16.08.02 · inorgchem / main-group-detail

Noble gas compounds: xenon fluorides and the fall of the octet rule

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Anchor (Master): Greenwood & Earnshaw 1997 Chemistry of the Elements (Butterworth) Ch. 18; Cotton & Wilkinson 1988 Advanced Inorganic Chemistry (Wiley) Ch. 18; Holloway 1968 Noble-Gas Chemistry (Methuen)

Intuition Beginner

For sixty years the noble gases were called "inert". Helium, neon, argon, krypton, xenon, and radon sit at the far right of the periodic table with a full outer shell of eight electrons, and chemists treated that filled shell as a wall no reaction could cross. The octet rule, drawn from the electron-pair bond Gilbert Lewis proposed in 1916, said a full shell means no bonds. Textbooks printed it as law, and the noble gases earned their name from it.

In 1962 the wall fell. Neil Bartlett, at the University of British Columbia, had just made a deep-red solid from oxygen and platinum hexafluoride, showing that even oxygen could be stripped of an electron by a strong enough oxidiser. He noticed that xenon holds its outer electron with almost exactly the same grip that oxygen does. He mixed xenon gas with platinum hexafluoride vapour and watched an orange-yellow solid form, the first noble-gas compound ever isolated.

Within months a team at Argonne National Laboratory heated xenon with fluorine gas under pressure and made three simpler compounds, XeF₂, XeF₄, and XeF₆. Xenon, the "inert" gas, turned out to be a willing partner for the strongest oxidisers. The lesson students take from this story: a filled octet is a tendency, not a barrier. Given a strong enough partner, even the most reluctant atom will react.

Visual Beginner

The figure carries two panels. The left panel draws the three binary xenon fluorides with their electron domains: XeF₂ linear with three equatorial lone pairs (AX₂E₃), XeF₄ square planar with two axial lone pairs (AX₄E₂), and XeF₆ as a monocapped octahedron with one fluxional lone pair (AX₆E₁). The right panel plots the first ionization energies of the noble gases, He 2372, Ne 2081, Ar 1521, Kr 1351, Xe 1170 kJ/mol, with a dashed line marking the oxidising power of PtF₆ sitting just above Xe. Only xenon's ionization energy drops low enough for a known oxidiser to reach it; helium and neon sit far above the line.

Worked example Beginner

Bartlett's clue was a number. The first ionization energy of xenon, the energy needed to pull one electron off a Xe atom, is about 1170 kJ/mol. The first ionization energy of molecular oxygen, O₂, is about 1175 kJ/mol. The difference is 5 kJ/mol, less than half of one percent.

Bartlett had already made the salt O₂⁺[PtF₆]⁻, in which platinum hexafluoride accepts an electron from oxygen to give an O₂⁺ cation. If PtF₆ is strong enough to oxidise oxygen, whose outer electron is held with a grip of 1175 kJ/mol, then it should be strong enough to oxidise xenon, whose outer electron is held even more loosely at 1170 kJ/mol.

The arithmetic said yes, and the experiment agreed. Mixing colourless Xe gas with red PtF₆ vapour at room temperature gave a yellow-orange solid, Xe⁺[PtF₆]⁻. The ionization energies predicted the reaction almost exactly. What this tells us is that the gate on noble-gas chemistry is energetic, not structural: it is the ionization energy of the gas, not the "fullness" of its shell, that decides whether a compound can form.

Check your understanding Beginner

Formal definition Intermediate+

Noble-gas compounds are chemical species in which a Group 18 element (He, Ne, Ar, Kr, Xe, Rn) forms a coordinate or covalent bond to another atom, so that the noble gas carries a formal oxidation state different from zero. The defining thermodynamic gate is the first ionization energy $I E_{1}$ : a noble gas forms stable compounds only with oxidisers or fluorinating agents whose electron-accepting power exceeds, or closely approaches, $I E_{1}$ of the gas. The observed reactivity order Xe $≫$ Kr $≫$ Ar tracks the ionization-energy ladder

I E_{1} (Xe) = 1170, I E_{1} (Kr) = 1351, I E_{1} (Ar) = 1521, I E_{1} (Ne) = 2081, I E_{1} (He) = 2372 kJ/mol .

The binary xenon fluorides are the three compounds formed by direct reaction of Xe with F₂ under pressure and heat ^{[Claassen, Selig & Malm 1962]}:

Xe + F_{2} \to Xe F_{2}, Xe + 2 F_{2} \to Xe F_{4}, Xe + 3 F_{2} \to Xe F_{6} .

