Final answer:
Electronegativity is an atom's ability to attract electrons, which is affected by the number of protons and the degree of electron shielding. Effective nuclear charge increases across a period due to the limited increase in shielding and the added protons, making elements more electronegative as they attract electrons more strongly.
Step-by-step explanation:
Understanding Electronegativity and Electron Shielding
Electronegativity is a measure of an atom's ability to attract and hold electrons and is influenced by both the number of protons in the nucleus and the degree of electron shielding. An important concept to grasp is that not all electrons shield equally. Electrons in filled inner shells are highly effective at shielding outer shell electrons from the full charge of the nucleus, leading to a decrease in effective nuclear charge (Zeff). On the other hand, electrons in the same principal shell are not very effective at shielding each other due to electron-electron repulsions.
The increase in protons (positive charge) as we move across a period is not fully counteracted by an increase in electron shielding, which only slightly rises, resulting in a higher Zeff. This stronger pull draws electrons closer to the nucleus.
Consequently, the most electronegative elements are found on the right side of the periodic table, as they experience a greater nuclear attraction due to lower electron shielding effect. Elements with an unfilled electron shell tend to have higher electronegativity because their nucleus can exert a stronger pull on the valence electrons.