Electron Configuration & The Periodic Table

• Ground state 


  This is the lowest possible electronic configuration the electrons in an atom can adopt
• Oxidation number 


  The formal charge assigned to each atom in a compound according to certain rules
• What is the rule for filling and emptying the 4s and 3d subshells?

  The 4s subshell is always filled before the 3d and the 4s subshell is always emptied before the 3d.
• What is the exception to the Aufbau principle when filling the 3d orbitals?

  Chromium and Copper
• Why are Chromium and Copper the exception to the rule?

  There is extra stability associated with a half-filled (Chromium) and a filled (Copper) d-subshell
• The Aufbau principle 


  Electrons fill orbitals in order of increasing energy (‘Aufbau means ‘building up’ in German)
• Hund’s rule


  When degenerate orbitals are available, electrons fill each singly, keeping their spins parallel before spin pairing starts.
• The Pauli Exclusion Principle


  No two electrons in one atom can have the same set of four quantum numbers, therefore, no orbital can hold more than two electrons and these two electrons must have opposite spins.
• In an isolated atom the orbitals within each subshell are degenerate.
• The periodic table is subdivided into four blocks (s, p, d and f) corresponding to the outer electronic configurations of the elements within these blocks.
• The variation in first, second and subsequent ionisation energies with increasing atomic number for the first 36 elements can be explained in terms of the relative stability of different subshell electronic configurations. This provides evidence for these electronic configurations. Anomalies in the trends of ionisation energies can be explained by considering the electronic configurations.
• There is a special stability associated with half-filled and full subshells. The more stable the electronic configuration, the higher the ionisation energy.