The atomic number of a transition metal is typically less than its group number due to the electronic configuration of its atoms. In the periodic table, the group number indicates the number of valence electrons an element has, which determines its chemical properties and reactivity.
Transition metals are characterized by their partially filled d orbitals, which are involved in bonding and determining their unique properties. When filling electrons into the orbitals, the 3d orbitals of the transition metals are filled after the 4s orbital.
For example, let's consider the element copper (Cu) as an illustration. Copper is located in Group 11 of the periodic table, but its atomic number is 29. The electronic configuration of copper is [Ar] 3d^10 4s^1. As per the Aufbau principle, the 4s orbital is filled before the 3d orbital. However, in the case of copper, one electron from the 4s orbital "jumps" to the 3d orbital to achieve a more stable configuration with completely filled 3d orbitals (d^10). This arrangement is energetically favorable, resulting in a configuration of [Ar] 3d^10 4s^1 instead of [Ar] 3d^9 4s^2.
Therefore, the atomic number of copper (29) is less than its group number (11) because of the electron configuration and the arrangement of electrons in the d orbitals of the transition metal. This phenomenon is observed in other transition metals as well.