Electronegativity is a chemical property used to quantify an atom’s tendency to attract electrons when involved in a chemical bond. Elements in Group 18, the noble gases, are unique because they are generally unreactive and do not readily form the bonds necessary for this property to be measured directly. To fully understand their status, it is necessary to examine the definition of electronegativity, the inherent stability of the noble gases, and the exceptions that occur with their heavier members.
Understanding Electronegativity and Chemical Bonding
Electronegativity is defined as the measure of an atom’s ability to attract a shared pair of electrons toward itself within a chemical bond. This is a relative scale that compares the electron-attracting power of different elements when they are bonded together. The difference in electronegativity between two atoms determines the nature of the bond, classifying it as nonpolar covalent, polar covalent, or ionic.
The most widely used system is the Pauling scale, which calculates electronegativity based on the energy difference between a heterodiatomic bond (A-B) and the average energy of the two homodiatomic bonds (A-A and B-B). Since the calculation depends on bond energies, an element must form a stable chemical bond for its Pauling value to be directly measured. Without a shared pair of electrons, the Pauling method cannot be applied, which historically excluded most of the noble gases from the scale.
The Stability of Noble Gases
The primary reason most noble gases do not have a standard electronegativity value is their unique electron configuration. All noble gases, except for helium, possess a full octet of eight valence electrons. This complete outer shell configuration provides exceptional chemical stability, meaning these atoms do not have a strong tendency to gain, lose, or share electrons under normal conditions.
This inherent stability results in high ionization energies (the energy required to remove an electron) and electron affinities that are near zero or negative. Because electronegativity describes the pull on electrons in a bond, and these atoms are so reluctant to bond, the property becomes chemically irrelevant for the lighter noble gases. Their inertness is a direct consequence of this stable electronic structure.
Assigning Theoretical Electronegativity Values
Despite their general lack of bonding, chemists have developed theoretical methods to assign numerical electronegativity values to noble gases. One such approach is the Mulliken scale, which defines electronegativity as the average of the first ionization energy and the electron affinity. This approach does not require the atom to be bonded, making it suitable for elements like helium and neon.
Since the lighter noble gases—helium, neon, and argon—have the highest ionization energies of their respective periods, this theoretical calculation yields extremely high electronegativity values. For example, the Mulliken method can place noble gases above fluorine, the most electronegative element on the Pauling scale. These high theoretical values reflect the strong resistance of the noble gas nucleus to giving up an electron, but they do not translate to a strong ability to attract electrons in a bond that an atom would actually form.
Noble Gas Compounds: Xenon and Krypton
The concept of noble gas electronegativity shifts when considering the heavier elements: xenon and krypton. As atomic size increases down the group, the valence electrons are farther from the nucleus, leading to a significant decrease in the ionization energy. This lower ionization energy makes the outer electrons of xenon and krypton easier to engage in chemical bonding, particularly with highly electronegative elements like fluorine and oxygen.
The discovery of the first true noble gas compound, xenon hexafluoroplatinate (\(\text{XePtF}_6\)) in 1962, proved that these elements are not entirely inert. Since then, numerous compounds like xenon difluoride (\(\text{XeF}_2\)) and krypton difluoride (\(\text{KrF}_2\)) have been synthesized. Because these bonds exist, the Pauling method can be applied to estimate the electronegativity of xenon and krypton.
Calculations for bonded xenon place its electronegativity value in a measurable range, often comparable to elements like iodine, which has a Pauling value of 2.66. Specific estimates for xenon range from approximately 2.6 to 4.4 on the Pauling scale, depending on the calculation method and the compound studied. The ability of xenon and krypton to form stable bonds, driven by their lower ionization energy, is the only scenario where the property of electronegativity can be measured meaningfully for a noble gas.