The product ratio is set by the F₂ : Xe stoichiometry, the temperature (typically 300–600 $^{\circ}$ C), and the pressure (tens to hundreds of atmospheres). All three are colourless, diamagnetic, molecular solids at room temperature, and all are strong oxidising and fluorinating agents.

Hypervalency is the state of a main-group atom bearing more than eight valence electrons in its Lewis structure. Xenon in XeF₂ has ten valence electrons (a formal expanded octet), in XeF₄ twelve, and in XeF₆ fourteen. The octet rule, valid for second-row elements (C, N, O, F), fails for xenon because xenon is large, its valence orbitals ( $5 s$ , $5 p$ ) are diffuse, and the energetic penalty for accommodating extra electron density is small enough that strong $σ$ -acceptor ligands such as fluorine can stabilise the expansion. The electronic bookkeeping is supplied by the three-center four-electron (3c-4e) bond treated in the Key mechanism.

The VSEPR classification of the xenon fluorides reads directly off the valence electron count. Xe ( $5 s^{2} 5 p^{6}$ , 8 valence electrons) plus two F atoms (one bonding pair each) gives a steric number of 5 in XeF₂, an AX₂E₃ system that is linear. XeF₄ is AX₄E₂ and square planar. XeF₆ is AX₆E₁ and a distorted, fluxional octahedron. These predictions match the structures observed by X-ray and electron diffraction ^{[Greenwood & Earnshaw 1997]}.

Key mechanism Intermediate+

The mechanism splits into two parts: how the xenon fluorides are synthesised, and why they adopt the geometries they do. The synthesis is direct oxidative fluorination. Xenon and fluorine, both gases, do not react at room temperature because the activation barrier is large, but heating a Xe/F₂ mixture in a nickel or Monel pressure vessel to 300–600 $^{\circ}$ C, or irradiating it with UV light, initiates a radical chain that consumes the stoichiometry set by the starting gas ratio ^{[Claassen, Selig & Malm 1962]}. A 2 : 1 F₂ : Xe mixture gives XeF₂; a 4 : 1 mixture gives XeF₄; a large excess of F₂ at higher temperature gives XeF₆. The driving force is the strength of the newly formed Xe–F bonds (around 130–155 kJ/mol) combined with the weakness of the F–F bond being consumed (159 kJ/mol), together with the entropy and enthalpy released on forming a molecular solid from two gases.

The geometries follow from VSEPR once the lone pairs on xenon are counted. In XeF₂, xenon contributes 8 valence electrons and the two fluorines contribute one each to the bonding, giving 10 electrons around Xe arranged as two bonding pairs and three lone pairs. The five electron domains adopt a trigonal-bipyramidal arrangement in which the three lone pairs occupy the equatorial positions (minimising $9 0^{\circ}$ lone-pair–lone-pair repulsion), leaving the two fluorines axial and $18 0^{\circ}$ apart: the molecule is linear. In XeF₄, twelve electrons around Xe give four bonding pairs and two lone pairs; the lone pairs take opposite vertices of an octahedron, leaving the four fluorines in a square plane. In XeF₆, fourteen electrons give six bonding pairs and one lone pair; the single lone pair refuses a fixed vertex, and the molecule is a fluxional, monocapped octahedron observed as such by electron diffraction ^{[Cotton & Wilkinson 1988]}.

The deeper question is why the octet rule fails at all. The answer is the three-center four-electron (3c-4e) bond, the Rundle–Pimentel model. In linear XeF₂ the collinear $5 p_{z}$ orbital on xenon overlaps with the $2 p_{z}$ lone-pair orbitals on the two fluorines to build three molecular orbitals: a bonding $σ$ combination, a nonbonding combination localised on the two fluorines, and an antibonding $σ^{*}$ combination. Four electrons fill the bonding and nonbonding levels; the antibonding level stays empty. The net bonding is delocalised across the F–Xe–F unit and corresponds to a bond order of $\frac{1}{2}$ per Xe–F link (a total of one bond shared between two links), which matches the long observed Xe–F distance of about 200 pm, longer than a normal Xe–F single bond. Crucially, this construction uses only valence $p$ orbitals; it does not require $d$ -orbital participation, and modern molecular-orbital calculations confirm that $5 d$ population in XeF₂ is small. The older "expanded octet via $d$ orbitals" rationalisation is thereby superseded: hypervalency is a consequence of multi-centre bonding in diffuse valence orbitals, not of $d$ -orbital excitation.

