Final answer:
Effective nuclear charge (Zeff) is the net positive charge experienced by valence electrons, influencing the atoms' ability to attract electrons in a bond. It explains the increase in electronegativity across a period due to increased nuclear charge with consistent shielding, and the decrease down a group due to added electron shells and increased shielding.
Step-by-step explanation:
The effective nuclear charge (Zeff) explains the trend in electronegativity across the periodic table. Electronegativity increases from left to right and decreases from top to bottom. This happens because, as you move across a period, the number of protons (nuclear charge, Z) increases, but the inner electron shells provide a consistent shielding effect, making the effective nuclear charge felt by valence electrons stronger. Therefore, the nucleus can pull the bonding electrons closer more efficiently. Conversely, as you go down a group, the number of electron shells increases which results in greater shielding and a larger atomic radius, thus reducing the effective nuclear charge and the atom's ability to attract bonding electrons.
Pauling's electronegativity values follow these predictable periodic trends, with values increasing towards the upper right of the periodic table. The higher electronegativity in this area is due to the greater effective nuclear charge experienced by the electrons involved in bonding.