Sodium compounds, especially table salt (sodium chloride, NaCl), have been known and used since antiquity for food preservation and as a medium of exchange (hence the word "salary" from the Latin salarium). However, metallic sodium was only isolated in the early 19th century thanks to advances in electrochemistry.
In 1807, the British chemist Humphry Davy (1778-1829) isolated metallic sodium for the first time by electrolysis of molten caustic soda (sodium hydroxide, NaOH). A few days after isolating potassium by the same method, Davy succeeded in producing shiny and extremely reactive sodium globules. He observed that this metal oxidized rapidly in air and reacted violently with water, releasing hydrogen gas.
Davy named this element sodium (from the English word "soda" derived from the Arabic suwwad or suda, referring to certain plants from which soda was extracted). The chemical symbol Na comes from the Latin natrium, derived from natron (natural hydrated sodium carbonate) used since ancient Egypt. This discovery marked the beginning of the systematic study of alkali metals and revolutionized the understanding of the chemistry of elements.
Sodium (symbol Na, atomic number 11) is an alkali metal of group 1 of the periodic table, consisting of eleven protons, usually twelve neutrons (for the most common isotope), and eleven electrons. The only natural stable isotope is sodium-23 \(\,^{23}\mathrm{Na}\) (100% natural abundance).
At room temperature, sodium is a soft, silvery-white metal, soft enough to be cut with a knife. It has a relatively low density (≈ 0.968 g/cm³), lower than that of water, which means it would float on water if it did not react violently with it. Sodium is highly reactive, oxidizing rapidly in air and reacting vigorously with water to produce sodium hydroxide and hydrogen gas, a reaction sufficiently exothermic to ignite the hydrogen produced.
Sodium must be stored under mineral oil or in an inert atmosphere (argon) to protect it from oxidation. It has excellent electrical and thermal conductivity, typical characteristics of alkali metals.
The temperature at which the liquid and solid states can coexist (melting point): 370.944 K (97.794 °C). The temperature at which it transitions from liquid to gaseous state (boiling point): 1156.090 K (882.940 °C).
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Sodium-22 — \(\,^{22}\mathrm{Na}\,\) | 11 | 11 | 21.994437 u | Cosmogenic | 2.602 years | Radioactive β\(^+\) and electron capture yielding \(\,^{22}\mathrm{Ne}\) ; produced by cosmic rays; used in nuclear medicine. |
| Sodium-23 — \(\,^{23}\mathrm{Na}\,\) | 11 | 12 | 22.989769 u | 100 % | Stable | Only stable isotope; essential for biological functions; basis of all sodium compounds. |
| Sodium-24 — \(\,^{24}\mathrm{Na}\,\) | 11 | 13 | 23.990963 u | Non-natural | 14.997 hours | Radioactive β\(^-\) yielding \(\,^{24}\mathrm{Mg}\) ; used as a tracer in medicine and industry; strong gamma emission. |
| Sodium-25 — \(\,^{25}\mathrm{Na}\,\) | 11 | 14 | 24.989954 u | Non-natural | 59.1 s | Radioactive β\(^-\) ; produced in nuclear reactors. |
| Sodium-26 — \(\,^{26}\mathrm{Na}\,\) | 11 | 15 | 25.992633 u | Non-natural | 1.071 s | Radioactive β\(^-\) ; short half-life. |
| Other isotopes — \(\,^{18}\mathrm{Na}-\,^{21}\mathrm{Na},\,^{27}\mathrm{Na}-\,^{37}\mathrm{Na}\) | 11 | 7-10, 16-26 | — (resonances) | Non-natural | \(10^{-21}\) — 0.301 s | Very unstable states observed in nuclear physics; some have neutron halo structures. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Sodium has 11 electrons distributed across three electron shells. Its full electronic configuration is: 1s² 2s² 2p⁶ 3s¹, or simplified as: [Ne] 3s¹. This configuration can also be written as: K(2) L(8) M(1).
K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 8 electrons distributed as 2s² 2p⁶. This shell is also complete, forming a noble gas configuration (neon).
M Shell (n=3): contains only 1 electron in the 3s subshell. This single valence electron is very weakly bound to the nucleus and very easily lost during chemical reactions.
The single electron in the outer shell (3s¹) is the valence electron of sodium. This configuration explains its chemical properties:
By losing its 3s electron, sodium forms the Na⁺ ion (oxidation state +1), its unique and systematic oxidation state in all its compounds.
The Na⁺ ion then adopts an electronic configuration identical to that of neon [Ne], a noble gas, which gives this ion maximum stability.
Sodium exhibits no other oxidation state; only the +1 state is observed in chemistry.
The electronic configuration of sodium, with its valence shell containing a single 3s electron, classifies it among the alkali metals (Group 1 of the periodic table). This structure gives it characteristic properties: very high chemical reactivity (it reacts violently with water, releasing hydrogen gas and heat, and spontaneously ignites in humid air), very low ionization energy (the valence electron is extremely easy to remove), and exclusive formation of ionic compounds with an oxidation state of +1. Sodium is a soft, silvery metal that must be stored under mineral oil or in an inert atmosphere to protect it from oxidation. Its extremely strong tendency to lose its valence electron makes sodium one of the most powerful reducing agents and one of the most reactive metals. Its biological importance is fundamental: the Na⁺ ion plays a crucial role in regulating the water balance of organisms, transmitting nerve impulses, and maintaining the membrane potential of cells. In chemistry, metallic sodium is used as a powerful reducing agent, while its compounds are ubiquitous: sodium chloride NaCl (table salt), sodium hydroxide NaOH (caustic soda), sodium carbonate Na₂CO₃, and sodium bicarbonate NaHCO₃ are among the most widely used industrial chemicals.
