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Lithium hydride
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 Lithium cation, Li+ Hydrogen anion, H− | |
 __H− __Li+
Structure of lithium hydride. | |
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| Identifiers | |
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.028.623 |
PubChem CID
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| RTECS number |
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| UNII | |
CompTox Dashboard (EPA)
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| Properties | |
| LiH | |
| Molar mass | 7.95 g·mol−1 |
| Appearance | colorless to gray solid[1] |
| Density | 0.78 g/cm3[1] |
| Melting point | 688.7 °C (1,271.7 °F; 961.9 K)[1] |
| Boiling point | 900–1,000 °C (1,650–1,830 °F; 1,170–1,270 K) (decomposes)[2] |
| reacts | |
| Solubility | slightly soluble in dimethylformamide reacts with ammonia, diethyl ether, ethanol |
| −4.6·10−6 cm3/mol | |
Refractive index (nD)
|
1.9847[3]: 43 |
| Structure | |
| fcc (NaCl-type) | |
a = 0.40834 nm[3]: 56
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| 6.0 D[3]: 35 | |
| Thermochemistry | |
Heat capacity (C)
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3.51 J/(g·K) |
Std molar
entropy (S⦵298) |
170.8 J/(mol·K) |
Std enthalpy of
formation (ΔfH⦵298) |
−90.65 kJ/mol |
Gibbs free energy (ΔfG⦵)
|
−68.48 kJ/mol |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards
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extremely strong irritant, highly toxic, highly corrosive |
| GHS labelling: | |
  
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| Danger | |
| H260, H301, H314 | |
| P223, P231+P232, P260, P264, P270, P280, P301+P316, P301+P330+P331, P302+P335+P334, P302+P361+P354, P304+P340, P305+P354+P338, P316, P321, P330, P363, P370+P378, P402+P404, P405, P501 | |
| NFPA 704 (fire diamond) | |
| 200 °C (392 °F; 473 K) | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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77.5 mg/kg (oral, rat)[5] |
LC50 (median concentration)
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22 mg/m3 (rat, 4 h)[6] |
| NIOSH (US health exposure limits): | |
PEL (Permissible)
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TWA 0.025 mg/m3[4] |
REL (Recommended)
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TWA 0.025 mg/m3[4] |
IDLH (Immediate danger)
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0.5 mg/m3[4] |
| Safety data sheet (SDS) | ICSC 0813 |
| Related compounds | |
Other cations
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Sodium hydride Potassium hydride Rubidium hydride Caesium hydride |
Related compounds
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Lithium borohydride Lithium aluminium hydride |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |
Lithium hydride is an inorganic compound with the formula LiH. This alkali metal hydride is a colorless solid, although commercial samples are grey. Characteristic of a salt-like (ionic) hydride, it has a high melting point, and it is not soluble but reactive with all protic organic solvents. It is soluble and nonreactive with certain molten salts such as lithium fluoride, lithium borohydride, and sodium hydride. With a molar mass of 7.95 g/mol, it is the lightest ionic compound.
Physical properties
[edit]LiH is diamagnetic and an ionic conductor with an electric conductivity gradually increasing from 2×10−5 Ω−1cm−1 at 443 °C to 0.18 Ω−1cm−1 at 754 °C; there is no discontinuity in this increase through the melting point.[3]: 36 The dielectric constant of LiH decreases from 13.0 (static, low frequencies) to 3.6 (visible-light frequencies).[3]: 35 LiH is a soft material with a Mohs hardness of 3.5.[3]: 42 Its compressive creep (per 100 hours) rapidly increases from < 1% at 350 °C to > 100% at 475 °C, meaning that LiH cannot provide mechanical support when heated.[3]: 39
The thermal conductivity of LiH decreases with temperature and depends on morphology: the corresponding values are 0.125 W/(cm·K) for crystals and 0.0695 W/(cm·K) for compacts at 50 °C, and 0.036 W/(cm·K) for crystals and 0.0432 W/(cm·K) for compacts at 500 °C.[3]: 60 The linear thermal expansion coefficient is 4.2×10−5/°C at room temperature.[3]: 49
Synthesis and processing
[edit]LiH is produced by treating lithium metal with hydrogen gas:
- 2 Li + H2 → 2 LiH
