Sodium

Category: alkali metal Atomic number: 11 Colour: silvery white Melting point: 98°C (208°F) Boiling point: 883°C (1,621°F) First identified: 1807

At least two important sodium compounds have been used since the earliest civilizations. The ancient Egyptians harvested sodium carbonate (‘soda’) from dried-out flood plains near the River Nile – the crystals, which were used as a cleaning agent, are also mentioned in the Bible. And salt (sodium chloride), which was harvested from salt flats (or underground deposits), has always been an important part of our diet, whether added to meals or absorbed from animal-based food.

We have about 100 grams of sodium in our body – it is an electrolyte, like potassium and calcium, and thus is crucial in regulating inflows to and outflows from cells. It helps cells to transmit nerve signals, and to regulate the water levels in our body. However, too much sodium can dangerously raise our blood pressure, which is why high-blood-pressure patients are advised to reduce their salt intake.

Salt taxes have caused unrest throughout history – for instance, they were a contributory factor to the French Revolution. And, when the British Empire imposed a salt tax on India, where the largest part of the population was vegetarian, Mahatma Gandhi’s salt march in protest became a key moment in the Indian independence movement.

Chemists in the Kitchen

Salt and soda have been with us for millennia, and caustic soda was first prepared by thirteenth-century soapmakers, but baking soda (sodium bicarbonate) is a relatively recent innovation. In 1843, Alfred Bird, a British chemist, created a batch to help his wife, who had a yeast allergy. It works by releasing carbon dioxide bubbles into the batter or dough through an acid-base reaction.

In spite of sodium being the sixth most common element on the planet and one whose compounds were extensively used by humans, it was only in the nineteenth century that its true nature was understood. It is extremely reactive, so it is never found naturally in its pure form (it tarnishes immediately if exposed to the air and can only be preserved beneath particular oils).

The first pure sodium was extracted by Sir Humphry Davy at London’s Royal Institution – he ran an electric current through caustic soda (sodium hydroxide) and produced small globules of the metal. (Today it is generally produced by electrolysis of dry molten sodium chloride.)

It has many practical uses, including as a coolant in nuclear reactors, for de-icing roads in the winter (in the form of salt), clearing drains (in the form of caustic soda), and acting as a reagent (a substance that triggers reactions) in the biochemicals industry. However, most chemists know that the best fun you can have with a freshly sliced lump of sodium is to drop it into water and watch the resulting reaction (it bursts into flames before exploding) from a safe distance – the smartest way to see this is to watch a video on the internet, rather than trying it at home!