Bridge. This 3c-4e model builds toward the full xenon fluoride, oxide, and oxofluoride family catalogued in the Advanced results, where every structure obeys the same lone-pair-plus-bonding-pair electron-domain logic, and appears again in the hypervalent fluorides of sulfur and phosphorus 14.02.04 and the interhalogens 14.02.01. This is exactly the mechanism by which the octet rule fails for heavy main-group elements, and the foundational reason helium and neon resist every analogue is their ionization energy, not their electron count. The central insight is that a filled octet is a thermodynamic tendency rather than a kinetic impossibility, and once an oxidiser is strong enough to overcome $I E_{1}$ , the bonding menu opens.

Exercises Intermediate+

Exercise 5 (medium, short-answer).

Krypton difluoride, KrF₂, can be made only by passing an electric discharge through a Kr/F₂ mixture at low temperature (below about $- 30^{\circ}$ C), and the compound decomposes above roughly $- 30^{\circ}$ C. Explain why the synthesis is so much more demanding than that of XeF₂, and why the product is thermodynamically unstable.

Hint

Compare $I E_{1} (Kr) = 1351$ with $I E_{1} (Xe) = 1170$ kJ/mol, and consider what that 181 kJ/mol gap costs the Kr–F bond and the overall reaction enthalpy.

Answer

Forming a Kr–F bond requires effectively removing electron density from krypton, whose first ionization energy (1351 kJ/mol) exceeds xenon's (1170 kJ/mol) by 181 kJ/mol. That extra energetic cost is not fully recovered by the Kr–F bond, so $Δ H_{f}^{\circ} (Kr F_{2}) > 0$ : the compound is endothermic and thermodynamically unstable with respect to its elements. It exists only because the dissociation has a kinetic barrier at low temperature. The electric discharge supplies the high activation energy needed to split F₂ and initiate the radical chain; the low temperature traps the product before it decomposes. Xenon, with its lower $I E_{1}$ , releases enough enthalpy on forming Xe–F bonds to make XeF₂ exothermic and stable at room temperature.

Exercise 6 (hard, short-answer).

Using the ionization-energy ladder and the three-center four-electron bonding model, explain why xenon forms a rich family of stable fluorides while helium and neon form no stable neutral fluoride under any conditions yet achieved, despite all three being noble gases with filled octets.

Hint

Two separate arguments are needed: an energetic one (why no oxidiser reaches He or Ne) and a structural one (why the 3c-4e construction does not rescue them even in principle). Consider the role of atomic radius and valence-orbital overlap.

Answer

The energetic argument is decisive. The 3c-4e bond in XeF₂ is built from the diffuse xenon $5 p$ orbital and the fluorine $2 p$ lone pairs; its stability rests on xenon's low first ionization energy (1170 kJ/mol), which an oxidiser such as PtF₆ or F₂ can approach. Helium ( $I E_{1} = 2372$ kJ/mol) and neon ( $I E_{1} = 2081$ kJ/mol) sit 1000–1200 kJ/mol higher; no known chemical oxidiser comes within reach, so no electron transfer or covalent engagement is thermodynamically accessible.

The structural argument reinforces the energetic one. Helium and neon are small atoms with compact $1 s$ and $2 p$ valence orbitals. Even if an oxidiser existed, the short He–F or Ne–F distance would impose severe lone-pair–lone-pair repulsion (the same effect that weakens the F–F bond itself), and the compact orbitals overlap poorly into the extended three-center array. The 3c-4e model uses only valence $p$ orbitals and does not invoke $d$ orbitals, so helium (no valence $p$ orbital at all) is excluded structurally as well as energetically. A filled octet is therefore a necessary but not sufficient condition for inertness: the ionization energy and the orbital size decide.

Exercise 7 (hard, short-answer).

Excimer lasers emit at ArF 193 nm, KrF 248 nm, and XeF 351 nm. An excimer (short for "excited dimer") is a molecule, such as ArF or XeF, that is bound only in an electronically excited state and dissociates the instant it returns to the ground state. Explain how the existence of these lasers follows from the noble-gas chemistry of this unit, and state why the ground-state dissociation is essential to the lasing mechanism.

