Advances in catalysing the hydrogen storage in main group metals and their tetrahydroborates and tetrahydroaluminates

Abstract

Hydrogen is a promising clean and renewable energy source; however, its efficient storage is one of the key challenges in establishing the sustainable hydrogen economy. The light main group metals and their tetrahydroborates and tetrahydroaluminates show great potential for high hydrogen storage capacity close to ambient conditions; however, their high hydrogenation and dehydrogenation temperatures, sluggish kinetics, and limited reversibility have always been an obstacle for practical applications. Large efforts have been devoted to modifying the thermodynamic and kinetic properties of these systems, and reviewing these efforts and highlighting future directions are the aims of the present review. Based on recent research, the application of multicomponent systems utilizing multiple modification methods, such as catalysis, nanoconfinement, alloying, and structure engineering, is essential for enhancing the storage conditions. The synergistic effect of multiple catalysts is now a key requirement to address various steps of the overall process, including forming/breaking the H–H and metal–H bonds, transporting hydrogen and heat, and suppressing the formation of side products. Compared to pristine systems, tremendous improvement has been achieved. Catalysed AlH₃ decomposition can now operate as a one-way hydrogen source below 100 °C and the Mg/MgH₂ hydrogen storage system exhibits good cyclic performance at elevated temperatures. Metal hydrides, tetrahydroborates, tetrahydroaluminates, and their composite systems face challenges in achieving close to ambient operating conditions and cyclic stability. As the demand for improved hydrogen energy storage is expected to grow, further research for the enhancement of these systems will continue to advance the state of hydrogen storage technology.

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Broader context

Development of effective storage technologies for the hydrogen industry will support decarbonisation and shifting away from fossil fuels. Mobile applications require safe and robust technologies operating close to ambient conditions. To address this, storage materials capable of delivering and absorbing hydrogen under mild conditions need to be explored. Promising material groups for mobile hydrogen storage are the light main group metals and their tetrahydroborates and tetrahydroaluminates. We demonstrate in our present study the possibility of improving the storage performances of these compounds by tuning their thermodynamic and kinetic properties utilizing multiple modification methods, such as multi-component catalysis, nanoconfinement and alloying. Given the potential of these techniques, findings presented in this review may enable the much-needed improvements in solid-phase hydrogen storage to develop high-capacity storage systems operating close to ambient conditions.

1. Introduction

Three of the most difficult challenges facing mankind today are global warming caused by air pollution, increasing energy demand, and depleting sources of fossil fuels. The three are interconnected as current energy production is heavily based on fossil fuels,¹ the combustion of fossil fuels leads to about three quarters of global greenhouse gas emissions,¹ and depleting fossil fuel sources endanger energy supply.² 76.5% of the global primary energy consumed in 2023 was based on fossil fuels,¹ and the mass of carbon in the atmosphere has increased by 48% from eighteenth century pre-industrial levels.³ The world therefore needs an alternative energy source and technology to shift away from fossil fuels. Water is abundant on Earth and splitting water into hydrogen and oxygen and using hydrogen to generate energy if needed would provide a sustainable and carbon-free energy cycle if hydrogen is generated by utilizing renewable energy (Fig. 1). Several governments around the globe have realised the need for technology change and abandoning dependence on fossil fuels and started to build a ‘hydrogen-powered economy’.^4–8 Key enabling technologies to realize these plans are needed for the production, storage, distribution, and application of hydrogen. The storage of hydrogen, however, is challenging due to its inherent nature; hydrogen is a low-density gas under ambient conditions (0.0813 kg m⁻³ at 1 bar and 25 °C)⁹ and storage, transport and utilization, especially in mobile applications, require high storage densities. The U.S. Department of Energy (DOE) technical system targets for onboard hydrogen storage for light-duty fuel cell vehicles highlight these challenges and formulate targets for industry and academic research (Table 1).¹⁰ Considering these, a 615-fold increase of the density of hydrogen in the storage system must be achieved compared to that of ambient gas, which is even more demanding for the storage material as the storage vessel also has volume and mass.

Fig. 1 The carbon-free hydrogen energy cycle.

Table 1 DOE technical system targets for onboard hydrogen storage for light-duty fuel cell vehicles¹⁰

	System properties
	Gravimetric capacity (wt%)	Volumetric capacity (kg m^−3)	Min/max delivery temperature (°C)	System fill time (min)	Operational cycle life	Dormancy^a (days)
a Considering that the vehicle is parked at an ambient temperature of 35 °C; minimum until first release from initial 95% usable capacity.
2025 target	5.5	40	−40/85	3–5	1500	10
Ultimate target	6.5	50	−40/85	3–5	1500	14

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The hydrogen storage system needs to fulfill several requirements, including safety, high gravimetric and volumetric storage capacity, sufficient rate of hydrogen delivery and recharging, close to ambient operating conditions, easy handling, abundant resources, and competitive price. None of the current systems can fulfill all these requirements. There are several research directions reflecting application needs, and hydrogen may be stored in molecular form or in compounds. The current most advanced molecular hydrogen storage methods are storage of hydrogen as a pressurized gas at ambient temperatures (technology readiness level¹¹ (TRL) = 9) and as liquified hydrogen at cryogenic temperatures (TRL = 9). Drawbacks of these methods include the safety issues and cost of compressing hydrogen to pressures up to 700 bar, the low volumetric density of the pressurized gas and the high production cost of high pressure storage containers for high pressure gaseous storage, and the cost of liquefaction, bulky insulation to maintain low temperatures (<33 K) and the control of the boil off for liquid hydrogen storage.^12–15 Large efforts have been devoted to developing hydrogen storage based on molecular hydrogen adsorption on solid materials such as doped carbons, metal organic frameworks (MOFs), covalent organic frameworks (COFs) and porous aromatic frameworks (PAFs).^16–19 However, high storage capacity can only be achieved at cryogenic temperatures (TRL = 5); ambient temperature storage satisfying DOE targets is not yet feasible for storage of hydrogen via adsorption.^16–19 Storing hydrogen in compounds is an alternative to molecular hydrogen storage and has several advantages such as the stability of many hydrogen compounds under ambient conditions, the high gravimetric and volumetric hydrogen capacity (up to about 20 wt% or 150 kg m⁻³), and the relatively easy storage and transport.^20–23 Hydrogen compounds for storage may be classified into three groups, namely (1) molecular hydrides (e.g., ammonia, perhydro-dibenzyltoluene, methylcyclohexane), (2) interstitial hydrides of transition metals and their alloys (e.g., TiFe, Nd₂Ni₇, high entropy alloys such as V₃₅Ti₃₀Cr₂₅Fe₅Mn₅), or (3) ionic and polymeric hydrides (e.g., LiH, MgH₂). An advantage of currently considered liquid organic molecular hydrides is that they closely resemble conventional fuels and existing infrastructures could be used to store and transport them safely.^24–27 However, industrial environments are required for the hydrogenation and purification of the hydrogen-lean and hydrogen-rich forms of storage compounds and mobile applications can consider these hydrogen carriers only as “off-board” regenerable materials. None of the currently considered molecular liquid hydrogen carriers fulfill all storage requirements, namely low melting (<−30 °C) and high boiling (>300 °C) point, high volumetric and gravimetric hydrogen storage capacities, stability for long lifespan, low dehydrogenation temperatures, low production cost, and low toxicity. The toluene/methylcyclohexane system is technologically the most advanced (TRL = 9); a demonstration project has been carried out by hydrogenating toluene in a chemical plant in Brunei, shipping methylcyclohexane to Japan, dehydrogenating methylcyclohexane in Japan, and shipping toluene back to Brunei.^26,28 Storing hydrogen in interstitial, ionic or polymeric metal hydrides solves the problem of volatility of molecular hydrides; however, it introduces new challenges. The hydrogen storage in interstitial hydrides has several advantages such as reversibility, fast kinetics, and the ability of several transition metals and alloys to absorb and desorb hydrogen under ambient and near ambient conditions, as well as the tunability of operating pressure and temperature, including the region in which electrolysers and fuel cells operate (between 1 and 10 bar pressure and below 100 °C) (TRL = 9).^22,29,30 Disadvantages, however, are the high cost and low gravimetric storage capacity, well below the DOE targets, which prevent their wide scale use. Nonetheless, interstitial hydrides are successfully applied as hydrogen storage materials in vehicles where the heavy storage material is not a weight penalty, for example in forklifts, submarines, and canal boats.^31–33

Promising groups of hydride compounds for hydrogen storage, that may fulfill storage requirements in the future and bypass the limitations of storage methods mentioned above, are the ionic and polymeric binary hydrides of light main group metals (alkali and alkali earth elements and aluminium) and their tetrahydroborates and tetrahydroaluminates (which can be considered to be complexes of metal hydrides and BH₃ and AlH₃, respectively), as well as their mixtures (reactive hydride composites, RHCs) due to their high gravimetric and volumetric hydrogen storage capacities (Table 2). These compounds are stable under ambient conditions in an inert atmosphere and several of the relevant metals, e.g. Mg and Al, are cheap, light, and abundant on Earth, which makes them very attractive for both mobile and stationary storage systems. Inherent properties of these compounds, namely the high temperatures and pressures needed for hydrogen uptake and release, the sluggish kinetics, and the partial- or non-reversibility of hydrogen release and uptake, as well as the formation of side products, however, have impeded their wide scale application to date (TRL = 3–5 depending on the compound). Due to their potential in filling the gap in reversible and high-capacity hydrogen storage under close to ambient conditions, large efforts have been devoted during the last few decades to bring operation conditions closer to ambient conditions by improving the kinetics and thermodynamics of the hydrogenation and dehydrogenation reactions using additives and catalysts. Considering this and the increasing demand for effective storage materials, as well as the rapidly changing hydrogen storage industry, the aim of the present work is to review advances in improving hydrogen storage properties of binary hydrides of light main group metals and their tetrahydroborates and tetrahydroaluminates, and to highlight their advantages, limitations, and challenges, as well as to point out future trends in developing storage materials for the hydrogen industry. We consider DOE initiatives and the operation conditions of fuel cells as research targets, namely storage capacities of 6.5 wt% and 50 kg m⁻³ and lower than 10 bar pressure and 100 °C temperature (note that the DOE target is the storage system target and the storage material must perform better, but to keep discussion simple we use these target numbers for the storage material in this article). Additives to hydrides which do not contain hydrogen decreases the storage capacity; therefore large amounts of additives used to alter reaction conditions are not considered in this review. Considering gravimetric capacities in Table 2, pristine NaH, KH, CaH₂, and KAlH₄ do not satisfy DOE targets (Table 1), but can be considered as one of the components of RHCs. Heavier I–III group metals are not relevant for hydrogen storage due to their weight, toxicity, and/or low abundance. Although beryllium compounds have exceptionally high hydrogen storage capacity (Table 2), they are not considered for hydrogen storage because beryllium ores are not abundant, beryllium is expensive, and beryllium salts, including the hydride, are extremely toxic.^34,35 Al(BH₄)₃ is a volatile pyrophoric liquid at room temperature, explosive in air, has poor storage stability, and releases also diborane upon heating; therefore, its practical use in hydrogen storage is restricted.^36–38 To discuss the highlighted topic, this review is organised in subchapters according to the chemical groups of hydrogen compounds.

Table 2 Gravimetric storage capacity of main group metal hydrides and alkali and alkali earth tetrahydroborates and tetrahydroaluminates^ab

a Calculated in this work using atomic masses and the density of hydrides.^35,39–44b BeH₂, MgH₂, and AlH₃ have a polymeric structure with three-centre M–H–M bonds (M = Be, Mg, or Al). Colour codes: above the DOE target (green and yellow, green is higher than yellow and primary target), below the DOE target (pink), and not considered for hydrogen storage (no colour).c Al(AlH₄)₃ = AlH₃.

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2. Results and discussion

2.1. General considerations and experimental determination of thermodynamic and kinetic parameters of hydrogen absorption and release

The hydrogen absorption by metals is a multiple-step process, and involves several consecutive steps, namely (a) physisorption of hydrogen, (b) dissociation of hydrogen into atoms on the surface, (c) penetration of hydrogen atoms in the subsurface, (d) hydrogen diffusion in the bulk material, and (e) nucleation and growth of the hydride phase. The reverse process, hydrogen release, follows the opposite multistep process, starting with the dissociation of the metal–hydrogen bond. Each step of both absorption and desorption processes must overcome a range of activation energy barriers, and decreasing these energy barriers, especially that of the rate determining step, is of crucial importance for improving the rate and/or decreasing the reaction temperature of hydrogen absorption and desorption. Covalent and ionic bonds are strong; therefore, dissociation of molecular hydrogen and the metal–hydrogen bond is, in general, rate determining and thermodynamically limiting. The overall thermodynamics and kinetics govern the hydrogen uptake and release and altering these is an effective strategy for improving hydrogen storage properties in metals.

