Electronic configuration and orbitalsEdexcel A-Level Chemistry: Revision notes

Ideas about electronic configuration came from experimental evidence.

• Atomic emission spectra show discrete lines, not a continuous band. Electrons can only have certain energies, so energy is quantised and electrons are arranged in quantum shells.
• Successive ionisation energies show large jumps when a new, inner shell is reached. The number of electrons before the first large jump gives the group, and the sizes of the groups of ionisation energies give the number of electrons in each shell.
• First ionisation energies of successive elements show small dips that cannot be explained by shells alone. For example, the fall from beryllium (900) to boron (801 kJ mol⁻¹) is evidence for sub-shells: the 2p electron of boron is higher in energy than the 2s electrons of beryllium.

The first four quantum shells hold up to 2, 8, 18 and 32 electrons. Each shell is divided into sub-shells.

• s sub-shell: 1 orbital, 2 electrons
• p sub-shell: 3 orbitals, 6 electrons
• d sub-shell: 5 orbitals, 10 electrons

An orbital is a region within an atom that can hold up to two electrons with opposite spins. An s orbital is spherical; a p orbital is dumb-bell shaped, with three p orbitals at right angles in each p sub-shell. Shell n = 1 has 1s only; n = 2 has 2s and 2p; n = 3 has 3s, 3p and 3d; n = 4 has 4s, 4p, 4d and 4f.

Electrons occupy the lowest energy orbitals available. The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. The 4s sub-shell fills before the 3d sub-shell.

Two rules apply in each sub-shell: electrons fill orbitals of equal energy singly before pairing, because this minimises repulsion, and two electrons in the same orbital have opposite spins.

In an electrons-in-boxes diagram each orbital is a box and each electron an arrow. Nitrogen, 1s²2s²2p³, has three single (unpaired) arrows with the same spin in its three 2p boxes.

Atoms up to Z = 36 are written in 1s notation, for example sulfur 1s²2s²2p⁶3s²3p⁴ and iron 1s²2s²2p⁶3s²3p⁶3d⁶4s². Two exceptions arise because half-full and full 3d sub-shells are stable: chromium is [Ar]3d⁵4s¹ and copper is [Ar]3d¹⁰4s¹.

s- and p-block ions are written by adding or removing electrons from the outer sub-shell: S²⁻ is 1s²2s²2p⁶3s²3p⁶, K⁺ is 1s²2s²2p⁶3s²3p⁶, Br⁻ is [Ar]3d¹⁰4s²4p⁶ and Ca²⁺ is 1s²2s²2p⁶3s²3p⁶.

Elements are classified by the sub-shell being filled by the outer electron: s block (Groups 1 and 2), p block (Groups 3 to 0), and d block (the transition metals between Groups 2 and 3). The outer configuration determines chemical properties: Group 1 elements, ns¹, lose one electron, while Group 7 elements, ns²np⁵, gain one. Elements in the same group have similar chemical properties because they share the same outer configuration.