH. Table 20.1 summarizes our results for the

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specific exothermic chemical reactions that take place for H2 desorption in different binary and ternary hydrides. We observe that ∆H decreases from 108.8 kJ/ mol (1.13 eV) H2 in Li-imide to 102.6 kJ/mol (1.06 eV) H2 and to 82.8 kJ/mol (0.86 eV) H2 for Ca and Mg ternary imides, respectively (Figure 20.4 and 20.5). The corresponding ∆H value estimated by Araujo et al. [41] is 118 kJ/mol H2 for Li2NH assuming Pnma space group and is 84 kJ/mol H2 for Li2Mg(NH)2. It is interesting to note that the N–H bond lengths increase on ternary addition, indicating the weakening of the N–H bonds. Hydrogen removal energy ∆EH for Li2Ca(NH)2 has been calculated using the relation

∆EH [Li 2Ca ( NH)2 ] = ET [Li 6Ca 3 N 6H 5 ] +

1 ET [H 2 ] − ET [Li 6Ca 3 N 6H 6 ], (20.5) 2

Hydrogen desorption (a.u.)

0.1 MPa argon, 10 K min–1 MNH2 (M = Li – x at. %Mg)

300

x = 30

x = 10

x=0

400 500 Temperature (K)

600

FIGURE 20.4 Hydrogen desorption reaction of lithium imide as a function of ternary addition (x% Mg). For x = 30, the desorption starts at around 370 K. (Reproduced from Orimo, S.-I. et al., Chem. Rev., 107, 4111, 2007. With permission.)

6

LiNH2 Li 2NH Li 2Mg(NH)2 Li 2Ca(NH)2

H 2 removal energy in eV

5

4 3 2 1 0

FIGURE 20.5 Estimated hydrogen removal energies for undoped and doped Li imides.

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where ET[Li6Ca3N6H5], ET[Li6Ca3N6H6], and ET[H2] are the ground state total energies of Li6Ca3N6H5, Li6Ca3N6H6 cell, and H2 molecule in the gas phase, respectively. For Li2Ca(NH)2, we have used the optimized structure shown in Figure 20.2. It is interesting to note that ∆EH reduces by ~5.5% for the ternary Ca imide. The enthalpy of reaction ∆H = T∆S for pure lithium imide decreases on ternary addition. Assuming the entropy change ∆S to remain more or less constant during the reactions, the dehydrogenation temperature T is expected to come down to a desirable range. The H removal energy correspondingly decreases by about 5.5% with a concomitant increase in the N–H bond length by about 0.01 Å for the ternary Ca imide system.

20.5 Chemical Hydride: Case Study of Monoammoniated Lithium Amidoborane Ammonia borane (NH3BH3), AB for short, complexes have emerged as attractive candidates for solid-state hydrogen-storage materials because of their high percentage of available hydrogen (19.6 wt.%). However, relatively poor kinetics and high temperature of dehydrogenation as well as release of volatile contaminants, such as borazine, are posing big challenges for practical application of AB [42,43]. When one H atom in AB is replaced by an alkali or alkaline earth metal (M), a new class of materials called metal amidoboranes (MABs) is formed, which in turn can be used for efficient storage of molecular hydrogen [44,45]. These materials were highlighted as some of the best potential hydrogen storage materials in the 2008 Department of Energy (DOE) hydrogen program annual progress report. For example, LiAB and NaAB provide high storage capacity of 10.9 and 7.5 wt.%, respectively [46,47] at easily accessible temperatures without the unwanted release of borazine. LiNH2BH3 is environmentally harmless and stable in solid state at ambient temperature and normal pressure. The bonds get distorted as compared with those in pristine ammonia borane, as can be seen from the ball-and-stick model optimized geometries and corresponding bond lengths (Figure 20.6). However, to improve the operating properties of these materials, such as rapid H2

H

H H

Li

B

N H

H

FIGURE 20.6 Optimized structure of LiNH2BH3. Equilibrium bond distances estimated from first-principles calculations are B–N = 1.61 Å, Li–N = 1.85 Å, Li–B = 2.09 Å, B–H = 1.25 Å (av), and N–H = 1.01 Å (av).

