 2 crystals. 1962. | |
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| Names | |
|---|---|
| IUPAC names Xenon difluoride Xenon(II) fluoride | |
| Identifiers | |
3D model (JSmol) | |
| ChemSpider |
|
| ECHA InfoCard | 100.033.850 |
| UNII | |
| |
| |
| Properties | |
| F2Xe | |
| Molar mass | 169.290 g·mol−1 |
| Appearance | White solid |
| Density | 4.32 g/cm3, solid |
| Melting point | 128.6 °C (263.5 °F; 401.8 K)[2] |
| 25 g/L (0 °C, slow decomposition) | |
| Vapor pressure | 6.0×102 Pa[1] |
| Structure | |
| parallel linear XeF2 units | |
| Linear | |
| 0D | |
| Thermochemistry | |
Std molar entropy(S⦵298) | 254 J·mol−1·K−1[3] |
Std enthalpy of formation(ΔfH⦵298) | −108 kJ·mol−1[3] |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards | Corrosive to exposed tissues. Releases toxic compounds on contact with moisture.[5] |
| GHS labelling: | |
   | |
| Danger | |
| H272,H301,H314,H330 | |
| P210,P220,P221,P260,P264,P270,P271,P280,P284,P301+P310+P330,P303+P361+P353,P304+P340+P310,P305+P351+P338,P331,P363,P370+P378,P403+P233,P405,P501[4] | |
| NFPA 704 (fire diamond) | |
| Safety data sheet (SDS) | PELCHEM MSDS |
| Related compounds | |
Otheranions | Xenon dichloride Xenon dibromide |
Othercations | Krypton difluoride Radon difluoride |
Related compounds | Xenon tetrafluoride Xenon hexafluoride |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Xenon difluoride is a powerfulfluorinating agent with the chemical formulaXeF
2, and one of the most stablexenon compounds. Like mostcovalent inorganicfluorides, it is moisture-sensitive. It graduallydecomposes on contact withwater vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourlesscrystalline solid.
It has a nauseating odour and lowvapor pressure.[6]
Xenon difluoride is alinear molecule with an Xe–F bond length of197.73±0.15 pm in the vapor stage, and 200 pm in the solid phase. The packing arrangement in solidXeF
2 shows that the fluorine atoms of neighbouring molecules avoid the equatorial region of eachXeF
2 molecule. This agrees with the prediction ofVSEPR theory, which predicts that there are 3 pairs of non-bonding electrons around the equatorial region of the xenon atom.[1]
At high pressures, novel, polymeric forms of xenon difluoride can be obtained. Under a pressure of ~50 GPa,XeF
2 transforms into a semiconductor consisting ofXeF
4 units linked in a two-dimensional structure, likegraphite. At even higher pressures, above 70 GPa, it becomes metallic, forming a three-dimensional structure containingXeF
8 units.[7] A 2011 theoretical study cast doubt on these experimental results, suggesting that xenon difluoride remains stable up to 200 GPa, at which point itdissociates into an ionic solid.[8]
The Xe–F bonds are weak. XeF2 has a total bond energy of 267.8 kJ/mol (64.0 kcal/mol), with first and second bond energies of 184.1 kJ/mol (44.0 kcal/mol) and 83.68 kJ/mol (20.00 kcal/mol), respectively. However, XeF2 is much more robust than KrF2, which has a total bond energy of only 92.05 kJ/mol (22.00 kcal/mol).[9]
Synthesis proceeds by the simple reaction:
The reaction needs heat, irradiation, or an electrical discharge. The product is a solid. It is purified byfractional distillation or selective condensation using a vacuum line.[10]
The first published report of XeF2 was in October 1962 by Chernick, et al.[11] However, though published later,[12] XeF2 was probably first created byRudolf Hoppe at theUniversity of Münster, Germany, in early 1962, by reacting fluorine and xenon gas mixtures in an electrical discharge.[13] Shortly after these reports, Weeks, Chernick, and Matheson ofArgonne National Laboratory reported the synthesis of XeF2 using an all-nickel system with transparentalumina windows, in which equal parts xenon and fluorine gases react at low pressure upon irradiation by anultraviolet source to give XeF2.[14] Williamson reported that the reaction works equally well at atmospheric pressure in a dryPyrex glass bulb using sunlight as a source. It was noted that the synthesis worked even on cloudy days.[15]
