electronegativities

Electronegativities are a set of related concepts that describe how strongly an atom in a chemical bond attracts the electrons shared with another atom. The term and its practical scale were developed to explain bond polarity and reaction behavior. Electronegativities are not directly measurable as a single physical property; they are relative measures derived from experimental data such as bond energies, ionization energies, and electron affinities.

The most widely used scale is the Pauling scale, which assigns higher numbers to more electronegative elements. Other common scales include the Mulliken scale, which averages ionization energy and electron affinity, and the Allred–Rochow scale, which relates effective nuclear charge to covalent radius. On the Pauling scale, fluorine has the highest value (about 3.98), with oxygen, nitrogen, and chlorine also showing high values. Values for hydrogen and carbon are lower, reflecting weaker attraction of shared electrons.

General trends show electronegativity increasing across a period and decreasing down a group in the periodic table. The increase across a period is due to increasing effective nuclear charge and relatively constant shielding, while the decrease down a group results from greater atomic size and shielding, which reduce the nucleus’s pull on bonding electrons. Transition metals display more variable electronegativities that depend on oxidation state and coordination.

Electronegativity is a key predictor of bond polarity; large differences between bonded atoms indicate ionic character, while small differences yield nonpolar or polar covalent bonds. It also helps explain acidity, basicity, and reaction pathways. Limitations include dependence on the chosen scale, environment, and molecular context, and the fact that electronegativity is a relative, not intrinsic, property.