Determining Electronic Configuration (CIE A Level Chemistry)

• Writing out the electronic configuration tells us how the electrons in an atom or ion are arranged in their shells, sub-shells and orbitals
• This can be done using the full electron configuration or the shorthand version
  • The full electron configuration describes the arrangement of all electrons from the 1s sub-shell up
  • The shorthand electron configuration includes using the symbol of the nearest preceding noble gas to account for however many electrons are in that noble gas
• Ions are formed when atoms lose or gain electrons
  • Negative ions are formed by adding electrons to the outer sub-shell
  • Positive ions are formed by removing electrons from the outer sub-shell
  • The transition metals fill the 4s sub-shell before the 3d sub-shell but lose electrons from the 4s first and not from the 3d sub-shell
    • Remember: The 4s sub-shell is lower in energy
• The Periodic Table is split up into four main blocks depending on their electronic configuration:
  • s block elements
    • Have their valence electron(s) in an s orbital
  • p block elements
    • Have their valence electron(s) in a p orbital
  • d block elements
    • Have their valence electron(s) in a d orbital
  • f block elements
    • Have their valence electron(s) in an f orbital

The blocks of the Periodic Table

The elements can be divided into the s, p, d or f block, according to their outer shell electron configuration

Exceptions

• Chromium and copper have the following electron configurations, which are different to what you may expect:
  • 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 stable