In the previous syntheses the fluorine gas reactant had been purified to removehydrogen fluoride. Šmalc and Lutar found that if this step is skipped the reaction rate proceeds at four times the original rate.[16]
In 1965, it was also synthesized by reacting xenon gas withdioxygen difluoride.[17]
XeF
2 issoluble ininterhalogensolvents such asBrF
5,BrF
3,IF
5, and others like anhydroushydrogen fluoride, andacetonitrile, without reduction or oxidation. Solubility in hydrogen fluoride is high, at 167 g per 100 g HF at 29.95 °C.[1]
Other xenon compounds may be derived from xenon difluoride. The unstableorganoxenon compoundXe(CF
3)
2 can be made by irradiatinghexafluoroethane to generateCF•
3radicals and passing the gas overXeF
2. The resulting waxy white solid decomposes completely within 4 hours at room temperature.[18]
The XeF+ cation is formed by combining xenon difluoride with a strong fluoride acceptor, such as an excess of liquidantimony pentafluoride (SbF
5):
Adding xenon gas to this pale yellow solution at a pressure of 2–3atmospheres produces a green solution containing the paramagneticXe+
2 ion,[19] which contains a Xe−Xe bond: ("apf" denotes solution in liquidSbF
5)
This reaction is reversible; removing xenon gas from the solution causes theXe+
2 ion to revert to xenon gas andXeF+
, and the color of the solution returns to a pale yellow.[20]
In the presence of liquidHF, dark green crystals can be precipitated from the green solution at −30 °C:
X-ray crystallography indicates that the Xe–Xe bond length in this compound is 309 pm, indicating a very weak bond.[18] TheXe+
2 ion isisoelectronic with theI−
2 ion, which is also dark green.[21][22]
Bonding in the XeF2 molecule is adequately described by thethree-center four-electron bond model.
XeF2 can act as aligand incoordination complexes of metals.[1] For example, in HF solution:
Crystallographic analysis shows that the magnesium atom is coordinated to 6 fluorine atoms. Four of the fluorine atoms are attributed to the four xenon difluoride ligands while the other two are a pair ofcis-AsF−
6 ligands.[23]
A similar reaction is:
In the crystal structure of this product the magnesium atom isoctahedrally-coordinated and the XeF2 ligands are axial while theAsF−
6 ligands are equatorial.
Many such reactions with products of the form [Mx(XeF2)n](AF6)x have been observed, where M can becalcium,strontium,barium,lead,silver,lanthanum, orneodymium and A can bearsenic,antimony orphosphorus. Some of these compounds feature extraordinarily highcoordination numbers at the metal center.[24]
In 2004, results of synthesis of a solvate where part of cationic centers were coordinated solely by XeF2 fluorine atoms were published.[25] Reaction can be written as:
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 andAsF−
6.
Xenon difluoride is a strong fluorinating and oxidizing agent.[26][27] With fluoride ion acceptors, it formsXeF+
andXe
2F+
3 species which are even more powerful fluorinators.[1]
Among the fluorination reactions that xenon difluoride undergoes are:
XeF
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 intrimethylarsine, but leaves themethyl groups untouched:[30]
XeF2 can similarly be used to prepareN-fluoroammonium salts, useful as fluorine transfer reagents in organic synthesis (e.g.,Selectfluor), from the corresponding tertiary amine:[31]
XeF
2 will also oxidatively decarboxylatecarboxylic acids to the correspondingfluoroalkanes:[32][33]
Silicon tetrafluoride has been found to act as a catalyst in fluorination byXeF
2.[34]
Xenon difluoride is also used as an isotropic gaseousetchant forsilicon, particularly in the production ofmicroelectromechanical systems (MEMS), as first demonstrated in 1995.[35] Commercial systems use pulse etching with an expansion chamber[36]Brazzle, Dokmeci, et al. describe this process:[37]
The mechanism of the etch is as follows. First, the XeF2 adsorbs and dissociates to xenon and fluorine atoms on the surface of silicon. Fluorine is the main etchant in the silicon etching process. The reaction describing the silicon with XeF2 is
XeF2 has a relatively high etch rate and does not requireion bombardment or external energy sources in order to etch silicon.
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