**2. Basics of hydrogen storage in metal hydrides**

Many metals and alloys react reversibly with hydrogen to form metal hydrides according to the reaction (1):

$$\text{Me} + \text{x} / 2\text{ H}\_2 \leftrightarrow \text{MeH}\_x + \text{Q.} \tag{1}$$

Here, Me is a metal, a solid solution, or an intermetallic compound, MeHx is the respective hydride and x the ratio of hydrogen to metal, x=cH [H/Me], Q the heat of reaction. Since the entropy of the hydride is lowered in comparison to the metal and the gaseous hydrogen phase, at ambient and elevated temperatures the hydride formation is exothermic and the reverse reaction of hydrogen release accordingly endothermic. Therefore, for hydrogen release/desorption heat supply is required.

Metals can be charged with hydrogen using molecular hydrogen gas or hydrogen atoms from an electrolyte. In case of gas phase loading, several reaction stages of hydrogen with the metal in order to form the hydride need to be considered. Fig. 1 shows the process schematically.

The first attractive interaction of the hydrogen molecule approaching the metal surface is the Van der Waals force, leading to a physisorbed state. The physisorption energy is typically of the order EPhys ≈ 6 kJ/mol H2. In this process, a gas molecule interacts with several atoms at the surface of a solid. The interaction is composed of an attractive term, which diminishes with the distance of the hydrogen molecule and the solid metal by the power of 6, and a repulsive term diminishing with distance by the power of 12. Therefore, the potential energy of the molecule shows a minimum at approximately one molecular radius. In addition to hydrogen storage in metal hydrides molecular hydrogen adsorption is a second technique to store hydrogen. The storage capacity is strongly related to the temperature and the specific

free-form, ability to store hydrogen for longer times without any hydrogen losses, cyclability as well as recyclability and costs. Further research and development is strongly required. One major advantage of hydrogen storage in metal hydrides is the ability to store hydrogen in a very energy efficient way enabling hydrogen storage at rather low pressures without further need for liquefaction or compression. Many metals and alloys are able to absorb large amounts of hydrogen. The metal-hydrogen bond offers the advantage of a very high volumetric hydrogen density under moderate pressures, which is up to 60% higher than

Depending on the hydrogen reaction enthalpy of the specific storage material during hydrogen uptake a huge amount of heat (equivalent to 15% or more of the energy stored in hydrogen) is generated and has to be removed in a rather short time, ideally to be recovered and used as process heat for different applications depending on quantity and temperature. On the other side, during desorption the same amount of heat has to be applied to facilitate the endothermic hydrogen desorption process – however, generally at a much longer time scale. On one side this allows an inherent safety of such a tank system. Without external heat supply hydrogen release would lead to cooling of the tank and finally hydrogen desorption necessarily stops. On the other side it implies further restrictions for the choice of suitable storage materials. Highest energy efficiencies of the whole tank to fuel combustion or fuel cell system can only be achieved if in case of desorption the energy required for hydrogen release can be supplied by the waste heat generated in case of mobile applications

Many metals and alloys react reversibly with hydrogen to form metal hydrides according to

Here, Me is a metal, a solid solution, or an intermetallic compound, MeHx is the respective hydride and x the ratio of hydrogen to metal, x=cH [H/Me], Q the heat of reaction. Since the entropy of the hydride is lowered in comparison to the metal and the gaseous hydrogen phase, at ambient and elevated temperatures the hydride formation is exothermic and the reverse reaction of hydrogen release accordingly endothermic. Therefore, for hydrogen

Metals can be charged with hydrogen using molecular hydrogen gas or hydrogen atoms from an electrolyte. In case of gas phase loading, several reaction stages of hydrogen with the metal in order to form the hydride need to be considered. Fig. 1 shows the process

The first attractive interaction of the hydrogen molecule approaching the metal surface is the Van der Waals force, leading to a physisorbed state. The physisorption energy is typically of the order EPhys ≈ 6 kJ/mol H2. In this process, a gas molecule interacts with several atoms at the surface of a solid. The interaction is composed of an attractive term, which diminishes with the distance of the hydrogen molecule and the solid metal by the power of 6, and a repulsive term diminishing with distance by the power of 12. Therefore, the potential energy of the molecule shows a minimum at approximately one molecular radius. In addition to hydrogen storage in metal hydrides molecular hydrogen adsorption is a second technique to store hydrogen. The storage capacity is strongly related to the temperature and the specific

Me + x/2 H2 MeHx + Q. (1)

on-board by the hydrogen combustion process and the fuel cell respectively.

**2. Basics of hydrogen storage in metal hydrides** 

release/desorption heat supply is required.

the reaction (1):

schematically.

that of liquid hydrogen (Reilly & Sandrock, 1980).

surface areas of the chosen materials. Experiments reveal for carbon-based nanostructures storage capacities of less than 8 wt.% at 77 K and less than 1wt.% at RT and pressures below 100 bar (Panella et al., 2005; Schmitz et al., 2008).

