Smaller atoms generally have higher electronegativity because electronegativity is a measure of an atom's ability to attract electrons towards itself in a chemical bond. Several factors contribute to this trend:
Effective nuclear charge: Smaller atoms have fewer electron shells, which means the electrons are closer to the positively charged nucleus. This results in a higher effective nuclear charge, meaning the attractive force between the protons in the nucleus and the valence electrons is stronger. As a result, smaller atoms can exert a greater pull on electrons, increasing their electronegativity.
Atomic size: Smaller atoms have a smaller atomic radius, which leads to a higher electron density. The valence electrons are closer to the nucleus and are more strongly attracted to it. This proximity allows smaller atoms to exert a greater influence on shared electrons in a chemical bond, making them more electronegative.
Shielding effect: Electrons in inner energy levels shield the outer electrons from the full positive charge of the nucleus. In smaller atoms, there are fewer inner electrons to shield the outer electrons, so the attractive force from the nucleus is less diminished. This lack of shielding enhances the attractive force between the nucleus and valence electrons, resulting in higher electronegativity.
It's important to note that electronegativity is a periodic property, meaning it generally increases from left to right across a period in the periodic table and decreases from top to bottom within a group. However, there can be exceptions and variations due to other factors such as electron configuration and bonding patterns.