Transition metal: Encyclopedia II - Transition metal - Electronic configuration

Transition metal - Electronic configuration

Main group elements prior to the appearance of the transition group elements in the periodic chart (ie, elements number 1 through 20) have no electrons in d orbitals, but only in the s and p orbitals. The 3rd period p block elements have empty d orbitals. In the fourth period from scandium to zinc, d-block elements fill up their d orbitals across the period. With the exception of the copper group and the chromium group, all d-block elements in the ground state have two electrons in their outer s orbital. The electronic configuration of the d-block elements is ns²(n-1)d^1-10, where n is the ground state principal quantum number.

The outer s orbitals in the d-block elements are at lower energy states than the d orbitals of the n−1 levels. As atoms always strive to be in states of lowest energy, s orbitals are filled up first. The copper (4s¹3d¹⁰) and chromium (4s¹3d⁵) exceptions, which have one electron in their outer orbital, occur because half- and fully-filled orbitals are more stable than any other configurations (this occurs when there are 5 or 10 electrons in the d-orbitals).

Scandium has one electron in its d orbital, and 2 electrons in its outer s orbital. As scandium's only ion (Sc³⁺) has no electrons in its d orbital it is clear that it does not have a 'partially filled d orbital', and is not a transition metal in the stricter sense. Similarly, zinc is not a transition metal in the stricter sense because its only ion, Zn²⁺, has a full d orbital, which does not participate in bonding.

Adapted from the Wikipedia article "Electronic configuration", under the G.N U Free Docmentation License. Please also see http://en.wikipedia.org/wiki