This reaction is especially rapid at temperatures above 600 °C. Addition of 0.001–0.003% carbon, and/or increasing temperature/pressure, increases the yield up to 98% at 2-hour residence time.[3]: 147 However, the reaction proceeds at temperatures as low as 29 °C. The yield is 60% at 99 °C and 85% at 125 °C, and the rate depends significantly on the surface condition of LiH.[3]: 5
Less common ways of LiH synthesis include thermal decomposition of lithium aluminium hydride (200 °C), lithium borohydride (300 °C), n-butyllithium (150 °C), or ethyllithium (120 °C), as well as several reactions involving lithium compounds of low stability and available hydrogen content.[3]: 144–145
Chemical reactions yield LiH in the form of lumped powder, which can be compressed into pellets without a binder. More complex shapes can be produced by casting from the melt.[3]: 160 ff. Large single crystals (about 80 mm long and 16 mm in diameter) can be then grown from molten LiH powder in hydrogen atmosphere by the Bridgman–Stockbarger technique. They often have bluish color owing to the presence of colloidal Li. This color can be removed by post-growth annealing at lower temperatures (~550 °C) and lower thermal gradients.[3]: 154 Major impurities in these crystals are Na (20–200 ppm), O (10–100 ppm), Mg (0.5–6 ppm), Fe (0.5-2 ppm) and Cu (0.5-2 ppm).[3]: 155
Bulk cold-pressed LiH parts can be easily machined using standard techniques and tools to micrometer precision. However, cast LiH is brittle and easily cracks during processing.[3]: 171
A more energy efficient route to form lithium hydride powder is by ball milling lithium metal under high hydrogen pressure. To prevent cold welding of lithium metal (due to its high ductility), small amounts of lithium hydride powder are added during this process.[7]
Reactions
[edit]LiH powder reacts rapidly with air of low humidity, forming LiOH, Li2O and Li2CO3. In moist air the powder ignites spontaneously, forming a mixture of products including some nitrogenous compounds. The lump material reacts with humid air, forming a superficial coating, which is a viscous fluid. This inhibits further reaction, although the appearance of a film of "tarnish" is quite evident. Little or no nitride is formed on exposure to humid air. The lump material, contained in a metal dish, may be heated in air to slightly below 200 °C without igniting, although it ignites readily when touched by an open flame. The surface condition of LiH, presence of oxides on the metal dish, etc., have a considerable effect on the ignition temperature. Dry oxygen does not react with crystalline LiH unless heated strongly, when an almost explosive combustion occurs.[3]: 6
LiH is highly reactive towards water and other protic reagents:[3]: 7
- LiH + H2O → Li+ + H2 + OH−
LiH is less reactive with water than Li and thus is a much less powerful reducing agent for water, alcohols, and other media containing reducible solutes. This is true for all the binary saline hydrides.[3]: 22
LiH pellets slowly expand in moist air, forming LiOH; however, the expansion rate is below 10% within 24 hours in a pressure of 2 Torr of water vapor.[3]: 7 If moist air contains carbon dioxide, then the product is lithium carbonate.[3]: 8 LiH reacts with ammonia, slowly at room temperature, but the reaction accelerates significantly above 300 °C.[3]: 10 LiH reacts slowly with higher alcohols and phenols, but vigorously with lower alcohols.[3]: 14
LiH reacts with sulfur dioxide to give the dithionite:
- 2 LiH + 2 SO2 → Li2S2O4 + H2
though above 50 °C the product is lithium sulfide instead.[3]: 9
LiH reacts with acetylene to form lithium carbide and hydrogen. With anhydrous organic acids, phenols and acid anhydrides, LiH reacts slowly, producing hydrogen gas and the lithium salt of the acid. With water-containing acids, LiH reacts faster than with water.[3]: 8 Many reactions of LiH with oxygen-containing species yield LiOH, which in turn irreversibly reacts with LiH at temperatures above 300 °C:[3]: 10
- LiH + LiOH → Li2O + H2
Lithium hydride is rather unreactive at moderate temperatures with O2 or Cl2. It is, therefore, used in the synthesis of other useful hydrides,[8] e.g.,
Applications
[edit]Hydrogen storage and fuel
[edit]With a hydrogen content in proportion to its mass three times that of NaH, LiH has the highest hydrogen content of any hydride. LiH is periodically of interest for hydrogen storage, but applications have been thwarted by its stability to decomposition. Thus removal of H2 requires temperatures above the 700 °C used for its synthesis, such temperatures are expensive to create and maintain. The compound was once tested as a fuel component in a model rocket.[9][10]