Fig. 1. Reaction of a H2 molecule with a storage material: a) H2 molecule approaching the metal surface. b) Interaction of the H2 molecule by Van der Waals forces (physisorbed state). c) Chemisorbed hydrogen after dissociation. d) Occupation of subsurface sites and diffusion into bulk lattice sites.

In the next step of the hydrogen-metal interaction, the hydrogen has to overcome an activation barrier for the formation of the hydrogen metal bond and for dissociation, see Fig. 1c and 2. This process is called dissociation and chemisorption. The chemisorption energy is typically in the range of EChem ≈ 20 - 150 kJ/mol H2 and thus significantly higher than the respective energy for physisorption which is in the order of 4-6 kJ/mol H2 for carbon based high surface materials (Schmitz et al., 2008).

Fig. 2. Schematic of potential energy curves of hydrogen in molecular and atomic form approaching a metal. The hydrogen molecule is attracted by Van der Waals forces and forms a physisorbed state. Before diffusion into the bulk metal, the molecule has to dissociate forming a chemisorbed state at the surface of the metal (according to Züttel, 2003).

Thermodynamics of Metal Hydrides: Tailoring Reaction Enthalpies of Hydrogen Storage Materials 895

In the equilibrium the chemical potentials of the hydrogen in the gas phase and the

 

Since the internal energy of a hydrogen molecule is 7/2 k*T* the enthalpy and entropy of a

<sup>7</sup> k E

<sup>7</sup> <sup>8</sup> <sup>π</sup>kM r k k ln with ( ) <sup>2</sup> <sup>h</sup> *gas p T*

> k ln E k T ln ( ) p *gas gas p p <sup>T</sup>*

Here k is the Boltzmann constant, *T* the temperature, *p* the applied pressure, EDiss the dissociation energy for hydrogen (EDiss = 4.52 eV eV/H2), MH-H the mass of the H2 molecule,

0 0

conf vibr,electr *h Ts s s s* mit

Here, *s*,conf is the configuration entropy which is originating in the possible allocations of

N ! k ln


with nis being the number of interstitial sites per metal atom: nis = Nis/NMe and *c*H the

Therefore the chemical potential of hydrogen in the solid solution (-phase) is given by

vibr,electr

*<sup>c</sup> h Ts*

 

is

H is H

H

is H

*c*

n *c*

<sup>H</sup> <sup>α</sup>

k ln n

is H

*c*

(9)

N !(N -N )!

Diss

0 5

Diss <sup>0</sup>

 . (6)

(7)

(8)

(4)

*gas metal* . (2)

<sup>2</sup> *gas h T* (3)

<sup>7</sup> <sup>5</sup> <sup>2</sup> <sup>2</sup> <sup>2</sup>

 (5)

H-H H-H

1 2 

0

In the solid solution (-phase) the chemical potential is accordingly

rH-H the interatomic distance of the two hydrogen atoms in the H2 molecule. Consequently the chemical potential of the hydrogen gas is given by

*p T*

 

,conf

*s*

,conf

and accordingly for small cH using the Stirling approximation we get

*S*

number of hydrogen atoms per metal atom: *c*H = *N*H/NMe.

 

*<sup>s</sup> p T p (T)*

 

NH hydrogen atoms on Nis different interstitial sites:

hydrogen absorbed in the metal are the same:

hydrogen molecule are

with p0 = 1.01325 105 Pa.

and

After dissociation on the metal surface, the H atoms have to diffuse into the bulk to form a M-H solid solution commonly referred to as -phase. In conventional room temperature metals / metal hydrides, hydrogen occupies interstitial sites - tetrahedral or octahedral - in the metal host lattice. While in the first, the hydrogen atom is located inside a tetrahedron formed by four metal atoms, in the latter, the hydrogen atom is surrounded by six metal atoms forming an octahedron, see Fig. 3.

Fig. 3. Octahedral (O) and tetrahedral (T) interstitial sites in fcc-, hcp- and bcc-type metals. (Fukai, 1993).

In general, the dissolution of hydrogen atoms leads to an expansion of the host metal lattice of 2 to 3 Å3 per hydrogen atom, see Fig. 4. Exceptions of this rule are possible, e.g. several dihydride phases of the rare earth metals, which show a contraction during hydrogen loading for electronic reasons.

