Writing Electron Configurations (DP IB Chemistry): Revision Note

Writing Electron Configurations

• The electron configuration gives information about the number of electrons in each shell, subshell and orbital of an atom

• The subshells are filled in order of increasing energy

Electron Configuration Key

• Electrons can be imagined as small spinning charges which rotate around their own axis in either a clockwise or anticlockwise direction

Spin pair repulsion diagram

• Electrons with the same spin repel each other (spin–pair repulsion):

  • Electrons occupy separate orbitals within a subshell first to minimise repulsion, with spins aligned

  • They pair only after all orbitals are singly occupied, with opposite spins

• This is known as Hund’s Rule:

  • In a p subshell, electrons fill px, py, and pz orbitals singly before pairing.

 Hund's Rule

• Hund’s Rule:

  • Electrons fill degenerate orbitals (same energy) singly first, with parallel spins, to minimise repulsion

• The principal quantum number (n) indicates the energy level of an electron:

  • For example, 2p electrons are in the second shell, n = 2

• Pauli Exclusion Principle:

  • An orbital holds two electrons with opposite spins only

  • Electrons pair only when no empty orbital of the same energy is available

  • Pairing costs less energy than jumping to a higher orbital

• Each box represents an atomic orbital

• The boxes are arranged in order of increasing energy from lower to higher (i.e. starting from closest to the nucleus)

• The electrons are represented by opposite arrows to show the spin of the electrons

  • E.g. the box notation for titanium is shown below

Electron box notation for titanium diagram

How to write 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

s, p, d and f blocks in 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