Thermodynamically, the hydrogenation reaction occurs if the Gibbs free energy (ΔG) of the reaction, related to the enthalpy of formation (ΔH), entropy of formation (ΔS), and temperature (T), is negative (ΔG = ΔH−T·ΔS). Since heat is released and entropy is reduced during the hydrogenation reaction (as hydrogen is moving from the gas phase into its bonded solid phase), the enthalpy and entropy both have negative signs; consequently, thermodynamically, and depending on the concrete value of ΔH, the hydrogenation is favoured at lower temperatures, but the reverse reaction, dehydrogenation, is favoured at higher temperatures when ΔG becomes positive. Considering an ideal hydrogen storage system using temperature swing, a right balance between ΔH and T·ΔS must exist where |ΔH| > |T·ΔS| at charging (hydrogenation at lower temperature, ideally under ambient conditions) and |ΔH| < |T·ΔS| at discharging (hydrogen release at elevated temperature). With very large ΔH, which is related to the strength of the metal–hydrogen bond, hydrogen release requires high temperatures. Considering that the entropy change during hydrogen absorption corresponds mostly to the change from molecular hydrogen gas to bonded hydrogen in solid hydrides, ΔS is approximately equal to the standard entropy of hydrogen for all metal-hydrogen systems, namely −130 J K⁻¹ mol⁻¹H₂.^45,46 The thermodynamic properties of metal-hydrogen systems are therefore usually characterized by the reaction enthalpy. (The same principle can be applied to tetrahydroborates and tetrahydroaluminates, too.) One way to alter thermodynamics is the application of additives which form alloys during dehydriding (see Fig. 2a). Non-hydriding additives decrease hydrogen storage capacity, as they dilute the metal; however, hydriding additives can preserve or even increase hydrogen storage capacity (Fig. 2a), which highlights the advantage of RHCs. In this sense, the formation of metal borides and alloys during the dehydrogenation reactions of tetrahydroborates, tetrahydroaluminates, and RHCs (formally the mixtures of two or more binary hydrides) is crucial for the desired thermodynamics of the system; the separation of metallic and boron phases is unfavourable (as will be discussed below). Thermodynamic parameters of hydrogen absorption and desorption are usually determined using the pressure–composition–temperature (PCT) phase diagram and the van’t Hoff equation (Fig. 2b).^45–48 Considering a metal–hydrogen system and the Gibbs phase rule, the system has two degrees of freedom (due to the presence of two phases and two components) when hydrogen starts to absorb. Hydrogen first dissolves in the metal lattice and forms a solid solution phase (α-phase), which has the same crystal structure as the bare metal and has low hydrogen concentration. In the α-phase region the concentration of hydrogen in the solid phase increases as the hydrogen pressure increases. At a certain point, nucleation of a higher concentration phase (β-phase) occurs. Due to the occurrence of an additional phase and thus the presence of three phases (α, β and gas), the system's degree of freedom is decreased to one, and the pressure stays constant (equilibrium pressure, p_eq) until the α- and β-phases coexist. After the disappearance of the α-phase, hydrogen enters the β-phase (solid solution) and the system's degree of freedom is two again and the concentration of hydrogen increases as the pressure is increases. As Fig. 2b shows, there are three phase regions in the PCT diagram, the α, α + β, and β regions separated by the transition dome curve (dotted line). The two-phase region (α + β) ends in a critical point, characterized by the critical temperature (T_cr) and pressure (p_cr), above which the transition from the α- to β-phase is continuous. When α and β phases coexist, there is a plateau in the isotherms, the pressure hardly changes during the α → β phase transition, therefore a significant amount of hydrogen is absorbed with small pressure increases. The plateau region (α + β isotherm), however, is not entirely flat due to the expansion of the lattice and small changes of ΔH as a function of hydrogen concentration. In addition, the absorption pressure is slightly higher than the desorption pressure due to unequal accommodation of elastic and plastic energies, which causes a ‘hysteresis’.⁴⁸ The hysteresis, in general, decreases with increasing temperature.⁴⁹ Typical hysteresis magnitudes (Δp/p_des, where Δp = p_abs − p_des) for selected systems as examples are 0.015 (at 401 °C) and 0.007 (at 497 °C) for Mg,⁵⁰ 0.51 (at 0 °C) and 0.24 (at 80 °C) for LaNi₅,⁴⁹ and 1.43 (at 25 °C) and 0.97 (at 85 °C) for TiFe.⁵¹ The pressure–temperature–concentration relationship can be expressed for both absorption and desorption by the extended van’t Hoff equation; by setting the hydrogen concentration to the value at the middle of the α + β phase region, the equation is simplified and the pressure–temperature relationship can be used to determine ΔH values experimentally (see Fig. 2b).

Fig. 2 (a) Generalized enthalpy diagram showing the effect of a catalyst on the kinetics and the effect of alloy formation on the thermodynamics for a hypothetical metal/metal hydride system (A = Li, Na, K, Mg, Ca or Al; B = non-hydriding alloying element; C = hydriding alloying element), (b) schematic PCT phase diagram on the left hand side and the van’t Hoff equation and the corresponding van't Hoff lines on the right hand side for hydrogen absorption and desorption on a typical metal (c_H = hydrogen concentration in the solid phase, p = hydrogen pressure, p_eq = mid-plateau equilibrium pressure, p_ref = reference pressure of 1 bar, p_cr = critical pressure, R = universal gas constant, T = absolute temperature, T_cr = critical temperature), and (c) schematic representation of a thermal event at various heating rates on the left hand side and the Kissinger line on the right hand side (β = heating rate, T_p = peak temperature, E_a = apparent activation energy, A = pre-exponential factor).

The rates of the hydrogenation and dehydrogenation reactions are of crucial importance for practical storage applications. They depend on the activation energy barrier of the absorption and desorption processes and on the temperature and hydrogen pressure, as well as on the surface area of the solid. Nanosizing metals and alloys to increase the rate of the heterogeneous hydrogenation reaction is an obvious step to increase the rate. Current research on improving the kinetics of hydrogen absorption/desorption aims to reduce the activation energy barrier using catalysts to allow milder reaction conditions (temperature and pressure) for practical applications. The Kissinger method is currently overwhelmingly popular to estimate the apparent activation energy (E_a) for both hydrogen absorption and desorption due to its simplicity (even if complex processes may not be well represented by a single activation energy).^52,53 The Kissinger equation is shown in Fig. 2c. The method is based on measuring a chemical (or physical) thermal event at several non-isothermal tests conducted at constant heating rates (β), using e.g. DSC, DTG, DTA, or MS, and measuring the peak (or onset) temperature of the event (T_p). A plot of ln(β/T_p²) versus 1/T_p leads to a straight line with the slope equal to −E_a/R (Fig. 2c). Although this method is simple, it oversimplifies the process by assuming single-step kinetics with first-order reaction. If the reaction model is not first order, its dependence on the heating rate introduces an inaccuracy which increases with decreasing E_a/RT. This error has been shown to be below 5% as long as E_a/RT is larger than 10;^53,54 therefore the E_a/RT > 10 relationship may be considered as a prerequisite for using the Kissinger method.⁵³ In addition, the linear heating rate is an integral part of the Kissinger equation; therefore nonlinear heating programs can be applied only if the extent of conversion of the reactant to products doesn’t vary significantly with the heating rate at the event peak.⁵³ Further drawbacks of the method are its inability to detect complexity of processes (namely multiple steps covered by one event peak) and inapplicability to any data obtained on cooling.⁵³ The appearance of multiple peaks in experimental data requires treating them separately (note that a single peak doesn’t necessarily mean a single-step process). If peaks are overlapping, they need to be deconvoluted and the Kissinger method should be applied separately to each individual peak to estimate activation energies of specific processes. Considering its limitations, the Kissinger method should be used for quick estimation of apparent activation energies in comparative studies of simple transformations, rather than providing details of process kinetics. An alternative, although less frequently used and more time consuming, way to obtain activation energies is based on calculating rate constants at several temperatures using the Johnson–Mehl–Avrami–Kolmogorov (JMAK)^55,56 equation [“linearized” form: ln[−ln(1 − α)] = η ln(k) + η ln(t), where k = rate constant, α = converted reaction fraction, η = growth dimensionality, t = time] by fitting (de)hydrogenation curves and calculating the activation energy using the Arrhenius equation (the slope of ln(k) versus 1/T is equal to −E_a/R). The JMAK model includes simultaneous nucleation, growth and impingement and offers a more detailed model of phase transformation kinetics than the Kissinger method. Limitations may arise from model specifics, namely assuming phase transformation by random nucleation and growth, a zero critical nucleus size, requirement of a large number of nuclei for accurate prediction, and nucleation and growth following a predetermined law.^56–59 It is important to note that although the limitations of the Kissinger model have been stretched, its popularity continues;^53,60 however, this may change in the future by wider application of more sophisticated, currently underutilised kinetic models for hydrogen storage materials.⁵⁶

Macroscopic PCT and thermal analyses are the most widely used methods for obtaining information about the kinetics and thermodynamics of the system due to easy accessibility and cost of the instrumentation, but modern operando tools (e.g., synchrotron radiation X-ray power diffraction (SR-XRPD), neutron scattering (NS), and in situ transmission electron microscopy (TEM)) can allow direct observation of phase evolution and kinetics. The SR-XRPD light source is more intense than those of laboratory XRPD instruments, highly collimated, tuneable in a wide energy spectrum, and has a short pulse time structure; therefore SR-XRPD allows high-resolution and time-resolved studies to investigate microstructures of materials, including crystal structure, crystallite size, phase composition, defect and strain.^61,62 Recording in situ SR-XRPD data throughout the whole hydrogenation/dehydrogenation reaction can provide information on the reaction mechanism, transient phases, and kinetics of phase transitions. Hydrogen has a high scattering cross section with neutrons; therefore, neutron scattering techniques, including neutron diffraction (ND), total neutron scattering (TNS), inelastic neutron scattering (INS), quasi-elastic neutron scattering (QENS) and small-angle neutron scattering (SANS), are valuable tools for studying hydrogen storage materials.^63–68 ND and TNS can locate hydrogen atoms in the crystal structure which is practically impossible with XRD. TNS is also powerful for studying amorphous materials. Using neutrons also has the benefit of being able to put the sample under a greater range of pressures, temperatures and gas environments, compared to X-ray instruments, as the neutrons can pass through thick stainless-steel cells and instruments. This makes them particularly useful for operando measurements. INS and QENS are conceptually similar to infrared and Raman spectroscopy and provide information on molecular and crystal vibrations as well as diffusion and liberation of hydrogen, but in contrast to IR and Raman all vibrations are visible (no selection rules) and the spectrum is dominated by hydrogen displacements due to the approximately order of magnitude higher scattering cross section of neutrons for hydrogen as compared to other elements. The SANS pattern contains information on atomic positions and shapes of particles and can therefore be used to obtain information on the size, shape, and interaction of nanoparticles. Both SR-XPRD and NS facilities are uncommon and accessing them can be expensive and highly competitive. In situ TEM is a much cheaper and more widely available technique suited to follow phase transformations and kinetics.^69–71 Real-time observations during dehydrogenation can be readily performed in the high vacuum environment of the instrument; however, studying hydrogenation (which requires a high-pressure hydrogen atmosphere) is challenging. The reactant and hydrogen must be sealed within a tiny space of a gas cell equipped with electron-transparent windows, allowing a gas layer only a few micrometres thick and separated from the other parts of the microscope. An advantage of the sealed-gas-cell approach is that normal TEM instruments can be modified to run in situ hydrogen absorption experiments by modifying the sample holder.⁷¹

The hydrogen release from binary hydrides is a one-step chemical reaction, but it is multi-step for tetrahydroborates, tetrahydroaluminates, and RHCs involving the formation and decomposition of alkali and earth alkali hydrides in the last step. Fig. 3 shows a simplified decomposition scheme for hydrogen release from these compounds and a more complex scheme for Mg(BH₄)₂⁷² to illustrate the complexity of the problem. Phase separation of metals or metals and boron is an issue for reversibility, and the formation of alloys or borides is favourable for the overall thermodynamics of the system (see Fig. 2a). In addition, passivation of metal surfaces during hydrogenation influences kinetics negatively. The diffusion of hydrogen in LiH and MgH₂ was found to be much slower than in respective metals (there is a three–five order of magnitude difference in diffusion coefficients);^73–75 therefore the hydride layer passivates the surface. As a result, for example, only partial hydriding was observed on magnesium particles larger than 50 μm due to fast hydride formation.⁷⁵ Downsizing metals and alloys has a positive effect in solving this issue and improves kinetics as mentioned above. Since the hydrogen absorption/desorption leads to nanocrystalline solids, due to volume expansion and contraction,⁴⁵ downsizing is a natural phenomenon after one hydrogen absorption/desorption cycle. Nanosizing and activation by removing the passivating oxide layer is therefore essential before the first hydrogen absorption step. General strategies to activate metal surfaces, in addition to nanosizing, are thermal annealing and surface functionalization and doping. Thermal annealing creates cracks on the passivating oxide film and surface metal vacancies, thereby exposing the metal to hydrogen and reducing the energy barrier of hydrogen dissociation, respectively.⁷⁶ The most widely used method for surface modification and functionalization to create more active sites for molecular hydrogen dissociation and hydrogen atom diffusion on the material surface is currently ball milling with additives and/or catalysts. In general, ball milling introduces mechanical stress and modifies the structure, additives open diffusion channels and catalysts facilitate hydrogen dissociation.^77–79

Fig. 3 (a) Simplified decomposition scheme for metal hydrides, tetrahydroborates, and tetrahydroaluminates, (b) volumetric temperature programmed desorption curve of Mg(BH₄)₂ and (c) decomposition pathways of Mg(BH₄)₂ (observed hydrogen evolution steps, see panel (b), marked by dashed lines; amorphous phases denoted by asterisks). Reproduced from ref. 72 with permission from Elsevier Ltd, copyright 2008.