Fig. 4. Volume expansion of the Nb host metal with increasing H content. (Schober & Wenzl, 1978)

In the equilibrium the chemical potentials of the hydrogen in the gas phase and the hydrogen absorbed in the metal are the same:

$$\frac{1}{2}\mu\_{\text{gas}} = \mu\_{\text{metal}}\,. \tag{2}$$

Since the internal energy of a hydrogen molecule is 7/2 k*T* the enthalpy and entropy of a hydrogen molecule are

$$\mathbf{h}\_{\text{gas}} = \frac{\mathbf{T}}{\mathbf{2}} \cdot \mathbf{k} \cdot \mathbf{T} - \mathbf{E}\_{\text{Diss}} \tag{3}$$

and

894 Thermodynamics – Interaction Studies – Solids, Liquids and Gases

After dissociation on the metal surface, the H atoms have to diffuse into the bulk to form a M-H solid solution commonly referred to as -phase. In conventional room temperature metals / metal hydrides, hydrogen occupies interstitial sites - tetrahedral or octahedral - in the metal host lattice. While in the first, the hydrogen atom is located inside a tetrahedron formed by four metal atoms, in the latter, the hydrogen atom is surrounded by six metal

Fig. 3. Octahedral (O) and tetrahedral (T) interstitial sites in fcc-, hcp- and bcc-type metals.

In general, the dissolution of hydrogen atoms leads to an expansion of the host metal lattice of 2 to 3 Å3 per hydrogen atom, see Fig. 4. Exceptions of this rule are possible, e.g. several dihydride phases of the rare earth metals, which show a contraction during hydrogen

Fig. 4. Volume expansion of the Nb host metal with increasing H content. (Schober & Wenzl,

atoms forming an octahedron, see Fig. 3.

(Fukai, 1993).

1978)

loading for electronic reasons.

$$\mathbf{s}\_{\text{gas}} = \frac{7}{2} \cdot \mathbf{k} - \mathbf{k} \cdot \ln \frac{p}{p\_0(T)} \quad \text{with} \quad p\_0(T) = \frac{8 \left( \text{rk} \, T \right)^{\frac{1}{2}} \cdot \mathbf{M}\_{\text{H+H}} \stackrel{\frac{2}{2}}{\cdot} \cdot \mathbf{r}\_{\text{H+H}}}{\text{h}^5} \tag{4}$$

Here k is the Boltzmann constant, *T* the temperature, *p* the applied pressure, EDiss the dissociation energy for hydrogen (EDiss = 4.52 eV eV/H2), MH-H the mass of the H2 molecule, rH-H the interatomic distance of the two hydrogen atoms in the H2 molecule. Consequently the chemical potential of the hydrogen gas is given by

$$
\mu\_{\rm gas} = \mathbf{k} \cdot T \cdot \ln \frac{p}{p\_0(T)} - \mathbf{E}\_{\rm Diss} = \mathbf{k} \cdot \mathbf{T} \cdot \ln \frac{p}{\mathbf{P}\_0} + \mu\_{\rm gas\_0} \tag{5}
$$

with p0 = 1.01325 105 Pa.

In the solid solution (-phase) the chemical potential is accordingly

$$
\mu\_{\alpha} = \mathbf{h}\_{\alpha} - \mathbf{T} \mathbf{s}\_{\alpha} \quad \text{mit} \quad \mathbf{s}\_{\alpha} = \mathbf{s}\_{a\_{\text{conf}}} + \mathbf{s}\_{a\_{\text{vibr,loctr}}} \,. \tag{6}
$$

Here, *s*,conf is the configuration entropy which is originating in the possible allocations of NH hydrogen atoms on Nis different interstitial sites:

$$S\_{a, \text{conf}} = \mathbf{k} \cdot \ln \frac{\mathbf{N}\_{\text{is}}!}{\mathbf{N}\_{\text{H}}! (\mathbf{N}\_{\text{is}} \cdot \mathbf{N}\_{\text{H}})!} \tag{7}$$

and accordingly for small cH using the Stirling approximation we get

$$s\_{\alpha, \text{conf}} = \text{-k} \cdot \ln \frac{c\_{\text{H}}}{\text{n}\_{\text{is}} \cdot c\_{\text{H}}} \tag{8}$$

with nis being the number of interstitial sites per metal atom: nis = Nis/NMe and *c*H the number of hydrogen atoms per metal atom: *c*H = *N*H/NMe.

Therefore the chemical potential of hydrogen in the solid solution (-phase) is given by

$$
\mu\_{\alpha} = h\_{\alpha} - T \cdot \left( s\_{\mathbf{a}\_{\text{vhr,electr}}} - \mathbf{k} \cdot \ln \frac{\mathbf{c\_H}}{\mathbf{n}\_{\text{is}} - \mathbf{c\_H}} \right) \tag{9}
$$

Thermodynamics of Metal Hydrides: Tailoring Reaction Enthalpies of Hydrogen Storage Materials 897

ln 2 1

commonly is referred to as -phase.

as the pressure increases, see Fig. 5.

Hoff plot) as seen in Fig. 6.

solution and the hydride phase coexist.

0

p R R

The temperature dependent plateau pressure of this two phase field is the equilibrium dissociation pressure of the hydride and is a measure of the stability of the hydride, which

After complete conversion to the hydride phase, further dissolution of hydrogen takes place

Fig. 5. Schematic Pressure/Composition Isotherm. The precipitation of the hydride phase starts when the terminal solubility of the -phase is reached at the plateau pressure.