Precursor to complex metal hydrides
[edit]LiH is not usually a hydride-reducing agent, except in the synthesis of hydrides of certain metalloids. For example, silane is produced in the reaction of lithium hydride and silicon tetrachloride by the Sundermeyer process:
- 4 LiH + SiCl4 → 4 LiCl + SiH4
Lithium hydride is used in the production of a variety of reagents for organic synthesis, such as lithium aluminium hydride (Li[AlH4]) and lithium borohydride (Li[BH4]). Triethylborane reacts to give superhydride (Li[BH(CH2CH3)3]).[11]
In nuclear chemistry and physics
[edit]Lithium hydride (LiH) is sometimes a desirable material for the shielding of nuclear reactors, with the isotope lithium-6 (Li-6), and it can be fabricated by casting.[12][13]
Lithium deuteride
[edit]Lithium deuteride, in the form of lithium-7 deuteride (7Li2H or 7LiD), is a good moderator for nuclear reactors, because deuterium (2H or D) has a lower neutron absorption cross-section than ordinary hydrogen or protium (1H) does, and the cross-section for 7Li is also low, decreasing the absorption of neutrons in a reactor. 7Li is preferred for a moderator because it has a lower neutron capture cross-section, and it also forms less tritium (3H or T) under bombardment with neutrons.[14]
The corresponding lithium-6 deuteride (6Li2H or 6LiD) is the primary fusion fuel in thermonuclear weapons.[citation needed] In hydrogen warheads of the Teller–Ulam design, a nuclear fission trigger explodes to heat and compress the lithium-6 deuteride, and to bombard the 6LiD with neutrons to produce tritium in an exothermic reaction:
- 6LiD + n → 4He + T + D
The deuterium and tritium then fuse to produce helium, one neutron, and 17.59 MeV of free energy in the form of gamma rays, kinetic energy, etc. Tritium has a favorable reaction cross section. The helium is an inert byproduct.[citation needed]
- 3
1H + 2
1H → 4
2He + n.
Before the Castle Bravo nuclear weapons test in 1954, it was thought that only the less common isotope 6Li would breed tritium when struck with fast neutrons. The Castle Bravo test showed (accidentally) that the more plentiful 7Li also does so under extreme conditions, albeit by an endothermic reaction.
Safety
[edit]LiH reacts violently with water to give hydrogen gas and LiOH, which is caustic. Consequently, LiH dust can explode in humid air, or even in dry air due to static electricity. At concentrations of 5–55 mg/m3 in air the dust is extremely irritating to the mucous membranes and skin and may cause an allergic reaction. Because of the irritation, LiH is normally rejected rather than accumulated by the body.[3]: 157, 182
Some lithium salts, which can be produced in LiH reactions, are toxic. LiH fire should not be extinguished using carbon dioxide, carbon tetrachloride, or aqueous fire extinguishers; it should be smothered by covering with a metal object or graphite or dolomite powder. Sand is less suitable, as it can explode when mixed with burning LiH, especially if not dry. LiH is normally transported in oil, using containers made of ceramic, certain plastics or steel, and is handled in an atmosphere of dry argon or helium.[3]: 156 Whilst nitrogen can be used, it will react with lithium at elevated temperatures.[3]: 157 LiH normally contains some metallic lithium, which corrodes steel or silica containers at elevated temperatures.[3]: 173–174, 179
References
[edit]- ^ a b c Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton, Florida: CRC Press. p. 4.70. ISBN 0-8493-0486-5.
- ^ David Arthur Johnson; Open University (12 August 2002). Metals and chemical change. Royal Society of Chemistry. pp. 167–. ISBN 978-0-85404-665-2. Retrieved 1 November 2011.
- ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad Smith, R. L.; Miser, J. W. (1963). Compilation of the properties of lithium hydride. NASA.
- ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0371". National Institute for Occupational Safety and Health (NIOSH).
- ^ Chambers, Michael. "ChemIDplus - 7580-67-8 - SIAPCJWMELPYOE-UHFFFAOYSA-N - Lithium hydride - Similar structures search, synonyms, formulas, resource links, and other chemical information". chem.sis.nlm.nih.gov. Retrieved 10 April 2018.
- ^ "Lithium hydride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- ^ Solvent-and catalyst-free mechanochemical synthesis of alkali metal monohydrides IZ Hlova, A Castle, JF Goldston, S Gupta, T Prost… - Journal of Materials Chemistry A, 2016
- ^ "NCERT Chemistry Textbook" (PDF).