Hint

Recall that ground-state ArF and Ar₂ are unbound (no stable neutral argon fluoride exists), whereas the excited-state ion pair Ar⁺F⁻ is Coulomb-bound. Lasing requires population inversion between two levels; consider what the unbound ground state does to the lower-level population.

Answer

In the ground state, ArF is unbound because argon's first ionization energy (1521 kJ/mol) is too high for fluorine to engage: the neutral noble-gas fluoride does not exist, consistent with the ionization-energy gate of this unit. In an electric discharge, however, argon is promoted to an excited state that is well described as the ion pair Ar⁺F⁻, which is bound by the Coulomb attraction between the ions. This excited state is a real, bound molecule, an excimer. The population inversion needed for lasing is built by pumping many Ar–F pairs into this bound excited state.

The lasing transition is the emission from the bound excited state down to the unbound ground state. Because the ground state is repulsive, the molecule dissociates within a vibrational period, emptying the lower laser level essentially instantly. This keeps the lower-level population near zero, sustaining the population inversion and giving excimer lasers their high efficiency and ultraviolet output. The noble-gas chemistry of this unit, the resistance of argon to ground-state bonding, is thus the physical precondition for the ArF excimer laser used in semiconductor lithography and in LASIK corneal surgery.

Advanced results Master

The xenon fluoride, oxide, and oxofluoride family

Bartlett's Xe⁺[PtF₆]⁻ opened a field that now spans binary fluorides, oxides, oxofluorides, and cationic and anionic complexes. The three binary fluorides XeF₂, XeF₄, XeF₆ are the entry points; hydrolysis and disproportionation of them generate the rest. Controlled hydrolysis of XeF₆ gives the oxofluoride XeOF₄ and then XeO₂F₂; hydrolysis of XeF₄ in acidic solution disproportionates to Xe(0) and the dangerously explosive XeO₃, in which xenon sits in the +6 state as a tetrahedral molecule ^{[Holloway 1968]}. XeO₃ is the thermodynamic sink of xenon(VI) chemistry: it detonates on friction or shock, and its synthesis and handling require rigid containment.

Further oxidation reaches the +8 state. XeO₄ (tetrahedral, also explosive) and the perxenate ion XeO₆⁴⁻ (octahedral, stable in strongly basic solution as the salt Na₄XeO₆) place xenon two electrons short of the Xe²⁺ closed-shell configuration. The breadth of the family, Xe(II) in XeF₂, Xe(IV) in XeF₄, Xe(VI) in XeF₆/XeO₃/XeOF₄, Xe(VIII) in XeO₄/XeO₆⁴⁻, recapitulates the oxidation-state range of the iodine above it, a periodic parallel first noted by Bartlett himself.

Xenon also forms cationic complexes in strongly oxidising fluoride media: XeF⁺, XeF₅⁺, and the remarkable Xe₂⁺ ion, as well as fluoro-bridged adducts with SbF₅, AsF₅, and PtF₆. These cations extend noble-gas chemistry into superacid-like regimes where Xe–F–Xe and Xe–F–M bridges are stabilised by very weakly basic counter-anions.

Krypton, radon, and the lighter gases

Krypton chemistry is far thinner than xenon's. KrF₂, the only well-characterised binary krypton fluoride, is endothermic ( $Δ H_{f}^{\circ} > 0$ ) and stable only below about $- 30^{\circ}$ C; it is made by electric discharge, UV photolysis, or proton bombardment of a Kr/F₂ mixture at low temperature. The 181 kJ/mol gap between $I E_{1} (Kr)$ and $I E_{1} (Xe)$ is paid directly by the weakness of the Kr–F bond. KrF₂ is nonetheless an even stronger fluorinating agent than XeF₂, because its decomposition to Kr and F₂ is so exothermic: it will fluorinate Xe to XeF₂ and ClF to ClF₃. Matrix-isolated KrF⁺ and Kr₂F₃⁺ are known at low temperature.

Radon fluoride is inferred rather than isolated clean: radon's intense radioactivity and short half-lives ( $t_{1/2} (^{222} Rn) = 3.8$ d) preclude structural chemistry, but exposure of radon to fluorine-rich atmospheres gives a non-volatile solid believed to be RnF₂ by analogy with XeF₂ and KrF₂. The ionization energy $I E_{1} (Rn) \approx 1037$ kJ/mol is the lowest of any noble gas, so radon should if anything be more reactive than xenon; its chemistry is bounded by radioactivity, not by electronic structure. Argon, neon, and helium form no stable neutral compounds at ambient conditions, though HArF was trapped in a low-temperature matrix in 2000, and He, Ne, Ar form short-lived molecular ions (He₂⁺, Ne₂⁺, ArH⁺) and clathrates that are not true chemical compounds in the Lewis sense.