Multiple plateaus are possible and frequently observed in composite materials consisting of two hydride forming metals or alloys. The equilibrium dissociation pressure is one of the

If the logarithm of the plateau pressure is plotted vs 1/T, a straight line is obtained (van't

Fig. 6. Schematic pcT-diagram and van't Hoff plot. The -phase is the solid solution phase, the -phase the hydride phase. Within the two phase region both the metal-hydrogen

most important properties of a hydride storage material.

*T*

*S*

*<sup>p</sup> <sup>H</sup>* (16)

Taking into account the equilibrium condition (2) the hydrogen concentration *c*H can be determined via

$$\frac{c\_{\rm H}}{\mathbf{n}\_{\rm is} - c\_{\rm H}} = \sqrt{\frac{p}{p\_0(T)}} \cdot \mathbf{e}^{\frac{\Delta \mathbf{g}\_{\rm s}}{\mathbf{k} \cdot T}} \qquad \text{with} \quad \Delta \mathbf{g}\_{\rm s} = h\_{\rm a} - T \cdot s\_{\rm a\_{\rm vir}} + \frac{1}{2} \mu\_{\rm g\_0} \tag{10}$$

or

$$\frac{c\_{\rm H}}{\mathbf{m\_{is}} - c\_{\rm H}} = \sqrt{\frac{p}{p\_0(T)}} \cdot \mathbf{e}^{\frac{\Lambda \mathbf{G\_s}}{\mathbf{R} \cdot T}} \qquad \text{with} \qquad \Delta \mathbf{G\_s} = \Delta \mathbf{H\_s} - T \Delta \mathbf{S} \tag{11}$$

Here g0 is the chemical potential of the hydrogen molecule at standard conditions and R being the molar gas constant.

For very small hydrogen concentrations *c*H*c*H << nis in the solid solution phase the hydrogen concentration is directly proportional to the square root of the hydrogen pressure in the gas phase. This equation is also known as the **Sievert's law**, i.e.

$$\mathbf{c}\_{\rm H} = \frac{1}{\mathbf{K}\_{\rm S}} \sqrt{p} \tag{12}$$

with KS being a temperature dependent constant. As the hydrogen pressure is increased, saturation occurs and the metal hydride phase MeH*<sup>c</sup>* starts to form.

For higher hydrogen pressures/concentrations metal hydride formation occurs.

The conversion from the saturated solution phase to the hydride phase takes place at constant pressure p according to:

$$\text{Me-H}\_{\text{c}\_a}\big|\_{\text{a}} + \frac{1}{2}(\mathfrak{c}\_{\beta} - \mathfrak{c}\_a)\text{H}\_2 \leftrightarrow \text{MeH}\_{\text{c}\_{\beta}}\big|\_{\beta} + \text{Q}\_{a \to \beta} \text{ }. \tag{13}$$

In the equilibrium the chemical potentials of the gas phase, the solid solution phase and the hydride phase coincide:

$$
\mu\_a(p, T, \mathbf{c}\_a) = \mu\_\beta \left( p\_\prime T, \mathbf{c}\_\beta \right) \\
\quad = \frac{1}{2} \mu\_\text{gas} \left( p\_\prime T \right) = \frac{1}{2} \cdot \mathbf{k} \cdot T \cdot \ln \left( \frac{p\_{\text{eq}}(T)}{p\_0} \right) + \frac{1}{2} \mu\_\text{gas} \,, \tag{14}
$$

Following Gibb's Phase Rule *f*=*c*-*p*+2 with *f* being the degree of freedom, *k* being the number of components and *p* the number of different phases only one out of the four variables *p, T*, *c*, *c* is to be considered as independent. Therefore for a given temperature all the other variables are fixed.

Therefore the change in the chemical potential or the Gibbs free energy is just a function of one parameter, i.e. the temperature *T*:

$$
\Delta G = \frac{1}{2} \cdot \mathbb{R} \cdot T \cdot \ln\left(\frac{p(T)}{p\_0}\right) \tag{15}
$$

From this equation follows the frequently-used **Van't Hoff equation** (16):

Taking into account the equilibrium condition (2) the hydrogen concentration *c*H can be

vibr 0

*c p* (12)

eq

*p T*

0

. (15)

gas gas

. (13)

. (14)

*g*

s

*T*

is H 0

*c pT*

is H 0

n ()

*c pT*

in the gas phase. This equation is also known as the **Sievert's law**, i.e.

saturation occurs and the metal hydride phase MeH*<sup>c</sup>* starts to form.

 *pTc pTc pT T*

From this equation follows the frequently-used **Van't Hoff equation** (16):

<sup>g</sup> - <sup>H</sup> <sup>k</sup> <sup>s</sup> α α

*<sup>c</sup> <sup>p</sup> h Ts*

H

For higher hydrogen pressures/concentrations metal hydride formation occurs.

