Xenon difluoride

Last updated
Xenon difluoride
XeF 2 crystals. 1962. XeF2 kristalle.png
XeF
2 crystals. 1962.
Xenon difluoride Xenon-difluoride-2D.png
Xenon difluoride
Xenon difluoride Xenon-difluoride-3D-vdW.png
Xenon difluoride
Xenon-difluoride-xtal-3D-vdW.png
Names
IUPAC names
Xenon difluoride
Xenon(II) fluoride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.033.850 OOjs UI icon edit-ltr-progressive.svg
PubChem CID
UNII
  • InChI=1S/F2Xe/c1-3-2 Yes check.svgY
    Key: IGELFKKMDLGCJO-UHFFFAOYSA-N Yes check.svgY
  • InChI=1/F2Xe/c1-3-2
    Key: IGELFKKMDLGCJO-UHFFFAOYAE
  • F[Xe]F
Properties
F2Xe
Molar mass 169.290 g·mol−1
AppearanceWhite solid
Density 4.32 g/cm3, solid
Melting point 128.6 °C (263.5 °F; 401.8 K) [1]
25 g/L (0 °C)
Vapor pressure 6.0×102 Pa [2]
Structure
parallel linear XeF2 units
Linear
0 D
Thermochemistry
Std molar
entropy (S⦵298)
254 J·mol−1·K−1 [3]
−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. [4]
GHS labelling:
GHS-pictogram-acid.svg GHS-pictogram-skull.svg GHS-pictogram-rondflam.svg
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 [5]
NFPA 704 (fire diamond)
NFPA 704.svgHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
3
0
1
OX
Safety data sheet (SDS) PELCHEM MSDS
Related compounds
Other anions
Xenon dichloride
Xenon dibromide
Other cations
Krypton difluoride
Radon difluoride
Related compounds
 Xenon tetrafluoride
Xenon hexafluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Yes check.svgY  verify  (what is  Yes check.svgYX mark.svgN ?)

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 water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.

Contents

It has a nauseating odour and low vapor pressure. [6]

Structure

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. [2]

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. [7] However, a recent theoretical study has cast doubt on these experimental results. [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]

Chemistry

Synthesis

Synthesis proceeds by the simple reaction:

Xe + F2 → XeF2

The reaction needs heat, irradiation, or an electrical discharge. The product is a solid. It is purified by fractional 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 by Rudolf Hoppe at the University 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 of Argonne National Laboratory reported the synthesis of XeF2 using an all-nickel system with transparent alumina windows, in which equal parts xenon and fluorine gases react at low pressure upon irradiation by an ultraviolet source to give XeF2. [14] 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. [15]

In the previous syntheses the fluorine gas reactant had been purified to remove hydrogen 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 with dioxygen difluoride. [17]

Solubility

XeF
2 is soluble in solvents such as BrF
5, BrF
3, IF
5, anhydrous hydrogen fluoride, and acetonitrile, without reduction or oxidation. Solubility in hydrogen fluoride is high, at 167 g per 100 g HF at 29.95 °C. [2]

Derived xenon compounds

Other xenon compounds may be derived from xenon difluoride. The unstable organoxenon compound Xe(CF
3)
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. [18]

The XeF+ cation is formed by combining xenon difluoride with a strong fluoride acceptor, such as an excess of liquid antimony pentafluoride (SbF
5):

XeF
2 + SbF
5 → XeF+
 + SbF−
6

Adding xenon gas to this pale yellow solution at a pressure of 2–3 atmospheres produces a green solution containing the paramagnetic Xe+
2 ion, [19] which contains a Xe−Xe bond: ("apf" denotes solution in liquid SbF
5)

3 Xe(g) + XeF+
(apf) + SbF
5(l)⇌ 2 Xe+
2(apf) + SbF−
6(apf)

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. [20]

In the presence of liquid HF, dark green crystals can be precipitated from the green solution at −30 °C:

Xe+
2(apf) + 4 SbF−
6(apf) → Xe+
2Sb
4F−
21(s) + 3 F−
(apf)

X-ray crystallography indicates that the Xe–Xe bond length in this compound is 309  pm, indicating a very weak bond. [18] The Xe+
2 ion is isoelectronic with the I−
2 ion, which is also dark green. [21] [22]

Coordination chemistry

Bonding in the XeF2 molecule is adequately described by the three-center four-electron bond model.

