Determining Electronic Configuration (Cambridge (CIE) AS Chemistry)

Determining Electronic Configurations

• 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] 4s¹ 3d⁵ not [Ar] 4s² 3d⁴ 

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

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