 

<sup>1</sup> e with g n () <sup>2</sup>

Gs - <sup>H</sup> <sup>R</sup> s s

*<sup>T</sup> <sup>c</sup> <sup>p</sup> <sup>T</sup>*

For very small hydrogen concentrations *c*H*c*H << nis in the solid solution phase the hydrogen concentration is directly proportional to the square root of the hydrogen pressure

> S 1 K

with KS being a temperature dependent constant. As the hydrogen pressure is increased,

The conversion from the saturated solution phase to the hydride phase takes place at

In the equilibrium the chemical potentials of the gas phase, the solid solution phase and

<sup>0</sup>

11 1 , , , , , k ln 2 2 p2

Following Gibb's Phase Rule *f*=*c*-*p*+2 with *f* being the degree of freedom, *k* being the number of components and *p* the number of different phases only one out of the four variables *p, T*, *c*, *c* is to be considered as independent. Therefore for a given temperature all

Therefore the change in the chemical potential or the Gibbs free energy is just a function of

<sup>1</sup> ( ) R ln 2 p *<sup>p</sup> <sup>T</sup> G T*

0

 

 <sup>α</sup> <sup>α</sup> <sup>2</sup> β αβ <sup>1</sup> Me-H H MeH Q 2 *c c c c*

g0 is the chemical potential of the hydrogen molecule at standard conditions and R

(10)

e with G H S

. (11)

determined via

or

Here 

being the molar gas constant.

constant pressure p according to:

the hydride phase coincide:

the other variables are fixed.

 

one parameter, i.e. the temperature *T*:

$$\frac{1}{2} \cdot \ln \frac{p}{\mathbf{p}\_0} = \frac{\Delta H}{\mathbf{R}T} - \frac{\Delta S}{\mathbf{R}} \tag{16}$$

The temperature dependent plateau pressure of this two phase field is the equilibrium dissociation pressure of the hydride and is a measure of the stability of the hydride, which commonly is referred to as -phase.

After complete conversion to the hydride phase, further dissolution of hydrogen takes place as the pressure increases, see Fig. 5.

Fig. 5. Schematic Pressure/Composition Isotherm. The precipitation of the hydride phase starts when the terminal solubility of the -phase is reached at the plateau pressure.

Multiple plateaus are possible and frequently observed in composite materials consisting of two hydride forming metals or alloys. The equilibrium dissociation pressure is one of the most important properties of a hydride storage material.

If the logarithm of the plateau pressure is plotted vs 1/T, a straight line is obtained (van't Hoff plot) as seen in Fig. 6.

Fig. 6. Schematic pcT-diagram and van't Hoff plot. The -phase is the solid solution phase, the -phase the hydride phase. Within the two phase region both the metal-hydrogen solution and the hydride phase coexist.

Thermodynamics of Metal Hydrides: Tailoring Reaction Enthalpies of Hydrogen Storage Materials 899

58 mol H2/l or 116 g/l and has to be compared with the hydrogen density in liquid hydrogen (20 K): 4.2 1022 (35 mol H2/l or 70 g/l) and in compressed hydrogen (350 bar / 700 bar): 1.3 / 2.3 1022 atoms/cm3 ( 11 mol H2/l or 21 g/l and 19 mol H2/l or 38 g/l respectively) . The hydrogen density varies a lot between different hydrides. VH2 for example has an even higher hydrogen density which amounts to 11.4 1022 hydrogen atoms per cubic centimetre and accordingly 95 mol H2/l or 190 g/l. As in the case of many other transition metal hydrides Zr has a number of different hydride phases ZrH2-x with a wide variation in the stoichiometry (Hägg, 1931). Their compositions extend from about ZrH1.33 up to the saturated hydride ZrH2. Because of the limited gravimetric storage density of only about 2 wt.% and the negligibly low plateau pressure within the temperature range of 0 – 150 °C Zr as well as Ti and Hf are not suitable at all as a reversible hydrogen storage material. Thus, they are not useful for reversible hydrogen storage if only the pure binary hydrides are considered (Dornheim & Klassen, 2009). Libowitz et al. (Libowitz et al., 1958) could achieve a breakthrough in the development of hydrogen storage materials by discovering the class of reversible intermetallic hydrides. In 1958 they discovered that the intermetallic compound ZrNi reacts reversibly with gaseous hydrogen to form the ternary hydride ZrNiH3. This hydride has a thermodynamic stability which is just in between the stable high temperature hydride ZrH2 (fH0= -169 kJ/mol H2) and the rather unstable NiH (fH0= -8.8 kJmol-1H2). Thus, the intermetallic Zr-Ni bond exerts a strong destabilizing effect on the Zr-hydrogen bond so that at 300°C a plateau pressure of 1bar is achieved which has to be compared to 900°C in case of the pure binary hydride ZrH2. This opened up a completely new research field. In the following years hundreds of new storage materials with different thermodynamic properties were discovered which generally follow

the well-known semi-empirical rule of Miedema (Van Mal et al., 1974):

for NiH are Hf,NiH = 8.8 kJmol-1H2 and Pdiss,NiH,RT=3400 bar.

