Noble gas compound

Noble gas compounds are chemical compounds that include an element from the noble gases, group 18 of the periodic table. Many such compounds have been observed, particularly involving the element xenon.

History and background

It was initially believed that the noble gases could not form compounds due to their full valence shell of electrons that rendered them very chemically stable and nonreactive.

All noble gases have full s and p outer electron shells (except helium, which has no p sublevel), and so do not form chemical compounds easily. Because of their high ionization energy and almost zero electron affinity, they were not expected to be reactive.

In 1933, Linus Pauling predicted that the heavier noble gases would be able to form compounds with fluorine and oxygen. Specifically, he predicted the existence of krypton hexafluoride (KrF6) and xenon hexafluoride (XeF6), speculated that XeF8 might exist as an unstable compound, and suggested that xenic acid would form perxenate salts.[1][2] These predictions proved quite accurate, although subsequent predictions for XeF8 indicated that it would be not only thermodynamically unstable, but kinematically unstable.[3] As of 2013, XeF8 has not been made, although the octafluoroxenate(VI) anion (XeF2−
8) has been observed.

The heavier noble gases have more electron shells than the lighter ones. Hence, the outermost electrons are subject to a shielding effect from the inner electrons that makes them more easily ionized, since they are less strongly attracted to the positively charged nucleus. This results in an ionization energy low enough to form stable compounds with the most electronegative elements, fluorine and oxygen, and even with less electronegative elements such as nitrogen and carbon under certain circumstances.[4][5]

True noble gas compounds

These compounds are listed in order of decreasing order of the atomic weight of the noble gas, which generally reflects the priority of their discovery, and the breadth of available information for these compounds.

Radon compounds

Radon is not chemically inert, but its short half-life and the high energy of its radioactivity make it difficult to investigate its sole described fluoride, RnF2, and its reaction products.[6]

Xenon compounds

The first published report, in June 1962, of a noble gas compound was by Neil Bartlett, who noticed that the highly oxidising compound platinum hexafluoride ionised O2 to O+
2. As the ionisation energy of O2 to O+
2 (1165 kJ mol−1) is nearly equal to the ionisation energy of Xe to Xe+ (1170 kJ mol−1), he tried the reaction of Xe with PtF6. This yielded a crystalline product, xenon hexafluoroplatinate, whose formula was proposed to be Xe+
[PtF
6]−
.[2][7] It was later shown that the compound is actually more complex, containing both XeFPtF5 and XeFPt2F11.[8] This was the first real compound of any noble gas.

The first binary noble gas compounds were reported in 1962. Neil Bartlett synthesized Xenon tetrafluoride (XeF4) by subjecting a mixture of xenon and fluorine to high temperature.[9] Rudolf Hoppe, among other groups, synthesized xenon difluoride (XeF2) by the reaction of the elements.[10]

Since these initial studies, other xenon compounds that have been synthesized include other fluorides (XeF6), oxyfluorides (XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4) and oxides (XeO
2, XeO3 and XeO4). Xenon difluoride has been produced by the exposure of Xe and F2 gases to sunlight, a result which eluded observation for 50 years. Xenon fluorides react with several other fluorides to form fluoroxenates, such as sodium octafluoroxenate (Na+
2XeF2−
8), and fluoroxenonium salts, such as trifluoroxenonium hexafluoroantimoniate (XeF+
3SbF−
6).

In terms of other halide reactivity, short-lived excimers of noble gas halides such as XeCl2 are prepared in situ, and are used in the function of excimer lasers.

Recently, xenon has been shown to produce a wide variety of compounds of the type XeOnX2 where n is 1,2 or 3 and X is any electronegative group, such as CF3, C(SO2CF3)3, N(SO2F)2, N(SO2CF3)2, OTeF5, O(IO2F2), etc.; the range of compounds is impressive, similar to that seen with the neighbouring element iodine, running into the thousands and involving bonds between xenon and oxygen, nitrogen, carbon, boron and even gold, as well as perxenic acid, several halides, and complex ions.

A compound with an Xe-Xe bond has been reported, and is noteworthy: the compound Xe2Sb2F11 contains a Xe–Xe bond, and it is the longest element-element bond known (308.71 pm = 3.0871 Å).[11] Short-lived excimers of Xe2 are reported to exist as a part of the function of excimer lasers.