2.2. Tuning the hydrogen storage in lithium, magnesium, and aluminium

Li, Mg, and Al have high hydrogen storage capacities (Table 1, Be is not considered and the capacities of Na, K, and Ca are below the DOE target); however, their practical application faces different challenges. LiH is extremely stable, and the formation enthalpy is −181.2 kJ mol⁻¹ H₂ (−90.6 kJ mol⁻¹ LiH, Table 3); therefore the dehydrogenation of LiH requires very high temperatures (the main desorption peak is observed at 671 °C).⁷³ In addition, there is an activation energy of 52.8 kJ mol⁻¹ for the reaction of H₂ with molten lithium.⁸⁰ Consequently, lithium reacts with hydrogen at high temperatures or pressures (at temperatures above 700 °C under ambient pressure or at 500 bar pressure at room temperature).^73,81,82 LiH therefore cannot be considered as a practical hydrogen storage material as it is economically unviable for cyclic energy storage. Destabilizing the LiH phase and catalysing the hydrogenation and dehydrogenation reactions would be essential to reach milder reaction conditions. To underpin the idea, it is important to note that Wang et al. synthesized LiH nanoparticles with a Ni shell at the surface (composition 85.5 wt% Li, 14.0 wt% Ni and 0.5 wt% Ti, corresponding to an overall hydrogen capacity of 10.8 wt% H₂) and achieved full hydrogen release/uptake in less than 50 min at 350 °C.⁷³ Extensive cycling, however, could not be achieved at this temperature due to the formation of large particles behaving like bulk LiH and thus not cycling hydrogen at 350 °C; the reversible storage capacity of LiH/nano-Ni remained limited to 2 wt% instead of the potential 10.8 wt% H₂. The experiment, however, demonstrated the possibility of tuning the Li/LiH system for hydrogen storage. Thermodynamic tuning of the Li/LiH system was reported earlier,^83–85 but reversible hydrogen cycling at the level of DOE targets could not be achieved. Alloying Li with porous silicon nanowires (molar ratio: 4.4 : 1), for example, resulted in a reduced hydrogen absorption enthalpy of −118 kJ mol⁻¹ H₂ (−59 kJ mol⁻¹ LiH) and the alloy could reach a maximum observed H₂ absorption capacity of 3.95 wt% at 400 °C and 40 bar.⁸³ Hydrogen storage properties of LiH were demonstrated to be size dependent; LiH nanoconfined in high surface area graphite could release 1.9 wt% hydrogen with a peak at 340 °C and could reversibly absorb 1 wt% hydrogen at 300 °C and 60 bar hydrogen pressure.⁸¹

Table 3 Comparison of thermodynamic and kinetic properties of AlH₃, LiH and MgH₂

Hydride	ΔfH^a (s) (kJ mol^−1)	ΔrS°^b (s) (J mol^−1 K^−1)	ΔfG^c (kJ mol^−1)	Ea (des.)^d (kJ mol^−1 H2)	Cycling properties^e
a Enthalpy of formation from elements at 25 °C from ref. 9 and 86.b Entropy change of the hydride formation reaction at 25 °C, calculated using S° (298 K) values from ref. 9.c At 25 °C.d Overall activation energy of hydrogen release from ref. 87 and 88.e Number of stable cycles and capacity retention (in parenthesis) from ref. 89; n.a. = not available.
AlH3	−9.9	−194.3	48.0	102.2	n.a.
LiH	−90.6	−74.1	−68.5	n.a.	n.a.
MgH2	−76.2	−132.3	−36.7	160	515 (ca. 60%)

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AlH₃, in contrast to LiH, is metastable at room temperature; although the formation enthalpy is slightly negative (−9.9 kJ mol⁻¹ for α–AlH₃), the Gibbs free energy (ΔG) is positive (+48 kJ mol⁻¹ at 25 °C) (Table 3). AlH₃ fully decomposes in the 100–200 °C temperature region, and the activation energy of decomposition can be decreased by using alkali hydrides and transition metal compounds.^90–92 The hydrogenation of aluminium, however, requires extremely high pressures (above 25 000 bar);^90,91,93 therefore AlH₃ can be considered only as a one-way hydrogen storage material and cannot be regenerated on-board a vehicle.^90,93 The findings from recent research to bring the AlH₃ decomposition temperature below 100 °C by doping catalytic amounts of various metal compounds into the AlH₃ matrix through ball milling are summarized in Table 4. Substantial improvement in reaction kinetics was observed compared to pristine AlH₃, promoted by metal hydrides, metals, and solid solutions, formed in situ or by the accordion structure of MXene additives (Table 4). The catalytic effect of these materials is based on lattice distortion,⁹² destruction of the aluminium oxide layer by Ti₃C₂,⁹⁶ enhanced electron transfer and hydrogen diffusion by in situ formed transition metals or metal compound species,^94,95,97,98 and/or destabilization of AlH₃ through increased Al–H bond length.^98–100 As an example, the multivalent conversion of TiO₂ and Pr₆O₁₁ promoted the electron transfer of H⁻ in AlH₃, and the newly generated PrH_2.37 hydrogen diffusion pathway promoted the release of hydrogen (Fig. 4a). It seems that the DOE system targets considering storage capacity and hydrogen release temperature can be met with catalysed AlH₃ decomposition, which makes AlH₃ a promising one-way hydrogen source medium for mobile applications. Cheap production or regeneration of the hydride, however, is not solved yet.

Table 4 Effect of catalysts on the hydrogen release properties of AlH₃

Catalyst precursor	Active form^a	Capacity^b (wt%)	Decomp. temp.^c onset/peak (°C)	Isothermal H2 release wt%/°C/min	Ea (des.) (kJ mol^−1 H2)	Ref.
a Catalytically active component formed during ball milling of precursors; ss = solid solution.b Experimentally determined maximum storage capacity.c n.a. = not available.
TiB2 (2.5 wt%)	Al–Ti–B ss	8.7	78/n.a.	5.3/80/200	86	92
				7.3/110/60
NbF5 (1 mol%)	Nb, AlF3	6	60/122		96	94
LaNi5 (10 wt%)	Ni	8.6	86/n.a.	6.3/100/60	98	95
				7.7/120/30
Ti3C2 (4 wt%)	Ti3C2	8.1	61/n.a.	6.9/100/20	40	96
CeO2 (1 mol%)	CeH2+x	7.8	114/n.a.	6.5/100/100	69	97
TiO2–Pr6O11 (1 mol%)	PrH2.37	8.3	43/119	5.9/100/60	56	98
Bi@C (5 wt%)	Bi	8.0	80/n.a.	5.4/120/40	87	99
Fe@C (3 wt%)	Fe@C	8.5	94/n.a.	6.2/120/40	78	100

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Fig. 4 (a) Mechanism for the catalytic effect of TiO₂ and Pr₆O₁₁ on AlH₃. Reproduced from ref. 98 with permission from the American Chemical Society, copyright 2021, (b) mechanism for the hydrogenation of Mg catalysed by ZrMn₂ nanoparticles. Reproduced from ref. 118 with permission from the Royal Society of Chemistry, copyright 2019, and (c) dehydrogenation/hydrogenation cyclic curves of MgH₂ with 10 wt% TiO₂-ZrO₂-V₂O₅ catalyst at 300 °C. Reproduced from ref. 127 with permission from Elsevier Ltd, copyright 2024.

Considering the extreme stability of LiH and the thermodynamic instability of AlH₃, the Mg/MgH₂ system seems to be the golden middle road for reversible hydrogen storage. Hydrogen absorption in Mg and especially desorption from MgH₂, however, require high temperatures. Considering the van’t Hoff equation (Fig. 2b), the enthalpy of MgH₂ formation of −74 kJ mol⁻¹ H₂¹⁰¹ and an entropy change of 130 J K⁻¹ mol⁻¹ H₂, the temperature required to reach 1 bar hydrogen equilibrium pressure at MgH₂ desorption is about 300 °C. If MgH₂ could be destabilized to have a desorption enthalpy below 48.5 kJ mol⁻¹ (TΔS at 100 °C), hydrogen release would be thermodynamically favoured at 100 °C. Improving both thermodynamics and kinetics of the Mg/MgH₂ system is therefore essential for reversible hydrogen storage; however, there is a limit of 15 wt% for the additive concerning the ultimate DOE target of 6.5 wt% and the theoretical MgH₂ storage capacity of 7.66 wt%. Alloying Mg with Al was found to improve both kinetics and thermodynamics of the system;^102,103 however, Al served only as a diluent due to the instability of AlH₃ (see above). Despite improvements, the kinetics of hydrogen absorption in the Mg₉₀Al₁₀ alloy (11 wt% Al) was still slow, namely 2.7 wt% and 4.6 wt% hydrogen capacity could be achieved at 260 °C and 380 °C, respectively, at 30 bar hydrogen pressure within 2 hours.¹⁰² The Mg–Zn alloy containing 17 wt% Zn performed better and could absorb 5.0 wt% hydrogen in 20 min at 300 °C and had a hydrogen storage capacity of 6.1 wt% at this temperature (but below the DOE target); the enhanced hydrogen storage properties for this system were attributed to improved kinetics (low hydrogen absorption activation energy of 56.3 kJ mol⁻¹) rather than the change in enthalpy.¹⁰⁴ Mg has also been alloyed with Li; however, the hydrogenation of Mg–Li alloys always leads to a mixture of MgH₂ and LiH, from which MgH₂ could be thermally decomposed up to 450 °C, but LiH remained stable.¹⁰⁵ Considering this, Mg–Li alloys are not good candidates for reversible hydrogen storage because their reversibility and stability are very similar to those of pure Mg, but with lower storage capacity.