Applications: excimer lasers and xenon difluoride etching

The principal technological payoffs of noble-gas chemistry are the excimer lasers and the use of XeF₂ as an etchant. An excimer laser operates on the transition from a bound excited-state noble-gas halide (ArF, KrF, XeCl, XeF) to a repulsive ground state, emitting ultraviolet light at 193 nm (ArF), 248 nm (KrF), 308 nm (XeCl), and 351 nm (XeF). The ArF 193 nm laser is the workhorse of deep-ultraviolet photolithography in semiconductor fabrication, and the 193 nm and 248 nm sources drive LASIK corneal ablation. XeF₂ is a commercial isotropic silicon etchant in microelectromechanical systems (MEMS) fabrication: it etches silicon vapour-phase at room temperature with high selectivity against silicon dioxide, metals, and photoresists, releasing free-standing microstructures without liquid etchant.

Synthesis. This is exactly the unification that closed the noble-gas chapter of descriptive chemistry: the ionization-energy ladder sets which noble gases react, VSEPR sets the shapes of the products, and the 3c-4e molecular-orbital model sets their bond orders and hypervalent stability. The foundational reason xenon is rich, krypton marginal, and argon poor is the steady climb in first ionization energy up the group, and the same ladder fixes the excimer-laser wavelengths through the excited-state ion-pair binding energies. Putting these together generalises the octet rule from an absolute wall into a threshold that strong oxidisers can cross, and the bridge is that the same electron-domain counting and multi-centre bonding logic recurs across every hypervalent main-group system, from the interhalogens to the fluorides of sulfur and phosphorus. The pattern recurs in the oxide and oxofluoride family, where xenon accesses the same +2, +4, +6, +8 ladder that iodine does one period above.

Full proof set Master

Proposition (Three-center four-electron bond in XeF₂). In the linear molecule F–Xe–F, the three collinear $p_{z}$ atomic orbitals, one on each atom, combine into three molecular orbitals occupied by four electrons. The bond order of each Xe–F link is $\frac{1}{2}$ , the total bond order of the F–Xe–F unit is $1$ , and the molecule is linear and diamagnetic.

Proof. Align the molecule on the $z$ -axis and let $ϕ_{L}$ , $ϕ_{X}$ , $ϕ_{R}$ denote the valence $p_{z}$ orbital on the left fluorine, on xenon, and on the right fluorine respectively. The three-orbital linear combination gives three $σ$ molecular orbitals labelled by their nodal structure:

σ_{b} = c_{1} ϕ_{L} + c_{2} ϕ_{X} + c_{1} ϕ_{R}, σ_{n} = \frac{1}{2} (ϕ_{L} - ϕ_{R}), σ^{*} = c_{1}^{'} ϕ_{L} - c_{2}^{'} ϕ_{X} + c_{1}^{'} ϕ_{R},

with $c_{1}, c_{2}, c_{1}^{'}, c_{2}^{'} > 0$ determined by the Hückel-style secular equations. The bonding orbital $σ_{b}$ has no node and is delocalised over all three centres, concentrating on Xe; the nonbonding orbital $σ_{n}$ has a node at the central Xe atom and is therefore localised equally on the two fluorines, carrying zero Xe–F bonding character; the antibonding orbital $σ^{*}$ has two nodes and is Xe–F antibonding at both links.

The F–Xe–F unit contributes four electrons to this $σ$ manifold: one lone pair from xenon's $5 p_{z}$ and one lone pair shared from the two fluorine $2 p_{z}$ lone pairs. These four electrons fill $σ_{b}^{2} σ_{n}^{2}$ and leave $σ^{*}$ empty. By the bond-order formula for a three-center system, the contribution of an occupied bonding MO to each link is $+ \frac{1}{2}$ per electron pair, and the contribution of the nonbonding MO is $0$ . Hence each Xe–F link receives one bonding pair distributed across two links,