Ca, La, Rare

AB5

(A B H ) (A H ) (B H ) (A B ) *H HHH nm x<sup>y</sup> nx m <sup>y</sup> n m* (17)

Around 1970, hydrides with significantly lowered values of hydrogen reaction enthalpies, such as LaNi5 and FeTi but also Mg2Ni were discovered. While 1300 C are necessary to reach a desorption pressure of 2 bar in case of the pure high temperature hydride LaH2, in case of LaNi5H6 a plateau pressure of 2 bar is already reached at 20 C only. The value of the hydrogen reaction enthalpy is lowered to HLaNi5H6 = 30.9 kJmol-1H2. The respective values

In the meantime, several hundred other intermetallic hydrides have been reported and a number of interesting compositional types identified (table 1). Generally, they consist of a high temperature hydride forming element A and a non hydride forming element B, see fig. 8.

A2B Mg, Zr Ni, Fe, Co Mg2Ni, Mg2Co, Zr2Fe AB Ti, Zr Ni, Fe TiNi, TiFe, ZrNi

AB3 La, Y, Mg Ni, Co LaCo3,YNi3,LaMg2Ni9

Table 1. Examples of intermetallic hydrides, taken from Dornheim et al. (Dornheim, 2010).

AB2 Zr, Ti, Y, La V, Cr, Mn, Fe, Ni LaNi2, YNi2,YMn2, ZrCr2, ZrMn2,ZrV2,

Earth Ni, Cu, Co, Pt, Fe CaNi5, LaNi5, CeNi5, LaCu5, LaPt5,

TiMn2

LaFe5

COMPOSITION A B COMPOUNDS

## **2.1 Conventional metal hydrides**

Fig. 7 shows the Van't Hoff plots of some selected binary hydrides. The formation enthalpy of these hydrides H0 f determines the amount of heat which is released during hydrogen absorption and consequently is to be supplied again in case of desorption. To keep the heat management system simple and to reach highest possible energy efficiencies it is necessary to store the heat of absorption or to get by the waste heat of the accompanying hydrogen utilizing process, e.g. energy conversion by fuel cell or internal combustion system. Therefore the reaction enthalpy has to be as low as possible. The enthalpy and entropy of the hydrides determine the working temperatures and the respective plateau pressures of the storage materials. For most applications, especially for mobile applications, working temperatures below 100°C or at least below 150°C are favoured. To minimize safety risks and avoid the use of high pressure composite tanks the favourable working pressures should be between 1 and 100 bar.

Fig. 7. Van't Hoff lines (desorption) for binary hydrides. Box indicates 1-100 atm, 0-100 °C ranges, taken from Sandrock et al. (Sandrock, 1999).

However, the Van't Hoff plots shown in Fig. 7 indicate that most binary hydrides do not have the desired thermodynamic properties. Most of them have rather high thermodynamic stabilities and thus release hydrogen at the minimum required pressure of 1 bar only at rather high temperatures (T>300°C). The values of their respective reaction enthalpies are in the range of 75 kJ/(mol H2) (MgH2) or even higher. Typical examples are the hydrides of alkaline metals, alkaline earth metals, rare earth metals as well as transition metals of the Sc-, Ti- and V-group. The strongly electropositive alkaline metals like LiH and NaH and CaH2 form saline hydrides, i.e. they have ionic bonds with hydrogen. MgH2 marks the transition between these predominantly ionic hydrides and the covalent hydrides of the other elements in the first two periods.

Examples for high temperature hydrides releasing the hydrogen at pressures of 1 bar at extremely high temperatures (T > 700°C) are ZrH2 and LaH2 (Dornheim & Klassen, 2009). ZrH2 for example is characterized by a high volumetric storage density NH. NH values larger than 7 1022 hydrogen atoms per cubic centimetre are achievable. This value corresponds to

Fig. 7 shows the Van't Hoff plots of some selected binary hydrides. The formation enthalpy

absorption and consequently is to be supplied again in case of desorption. To keep the heat management system simple and to reach highest possible energy efficiencies it is necessary to store the heat of absorption or to get by the waste heat of the accompanying hydrogen utilizing process, e.g. energy conversion by fuel cell or internal combustion system. Therefore the reaction enthalpy has to be as low as possible. The enthalpy and entropy of the hydrides determine the working temperatures and the respective plateau pressures of the storage materials. For most applications, especially for mobile applications, working temperatures below 100°C or at least below 150°C are favoured. To minimize safety risks and avoid the use of high pressure composite tanks the favourable working pressures

Fig. 7. Van't Hoff lines (desorption) for binary hydrides. Box indicates 1-100 atm, 0-100 °C

