Electronegativity

Electronegativity is the relative tendency of an atom, when it is part of a covalent bond, to attract the shared pair of bonding electrons toward itself. Unlike ionization enthalpy or electron gain enthalpy, electronegativity is not a property of an isolated atom — it only has meaning for an atom bonded to another atom — and it has no single agreed unit; instead it is reported on a comparative scale.

The most widely used scale is the Pauling scale, proposed by Linus Pauling from the extra ("excess") bond-dissociation energy of a polar $A-B$ bond compared with the average of the $A-A$ and $B-B$ bond energies. On this scale, fluorine is assigned the highest electronegativity of any element, $\chi_F = 4.0$, and every other element's electronegativity is measured relative to it; caesium and francium have among the lowest values, close to $0.7$.

Trend across a period. Electronegativity increases from left to right across a period. This follows the same effective-nuclear-charge argument used for atomic radius and ionization enthalpy: as nuclear charge rises across a period while the bonding electrons remain in the same outer shell, the nucleus pulls a shared electron pair toward itself more strongly, so the tendency to attract bonding electrons increases.

Trend down a group. Electronegativity decreases on descending a group. The bonding electrons in a heavier element's valence shell lie farther from the nucleus and are more shielded by inner filled shells, so the nucleus attracts a shared pair less strongly despite its larger charge. …