Alkaline earth metals

The metals behind bones, chalk, fireworks and the brilliant white flare of burning magnesium.

Your bones and teeth are built around calcium. The green colour of every leaf depends on magnesium. The brilliant white light of a sparkler comes from burning magnesium ribbon, and the deep red of a firework is strontium. All of these elements belong to group 2, the alkaline earth metals: beryllium, magnesium, calcium, strontium, barium and radium.

The odd name is a leftover from early chemistry. Alchemists used the word earth for any solid that did not melt or change when heated in a fire. The oxides of these metals were earths that also happened to be alkaline, meaning they neutralised acids.

What they have in common

Each alkaline earth metal has two electrons in its outer shell. It loses both to form an ion with a charge of plus two, and that two-plus ion is the form in which the metal appears in nearly all of its compounds. Losing two electrons takes more energy than losing one, so these metals are noticeably less reactive than their neighbours in group 1. They can be handled in air, although they slowly tarnish, and none of them needs to be kept under oil.

They are harder, denser and higher melting than the alkali metals. Beryllium melts at about 1287 degrees Celsius and is stiff enough to be used in precision instruments. Calcium can still be scratched with a knife, but it does not smear like sodium.

Their compounds are typically white ionic solids, but with a twist compared with group 1. Many of them are poorly soluble in water. Calcium carbonate, the stuff of chalk, limestone and marble, hardly dissolves at all, which is why these rocks survive for millions of years. Barium sulfate is so insoluble that patients can safely swallow it before an X-ray even though barium ions on their own are poisonous.

How they change down the group

Reactivity increases down the group, for the same reason as in group 1. Larger atoms hold their outer electrons more loosely, so they are lost more easily. Beryllium barely reacts with water at all. Magnesium reacts very slowly with cold water but burns vigorously in steam. Calcium fizzes steadily in cold water, producing hydrogen and a cloudy suspension of calcium hydroxide, and strontium and barium react more quickly still.

The trend in melting points is less tidy than in group 1. Beryllium and magnesium melt at about 1287 and 650 degrees Celsius, calcium at about 842 degrees, strontium at about 777 degrees and barium at about 727 degrees. The general direction is downwards but the steps are uneven, partly because the metals pack their atoms in different crystal arrangements.

Density rises down the group, from beryllium and magnesium, which are among the lightest structural metals, to barium at around three and a half times the density of water and radium heavier still. Flame colours also shift: calcium gives an orange-red flame, strontium a crimson one and barium an apple green.

Where you meet them

Calcium is everywhere. It is in limestone hills, coral reefs, egg shells, cement, plaster and toothpaste, as well as in your skeleton. Magnesium is almost as familiar. Its alloys are light and strong, so they turn up in bicycle frames, laptop cases and aircraft parts. Magnesium sits at the centre of the chlorophyll molecule that lets plants capture sunlight, and it is a common ingredient in indigestion remedies.

Strontium and barium are best known for colouring fireworks and flares. Barium compounds also go into some glass and paints, and barium sulfate is the barium meal used to make the gut show up on medical X-rays. Beryllium is used in small amounts in copper alloys for springs and non-sparking tools, and in windows for X-ray equipment because X-rays pass through it easily. Its dust is toxic, so it is handled with care.

Odd ones out

Beryllium is the member that least fits the pattern. Its atoms are so small that they hold their two electrons unusually tightly, so beryllium tends to share electrons rather than form simple ions. Its compounds behave more like those of aluminium in group 13 than like those of calcium. It is also the only member of the group whose compounds are dangerously poisonous to breathe in as dust.

Radium is odd because it is intensely radioactive. It is formed as uranium decays in rocks, and it glows faintly in the dark as its radiation excites the air around it. In the early twentieth century it was painted onto watch dials and even added to health products before the dangers were understood. Today it has almost no uses, and its place in the story of chemistry is as a warning as much as a discovery.

The numbers

Melting points fall down the list, from Beryllium at 1287 °C to Radium at 960 °C.

Density in g/cm³ (g/L for gases). Values for synthetic elements are predictions.

Did you know?

• Calcium carbonate makes up chalk, limestone, marble, coral, egg shells and the shells of many sea creatures.
• Magnesium burns with a light so bright and white that it was once used in camera flashes.
• Beryllium melts at about 1287 degrees Celsius, far higher than any alkali metal.
• Strontium colours fireworks crimson and barium colours them green.
• Barium sulfate is so insoluble that it can be swallowed safely for X-rays, even though free barium ions are toxic.

Frequently asked questions

Written by Anthony. Data from Periodic-Table-JSON (CC BY-SA).