- ^ Lex Archived 2008-07-23 at the Wayback Machine. Astronautix.com (1964-04-25). Retrieved on 2011-11-01.
- ^ Empirical laws for hybrid combustion of lithium hydride with fluorine in small rocket engines. Ntrs.nasa.gov. Retrieved on 2011-11-01.
(password-protected)
- ^ Peter Rittmeyer, Ulrich Wietelmann "Hydrides" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a13_199
- ^ Peter J. Turchi (1998). Propulsion techniques: action and reaction. AIAA. pp. 339–. ISBN 978-1-56347-115-5. Retrieved 2 November 2011.
- ^ Welch, Frank H. (February 1974). "Lithium hydride: A space age shielding material". Nuclear Engineering and Design. 26 (3): 440–460. doi:10.1016/0029-5493(74)90082-X.
- ^ Massie, Mark; Dewan, Leslie C. "US 20130083878 A1, April 4, 2013, NUCLEAR REACTORS AND RELATED METHODS AND APPARATUS". U.S. Patent Office. U.S. Government. Retrieved 2 June 2016.
External links
[edit]
Look up lithium hydride in Wiktionary, the free dictionary.
Lithium hydride
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History
Discovery and early development
Lithium hydride (LiH) was first prepared in 1896 by the French chemist Moïse Guntz through the direct reaction of lithium metal with hydrogen gas at elevated temperatures, yielding a reasonably stoichiometric white solid.[10] [11] Guntz's synthesis involved heating lithium in a stream of dry hydrogen, initially motivated by attempts to form lithium nitride but resulting in preferential hydride formation due to trace hydrogen absorption. This marked the initial empirical confirmation of LiH as a stable alkali metal hydride, distinct from less stable analogs like sodium hydride. Early research following Guntz's work focused on verifying the compound's composition and basic thermodynamic properties. Guntz and collaborator R. Benoit conducted calorimetric measurements, determining the heat of formation to be approximately 21.6 kcal/mol, underscoring LiH's relative stability compared to other hydrides.[12] Experiments in the late 19th and early 20th centuries, including thermal dissociation studies, established that LiH decomposes above 700°C but reacts vigorously with water to liberate hydrogen gas (LiH + H₂O → LiOH + ½H₂), highlighting its potential as a hydrogen source while noting handling challenges due to moisture sensitivity.[13] By the mid-20th century, amid increasing availability of bulk lithium metal, LiH transitioned from laboratory-scale preparation to preliminary industrial interest for its low density (0.82 g/cm³) and high hydrogen content (12.7 wt%), positioning it as a candidate for lightweight hydrogen storage and reducing applications.[7] Initial scaling efforts in the 1940s involved optimized direct synthesis—reacting molten lithium with hydrogen at 500–700°C under pressure in continuous-flow reactors—yielding purer products via patents like U.S. Patent 2,408,748 (1946), which emphasized economical production from electrolytic lithium.[14] These developments laid groundwork for broader reactivity studies, though production remained limited to specialized needs until post-war advancements.[15]Structure and bonding
Crystal structure
Lithium hydride adopts a rock-salt crystal structure, classified as face-centered cubic with space group Fm\overline{3}m (No. 225). In this ionic lattice, Li⁺ cations occupy the (0,0,0) positions and H⁻ anions the (0.5,0.5,0.5) sites, forming an arrangement where each ion is octahedrally coordinated to six ions of the opposite type.[3][16] The experimental lattice parameter at room temperature is 4.083 Å, corresponding to a unit cell volume that yields a density of 0.82 g/cm³ for the pure compound.[17][5] Under standard conditions of temperature and pressure, this cubic phase represents the sole stable polymorph, with no transitions observed at ambient pressures.[16][18] Impurities such as lithium nitride (Li₃N) or oxides commonly present in commercial samples introduce lattice defects or secondary phases, leading to bluish-gray discoloration and potential deviations in lattice perfection or mechanical integrity compared to pure white single crystals.[1][17] These contaminants arise during synthesis and handling, compromising the structural homogeneity essential for applications requiring high purity.[1]Nature of the Li-H bond