XeF2 can act as a ligand in coordination complexes of metals. [2] For example, in HF solution:

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 fluorine atoms are attributed to the four xenon difluoride ligands while the other two are a pair of cis-AsF−
6 ligands. [23]

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 calcium, strontium, barium, lead, silver, lanthanum, or neodymium and A can be arsenic, antimony or phosphorus. Some of these compounds feature extraordinarily high coordination 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:

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−
6.

Applications

As a fluorinating agent

Xenon difluoride is a strong fluorinating and oxidizing agent. [26] [27] With fluoride ion acceptors, it forms XeF+
 and Xe
2F+
3 species which are even more powerful fluorinators. [2]

Among the fluorination reactions that xenon difluoride undergoes are:

Ph3TeF + XeF2 → Ph3TeF3 + Xe
2 CrO2F2 + XeF2 → 2 CrOF3 + Xe +O2
Fluor1.png
Fluor2.png
Fluor3.png
Xe decarboxylation.tif
Xe silanes.tif

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 in trimethylarsine, but leaves the methyl groups untouched: [30]

(CH
3)
3As + XeF
2 → (CH
3)
3AsF
2 + Xe

XeF2 can similarly be used to prepare N-fluoroammonium salts, useful as fluorine transfer reagents in organic synthesis (e.g., Selectfluor), from the corresponding tertiary amine: [31]

[R–+N(CH2CH2)3N:][BF−
4] + XeF2 + NaBF4 → [R–+N(CH2CH2)3+N–F][BF−
4]2 + NaF + Xe

XeF
2 will also oxidatively decarboxylate carboxylic acids to the corresponding fluoroalkanes: [32] [33]

RCOOH + XeF2 → RF + CO2 + Xe + HF

Silicon tetrafluoride has been found to act as a catalyst in fluorination by XeF
2. [34]

As an etchant

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. [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

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.

Related Research Articles

In chemistry, noble gas compounds are chemical compounds that include an element from the noble gases, group 8 or 18 of the periodic table. Although the noble gases are generally unreactive elements, many such compounds have been observed, particularly involving the element xenon.

In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

<span class="mw-page-title-main">Oxygen fluoride</span> Any binary compound of oxygen and fluorine

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:

<span class="mw-page-title-main">Xenon tetrafluoride</span> Chemical compound

Xenon tetrafluoride is a chemical compound with chemical formula XeF
4. It was the first discovered binary compound of a noble gas. It is produced by the chemical reaction of xenon with fluorine:

<span class="mw-page-title-main">Xenon hexafluoride</span> Chemical compound

Xenon hexafluoride is a noble gas compound with the formula XeF6. It is one of the three binary fluorides of xenon that have been studied experimentally, the other two being XeF2 and XeF4. All known are exergonic and stable at normal temperatures. XeF6 is the strongest fluorinating agent of the series. It is a colorless solid that readily sublimes into intensely yellow vapors.

<span class="mw-page-title-main">Platinum hexafluoride</span> Chemical compound

Platinum hexafluoride is the chemical compound with the formula PtF6, and is one of seventeen known binary hexafluorides. It is a dark-red volatile solid that forms a red gas. The compound is a unique example of platinum in the +6 oxidation state. With only four d-electrons, it is paramagnetic with a triplet ground state. PtF6 is a strong fluorinating agent and one of the strongest oxidants, capable of oxidising xenon and O2. PtF6 is octahedral in both the solid state and in the gaseous state. The Pt-F bond lengths are 185 picometers.