Krypton compounds

Following the first successful synthesis of xenon compounds, synthesis of krypton difluoride (KrF
2) was reported in 1963.[12] Krypton gas reacts with fluorine gas under extreme forcing conditions, forming KrF2 according to the following equation:

Kr + F2 → KrF2

KrF2 reacts with strong Lewis acids to form salts of the KrF+ and Kr
2F+
3 cations.[12] The preparation of KrF
4 reported by Grosse in 1963, using the Claasen method, was subsequently shown to be a mistaken identification.[13]

Krypton compounds with other than Kr-F bonds (compounds with atoms other than fluorine) have also been described. KrF
2 reacts with B(OTeF
5)
3 to produces the unstable compound, Kr(OTeF
5)
2, with a krypton-oxygen bond. A krypton-nitrogen bond is found in the cation [HC≡N–Kr–F]+
, produced by the reaction of KrF
2 with [HC≡NH]+
[AsF−
6] below −50 °C.[14]

Argon compounds

The discovery of HArF was announced in 2000.[15][16] The compound can exist in low temperature argon matrices for experimental studies, and has also been studied computationally.[16] Argon hydride ion ArH+
 was obtained in the 1970s.[17] This molecular ion has also been identified in the Crab nebula, based on the frequency of its light emissions.[18]

Neon and helium compounds

A neon compound has yet to be described or identified. There is some empirical and theoretical evidence for a few metastable helium compounds which may exist at very low temperatures or extreme pressures. The stable cation HeH+ was reported in 1925.[19]

Reports prior to xenon hexafluoroplatinate and xenon tetrafluoride

Clathrates

Kr(H2)4 and H2 solids formed in a diamond anvil cell. Ruby was added for pressure measurement.[20]
Structure of Kr(H2)4. Krypton octahedra (green) are surrounded by randomly oriented hydrogen molecules.[20]

Prior to 1962, the only isolated compounds of noble gases were clathrates (including clathrate hydrates); other compounds such as coordination compounds were observed only by spectroscopic means.[2] Clathrates (also known as cage compounds) are compounds of noble gases in which they are trapped within cavities of crystal lattices of certain organic and inorganic substances. The essential condition for their formation is that the guest (noble gas) atoms should be of appropriate size to fit in the cavities of the host crystal lattice; for instance, Ar, Kr, and Xe can form clathrates with crystalline β-quinol, but He and Ne cannot fit because they are too small. As well, Kr and Xe can appear as guests in crystals of melanophlogite.

Helium-nitrogen (He(N2)11) crystals have been grown at room temperature at pressures ca. 10 GPa in a diamond anvil cell.[21] Solid argon-hydrogen clathrate (Ar(H2)2) has the same crystal structure as the MgZn2 Laves phase. It forms at pressures between 4.3 and 220 GPa, though Raman measurements suggest that the H2 molecules in Ar(H2)2 dissociate above 175 GPa. A similar Kr(H2)4 solid forms at pressures above 5 GPa. It has a face-centered cubic structure where krypton octahedra are surrounded by randomly oriented hydrogen molecules. Meanwhile, in solid Xe(H2)8 xenon atoms form dimers inside solid hydrogen.[20]

Coordination compounds

Coordination compounds such as Ar·BF3 have been postulated to exist at low temperatures, but have never been confirmed. Also, compounds such as WHe2 and HgHe2 were reported to have been formed by electron bombardment, but recent research has shown that these are probably the result of He being adsorbed on the surface of the metal; therefore, these compounds cannot truly be considered chemical compounds.

Hydrates

Hydrates are formed by compressing noble gases in water, where it is believed that the water molecule, a strong dipole, induces a weak dipole in the noble gas atoms, resulting in dipole-dipole interaction. Heavier atoms are more influenced than smaller ones, hence Xe•5.75 H2O was reported to have been the most stable hydrate;[22] it has a melting point of 24 °C.[23] The deuterated version of this hydrate has also been produced.[24]

Fullerene adducts

Noble gases can also form endohedral fullerene compounds where the noble gas atom is trapped inside a fullerene molecule. In 1993, it was discovered that when C60 is exposed to a pressure of around 3 bar of He or Ne, the complexes He@C60 and Ne@C60 are formed.[25] Under these conditions, only about one out of every 650,000 C60 cages was doped with a helium atom; with higher pressures (3000 bar), it is possible to achieve a yield of up to 0.1%. Endohedral complexes with argon, krypton and xenon have also been obtained, as well as numerous adducts of He@C60.[26]

Applications

Most applications of noble gas compounds are either as oxidising agents or as a means to store noble gases in a dense form. Xenic acid is a valuable oxidising agent because it has no potential for introducing impurities—xenon is simply liberated as a gas—and so is rivalled only by ozone in this regard.[2] The perxenates are even more powerful oxidizing agents. Xenon-based oxidants have also been used for synthesizing carbocations stable at room temperature, in SO
2ClF solution.[27]

Stable salts of xenon containing very high proportions of fluorine by weight (such as tetrafluoroammonium heptafluoroxenate, NF4XeF7, and the related tetrafluoroammonium octafluoroxenate (NF4)2XeF8), have been developed as highly energetic oxidisers for use as propellants in rocketry.[28] [29]

The xenon fluorides are good fluorinating agents.

