Destabilization of lithium hydride and the thermodynamic assessment of the Li–Al–H system for solar thermal energy storage

Abstract

Lithium hydride destabilised with aluminium, LiH–Al (1 : 1 mole ratio) was systematically studied and its suitability as a thermal energy storage system in Concentrating Solar Power (CSP) applications was assessed. Pressure composition isotherms (PCI) measured between 506 °C and 652 °C were conducted to investigate the thermodynamics of H₂ release. Above the peritectic temperature (596 °C) of LiAl, PCI measurements were not consistently reproducible, possibly due to the presence of a molten phase. However, below 596 °C, the hydrogen desorption enthalpy and entropy of LiH–Al was ΔH_des = 96.8 kJ (mol H₂)⁻¹ and ΔS_des = 114.3 J (K mol H₂)⁻¹, respectively LiH_(s) at 956 °C, ΔH_des = 133.0 kJ (mol H₂)⁻¹ and ΔS_des = 110.0 J (K mol H₂)⁻¹. Compared to pure LiH, the Li–Al–H system has a reduced operating temperature (1 bar H₂ pressure at T ∼ 574 °C) that, combined with favourable attributes such as high reversibility, good kinetics and negligible hysteresis, makes the Li–Al–H system a potential candidate for solar thermal energy storage applications. Compared to pure LiH, the addition of Al can reduce the cost of the raw materials by up to 44%. This cost reduction is insufficient for next generation CSP but highlights the potential to improve the properties and cost of high temperature hydrides via destabilisation.

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Introduction

The essential properties of hydrogen powered devices are the requirement of high capacity hydrogen storage materials and reversible hydrogen sorption reactions. Particularly, complex hydrides of alkali metal (Li, Na, K) combined with Al have attracted significant attention due to the high hydrogen storage capacity of 5.5–7.4 wt%,^1–3 and the possibility of a reversible reaction by the addition of a catalyst such as Ti.^4,5 This paper describes the reaction between Li–Al with H₂ at elevated temperatures. The high thermal stability of the Li–Al–H system (i.e., LiH + Al) makes it a potential thermal energy storage material for next-generation concentrating solar power (CSP) applications that are expected to operate in the 600 to 800 °C temperature range.^6,7 To the best of the authors knowledge, the thermodynamic investigation of the Li–Al–H system has previously only been studied experimentally by Veleckis in 1980.⁸

There are three main methods for storing thermal energy; sensible, latent and chemical heat storage. Sensible heat is the energy released (absorbed) by a material as its temperature is reduced (increased). Sensible heat storage media can either be a solid (mainly high temperature concrete or castable ceramics) or a liquid (molten salts, mineral oils and synthetic oils). The simplest is comprised of binary nitrate liquid molten salt mixtures (60% NaNO₃; 40% KNO₃) which have potential corrosion problems, not to mention the larger volume required to store sufficient heat to operate the plant for several hours during insufficient solar radiation hours (these nitrate salt mixtures have a maximum useable temperature of ∼590 to 600 °C as they begin to decompose above this temperature).^6,9 The second form of thermal storage uses “latent” heat, which is associated with the phase change of materials (PCM) at isothermal conditions e.g. heat of phase change such as heat of vaporization (liquid–vapour transition) or heat of fusion (solid–liquid transition). The issue with PCM is sluggish heat transfer and low charge and discharging rates.¹⁰ The third storage mechanism is assigned to chemical reactions i.e. chemical heat storage (thermochemical energy storage). This type of heat storage relies on a completely reversible chemical reaction. Principally, heat from solar radiation received, is used to excite an endothermic chemical reaction and the necessary heat is available whenever desired. The advantages of this storage type are, for instance, the high energy storage densities and long storage durations at near ambient temperatures i.e. the heat storage capacity is 150 kJ kg⁻¹ (290 < T < 600 °C) for molten salt mixtures, 200–500 kJ kg⁻¹ (300 < T < 800 °C) for latent heat compounds and 1160–8400 kJ kg⁻¹ (250 < T < 1000 °C) for metal hydrides depending on hydride composite system.^9,11

