The s-Block Elements

The s-block is Groups 1 and 2 — the elements whose highest electrons sit in an s orbital. Group 1 (alkali metals) have one outer electron and form +1 ions; Group 2 (alkaline earth metals) have two and form +2 ions. All are reactive metals, soft (Group 1 especially), with low densities and low melting points for metals.

Because they so readily lose their outer electrons, they are never found free in nature — always as compounds — and Group 1 metals are stored under oil to keep air and water away.

Reactivity increases down both groups: the outer electrons sit further from the nucleus and are more shielded, so they are lost more easily. Group 1 is more reactive than Group 2 in the same period, because losing one electron is easier than losing two.

With water, Group 1 metals react vigorously to give a hydroxide and hydrogen (sodium fizzes, potassium ignites); Group 2 react less violently. With oxygen they form oxides, and many s-block compounds give characteristic flame colours — sodium yellow, potassium lilac, calcium brick-red — used to identify them.

Their compounds are mostly ionic, white, and soluble (Group 1 especially). Thermal stability of carbonates and nitrates increases down a group. A subtle point examiners like: the first element of each group is anomalous and resembles the element diagonally below-right — the diagonal relationship (lithium resembles magnesium, beryllium resembles aluminium) because their charge densities are similar.

Heat an s-block compound in a flame and it glows a characteristic colour — lithium red, sodium yellow, potassium lilac, calcium brick-red, barium green. The colour is a direct window into the atom: the flame's energy promotes electrons to higher levels, and as they fall back they emit light of a fixed wavelength set by the size of the energy gap.

Because each element has its own set of gaps, each gives its own colour — the basis of the flame test and of the sodium glow in street lamps.

Going down Group 2, the carbonates and nitrates become more thermally stable — they need a higher temperature to decompose. The reason is polarising power: a small, highly charged cation distorts the anion and helps it break up, and as the cations get larger down the group they polarise less, so the compounds hold together better.

Solubility trends run in opposite directions: Group 2 hydroxides become more soluble down the group, while the sulfates become less soluble — which is why barium sulfate is insoluble enough to use as the white precipitate in the sulfate test.