electronegativity

Electronegativity is a chemical property that describes an atom’s tendency to attract electrons toward itself when forming a chemical bond. It is not directly measurable for a single isolated atom; instead, it is inferred from observable data such as bond dissociation energies, ionization energies, and electron affinities. Several scales quantify electronegativity; the Pauling scale is the most widely used, with values ranging roughly from 0.7 for cesium to about 4.0 for fluorine. Other scales include Mulliken, which defines electronegativity as half the sum of ionization energy and electron affinity (in electronvolts), and Allred-Rochow, which relates it to effective nuclear charge and atomic radius.

In the periodic table, electronegativity generally increases across a period from left to right and decreases down a group. This pattern reflects competing factors: across a period, increasing effective nuclear charge and smaller atomic radii enhance the attraction for bonding electrons; down a group, greater atomic size and shielding reduce this attraction.

Electronegativity differences between bonded atoms influence bond character. Small differences yield nonpolar covalent bonds, moderate differences produce polar covalent bonds, and large differences lead to bonds with substantial ionic character. For example, the H–Cl bond is polar covalent, whereas the bond in NaCl is largely ionic. Electronegativity also helps explain molecular polarity, reactivity, and trends in acidity and basicity in many chemical systems.

Limitations include context dependence: electronegativity values vary with oxidation state, coordination, and environment, and scales are defined for specific reference conditions. Nevertheless, electronegativity remains a central concept for predicting bond type and molecular behavior.