The Li-H bond in lithium hydride exhibits predominantly ionic character, modeled as an electrostatic interaction between Li⁺ and H⁻ ions, consistent with the compound's high melting point of 680 °C and its adoption of a rock salt lattice.[19] This ionic description aligns with Pauling's electronegativity difference of 1.22 between lithium (0.98) and hydrogen (2.20), which falls in the borderline range but favors ionicity given the lattice stability.[20] However, first-principles quantum chemistry calculations, including density functional theory (DFT), indicate partial covalent contributions arising from orbital overlap, particularly in the gas-phase diatomic LiH molecule, where the bond dissociation energy is approximately 243 kJ/mol.[21] Spectroscopic evidence supports this mixed character: the gas-phase LiH vibrational frequency (ω_e) is 1415.4 cm⁻¹, higher than expected for a purely ionic bond and indicative of strengthened bonding from covalent admixture, as lower frequencies would reflect weaker electrostatic interactions.[22] The equilibrium bond length of 1.596 Å in diatomic LiH is significantly shorter than the sum of ionic radii (Li⁺ ≈ 0.76 Å, H⁻ ≈ 1.40 Å, totaling ~2.16 Å), evidencing partial wavefunction overlap between lithium 2s and hydrogen 1s orbitals.[22] Mulliken population analysis in ab initio studies of LiH clusters further quantifies this, showing charge transfers of ~0.7-0.8 e⁻ rather than full ionicity, with covalent effects influencing structural distortions under pressure or in solvated forms.[23] In comparison to other alkali hydrides, the Li-H bond displays uniquely pronounced covalent features due to lithium's highest group electronegativity and smallest ionic radius (76 pm vs. 102 pm for Na⁺), which enhances polarization and orbital hybridization not as evident in NaH or KH.[20] DFT models of NaH yield near-complete charge separation (>0.95 e⁻), correlating with its lower reactivity toward covalent substrates, whereas LiH's partial covalency (~20-30% admixture per hybrid orbital analyses) underlies its enhanced reducing power and solubility anomalies in select solvents.[23] These theoretical insights explain LiH's reactivity, such as faster kinetics in hydrogen desorption pathways compared to heavier analogs, without invoking bulk lattice effects.[24]Properties
Physical properties
Lithium hydride is obtained as a white to gray, odorless powder that darkens upon exposure to light.[2] Its density is 0.82 g/cm³ at 25 °C, with bulk densities for powdered forms ranging from 290 to 430 kg/m³.[2][25] The compound melts at 680 °C, with reported values ranging from 680 °C to 688.7 °C depending on measurement conditions.[2][25] It decomposes at higher temperatures without boiling, typically above 850–900 °C.[25][4]| Property | Value | Notes/Source |
|---|---|---|
| Solubility in organics | Insoluble in acetone, benzene, toluene; slightly soluble in dimethylformamide | Reacts with water[2] |
| Thermal conductivity | ~14.7 W/m·K at 300 K | For polycrystalline samples; varies with temperature and density[26] |
| Mohs hardness | 3.5 | Indicates relatively soft material |
Thermodynamic properties
The standard enthalpy of formation (ΔfH°298) of solid lithium hydride is −90.5 kJ/mol.[29] The standard Gibbs free energy of formation (ΔfG°) is negative, indicating thermodynamic stability relative to the elements, with values derived from electromotive force measurements yielding ΔG° ≈ −68 kJ/mol near 298 K.[30] The standard molar entropy (S°298) is 20.75 J/(mol·K) for the solid phase, reflecting its ionic lattice structure.[31] The molar heat capacity at constant pressure (Cp) of LiH at 298 K is 27.9 J/(mol·K), increasing with temperature due to lattice vibrations in the rock-salt crystal.[29] Differential scanning calorimetry measurements confirm this behavior up to 800 K, with no anomalous peaks indicative of phase transitions in that range under ambient pressure.[32] Lithium hydride decomposes thermally via the endothermic reaction 2LiH(s) → 2Li(s) + H2(g) above approximately 959 K at standard pressure, with the decomposition temperature rising linearly with deuterium content in LiH1−xDx mixtures up to 999 K for pure LiD.[33] The Li-LiH phase diagram exhibits peritectic decomposition, where LiH melts incongruently into liquid lithium and hydrogen gas, with partial molar enthalpies of solution around 64.7 kJ/mol.[34] Under elevated pressures, LiH maintains stability in its B1 (NaCl-type) phase up to at least 36 GPa, transitioning to a B2 (CsCl-type) phase at higher pressures near 200 GPa, as probed by ab initio computations and diffraction studies.[35][36]Chemical properties
Lithium hydride (LiH) is characterized by the presence of the hydride anion (H⁻), which endows the compound with strong reducing properties and pronounced basicity, stemming from the high proton affinity of H⁻ as the conjugate base of dihydrogen.[4] The oxidation states in LiH are +1 for lithium and -1 for hydrogen, consistent with the ionic formulation Li⁺H⁻.[37] Owing to the nucleophilic and basic nature of the hydride ion, LiH displays extreme sensitivity to atmospheric oxygen and moisture, rendering it pyrophoric with potential for spontaneous combustion upon exposure.[1][4] Consequently, it requires handling and storage exclusively under inert atmospheres, such as dry argon or nitrogen, to mitigate reactivity with protic or oxidizing species.[1] Incompatible with most common laboratory solvents due to gradual or rapid decomposition, LiH's chemical stability is preserved only in anhydrous, non-reactive environments.[38]Synthesis and production