Early long term hydrogen cycling experiments on Mg in 1984,⁸⁹ involving 515 cycles and using high purity hydrogen (99.9997%), an operating temperature of 390 °C, and 30 min charging and discharging time at 32 bar H₂ pressure and vacuum, respectively, highlighted the potential of reversible hydrogen storage in Mg and challenges as well, namely the high operating temperature and pressure, decreasing storage capacity by cycle numbers and the sensitivity of magnesium to oxygen and nitrogen impurities in repeated experiments.¹⁰⁶ The high activity of freshly prepared pure magnesium to oxygen, water and nitrogen was pointed out recently.^107,108 To bring operating conditions closer to ambient conditions and improve hydrogen cycling properties, a very wide range of metal compounds have been screened during the last few decades to find effective catalysts to improve kinetic and cycling properties for magnesium by reducing activation energies;^{74,101,109–115} these included transition and rare earth metals, alloys, oxides, halides, carbides, carbonitrides, nitrides, sulphides, hydrides, carbon materials, MXenes, and their mixtures. The catalytic activity of these compounds was based on various mechanisms,^114–144 namely absorbing and delivering hydrogen to Mg thereby acting as a “hydrogen pump”,^{114,116–118,123,129,130,133,135,137,139,142,144} providing channels for fast hydrogen diffusion,^{114,115,121,125,128,130,131,137,139,142} providing nucleation sites,^{115,119,121,125,135} providing abundant oxygen vacancies in oxides,^116,127,132 and weakening the Mg–H and/or H–H bond strength.^{114,120,123,124,136,138,141,143} The addition of carbon nanotubes (CNTs) or carbon was found to be advantageous for preventing the aggregation of Mg/MgH₂ particles, thereby preserving stable cycle properties,^{122,130,142,143} as well as aiding hydrogen diffusion and heat transfer.¹⁴¹ The catalytic mechanism can be structured into three groups: (1) kinetic facilitation of dehydrogenation/hydrogenation processes (e.g., LiCoO₂,¹¹⁴ TiO_2−x,¹¹⁵ V₂O₅,¹¹⁶ NiV₂O₆,¹¹⁷ ZrMn₂,¹¹⁸ TiFe,¹²² Ni₃ZnC_0.7,¹²³ Nb₂CTx MXene,¹²⁴ TiO₂@MnCo₂O_4.5,¹²⁵ TiMn₂,¹²⁶ TiO₂-ZrO₂-V₂O₅,¹²⁷ Ti₂VC₂ MXene,¹²⁸ shell-like Ni/Mo₂N,¹²⁹ Ni₃Fe/Ni/NiFe₂O₄@C,¹³⁰ Nb₂O₅-C/Ti₃C₂Tx MXene,¹³¹ MoO_3−x,¹³² MXene Ti₃C₂@NiO,¹³³ NiSe₂,¹³⁵ Ni/TiO₂,¹³⁶ Ce_0.6Zr_0.4O₂,¹³⁷ Co₃V₂O₈,¹³⁸ NaNbO₃,¹³⁹ Al₃V,¹⁴⁰ Ti_0.9Zr_0.1Mn_1.5V_0.3,¹⁴¹ NiS@C,¹⁴² Ni/Mo2C@C,¹⁴³ Ni₃Mn¹⁴⁴), (2) structural stabilization of nanoparticles (e.g., V₂O₅,¹¹⁶ CoFeB/CNT,¹²¹ CNT,¹²² carbon^123,142,143), and (3) thermodynamic tuning (e.g., TiH₂¹¹⁹). Most of the materials in catalyst groups 1 and 2 did little to reduce the thermodynamic stability of MgH₂. In general, transition metal oxides acted as precursors for in situ formation of catalytically active species (see Table 5).

Table 5 Effect of catalysts on the hydrogen absorption and desorption kinetics of Mg/MgH₂^a

Catalyst (active form)	Isothermal absorption				Non-isoth. des. onset/peak	Isothermal desorption				Cycles/capacity retention				Ea (abs/des)	Ref.
	wt%	°C	bar	min	°C/°C	wt%	°C	bar	min	wt%	°C	No.	%	kJ mol^−1 H2
a RT = room temperature, NS = nanosheet, CNT = carbon nanotube, n.a. = not available.b As-prepared MgH2.c Solar heating under a light density of 3.5 W cm^−2.
(no catalyst)^b	1.3	210	32	10	310/n.a.	1.1	325	0.003	10	n.a.	n.a.	n.a.	n.a.	81/156	114
	6.1	270	32	10		7.1	375	0.003	10
LiCoO2 (7 wt%)	2.2	125	32	10	180/n.a.	5.5	250	0.003	60	6.2	300	20	94	45/49	114
(Mg2Co–Mg2CoH5)	3.8	150	32	10		6.6	325	0.003	10
	6	200	32	10
5% K–TiO2−x	4	RT	80	6	194/247	6.7	300	0.001	10	6.5	300	30	100	n.a./69	115
(K2Ti2O3)	6	RT	80	180
	6.7	200	30	60
10% H–V2O5	2.4	30	30	60	185/n.a.	4.4	225	0.1	10	5.7	275	100	99	43/85	116
(VH2, V)	4.1	100	30	30		6	275	0.1	5
10% NiV2O6	2.2	75	30	50	n.a.	4.5	250	0.01	75	6.1	325	50	90	25/93	117
(Mg2Ni, V2O5)	5.6	150	30	50		5.2	300	0.01	10
10% ZrMn2	4.3	50	30	40	182/251	6.7	300	—	5	6.5	275	10	85	22/82	118
	5.3	100	30	10
10% TiH2	4	RT	20	240	180/n.a.	4	240	0.001	20	6	300	80	100	n.a./58	119
	5	240	20	1
5% Fe NS	6	200	32	10	182/223	5.4	300	0.03	10	6	300	50	83	n.a./41	120
10% CoFeB/CNT (Co3MgC, Fe, CoFe)	6.2	150	50	10	177/n.a.	6.5	300	vac.	30	6.4	300	20	99	n.a./83.2	121
TiFe + CNT (10 + 5 wt%)	4.5	125	30	60	210/n.a.	6.1	275	0.03	30	6.2	300	10	100	61/n.a.	122
						6.2	300	0.03	14
Ni3ZnC0.7@C (2.5 wt%) (Mg2Ni, MgZn2)	5.2	150	50	3	312/360	5.7	400	0.001	5	4.5	400	20	95	42/101	123
Nb2CTx–5Na (5 wt%)	4	100	30	1	n.a.	5.6	250	0.01	12	6.2	300	40	90	n.a./70	124
	5.5	125	30	1		6.4	300	0.01	3
TiO2@MnCo2O4.5 (6 wt%)	2	30	30	30	n.a./290	5	225	0.01	120	6.2	325	30	79	n.a./76	125
	5.1	150	30	3		6.4	300	0.01	30
TiMn2 (10 wt%)	5.1	225	10	2	n.a.	5.1	225	0.2	8	5.8	225	414	100	19/83	126
TiO2–ZrO2–V2O5 (10 wt%)	5.1	50	40	30	220/n.a.	6.5	300	0.001	5	6.4	300	50	98	n.a./66	127
Ti2VC2 (10 wt%)	3.2	25	60	60	170/n.a.	4.5	225	0.1	60	6.2	325	100	98	n.a./71	128
	5.5	150	60	10		6.2	300	0.1	5
Ni/Mo2N (6 wt%)	5.1	125	20	10	186/n.a.	5.9	265	—	30	6.7	330	10	98	n.a./76	129
(Mg2Ni)	5.5	150	20	15		6.8	325	—	5
Ni3Fe/Ni/NiFe2O4@C (10 wt%)	4.2	100	30	10	n.a./290	3.6	250	0.2	60	6.4	300	10	100	n.a./91	130
	5.8	150	30	1		5.8	325	0.2	5
Nb2O5–C/Ti3C2 (10 wt%)	5.5	100	40	15	n.a.	5.8	300	0.001	3	6.3	300	50	99	n.a./63	131
MoO3−x (10 wt%)	4.6	175	30	8	200/n.a.	5.9	300	0.05	8	6.1	300	50	99	n.a./71	132
Ti3C2@NiO (9 wt%) (Mg2Ni, Ti)	4	100	30	10	184/n.a.	6.7	300	0.01	10	6.7	300	20	96	n.a./53	133
	5.2	150	30	10
Cr,Mn,Fe,Co,Ni–oxide (10 wt%)	5.2	75	30	50	184/n.a.	6.7	250	n.a.	12	6.5	250	50	98	17/55	134
NiSe2 (5 wt%) (MgSe, Mg2NiH4)	5.2	100	30	80	n.a./319	6	300	0.2	20	6.2	360	20	100	n.a./71	135
	5.7	200	30	4
Ni/TiO2 (7 wt%) (Mg2Ni)	5.4	150	25	1	181/n.a.	6	250	0.001	30	5.5	200	20	95	n.a./78	136
						6.3	300	0.001	5
Ce0.6Zr0.4O2 (7 wt%) (CeH2.73/CeO2−x)	4.4	100	50	5	201/n.a.	3.9	230	0.001	40	5.8	270	20	99	n.a./69	137
	6	150	50	3		6.2	270	0.001	8
Co3V2O8 (6 wt%) (Co, V2O3)	1.3	RT	30	30	n.a.	3.4	225	0.09	120	6	300	30	73	n.a./69	138
	3.9	100	30	30		5.9	300	0.09	5
	5.9	300	30	1		6.4	325	0.09	3
NaNbO3 (5 wt%) (NbHx, Nb, NaMgH3)	3.5	30	30	90	n.a.	6.2	260	0.01	12	6.3	300	40	100	n.a./85	139
	4.7	125	30	5
Al3V (5 wt%)	6.3	300	50	10	291/n.a.	6.5	400	0.001	6	6	400	20	97	n.a./86	140
Ti0.9Zr0.1Mn1.5V0.3/CNT (10 + 1 wt%)	2.4	25	30	60	195/n.a.	6.1	300	0.001	5	6.2	325	100	97	n.a./86	141
	4.8	150	30	10
NiS/C (5 wt%) (Mg2Ni, MgS)	3	180	50	60	235/n.a.	3.5	300	0	60	6.5	400	20	97	24/67	142
Ni/Mo2C@C (7.5 wt%) (Mg2Ni)	n.a.	n.a.	35	n.a.	240/318	6.4	300	0.001	20	6.5	300	20	98	n.a./97	143
Ni3Mn (9 wt%) (Mg2Ni)	3.5	75	30	60	190/n.a.	4.4	235	0.03	60	6.6	295	20	94	39/87	144
	6.1	150	30	10		6.6	295	0.03	5
CuNi (10 wt%)^c (Mg2Ni(Cu))	6.1	265	10	10	n.a./245	6.1	277	n.a.	15	6.1	277	15	95	n.a./81	146

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Each step of the multistep hydrogenation and dehydrogenation processes must be addressed for effective catalysis, including bond dissociation, hydrogen transfer, bond formation, phase transformation, and heat transfer. In addition, uniform distribution and prevention of phase separation of catalysts and controlling the surface area of MgH₂/Mg particles by blocking large particle growth and aggregation are also of key importance. In situ formed hydride/metal or hydride/alloy couples that more easily release and accept hydrogen than MgH₂/Mg in the material matrix can serve as “hydrogen pumps” by removing/delivering hydrogen to MgH₂/Mg; examples are Mg₂Co/Mg₂CoH₅,¹¹⁴ VH₂/V,¹¹⁶ Mg₂Ni/Mg₂NiH₄,^{117,123,129,130,133,135,136,142–144} and ZrMn₂/Zr₄Mn₈H_10.99 (Fig. 4b).¹¹⁸ Another “hydrogen pump” mechanism is found to be the facilitation of electron transfer between Mg²⁺ and H⁻ by transition metal compounds, which promotes the dissociation and recombination of H₂, observed e.g. for TiFe,¹²² and multivalent transition–metal related species formed during hydrogen cycling from Nb₂CTx MXene,¹²⁴ TiO₂@MnCo₂O_4.5,¹²⁵ Nb₂O₅–C/Ti₃C₂Tx MXene,¹³¹ Ni/TiO₂,¹³⁶ Ce_0.6Zr_0.4O₂,¹³⁷ or NaNbO₃.¹³⁹ In addition, transition metals promote the dissociation of H₂ molecules¹²¹ and rupture of the Mg–H bond by weakening the interaction and elongating the bond length (see e.g., the ZrMn₂–MgH₂,¹¹⁸ Fe–MgH₂,¹²⁰ Co–H₂,¹³⁸ V₂O₃–MgH₂,¹³⁸ Ti_0.9Zr_0.1Mn_1.5V_0.3–MgH₂ interactions¹⁴¹). Oxygen vacancies of transition metal oxide catalysts have been found to have a promoting effect on dehydrogenation/rehydrogenation by accelerating electron transfer and creating a “hydrogen pump” (e.g. V₂O₅)¹¹⁶ or by enhancing the hydrogen adsorption ability of the oxide (e.g. MoO_3−x).¹³² There are multiple effects and processes during hydrogenation and dehydrogenation of magnesium, and to utilise all of these, synergistic use of multiple catalytic phases is required. The in situ generated phases can collectively constitute an efficient catalytic system. For example the coexistence of Mg₆MnO₈, Mn, Ti₂O₃, and CoO phases generated in situ from TiO₂@MnCo₂O_4.5¹²⁵ exhibits abundant diffusion pathways and nucleation sites for hydrogen absorption and desorption reactions.¹²⁵ Another example is the Ni₃Fe/Ni/NiFe₂O₄@C catalyst system where the in situ formed Mg₂Ni/Mg₂NiH₄ acts as a ‘‘hydrogen pump”, Fe⁰ facilitates electron transfer and creates hydrogen diffusion channels, and carbon effectively prevents particle agglomeration.¹³⁰ The catalyst Mg/MgH₂ interfaces play a key role in hydrogen transfer. Catalysts can cause morphological defects and could boost the nucleation by reducing the nucleation energy of Mg/MgH₂ and thereby improve the hydrogen diffusion rate.¹¹⁵ 2D nanomaterial sheets and layered materials, e.g. MXenes, are effective in creating abundant active sites and enabling faster hydrogen boundaries diffusion along the interfaces.^120,128 They are also effective in hindering powder agglomeration.¹³⁴

