In 'd' block, electrons are added to an inner 'd' orbital and this shields the outer 's' electrons from the increased nuclear charge. Therefore, the atomic radius decreases only slightly and electronegativity and ionisation energies increase only slightly.

The difference in energy between the second and third shells is less than between the first and second. By the time the fourth shell is reached, there is, in fact an overlap between the third and fourth shells. In other words, from scandium onwards, the orbitals of highest energy in the third shell (the 3d orbital) have higher energy than those of the lowest energy in the fourth shell (the 4s orbital). Hence, when writing the electronic configurations of these 'd' block elements we fill the 4s then the 3d orbitals.

Exceptions do occur:

Cr [Ar] 4s¹3d⁵

Cu [Ar] 4s¹3d¹⁰

This can be explained by the extra stability offered by full and half-filled 'd' orbitals.

This property can be illustrated by the examples below:

Fe²⁺ (3d⁶) is readily oxidised to Fe³⁺ (3d⁵)

Mn²⁺ (3d⁵) is not readily oxidised to Mn³⁺ (3d⁴)

When transition metals form ions, electrons are lost first from the 4s sub-shell rather than the 3d sub-shell. Thus Fe²⁺ ions have the electronic structure [Ar]3d⁶ rather than [Ar]4s²3d⁴.

This occurs due to the presence of electrons in the 3d level, these repel the 4s electrons even further from the nucleus. Therefore, the 4s electrons are pushed to a higher energy level, higher than 3d. Consequently, when transition atoms become ions, the electrons from the 4s level before the 3d.

This means that all transition metals will have similar chemical properties which are dictated by the behaviour of the 4s outer electrons.