Clathrates have been used for separation of He and Ne from Ar, Kr, and Xe, and also for the transportation of Ar, Kr, and Xe. (For instance, radioactive isotopes of krypton and xenon are difficult to store and dispose, and compounds of these elements may be more easily handled than the gaseous forms.[2]) In addition, clathrates of radioisotopes may provide suitable formulations for experiments requiring sources of particular types of radiation; hence. 85Kr clathrate provides a safe source of beta particles, while 133Xe clathrate provides a useful source of gamma rays.

References

  1. ↑ Pauling, Linus (June 1933). "The Formulas of Antimonic Acid and the Antimonates". J. Am. Chem. Soc. 55 (5): 1895–1900. doi:10.1021/ja01332a016.
  2. 1 2 3 4 5 Holloway, John H. (1968). Noble-Gas Chemistry. London: Methuen. ISBN 0-416-03270-2.
  3. ↑ Seppelt, Konrad (June 1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research 12 (6): 211–216. doi:10.1021/ar50138a004.
  4. ↑ Smith GL, Mercier HP, Schrobilgen GJ (February 2007). "Synthesis of [F3S≡NXeF][AsF6] and structural study by multi-NMR and Raman spectroscopy, electronic structure calculations, and X-ray crystallography". Inorganic Chemistry 46 (4): 1369–78. doi:10.1021/ic061899. PMID 17256847.
  5. ↑ Smith GL, Mercier HP, Schrobilgen GJ (May 2008). "F5SN(H)Xe+; a rare example of xenon bonded to sp3-hybridized nitrogen; synthesis and structural characterization of [F5SN(H)Xe][AsF6]". Inorganic Chemistry 47 (10): 4173–84. doi:10.1021/ic702039f. PMID 18407626.
  6. ↑ Kenneth S. Pitzer (1975). "Fluorides of radon and element 118". J. Chem. Soc., Chem. Commun. (18): 760b – 761. doi:10.1039/C3975000760b.
  7. ↑ Bartlett, N. (1962). "Xenon hexafluoroplatinate Xe+[PtF6]−". Proceedings of the Chemical Society of London (6): 218. doi:10.1039/PS9620000197.
  8. ↑ Graham, L.; Graudejus, O., Jha N.K., and Bartlett, N. (2000). "Concerning the nature of XePtF6". Coordination Chemistry Reviews 197: 321–334. doi:10.1016/S0010-8545(99)00190-3. Cite uses deprecated parameter |coauthors= (help)
  9. ↑ Claassen, H. H.; Selig, H.; Malm, J. G. (1962). "Xenon Tetrafluoride". J. Am. Chem. Soc. 84 (18): 3593. doi:10.1021/ja00877a042.
  10. ↑ Hoppe, R. ; Daehne, W. ; Mattauch, H. ; Roedder, K. (1962-11-01). "FLUORINATION OF XENON". Angew. Chem. Int. Ed. Engl. 1 (11): 599. doi:10.1002/anie.196205992.
  11. ↑ Li, Wai-Kee; Zhou,, Gong-Du; Mak, Thomas C. W. (2008). Advanced Structural Inorganic Chemistry. Oxford University Press. p. 674. ISBN 0-19-921694-0.
  12. 1 2 Lehmann, J (2002). "The chemistry of krypton". Coordination Chemistry Reviews. 233-234: 1. doi:10.1016/S0010-8545(02)00202-3.
  13. ↑ Prusakov, V. N.; Sokolov, V. B. (1971). "Krypton difluoride". Soviet Atomic Energy 31 (3): 990–999. doi:10.1007/BF01375764.
  14. ↑ John H. Holloway; Eric G. Hope (1998). A. G. Sykes, ed. Advances in Inorganic Chemistry. Academic Press. p. 57. ISBN 0-12-023646-X.
  15. ↑ Khriachtchev, L., Pettersson, M., Runeberg, N., Lundell, J., Räsänen, M. (2000). "A stable argon compound". Nature 406 (6798): 874–876. doi:10.1038/35022551. PMID 10972285.
  16. 1 2 Bochenkova, Anastasia V.; Bochenkov, Vladimir E.; Khriachtchev, Leonid (2 July 2009). "HArF in Solid Argon Revisited: Transition from Unstable to Stable Configuration". The Journal of Physical Chemistry A 113 (26): 7654–7659. doi:10.1021/jp810457h.
  17. ↑ Wyatt, J. R.; Strattan, L. W.; Snyder, S. C.; Hierl, P. M. (1975). "Chemical accelerator studies of reaction dynamics: Ar+