425


release near room temperature, it is vital to understand the underlying mechanism for the release of H2. Recent experimental and computational studies have shown that NH3 reacts with LiAB to yield H2, and the dehydrogenation takes place in three different stages [48], each time resulting in an intermediate metastable product (adduct). The steps involved in the first dehydrogenation reactions along with the transition states and the intermediate products are shown in Figure 20.7 (similar things have been determined for the second and third dehydrogenation processes, which are not shown), whereas the estimated values from our first-principles calculations are summarized in Table 20.2. The first dehydrogenation from monomer occurs with an activation barrier of 78 kJ/mol (1 eV ≈ 100 kJ/mol) followed by H2 removal energy of 0.16 eV/ H2, leaving a metastable product Li(NH2)NH2BH2. We have found a transition state (Figure 20.8) where the hydric B–H bond in the [NH2BH3] unit interacts with the protic N–H bond of NH3, which in turn leads to H2 release from the system as a first dehydrogenation process. The reaction pathway having the minimum activation barrier has been estimated using transition state calculations and is shown schematically in Figure 20.8. Similarly, we have determined using our first-principles approach the second and third dehydrogenation processes that result in relatively high activation barriers and H2 removal energies (shown in Table 20.2), whereas the metastable products left behind are Li(NH)NH2BH and Li(NH)NBH, respectively, with the final product matching with the available experimental results. First dehydrogenations: Equation: Li-NH2-BH3 + NH3= LiNH2—NH2BH2 + H2 Step-1 Reactant: Li-NH2-BH3 + NH3

Step-2 Complex: Li-NH2-BH3 + NH3

Step-3 Translate state: Li-NH2—NH2−H2

Step-4 Product= LiNH2—NH2BH2+H2

Energy in Hartree

−146.6

Step 1: Reaction between LiAB and NH3 Step 2: Stable complex formation of LiAB-NH3 Step 3: Transition state before first dehydrogenation of LiAB-NH3 complex Step 4: Products after first dehydrogenation

−146.623 78 KJ/mole

−146.645

−146.668

−146.69 Reactant

Complex

Transition state

Product

Activation barrier = 78 KJ/mole

FIGURE 20.7 First dehydrogenation mechanism from LiAB monomer and NH3 interaction and the corresponding reaction barrier.

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TABLE 20.2 Activation Barriers and Hydrogen Removal Energies for First, Second, and Third Dehydrogenation Energies of Monoammoniated LiAB Activation Barrier (kJ/mol H2) Dehydrogenation First Second Third Second [1] Second [2] Second [3] Third [1] Third [2] Third [3]

Monomer of LiAB + NH3

H2 Removal Energy (eV/H2)

78 105 353 [LiNH2–BH2–NH2]3 No barrier No barrier No barrier 230 236 (average ~ 236) 243

0.16 0.27 1.3 0.14 0.20 (average ~ 0.25) 0.40 0.14 0.20 (average ~ 0.67) 0.40

Source: Bhattacharya, S. et al., J. Phys. Chem. C, 116, 8859, 2012. With permission.

The reactions are, however, not reversible. The H2 removal energy defined as

EH2-removal = Etotal[LiNH2BH2NH2] − Etotal[LiNH2BHNH] − Etotal[H2] (20.6)

is found to be 0.16, 0.27, and 1.30 eV/H2 for the first, second, and third dehydrogenation processes, respectively. This increasing trend, especially the high activation barrier for the third dehydrogenation, rules out, in principle, the possibility of a spontaneous evolution of molecular hydrogen under normal conditions.

Li(NH2)NH2BH2(H-H) −146.60 −146.62 Energy in Ha

−146.64

LiAB + NH3

(NH2)LiNH2BH2 + H2

c d

a

−146.66 −146.68

b

Li(NH3)NH2BH3

e

Li(NH2)NH2BH2 + H2

−146.70

FIGURE 20.8 Three-stage dehydrogenation reaction path of NH3 reacting with LiAB, with monomer. (From Bhattacharya, S. et al. J. Phys. Chem. C, 116, 8859, 2012. With permission.)

