Krypton difluoride

Krypton difluoride
Skeletal formula of krypton difluoride with a dimension	Spacefill model of krypton difluoride
Names
	IUPAC name Krypton difluoride
	Other names Krypton fluoride; Krypton(II) fluoride
Identifiers
CAS Number	13773-81-4 ;
3D model (JSmol)	Interactive image;
ChemSpider	75543 ;
PubChem CID	83721;
UNII	A91DJL4OJC ;
CompTox Dashboard (EPA)	DTXSID7065623 ;
	InChI InChI=1S/KrF2/c1-3-2 Key: QGOSZQZQVQAYFS-UHFFFAOYSA-N ; InChI=1/F2Kr/c1-3-2 Key: QGOSZQZQVQAYFS-UHFFFAOYAJ;
	SMILES F[Kr]F;
Properties
Chemical formula	F2Kr
Molar mass	121.795 g·mol−1
Appearance	Colourless crystals (solid)
Density	3.24 g cm−3 (solid)
Solubility in water	Reacts
Structure
Crystal structure	Body-centered tetragonal
Space group	P42/mnm, No. 136
Lattice constant	a = 0.4585 nm, c = 0.5827 nm
Molecular shape	Linear
Dipole moment	0 D
Related compounds
Related compounds	Xenon difluoride
	Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). verify (what is ?) Infobox references

Krypton difluoride, KrF₂ is a chemical compound of krypton and fluorine. It was the first compound of krypton discovered.^[2] 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.^[3]

The atomization energy of KrF₂ (KrF_2(g) → Kr_(g) + 2 F_(g)) is 21.9 kcal/mol, giving an average Kr–F bond energy of only 11 kcal/mol,^[4] the weakest of any isolable fluoride. In comparison, the dissociation of difluorine to atomic fluorine requires cleaving a F–F bond with a bond dissociation energy of 36 kcal/mol. Consequently, KrF₂ is a good source of the extremely reactive and oxidizing atomic fluorine. It is thermally unstable, with a decomposition rate of 10% per hour at room temperature.^[5] The formation of krypton difluoride is endothermic, with a heat of formation (gas) of 14.4 ± 0.8 kcal/mol measured at 93 °C.^[5]

Synthesis

Krypton difluoride can be synthesized using many different methods including electrical discharge, photoionization, hot wire, and proton bombardment. The product can be stored at −78 °C without decomposition.^[6]

Electrical discharge

Electric discharge was the first method used to make krypton difluoride. It was also used in the only experiment ever reported to produce krypton tetrafluoride, although the identification of krypton tetrafluoride was later shown to be mistaken. The electrical discharge method involves having 1:1 to 2:1 mixtures of F₂ to Kr at a pressure of 40 to 60 torr and then arcing large amounts of energy between it. Rates of almost 0.25 g/h can be achieved. The problem with this method is that it is unreliable with respect to yield.^[3]^[7]

Proton bombardment

Using proton bombardment for the production of KrF₂ has a maximum production rate of about 1 g/h. This is achieved by bombarding mixtures of Kr and F₂ with a proton beam operating at an energy level of 10 MeV and at a temperature of about 133 K. It is a fast method of producing relatively large amounts of KrF₂, but requires a source of high-energy protons, which usually would come from a cyclotron.^[3]^[8]

Photochemical

The successful photochemical synthesis of krypton difluoride was first reported by Lucia V. Streng in 1963. It was next reported in 1975 by J. Slivnik.^[9]^[10]^[3] The photochemical process for the production of KrF₂ involves the use of UV light and can produce under ideal circumstances 1.22 g/h. The ideal wavelengths to use are in the range of 303–313 nm. Harder UV radiation is detrimental to the production of KrF₂. Using Pyrex glass or Vycor or quartz will significantly increase yield because they all block harder UV light. In a series of experiments performed by S. A Kinkead et al., it was shown that a quartz insert (UV cut off of 170 nm) produced on average 158 mg/h, Vycor 7913 (UV cut off of 210 nm) produced on average 204 mg/h and Pyrex 7740 (UV cut off of 280 nm) produced on average 507 mg/h. It is clear from these results that higher-energy ultraviolet light reduces the yield significantly. The ideal circumstances for the production KrF₂ by a photochemical process appear to occur when krypton is a solid and fluorine is a liquid, which occur at 77 K. The biggest problem with this method is that it requires the handling of liquid F₂ and the potential of it being released if it becomes overpressurized.^[3]^[7]

Hot wire

The hot wire method for the production of KrF₂ uses krypton in a solid state with a hot wire running a few centimeters away from it as fluorine gas is then run past the wire. The wire has a large current, causing it to reach temperatures around 680 °C. This causes the fluorine gas to split into its radicals, which then can react with the solid krypton. Under ideal conditions, it has been known to reach a maximum yield of 6 g/h. In order to achieve optimal yields the gap between the wire and the solid krypton should be 1 cm, giving rise to a temperature gradient of about 900 °C/cm. A major downside to this method is the amount of electricity that has to be passed through the wire. It is dangerous if not properly set up.^[3]^[7]

Structure

Krypton difluoride can exist in one of two possible crystallographic morphologies: α-phase and β-phase. β-KrF₂ generally exists at above −80 °C, while α-KrF₂ is more stable at lower temperatures.^[3] The unit cell of α-KrF₂ is body-centred tetragonal.

