Electronic configuration and the blocks

The periodic table follows directly from the electronic configurations of the elements. A simple tip: the number of shells tells you the period, and the outer electrons tell you the group.

Each new period starts filling a new shell. The number of elements in a period equals twice the number of orbitals being filled. Period 5 fills 5s, 4d and 5p: 1 + 5 + 3 = 9 orbitals, so 18 elements. Period 6 also fills the 4f orbitals: 6s, 4f, 5d, 6p, giving 1 + 7 + 5 + 3 = 16 orbitals.

The subshell that receives the last electron names the element's block:

• s-block (groups 1-2): outer configuration ns¹ or ns². Reactive metals with low ionisation enthalpy. Sodium, whose last electron enters 3s, is an s-block element in period 3.
• p-block (groups 13-18): ns² np¹ to ns² np⁶. It contains metals, non-metals and metalloids, and ends with the noble gases, whose filled shells make them very unreactive.
• d-block (groups 3-12): (n – 1)d¹⁻¹⁰ ns⁰⁻². The transition metals: they often form coloured ions, show variable valence, and are good catalysts.
• f-block: (n – 2)f¹⁻¹⁴ (n – 1)d⁰⁻¹ ns². The lanthanoids and actinoids, the 'inner transition' elements.

Elements can also be grouped by type:

• Metals, more than 78% of all elements, are on the left side of the table.
• Non-metals are at the top right.
• Metalloids (semi-metals), such as silicon, germanium, arsenic, antimony and tellurium, lie along the zig-zag line between them, with properties of both.

Metallic character increases down a group and decreases across a period from left to right. So P < Si < Be < Mg < Na in metallic character.