Sodium has a single valence electron in its outer shell, which it readily donates to form the sodium ion (Na⁺) with a stable noble gas electronic configuration. This high electropositivity makes sodium a powerful reducing agent and an extremely reactive metal.
Sodium reacts vigorously with water according to the reaction: 2 Na + 2 H₂O → 2 NaOH + H₂ (gas). This reaction is sufficiently exothermic to melt sodium and ignite the hydrogen produced, creating a characteristic yellow flame due to the excitation of sodium atoms. With oxygen in the air, sodium rapidly forms a layer of sodium oxide (Na₂O) and sodium peroxide (Na₂O₂) that dulls its shiny surface.
Sodium reacts with halogens (fluorine, chlorine, bromine, iodine) to form halides, the most well-known and biologically important of which is sodium chloride (NaCl, table salt). It also reacts with hydrogen to form sodium hydride (NaH), a powerful reducing agent used in organic chemistry, and with liquid ammonia to form characteristic blue solutions containing solvated electrons.
In living organisms, the sodium ion (Na⁺) plays an absolutely essential physiological role. It is the main extracellular cation and maintains osmotic balance, regulates blood volume and blood pressure. Sodium is crucial for the transmission of nerve impulses and muscle contraction via sodium-potassium pumps (Na⁺/K⁺-ATPase) that maintain electrochemical gradients across cell membranes. An imbalance of sodium (hyponatremia or hypernatremia) can have serious, even fatal, consequences.
Sodium is a relatively abundant element in the universe, ranking approximately 14th in cosmic abundance. It is entirely produced by stellar nucleosynthesis, primarily in massive stars.
Sodium-23 is synthesized in massive stars through several nuclear processes. The main mechanism is carbon burning at temperatures of about 600-800 million kelvins, where the fusion of two carbon-12 nuclei can produce sodium-23 plus a proton (¹²C + ¹²C → ²³Na + p). Sodium can also be produced during neon burning by photodisintegration of neon-20 followed by alpha particle captures, or by processes involving magnesium in the burning stellar layers.
During supernova explosions, sodium is produced in significant quantities and ejected into the interstellar medium. Type II supernovae (gravitational collapse of massive stars) and Type Ia supernovae (thermonuclear explosion of white dwarfs) both contribute to the galactic enrichment of sodium, although through different mechanisms and with varying yields.
In the interstellar medium, neutral atomic sodium (Na I) exhibits characteristic absorption lines in the visible spectrum, notably the sodium D lines (doublet at 589.0 and 589.6 nm). These lines, discovered by Joseph von Fraunhofer in the solar spectrum in 1814, are among the strongest and most easily observable. They serve as tracers of the structure and dynamics of the interstellar medium, allowing astronomers to map gas clouds between stars and study their velocities via the Doppler effect.
The sodium D lines are also observed in the spectra of stars, providing information on temperature, chemical composition, and the movements of stellar atmospheres. The intensity of these lines varies considerably from star to star, reflecting differences in chemical abundance related to stellar metallicity and the chemical evolution of the galaxy.
In the solar system, sodium has been detected in several surprising environments. A tenuous sodium exosphere surrounds the planet Mercury, created by the sputtering of the surface by the solar wind and micrometeorite impacts. This sodium exosphere extends tens of thousands of kilometers and exhibits a comet-like tail. Sodium has also been detected in the exosphere of the Moon, in the geysers of Enceladus (Saturn's moon), and in the tails of comets.
The atmospheres of exoplanets have revealed the presence of sodium through transit spectroscopy. When an exoplanet passes in front of its star, the absorption of starlight by the planetary atmosphere creates a characteristic spectral signature. Sodium lines were among the first detected in the atmospheres of hot Jupiter-type exoplanets, providing crucial information on the composition, structure, and meteorology of these distant worlds.
In adaptive optical astronomy, sodium lasers are used to create artificial guide stars. These lasers excite sodium atoms in the Earth's mesospheric layer (at about 90 km altitude), creating a point light source that allows real-time measurement and correction of atmospheric distortions, significantly improving the resolution of ground-based telescopes.
N.B. :
The salt paradox illustrates the complex relationship between sodium and human health. Sodium is absolutely essential for life: without it, nerve impulses could not propagate, muscles could not contract, and the body's water balance would collapse. Historically, salt was so valuable that it served as a medium of exchange and wars were fought to control salt routes. However, in modern societies, the overconsumption of sodium (mainly in the form of salt) has become a major public health issue, contributing to high blood pressure, cardiovascular diseases, and strokes. The World Health Organization recommends less than 5 grams of salt per day, but average consumption in many developed countries exceeds 9-12 grams. This situation reflects a mismatch between our biology, which evolved in salt-poor environments (where conserving sodium was crucial), and our modern food environment rich in processed foods containing excessive amounts of added salt. Sodium thus embodies the fundamental toxicological rule formulated by Paracelsus: "Everything is poison, nothing is poison, it is the dose that makes the poison."
