Historically, the periodic table has been been into blocks named after the orbital occupied by the elements' valence electrons. 

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The elements that lie between group 2 and group 13 are often referred to as the d-block elements. In this topic, we will concern ourselves with the elements of the third period: scandium to zinc.

The electron configurations of the 3rd period d-block elements are shown below:

Which electron configurations appear to be anomalous?

The electron configurations for chromium and copper seem anomalous as the 4s orbital is not doubly occupied. However, we know that electrons adopt the most energetically stable configuration and we must conclude that though appearing anomalous, they are more energetically stable. 

A further quandry arises when we examine the ionisation energies of the d-block elements. We find that it is the 4s orbital electrons that are ionised first before the 3d orbitals. We have to conclude that, ironically for the d-block elements, the 4s orbitals are in fact higher in energy than the 3d orbitals and the valence electrons are in the 4s orbtial.This raises another question.

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If the 4s orbitals are higher in energy, then why do electrons enter the 4s orbital before the 3d orbital is completely full? Why does scandium have the electronic configuration [Ar] 3d¹ 4s² and not [Ar] 3d³?

 We will go into this more deeply in Exploring Understanding, The Aufbau Principle.