The cheap and abundant metal hydrides MgH₂ and Mg₂FeH₆ (ref. 12 and 13) have been previously considered for thermal energy storage at temperatures above ∼350 °C. The drawback with MgH₂ is the narrow operating temperatures ranging between 400 to 480 °C due to its high equilibrium pressure, which means that the energy efficiency involved between conversions of heat to electricity is too low.^14,15

More recent work suggests substitution of fluorine in NaMgH₃ forming NaMgH₂F which enhances the stability relative to pure NaMgH₃. Furthermore, cost assessment based on NaMgH₂F suggests the metal hydride system to be suitable for concentrated solar thermal storage.^16–18

The advantages of LiH for concentrated solar thermal storage applications is the high hydrogen content of ρ_m = 12.7 wt% and high theoretical heat storage capacity of 8397 kJ kg⁻¹. However LiH suffers from a high operating temperature of above 850 °C, a high temperature to reach a hydrogen equilibrium pressure of 1 bar, denoted T(1 bar) ∼ 956 °C,¹⁹ and a relatively high cost of the raw material.

The advantages of adding another element, such as Al, are destabilization of the system, potentially reducing the cost and decreasing the 1 bar H₂ equilibrium temperature in comparison to pure LiH. Moreover, it is important to understand the thermodynamics and to evaluate beneficial reactions involved in the Li–Al–H system. In addition LiH + Al are often end products of complex metal hydrides that contain Li and Al (e.g. LiAlH₄ and LiBH₄ + Al^20–22). Of note, the destabilization of LiH has previously been attempted with the addition of various elements such as Ge, Si and Sn.^23–26 A similar alternative system such as Li–Mg–H system has been investigated for energy storage.²⁷ However, the drawback of this system, according to the Mg–Li phase diagram,²⁸ is that no line compounds form between Li–Mg and this limits the ability of the system to produce a hydrogen desorption/absorption plateau at near constant pressure.

This study investigates the thermodynamic and kinetic properties for hydrogen desorption of the Li–Al–H system using pressure composition isotherm (PCI) data and kinetic measurements. The potential application of this system for thermochemical heat storage in CSP is also assessed.

Experimental details

The samples were prepared from commercially available LiH (Aldrich, 95%) and commercially available Al flakes (Aldrich, 99.5%) in the molar ratio 0.51LiH–0.49Al (resulting in an approximate Li : Al ratio of 1 : 1 once purity levels of the reagents were taken into account). We note here that we use the terms LiAl and Li_0.5Al_0.5 interchangeably throughout as the former is usually used in relation to the phase and phase diagram while the latter is more convenient in manipulating chemical formulae. The hydride mixture was ball milled using a stainless steel bowl (80 mL) and balls (Ø: 10 mm), with a planetary mill and a powder to ball ratio of 1 : 24. The milling was performed at 380 rpm with a total milling time of 2 h uninterrupted.

The pressure-composition isotherms and kinetic data were collected using a Sieverts type apparatus (PCTpro 2000 E&E). The samples were sealed in a specially treated (see further details below) 316L stainless steel autoclave under argon and attached to the Sieverts' apparatus. Hydrogen desorption and absorption PCI's were collected at the temperatures 506, 540, 560, 573, 585, 608, 619, 634 and 652 °C, using the same sample. Desorption and absorption measurements were conducted between 0 to 15 bar and the time to reach equilibrium varied but were typically around 2 or 3 hours for each PCI equilibrium data point.

Hydrogen permeability through stainless steel increases exponentially with increasing temperature. The stainless steel autoclave was specially treated via an aluminium dip coating process that was followed by oxidation to produce a coating of Al₂O₃.²⁹ The purpose of the Al₂O₃ coating is to reduce the diffusion of H₂ but, due to technical limitations, the coating only covered the inner and outer surface of the 1/2′′ diameter stainless steel tube and not the 1/2′′ end cap that seals the tube. Since H₂ permeated through the endcap of the autoclave, all of the collected data was corrected for the hydrogen loss by calculating the amount of hydrogen leaked per desorption step using the permeability of steel at the measured temperatures and pressures.³⁰