Reactions

Krypton difluoride is primarily a powerful oxidising and fluorinating agent, more powerful even than elemental fluorine because Kr–F has less bond energy. It has a redox potential of +3.5 V for the KrF₂/Kr couple,^{[citation needed]} making it the most powerful known oxidising agent. However, the hypothetical KrF
₄ could be even stronger^[11] and nickel tetrafluoride comes close.

For example, krypton difluoride can oxidise gold to its highest-known oxidation state, +5:

7 KrF₂ + 2 Au → 2 KrF⁺AuF−6 + 5 Kr

KrF⁺
AuF⁻
₆ decomposes at 60 °C into gold(V) fluoride and krypton and fluorine gases:^[12]

[KrF⁺][AuF−6] → AuF₅ + Kr + F₂

KrF
₂ can also directly oxidise xenon to xenon hexafluoride:^[11]

3 KrF₂ + Xe → XeF₆ + 3 Kr

KrF
₂ is used to synthesize the highly reactive BrF⁺
₆ cation.^[6] KrF
₂ reacts with SbF
₅ to form the salt KrF⁺
SbF⁻
₆; the KrF⁺
cation is capable of oxidising both BrF
₅ and ClF
₅ to BrF⁺
₆ and ClF⁺
₆, respectively.^[13]

KrF
₂ can also react with elemental silver to produce AgF
₃.^[14]^[15]

Irradiation of a crystal of KrF₂ at 77 K with γ-rays leads to the formation of the krypton monofluoride radical, KrF•, a violet-colored species that was identified by its ESR spectrum. The radical, trapped in the crystal lattice, is stable indefinitely at 77 K but decomposes at 120 K.^[16]

References

^ R. D. Burbank, W. E. Falconer and W. A. Sunder (1972). "Crystal Structure of Krypton Difluoride at −80 °C". Science. 178 (4067): 1285–1286. doi:10.1126/science.178.4067.1285. PMID 17792123. S2CID 96692996.
^ Grosse, A. V.; Kirshenbaum, A. D.; Streng, A. G.; Streng, L. V. (1963). "Krypton Tetrafluoride: Preparation and Some Properties". Science. 139 (3559): 1047–8. Bibcode:1963Sci...139.1047G. doi:10.1126/science.139.3559.1047. PMID 17812982.
^ ^a ^b ^c ^d ^e ^f ^g Lehmann, J (1 November 2002). "The chemistry of krypton". Coordination Chemistry Reviews. 233–234: 1–39. doi:10.1016/S0010-8545(02)00202-3.
^ The values of D_e(F–KrF) and D_e(F–Kr•) are estimated to be comparable, at ~10-12 kcal/mol, while ΔH(KrF⁺ → Kr⁺ + F•) is estimated to be ~42 kcal/mol.
^ ^a ^b Cockett, A. H.; Smith, K. C.; Bartlett, Neil (1973). The Chemistry of the Monatomic Gases: Pergamon Texts in Inorganic Chemistry. Pergamon Press. ISBN 978-0-08-018782-2.
^ ^a ^b Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
^ ^a ^b ^c Kinkead, S. A.; Fitzpatrick, J. R.; Foropoulos, J. Jr.; Kissane, R. J.; Purson, D. (1994). "3. Photochemical and thermal Dissociation Synthesis of Krypton Difluoride". Inorganic Fluorine Chemistry: Toward the 21st Century. San Francisco, California: American Chemical Society. pp. 40–54. doi:10.1021/bk-1994-0555.ch003. ISBN 978-0-8412-2869-6.
^ MacKenzie, D. R.; Fajer, J. (1966). "Synthesis of Noble Gas Compounds by Proton Bombardment". Inorganic Chemistry. 5 (4): 699–700. doi:10.1021/ic50038a048.
^ Xu, Ruren; Pang, Wenqin; Huo, Qisheng (2010). Modern Inorganic Synthetic Chemistry. Burlington: Elsevier Science. p. 54. ISBN 9780444536006. Retrieved 8 April 2017.
^ Jaffe, Mark (April 30, 1995). "Lucia V. Streng, 85; Innovative Chemist At Temple University". The Philadelphia Inquirer. Archived from the original on 16 March 2016. Retrieved 24 August 2016.
^ ^a ^b W. Henderson (2000). Main group chemistry. Great Britain: Royal Society of Chemistry. p. 149. ISBN 0-85404-617-8.
^ Charlie Harding; David Arthur Johnson; Rob Janes (2002). Elements of the p block. Great Britain: Royal Society of Chemistry. p. 94. ISBN 0-85404-690-9.
^ John H. Holloway; Eric G. Hope (1998). A. G. Sykes (ed.). Advances in Inorganic Chemistry. Academic Press. pp. 60–61. ISBN 0-12-023646-X.
^ A. Earnshaw; Norman Greenwood (1997). Chemistry of the Elements (2nd ed.). Elsevier. p. 903. ISBN 9780080501093.
^ Bougon, Roland (1984). "Synthesis and properties of silver trifluoride AgF3". Inorganic Chemistry. 23 (22): 3667–3668. doi:10.1021/ic00190a049.
^ Falconer, W. E.; Morton, J. R.; Streng, A. G. (1964-08-01). "Electron Spin Resonance Spectrum of KrF". The Journal of Chemical Physics. 41 (3): 902–903. Bibcode:1964JChPh..41..902F. doi:10.1063/1.1725990. ISSN 0021-9606.