How do the electron configurations of transition metals differ from those of other elements?

The electron configurations of many transition metals will differ slightly from those predicted by Hund's rule, the Aufbau principle, and the Pauli exclusion principle.

Applying these rules, for instance, would result in the electron configuration [Ar] 4s2 3d9 for copper. However, the electron configuration of copper actually moves one electron from the 4s sublevel into the 3d sublevel, resulting in [Ar] 4s1 3d10. The electron configuration of the copper 1+ ion is [Ar] 4s0 3d10, and the electron configuration of the copper 2+ ion is [Ar] 4s0 3d9. Numerous other transition elements exhibit electrons moving from the highest energy s sublevel into the d sublevel one energy below.

Transition metals have electron configurations that involve filling the d sublevel before completing the s sublevel, unlike other elements where the s sublevel is filled first. They also tend to exhibit a greater variety of oxidation states compared to other elements. Transition metals differ from other elements in that they have partially filled d orbitals in their valence shells, which allows them to exhibit multiple oxidation states and form complex ions.