Electronegativity is a fundamental chemical property that quantifies an atom’s tendency to attract electrons toward itself within a chemical bond. It is a dimensionless value on a relative scale that compares the electron-attracting power of different elements. The most widely recognized system is the Pauling scale, which assigns values from approximately 0.7 to 3.98. On this scale, a higher number signifies a stronger pull on a shared electron pair.
Electronegativity generally increases moving left to right across a period because the nuclear charge increases, pulling the valence electrons closer. Conversely, it decreases moving down a group because atoms become larger, and the outer electrons are farther from the nucleus. Fluorine, positioned in the upper right corner of the main group elements, is the most electronegative element with a Pauling value of 3.98.
The Stable Structure of Noble Gases
Noble gases, found in Group 18, are generally considered non-electronegative and are assigned a value of zero on the standard Pauling scale. This stems from their unique and highly stable electron configuration. With the exception of helium (two valence electrons), all noble gases possess a full outer shell containing eight electrons, a state described by the octet rule. This complete valence shell configuration represents the lowest possible energy state, granting them exceptional chemical inertness.
Since they already have a full set of outer electrons, noble gases have virtually no energetic incentive to gain electrons from other atoms. Electronegativity concerns an atom’s ability to attract electrons in a chemical bond, but noble gases rarely form bonds under normal conditions. Their stability means they do not participate in the electron sharing or transfer reactions that the Pauling scale is designed to measure. Consequently, for lighter noble gases like neon and argon, the standard electronegativity concept is largely irrelevant.
Bonding Exceptions and Theoretical Values
The traditional view of noble gases as completely non-reactive began to change with the synthesis of xenon compounds in the 1960s. While light noble gases like helium and neon remain chemically aloof, the heavier elements, such as krypton, xenon, and radon, can be coaxed into forming compounds with highly electronegative partners, primarily fluorine and oxygen. This slight increase in reactivity is due to their larger atomic radii, which means their outermost electrons are held less tightly by the nucleus and are thus more available for sharing.
Since bonding is possible for these heavier elements, scientists have assigned non-zero electronegativity values to them, moving away from the traditional Pauling zero. Xenon, for instance, has a Pauling electronegativity value of approximately 2.6, which is comparable to that of sulfur or iodine. Krypton is sometimes assigned a value near 3.0. These values are often derived from theoretical calculations or alternative scales, like the Mulliken or Allen scales. These scales define electronegativity based on an atom’s ionization energy and electron affinity rather than bond energies. This theoretical assignment reflects their power to attract electrons when they are forced into a chemical partnership.