Krypton difluoride

Krypton difluoride  |

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Names
Other names

Krypton fluoride
Krypton(II) fluoride

Identifiers

CAS Number

* 13773-81-4 Y

3D model (JSmol)

* Interactive image

ChemSpider |

* 75543 Y

PubChem CID

* 83721

UNII |

* A91DJL4OJC Y

InChI

* InChI=1S/KrF2/c1-3-2 Y

Key: QGOSZQZQVQAYFS-UHFFFAOYSA-N Y

* 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[1]

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).

N verify (what is YN ?)
Infobox references

Tracking categories (test):

* PairSizeSet

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 Kr2F+3 cations.[3]

The atomization energy of KrF2 (KrF2(g) → Kr(g) \+ 2F(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, difluorine is held together by a bond 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] Krypton difluoride is endothermic, with a heat of formation of 14.4 ± 0.8 kcal/mol measured at 93 °C.[5]

## Contents

* 1 Synthesis
* 1.1 Electrical discharge
* 1.2 Proton bombardment
* 1.3 Photochemical
* 1.4 Hot wire
* 2 Structure
* 3 Chemistry
* 4 See also
* 5 References
* 6 General reading
* 7 External links

## 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 α-particles, 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

β-KrF2

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.

## Chemistry

Krypton difluoride is primarily a powerful oxidising and fluorinating agent: for example, it can oxidise gold to its highest-known oxidation state, +5. It is more powerful even than elemental fluorine due to the even lower bond energy of Kr–F compared to F–F, with a redox potential of +3.5 V for the KrF2/Kr couple, making it the most powerful known oxidising agent, though KrF4 could be even stronger:[11]

7 KrF2 (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:[12]

KrF+AuF−6 → AuF5 (s) + Kr (g) + F2 (g)

KrF2 can also directly oxidise xenon to xenon hexafluoride:[11]

3 KrF2 \+ Xe → XeF6 \+ 3 Kr

KrF2 is used to synthesize the highly reactive BrF+6 cation.[6] KrF2 reacts with SbF5 to form the salt KrF+SbF−6; the KrF+ cation is capable of oxidising both BrF5 and ClF5 to BrF+6 and ClF+6, respectively.[13]

KrF2 is able to oxidise silver to its +3 oxidation state, reacting with elemental silver or with AgF to produce AgF3.[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]

## See also

* Krypton fluoride laser

## References

1. ↑ 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.
2. ↑ 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. Bibcode: 1963Sci...139.1047G.
3. ↑ 3.0 3.1 3.2 3.3 3.4 3.5 3.6 Lehmann, J (1 November 2002). "The chemistry of krypton". Coordination Chemistry Reviews 233–234: 1–39. doi:10.1016/S0010-8545(02)00202-3.
4. ↑ The values of De(F–KrF) and De(F–Kr•) are estimated to be comparable, at ~10-12 kcal/mol, while ΔH(KrF+ → Kr+ \+ F•) is estimated to be ~42 kcal/mol.
5. ↑ 5.0 5.1 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.
6. ↑ 6.0 6.1 Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils, ed., Inorganic Chemistry, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
7. ↑ 7.0 7.1 7.2 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.
8. ↑ MacKenzie, D. R.; Fajer, J. (1966). "Synthesis of Noble Gas Compounds by Proton Bombardment". Inorganic Chemistry 5 (4): 699–700. doi:10.1021/ic50038a048.
9. ↑ Xu, Ruren; Pang, Wenqin; Huo, Qisheng (2010). Modern Inorganic Synthetic Chemistry. Burlington: Elsevier Science. p. 54. ISBN 9780444536006. https://books.google.com/books?id=mJBW5pLF5R8C&pg=PA54. Retrieved 8 April 2017.
10. ↑ Jaffe, Mark (April 30, 1995). "Lucia V. Streng, 85; Innovative Chemist At Temple University". The Philadelphia Inquirer. http://articles.philly.com/1995-04-30/news/25687426_1_krypton-painting-china-chemist.
11. ↑ 11.0 11.1 W. Henderson (2000). Main group chemistry. Great Britain: Royal Society of Chemistry. p. 149. ISBN 0-85404-617-8. https://archive.org/details/maingroupchemist00hend_891.
12. ↑ 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.
13. ↑ 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. https://archive.org/details/isbn_0120236451.
14. ↑ A. Earnshaw; Norman Greenwood (1997). Chemistry of the Elements (2nd ed.). Elsevier. p. 903. ISBN 9780080501093.
15. ↑ Bougon, Roland (1984). "Synthesis and properties of silver trifluoride AgF3". Inorganic Chemistry 23 (22): 3667–3668. doi:10.1021/ic00190a049.
16. ↑ 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. doi:10.1063/1.1725990. ISSN 0021-9606. Bibcode: 1964JChPh..41..902F.

