Calcium

Silvery, soft and never alone

Calcium is a silvery metal, harder than the alkali metals next door but still cut by a determined knife, and lighter than any other alkaline earth metal — lighter than magnesium above it. It has two electrons in its outer shell and gives both up readily to become Ca²⁺, the only form in which nature holds it. In air the fresh metal dulls within hours under a grey coat of oxide and nitride; in water it fizzes steadily, releasing hydrogen without sodium's violence; heated in air it burns with a brick-red flame, the colour calcium salts lend to fireworks. Its compounds are what we meet: the carbonate in limestone, chalk and marble, the sulphate in gypsum and plaster, the fluoride in fluorite, the phosphate in apatite and in bone, and the oxide and hydroxide that have been called lime since Roman times.

Limestone, shells and lime

Some 4.15 % of the crust is calcium, the fifth most abundant element in it. The calcium that matters lies in limestone: thousands of metres of it, laid down over hundreds of millions of years, nearly all of it once alive. Corals, shellfish and single-celled plankton draw calcium and carbonate out of seawater to build their skeletons; when they die the skeletons settle, and chalk, limestone and marble are what remains. Rain, made faintly acid by carbon dioxide, dissolves limestone again and hollows out caves, and the water that runs off it is “hard”.

Burn limestone at 900 °C and it gives up carbon dioxide, leaving quicklime, calcium oxide, which hisses and swells in water to become slaked lime; mixed with sand, slaked lime sets as mortar by taking the carbon dioxide back from the air. Heat limestone with clay at 1,450 °C and the product is Portland cement, of which the world now makes over four billion tonnes a year. The metal itself is made in small amounts by heating lime with aluminium under vacuum and condensing the calcium vapour; it serves mainly to strip oxygen and sulphur from steel and to reduce uranium from its compounds.

Bones and signals

An adult body holds about a kilogram of calcium, 99 % of it as hydroxyapatite, a calcium phosphate, in bone and teeth. Bone is not a fixed store: it is dissolved and rebuilt throughout life, and when the diet runs short the body draws on it, which is how osteoporosis begins. The remaining 1 % does the more surprising work. Hormones hold the calcium in the blood within narrow limits because the ion is a signal: a pulse of calcium released inside a muscle fibre makes it contract, calcium entering a nerve ending makes it release its transmitter, and calcium sets off the clotting of blood and the first divisions of a fertilised egg. Milk gives about 120 mg in every 100 mL; an adult needs close to a gram a day.

Discovery and name

Lime, mercury and a battery

Lime was known to the Egyptians, described by Vitruvius and used by the Romans, whose concrete of lime and volcanic ash still holds up the dome of the Pantheon. Lavoisier in 1789 listed lime among the “earths” he suspected of being oxides of metals that no fire could reduce. The proof came in 1808, once Davy's battery had taken potassium and sodium out of their alkalis. Berzelius and Pontin in Stockholm electrolysed moist lime against a pool of mercury and obtained an amalgam; Davy did the same in London, distilled the mercury off and was left with the metal. He named it calcium, from the Latin calx, lime — the word behind “calculus”, a pebble used for counting. Pure calcium came only in 1898, from Henri Moissan.