Determining Electronic Configuration (Cambridge (CIE) A Level Chemistry): Revision Note

Determining Electronic Configurations

• Electron configuration shows how electrons are arranged in shells, subshells, and orbitals.

• There are two formats:

  • Full configuration; lists all electrons from 1s onward

  • Shorthand configuration; uses the symbol of the nearest noble gas in brackets to represent inner electrons (e.g. [Ar])

• Ions form when atoms gain or lose electrons:

  • Anions (negative) form by adding electrons to the outer shell

  • Cations (positive) form by removing electrons from the outer shell

• Transition metals:

  • Fill the 4s before 3d when neutral

  • Lose electrons from 4s first, not 3d, when forming ions

• In the Periodic Table the elements are grouped into blocks based on their valence subshell:

  • s-block: valence electrons in an s orbital

  • p-block: valence electrons in a p orbital

  • d-block: valence electrons in a d orbital

  • f-block: valence electrons in an f orbital

The blocks of the Periodic Table

Examples

• Electronic configuration of Fe

  • Atomic number = 26 so there are 26 electrons

  • Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

  • Shorthand: [Ar] 4s² 3d⁶

• Electronic configuration of Fe²⁺

  • Atomic number = 26 so there are 26 electrons, but the Fe²⁺ ion only has 24 electrons

  • Electrons are removed from the 4s orbital before the 3d

  • Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶

  • Shorthand: [Ar] 3d⁶

Exceptions to the Aufbau principle

• Chromium and copper have the following electron configurations:

  • Cr is [Ar] 3d⁵ 4s¹ not [Ar] 3d⁴ 4s²

  • Cu is [Ar] 3d¹⁰ 4s¹ not [Ar] 3d⁹ 4s²

• This is because the [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹ configurations are energetically favourable

• By promoting an electron from 4s to 3d, these atoms achieve a half full or full d-subshell, respectively