Direct synthesis from elements
Lithium hydride is synthesized directly from its elements via the exothermic reaction of molten lithium metal with hydrogen gas: $ 2 \mathrm{Li} + \mathrm{H_2} \rightarrow 2 \mathrm{LiH} $. This method, the primary route for producing high-purity LiH, requires heating lithium to its molten state and exposing it to a stream of dry hydrogen at atmospheric pressure. The reaction proceeds rapidly at temperatures of 700–900 °C, typically conducted in low-carbon iron crucibles to minimize contamination from reactive vessel materials.[39] Lower temperatures, below approximately 500 °C, result in impractically slow kinetics due to the formation of a passivating LiH overlayer on the lithium surface, which hinders hydrogen diffusion to unreacted metal.[18] Optimization of yields involves using high-purity hydrogen gas, often predried over phosphorus pentoxide and further purified by passage through molten alkali metals like sodium or potassium to remove trace oxygen and moisture that could induce side reactions forming lithium oxide or hydroxide. Near-quantitative yields are achievable under controlled conditions, with the reaction's exothermicity necessitating efficient heat dissipation to prevent localized overheating and incomplete conversion. Early laboratory-scale syntheses highlighted scaling challenges, including uniform gas distribution in larger reactors and management of the solid product buildup, which could impede heat transfer and gas flow.[40] Commercial adoption of this direct synthesis accelerated during World War II, when LiH production ramped up to supply hydrogen generation for inflating emergency rescue flares and life rafts via hydrolysis. Post-war, the method persisted for industrial-scale output, supporting nuclear applications such as neutron moderation in reactors, where high-purity LiH was essential; U.S. production facilities, including those at Oak Ridge Y-12, adapted the process for enriched lithium compounds amid Cold War demands. Purification of the crude LiH, obtained as a brittle solid upon cooling, typically entails inert-atmosphere crushing to powder form, followed by vacuum heating to volatilize residual hydrogen or impurities, ensuring minimal contamination for downstream uses.[41][42]Alternative production methods
One alternative method involves high-pressure compression of elemental lithium and hydrogen gas in a diamond anvil cell, achieving synthesis at pressures as low as 50 MPa and room temperature, contrasting with higher-temperature direct routes.[43] This approach, demonstrated in 2012, enables LiH formation under controlled quasihydrostatic conditions using laser heating if needed, though scalability remains limited due to equipment constraints.[44] Mechanochemical synthesis via reactive ball milling of lithium metal under hydrogen atmosphere offers a catalyst-free, room-temperature alternative operable at ambient pressures.[45] In one variant, organic solvent assistance facilitates hydride formation at mild hydrogen pressures (e.g., 0.1-1 MPa), yielding high-purity LiH with potential for large-scale production due to its avoidance of extreme temperatures or pressures.[46] Solvent- and catalyst-free milling at 700°C or below with 0.02 MPa hydrogen has also been reported, emphasizing energy efficiency for hydrogen storage applications.[47] Electrochemical routes, such as pulsed-potentiometric reduction under 7 bar hydrogen at room temperature, provide another pathway, though primarily explored in research settings for hydride precursors rather than bulk production.[48] These methods prioritize nanostructuring or destabilization, as in hydrogenation of lithium-intercalated graphite to form LiH-carbon composites with altered thermodynamics for enhanced reactivity.[49] Recycling-specific processes for LiH from lithium waste remain underdeveloped, with focus instead on broader lithium recovery techniques not yielding hydride directly.[50]Reactions
Reactions with protic compounds
