Why does electronegativity increase across a period but decrease down a group?

Electronegativity increases across a period due to increasing nuclear charge and decreases down a group due to increasing atomic radius.

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. As you move across a period from left to right on the periodic table, the number of protons in the nucleus (nuclear charge) increases. This increased positive charge attracts the electrons in the bond more strongly, thus increasing the atom's electronegativity. At the same time, the number of energy levels (shells) remains the same, so the increase in nuclear charge is not shielded from the bonding electrons. This results in a stronger pull on the bonding electrons, increasing the atom's electronegativity.

On the other hand, as you move down a group on the periodic table, the atomic radius (the distance from the nucleus to the outermost shell of electrons) increases. This is due to the addition of more energy levels or shells. The increased distance between the nucleus and the bonding electrons reduces the attractive force exerted by the nucleus on the electrons. This effect is known as shielding. The increased shielding and distance reduce the ability of the atom to attract bonding electrons, thus decreasing its electronegativity.

To better understand these concepts, you can explore how trends in ionisation energy affect electron attraction, how the periodic trend of electron affinity impacts an atom's ability to gain electrons, and the role of atomic radius in influencing electronegativity.

IB Chemistry Tutor Summary: Electronegativity measures how strongly an atom attracts bonding electrons. Across a period, electronegativity increases because the nuclear charge rises without additional shielding, pulling electrons closer. Down a group, it decreases as atomic radius grows and shielding reduces the nucleus's pull on electrons. Understanding these changes helps grasp chemical bonding dynamics.