Electronegativity is the ability of an atom in a molecule to attract shared electrons. As we move down a group in the periodic table, electronegativity tends to decrease. This trend can be attributed to a couple of important factors.
First, as we progress down a group, the number of electron shells increases. For example, elements in the same group have the same number of valence electrons but different numbers of total electron shells. This leads to greater distance between the positively charged nucleus and the valence electrons. Consequently, the attraction between the nucleus and the shared electrons decreases, resulting in lower electronegativity.
Secondly, there is an effect known as shielding or screening. The inner shells of electrons can shield the valence electrons from the full force of the nucleus’ positive charge. As more electron shells are added, the shielding effect becomes more pronounced. This further diminishes the nucleus’s ability to attract additional electrons, contributing to the decrease in electronegativity down a group.
In summary, the decrease in electronegativity down a group is primarily due to the increased distance between the nucleus and the valence electrons along with the increased shielding effect from inner electrons. These factors together make it less favorable for atoms lower in a group to attract shared electrons.