Electron configuration and the periodic table

Filling the shells

Electrons occupy shells (energy levels) around the nucleus. The first three shells hold:

• Shell 1: up to 2 electrons
• Shell 2: up to 8 electrons
• Shell 3: up to 8 electrons (for the first 20 elements at GCSE)

Electrons fill from the lowest shell upwards.

Examples:

• Sodium (Z=11): 2,8,1
• Chlorine (Z=17): 2,8,7
• Calcium (Z=20): 2,8,8,2

Write configurations with commas: 2,8,1. Each comma is a shell boundary.

Periodic table layout

• Group number = number of electrons in the outer shell (for groups 1-7 and 0).
• Period number = number of occupied shells.

Sodium (2,8,1) is in Group 1 (one outer electron) and Period 3 (three shells).

Trends down a group

Going down a group, atoms gain an extra shell each period. The outer electrons are:

• Further from the nucleus
• Shielded by more inner electrons

Consequences:

• Group 1 (alkali metals): reactivity increases down the group — easier to lose the outer electron.
• Group 7 (halogens): reactivity decreases down the group — harder to gain an electron.
• Group 0 (noble gases): full outer shell, very unreactive.

WJEC exam tip

When predicting reactivity, always reference the outer-shell electron AND the distance/shielding. "Lithium reacts less vigorously than sodium because its outer electron is closer to the nucleus and less shielded, so it is held more strongly and harder to lose."