Lithium was discovered in 1817 by the Swedish chemist Johan August Arfwedson (1792-1841) while analyzing the mineral petalite from the island of Utö in Sweden. Arfwedson identified the presence of a new alkaline element but was unable to isolate it in metallic form. His mentor, Jöns Jacob Berzelius (1779-1848), named this element lithium (from the Greek lithos = stone), as it was the first alkali metal discovered in a mineral rather than in plant matter. It was not until 1821 that the British chemist William Thomas Brande (1788-1866) and independently the Swedish chemist Johan August Arfwedson succeeded in isolating metallic lithium by electrolysis of lithium oxide.
Lithium (symbol Li, atomic number 3) is the first alkali metal in the periodic table, consisting of three protons, usually four neutrons (for the most common isotope), and three electrons. The two stable isotopes are lithium-7 \(\,^{7}\mathrm{Li}\) (≈ 92.5%) and lithium-6 \(\,^{6}\mathrm{Li}\) (≈ 7.5%).
At room temperature, lithium is a soft, silvery-white metal, extremely light (density ≈ 0.534 g/cm³), making it the least dense of all metals. It is highly reactive, particularly with water and oxygen, and must be stored under mineral oil or in an inert atmosphere. The temperature at which the liquid and solid states can coexist (melting point): 453.65 K (180.50 °C). The temperature at which it transitions from liquid to gas (boiling point): 1615 K (1341.85 °C).
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Lithium-6 — \(\,^{6}\mathrm{Li}\,\) | 3 | 3 | 6.015122 u | ≈ 7.5 % | Stable | Used in nuclear fusion to produce tritium; absorbs thermal neutrons. |
| Lithium-7 — \(\,^{7}\mathrm{Li}\,\) | 3 | 4 | 7.016003 u | ≈ 92.5 % | Stable | Major isotope; used in lithium-ion batteries and industrial applications. |
| Lithium-8 — \(\,^{8}\mathrm{Li}\,\) | 3 | 5 | 8.022487 u | Unnatural | 0.838 s | Radioactive β\(^-\) decay to \(\,^{8}\mathrm{Be}\), which immediately decays into two alpha particles. |
| Lithium-9 — \(\,^{9}\mathrm{Li}\,\) | 3 | 6 | 9.026790 u | Unnatural | 0.178 s | Radioactive β\(^-\); artificially produced in particle accelerators. |
| Lithium-4, 5 — \(\,^{4}\mathrm{Li},\,^{5}\mathrm{Li}\,\) | 3 | 1 — 2 | — (resonances) | Unnatural | \(10^{-22}\) s | Very unstable states observed in nuclear physics; immediate decay. |
| Heavy isotopes — \(\,^{10}\mathrm{Li},\,^{11}\mathrm{Li},\,^{12}\mathrm{Li}\) | 3 | 7 — 9 | — (resonances) | Unnatural | \(10^{-21}\) — 0.009 s | Neutron halos; \(\,^{11}\mathrm{Li}\) has two very weakly bound neutrons forming a halo around the nucleus. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Lithium has 3 electrons distributed across two electron shells. Its full electronic configuration is: 1s² 2s¹, or simplified as: [He] 2s¹. This configuration can also be written as: K(2) L(1).
K Shell (n=1): Contains 2 electrons in the 1s sub-shell. This inner shell is complete and highly stable, forming a configuration identical to that of helium.
L Shell (n=2): Contains only 1 electron in the 2s sub-shell. This single valence electron is relatively weakly bound to the nucleus and easily lost during chemical reactions. The 2p orbitals remain completely empty.
The single electron in the outer shell (2s¹) is the valence electron of lithium. This configuration explains its chemical properties:
By losing its 2s electron, lithium forms the Li⁺ ion (oxidation state +1), its unique and systematic oxidation state in all its compounds.
The Li⁺ ion then adopts an electronic configuration identical to that of helium [He], a noble gas, which gives this ion maximum stability.
Lithium does not exhibit any other oxidation state; only the +1 degree is observed in chemistry.
The electronic configuration of lithium, with its valence shell containing a single 2s electron, classifies it among the alkali metals (group 1 of the periodic table) and makes it the lightest of all metals. This structure gives it characteristic properties: high chemical reactivity (it reacts with water, oxygen, and most non-metals), low ionization energy (the valence electron is easily removed), and exclusive formation of ionic compounds with an oxidation state of +1.
Lithium is a soft, silvery metal with a very low density (0.53 g/cm³, the lightest metal), which must be stored under mineral oil or in an inert atmosphere to protect it from oxidation. It reacts slowly with water at room temperature, unlike sodium and potassium, which react violently. This moderate reactivity compared to other alkali metals is explained by its small atomic size and relatively stronger bond energy.
The importance of lithium has become crucial in the modern world: lithium-ion batteries have become essential for portable electronics (smartphones, computers) and electric vehicles, making lithium a strategic element for the energy transition; lithium carbonate Li₂CO₃ is used in psychiatry to treat bipolar disorders; aluminum-lithium alloys are used in aerospace for their exceptional lightness; lithium is used as a flux in welding and brazing processes; lithium hydride LiH is a powerful reducing agent and a potential hydrogen storage medium; organolithium compounds (such as butyllithium) are important reagents in organic chemistry. Lithium-6 is used in nuclear technology to produce tritium. Global demand for lithium is growing exponentially with the development of battery technologies, making its extraction and recycling major economic and environmental challenges.
Lithium is an extremely reactive alkali metal. It has a single valence electron that it readily donates, forming the Li⁺ ion. It reacts vigorously with water to produce lithium hydroxide (LiOH) and hydrogen gas. In contact with air, lithium rapidly oxidizes to form lithium oxide (Li₂O) and lithium nitride (Li₃N), the latter being an unusual reaction among alkali metals. Lithium also forms compounds with halogens (lithium fluoride, chloride, bromide) and reacts with carbon to produce lithium carbide (Li₂C₂). Its strong electropositivity makes it an excellent reducing agent in organic and inorganic chemical reactions.
Lithium holds a unique place in cosmology as it is one of the only three elements (along with hydrogen and helium) synthesized in significant quantities during primordial nucleosynthesis, a few minutes after the Big Bang. However, the observed abundance of lithium in the current universe poses a major problem known as the "cosmological lithium problem". Big Bang models predict an abundance of lithium-7 about three times higher than that observed in the old stars of our galaxy.
In stars, lithium is quickly destroyed by nuclear fusion at relatively low temperatures (about 2.5 million kelvin), well below the temperatures required to burn hydrogen. This destruction makes lithium an excellent tracer for studying the internal mixing processes of stars and their evolution. Measuring the abundance of lithium in different types of stars allows astrophysicists to constrain stellar models and understand the chemical history of the galaxy.
Lithium-6, although rare, can be produced by cosmic ray reactions in the interstellar medium. Its ratio to lithium-7 provides valuable information about the intensity of cosmic rays in the past of our galaxy and about galactic nucleosynthesis processes.
Spectroscopic study of lithium in the atmospheres of exoplanets and brown dwarfs also helps determine their age and thermal history, as the presence or absence of lithium indicates whether the object has reached internal temperatures sufficient to destroy it.
N.B.:
The "cosmological lithium problem" remains one of the unsolved mysteries of modern cosmology. Several hypotheses have been proposed to explain this discrepancy: destruction of lithium in the first stars, errors in primordial nucleosynthesis models, physics beyond the standard model, or observational biases in measuring lithium abundance. This enigma illustrates that even the simplest elements can reveal deep and mysterious aspects of the evolution of our universe, and its resolution could have major implications for our understanding of fundamental physics and cosmology.