<span class="mw-page-title-main">Silver(II) fluoride</span> Chemical compound

Silver(II) fluoride is a chemical compound with the formula AgF2. It is a rare example of a silver(II) compound - silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.

Xenon compounds are compounds containing the element xenon (Xe). After Neil Bartlett's discovery in 1962 that xenon can form chemical compounds, a large number of xenon compounds have been discovered and described. Almost all known xenon compounds contain the electronegative atoms fluorine or oxygen. The chemistry of xenon in each oxidation state is analogous to that of the neighboring element iodine in the immediately lower oxidation state.

<span class="mw-page-title-main">Selenium tetrafluoride</span> Chemical compound

Selenium tetrafluoride (SeF4) is an inorganic compound. It is a colourless liquid that reacts readily with water. It can be used as a fluorinating reagent in organic syntheses (fluorination of alcohols, carboxylic acids or carbonyl compounds) and has advantages over sulfur tetrafluoride in that milder conditions can be employed and it is a liquid rather than a gas.

<span class="mw-page-title-main">Xenon oxytetrafluoride</span> Chemical compound

Xenon oxytetrafluoride is an inorganic chemical compound. It is an unstable colorless liquid with a melting point of −46.2 °C that can be synthesized by partial hydrolysis of XeF
6, or the reaction of XeF
6 with silica or NaNO
3:

<span class="mw-page-title-main">Krypton difluoride</span> Chemical compound

Krypton difluoride, KrF2 is a chemical compound of krypton and fluorine. It was the first compound of krypton discovered. It is a volatile, colourless solid at room temperature. The structure of the KrF2 molecule is linear, with Kr−F distances of 188.9 pm. It reacts with strong Lewis acids to form salts of the KrF+ and Kr
2F+
3 cations.

<span class="mw-page-title-main">Dioxygenyl</span> Chemical compound

The dioxygenyl ion, O+
2, is a rarely-encountered oxycation in which both oxygen atoms have a formal oxidation state of +⁠1/2⁠. It is formally derived from oxygen by the removal of an electron:

A hexafluoride is a chemical compound with the general formula QXnF6, QXnF6m−, or QXnF6m+. Many molecules fit this formula. An important hexafluoride is hexafluorosilicic acid (H2SiF6), which is a byproduct of the mining of phosphate rock. In the nuclear industry, uranium hexafluoride (UF6) is an important intermediate in the purification of this element.

<span class="mw-page-title-main">Tetrafluoroammonium</span> Chemical compound

The tetrafluoroammonium cation is a positively charged polyatomic ion with chemical formula NF+
4. It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine. Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF
4, trifluoramine oxide ONF
3, tetrafluoroborate BF−
4 anion and the tetrafluoroberyllate BeF2−
4 anion.

<span class="mw-page-title-main">Chromyl fluoride</span> Chemical compound

Chromyl fluoride is an inorganic compound with the formula CrO2F2. It is a violet-red colored crystalline solid that melts to an orange-red liquid.

Organoxenon chemistry is the study of the properties of organoxenon compounds, which contain carbon to xenon chemical bonds. The first organoxenon compounds were divalent, such as (C6F5)2Xe. The first tetravalent organoxenon compound, [C6F5XeF2][BF4], was synthesized in 2004. So far, more than one hundred organoxenon compounds have been researched.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

<span class="mw-page-title-main">Xenon fluoride nitrate</span> Chemical compound

Xenon fluoride nitrate, also known as fluoroxenonium nitrate, is the chemical compound with formula FXeONO2.

Radical fluorination is a type of fluorination reaction, complementary to nucleophilic and electrophilic approaches. It involves the reaction of an independently generated carbon-centered radical with an atomic fluorine source and yields an organofluorine compound.

<span class="mw-page-title-main">Radon compounds</span>

Radon compounds are chemical compounds formed by the element radon (Rn). Radon is a noble gas, i.e. a zero-valence element, and is chemically not very reactive. The 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer. Because radon is a gas under normal circumstances, and its decay-chain parents are not, it can readily be extracted from them for research.