However, the Van't Hoff plots shown in Fig. 7 indicate that most binary hydrides do not have the desired thermodynamic properties. Most of them have rather high thermodynamic stabilities and thus release hydrogen at the minimum required pressure of 1 bar only at rather high temperatures (T>300°C). The values of their respective reaction enthalpies are in the range of 75 kJ/(mol H2) (MgH2) or even higher. Typical examples are the hydrides of alkaline metals, alkaline earth metals, rare earth metals as well as transition metals of the Sc-, Ti- and V-group. The strongly electropositive alkaline metals like LiH and NaH and CaH2 form saline hydrides, i.e. they have ionic bonds with hydrogen. MgH2 marks the transition between these predominantly ionic hydrides and the covalent hydrides of the

Examples for high temperature hydrides releasing the hydrogen at pressures of 1 bar at extremely high temperatures (T > 700°C) are ZrH2 and LaH2 (Dornheim & Klassen, 2009). ZrH2 for example is characterized by a high volumetric storage density NH. NH values larger than 7 1022 hydrogen atoms per cubic centimetre are achievable. This value corresponds to

ranges, taken from Sandrock et al. (Sandrock, 1999).

other elements in the first two periods.

f determines the amount of heat which is released during hydrogen

**2.1 Conventional metal hydrides** 

should be between 1 and 100 bar.

of these hydrides H0

58 mol H2/l or 116 g/l and has to be compared with the hydrogen density in liquid hydrogen (20 K): 4.2 1022 (35 mol H2/l or 70 g/l) and in compressed hydrogen (350 bar / 700 bar): 1.3 / 2.3 1022 atoms/cm3 ( 11 mol H2/l or 21 g/l and 19 mol H2/l or 38 g/l respectively) . The hydrogen density varies a lot between different hydrides. VH2 for example has an even higher hydrogen density which amounts to 11.4 1022 hydrogen atoms per cubic centimetre and accordingly 95 mol H2/l or 190 g/l. As in the case of many other transition metal hydrides Zr has a number of different hydride phases ZrH2-x with a wide variation in the stoichiometry (Hägg, 1931). Their compositions extend from about ZrH1.33 up to the saturated hydride ZrH2. Because of the limited gravimetric storage density of only about 2 wt.% and the negligibly low plateau pressure within the temperature range of 0 – 150 °C Zr as well as Ti and Hf are not suitable at all as a reversible hydrogen storage material. Thus, they are not useful for reversible hydrogen storage if only the pure binary hydrides are considered (Dornheim & Klassen, 2009). Libowitz et al. (Libowitz et al., 1958) could achieve a breakthrough in the development of hydrogen storage materials by discovering the class of reversible intermetallic hydrides. In 1958 they discovered that the intermetallic compound ZrNi reacts reversibly with gaseous hydrogen to form the ternary hydride ZrNiH3. This hydride has a thermodynamic stability which is just in between the stable high temperature hydride ZrH2 (fH0= -169 kJ/mol H2) and the rather unstable NiH (fH0= -8.8 kJmol-1H2). Thus, the intermetallic Zr-Ni bond exerts a strong destabilizing effect on the Zr-hydrogen bond so that at 300°C a plateau pressure of 1bar is achieved which has to be compared to 900°C in case of the pure binary hydride ZrH2. This opened up a completely new research field. In the following years hundreds of new storage materials with different thermodynamic properties were discovered which generally follow the well-known semi-empirical rule of Miedema (Van Mal et al., 1974):

$$
\Delta \mathbf{H} (\mathbf{A}\_n \mathbf{B}\_m \mathbf{H}\_{x+y}) = \Delta \mathbf{H} (\mathbf{A}\_n \mathbf{H}\_x) + \Delta \mathbf{H} (\mathbf{B}\_m \mathbf{H}\_y) - \Delta \mathbf{H} (\mathbf{A}\_n \mathbf{B}\_m) \tag{17}
$$

Around 1970, hydrides with significantly lowered values of hydrogen reaction enthalpies, such as LaNi5 and FeTi but also Mg2Ni were discovered. While 1300 C are necessary to reach a desorption pressure of 2 bar in case of the pure high temperature hydride LaH2, in case of LaNi5H6 a plateau pressure of 2 bar is already reached at 20 C only. The value of the hydrogen reaction enthalpy is lowered to HLaNi5H6 = 30.9 kJmol-1H2. The respective values for NiH are Hf,NiH = 8.8 kJmol-1H2 and Pdiss,NiH,RT=3400 bar.

In the meantime, several hundred other intermetallic hydrides have been reported and a number of interesting compositional types identified (table 1). Generally, they consist of a high temperature hydride forming element A and a non hydride forming element B, see fig. 8.


Table 1. Examples of intermetallic hydrides, taken from Dornheim et al. (Dornheim, 2010).

