Alkaline Earth Metals

Have you ever seen a hard seashell, a marble column, or fireworks lighting up the night sky? Early chemists used the word "earth" for natural substances that, like soil or stone, don't burn and barely dissolve in water.

Substances like lime and magnesia turn basic (alkaline) when they meet water, so these "earths" were called "alkaline earths." Then, around 1808, it was discovered that these earths were really compounds of a metal and oxygen. Once the metals inside were separated out, they were named alkaline earth metals.

Calcium (Ca), which makes up the bones and teeth in your body, and magnesium (Mg), which is light and strong enough to be used in smartphone casings, are the star members of this family. They're also key ingredients in cement and limestone, and they've helped lay the foundations of human civilization.

So alkaline earth metals show up everywhere, from rocks and soil underground to the bodies of living things and cutting-edge industry. They may look like plain, rough dirt, but they're actually helpful elements that quietly hold up our everyday lives.

Atoms like to fill their outer shell with 8 electrons to reach a stable state (the octet rule). Alkaline earth atoms are born with just two valence electrons in their outermost shell.

To reach 8, they'd have to hunt down 6 more electrons from somewhere. It's far easier and more stable to simply hand over the 2 they already have. So given the chance, they happily shed those 2 electrons and turn into positive ions (+2).

They react a bit less than the Group 1 alkali metals right next door, but they still react very easily with air and water. That's why they never show up in nature as plain lumps of pure metal. They're always found locked up in minerals, bonded with oxygen, carbonate, and other partners.

Because of this, getting pure magnesium or calcium metal takes special smelting methods, such as running electricity through a molten compound to separate it, or using another substance at high temperature to pull the oxygen away. Their urge to give up electrons and pair with other elements is that strong.

As you move down a column of the periodic table, atoms get bigger and bigger. The outer electrons sit farther from the nucleus, so the nucleus holds on to them more weakly and they pop off much more easily.

That's why chemical reactivity gets noticeably fiercer as you go down, from beryllium (Be) to radium (Ra). Beryllium at the top barely reacts with water at room temperature. Calcium, though, dissolves in cold water while bubbling out hydrogen gas, and barium reacts with oxygen in the air so quickly that it's stored under oil.

Compared with the Group 1 alkali metals, alkaline earth metals give up 2 electrons when they bond, so the metallic bonds holding their atoms together are stronger. As a result, they have much higher melting points and are harder than alkali metals. Even so, calcium, strontium, and barium near the bottom are still soft enough to cut with a knife.

The "flame test," where a metal gives off its own color in a flame, is another charm of alkaline earth metals. Calcium burns orange, strontium deep red, and barium yellow-green, which makes them the stars behind the colorful fireworks at night festivals.