References

  1. ↑ Hindermann, D. K., Falconer, W. E. (1969). "Magnetic Shielding of 19F in XeF2". J. Chem. Phys. 50 (3): 1203. Bibcode:1969JChPh..50.1203H. doi:10.1063/1.1671178.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  2. 1 2 3 4 5 Melita Tramšek; Boris Žemva (2006). "Synthesis, Properties and Chemistry of Xenon(II) Fluoride" (PDF). Acta Chim. Slov. 53 (2): 105–116. doi:10.1002/chin.200721209.
  3. 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN   978-0-618-94690-7.
  4. ↑ "MSDS: xenon difluoride" (PDF). BOC Gases. Retrieved 2010-06-01.
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  7. ↑ Kim, M.; Debessai, M.; Yoo, C. S. (2010). "Two- and three-dimensional extended solids and metallization of compressed XeF2". Nature Chemistry. 2 (9): 784–788. Bibcode:2010NatCh...2..784K. doi:10.1038/nchem.724. PMID   20729901.
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  9. ↑ Cockett, A. H.; Smith, K. C.; Bartlett, Neil (2013). The Chemistry of the Monatomic Gases. Pergamon Texts in Inorganic Chemistry. St. Louis, MO: Elsevier Science. ISBN   9781483157368. OCLC   953379200.
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  13. ↑ Hoppe, R. (1964). "Die Valenzverbindungen der Edelgase". Angewandte Chemie. 76 (11): 455. Bibcode:1964AngCh..76..455H. doi:10.1002/ange.19640761103. First review on the subject by the pioneer of covalent noble gas compounds.
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  29. ↑ Lothian, A. P.; Ramsden, C. A. (1993). "Rapid fluorodesilylation of aryltrimethylsilanes using xenon difuoride: An efficient new route to aromatic fluorides". Synlett. 1993 (10): 753–755. doi:10.1055/s-1993-22596. S2CID   196734038.
  30. ↑ W. Henderson (2000). Main group chemistry . Great Britain: Royal Society of Chemistry. p.  150. ISBN   978-0-85404-617-1.
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  32. ↑ Patrick, Timothy B.; Johri, Kamalesh K.; White, David H.; Bertrand, William S.; Mokhtar, Rodziah; Kilbourn, Michael R.; Welch, Michael J. (1986). "Replacement of the carboxylic acid function with fluorine". Can. J. Chem. 64: 138–141. doi:10.1139/v86-024.
  33. ↑ Grakauskas, Vytautas (1969). "Aqueous fluorination of carboxylic acid salts". J. Org. Chem. 34 (8): 2446–2450. doi:10.1021/jo01260a040.
  34. ↑ Tamura Masanori; Takagi Toshiyuki; Shibakami Motonari; Quan Heng-Dao; Sekiya Akira (1998). "Fluorination of olefins with xenon difluoride-silicon tetrafluoride". Fusso Kagaku Toronkai Koen Yoshishu (in Japanese). 22: 62–63. Journal code: F0135B; accession code: 99A0711841.
  35. ↑ Chang, F.; Yeh, R.; G., Lin; Chu, P.; Hoffman, E.; Kruglick, E.; Pister, K.; Hecht, M. (1995). Bailey, Wayne; Motamedi, M. Edward; Luo, Fang-Chen (eds.). "Gas-phase silicon micromachining with xenon difluoride". SPIE Proceedings. 2641: 117–128. Bibcode:1995SPIE.2641..117C. doi:10.1117/12.220933. S2CID   39522253.
  36. ↑ Chu, P.; Chen, J.; Chu, P.; Lin, G.; Huang, J.; Warneke, B; Pister, K. (1997). Controlled Pulse-Etching with Xenon Difluoride. Int. Conf. Solid State Sensors and Actuators (Transducers 97). pp. 665–668.
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