Aside from lithium’s well-known use in batteries (rechargeable and otherwise) it seems like it might be useful as a fuel, similar to the tamer magnesium, or alloyed with it.
| Property | Li | Li₂O | LiH | Li₂S | O | S | H₂O | SO₂ |
|---|---|---|---|---|---|---|---|---|
| Molar mass, g/mol | 6.94 | 29.88 | 7.95 | 45.95 | 15.999 | 32.06 | 18.015 | 64.066 |
| Density, g/cc | 0.534 | 2.013 | 0.78 | 1.67 | .001429 | 2.07 | 1.00 | .00263 |
| Enthalpy of formation, kJ/mol | 0 | -595.8 | -90.65 | -447 | 0 | 0 | -285.83 | -296.81 |
| ", MJ/kg | 0 | -20.01 | -11.4 | -9.401? | 0 | 0 | -15.87 | -4.63 |
| Heat capacity (room temp) J/mol/K | 24.860 | 54.1 | 27.9 | ?? | 29.378 | 22.75 | 75.39 | 39.87 |
| Specific heat, J/g/K | 3.58 | 1.81 | 3.51 | ?? | 0.92 | 0.71 | 4.184 | 0.622 |
| Extrapolated ΔT, °C | 0 | 11000 | 3200 | ?? | 0 | 0 | 3800 | 7400 |
| Melting point, °C | 181 | 1438 | 689 | 938 | -219 | 115 | 0 | -72 |
| Boiling point, °C | 1330 | 2600 | 900 | 1372 | -183 | 445 | 100 | -10 |
The extrapolated temperature changes here are what the temperature would be in forming the compound from the elements, if its specific heat stayed the same as that at room temperature and it underwent no phase transitions. But of course specific heats typically rise with temperature (at 1200 K SO₂’s heat capacity is 55.81 J/mol/K) and several of these materials do in fact undergo phase transitions.
The lithium combustion equation should be 4Li + O₂ → 2Li₂O, producing 1192 kJ per mole of O₂ and per four moles of Li, thus 298 kJ per mole of Li; that’s 42.9 MJ/kg, quite impressive, almost twice magnesium’s 24.7 MJ/kg. The energy density then would be 22.9 MJ/ℓ, lower than magnesium’s 43 MJ/ℓ.
So it turns out that lithium as a fuel burned with oxygen has a specific energy even higher than magnesium, though much lower density. Because of that and because magnesia’s specific heat is higher than lithia’s, burning lithium probably will reach lower temperatures.
At 2%, magnesium is much more abundant in Earth’s crust than lithium — despite lithium’s name, it’s only found at 20 ppm. This is a disadvantage for lithium as an energy carrier. Aside from lithium’s higher specific energy, though, it also melts at a reasonable temperature which might make it possible to use as a liquid fuel in some kinds of engines, though for some of them the refractory nature of lithia would be a drawback.
Lithium, like magnesium, is produced by molten-salt electrolysis of the chloride, in this case fluxed with KCl, and at a very friendly temperature of 450°. I don’t know of a lithium equivalent to the Pidgeon silicothermic reduction process for magnesium.
Reacting lithium instead to the hydride or the poorly characterized sulfide would produce about half as much heat, although that’s still a pretty acceptable energy density, and those reactions also produce no gases, just solid or liquid salts. It may be feasible to produce the sulfide with a reaction between molten lithium and molten sulfur, producing molten lithium sulfide, although sulfur’s tendency to polymerize around lithium’s melting point may be an obstacle.
Also, though, those salts can themselves be used as fuels!
Either of these can be burned as a fuel with oxygen to produce the other half of the heat, plus additional energy from oxidizing the anion. They may be more or less convenient than the lithium metal due to their higher stability in air (?) and melting points. The sulfide has a higher boiling point than lithium as well.
As for the hydride’s stability in air, WP says that lumps of it form a protective tarnish in humid air, inhibiting further reaction, doesn’t ignite “in a metal dish” (?) until heated past 200°, and is “less reactive with water” than Li, but still “highly reactive” and “reacts violently with water”. Furthermore you cannot extinguish its fires with ordinary (presumably quartz) sand.
The sulfide is reported to be deliquescent in air and, like other metal sulfides, hydrolyzes to produce sulfuretted hydrogen. Unlike lithium or the hydride, it does not seem to pose an explosion risk with water, just a poison gas risk.
The hydride combustion equation would be something like 2LiH + O₂ → Li₂O + H₂O, yielding steam and 595.8 + 285.83 - 2×90.65 kJ per two moles of the hydride, 700 kJ, which is 350 kJ per mole of the hydride or 44 MJ/kg. This is almost the same as the specific energy of lithium itself! It’s an energy density of 34 MJ/ℓ, the same as gasoline or diesel fuel, much better than lithium, and nearly as high as magnesium.
The sulfide combustion equation would be 2Li₂S + 3O₂ → 2Li₂O + 2SO₂, yielding SO₂ gas and (2×595.8 + 2×296.81 - 2×447) kJ per two moles of the sulfide; that’s 892 kJ per two moles or 446 kJ per mole, which is 9.7 MJ/kg and 5.8 MJ/ℓ. That’s still usable as a fuel, but it’s at the low end of what’s usable, and the corrosive gas is probably a killer drawback.
The hydride may also be interesting as a potential working fluid for high-temperature heat engines due to its large expansion when heated; it decomposes to liquid lithium and gaseous hydrogen at 900°–1000°, and at 1330° the lithium boils, I think to individual atoms rather than diatomic molecules. So each mole of hydride will, I think, produce three moles of gas. At room temperature the hydride’s 0.78 g/cc is 0.098 mol/cc, 98 moles per liter, while at 1330° the 294 moles of gas produced from that liter would ideally each occupy some 132 liters, a total of nearly 39 cubic meters.
Whether the sulfide would act similarly is, I think, anybody’s guess.
The borohydride of lithium has been discussed as a possible fuel, though no boron-containing fuel is in use today for excellent reasons.
There is also a metastable lithium aluminum hydride LiAlH₄, containing a “tetrahydroaluminumate” or “alanate” or “tetrahydridoaluminate(III)” or “alumanuide” ion; it decomposes to lithium hydride and a different lithium aluminum hydride Li₃AlH₆. It contains nearly as much hydrogen by weight as lithium hydride itself, and is even denser at 0.917 g/cc, so it contains more hydrogen by volume.