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Talk:Fluorine/Comparison between the highest oxidation states of oxides and fluorides

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Comparison between the highest oxidation states of oxides and fluorides

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For groups 1—6 and 13—16 the highest oxidation states of oxides and fluorides are always equal, and differences are only seen in groups 7—11, mercury, halogens, and the noble gases. The general trend is fluorination allows to achieve relatively low[note 1] but hardly achievable oxidation states; for example, no binary oxide is known for krypton, but krypton difluoride is well-studied.[1] However, very high oxidation states of several elements are known for oxygen only; for example, none has shown that existence of ruthenium octafluoride is possible yet, while ruthenium tetroxide is well-studied.[2]

Later transition metals
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central gray sphere with 4 others attached by spokes in a plane, 90 degrees apart
Square planar structure of mercury(IV) fluoride

With the exceptions of the +7 and +8 oxidation states, fluorine is the key in achieving many rare high oxidation states of the transition metals. For instance, direct reaction of the respective metals with fluorine gives rise to palladium(VI)[3] and platinum(VI).[4] The only occurrence of mercury(IV)[5] is binary mercury(IV) fluoride, synthesized at temperatures close to absolute zero. Gold(V)[6] is only known in the hexafluoroaurate(V) ion, which can be synthesized indirectly under extreme conditions, and the gold(V) fluoride, which is obtained during hexafluoroaurate(V) decomposition. The high oxidizing potential of fluorine has led to the claim of the gold(VII) existence in gold heptafluoride,[7] but current calculations show that the claimed AuF7 molecule was AuF5·F2.[8] The great oxidizing power of fluorine is also illustrated by the fluorine-containing complexes of copper(IV),[9] silver(IV),[10] nickel(IV),[11] iridium(VI),[12] and others. It is possible that the element 113, ununtrium, will be the first element in boron group to form a species in the +5 oxidation state, the fluorine-based hexafluoroununtrate(V), UutF−
6;[13] the possibility of a +5 oxygen-based species is not known to be calculated.

Fluorine is a generally stronger oxidizer than oxygen; however, this strength does not apply for every case. Dinitrogen pentoxide, with nitrogen in the oxidation state of +5, is known; but creating nitrogen pentafluoride would need to squeeze five fluorine atoms attached to the central atom. This is hard to perform, as a nitrogen atom is smaller than most other atoms. It is not known whether the molecule is possible to produce or not, and if possible, whether it is stable or not.[14] Similarly, the highest oxidation states of several late transition metals may be achieved in oxides only: for example, even though only gold(V) is known now, and only in the form of a fluoride, calculations provided by Dementyev et al. in 1997 show that the element may be oxidized up to gold(IX), in the form of the tetroxoaurlyl(IX) ion, [AuO4]+, but not as a fluorine-based compound or ion.[15] Similarly, this and another (Rother et al., 1969) calculations revealed that oxygen-based complexes that contain iridium(IX),[16] platinum(X),[15] and mercury(VIII)[15] might be possible. However, these species were denied by the University of Würzburg in a 2006 paper; it expects platinum(VI), gold(V), and mercury(IV), known in binary fluorides, to be the highest for the elements.[17] It has been shown osmium and iridium may form heptafluorides;[17] for osmium, even an octafluoride may be possible.[17]

Halogens and noble gases
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Among halogens, chlorine and bromine form perchlorates[18] and perbromates,[19] both oxygen-based and with the representative halogen in +7 state; however, chlorine, unlike bromine, forms a binary heptoxide.[20] Out of their stable fluorinated species, pentafluorides are the species in the highest oxidation state achieved; however, bromine hexafluoride, BrF6•, is known as well.[21] Iodine shows the reverse picture: no heptoxide is known, unlike heptafluoride, a well-known stable compound; however, periodic acid, containing iodine(VII), is known as well.[22]

Noble gases do not show a trend as well: as noted above, krypton has no known binary oxides, but has a well-studied difluoride. Xenon forms a tetroxide of oxygen-based species,[23] but only a hexafluoride of fluorine-based ones. Neutral xenon octafluoride is not known nor expected to be stable,[24] but octafluoroxenate(VI), XeF2−
8, has been synthesized.[25] Contradictory data is known about fluorides and especially oxides of radon; no binary fluoride or oxide of lighter noble gases are known.