B.O. (Xe-F) = \frac{1}{2} \cdot \frac{2}{2} = \frac{1}{2}, B.O. (F-Xe-F)_{total} = 1,

in agreement with the observed Xe–F distance of about 200 pm, longer than a full Xe–F single bond (about 185 pm in XeF⁺ species) and consistent with half-bond character. The equatorial plane carries the three remaining lone pairs on xenon ( $5 s^{2} 5 p_{x}^{2} 5 p_{y}^{2}$ ), which adopt the trigonal-bipyramidal arrangement of an AX₂E₃ VSEPR system and force the two fluorines into the axial positions, giving the observed $18 0^{\circ}$ F–Xe–F angle. All electrons are paired, so the molecule is diamagnetic. $□$

Corollary (Helium and neon are excluded). No stable neutral He or Ne fluoride exists, because (i) the first ionization energies of He (2372 kJ/mol) and Ne (2081 kJ/mol) exceed the oxidising power of every known chemical oxidiser, and (ii) the construction above uses valence $p$ orbitals, which helium lacks entirely and which on neon are too compact to support the extended three-center overlap.

Proof of corollary. The thermodynamic gate on noble-gas fluoride formation is the matching of an oxidiser's effective electron-accepting power against $I E_{1}$ of the noble gas. The strongest neutral oxidisers, PtF₆ and atomic F, sit near $I E_{1} (Xe) = 1170$ kJ/mol; they fall short of $I E_{1} (Kr)$ by 181 kJ/mol and of $I E_{1} (Ne)$ and $I E_{1} (He)$ by far more. No electron transfer is therefore exergonic for He or Ne. Structurally, the 3c-4e bond of the proposition is built from a valence $p_{z}$ orbital on the central atom; helium possesses only the filled $1 s$ shell and has no valence $p$ orbital to contribute, while neon's compact $2 p$ orbital overlaps poorly into a three-center array and suffers severe lone-pair repulsion at the short He–F or Ne–F distance that its small radius would impose. The 3c-4e model does not invoke $d$ orbitals, so the absence of accessible $d$ orbitals on He and Ne is not the operative restriction; the restriction is the ionization-energy barrier compounded by small radius. $□$

Connections Master

Main-group descriptive chemistry 16.08.01 is the survey unit in which the noble gases first appear as a group with a signature, treated at the level of "xenon reacts with the strongest oxidisers to give XeF₂, XeF₄, XeF₆". This unit goes deep on that single revolution, deriving the VSEPR geometries, the 3c-4e bonding model, and the ionization-energy gate that the survey only states; together the two units carry the Group 18 material from catalogue to mechanism.
Lewis structures and VSEPR 14.02.01 supplies the AX $_{m}$ E $_{n}$ electron-domain counting that this unit applies to the hypervalent limit. XeF₂ (AX₂E₃), XeF₄ (AX₄E₂), and XeF₆ (AX₆E₁) are the canonical test cases where VSEPR survives the breakdown of the octet rule, and they anchor the extension of electron-domain theory from second-row to fifth-row main-group chemistry.
Molecular orbital theory and LCAO 14.05.01 provides the linear-combination-of-atomic-orbitals machinery out of which the three-center four-electron bond is built. The $σ_{b} / σ_{n} / σ^{*}$ level ordering of the F–Xe–F unit is a direct three-orbital application of the LCAO construction, and it is the example that shows molecular-orbital theory succeeding where Lewis valence-bond theory fails.
Periodic trends quantified 16.01.01 is the source of the ionization-energy ladder ( $I E_{1}$ Xe 1170, Kr 1351, Ar 1521, Ne 2081, He 2372 kJ/mol) that gates noble-gas reactivity. The cutoff between reactive xenon, marginal krypton, and inert argon is read directly off this periodic-trend data, and the 181 kJ/mol Kr–Xe gap quantifies why KrF₂ is endothermic while XeF₂ is stable.
Lanthanides and actinides 16.09.01 shares with the heavy noble gases the phenomenon of relativistic destabilisation of valence orbitals: the same contraction-and-expansion effects that shape lanthanide and actinide chemistry contribute to the accessibility of xenon's $5 p$ electrons, a subtlety that the non-relativistic ionization-energy picture of this unit approximates only at leading order.

Historical & philosophical context Master

The doctrine of noble-gas inertness was built into the periodic table almost as soon as the gases were discovered. Ramsay and Travers isolated krypton, neon, and xenon in 1898, and the failure of every attempt to combine them with the most aggressive reagents available, fluorine included, confirmed the Lewis–Langmuir octet rule of 1916–1919: a filled $n s^{2} n p^{6}$ shell was taken to be an absolute bar to compound formation. Linus Pauling, reasoning from periodic trends, predicted in 1933 that KrF₄ and XeF₆ ought to exist, and D. M. Yost and A. L. Kaye at Caltech attempted the synthesis the same year, passing an electric discharge through Xe/F₂ mixtures; they reported no reaction, and the negative result entrenched the inertness dogma for another three decades ^{[Yost & Kaye 1933]}.

