Group 18 (Noble Gases)
Group 18 elements — helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) — were long called 'noble' or 'inert' gases because chemists believed their completely filled valence shells made chemical bonding impossible. This changed dramatically in 1962 when Neil Bartlett synthesized the first noble gas compound, [XePtF₆], proving that heavier noble gases with lower ionization energies could indeed form bonds under the right conditions. Xenon is the most chemically versatile noble gas. Its relatively large atomic radius and comparatively low ionization energy (1170 kJ/mol) allow it to form stable compounds — predominantly with the highly electronegative fluorine and oxygen atoms. The principal xenon fluorides are XeF₂, XeF₄, and XeF₆, while important oxo-compounds include XeO₃ and XeO₄. These compounds exhibit fascinating geometries predicted by VSEPR theory and undergo characteristic hydrolysis reactions that reveal important aspects of oxidation-state chemistry. Understanding xenon compounds deepens your grasp of VSEPR theory, hybridization beyond the octet rule (expanded octets), and the interplay between electronegativity and reactivity. Xenon compounds are also practically significant: XeF₂ is used in the semiconductor industry as a mild fluorinating agent and for isotropic silicon etching, while xenon compounds serve as model systems for studying chemical bonding in heavier p-block elements.
- XeF₂ — Linear Geometry (sp³d Hybridization)
- XeF₄ — Square Planar Geometry (sp³d² Hybridization)
- XeO₃ — Pyramidal Geometry and Explosive Hazard
- Oxidation States and Reactivity Trends
XeF₂ — Linear Geometry (sp³d Hybridization)
XeF₂ has 8 valence electrons from Xe plus 1 from each F, giving 10 electrons (5 pairs) around Xe. By VSEPR, 5 electron pairs adopt a trigonal bipyramidal arrangement. The 3 lone pairs occupy equatorial positions (minimizing LP–LP repulsion), and the 2 bond pairs occupy axial positions, giving a linear molecular geometry with F–Xe–F bond angle = 180°. Xe is sp³d hybridized. XeF₂ hydrolyzes slowly in water: \ce{2XeF2 + 2H2O -> 2Xe + 4HF + O2}. Xe is reduced from +2 to 0 (and O is oxidized from −2 in water to 0 in O₂), making this a disproportionation-type redox hydrolysis.
XeF₄ — Square Planar Geometry (sp³d² Hybridization)
XeF₄ has 12 electrons around Xe (4 bond pairs + 2 lone pairs = 6 total electron pairs). Six electron pairs form an octahedral electron geometry (sp³d²). The 2 lone pairs occupy opposite axial positions, giving a square planar molecular geometry with F–Xe–F bond angles of 90°. Hydrolysis of XeF₄ is more complex and potentially explosive: \ce{6XeF4 + 12H2O -> 4Xe + 2XeO3 + 24HF + 3O2}. This is a disproportionation: two-thirds of the Xe(+4) is reduced to Xe(0) and one-third is oxidized to Xe(+6) in XeO₃.
XeO₃ — Pyramidal Geometry and Explosive Hazard
XeO₃ (xenon trioxide) is formed by the hydrolysis of XeF₄ or XeF₆. Xe is in the +6 oxidation state. The molecule has 3 bond pairs and 1 lone pair around Xe, giving a tetrahedral electron geometry (sp³) but a trigonal pyramidal molecular shape (like NH₃). The Xe=O bonds have partial double-bond character due to π-backdonation from O lone pairs into Xe d-orbitals, making the bonds shorter (~173 pm) than expected for a single bond. XeO₃ is a powerful oxidizing agent and is dangerously explosive in its dry crystalline form. Hydrolysis of XeF₆ is the cleanest route: \ce{XeF6 + 3H2O -> XeO3 + 6HF}.
Oxidation States and Reactivity Trends
Xenon shows oxidation states of 0, +2, +4, +6, and +8 (in XeO₄). The stability of higher oxidation states is governed by the fluorine/oxygen ligands stabilizing the expanded octet. Higher oxidation state compounds are stronger oxidizing agents: XeO₃ and XeF₆ are extremely reactive, while XeF₂ is relatively stable and mild. All xenon compounds are thermodynamically unstable with respect to the elements (ΔG°formation > 0), meaning they are formed kinetically, not thermodynamically — synthesis typically requires UV photolysis or high temperatures. The high Xe–F bond energy (~133 kJ/mol in XeF₂) provides kinetic stability despite the unfavorable thermodynamics.
The Key Equation
- \ce{XeF4} — Xenon tetrafluoride (mol): Reactant; Xe in +4 oxidation state, square planar geometry
- \ce{XeO3} — Xenon trioxide (mol): Product; Xe in +6 oxidation state, pyramidal geometry, powerful oxidizer
- \ce{HF} — Hydrofluoric acid (mol): Corrosive acid produced as fluorine atoms are released and combine with water's hydrogen
- \ce{O2} — Oxygen gas (mol): Oxygen oxidized from −2 in water to 0 in O₂ (disproportionation byproduct)
Worked Example
Problem. Determine the oxidation state of xenon in XeF₂, XeF₄, and XeO₃. Then, using VSEPR theory, predict the molecular geometry of XeF₂ and justify your answer by counting electron pairs.
Answer. Oxidation states: Xe is +2 in XeF₂, +4 in XeF₄, +6 in XeO₃. XeF₂ has a linear molecular geometry (180° F–Xe–F) with 3 lone pairs in equatorial positions of a trigonal bipyramidal electron arrangement.
Origins · 1 minNeil Bartlett and how this was discovered
Neil Bartlett shattered the dogma of noble gas inertness in 1962 by synthesizing xenon hexafluoroplatinate, [Xe⁺][PtF₆⁻], the first compound of a noble gas. His key insight was noticing that the first ionization energy of xenon (1170 kJ/mol) is almost identical to that of dioxygen O₂ (1165 kJ/mol), and since he had already made O₂⁺PtF₆⁻, he reasoned Xe should do the same. Within a year, other chemists had synthesized XeF₂, XeF₄, and XeF₆ by direct combination, igniting a new field of noble gas chemistry. Bartlett made his historic discovery at the University of British Columbia using a simple glass vacuum line. When he mixed colorless XeF gas with PtF₅, an immediate orange-yellow solid formed — the reaction was so fast it was visible to the naked eye in real time.
Related Concepts
- VSEPR Theory and Molecular Geometry — VSEPR theory is the primary tool for predicting the linear, square planar, and pyramidal shapes of XeF₂, XeF₄, and XeO₃ respectively.
- Hybridization and Expanded Octets — Xenon compounds involve sp³d and sp³d² hybridization, requiring an understanding of how d-orbitals participate in bonding beyond the standard octet rule.
- Oxidation States and Redox Reactions — Hydrolysis of XeF₄ is a disproportionation reaction where Xe undergoes both oxidation (+4→+6) and reduction (+4→0), making redox concepts essential.
- Group 17: The Halogens — Fluorine's extreme electronegativity is the key reason it can force xenon into bonding; comparing halogen properties helps explain why only F and Cl form stable noble gas compounds.