427


Based on these observations, we argued [48] that the LiAB dehydrogenation in presence of ammonia does not occur through single-stage reaction but possibly undergoes a combined reaction mechanism. We explored the possibility of forming a higher-order cluster, especially after the first dehydrogenation when the metastable product LiNH2—BH2NH2 is reached. For the subsequent (i.e., after the first) dehydrogenation reactions, we have studied the stability of [LiNH2–BH2–NH2]n clusters with n varying from 2 to 6. The stability of these clusters can be defined as

S=

Etotal [LiNH 2 BH 2 NH 2 ]n − n × Etotal [LiNH 2  BH 2 NH 2 ]metastable . n

(20.7)

The results (Figure 20.9) reveal that as the cluster size goes up from monomer to dimer to trimer (n = 3), the relative stability keeps on increasing (i.e., more negative) and tends to saturate for n ≥ 6. The estimated activation barriers for three consecutive H2 releases for the second dehydrogenation from the residual complex cluster (n = 3) are shown in Figure 20.10. The same trend has been observed for n = 1 – 3 in the case of the third dehydrogenation. The detailed reactions and their pathways are given in reference 48. It is this reduction in the activation barrier as a function of –0.050

–0.060 –0.065 –0.070 –0.075 –0.080 2

n=2

n=1

1

4 3 Cluster size (n)

5

6

n=6

–0.085

n=3

Cluster stability (S)

–0.055

FIGURE 20.9 Relative stability of [Li(NH2)–NH2BH3]n as a function of cluster size (n).

428


−436.55

Energy in Ha

−436.60 −436.65 −436.70 −436.75 −436.80 −436.85

[(NH2)LiNH2BH2]3 ∆H = 42.9 KCal/molH2∆H = 43.5 KCal/molH2 150.6 KCal/molH2

−436.50

∆H = 40.5 KCal/molH2 ts2

ts3

ts1

[Li(NH)NH2BH]3+H2 Li3(N3H4)N3H6B3H4+H2 Li3(N3H5)N3H6B3H5+H2 [Li(NH2)NH2BH2]3

−436.90

FIGURE 20.10 Reaction path for second dehydrogenation from the [Li(NH2)–NH2BH3]3 cluster.

increasing cluster size that provides an explanation for the dehydrogenation mechanism in the monoammoniated LiAB system.

20.6 Concluding Remarks The first-principles density functional approach has been used to design efficient hydrogen storage materials, such as complex hydrides, viz., lithium imides, and chemical hydrides, viz., lithium amidoboranes, with improved dehydrogenation behavior at near ambient conditions. The enthalpy of reaction ∆H = T∆S for pure lithium imide decreases on ternary addition. Assuming the entropy change ∆S to remain more or less constant during the reactions, the dehydrogenation temperature T is expected to come down to a desirable range. The H removal energy correspondingly decreases by about 5.5% with a concomitant increase in the N–H bond length for the ternary Ca imide system. Another promising material for chemical storage of hydrogen is ammonia borane (NH3–BH3). Lithium amidoborane LiAB, in particular, provides high storage capacity ~10.9 wt.% of hydrogen at easily accessible temperature, without the release of any unwanted borazine. Our first-principles result suggests that this reaction is a three-step process, and each stage is combined with the evolution of hydrogen from the system. However, the first dehydrogenation occurs between the interactions of LiAB monomer and NH3 molecule, and the second and third dehydrogenations are multicluster interactions.

Acknowledgments The authors have been collaborating on hydrogen energy with different groups in India and abroad. They would like to thank Chiranjib Majumder, Prasenjit Sen, Yuan Ping Feng, Ping Chen, Puru Jena, and Yoshiyuki Kawazoe for many useful


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discussions in the course of the work included in this article and Sonali Barman and Amrita Bhattacharya for their contributions.

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