It is becoming clearer that effective hydrogen storage in magnesium requires both addressing each step of the multistep hydrogenation/dehydrogenation processes by benefiting from the synergistic catalytic effect of multiple phases of multiple compounds and limiting the growth of Mg and MgH₂ particles. Developing composite materials seems to be the best strategy to achieve long term cyclic goals. Significant, 50–70%, reduction of the overall activation energies could be achieved in dehydrogenation/rehydrogenation considering recent efforts; good performers, also tested for cycling, are provided in Table 5. It is now possible to recharge Mg below 100 °C; however, hydrogen release requires much higher temperatures (Table 5). Although the reduction of operation temperatures is a fundamental aim of research efforts, the excellent cycling stability is also crucial for practical applications; therefore, recent research is increasingly focusing on cycling stability (Table 5, Fig. 4c). To date, cycling tests have been conducted at relatively high temperatures between 275 and 400 °C due to slow hydrogen release at lower temperatures. Capacity retention close to 100% could be achieved for cycles ranging from 10 up to 414 (Table 5), but long-term cyclic stability evaluation as per DOE requirements is yet to be conducted. The high temperature of the hydrogen release of MgH₂, however, remains problematic for practical applications; therefore the Mg/MgH₂ hydrogen storage system still cannot work in the operational range of PEM fuel cells. Its application is currently limited to high-temperature fuel cells or stationary applications, or the released hydrogen must be cooled for PEM fuel cells. An example for high-temperature MgH₂-based storage is the application of the 0.9MgH₂–0.1TiH₂–5 wt% C nanocomposite in a storage tank which could deliver 185 normal litre hydrogen in temperature swing operation involving charging at 150 °C and discharging at 370 °C.¹⁴⁵

It is important to note the recent novel concept of solar-driven reversible hydrogen storage for magnesium, which integrates the photothermal localized heat and catalytic action by a single-component phase that simultaneously exhibits photothermal and catalytic effects.¹⁴⁶ Incorporating 10 wt% NiCu into Mg resulted in the formation of Mg₂Ni(Cu) which acted as a Mg₂Ni(Cu)/Mg₂Ni(Cu)H₄ hydrogen pump in cycling and also weakened the Mg–H bond. Reversible hydrogen storage was achieved with 6.1 wt% capacity and 95% retention.¹⁴⁶ The photodecomposition and rehydrogenation of the LiH surface under illumination at ambient temperature was also observed recently (dehydrogenated LiH could uptake 0.08 wt% hydrogen at 1 bar hydrogen pressure),¹⁴⁷ highlighting a new path for the development of hydrogen storage methods.

Among Li, Mg, and Al, which exhibit hydrogen storage properties, currently, only magnesium has the potential to become a reversible hydrogen storage material. Although it is still not satisfying all DOE targets, it seems to be on track to meet long-term cycling goals under close to ambient conditions. Strategies that might enable this are expected to use multielement and multiphase systems which can address kinetics of all steps of the multistep hydrogenation and dehydrogenation process and can control phase transfer and growth. A key advantage of magnesium is that it is abundant, non-poisonous, recyclable and a relatively cheap material. Mg is produced on a large scale in the smelting industry (1 million tonnes in 2024)¹⁴⁸ and can be hydrogenated directly in a rotary autoclave at 300–400 °C and 100–150 bar hydrogen pressure, or in THF solvent at 60–70 °C and 80 bar in the presence of catalysts.¹⁴⁹ A major disadvantage is its sensitivity to oxygen and water, with Mg nanopowders reacting violently with air and presenting the hazard of dust explosions.¹⁵⁰ Pure and finely divided MgH₂ is also pyrophoric.¹⁴⁹ The application of Li as a hydrogen storage material is not practical due to the high temperatures required for hydrogen absorption and release and competing demand for Li for batteries. LiH is produced in the industry by reacting molten lithium with hydrogen at temperatures above 700 °C.¹⁴⁹ Although Al is highly abundant and AlH₃ releases hydrogen upon moderate heating (temperatures below 100 °C if catalysed), regeneration of AlH₃ is not feasible by the direct Al + H₂ reaction due to thermodynamic instability. AlH₃ is currently not produced in industry; the main small scale laboratory production methods are based on the reaction of LiAlH₄ with AlCl₃ in a diethylether solvent (Scheme 1).^149,151 The energy intensive nature of the desolvation process and conversion of the LiCl by-product to lithium limits the potential of this method for large scale production. An alternative to this synthetic route could be the thermal decomposition of Al(Et)₃ (triethylaluminium), produced on an industrial scale from Al, H₂ and ethylene in a two-step process;¹⁵² however, the current synthesis at the laboratory scale is still suffering from low yields (19%).¹⁵³ Due to the lack of convenient synthesis/regeneration of AlH₃, currently, AlH₃ cannot compete with other potential one-way storage systems such as ammonia and LOHCs. In addition, the handling of Al and AlH₃ nanopowders is more problematic compared to that of LOHCs, as these nanopowders present the same dust explosion hazards as finely divided/powdered magnesium.

Scheme 1 Synthesis of AlH₃ and LiH (industrial processes are indicated by the boxes).^149–153

2.3. Alkali and alkali earth metal tetrahydroborates

Tetrahydrobotates exhibit higher hydrogen storage capacity than metal hydrides due to the light boron atom constituent but have high decomposition temperatures (above 400 °C), the kinetics of hydrogen desorption are sluggish, and the multi-step decomposition reaction is not or only partially reversible and only at high hydrogen pressures.^72,154–167 In addition, the tetrahydroborate decomposition may generate stable intermediates, such as borate cluster anions, or volatile by-products, like diborane,^72,167 and the evaporation of volatile alkali metals, namely Na and K, is an issue for practical applications. The latter issue might explain the paucity of experimental work on the thermal decomposition of sodium and potassium tetrahydroborates. Alkali and alkali earth tetrahydroborates, therefore, are not suited for reversible hydrogen storage without modifying the thermodynamics and kinetics of hydrogen absorption/desorption and establishing lower operating temperatures and reversibility for these reactions. It is, however, an advantage of tetrahydroborates compared to metal hydrides that they can tolerate larger amounts of additives/catalysts to satisfy DOE targets due to larger hydrogen storage capacities; the amount of additive in an additive + LiBH₄ mixture can be as high as 75 wt% to satisfy the DOE target of 6.5 wt% hydrogen storage capacity, but it is only 13 wt% for KBH₄ (see Tables 1 and 2). The possible higher additive load and storage capacity therefore opens the door for storage techniques that are not feasible for the Mg/MgH₂ system considering the DOE target, such as combining (nano)confinement and catalysis or applying only partial dehydrogenation by limiting the temperature and stopping before full decomposition, typically neglecting the final high temperature hydrogen release. An example is LiBH₄ decomposition due to the formation of thermally stable LiH whose hydrogen release proceeds only at very high temperatures (LiBH₄ → LiH + B + 3/2H₂).

The role of additives/catalysts is multiple in the tetrahydroborate systems as they should decrease the overall activation energy barrier of decomposition and hydrogen absorption, alter thermodynamics, and prevent the formation of stable intermediates and the separation of elemental boron by establishing an alternative lower energy pathway to metal borides (which are thermodynamically more stable reaction products than elemental boron). Considerable work has been dedicated to date to tailor the hydrogen absorption and desorption properties of tetrahydroborates using transition metal oxides, halides, and carbides, transition, rare earth and main group metals, and carbon-based materials.^46,167–210 Achieving reversible cycling is seemingly the most challenging (Table 6). The formation and accumulation of stable intermediates, for example Li₂B₁₂H₁₂, loss of nanoconfinement effects due to increasing particle size, and segregation of B and metal hydride phases are possibly the main reasons for the decrease in cyclic hydrogen capacity.^185,186,189 The mechanism of catalytic action of various additives may involve weakening B–H bonds,^{181,186,188,191,193} facilitating hydrogen diffusion,^{181,187,194,200,201,204–206} inhibiting B₁₂H₁₂²⁻ formation,^{185,188,191,192,194,198} suppressing the formation of B₂H₆,^192,203 enhancing the effect of nanoconfinement,^185,186 providing nucleation sites,^189,204 establishing B reservoirs by boride formation,¹⁹⁷ and limiting particle size growth.¹⁹¹ Carbonaceous materials, e.g. CNTs, can facilitate heat management as they effectively increase the thermal conductivity of the confined system.¹⁸⁷ As a new approach, the light-driven dehydrogenation of LiBH₄, without thermal heating, was achieved recently by incorporating TiO₂ for the photocatalytic effect and TiF₃ for enhanced dehydrogenation; the LiBH₄–8TiF₃–2TiO₂ composite could release about 3.12 wt% hydrogen in 30 min under a light intensity of 0.76 W cm⁻².²⁰⁹ LiBH₄ nanoconfined in ultralow-content Ni-doped structured mesoporous carbon could release 11.02 wt% hydrogen of LiBH₄ (not considering the template matrix) in 150 min under 3.5 W cm⁻² light irradiation and maintain a capacity retention of 85% after 5 cycles due to the nanoconfinement effect, trace amounts of Ni, and photogenerated vacancies in the LiH product which thermodynamically destabilize B−H bonds.²¹⁰ Transition metals and alloys, such as Ti,^201,204,209 Ni,^192,205 and CoNi,¹⁹³ and transition metal compounds, such as NbB₂,^186,202 LiTiO₂ and TiB₂,¹⁸⁷ FeB and Fe₂B,^188,190 Ni₂B,¹⁹¹ Li₃BO₃ and NbH,¹⁹⁴ Co₂B/Ni₂B,¹⁹⁷ V_xB_y and Ni_xB_y,¹⁹⁸ VH_2.01,²⁰¹ TiH₂,²⁰³ TiH_1.924, TiB, and TiB₂,²⁰⁴ and NbVH,²⁰⁶ generated in situ from oxide, carbide and halide precursors, as well as additives, such as MXenes (Ti₃C₂, Ti₂C, Nb₂C),¹⁸¹ were found to destabilize the BH₄⁻ anion by weakening the B–H bond and decreasing thereby hydrogen dissociation energy for hydrogen desorption due to electronic interactions/electron transfer. NbVH nanoparticles, for example, operated as “bidirectional hydrogen pumps” combining both hydrogen spillover and hydrogen diffusion.²⁰⁶ The bidirectional charge-transfer catalytic mechanisms of FeB and Fe₂B were revealed wherein they act as electron acceptors accepting electrons from BH₄⁻ and electron donors donating electrons to boron, which enhances their reactivity for dehydrogenation and hydrogenation, respectively.^188,190 In addition to kinetic facilitation of dehydrogenation/hydrogenation processes, in situ formed borides contribute to thermodynamic tuning of the system. An example is the reversible transformation between NiB/CoB and Ni₂B/Co₂B, acting as boron reservoirs, in the dehydrogenation and hydrogenation reactions of NaBH₄.¹⁹⁷ A problematic side product of tetrahydroborate decomposition is B₂H₆, a poisonous and flammable gas, which results in boron loss and/or in the formation of a stable B₁₂H₁₂⁻ dodecaborate anion via the B₂H₆ + BH₄⁻ reaction at elevated temperatures.¹⁹² The accumulation of B₁₂H₁₂²⁻ or boron loss during cycling gradually decreases storage capacity. Strategies to prevent these reactions include lowering the temperature below the temperature required to generate B₂H₆ (below 300 °C)¹⁹² or catalytically decomposing B₁₂H₁₂⁻. The nanoconfinement effect has also been found to suppress the release of B₂H₆ and inhibit B₁₂H₁₂⁻ formation during the dehydrogenation process.^185,210 Concerning the stable dodecaborate anion, reactive lithium ferrites and catalysis from iron borides were shown to effectively promote the decomposition of Li₂B₁₂H₁₂.¹⁸⁸ In situ formed Ni_xB or the synergistic action of Li₃BO₃ and NbH were also found to inhibit the formation of Li₂B₁₂H₁₂ during hydrogenation.^194,210 Carbonaceous materials (e.g., graphene, CNTs) and MXenes have multiple roles, as they provide structural support, interfaces for electron conduction and hydrogen diffusion, and also have a catalytic effect on hydrogenation and dehydrogenation processes.^181,201 The decomposition of LiBH₄, for example, is believed to preferably occur at the graphene–LiBH₄ interface, than within bulk LiBH₄, as an interface reaction.¹⁸⁵ Graphene was also shown to weaken both the Mg–H and B–B bonds of MgH₂ and B₂H₆, respectively.²⁰⁰ Nanoconfinement shortens the hydrogen diffusion pathway, and this effect is believed to be responsible for the decrease of activation energy.¹⁸⁵ Table 6 lists recent examples, focusing on those where hydrogen cycling has also been investigated, and highlights challenges for the application of alkali and alkali earth tatrahydroborates in hydrogen storage, namely the high temperature for hydrogen absorption/desorption and poor cycling stability. Obviously, the utilization of the synergistic effects of catalysts, nanoconfinement, improved surface interaction between various phases in the systems, and enhanced thermal transfer has already resulted in marked improvement of hydrogen storage properties. Fig. 5 shows, as an example, the fabrication of a multi-metal catalyst system for enhancing hydrogen release from NaBH₄ and the energy diagram illustrating altered thermodynamics and kinetics compared to pristine NaBH₄. There is a large variability for combining catalyst, confinement, and support; the most promising results, in terms of cycle number and capacity retention, are obtained by the nano-synergy of carbon materials and transition metal compounds, such as nitrogen-doped carbon nanocages/NbF₅,¹⁸⁶ few-layer graphene,¹⁸⁹ Ni nanocrystal decorated graphene (which has been reported to have the best performance in terms of cycle number),¹⁹² and bimetallic (NiCo) nanoporous carbon nanosheets.¹⁹⁷ Nanocarbon provides a platform for heat and electron transport, structural support, and interface for catalysts and BH₄⁻ ions. Transition metals and in situ formed transition metal borides catalyse the hydrogen uptake and release.

