Periodic trends in chemical properties

The valence (combining capacity) of main-group elements follows a simple pattern:

• Groups 1, 2, 13 and 14: valence = number of valence electrons (1, 2, 3, 4).
• Groups 15, 16 and 17: valence = 8 minus the number of valence electrons (3, 2, 1).

To write a formula, cross over the valences. Silicon (4) and bromine (1) give SiBr₄. Magnesium (2) and nitrogen (3) give Mg₃N₂.

Oxidation state and the number of bonds (covalency) can differ. In the ion [AlCl(H₂O)₅]²⁺, aluminium has an oxidation state of +3, but it forms 6 bonds.

The second-period elements (lithium to fluorine) often behave differently from the rest of their groups (anomalous behaviour). They are small, highly electronegative, and have only four valence orbitals (2s and 2p), so their maximum covalency is 4: boron forms BF₄⁻, while aluminium, with d orbitals available, forms AlF₆³⁻. They also form strong pπ–pπ multiple bonds (C=C, N≡N, O=O). Some second-period elements resemble the element diagonally below them in the next group: lithium resembles magnesium, and beryllium resembles aluminium (the diagonal relationship).

Reactivity is highest at the two ends of a period: group 1 metals lose an electron very easily, and group 17 non-metals gain one very easily. Down group 1, reactivity increases (the outer electron is farther away and easier to lose). Down group 17, it decreases (the incoming electron is held less strongly).

The oxides show the same left-to-right change:

• On the left, oxides are basic: Na₂O + H₂O → 2NaOH.
• On the right, they are acidic: Cl₂O₇ + H₂O → 2HClO₄.
• In the middle, some are amphoteric (reacting with both acids and bases), such as Al₂O₃, and some are neutral (neither acidic nor basic), such as CO, NO and N₂O.