Lithium hydride reacts violently with water, undergoing hydrolysis to produce lithium hydroxide and hydrogen gas according to the equation LiH + H₂O → LiOH + H₂.[10] This reaction is highly exothermic, releasing approximately 114 kJ/mol of heat, which can lead to ignition of the evolved hydrogen if not controlled.[27] The kinetics follow a nonequilibrium thermodynamic model, with initial surface adsorption of water preceding the bulk reaction; under low relative humidity (e.g., below 0.04%), significant hydrolysis is minimized, but exposure to liquid water or high humidity results in rapid, sustained progression potentially forming additional products like Li₂O or LiOH·H₂O over time.[51] [52] [10] Analogous reactions occur with alcohols, where LiH acts as a strong base, deprotonating the protic solvent to yield lithium alkoxides and hydrogen gas; lower alcohols like methanol react vigorously, while higher alcohols and phenols proceed more slowly.[53] With acids, the reaction is even more rapid and exothermic, liberating hydrogen and forming lithium salts, such as LiCl from HCl, necessitating inert handling to prevent uncontrolled gas evolution.[4] Lithium hydride also reacts with ammonia, albeit slowly at room temperature and accelerating above 300 °C, via LiH + NH₃ → LiNH₂ + H₂, an exothermic process with ΔH ≈ -43.1 kJ/mol H₂.[54] These reactions underscore the compound's extreme moisture sensitivity, requiring storage and manipulation under anhydrous, inert atmospheres to avoid spontaneous ignition or explosion from hydrogen buildup; quenching excess LiH typically involves controlled addition to protic media in ventilated systems to manage heat and gas release.[10] [4]Reducing agent applications
Lithium hydride (LiH) functions as a hydride donor in reduction reactions, primarily in niche organic and inorganic contexts where its ionic character and insolubility in aprotic solvents limit broader utility compared to soluble reagents like lithium aluminum hydride (LiAlH₄), which enables faster, more homogeneous reductions of carbonyls to alcohols. LiH's heterogeneous reactions often proceed slowly at elevated temperatures, potentially yielding lower selectivity and side products such as over-reduced hydrocarbons or decomposition-derived gases.[55] In organic synthesis, commercial LiH requires activation—typically via in situ generation of lithium alkoxides with nickel salts—to enhance reactivity, enabling reductions of ketones to secondary alcohols, alkyl or aryl halides to hydrocarbons, and ethylenic or sulfurated compounds.[55] For example, activated LiH effects reductive silylation of carbonyls when combined with chlorotrimethylsilane, providing protected alcohols.[56] A specialized application involves heating LiH with acid chlorides (RCOCl) or thioesters (RCOSR') in boiling benzene, toluene, or xylene to afford aldehydes (RCHO), contrasting with LiAlH₄'s tendency for over-reduction to primary alcohols and offering selectivity akin to milder agents.[27] In inorganic chemistry, LiH reduces certain metal salts to lower oxidation states or facilitates adduct formation, as in the synthesis of binary hydrides from chlorides, though specific yields and conditions vary with the substrate's reactivity. These applications leverage LiH's high hydride content but are constrained by handling requirements to prevent premature hydrolysis or ignition.[43]Thermal decomposition and other reactions
Lithium hydride undergoes thermal decomposition at high temperatures, dissociating into elemental lithium and hydrogen gas via the endothermic reaction . This process becomes significant when the equilibrium hydrogen vapor pressure exceeds approximately 30 torr, which occurs around 750 °C under vacuum or low-pressure conditions.[57] Complete decomposition to lithium metal is achieved at 959 K (686 °C) for protium-containing LiH, while the deuterated isotopologue LiD requires a higher temperature of 999 K due to isotopic mass differences affecting vibrational energies and decomposition kinetics.[33] The decomposition pathway exhibits isotope effects, with mixed LiHD systems showing variations in stoichiometry and decomposition behavior influenced by the hydrogen-to-deuterium ratio, as predicted by thermodynamic models and experimental measurements of dissociation pressures from 450 to 750 °C.[58] [33] These differences arise from zero-point energy variations, leading to distinct plateau pressures and temperatures for hydrogen release in LiH versus LiD.[33] Beyond decomposition, lithium hydride engages in solid-state reactions with certain metals, particularly transition elements, to form addition compounds or alloys under hydrogen atmospheres at elevated temperatures. For instance, second- and third-period transition metals react with LiH and H₂ to yield novel intermetallic phases not obtainable from the elements alone.[59] Similarly, LiH reacts with aluminum in solid-state conditions to produce aluminohydrides or related composites, highlighting its role in forming metal hydride alloys.[60] Isotopic exchange represents another key reaction pathway, enabling the preparation of lithium deuteride (LiD) through gas-solid interactions. Single crystals of LiH undergo deuterium exchange with D₂ gas, with the rate dependent on pressure and facilitating H/D substitution via surface-mediated diffusion and hydrogen transport mechanisms.[61] This process has been quantified as a function of deuterium pressure, underscoring LiH's utility in isotope separation and production of heavy hydride isotopologues for applications requiring specific nuclear properties.[61]Applications