Results and discussion

According to literature and the phase diagram of Li–Al, the system is known to have four homogenous solid solution (ss) phases; α: which is a dilute solution of Li in Al, β: “LiAl”, γ: “Li₃Al₂” and δ: “Li₉Al₄”.³¹ As the focus of our study is to investigate the thermodynamic properties of the α and β coexisting phases i.e. plateau pressure region, the γ and δ line compounds are irrelevant to this study as they are unlikely to ever be present based on the selected starting ratios of Li and Al.³² The phase diagram of Li and Al suggests that LiH should react with Al to form a solid solution of Li in Al up to a temperature-dependent composition limit according to reaction (1) below the peritectic line at 596 °C.

xLiH + (1 − x)Al → Al_1−xLi_x + x/2H₂, where 0 < x < limit of Li solubility in Al (1)

Provided that the system is maintained below the peritectic temperature (596 °C), this solid solution of Li in Al can then further react with LiH at constant hydrogen pressure to yield an equilibrium plateau described by reaction (2).

2LiH + 3.273Al_0.87Li_0.13 → 5.273Al_0.54Li_0.46 + H₂, 2.11 wt% H₂ (2)

A generic equation reaction describing the reaction process after the plateau that depends on the composition, is described as the following:

yLiH + (1 − y)Al_{y_min}Li_{y_min} → Al_{(1−y)y_min}Li_{y(1−y_min+y_min y⁻¹)} + y/2H₂, where y_min < y < y_max (3)

Above 596 °C, a peritectic reaction between the Al–Li solid solution (Al, Li)_ss and LiH takes place. According to the phase diagram of Li–Al, two co-existing phases are prevailing; one with Al and a liquid phase, and the other with a solid solution of LiAl and the liquid phase. Two reactions are taking place across these regions, which can be described by reaction (4) and (5) at 627 °C.

2LiH + 20.25Al_0.89Li_0.11 → 22.25Al_0.81Li_0.19 + H₂, 0.4 wt% H₂ (4)

2LiH + 6.75Al_0.70Li_0.30 → 8.75Al_0.54Li_0.46 + H₂, 1.28 wt% H₂ (5)

A complete set of H₂ sorption pressure-composition isotherms (PCIs) have been conducted at the following temperatures; 506.3, 540.0, 559.9, 573.3, 584.8, 608.6, 618.9, 634.2 and 652.4 °C. A sufficient amount of dwell time is crucial during data collection in order to reach true equilibrium, as prematurely halting time can lead to erroneous equilibrium pressure, artificial hysteresis and the incorrect calculation of thermodynamic quantities.³³ Often the effect of severe kinetic limitation is preventing true equilibrium to be reached and if true equilibrium is not met, PCI curves often display large amounts of hysteresis between absorption and desorption isotherms. Furthermore, measured absorption pressures will be higher than actual true equilibrium and measured desorption pressures will be lower than actual true equilibrium, resulting in apparent hysteresis. Insufficient measurement time will also lead to false plateau and equilibrium pressure, and hence incorrect thermodynamic calculations. Furthermore, the equilibrium plateau can also by shortened due to short measurement times at low temperatures, indicating equilibrium has not been reached. These features can be due to kinetic limitations, especially in regions before and after the plateau region. Increasing the temperature of measurement leads to improved kinetics but it must be emphasised that kinetic data should be collected during PCI measurements to ensure sufficient time for equilibrium to occur. The kinetics of decomposition in the Li–Al–H system are relatively fast, thus equilibrium can be reached within 2 hours.

