Sodium

Sodium is an alkali metal so reactive that it never occurs free in nature. A piece dropped into water melts, skitters across the surface and sets fire to the hydrogen it is releasing. It is stored under kerosene.

And yet people eat sodium compounds every day: ordinary table salt is sodium chloride.

In seawater, in rock salt deposits, in soda, and in the silicates of the crust. Sodium is the sixth most abundant element in the crust and the second most abundant of the ions dissolved in the ocean.

Large rock salt deposits are dried-up ancient seas. They gave rise to salt roads, salt taxes and more than one war: before refrigeration, salt was the only preservative available at scale.

Sodium compounds were used in antiquity, but the metal resisted isolation for a long time: it cannot be reduced with charcoal the way copper or iron can.

Humphry Davy solved it in 1807 using electrolysis, then a new technique. Passing a current through molten caustic soda, he obtained metallic globules that caught fire in air. By the same method and in the same year he isolated potassium, and a year later calcium, magnesium, barium and strontium. Electrolysis gave chemistry more new elements in two years than the preceding century had.

The symbol Na comes from the Latin natrium, itself from the Egyptian name for natural soda.

Sodium has one electron in its outer shell and parts with it readily: its ionisation energy is low and its electronegativity among the lowest there is. Hence a single oxidation state, +1, and high reactivity.

With water it reacts violently, giving an alkali and hydrogen. In air it oxidises instantly — a fresh cut shines like mercury and dulls within seconds.

One practically important feature is the characteristic yellow it gives a flame. The test is so sensitive that traces of sodium show up in almost any sample, and the same emission is behind the unmistakable yellow of sodium street lamps.

Metallic sodium serves as a coolant in some nuclear reactors: it conducts heat well and stays liquid over a wide range of temperatures. It is also used as a reducing agent in organic synthesis.

Its compounds are used incomparably more widely: the chloride in food and as feedstock for chlorine and caustic soda; the hydroxide in making paper, soap and aluminium; the carbonate in glassmaking; the bicarbonate in the kitchen and in medicine.

Sodium ions are the principal cations of extracellular fluid. The difference in sodium and potassium concentrations across a cell membrane creates an electrical potential, and the abrupt change in it is a nerve impulse.

That difference is maintained by the sodium–potassium pump, a protein that pumps sodium out of the cell continuously. A noticeable share of all the energy an organism uses at rest goes into running it.