The noble gases, found in Group 18 of the periodic table, hold a unique place in chemistry due to their extreme stability. For centuries, these elements—Helium, Neon, Argon, Krypton, Xenon, and Radon—were known as “inert gases” because they resisted forming chemical bonds. This non-reactivity suggested a zero chemical character, which is quantified by the oxidation number. An oxidation number is a formal concept used to track the hypothetical charge an atom would possess if all its bonds were purely ionic.
Defining Oxidation Numbers
An oxidation number, or oxidation state, is a value assigned to an atom that reflects its electron status within a molecule or ion. It is a tool for monitoring electron movement during chemical reactions, specifically oxidation and reduction. A positive oxidation number suggests a hypothetical loss of electrons, while a negative number implies a hypothetical gain. Tracking these changes is fundamental to understanding how atoms interact and combine.
Several basic rules govern the assignment of these numbers across various chemical species. The oxidation number of any element existing in its pure, uncombined form is always zero, such as the atoms in gaseous oxygen (\(\text{O}_2\)) or metallic iron (\(\text{Fe}\)). For a single atom that has become an ion, the oxidation number is simply the charge of that ion, like \(+1\) for \(\text{K}^+\). In any neutral compound, the sum of the oxidation numbers for all the atoms must equal zero.
Why Noble Gases Are Typically Zero
The primary reason noble gases are typically assigned an oxidation number of zero relates directly to their electron configuration. All noble gases, except Helium, possess a full outer shell of eight valence electrons (\(ns^2np^6\)). Helium is similarly stable with its full \(1s^2\) shell. This complete valence shell configuration makes these atoms chemically saturated and highly reluctant to participate in bonding by losing, gaining, or sharing electrons.
This inherent stability is quantified by two key properties: ionization energy and electron affinity. Noble gases exhibit the highest first ionization energies in their respective periods, meaning significant energy is required to strip away an electron. Conversely, they possess electron affinities close to zero, showing almost no tendency to accept an extra electron. This powerful resistance to changing their electron count is why their default, elemental oxidation number remains zero.
The Synthesis of Noble Gas Compounds
The long-held belief in the complete inertness of the noble gases was overturned by a landmark discovery in the 1960s. This shift occurred because the stability of noble gases decreases as atomic size increases down the group, particularly for heavier elements like Krypton, Xenon, and Radon. The outermost valence electrons in these larger atoms are further from the nucleus and shielded by inner shells, resulting in a lower ionization energy.
The first true chemical compound of a noble gas was created in 1962 by chemist Neil Bartlett. Bartlett observed that platinum hexafluoride (\(\text{PtF}_6\)) reacted with oxygen gas. Realizing that Xenon (Xe) had an ionization energy nearly identical to oxygen, he hypothesized that Xenon could also be oxidized.
Mixing Xenon gas with platinum hexafluoride formed a yellow-orange solid, marking the first successful isolation of a noble gas compound. This initial compound, xenon hexafluoroplatinate, was quickly followed by the synthesis of simple binary compounds like xenon tetrafluoride (\(\text{XeF}_4\)) and xenon difluoride (\(\text{XeF}_2\)). While Xenon is the most prolific compound-former, Krypton and Radon also form compounds. Helium, Neon, and Argon largely remain resistant to stable compound formation.
Determining Non-Zero Oxidation States
The existence of noble gas compounds directly implies that the noble gas atom has adopted a positive, non-zero oxidation state. These compounds are formed almost exclusively with the most electronegative elements, primarily Fluorine and Oxygen. The noble gas atom is thus forced to share its electrons and assume a positive charge. To determine the oxidation number, standard assignment rules are applied using the known, fixed values for the highly electronegative partner atoms.
Fluorine is always assigned an oxidation number of \(-1\) in compounds, and Oxygen is nearly always \(-2\). In the neutral compound Xenon difluoride (\(\text{XeF}_2\)), the two Fluorine atoms contribute a total charge of \(-2\). The Xenon atom must balance this to ensure a net charge of zero, resulting in a \(+2\) oxidation state for Xenon. Similarly, in Xenon tetrafluoride (\(\text{XeF}_4\)), the four Fluorine atoms total \(-4\), making the Xenon oxidation number \(+4\). Xenon can reach even higher oxidation states, such as in Xenon tetroxide (\(\text{XeO}_4\)), where the four Oxygen atoms total \(-8\), forcing the Xenon atom to adopt its maximum known oxidation state of \(+8\).