Alkaline earth metals are the steadier cousins of the alkali metals. Sitting in Group 2, they have two electrons in their outer shell rather than one, which makes them reactive — but not wildly so. You won't see magnesium exploding in your water glass, though it does burn with a brilliant white flame so intense it was once used in photographic flash powder. These elements got their unusual name from a historical quirk: their oxides ("earths" in old chemistry) were found to be alkaline, and unlike the alkali metal oxides, they didn't dissolve easily in water. The name stuck even after chemists isolated the pure metals.
The range of properties across this group is fascinating. Beryllium is a lightweight, steel-gray metal so stiff that it's used in aerospace components and X-ray windows. Magnesium is the lightest structural metal, alloyed with aluminum in aircraft and laptop casings. Calcium, the most abundant alkaline earth metal in the human body, is the building block of bones and teeth — your skeleton contains about 1 kg of calcium. Strontium compounds produce the deep red color in fireworks and flares. Barium gives us the chalky "barium meal" that makes your digestive tract visible on X-rays. Radium, discovered by Marie Curie in 1898, is intensely radioactive and glows faintly blue in the dark.
Chemically, alkaline earth metals form +2 ions by losing both valence electrons. Their compounds are everywhere: calcium carbonate forms limestone, chalk, marble, and seashells; magnesium hydroxide is the active ingredient in milk of magnesia; and calcium sulfate is the plaster of Paris used in casts and construction. Globally, about 4.5 billion tonnes of cement (largely calcium compounds) are produced annually — making calcium chemistry one of humanity's biggest industrial operations.