Alkaline Earth Metals

Group 2 is the column that has to spend more to get less. Each member carries two electrons in an outer s subshell, and removing both costs far more than removing one: magnesium's first ionisation energy is 738 kJ/mol and its second is 1,450, so an Mg²⁺ ion is over two thousand kilojoules per mole more expensive to make than an Mg⁺ ion. Yet Mg⁺ compounds do not exist and Mg²⁺ compounds are everywhere.

The reason is that ionisation energy is only one line of the ledger. A doubly charged cation attracts anions four times as strongly in the lattice as a singly charged one of the same size, and that extra lattice energy — or hydration energy, in solution — more than covers the second ionisation. Group 2 therefore has exactly one working oxidation state, arrived at by a route that looks uneconomic until you finish the accounting. It is the clearest demonstration in the table that stable compounds are settled by energy cycles rather than by counting electrons.

Why "earth", and why "alkaline"

The name is a fossil of eighteenth-century laboratory vocabulary. An "earth" was a substance that refused to melt in the hottest available furnace and would not dissolve in water — the oxides lime, magnesia, baryta and strontia all qualified. They were called alkaline because their solutions, weak as they were, turned indicators the way soda and potash did.

Humphry Davy separated four of them into metals during 1808 using the electrolytic approach that had already given him potassium and sodium, though the group 2 earths fought harder and his products were often amalgams rather than clean metal. Beryllium was recognised earlier, in 1798, when Louis-Nicolas Vauquelin found a new earth in beryl, but it took until 1828 for Friedrich Wöhler and Antoine Bussy, working independently, to reduce it to the metal. For most of the nineteenth century that element had two names in circulation: glucinium, from the sweet taste of its salts, was preferred in France, while beryllium was preferred in Germany. IUPAC did not settle the matter formally until 1949, and the French name survived in some journals for decades after.

Two solubility trends that run in opposite directions

The single most useful thing to know about group 2 is that its salts do not behave uniformly as you descend, and the direction depends on the anion.

• Hydroxides become more soluble down the column. Magnesium hydroxide is barely soluble; calcium hydroxide gives the faintly alkaline solution known as limewater; barium hydroxide dissolves well enough to be used as a strong base.
• Sulfates become less soluble down the column. Magnesium sulfate dissolves freely, calcium sulfate only sparingly, and barium sulfate is one of the most stubbornly insoluble common salts there is.

Both trends come out of the same competition, run with different numbers. Dissolving a salt means paying its lattice energy and recovering hydration energy for the separated ions. Lattice energy depends on the sum of the ionic radii and hydration energy on each ion separately, so when the anion is small — hydroxide — the cation's size dominates the lattice term and larger cations give weaker lattices. When the anion is large — sulfate — the lattice energy hardly changes down the group while the cation's hydration energy falls away steeply, and solubility collapses with it.

That is not an abstract result. Barium is an element that would poison a person outright as a soluble salt, and barium sulfate is swallowed by the litre as radiographic contrast medium precisely because its solubility is low enough that essentially none of the barium enters the body. The trend is the safety margin.

Beryllium barely belongs

Every column has an odd member and group 2's is unusually far out of line. Beryllium's ion would be so small and so highly charged that it distorts the electron cloud of anything it approaches, which means it rarely forms a clean Be²⁺ ion at all.

The consequences run right through its chemistry. Beryllium chloride is a covalent, chain-polymeric solid that sublimes rather than a high-melting ionic salt. Beryllium oxide is amphoteric, reacting with acids and with strong alkali, where magnesium oxide and everything below it are simply basic. Beryllium consistently takes a coordination number of four, while calcium happily reaches six or eight. And beryllium forms a strongly bonded, adherent surface oxide that makes the metal far less reactive than its position predicts.

Beryllium also has a nuclear career unrelated to any of that. Because it is the lightest element that is both a solid and reasonably transparent to X-rays, it makes the windows of X-ray tubes and detectors, a job for which nothing else is a close substitute. And it was beryllium bombarded with alpha particles that produced the penetrating neutral radiation James Chadwick correctly identified as the neutron in 1932 — Walther Bothe and Herbert Becker had generated the same radiation two years earlier and read it as gamma rays, and the Joliot-Curies in Paris repeated that misreading. The particle that completed the model of the atom came out of the top of group 2.

