Electronegativity and bonding

A greater nuclear charge (more protons) leads to a stronger attraction for electrons, thus increasing electronegativity. The positive charge of the nucleus is more effective at attracting electrons.

A larger atomic radius means the valence electrons are further from the nucleus, experiencing weaker attraction. This decreases the atom's ability to attract electrons in a bond, thus reducing electronegativity.

Increased shielding from inner electron shells reduces the effective nuclear charge felt by valence electrons. This weakens the attraction for electrons in a bond, thus decreasing electronegativity.

Electronegativity generally increases across a period (left to right). This is because nuclear charge increases while shielding remains relatively constant, leading to a stronger attraction for bonding electrons.

Electronegativity generally decreases down a group. This is because atomic radius and shielding increase, reducing the effective nuclear charge felt by bonding electrons and thus weakening the attraction.

A large difference in electronegativity (typically > 1.7) between two atoms indicates a likely ionic bond, where one atom strongly attracts the electrons from the other. A small difference indicates a covalent bond with shared electrons.