Noble gases have an oxidation number of zero in their natural, uncombined state. This is true for all six noble gases: helium, neon, argon, krypton, xenon, and radon. Because they have completely filled outer electron shells, they don’t need to gain or lose electrons, which is the very definition of chemical stability. In fact, other elements on the periodic table are often assigned oxidation numbers based on how many electrons they need to gain or lose to reach a noble gas configuration.
Why Zero Is the Default
Every free, uncombined element has an oxidation number of zero. But noble gases are special because zero remains their most common oxidation number even in chemistry discussions about compounds. Their outer electron shells are completely full, giving them extraordinarily high ionization energies, meaning it takes a tremendous amount of energy to pull an electron away. Helium has the highest ionization energy of any element on the periodic table, and the lighter noble gases (helium, neon, argon) are so resistant to losing electrons that they form almost no compounds under normal conditions.
Noble gases also lack a standard electronegativity value. Electronegativity measures how strongly an atom attracts electrons in a bond, and because noble gases so rarely form bonds, chemists don’t assign them one. This reinforces the point: under everyday chemistry, noble gases sit at zero and stay there.
When Noble Gases Have Positive Oxidation States
The heavier a noble gas is, the easier it becomes to coax electrons away from it. Larger atoms have more layers of inner electrons that shield the outer electrons from the pull of the nucleus, lowering ionization energy. This is why xenon, the heaviest commonly studied noble gas, has the richest chemistry of the group.
Xenon’s chemistry is dominated by four oxidation states: +2, +4, +6, and +8. These correspond to real, well-characterized compounds:
- +2: Xenon difluoride (XeF₂), a crystalline solid used in semiconductor manufacturing.
- +4: Xenon tetrafluoride (XeF₄), one of the first xenon compounds studied in detail.
- +6: Xenon trioxide (XeO₃), a powerful and dangerously explosive oxidizing agent.
- +8: Perxenate salts and xenon tetroxide (XeO₄), where xenon reaches its maximum oxidation state.
Notice the pattern: xenon’s positive oxidation states are always even numbers. This reflects the way electrons are removed in pairs from its filled outer shell when it bonds with highly electronegative elements like fluorine and oxygen.
The First Noble Gas Compound
Until 1962, chemistry textbooks stated flatly that noble gases could not form compounds. That changed when Neil Bartlett, working at the University of British Columbia, combined xenon with a platinum fluoride to create an orange-yellow solid, xenon hexafluoroplatinate. Later analysis refined the formula to [XeF]⁺[PtF₅]⁻, showing that xenon had an oxidation state of +1 in the compound (though this specific oxidation state turned out to be unusual). The broader significance was immediate: noble gases could bond, and the race to synthesize more compounds began. Within months, other labs had produced XeF₂ and XeF₄.
Krypton, Radon, and the Lighter Gases
Krypton can form compounds, but far fewer than xenon. Krypton difluoride (KrF₂) is the best-known example, giving krypton a +2 oxidation state. The compound is unstable at room temperature and decomposes readily, which illustrates how much harder it is to pull electrons from krypton compared to xenon.
Radon sits below xenon on the periodic table, so in principle it should form compounds even more easily. However, radon is intensely radioactive with very short half-lives, making laboratory work extremely difficult. For practical purposes, its most commonly listed oxidation number is simply 0, though radon fluoride has been detected.
Argon, neon, and helium are far more resistant to compound formation. Argon fluorohydride (HArF) has been created, but only by breaking apart hydrogen fluoride inside a frozen argon matrix at cryogenic temperatures. The compound exists under those extreme conditions, giving argon what can be loosely described as a non-zero oxidation state, but it has no stability at room temperature or pressure.
Helium and neon have no confirmed conventional compounds. A 2017 study did demonstrate that helium and sodium can form a stable structure (Na₂He) at pressures above 113 gigapascals, roughly a million times atmospheric pressure. Even in this exotic material, helium doesn’t truly donate or accept electrons. Instead, electron pairs become trapped in the spaces between atoms, making the material an electride. Helium’s oxidation state in this structure is still effectively zero.
How to Assign Oxidation Numbers in Noble Gas Compounds
If you encounter a noble gas compound on a chemistry assignment, the rules are the same as for any molecule. Fluorine is always assigned -1. Oxygen is typically -2. The oxidation number of the noble gas is whatever value makes the total charge of the molecule add up correctly.
For XeF₂: each fluorine is -1, so the two fluorines contribute -2. For the molecule to be neutral, xenon must be +2.
For XeO₃: each oxygen is -2, contributing -6 total. Xenon must be +6.
For XeF₄: four fluorines at -1 each give -4. Xenon is +4.
This arithmetic approach, rooted in the IUPAC definition of oxidation state (assign each bond’s electrons to the more electronegative atom), works reliably for all noble gas compounds you’re likely to encounter.
Quick Reference by Element
- Helium: 0 only
- Neon: 0 only
- Argon: 0 under normal conditions
- Krypton: 0 and +2
- Xenon: 0, +2, +4, +6, +8
- Radon: 0 (limited data due to radioactivity)