## General reading

## External links

* NIST Chemistry WebBook: krypton difluoride

* v
* t
* e

Fluorine compounds

| HF |  |  |  |  |  |  |  |  |  |  |  |  |  |  |  |  | He

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LiF | BeF2 |  |  |  |  |  |  |  |  |  |  | BF3
B2F4 | CF4 | NF3 | OF2 | F− | Ne
NaF | MgF2 |  |  |  |  |  |  |  |  |  |  | AlF3 | SiF4 | PF3
PF5 | SF2
SF4
SF6 | ClF
ClF3
ClF5 | Ar
KF | CaF2 | ScF3 | TiF3
TiF4 | VF3
VF4
VF5 | CrF2
CrF3
CrF4
CrF5 | MnF2
MnF3 | FeF2
FeF3 | CoF2
CoF3 | NiF2 | CuF
CuF2 | ZnF2 | GaF3 | GeF4 | AsF3
AsF5 | SeF4
SeF6 | BrF3
BrF5 | KrF2
RbF | SrF2 | YF3 | ZrF4 | NbF4
NbF5 | MoF5
MoF6 | TcF6 | RuF3
RuF4
RuF5
RuF6 | RhF6 | PdF2 | AgF
AgF2
Ag2F | CdF2 | InF3 | SnF2
SnF4 | SbF3
SbF5 | TeF4 | IF
IF3
IF5
IF7 | XeF2
XeF4
XeF6
CsF | BaF2 | *  | HfF4 | TaF5 | WF4
WF6 | ReF4
ReF5
ReF6
ReF7 | OsF4
OsF5
OsF6 | IrF3
IrF5
IrF6 | PtF4
PtF6 | AuF3
Au2F10
AuF5·F2
| HgF2
Hg2F2 | TlF
TlF3 | PbF2
PbF4 | BiF3
BiF5 | Po | At | RnF2
Fr | Ra | **  | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og
|  | ↓ |  |  |  |  |  |  |  |  |  |  |  |  |  |  |
| * | LaF3 | CeF3 | Pr | NdF3 | PmF3 | SmF3
| Eu | Gd | Tb | DyF3 | Ho | Er | Tm | YbF3
| Lu |
| ** | AcF3 | ThF4 | Pa | UF3
UF4
UF5
UF6 | NpF3
NpF4
NpF5
NpF6
| PuF3
PuF4
PuF5
PuF6 | AmF3
AmF4 | CmF3 | Bk | Cf | Es | Fm | Md | No | Lr |

Other|

* AgBF4
* AgPF6
* Cs2AlF5
* K3AlF6
* Na3AlF6
* KAsF6
* LiAsF6
* NaAsF6
* HBF4
* KBF4
* LiBF4
* NaBF4
* RbBF4
* Ba(BF4)2
* Ni(BF4)2
* Pb(BF4)2
* Sn(BF4)2
* BaClF
* BaSiF6
* BaGeF6
* BrOF3
* BrO2F
* CBrF3
* CBr2F2
* CBr3F
* CClF3
* CCl2F2
* CCl3F
* CFN
* CF2O
* CF3I
* CHF3
* CH2F2
* CH3F
* C2Cl3F3
* C2H3F
* C6H5F
* C7H5F3
* C15F33N
* ClFO2
* CrFO4
* CrF2O2
* CsBF4
* NH4F
* FNO
* FNO2
* FNO3
* KHF2
* NaHF2
* ThOF2
* NH5F2
* (NH4)2SiF6
* F2OS
* F3OP
* PSF3
* HPF6
* HSbF6
* NH4PF6
* KPF6
* KSbF6
* LiPF6
* NaPF6
* NaSbF6
* Na2SiF6
* Na2TiF6
* Na2ZrF6
* TlPF6
* IOF3
* K2NbF7
* K2TaF7
* IO3F
* UO2F2

Chemical formulas

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