Table 6 Effect of catalysts on the hydrogen absorption and desorption kinetics of tetrahydroborates^a

Catalyst (active form)	Isothermal absorption				Non-isoth. des. onset/peak	Isothermal desorption				Cycles/capacity retention				Ea (des)	Ref.
	wt%	°C	bar	min	°C/°C	wt%	°C	bar	min	wt%	°C	No.	%	kJ mol^−1 H2
a RT = room temperature, GMF = wrinkled graphene microflowers, FL-GR = few-layer graphene, GR = graphene, NS = nanosheet, CNT = carbon nanotube, NC = nitrogen-doped carbon nanocages, vac = vacuum, n.a. = not available.b Desorption and absorption at 320 and 400 °C, respectively.c Desorption and absorption at 400 and 450 °C, respectively.d Desorption and absorption at 350 and 400 °C, respectively.e Desorption and absorption at 400 and 500 °C, respectively.f Dehydrogenation and rehydrogenation at 550 and 400 °C, respectively.
LiBH4
GMF (30 wt%)	5.8	400	100	55	n.a./327	8.0	320	vac	100	8.0	^b	10	54	107	185
NC-NbF5 (50 wt%) (NbB2)	6.3	300	120	96	141/311	5.9	280	vac	159	6.3	300	20	93	83	186
						7.6	320	vac	39
(CNT@C@TiO2) (40%) (LiTiO2, TiB2)	6.8	400	100	300	224/318	7.3	320	vac	60	8.4	^b	20	61	96	187
Fe3O4 (30 wt%) (FeB, Fe2B)	6.4	450	100	40	183/n.a.	7.9	375	vac	100	6.4	^c	10	70	83	188
						8.3	400	vac	50
FL-GR (30 wt%)	7	350	100	400	230/359	5.5	300	vac	400	7.1	350	10	100	116	189
						7.2	350	vac	60
FeF2/FeOX@GR (30 wt%) (FeB, Li3BO3, Fe2B)	9.4	400	100	150	100/375	7.0	350	vac	80	9.4	^c	10	59	99	190
Ni/C (20 wt%) (Ni2B)	n.a.	n.a.	n.a.	n.a.	200/340	12.2	320	0.01	180	8.9	320	15	72	99	191
Ni@GR (10 + 20 wt%)	9.2	300	100	200	130/285	7.5	200	vac	600	9.2	300	100	92	106	192
						9.2	300	vac	175
CoNi/C (40 wt%)	11.9	370	100	50	130/355	11.9	350	vac	100	9.4	^d	10	82	95	193
Nb(OEt)5 (4 mol%) (Li3BO3, NbH)	6.3	350	50	60	190/340	4.5	350	vac	20	7.9	^e	30	91	127	194
	7.9	500	50	20		7.0	350	vac	60
Ni/GR (20 wt%)	n.a.	n.a.	n.a.	n.a.	180/275, 465	12.8	450	1	45	15.2	^f	30	64	n.a.	195
NaBH4
GdF3 (25 mol%)	3.5	400	40	90	112/416	n.a.	n.a.	n.a.	n.a.	3.5	400	4	100	176	196
NiCo/C (50 wt%) (Ni2B, Co2B)	4.5	350	60	600	308/356	4.2	440	vac	60	4.5	400	5	100	37	197
V, Ni (VxBy, NixBy)	0.6	350	80	500	50/355	n.a.	n.a.	n.a.	n.a.	0.6	350	2	6	n.a.	198
KBH4
ZrCl4 (10 wt%)	n.a.	n.a.	n.a.	n.a.	350/500	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	162	199
Mg(BH4)2
GR (40 wt%)	n.a.	n.a.	n.a.	n.a.	154/199	5.7	175	n.a.	80	9.0	250	6	36	29	200
						9.0	250		50
Ti2C (30 wt%) (Ti)	10.2	300	100	480	132/281	10.2	260	n.a.	360	10.2	260	3	32	42	181
VF4@Ti3C2 (20 wt%) (VH2.01, Ti)	2.7	275	90	360	90/n.a.	5	275	n.a.	60	2.7	275	4	65	172	201
						8.2	275	n.a.	300
NaCl + NbF5 (48 + 10 wt%) (NbB2)	n.a.	n.a.	n.a.	n.a.	135/n.a.	5.6	500	n.a.	250	5.6	500	3	71	n.a.	202
K2TiF6 (3 mol%) (KBH4, TiH2)	n.a.	n.a.	n.a.	n.a.	105/n.a.	6.4	280	n.a.	200	6.4	280	2	36	n.a.	203
K2NbF7 (3 mol%) (KBH4, TiH2)	n.a.	n.a.	n.a.	n.a.	118/n.a.	6.6	280	n.a.	200	6.6	280	2	41	n.a.	203
Ti (TiH1.924, TiB, TiB2)	n.a.	n.a.	n.a.	n.a.	n.a.	7.9	290	n.a.	2000	6.2	270	4	48	57	204
Ni/CNT (5 wt%)	n.a.	n.a.	n.a.	n.a.	93/n.a.	6.7	300	4	60	n.a.	n.a.	n.a.	n.a.	120	205
NbVH/Ti3C2 (30 wt%)	n.a.	n.a.	n.a.	n.a.	106/n.a.	9.4	230	vac	1800	6.0	250	10	70	104, 125	206
						7.4	270	vac	60
Ca(BH4)2
Ti(OEt)4 (10 mol%) (TiO2)	4.7	320	90	1400	n.a./318	4.7	270	n.a.	400	n.a.	n.a.	n.a.	n.a.	118	207
TiCl3 (24 mol%) (Ti)	1.7	300	90	500	200/260	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	n.a.	208

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Fig. 5 (a) Illustration of the synthesis of vanadium-doped cube-like NaBH₄ with nickel shells, (b) the hydrogen release via the electron transfer process, and (c) the energy diagram showing altered thermodynamics and kinetics compared to pristine NaBH₄. Reproduced from ref. 198 with permission from the American Chemical Society, copyright 2023.

Currently, in general, tetrahydroborates underperform compared to the Mg/MgH₂ system in hydrogen storage due to factors like higher temperature and hydrogen pressure needed for recharging, slower kinetics (longer charging and discharging time) and poor hydrogen cycling stability (fewer cycles and lower capacity retention) (compare Table 5 with Table 6 and see Fig. 6 for selected data). Further improvements to establish reversible long-term cycling at lower than the current operating temperatures are necessary for practical applications. The advantages of the use of tetrahydroborates, compared to the Mg/MgH₂ system, are the higher hydrogen storage capacity per unit weight, larger additive tolerance concerning DOE targets, tunability through chemical design, and safety. Their stability in air increases as the ionic character of the salt increases, namely increasing stability from Li toward K and from Mg to Ca.¹⁴⁹ KBH₄ is stable in dry air and not hygroscopic. NaBH₄ is stable in dry air and can be stored in an airtight container for a long time. It absorbs water from moist air and decomposes slowly with the liberation of hydrogen. LiBH₄ is extremely hygroscopic and decomposes rapidly in moist air.¹⁴⁹ Mg(BH₄)₂ is hygroscopic and hydrolyses in moist air.²¹¹ Ca(BH₄)₂ is stable in dry air and slowly hydrolyses in water. An application advantage of NaBH₄ is that it is produced on a large scale in the industry (global production reached around 826 000 tons in 2023),²¹² e.g., from NaH and B(OCH₃)₃.¹⁴⁹ All other tetrahydroborates can be synthesized by metathesis reaction from NaBH₄ (Scheme 2).^{149,211–215}

Fig. 6 Relationship between temperature and time for (a) hydrogen uptake of dehydrogenated systems, where uptake is higher than 6.0 wt% within 160 min, and (b) hydrogen release from investigated systems, where the amount of released hydrogen is above 6.5 wt% and release time is within 120 min.

Scheme 2 Synthesis of tetrahydroborates (industrial processes are indicated by boxes).^{149,211–215}

