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. 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]
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| Names | |||
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| IUPAC name
Krypton(II) fluoride
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| Other names
Krypton fluoride
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| Identifiers | |||
3D model (JSmol)
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PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |||
| F2Kr | |||
| Molar mass | 121.795 g·mol−1 | ||
| Appearance | Colourless crystals (solid) | ||
| Density | 3.24 g cm−3 (solid) | ||
| Reacts | |||
| Structure | |||
| Body-centered tetragonal[1] | |||
| P42/mnm, No. 136 | |||
a = 0.4585 nm, c = 0.5827 nm
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| Linear | |||
| 0 D | |||
| Related compounds | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Synthesis
Krypton difluoride can be synthesized using many different methods including electrical discharge, photochemical, hot wire, and proton bombardment. It can also be prepared by irradiating krypton with ultraviolet rays in a fluorine-argon gas mixture at liquid helium temperature. The product can be stored at −78 °C without decomposition.[4]
Electrical discharge
The first method used to make krypton difluoride and the only one ever reported to produce krypton tetrafluoride (although the identification of krypton tetrafluoride was later shown to be mistaken) was the electrical discharge method. 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][5]
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 that is 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 α-particles which usually would come from a cyclotron.[3][6]
Photochemical
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. It is important to note that harder UV radiation is detrimental to the production of KrF2. In order to avoid the harder wavelengths, simply 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 ultra violet 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 77K. 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 over pressurized.[3][5]
Hot wire
The hot wire method for the production of KrF2 involves having the 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. The only major downside to this method is the amount of electricity that has to be passed through the wire thus making it dangerous if not properly set up.[3][5]
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 the α-KrF2 is more stable at lower temperatures.[3] The unit cell of α-KrF2 is body-centred tetragonal.
Chemistry
Krypton difluoride is primarily a powerful oxidising and fluorinating agent. It can oxidise gold to its highest-known oxidation state, +5, it is more powerful even that elemental fluorine due to the reduced bond F-F to Kr-F with redox potencial of 3.5, making it the most powerful oxidising agent there is, though KrF
4 could be even stronger oxidising agent:[7]
- 7 KrF
2 (g) + 2 Au (s) → 2 KrF+
AuF−
6 (s) + 5 Kr (g)
KrF+
AuF−
6 decomposes at 60°C into gold(V) fluoride and krypton and fluorine gases:[8]
- KrF+
AuF−
6 → AuF
5 (s) + Kr (g) + F
2 (g)
KrF
2 can also directly oxidise xenon to xenon hexafluoride:[7]
- 3 KrF
2 + Xe → XeF
6 + 3 Kr
KrF
2 is used to synthesize the highly reactive BrF+
6 cation.[4] 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.[9]
KrF
2 is able to oxidise silver to its +3 oxidation state, reacting with elemental silver or with AgF to produce AgF
3.[10]
Related compounds
- Xenon difluoride, XeF2
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.
- ^ Grosse, A. V.; Kirshenbaum, A. D.; Streng, A. G.; Streng, L. V. (1963). "Krypton Tetrafluoride: Preparation and Some Properties". Science. 139 (3559): 1047–8. doi:10.1126/science.139.3559.1047. PMID 17812982.
- ^ a b c d e f Lehmann, J (2002). "The chemistry of krypton". Coordination Chemistry Reviews. 233–234: 1. doi:10.1016/S0010-8545(02)00202-3.
- ^ 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.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1021/ic50038a048, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1021/ic50038a048instead. - ^ 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.
General reading
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.