Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF
2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with light or water vapor but is otherwise stable to storage. Xenon difluoride is a dense, white crystalline solid.
3D model (JSmol)
|Molar mass||169.29 g·mol−1|
|Density||4.32 g/cm3, solid|
|Melting point||128.6 °C (263.5 °F; 401.8 K)|
|25 g/l (0 °C)|
|Vapor pressure||6.0×102 Pa|
|parallel linear XeF2 units|
Std enthalpy of
|Main hazards||Corrosive to exposed tissues. Releases toxic compounds on contact with moisture.|
|Safety data sheet||PELCHEM MSDS|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Xenon difluoride is a linear molecule with an Xe–F bond length of 197.73±0.15 pm in the vapor stage, and 200 pm in the solid phase. The packing arrangement in solid XeF
2 shows that the fluorine atoms of neighbouring molecules avoid the equatorial region of each XeF
2 molecule. This agrees with the prediction of VSEPR theory, which predicts that there are 3 pairs of non-bonding electrons around the equatorial region of the xenon atom.
At high pressures, novel, non-molecular forms of xenon difluoride can be obtained. Under a pressure of ~50 GPa, XeF
2 transforms into a semiconductor consisting of XeF
4 units linked in a two-dimensional structure, like graphite. At even higher pressures, above 70 GPa, it becomes metallic, forming a three-dimensional structure containing XeF
8 units. However, a recent theoretical study has cast doubt on these experimental results.
The Xe–F bonds are weak. XeF2 has a total bond energy of 64 kcal/mol, with first and second bond energies of 44 kcal/mol and 20 kcal/mol, respectively. However, XeF2 is much more robust than KrF2, which has a total bond energy of only 22 kcal/mol.
Synthesis proceeds by the simple reaction:
- Xe + F2 → XeF2
The first published report of XeF2 was in October 1962 by Chernick, et al. However, though published later, XeF2 was probably first created by Rudolf Hoppe at the University of Münster, Germany, in early 1962, by reacting fluorine and xenon gas mixtures in an electrical discharge. Shortly after these reports, Weeks, Cherwick, and Matheson of Argonne National Laboratory reported the synthesis of XeF2 using an all-nickel system with transparent alumina windows, in which equal parts Xe and F2 gases react at low pressure upon irradiation by an ultraviolet source to give XeF2. Williamson reported that the reaction works equally well at atmospheric pressure in a dry Pyrex glass bulb using sunlight as a source. It was noted that the synthesis worked even on cloudy days.
In the previous syntheses the F2 reactant had been purified to remove HF. Šmalc and Lutar found that if this step is skipped the reaction rate proceeds at four times the original rate.
Derived xenon compoundsEdit
Other xenon compounds may be derived from xenon difluoride. The unstable organoxenon compound Xe(CF
2 can be made by irradiating hexafluoroethane to generate CF
3· radicals and passing the gas over XeF
2. The resulting waxy white solid decomposes completely within 4 hours at room temperature.
The XeF+ cation is formed by combining xenon difluoride with a strong fluoride acceptor, such as an excess of liquid antimony pentafluoride (SbF
2 + SbF
5 → XeF+
Adding xenon gas to this pale yellow solution at a pressure of 2-3 atm produces a green solution containing the paramagnetic Xe+
2 ion, which contains a Xe−Xe bond: ("apf" denotes solution in liquid SbF
- 3 Xe (g) + XeF+
(apf) + SbF
5 (l) ⇌ 2 Xe+
2 (apf) + SbF−
This reaction is reversible; removing xenon gas from the solution causes the Xe+
2 ion to revert to xenon gas and XeF+
, and the color of the solution returns to a pale yellow.
In the presence of liquid HF, dark green crystals can be precipitated from the green solution at −30 °C:
2 (apf) + 4 SbF−
6 (apf) → Xe+
21 (s) + 3 F−
X-ray crystallography indicates that the Xe–Xe bond length in this compound is 309 pm, indicating a very weak bond. The Xe+
2 ion is isoelectronic with the I−
2 ion, which is also dark green.
Bonding in the XeF2 molecule is adequately described by the three-center four-electron bond model.
- Mg(AsF6)2 + 4 XeF2 → [Mg(XeF2)4](AsF6)2
Crystallographic analysis shows that the magnesium atom is coordinated to 6 fluorine atoms. Four of the fluorines are attributed to the four xenon difluoride ligands while the other two are a pair of cis-AsF−
A similar reaction is:
- Mg(AsF6)2 + 2 XeF2 → [Mg(XeF2)2](AsF6)2
In the crystal structure of this product the magnesium atom is octahedrally-coordinated and the XeF2 ligands are axial while the AsF−
6 ligands are equatorial.
Many such reactions with products of the form [Mx(XeF2)n](AF6)x have been observed, where M can be Ca, Sr, Ba, Pb, Ag, La, or Nd and A can be As, Sb or P.
Recently, a compound was synthesised where a metal atom was coordinated solely by XeF2 fluorine atoms:
- 2 Ca(AsF6 )2 + 9 XeF2 → Ca2(XeF2)9(AsF6)4.
This reaction requires a large excess of xenon difluoride. The structure of the salt is such that half of the Ca2+ ions are coordinated by fluorine atoms from xenon difluoride, while the other Ca2+ ions are coordinated by both XeF2 and AsF−
As a fluorinating agentEdit
Among the fluorination reactions that xenon difluoride undergoes are:
- Oxidative fluorination:
- Ph3TeF + XeF2 → Ph3TeF3 + Xe
- Reductive fluorination:
- 2 CrO2F2 + XeF2 → 2 CrOF3 + Xe +O2
- Aromatic fluorination:
- Alkene fluorination:
- Radical fluorination in radical decarboxylative fluorination reactions, in Hunsdiecker-type reactions where xenon difluoride is used to generate the radical intermediate as well as the fluorine transfer source, and in generating aryl radicals from aryl silanes:
2 is selective about which atom it fluorinates, making it a useful reagent for fluorinating heteroatoms without touching other substituents in organic compounds. For example, it fluorinates the arsenic atom in trimethylarsine, but leaves the methyl groups untouched:
3As + XeF
2 → (CH
2 + Xe
- [R–N+(CH2CH2)3N:][BF4−] + XeF2 + NaBF4 → [R–N+(CH2CH2)3N+–F][BF4−]2 + NaF + Xe
- RCOOH + XeF2 → RF + CO2 + Xe + HF
As an etchantEdit
Xenon difluoride is also used as an isotropic gaseous etchant for silicon, particularly in the production of microelectromechanical systems, (MEMS), as first demonstrated in 1995. Commercial systems use pulse etching with an expansion chamber  Brazzle, Dokmeci, et al., describe this process:
The mechanism of the etch is as follows. First, the XeF2 adsorbs and dissociates to xenon (Xe) and fluorine (F) on the surface of silicon. Fluorine is the main etchant in the silicon etching process. The reaction describing the silicon with XeF2 is
- 2 XeF2 + Si → 2 Xe + SiF4
XeF2 has a relatively high etch rate and does not require ion bombardment or external energy sources in order to etch silicon.
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