The PCI desorption data below 596 °C, i.e. reaction (2) are displayed in Fig. 1(a). The plateau curves are very flat, with less than a 0.2 bar pressure variation between the start and the end of the plateau, and a total of ∼2.3 wt% H₂ is desorbed along the plateau at all temperatures below 596 °C. In accordance with the phase diagram,³⁴ the amount of H₂ desorbed along the plateau gradually decreases as the temperature increases due to the increased solubility of Li in the (Al, Li)_ss phase.³³ The total hydrogen desorption capacity ranges between 2.5–2.6 wt%. Since the composition of the starting material in the dehydrogenated state is known to have the composition Al_0.54Li_0.46 at the end of the plateau,²⁸ the theoretical hydrogen storage capacity of a 1 : 1 molar mixture of LiH and Al is calculated to be 2.89 wt%. Thus at 506 °C the amount of desorbed hydrogen corresponds to 90% of the available theoretical content, whereas at 584.8 °C the amount of desorbed hydrogen corresponds to 87%. It is worth noting that there is a very slight slope evident in the plateaux and may be due to an impurity in the starting reagents i.e. LiOH was observed by XRD to be present in the as-received LiH.

Fig. 1 (a) Hydrogen desorption pressure – composition isotherms for reaction (1) performed at various temperatures below 596 °C. (b) Kinetic H₂ desorption data of reaction (1) performed at 573.3 °C. (c) van't Hoff plot of H₂ desorption equilibrium pressures and the linear fit to the data.

Fig. 1(b) displays the kinetics of H₂ during each desorption step measured at 573 °C and corresponds to the isotherm presented in Fig. 1(a). Each end point of the kinetic measurement is associated with a pressure point on the isotherm i.e. point A in Fig. 1(a) and (b) corresponding to the same value point.

Dehydrogenation isotherms collected at 506.3, 540.0, 559.9, 573.3 and 584.8 °C, were used to construct a van't Hoff plot, Fig. 1(c), using the equilibrium pressure values taken from the approximate midpoint of each plateau at the same hydrogen composition (−1.354 wt% H₂). The desorption enthalpy (ΔH_des) and entropy (ΔS_des) were determined to be 96.8 kJ mol⁻¹ H₂ and 114.3 J K⁻¹ mol⁻¹ H₂, respectively. These values are comparable to those obtained by Veleckis,⁸ ΔH_abs = −98.2 kJ mol⁻¹ H₂ and ΔS_abs = −117.2 J K⁻¹ mol⁻¹ H₂, derived from constructing a van't Hoff plot from their reported hydrogen equilibrium pressures.

It is worth noting that in the PCI desorption data below 596 °C a dip occurs at the beginning of the plateau (at H₂ wt% of between −0.2 and −0.3). This phenomenon appears to be similar to what was previously observed for the Ti–Fe–H³⁵ and U–H system.^36–38 However the desorption PCI's of the LiH–Al system measured above 596 °C do not display the abovementioned dip, indicating that at higher temperatures, the dip becomes less pronounced and at a certain temperature regime (possibly around ∼600 °C) it disappears completely. This observation is also in good agreement with previous work,³⁵ where it disappeared above a critical temperature and was suggested to be due to the supersaturation of hydrogen vacancies in the hydride phase.^35,37 Whether this is the case for the LiH–Al system requires further verification.

PCI desorption isotherms at selected temperatures (608.6, 618.9, 634.2 and 652.4 °C) above the peritectic temperature are shown in ESI Fig. S2(a).† The PCI isotherms measured at 608.6, 618.9 and 634.2 °C indicate clear equilibrium plateaux with a minor slope but we note here that is was difficult to obtain reproducible results on different samples above 596 °C. Measurements (652.4 °C) closer to the melting point of Al (660 °C) resulted in significant changes to the PCI curves with a drastic shortening of the equilibrium plateau. We note that two plateaux would be expected for the PCI measurements conducted at temperatures above the 596 °C but that due to the factors discussed, thermodynamic calculations were not possible and the exact decomposition process could not be determined (see Fig. S2(b)†).