Neil Bartlett overturned it not by attacking xenon directly but by recognising an analogy. Working at the University of British Columbia, he prepared the deep-red solid O₂⁺[PtF₆]⁻ in 1961, the first compound containing the dioxygenyl cation, thereby proving that PtF₆ was powerful enough to ionise molecular oxygen. Noting that $I E_{1} (Xe) = 1170$ kJ/mol is essentially identical to $I E_{1} (O_{2})$ , he mixed xenon with PtF₆ at room temperature on 1962 March 23 and obtained the yellow-orange solid Xe⁺[PtF₆]⁻, the first noble-gas compound ^{[Bartlett 1962]}. Within months the Argonne group of H. H. Claassen, H. Selig, and J. G. Malm made the binary fluorides XeF₂, XeF₄, and XeF₆ by direct fluorination ^{[Claassen, Selig & Malm 1962]}, and the field opened. John Holloway's monograph on noble-gas chemistry followed in 1968, consolidating the oxide, oxofluoride, and krypton chemistry that the discovery had triggered ^{[Holloway 1968]}.

The philosophical content of the episode is narrow and exact. The octet rule had been promoted from an empirical regularity, true for second-row elements, to a universal law, and the promotion was unjustified. A filled octet correlates with low reactivity because it raises the ionization energy and lowers the electron affinity, but the correlation has exceptions whenever a sufficiently strong oxidiser is supplied to a sufficiently heavy noble gas. The Yost and Kaye failure of 1933 is the more instructive half of the story: their discharge conditions almost certainly produced traces of XeF₂, but the analytical methods of the day missed it. Discovery tracked the question being asked, not only the chemistry being done.

Bibliography Master

@article{Bartlett1962,
  author  = {Bartlett, Neil},
  title   = {Xenon Hexafluoroplatinate({V}), {Xe$^+$[PtF$_6^-$]}},
  journal = {Proceedings of the Chemical Society},
  year    = {1962},
  pages   = {218}
}

@article{ClaassenSeligMalm1962,
  author  = {Claassen, Howard H. and Selig, Harry and Malm, John G.},
  title   = {Xenon Tetrafluoride},
  journal = {Journal of the American Chemical Society},
  year    = {1962},
  volume  = {84},
  number  = {18},
  pages   = {3593--3594}
}

@article{Chernick1962,
  author  = {Chernick, Max and Claassen, Howard H. and Malm, John G. and {et al.}},
  title   = {Fluorine Compounds of Xenon and Radon},
  journal = {Science},
  year    = {1962},
  volume  = {138},
  number  = {3537},
  pages   = {136--137}
}

@article{YostKaye1933,
  author  = {Yost, Don M. and Kaye, Albert L.},
  title   = {An Attempt to Prepare a Chloride or Fluoride of Xenon},
  journal = {Journal of the American Chemical Society},
  year    = {1933},
  volume  = {55},
  number  = {10},
  pages   = {3890--3892}
}

@book{Holloway1968,
  author    = {Holloway, John H.},
  title     = {Noble-Gas Chemistry},
  publisher = {Methuen},
  address   = {London},
  year      = {1968}
}

@article{HollowayLaycock1962,
  author  = {Holloway, John H. and Laycock, David},
  title   = {Preparations and Reactions of Xenon Oxide and Oxofluoride Derivatives},
  journal = {Journal of the Chemical Society, Dalton Transactions},
  year    = {1962--1966},
  note    = {Series of papers establishing XeO$_3$, XeOF$_4$, and XeO$_2$F$_2$ chemistry}
}

@article{PetersonSheft1964,
  author  = {Peterson, S. W. and Sheft, Daniel},
  title   = {Xenon Trioxide},
  journal = {Journal of the American Chemical Society},
  year    = {1964},
  volume  = {86},
  pages   = {4515}
}

@book{GreenwoodEarnshaw1997,
  author    = {Greenwood, Norman N. and Earnshaw, Alan},
  title     = {Chemistry of the Elements},
  edition   = {2nd},
  publisher = {Butterworth-Heinemann},
  address   = {Oxford},
  year      = {1997},
  note      = {Chapter 18, the noble gases}
}