Hydrogen storage
Lithium hydride (LiH) possesses one of the highest theoretical gravimetric hydrogen storage capacities among simple metal hydrides, at approximately 12.7 wt%, calculated from its molecular weight where hydrogen constitutes 1.008 g out of 7.948 g per formula unit.[62] This exceeds that of magnesium hydride (MgH₂) at 7.6 wt% and sodium hydride (NaH) at 4.2 wt%, positioning LiH favorably for applications demanding maximal hydrogen density by weight.[63] However, practical utilization for reversible storage is hindered by the compound's thermodynamic stability, stemming from the strong Li–H bond. Hydrogen release from LiH occurs via thermal decomposition: , an endothermic process with a reaction enthalpy of about 75–90 kJ/mol H₂, necessitating temperatures exceeding 600–900°C under ambient pressure for significant desorption rates.[62] [64] Rehydrogenation to reform LiH demands elevated pressures (hundreds of bar) and temperatures, resulting in sluggish kinetics and limited reversibility; empirical studies report peak desorption temperatures as low as 190°C in nanocrystalline or modified forms, but cycling leads to capacity fade due to sintering of lithium metal and incomplete reabsorption.[65] Compared to complex hydrides like LiBH₄ (18.5 wt% theoretical), LiH offers simpler decomposition but inferior low-temperature performance, with desorption onset >400°C even under optimized conditions, versus MgH₂'s ~300°C.[66] Efforts to mitigate these challenges include doping with catalysts such as transition metals or carbon nanostructures to enhance kinetics and lower desorption barriers, achieving partial reversibility in lab-scale cycles (e.g., ~4 wt% releasable at 400°C in silicon-incorporated variants).[67] [64] Despite such modifications, systemic issues persist: the energy penalty for high-temperature operation offsets gravimetric advantages in system-level efficiency, and irreversible losses accumulate over cycles, rendering LiH suboptimal for onboard vehicular storage relative to established options like LaNi₅ alloys (1.4 wt%, but superior cyclability at <100°C).[63] These limitations underscore LiH's niche rather than broad applicability in reversible hydrogen storage.Precursor to complex hydrides
Lithium hydride (LiH) functions as a fundamental building block in the synthesis of complex metal hydrides, including lithium aluminum hydride (LiAlH4) and lithium borohydride (LiBH4), enabling the formation of materials with elevated hydrogen capacities suitable for advanced storage systems. These syntheses typically proceed via metathesis reactions in ethereal solvents, where LiH reacts with metal halides or boron halides to displace chloride or fluoride ions while incorporating hydride ligands.[68][69] A primary example is the production of LiAlH4, achieved through the reaction of LiH with anhydrous aluminum chloride (AlCl3) in diethyl ether:4 LiH + AlCl3 → LiAlH4 + 3 LiCl. This process yields LiAlH4 with a theoretical hydrogen content of 10.6 wt%, and it has been adapted in post-1990s efforts to generate nanostructured variants via mechanochemical or solvent-free methods for enhanced reactivity in composite materials.[68] Similarly, LiBH4 is formed by reacting LiH with boron trifluoride (BF3) under specific molar ratios exceeding 4:1 in ether: 4 LiH + BF3 → LiBH4 + 3 LiF, resulting in a hydride boasting 18.5 wt% hydrogen and potential for thermodynamic reversibility through endothermic decomposition pathways.[69][70]
These complex hydrides, derived from LiH, facilitate lightweight composites that improve hydrogen release kinetics when doped or nanostructured, as explored in research since the late 1990s amid growing interest in solid-state storage exceeding 7 wt% targets.[71] Nonetheless, practical limitations persist, including the elevated cost of LiH (often exceeding $100/kg in bulk) and stringent purity demands to avoid hydrolysis or impurity-induced degradation during synthesis, which can compromise yield and material performance in scaled applications.[68][70]