One of the crucial properties of a metal hydride for thermal energy storage is its kinetics, hysteresis and reversibility of hydrogen release and uptake. Fig. 2(A) shows both the absorption and desorption PCT measurements performed below the peritectic temperature at 584.8 °C while Fig. 2(B) and (C) show the kinetic curves for the desorption and absorption, respectively. Fig. 2(A) shows that there is negligible hysteresis in the plateau region but that there is some minor hysteresis before and after the plateau where only the solid-solution α and β phases exist. Examination of the kinetic curves, Fig. 2(B) and (C), reveals that this hysteresis is actually an artefact of slower kinetics and thus a longer measurement time is required in these regions. In contrast, PCT absorption and desorption measurements in the plateau region reach equilibrium within 1 hour. An equivalent result was obtained for PCT measurements measured at 618 °C, above the peritectic point (see Fig. S1(A)–(C)†). In PCT measurements, the hydrogen absorption/desorption kinetics are a complex interplay between the PCT step size (large aliquot effect), the thermodynamic driving force, the hydride enthalpy of formation/decomposition, the sample size and the thermal conductivity of the sample.³³ Of these, the dominant limitation is the low thermal conductivity of the hydride powder, typically less than 1–2 W m⁻¹ K⁻¹. This means that heat accumulation and heat flow become the rate-limiting steps³⁹ and this is only exacerbated in high-temperature hydrides with large negative heats of formation. As a result, the direct quantitative comparison of sorption kinetics between different hydrides is challenging. However, a single-step hydrogen absorption measurement performed at 618 °C (not shown) using an applied hydrogen pressure of 20 bar (versus P(H₂)_eq at 618 °C of ∼2 bar) facilitates a hydrogen overpressure equivalent to an excess thermodynamic driving force of ΔG_exc, of −17.1 kJ mol⁻¹ H₂ (ΔG_abs = ΔG_eq + ΔG_exc = ΔG_eq + RT ln(P_eq/P_abs), ΔG = Gibbs free energy) and resulted in complete absorption within 4 hours. This easily meets the U.S. Department of Energy (U.S. D.O.E.) SunShot charge/discharge target of 6 hours (ref. 40) for a CSP thermal energy storage system and suitable engineering to improve the thermal conductivity of the hydride bed would only further enhance the kinetics. Though highly suggestive, such a result would need to be confirmed using single-step absorption measurements using a lower hydrogen overpressure. The negligible hysteresis between the absorption–desorption isotherms in conjunction with kinetics that easily meet U.S. D.O.E. targets makes Li–Al–H a potential high temperature solar thermal heat storage medium.

Fig. 2 (A) Desorption–absorption PCI at 585 °C. The respective desorption kinetics (B) and absorption kinetics (C) are also displayed. * is due to a sudden temperature fluctuation in the furnace set-point.

In order to confirm the hydrogen sorption reactions, X-ray diffraction was performed after both desorption and absorption at 618 °C. XRD performed after desorption (Fig. S2(A) in ESI†) reveals that the main product is LiAl, as expected, with a trace amount of residual LiH. Likewise, after hydrogen absorption at 618 °C, XRD (Fig. S2(B) in ESI†) reveals the main products to be LiH and Al, along with a small amount of LiAlO₂ that results from slight oxidation of the sample. It should be noted that trace amounts of an unknown phase were identified in the sample after both desorption and re-absorption.

Metal hydrides have the potential to be the next generation of heat storage materials to replace molten salts in CSP.^6,11 As such, the cost of the metal hydride is of crucial importance. The cost comparison of pure LiH compared to LiH–Al, based on both theoretical and practical capacities, is given in Table 1. The addition of Al improves the operating temperature by decreasing the 1 bar H₂ equilibrium temperature by ∼300 °C and decreases the enthalpy of formation. While the consequence of the additive is a decrease in the practical heat storage capacity by approximately 83%, the cost per kW h of thermal energy stored is reduced by ∼44%. For thermal energy storage, it has been proposed that the high-temperature metal hydride would operate at a near constant temperature. However, additional heat could be stored in the Li–Al–H system if its temperature is allowed to fluctuate to take advantage of the heat of fusion of LiAl at 702 °C. The thermodynamic cycle would include (Fig. 3):

Table 1 A comparison of the properties of pure LiH compared to the LiH + Al system