2.4. Alkali and alkali earth metal tetrahydroaluminates

Tetrahydroaluminates and tetrahydroborates are structural analogues, but tetrahydroaluminates have lower gravimetric hydrogen storage capacity due to the replacement of the light boron atom with the heavier aluminium. LiAlH₄ can tolerate a maximum 39 wt% of additive to fulfill DOE target requirements considering that all hydrogen is released from the storage material; these values are 13, 30, and 18 wt% for Na, Mg, and Ca-tetrahydroaluminates, respectively (see Tables 1 and 2). In general, tetrahydroaluminates decompose and release hydrogen at lower temperature than tetrahydroborates and the only volatile decomposition product is hydrogen, and the solid residue consists of metals and/or alloys. The last step of the multi-step hydrogen release is the decomposition of alkali or alkali earth binary hydrides, which is practically not accessible due to the high temperatures required. If hydrogen discharge is conducted below their decomposition temperatures, the maximum hydrogen release from Li, Na, Mg, and Ca-tetrahydroaluminates would be 8.0, 5.6, 7.0, and 5.9 wt%, respectively. This places NaAlH₄ and Ca(AlH₄)₂ on the list of unsatisfactory compounds for hydrogen storage, while LiAlH₄ and Mg(AlH₄)₂ can tolerate only 9% and 7% additive, respectively. One of the major challenges of practical application of tetrahydroaluminates in hydrogen storage is that regenerating dehydrogenated tetrahydroaluminates requires high temperatures and hydrogen pressures.^{46,171,173,178} LiAlH₄, Mg(AlH₄)₂ and Ca(AlH₄)₂ are metastable and their rehydrogenation is thermodynamically not favoured.^216,217 The reason for the thermodynamic instability of LiAlH₄ and Ca(AlH₄)₂ is the formation of stable intermediates, namely Li₃AlH₆ and CaAlH₅, respectively.^218,219 According to thermodynamic calculations, more than 1000 bar hydrogen partial pressure is required for the hydrogen absorption reaction of Li₃AlH₆ → LiAlH₄ and CaAlH₅ → Ca(AlH₄)₂ above room temperature.^218,219 Practical reversible hydrogen storage with LiAlH₄ and Ca(AlH₄)₂ is therefore not possible without thermodynamically stabilizing these materials (and/or destabilizing intermediates). As an attempt to achieve this, the thermodynamic stabilization of LiAlH₄ within the nanopores of N-doped ordered mesoporous carbon by lithium coordination to nitrogen binding sites was demonstrated recently.²¹⁶ The decomposition of LiAlH₄ bypassed the Li₃AlH₆ intermediate and more than 80% of LiAlH₄ could be regenerated under 1000 bar hydrogen pressure. Mg(AlH₄)₂ decomposes without the formation of a complex hydride intermediate; the first dehydrogenation step to MgH₂ and Al is exothermic, which makes the reverse process infeasible.²²⁰ It appears likely that this first dehydrogenation step consists of two events, the endothermic decomposition of Mg(AlH₄)₂ and a consecutive exothermic event which could be the re-crystallization of non-crystalline MgH₂. The second dehydrogenation step has a relatively low temperature of 282 °C due to the formation of the intermetallic phase Mg₂Al₃.²²⁰ Considering thermodynamic instability, LiAlH₄, Mg(AlH₄)₂ and Ca(AlH₄)₂ may be best applied as one-way hydrogen storage materials, similarly to AlH₃ discussed above (Table 7). This consideration leaves NaAlH₄ as the only realistic target for reversible hydrogen storage but only if full dehydrogenation/rehydrogenation can be realized (Table 8). This material has the advantage of consisting only abundant elements, namely Na and Al. To address challenges, a wide range of transition and rare earth metal compounds and carbon-based materials, including oxides, chlorides, fluorides, nitrides, and borides, have been investigated to promote lower reaction temperatures and improved kinetics for tetrahydroaluminates.^221–255 General methods are like those used for main group binary hydrides and tetrahydroborates, discussed above, namely nanoconfinement and using additives and catalysts. Catalyst precursors (see examples in Tables 7 and 8) play a critical role as in situ formed transition metals, e.g. Ti, Zn, and Ni,^224,244,248 transition metal oxides, e.g. Ti₂O₃, NaFeO₂, and NiO,^233,244,251 and Al-alloys, e.g. Al₃Zr,²²³ Al₃Ti,²²⁵ Al₃Ni,²³³ NiAl, Al₃Ti and Ni₂Al₃,²⁴⁵ AlCo,²⁴⁶ Al₁₃Co₄ and Al_0.9Ni_4.22,²⁴⁷ as well as synergic multivalent Ti-species (Ti⁰, Ti³⁺)^252,253 or Ni and Ti species (Ni, NiO, Ti, Ti₂O₃)²⁵¹ are believed to catalyse the hydrogen uptake and release. The mechanism of catalytic action involved weakening the Al–H bonds, e.g. NiTiO₃ and Ni,^251,255 enhancing hydrogen transfer by shortening the diffusion path,^{244,247,254,255} e.g. Ti shell,²⁵⁵ and/or providing active sites to promote the nucleation and growth of dehydrogenation products, e.g. NiAl, Al₃Ti, Ni₂Al₃,²⁴⁵ and AlCo.²⁴⁶ Ti₃C₂ was found to effectively interact with tetrahydroaluminates and destabilize the Al–H bond through interfacial charge transfer.²²⁴ Nanosizing and nanoconfinement are mainly used to control the hydride particle size, and their effect is kinetic rather than thermodynamic by creating nucleation sites and shortening diffusion paths. N-doped carbon, however, was shown to alter thermodynamics by Li adatom formation and charge redistribution between the metal hydride and the host.²¹⁶ As a new strategy of research, the concept of light-driven reversible hydrogen storage was validated recently on NaAlH₄, using the catalytic effect of TiO₂ and the photothermal effect of nanolayered carbon.²⁵⁶

Table 7 Effect of catalysts on the hydrogen release properties of LiAlH₄, Mg(AlH₄)₂, and Ca(AlH₄)₂

Catalyst precursor	Active form^a	Capacity^b (wt%)	Decomp. temp.^c onset/peak (°C)	Isothermal H2 release wt%/°C min^−1	Ea^c (des.) (kJ mol^−1 H2)	Ref.
a Catalytically active component formed during ball milling of precursors; ss = solid solution.b Experimentally determined maximum storage capacity.c Consecutive two dehydrogenation stages, n.a. = not available.d Ca(AlH4)2·2NaCl, synthesized by ball milling NaAlH4 and CaCl2.
LiAlH4
ZrF4 (10 wt%)	Al3Zr, LiF	n.a.	80, 130/219, 258	5.6/100/120	60, 80	223
Ti3C2 (5 wt%)	Ti, Ti^3+	6.5	59/n.a.	4.2/150/10	80, 100	224
				5.5/200/35
TiF3 (2 mol%)	Al3Ti	7.3	35/108	7.0/140/80	67, 88	225
				7.3/250/18
Ni0.6Zn0.4O (10 wt%)	NiO, Zn	5.6	124, 170/n.a.	3.1/90/120	73, 85	244
Ni–C@TiO2 (7 wt%)	NiAl, Al3Ti, Ni2Al3	7.2	56, 131/n.a.	6.3/200/20	59, 68	245
Co–B (5 wt%)	AlCo	7.7	n.a./135, 167	7.7/200/120	83	246
Co3O4@CoNi (7 wt%)	Al13Co4, Al0.9Ni4.22	6.9	n.a./85, 177	6.3/200/120	80, 90	247
Mg(AlH4)2
TiCl3 (1 mol%)	Ti	9	60/n.a	1.5/90/150	n.a.	248
				6/150/n.a.
No catalyst	—	9.0	125/300, 415	5.5/135/125	123	249
Ca(AlH4)2
TiCl3 (1 mol%)	n.a.	n.a.	n.a./231, 295	n.a.	n.a.	242
TiF3 (10 wt.)^d	Ti	n.a.	n.a./116, 218	n.a.	57	243

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Table 8 Effect of catalysts on the hydrogen absorption and desorption kinetics of NaAlH₄^a

Catalyst (active form)	Isothermal absorption				Non-isoth. des.^b onset/peak	Isothermal desorption				Cycles/capacity retention				Ea^b (abs/des)	Ref.
	wt%	°C	bar	min	°C/°C	wt%	°C	bar	min	wt%	°C	No.	%	kJ mol^−1 H2
a RT = room temperature, CNT = carbon nanotube, vac = vacuum, n.a. = not available.b Values for the first, second and third decomposition step are separated by comma.c Desorption at 160 °C and absorption at 120 °C.d Desorption at 170 °C and absorption at 120 °C.e Under the light intensity of 5 and 10 sun during absorption and desorption, respectively (1 sun = 100 mW cm^−2).
NiFe2O4 (3 mol%) (NaFeO2, Al3Ni)	3.7	150	50	30	93, 147, 316/n.a.	4.6	150	vac	60	3.5	150	5	95	—/54, 73, 113	233
CNT (8 mol%)	4.2	150	140	600	n.a.	3.3	160	vac	120	3.7	150	8	n.a.	n.a.	237
CeB6 (2 mol%)	4.9	120	120	20	n.a.	4.9	160	0.1	120	4.9	^c	15	92	—/69, 99	250
NiTiO3 (10 wt%)	3.7	60	100	90	89/139, 173	3.3	110	0.01	120	4.6	^d	5	95	12/104, 74	251
(Ni, NiO, Ti, Ti2O3)						4.7	170	0.01	30
C@TiO2/Ti3C2 (10 wt%)	n.a.	n.a.	n.a.	n.a.	85/120, 160	4.0	140	vac	13	5.0	200	10	93	—/72, 64	252
Ti3C2 (7 wt%) (Ti, Ti^3+)	4.3	80	100	180	100/140	3.2	100	vac	180	4.8	200	10	97	—/87. 88	253
	4.6	120	100	60		4.7	140	vac	100
CoTiO3 (10 wt%) (AlTi3, AlxCo)	n.a.	n.a.	n.a.	n.a.	n.a./135, 185	3.5	150	vac	30	n.a.	n.a.	n.a.	n.a.	—/86. 92	254
TiCl3 (Ti)	4.0	180	130	75	155/180, 250	4.0	180	0.1	150	4.0	180	3	100	n.a.	255
TiO2/C (10 wt%)^e	4	202	80	60	n.a.	4	202	vac.	7	4	202	10	85	n.a./56	256
CeCl3 (2 mol%)	4	120	100	10	n.a.	3.9	180	1.3	n.a.	3.9	180	36	100	n.a.	261

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Handling tetrahydroaluminates requires more caution than handling tetrahydroborates due to the flammability of tetrahydroaluminates when exposed to oxygen and humidity. Tetrahydroaluminates hydrolyse in moist air and react violently with liquid water.^149,257 Although they are stable and not pyrophoric in the absence of moisture under ambient conditions, they are considered highly explosive in air when finely divided. Considering the one-way hydrogen delivery of LiAlH₄, its performance is approaching that of AlH₃, although the kinetics of LiAlH₄ dehydrogenation are a little slower (compare Table 4 and Table 7 and see Fig. 6); production cost and sustainability are expected to be decisive in their application. LiAlH₄ is produced in the industry, which is not the case for AlH₃. In addition, producing AlH₃ from LiAlH₄ makes AlH₃ more expensive than LiAlH₄ (Scheme 1). Considering the economics, LiAlH₄ production can be improved by recycling the LiCl by-product as only 25% of Li is converted to LiAlH₄ (Scheme 3).¹⁴⁹ Although dehydrogenated LiAlH₄ cannot be re-hydrogenated in the solid phase, it can be hydrogenated in THF slurry to form an LiAlH₄·4THF adduct (ΔG = −1.65 kJ mol⁻¹ LiAlH₄);²⁵⁸ the latter can be desolvated in vacuum at a low temperature of 60 °C.^258–260 Dimethyl ether can be used at sub-ambient temperature for this conversion and vents naturally with the excess hydrogen.^259,260 Detailed studies for the hydrogen delivery of magnesium and calcium tetrahydroaluminates are yet to be undertaken. NaAlH₄ can reversibly deliver hydrogen in catalysed reactions (Fig. 7a, b and Table 8); however, the limited storage capacity, below the 6.5 wt% DOE target, and the cyclic instability are drawbacks for its application as a hydrogen storage material. It is worth noting that prototypes of hydrogen storage tanks using 0.087 kg,²⁶¹ 8 kg,²⁶² and 20 kg²⁵⁷ NaAlH₄ have been developed and material gravimetric capacity of 3.9 wt% (36 cycles) and 3.6 wt% (18 cycles), and system gravimetric capacity of 2 wt% (25 cycles) could be achieved, respectively. Industry, currently, is not producing NaAlH₄, Mg(AlH₄)₂ and Ca(AlH₄)₂ on a large scale; these materials can be synthesized using various laboratory techniques from elements, hydrides and aluminium halides.²¹⁷ Selected frequently used methods are summarized in Scheme 3.

Scheme 3 Synthesis of tetrahydroaluminates (industrial processes are indicated by a box, RT = room temperature).^{149,217,242,258}

Fig. 7 (a) Isothermal dehydrogenation cyclic curves of NiTiO₃ nanorod-doped NaAlH₄ and (b) cyclic capacity retention of undoped and NiTiO₃ nanorod-doped NaAlH₄. Reproduced from ref. 251 with permission from Elsevier B. V., copyright 2023. (c) and (d) Cyclic performance of the 2LiBH₄ + MgH₂/2LiH + MgB₂ system with and without Ti₃C₂ catalysts: (c) dehydrogenation curves and (d) cyclic curves. Reproduced from ref. 280 with permission from the American Chemical Society, copyright 2019.