Thermodynamics of Metal Hydrides: Tailoring Reaction Enthalpies of Hydrogen Storage Materials 901

Novel light weight hydrides show much higher gravimetric storage capacities than the conventional room temperature metal hydrides. However, currently only a very limited number of materials show satisfying sorption kinetics and cycling behaviour. The most prominent ones are magnesium hydride (MgH2) and sodium alanate (NaAlH4). In both cases a breakthrough in kinetics could be attained in the late 90s of the last century / the

Magnesium hydride is among the most important and most comprehensively investigated light weight hydrides. MgH2 itself has a high reversible storage capacity, which amounts to 7.6 wt.%. Furthermore, magnesium is the eighth most frequent element on the earth and thus comparably inexpensive. Its potential usage initially was hindered because of rather sluggish sorption properties and unfavourable reaction enthalpies. The overall hydrogen sorption kinetics of magnesium-based hydrides is as in case of all hydrides mainly determined by the slowest step in the reaction chain, which can often be deduced e.g. by modelling the sorption kinetics (Barkhordarian et al, 2006; Dornheim et al., 2006). Different measures can be taken to accelerate kinetics. One important factor for the sorption kinetics is the micro- or nanostructure of the material, e.g. the grain or crystallite size. Because of the lower packing density of the atoms, diffusion along grain boundaries is usually faster than through the lattice. Furthermore, grain boundaries are favourable nucleation sites for the formation and decomposition of the hydride phase. A second important parameter is the outer dimension of the material, e.g. in case of powdered material, its particle size. The particle size (a) determines the surface area, which is proportional to the rate of the surface reaction with the hydrogen, and (b) is related to the length of the diffusion path of the hydrogen into and out of the volume of the material. A third major factor by which hydrogen sorption is improved in many hydrogen absorbing systems is the use of suitable additives or catalysts. In case of MgH2 it was shown by Oelerich et al. (Oelerich et al., 2001; Dornheim et al., 2007) that already tiny amounts of transition metal oxides have a huge impact on the kinetics of hydrogen sorption. Using such additives Hanada et al. (Hanada et al., 2007) could show that by using such additives hydrogen uptake in Mg is possible already at room temperature within less than 1 min. The additives often do not just have one single function but multiple functions. Suitable additives can catalyze the surface reaction between solid and gas. Dispersions in the magnesium-based matrix can act as nucleation centres for the hydride or the dehydrogenated phase. Furthermore, different additives, such as liquid milling agents and hard particles like oxides, borides, etc. , can positively influence the particle size evolution during the milling process (Pranzas et al., 2006; Pranzas et al., 2007; Dornheim et al, 2007) and prevent grain i.e. crystallite growth. More detailed information about the function of such additives in MgH2 is given in (Dornheim et al., 2007). Beyond that, a preparation technique like high-energy ball milling affects both the evolution of certain particle sizes as well as very fine crystallite sizes in the nm range and is also used to intermix the hydride and the additives/catalysts. Thus, good interfacial contact with the light metal hydride as well as a fine dispersion of the additives

As in case of MgH2 dopants play also an important role in the sorption of Na-Al-hydride, the so-called Na-alanate. While hydrogen liberation is thermodynamically favorable at moderate temperatures, hydrogen uptake had not been possible until in 1997 Bogdanovic et al. demonstrated that mixing of NaAlH4 with a Ti-based catalyst leads to a material, which can be reversibly charged with hydrogen (Bogdanovic, 1997). By using a tube vibration mill

**2.2 Hydrogen storage in light weight hydrides** 

early 21st century.

can be achieved.

Fig. 8. Hydride and non hydride forming elements in the periodic system of elements.

Even better agreement with experimental results than by use of Miedema's rule of reversed stability is obtained by applying the semi-empirical band structure model of Griessen and Driessen (Griessen & Driessen, 1984) which was shown to be applicable to binary and ternary hydrides. They found a linear relationship of the heat of formation H = H0f of a metal hydride and a characteristic energy E of the electronic band structure of the host metal which can be applied to simple metals, noble metals, transition metals, actinides and rare earths:

$$
\Delta H = \mathbf{u} \cdot \Delta E + \oint \tag{18}
$$

with *E* = *E*F-*E*S (*E*F being the Fermi energy and *E*S the center of the lowest band of the host metal, = 59.24 kJ (eV mol H2)-1 and = -270 kJ (mol H2)-1 and *E* in eV.

As described above, most materials experience an expansion during hydrogen absorption, wherefore structural effects in interstitial metal hydrides play an important role as well. This can be and is taken as another guideline to tailor the thermodynamic properties of interstitial metal hydrides. Among others Pourarian et al. (Pourarian, 1982), Fujitani et al. (Fujitani, 1991) and Yoshida & Akiba (Yoshida, 1995) report about this relationship of lattice parameter or unit cell volume and the respective plateau pressures in different material classes.

Intensive studies let to the discovery of a huge number of different multinary hydrides with a large variety of different reaction enthalpies and accordingly working temperatures. They are not only attractive for hydrogen storage but also for rechargeable metal hydride electrodes and are produced and sold in more than a billion metal hydride batteries per year. Because of the high volumetric density, intermetallic hydrides are utilized as hydrogen storage materials in advanced fuel cell driven submarines, prototype passenger ships, forklifts and hydrogen automobiles as well as auxiliary power units.