		H2 capacity (wt%)	ΔHdes (kJ mol^−1 H2)	T(1 bar) in °C	Heat storage capacity (kJ kg^−1)	Cost^a^,^b US$ per kg of metal	US$ per kW hth
a Li raw material cost of US$70.00 per kg taken from ref. 17. The cost of hydrogen was not taken into account.b Al raw material cost of US$1.48 per kg taken from ref. 41. The cost of hydrogen was not taken into account.c Assuming the reaction is 2LiH + 3.40Li0.142Al0.858 → 5.40Li0.46Al0.54 + H2.d Using the reaction scheme from Fig. 3.
LiH	Theor.	12.68	133.5	956	8397.3	61.12	26.23
	Pract.	7.61	133.5	956	4198.6	61.12	52.46
LiH + Al	Theor.	2.89	96.8	573	1412.5	13.51	35.12
	Pract.	2.06^c	96.8	573	1150.3	13.51	49.20
	Pract.^c^,^d	2.06^c	96.8	573	1888.1	13.5	29.51

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Fig. 3 An alternative thermodynamic cycle that utilises both the heat of fusion of LiAl and its enthalpy of hydrogen absorption. Green arrows indicate “day-time”, heat absorbing reactions while red arrows indicate “night-time” heat releasing reactions.⁴¹

(1) Night-time: at night-time the heat of fusion of molten LiAl_(liq) (ΔH = −14.8 kJ mol⁻¹ of LiAl)⁴¹ is exploited at 702 °C to release ∼423.7 kJ kg⁻¹* (* relative to mass of LiH + Al).

(2) Night-time: cooling the, now solid, LiAl_(s) down to 565 °C (i.e. below the peritectic point) exploits the heat capacity of LiAl to release a further 236 kJ kg⁻¹* of heat.

(3) Night-time: by reacting LiAl_(s) at 565 °C with hydrogen to form LiH and Al, releasing 1385.1 kJ kg⁻¹ of heat, according to the reaction LiAl_(s) + 1/2H₂ = LiH_(s) + Al_(s).

(4) Day-time: the reverse of steps (2) and (3) can be combined to absorb 1621.1 kJ kg⁻¹ of thermal energy during heating from temperature 565 to 702 °C while desorbing H₂.

(5) Finally, at 702 °C the latent heat of melting of LiAl is utilised to absorb 423.7 kJ kg⁻¹.

According to our calculations, this cycle would decrease the practical cost of the Li–Al–H system by ∼44%. While the calculated cost for the raw materials using this cycle is high (∼US$29.5 per kilowatt hour of thermal energy (kW h_th) not including low temperature hydrogen storage, engineering and installation) compared to current state-of-the art molten salt systems (∼US$25–40 per kW h_th (ref. 42)), it does highlight a general method for reducing the cost of thermal energy storage based on metal hydrides.

Conclusion

Ball milling of LiH with Al has created a homogenously dispersed aluminium within the LiH phase, leading to reaction of LiH with Al. The thermodynamic assessment of the solid solution phase between LiAl and Al determined a desorption enthalpy and entropy of 96.8 kJ mol⁻¹ H₂ and 114.3 J K⁻¹ mol⁻¹ H₂. It was shown that at temperatures above 596 °C, the determination of thermodynamics parameters were difficult to establish, due to the co-existence of the solid solution of LiAl and a liquid phase. Below 596 °C, the Li–Al–H system has a low hydrogen equilibrium pressure (<2 bar), rapid kinetics, a flat desorption/absorption plateau with minimal hysteresis between absorption and desorption. This means that the Li–Al–H system shows many of the characteristics required by a metal hydride for use as a heat storage material (such as in concentrated solar thermal power plants). However, the relatively high cost of the raw material, even with the cost-reducing addition of aluminium, means that its use would probably be limited to niche applications.

Acknowledgements

CEB, DAS and PJ acknowledge the financial support of the Australian Research Council (ARC) for ARC Linkage Grants LP120100435 and LP150100730. CEB also acknowledges ARC LIEF grants LE0989180 and LE0775551, which enabled the XRD and gas sorption studies to be done. TRJ acknowledge the Danish National Research Foundation, Center for Materials Crystallography (DNRF93), The Innovation Fund Denmark (project HyFill-Fast), the Danish Research Council for Nature and Universe (Danscatt) and the Carlsberg Foundation.

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Footnote

† Electronic supplementary information (ESI) available. See DOI: 10.1039/c6ra16983j

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