The noble gases (helium, neon, argon, krypton, xenon, and radon) are elements positioned in the far-right column of the periodic table, known for their exceptional lack of reactivity. The question of whether these chemically aloof elements possess electronegativity is central to understanding their behavior, as this property is fundamental to the formation of chemical bonds. While traditional chemistry often cites a value of zero or undefined for these elements, a deeper look reveals a more complex answer rooted in the nature of measurement and modern chemical theory.
What Electronegativity Measures
Electronegativity is defined as an atom’s inherent tendency to attract a shared pair of electrons toward itself when it is part of a chemical bond. This property is not a directly measurable physical quantity like mass or temperature, but rather a relative scale that indicates an atom’s electron-pulling power within a molecule. The difference in electronegativity between two atoms determines the polarity of the bond they form, classifying it as nonpolar covalent, polar covalent, or ionic. On the periodic table, electronegativity generally increases as you move from left to right across a period due to a rising nuclear charge, and it decreases as you move down a group because the valence electrons are farther from the nucleus.
The most widely used reference for this property is the Pauling scale, developed by Linus Pauling, which sets fluorine as the most electronegative element with a value of 4.0. This scale is based on the difference in bond dissociation energies between a heteronuclear bond (A-B) and the average of the homonuclear bonds (A-A and B-B). Because the Pauling scale relies on data from actual chemical bonds, it encounters a significant limitation when applied to elements that rarely or never form bonds under normal conditions.
Why Noble Gases Resist Bonding
The traditional view that noble gases do not have electronegativity stems from their remarkably stable electron configuration. All noble gases, except helium, have eight electrons in their outermost shell, a configuration known as a complete octet. This filled valence shell represents the most stable state for an atom, giving noble gases little to no energetic incentive to gain, lose, or share electrons with other atoms. This inherent stability results in exceptionally high ionization energies and electron affinities that are near zero or even positive, indicating they do not readily accept an additional electron. The lighter elements are simply too small and their outer electrons are too tightly held by the nucleus to be forced into a bond.
Calculated Values and Real-World Compounds
Theoretical Calculation
Despite the limitations of the Pauling scale, noble gases do possess a measurable, non-zero tendency to attract electrons, which can be determined through theoretical methods. Electronegativity can be calculated using the Mulliken scale, which defines the property as the average of an atom’s first ionization energy and its electron affinity. This approach bypasses the need for existing bond data, allowing chemists to assign theoretical electronegativity values to all elements, including the noble gases. Calculated Mulliken values for the noble gases show a trend of decreasing electronegativity from helium down to radon.
Evidence from Xenon and Krypton
Further confirmation that noble gases possess this property comes from the chemistry of the heavier members of the group, specifically krypton and xenon. Because of their larger atomic size, the outermost electrons of these atoms are shielded from the nucleus by many inner electron shells, making them less tightly bound than those in neon or argon. This weakened nuclear attraction allows highly electronegative elements, most commonly fluorine, to induce a reaction and form stable compounds like xenon difluoride (\(\text{XeF}_2\)) and xenon tetrafluoride (\(\text{XeF}_4\)). The existence of these covalent compounds proves that xenon and krypton possess a measurable capacity to attract electrons, which is necessary to form a bond with fluorine. For instance, xenon’s calculated Pauling-equivalent electronegativity is approximately 2.6, which is comparable to elements like iodine.