2.5. Reactive hydride composites

There is a vast possibility to combine metal hydrides, tetrahydroborates, and tetrahydroaluminates to create RHCs, especially as not only two but several components in various ratios could be mixed (Table 9). The concept of RHCs is that components of this system mutually destabilize each other and enhance the hydrogen desorption/adsorption kinetics, but without losing hydrogen storage capacity as the additive/destabilizing agent is also a hydrogen storage material. The presence of a secondary phase introduces new phase boundaries where dehydrogenation can be preferential, compared to the pristine material, may initiate metathesis reaction between components or react to form intermediate phases altering thereby the reaction path, and modify the overall thermodynamics by alloy or metal boride formation. Experiments have confirmed lower onset decomposition temperature, faster dehydrogenation/rehydrogenation rate, and improved hydrogen yield for composites of tetrahydroborates, tetrahydroaluminates and metal hydrides compared to their pristine constituents.^263–289 In addition, nano-sizing, nanoconfinement, and catalysis have further improved hydrogen storage performances.^273–287 There are also challenges for RHC systems, such as the formation of stable intermediates, which block the reversibility of the dehydrogenation, phase separation of components during hydrogenation or rehydrogenation, or increased interdiffusion barriers, which must be prevented by using additives and catalysts. A simplified view of RHC hydrogen release considers RHC components confined in each other and consecutive decomposition steps; the first step is the decomposition of one of the components, forming fine and well dispersed reagent particles on the surface of the second component, and the second step is the dehydrogenation of the second component. The formation of alloys and borides can thermodynamically reduce the dehydrogenation enthalpy. An example is the 6LiBH₄ + Mg(AlH₄)₂ system, where Mg(AlH₄)₂ decomposes to MgH₂ and Al in the initiating step, MgH₂ and Al react further to form Mg₂Al₃, and Mg₂Al₃ and Al particles enhance the subsequent dehydrogenation of LiBH₄ to form MgAlB₄, AlB₂ and LiH.²⁷⁹ The products of this RHC decomposition, however, could be rehydrogenated only to LiBH₄ + MgH₂ + Al (due to the thermodynamic instability of Mg(AlH₄)₂ discussed above). The 2LiBH₄ + MgH₂ system, the second example, is a typical RHC system exhibiting a high storage capacity of 11.5 wt%.²⁸⁴ The RHC releases hydrogen in two steps, the first is the MgH₂ decomposition to Mg and H₂, and the second is the reaction of Mg with LiBH₄ to produce MgB₂ and LiH. The formation of MgB₂ provides thermodynamic stabilization for the system and is essential for the reversibility of hydrogen release. There is however a very long incubation period after the first step (more than 10 hours) which is attributed to the sluggish nucleation process of MgB₂.^284,290 Catalysing the nucleation by VB₂ nanoparticles (having a lattice constant close to that of MgB₂) resulted in an excellent cycling performance of 9.3 wt% over 10 cycles with 100% retention, and without the nucleation period.²⁸⁴ It was observed in a similar system that the formation of MgB₂ crystallites preferentially occurs on TiB₂ nanoparticle surfaces, generated in situ from a 3TiCl₃·AlCl₃ additive, due to reduced elastic strain energy density compared to that of the MgB₂/Mg interface.²⁹⁰ A similar action of NbB₂,²⁷⁵ MgNi₃B₂ and CoB²⁸¹ nanoparticles has also been confirmed, highlighting an important catalytic role of transition metal borides. In general, catalysts and additives have been confirmed to act as active “hydrogen pumps”,²⁷⁴ preventing aggregation of nanoparticles,^273,277 providing nucleation sites,^{274,275,280–284} enhancing diffusion,²⁸³ and destabilizing the BH₄⁻ anion by elongating the B–H bond.²⁸³ Nanoconfinement was found to inhibit the formation of B₁₂H₁₂²⁻ clusters.²⁷³ As an example, in situ generated Ni₂B nanoparticles were found to act as bidirectional electron transfer catalysts in tetrahydroborate systems, destabilize the tetrahydroborate anion and reduce the kinetic barrier for its formation.²⁷³ VH_2.01, formed from V₂C, exhibited multiple roles; it enhanced H₂ diffusion due to its open crystal structure, increased the hydrogen atom diffusion rate by providing pathways, and effectively weakened the B–H bonds and catalysed hydrogen release/absorption through a “hydrogen pump” effect.²⁷⁴ The effect of Ti₂C, V₂C and Ti₃CN MXenes was also found to be multiple as they prevented the aggregation of hydrides and generated active metal, e.g. Ti,²⁷⁷ metal hydride, e.g. VH_2.01,²⁷⁴ or metal boride, e.g. TiB₂,²⁸² species. The traditional focus of catalytic tuning of hydrogenation kinetics is on transition metal cations; however, the perspective of additional anion tuning was proposed recently by highlighting the role of S²⁻ anions in Li⁺ ion migration and destabilization of B–H bonds using TiS₂, MoS₂, or ZrS₂ for catalysing the dehydrogenation/hydrogenation of 2LiBH₄ + MgH₂.²⁸³

Table 9 Effect of catalysts on the hydrogen absorption and desorption kinetics of RHCs^a

Catalyst (active form)	Isothermal absorption				Non-isoth. Des.^b onset/peak	Isothermal desorption				Cycles/capacity retention				Ea^b (des)	Ref.
	wt%	°C	bar	min	°C/°C	wt%	°C	bar	min	wt%	°C	No.	%	kJ mol^−1 H2
a RT = room temperature, CNT = carbon nanotube, NC = nanoporous carbon, vac = vacuum, n.a. = not available.b Values for the first, second and third decomposition step are separated by comma.c Desorption at 300° and absorption at 400 °C.d Desorption at 400 °C and absorption at 350 °C.
LiBH4 + KBH4
NiCp2/C (40 wt%) (Ni2B)	6.7	400	100	120	n.a./238	7.0	300	vac	30	6.7	^c	50	94	101	273
LiBH4 + Mg(BH4)2
V2C (20 wt%) (VH2.01)	n.a.	300	100	360	80/280	4.5	275	vac	50	7.3	300	4	42	118, 125	274
						10.1	350	vac	50
Ti2C (30 wt%) (Ti)	n.a.	400	100	480	124/198	6.5	200		300	9.6	260	4	37	106, 95, 122	277
						9.6	260		300
KBH4 + 4LiAlH4
no cat.	6.5	420	33	60	82, 138, 484/n.a.	1.7	420	1	60	n.a.	n.a.	n.a.	n.a.	153	278
6LiBH4 + Mg(AlH4)2
No cat. (Mg2Al3, Al)	6.4	400	30	300	n.a.	11.8	400	1	200	n.a.	n.a.	n.a.	n.a.	n.a.	279
2LiBH4 + MgH2
Ti3C2 (5 wt%) (TiB2)	9.4	350	100	5	n.a.	9.5	400	4	30	9.0	^d	15	92	114	280
NiCo@NC (10 wt%) (MgNi3B2, CoB)	n.a.	n.a.	n.a.	n.a.	n.a.	5.5	380	4	360	9.4	400	10	96	n.a.	281
						9.4	400	4	270
Ti3CN (5 wt%) (TiB2)	9.4	350	75	5	236/342, 404, 435	9.0	400	2	60	8.9	^d	20	93	121, 119	282
TiS2 (10 wt%) (MgS, TiB2)	8.4	350	40	120	221/n.a.	8.7	425	4	60	9.3	420	50	91	82, 101	283
VB2 (10 wt%) Ca(BH4)2 + MgH2	9.3	350	60	n.a.	n.a.	9.23	400	4	120	9.3	^d	10	100	n.a.	284
No cat.	n.a.	n.a.	70	n.a.	339/n.a.	n.a.	n.a.	1	n.a.	9.1	400	10	32	n.a.	285
NbF5 (2 mol%)	n.a.	n.a.	70	n.a.	n.a.	n.a.	n.a.	1	n.a.	9.1	400	10	45	n.a.	285
6LiBH4 + CaH2
TiCl3 (1 mol%)	11.6	450	80	960	n.a.	11.6	450	1	240	11.6	450	10	61	n.a.	286
2Li[AlH4] + Mg[BH4]2
TiF3 (5 wt%) (MgF2, LIF, TiH2)	3.8	320	33	180	60/n.a.	1.7	320	1	60	n.a.	n.a.	n.a.	n.a.	80, 109, 137	287
2NaAlH4 + Ca(BH4)2
No cat.	3.5	420	30	60	125/n.a.	3.0	420	1	60	n.a.	n.a.	n.a.	n.a.	143, 147	288
Ca(BH4)2 + 2LiBH4 + 2MgH2
No cat. (CaMgH3.72)	7.6	350	90	1080	320/339	6.9	320	vac	110	8.1	350	10	77	n.a.	289

---

The concept of forming RHCs by combining two or more hydrides has been proven to work in practice and combining hydrides has significantly improved the hydrogen storage performance of these composites, compared to their pristine constituents (Fig. 7c, d and Table 9). Challenges, however, remain the same for practical applications as for their constituents, namely further reduction of the operating temperature, enhancement of the kinetics, and especially improving reversibility of the hydrogen release and cyclic stability (Fig. 6). The VB₂ catalysed 2LiBH₄ + MgH₂ system exhibits currently the best performance in terms of storage capacity and capacity retention during cycling (Table 9), which is the result of combining the relatively low decomposition temperature of MgH₂, the reactivity of Mg with LiBH₄, the thermodynamic stabilization by MgB₂ formation, and the catalytic activity of VB₂ nanoparticles on MgB₂ nucleation.²⁸⁴

3. Conclusion

Developing solid-phase hydrogen storage materials for practical applications is of high importance for boosting the hydrogen industry. Light main group metal hydrides (LiH, MgH₂, AlH₃), their tetrahydroborates (LiBH₄, NaBH₄, KBH₄, Ca(BH₄)₂ and Mg(BH₄)₂), and tetrahydroaluminates (LiAlH₄, NaAlH₄, Ca(AlH₄)₂ and Mg(AlH₄)₂), as well as their hydride composites have high gravimetric hydrogen storage capacities, capable of fulfilling the minimum storage capacity requirements targeted by the DOE. The practical application of these materials, however, suffers from high operating temperatures, poor reversibility, slow kinetics, and low cyclic performance. To address these issues, large efforts have been devoted to tune hydrogen storage properties and recent research has justified that enhanced hydrogen storage can be achieved by nanostructuring and nanoconfining storage materials and incorporating catalysts. It has been found that small particle size and interaction between solid particles are vital elements for these catalysed reactions to enhance hydrogen diffusion and mass transport, and preventing particle growth is a fundamental requirement for good cyclic performances. In addition, enhancement of storage properties is largely attributed to the in situ formation of active metal species or metal compounds during the synthesis and operation of storage material composites, as well as to high dispersion of catalytic species. It has become clear that multicomponent catalysts outperform conventional single catalyst systems due to their ability to address several steps of the multi-step processes of hydrogen release and absorption. Recent research has also highlighted that it is crucially important to prevent the crystallite growth and agglomeration, the separation of phases of storage components, and the redistribution of catalytic species or phases during hydrogen cycling. Incorporating additives for enhancing hydrogen and heat transport is also fundamentally important. All these requirements can only be fulfilled if the synergistic effect of multiple phases introduced into storage materials in situ and by mixing with additives is utilized.

Not all hydride and tetrahydroaluminate compounds are feasible targets for reversible hydrogen storage as AlH₃, LiAlH₄, Ca(AlH₄)₂ and Mg(AlH₄)₂ are metastable and their rehydrogenation is infeasible. They may serve as one-way hydrogen source media if their alternative economical synthesis is established. Their best performers can already operate with fast hydrogen release around and below 100 °C and could serve PEM fuel cells. Currently the Mg/MgH₂ system is the best performer concerning reversibility, kinetics and cyclic properties. MgH₂ can be discharged and recharged with relatively fast kinetics, but only at high temperatures of 300–400 °C and pressures of 20–80 bar, respectively. In addition, a stable cycling for at least 1500 cycles has not been achieved yet. The required pressure for recharging discharged tetrahydroborates and their composites is even higher (30–120 bar at 300–400 °C). The challenge for NaAlH₄ is the same and if only partially discharged the storage capacity is insufficient for DOE targets. Currently, none of the discussed hydrogen storage systems can fulfill the requirements of close to ambient working conditions, namely below 10 bar pressure and 100 °C. The improvement for practical application, however, has seen extensive progress and recently developed catalysed and confined systems operate at much lower temperatures and pressures, exhibit faster kinetics and much lower capacity decay during cycling than their corresponding pristine derivatives. There is still room for improvement as for example modifying template materials by doping to alter thermodynamics and anionic tuning to enhance kinetics has hardly been utilized. The novel concept of light-driven reversible hydrogen storage is validated recently on binary hydrides, tetrahydroborates and tetrahydroaluminates, and opens a new direction for the development of storage systems. There are reasons for optimism when it comes to solving storage issues and future research is expected to focus on exploiting the advantage of synergistic effects of catalysis using multiple catalysts, nanoconfinement, nanostructuring, and host material doping. Understanding the mechanism of hydrogen release/absorption is crucial for storage material design; therefore future work is expected to shift stronger toward full understanding of catalytic mechanisms by using high-resolution operando techniques. Research in this field will gradually move from storage materials toward developing engineered storage systems for real world applications. As current research is conducted at a small scale, scaling up for practical applications is expected to present new challenges especially in material and thermal transport.

Author contributions

Conceptualization and design: M. K.; acquisition of data: M. K. and T. P.; analysis and interpretation of data: M. K., T. P., S. S. and V. P. T.; writing – original draft: M. K. and T. P.; writing – review and editing: M. K., T. P., S. S. and V. P. T.; final version of the manuscript is approved by M. K., T. P., S. S. and V. P. T.

Conflicts of interest

The authors declare no conflicts of interest.

Data availability

No primary research results, software or code have been included, and no new data were generated or analysed as part of this review.

Acknowledgements

This research was funded by the Australian Renewable Energy Agency, grant number 2023/TRAC733 (PRO-1050).

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Footnote

† Current address: Research School of Chemistry, College of Science and Medicine, The Australian National University, Canberra, ACT 2601, Australia.

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