Its resemblance to aluminum — the diagonal partner one place right and one place down — is closer than its resemblance to magnesium. Beryllium is also the one member of the column that is seriously toxic on its own account rather than as a heavy-metal ion: chronic beryllium disease is a granulomatous lung condition caused by an immune response to inhaled beryllium compounds, and it occurs at exposures far below those that would trouble the other members of the group.

The column that builds skeletons and rocks

Group 2 is structural chemistry at planetary scale. Calcium carbonate is the mineral of limestone, chalk, marble, coral and shell, and it is the reason a significant fraction of the Earth's crust is sedimentary at all. Vertebrate bone is calcium phosphate as hydroxyapatite, and the same ion doubles as the trigger for muscle contraction and for a long list of signalling cascades — cytosolic calcium concentration is held roughly ten thousand times lower than extracellular, and releasing that gradient is how a cell says "now".

Magnesium sits at the centre of the chlorophyll molecule, holding the position that iron holds in haem, so every gram of carbon fixed by photosynthesis passes through a group 2 coordination complex. It is also the counter-ion for ATP: what enzymes actually bind is the Mg-ATP complex, not bare ATP.

Strontium and barium have no biological role, but strontium's chemical similarity to calcium gave it a dark twentieth-century career. Strontium-90 comes out of fission in quantity and takes some twenty-nine years to halve, and because the body handles it as if it were calcium it concentrates in bone rather than passing through. That single fact drove much of the public campaign against atmospheric nuclear testing in the late 1950s.

Radium's story is more direct still. Marie and Pierre Curie isolated it in 1898 from pitchblende residues, processing tonnes of ore for fractions of a gram. Its luminous salts were used to paint watch and instrument dials, and the young women employed to do that painting in the United States during and after the First World War developed severe bone disease — the litigation that followed established, in law, an employer's liability for occupational illness. Radium sits directly below barium and follows calcium into bone for the same chemical reason strontium does.

Getting the metals out

Nobody smelts group 2 in the ordinary sense; the oxides are among the most thermodynamically stable compounds known, which is why they were called earths.

Magnesium is the interesting case, because for much of the twentieth century it was extracted literally from the sea. Seawater carries magnesium as its third most abundant ion, and treating it with lime — itself obtained by roasting oyster shells at the original Texas plant — precipitates magnesium hydroxide, which is converted to the chloride and electrolysed. Most current production uses the Pidgeon process instead, thermally reducing calcined dolomite with ferrosilicon under vacuum, which is more energy-intensive but simpler to build; the balance between the two routes has shifted several times on economics alone.

Calcium metal is a niche product made by aluminothermic reduction of lime, and its main use is as a reducing agent for other metals. Barium comes from barite, which the world consumes overwhelmingly not as a source of the element but as drilling-mud weighting agent — its density and its insolubility being the two properties that matter, and both being consequences of the same heavy, poorly hydrated cation.

Where this group is misread

The commonest error is importing group 1 intuitions wholesale. Reactivity does increase down group 2 for the same shielding reason it increases down group 1, but the starting point is much lower: these metals are far less reactive than their neighbours one column to the left, because two electrons must be removed instead of one.

A second error is treating the solubility rules as arbitrary facts to memorise. Once the two opposing trends are seen as the same energy competition with a different anion size, the exceptions stop needing a list.

A third is assuming that flame colours — the crimson, the brick red, the apple green — say something about the metals' bonding. They are electronic transitions in the free atoms produced in the flame, and their usefulness as a group 2 identification test is unrelated to anything else in this article.

Finally, hard water is often described as a group 2 nuisance, which understates it. The calcium and magnesium ions responsible are dissolved out of exactly the carbonate rocks that group 2 chemistry built, and the scale they deposit inside a kettle is that carbonate being put back.