     + CH
    4     → ArH+
     + CH
    3    ". The Journal of Chemical Physics 62 (7): 2555. doi:10.1063/1.430836.
  18. ↑ Barlow, M. J.; Swinyard, B. M.; Owen, P. J.; Cernicharo, J.; Gomez, H. L.; Ivison, R. J.; Krause, O.; Lim, T. L.; Matsuura, M.; Miller, S.; Olofsson, G.; Polehampton, E. T. (12 December 2013). "Detection of a Noble Gas Molecular Ion, 36
    ArH+
    , in the Crab Nebula". Science 342 (6164): 1343–1345. doi:10.1126/science.1243582.
  19. ↑ Hogness, T.R.; Lunn, E.G. (1925). "The Ionization of Hydrogen by Electron Impact as Interpreted by Positive Ray Analysis". Phys. Rev. Lett. (The American Physical Society) 26 (1): 44–55. doi:10.1103/PhysRev.26.44. Retrieved 15 December 2013.
  20. 1 2 3 Kleppe, Annette K.; Amboage, Mónica; Jephcoat, Andrew P. (2014). "New high-pressure van der Waals compound Kr(H2)4 discovered in the krypton-hydrogen binary system". Scientific Reports 4. doi:10.1038/srep04989.
  21. ↑ Vos, W. L.; Finger, L. W.; Hemley, R. J.; Hu, J. Z.; Mao, H. K.; Schouten, J. A. (1992). "A high-pressure van der Waals compound in solid nitrogen-helium mixtures". Nature 358 (6381): 46. doi:10.1038/358046a0.
  22. ↑ Pauling, L. (1961). "A molecular theory of general anesthesia". Science 134 (3471): 15–21. Bibcode:1961Sci...134...15P. doi:10.1126/science.134.3471.15. PMID 13733483. Reprinted as Pauling, Linus; Kamb, Barclay, eds. (2001). Linus Pauling: Selected Scientific Papers 2. River Edge, New Jersey: World Scientific. pp. 1328–1334. ISBN 981-02-2940-2.
  23. ↑ Henderson, W. (2000). Main group chemistry. Great Britain: Royal Society of Chemistry. p. 148. ISBN 0-85404-617-8.
  24. ↑ Ikeda, Tomoko; Mae, Shinji; Yamamuro, Osamu; Matsuo, Takasuke; Ikeda, Susumu; Ibberson, Richard M. (November 23, 2000). "Distortion of Host Lattice in Clathrate Hydrate as a Function of Guest Molecule and Temperature". Journal of Physical Chemistry A 104 (46): 10623–10630. doi:10.1021/jp001313j.
  25. ↑ Saunders, M.; Jiménez-Vázquez, H. A.; Cross, R. J. and Poreda, R. J. (1993). "Stable compounds of helium and neon. He@C60 and Ne@C60". Science 259 (5100): 1428–1430. Bibcode:1993Sci...259.1428S. doi:10.1126/science.259.5100.1428. PMID 17801275.
  26. ↑ Saunders, Martin; Jimenez-Vazquez, Hugo A.; Cross, R. James; Mroczkowski, Stanley; Gross, Michael L.; Giblin, Daryl E. and Poreda, Robert J. (1994). "Incorporation of helium, neon, argon, krypton, and xenon into fullerenes using high pressure". J. Am. Chem. Soc. 116 (5): 2193–2194. doi:10.1021/ja00084a089.
  27. ↑ Mercier, H. P. A.; Moran, M. D.; Schrobilgen, G. J.; Steinberg, C.; Suontamo, R. J. (2004). "The Syntheses of Carbocations by Use of the Noble-Gas Oxidant, [XeOTeF
    5    ][Sb(OTeF
    5    )
    6    ]    : The Syntheses and Characterization of the CX+
    3     (X = Cl, Br, OTeF
    5    ) and CBr(OTeF
    5    )+
    2     Cations and Theoretical Studies of CX+
    3     and BX
    3     (X = F, Cl, Br, I, OTeF
    5    )". J. Am. Chem. Soc. 126 (17): 5533–5548. doi:10.1021/ja030649e. PMID 15113225.
  28. ↑ Christe, KO; Wilson, WW (Dec 1982). "Perfluoroammonium and alkali-metal salts of the heptafluoroxenate(VI) and octafluoroxenate(VI) anions". Inorganic Chemistry 21 (12): 4113–4117. doi:10.1021/ic00142a001.
  29. ↑ Karl O. Christe, William W. Wilson. Perfluoroammonium salt of heptafluoroxenon anion. U.S. Patent 4,428,913, June 24, 1982

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