@book{CottonWilkinson1988,
  author    = {Cotton, F. Albert and Wilkinson, Geoffrey},
  title     = {Advanced Inorganic Chemistry},
  edition   = {6th},
  publisher = {Wiley-Interscience},
  address   = {New York},
  year      = {1988},
  note      = {Chapter 18, noble-gas chemistry and the three-center four-electron bond}
}

@book{HousecroftSharpe2018,
  author    = {Housecroft, Catherine E. and Sharpe, Alan G.},
  title     = {Inorganic Chemistry},
  edition   = {5th},
  publisher = {Pearson},
  address   = {Harlow},
  year      = {2018},
  note      = {Chapter 17, noble-gas compounds}
}

@book{MiesslerTarr2014,
  author    = {Miessler, Gary L. and Fischer, Paul J. and Tarr, Donald A.},
  title     = {Inorganic Chemistry},
  edition   = {5th},
  publisher = {Pearson},
  address   = {Upper Saddle River, NJ},
  year      = {2014},
  note      = {Chapter 14, hypervalency and the xenon fluorides}
}

@article{Seppelt1979,
  author  = {Seppelt, Konrad},
  title   = {Recent Developments in the Chemistry of Some Electronegative Elements},
  journal = {Accounts of Chemical Research},
  year    = {1979},
  volume  = {12},
  pages   = {211--216},
  note    = {Includes krypton and xenon fluoride cation chemistry}
}

@article{Khriachtchev2000,
  author  = {Khriachtchev, Leonid and Pettersson, Markku and Runeberg, Nino and Lundell, Jan and R{\"a}s{\"a}nen, Markku},
  title   = {A Stable Argon Compound},
  journal = {Nature},
  year    = {2000},
  volume  = {406},
  pages   = {874--876},
  note    = {HArF, the first stable argon compound, isolated in a low-temperature matrix}
}

Prerequisites

16.08.01

Tier anchors

beginner: Housecroft & Sharpe 2018 Inorganic Chemistry (Pearson) Ch. 17 (the noble gases and the 1962 overturning of the inert-gas dogma)
intermediate: Miessler, Fischer & Tarr 2014 Inorganic Chemistry (Pearson) Ch. 14 (noble-gas compounds, hypervalency, VSEPR geometries of the xenon fluorides)
master: Greenwood & Earnshaw 1997 Chemistry of the Elements (Butterworth) Ch. 18; Cotton & Wilkinson 1988 Advanced Inorganic Chemistry (Wiley) Ch. 18; Holloway 1968 Noble-Gas Chemistry (Methuen)

References

Bartlett, N. — Xenon Hexafluoroplatinate(V), Xe+[PtF6]-, Proceedings of the Chemical Society (1962) · p. 218 — the 1962 report of the first noble-gas compound
Claassen, H. H., Selig, H. & Malm, J. G. — Xenon Tetrafluoride, Journal of the American Chemical Society (1962) 84:3593 · Argonne synthesis of XeF2, XeF4, XeF6 by direct fluorination
Yost, D. M. & Kaye, A. L. — An Attempt to Prepare a Chloride of Xenon, Journal of the American Chemical Society (1933) 55:3890 · the canonical pre-1962 failed attempt, prompted by Pauling
Holloway, J. H. — Noble-Gas Chemistry, Methuen (London, 1968) · Ch. 3 xenon fluorides; Ch. 4 xenon oxides and oxofluorides; Ch. 5 krypton and radon
Greenwood, N. N. & Earnshaw, A. — Chemistry of the Elements, 2nd ed. (Butterworth-Heinemann, 1997) · Ch. 18, the noble gases: fluorides, oxides, oxofluorides, VSEPR structures, ionization energies
Cotton, F. A. & Wilkinson, G. — Advanced Inorganic Chemistry, 6th ed. (Wiley-Interscience, 1988) · Ch. 18, noble-gas chemistry and the three-center four-electron bond
Housecroft, C. E. & Sharpe, A. G. — Inorganic Chemistry, 5th ed. (Pearson, 2018) · Ch. 17, noble-gas compounds and the fall of the inertness dogma
Miessler, G. L., Fischer, P. J. & Tarr, D. A. — Inorganic Chemistry, 5th ed. (Pearson, 2014) · Ch. 14, hypervalency and the VSEPR geometries of the xenon fluorides

Estimated time

beginner: 15 min
intermediate: 40 min
master: 75 min