Cite error: There are <ref group=note> tags on this page, but the references will not show without a {{reflist|group=note}} template (see the help page).

  1. ^ Cite error: The named reference bla was invoked but never defined (see the help page).
  2. ^ Singh, B.; Srivastava, S. (1991). "Kinetics and Mechanism of Ruthenium tetroxide Catalysed Oxidation of Cyclic Alcohols by Bromate in a Base". Transition Metal Chemistry. 16: 466–68. doi:10.1007/BF01129466.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  3. ^ Aullón, G.; Alvarez, S. (2007). "On the Existence of Molecular Palladium(VI) Compounds: Palladium Hexafluoride". Inorganic Chemistry. 46 (7): 2700–03. doi:10.1021/ic0623819. PMID 17326630.
  4. ^ Weinstock, B.; Claassen, H. H.; Malm, J. G. (1957). "Platinum Hexafluoride". Journal of the American Chemical Society. 79 (21): 5822–32. doi:10.1021/ja01578a073.
  5. ^ Wang, Xuefang; Andrews, Lester; Riedel, Sebastian; Kaupp, Martin (2007). "Mercury is a Transition Metal: The First Experimental Evidence for HgF4". Angewandte Chemie. 119 (44): 8523–27. doi:10.1002/ange.200703710.
  6. ^ Emeléus, H. J.; Sharpe, A. G. (1983). Advances in Inorganic Chemistry and Radiochemistry. Academic Press. p. 83. ISBN 0120236273.
  7. ^ (in Russian) Timakov, A. A.; Prusakov, V. N.; Drobyshevskii, Y. V. (1986). Doklad akademii nauk SSSR. USSR Academy of Sciences. p. 125. {{cite book}}: Unknown parameter |unused_data= ignored (help)
  8. ^ Riedel, S.; Kaupp, M. (2006). "Has AuF7 Been Made?". Inorganic Chemistry. 45 (3): 1228–34. doi:10.1021/ic051944y. PMID 16441134.
  9. ^ Popova, T. V.; Aksenova, N. V. (2003). "Complexes of Copper in Unstable Oxidation States". Russian Journal of Coordinated Chemistry. 29 (11): 743–65. doi:10.1023/B:RUCO.0000003432.39025.cc.
  10. ^ (in Russian)Popov, N. V.; Kiselev, Y. M.; Sukhoverkhov, V. F.; Spitsyn, V. I. (1987). Doklad akademii nauk SSSR. USSR Academy of Sciences. p. 628.
  11. ^ Taylor, J. C.; Wilson, P. W. (1973). "The Structures of Fluorides—IV : A Neutron Diffraction Study of K2NiF6". Journal of Inorganic and Nuclear Chemistry. 36 (7): 1561–63. doi:10.1016/0022-1902(74)80623-8.
  12. ^ Drews, T; Supeł, J; Hagenbach, A; Seppelt, K. (2006). "Solid State Molecular Structures of Transition Metal Hexafluorides". Inorganic Chemistry. 45 (9): 3782–88. doi:10.1021/ic052029f. PMID 16634614. {{cite journal}}: Unknown parameter |month= ignored (help)CS1 maint: multiple names: authors list (link)
  13. ^ Haire, Richard G. (2006). "Transactinides and the future elements". In Morss; Edelstein, Norman M.; Fuger, Jean (eds.). The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Springer Science+Business Media. p. 1723. ISBN 1-4020-3555-1.{{cite book}}: CS1 maint: ref duplicates default (link)
  14. ^ Lewars 2008, pp. 63–64.
  15. ^ a b c (in Russian)Dementyev, A. I.; Kuznetsov, M. L.; Kiselev, Y. M. (1997). "On extremal oxidation states of heavy 5d elements". Zhurnal Neorganicheskoy Khimiyi (42): 1167.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  16. ^ (in German)Rother, P.; Wagner, F.; Zahn, U. (1969). "Chemical consequences of the 193Os(β-)193Ir decay in osmium". Radiochimica Acta (11): 203.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  17. ^ a b c Bayerische Julius-Maximilians-Universität Würzburg (2006). The Highest Oxidation States of the 5d Transition Metals: a Quantum-Chemical Study (Report). Chemical Society. pp. 95–105. {{cite report}}: |access-date= requires |url= (help)
  18. ^ Wolff, J. (1998). "Perchlorate and the Thyroid Gland". Pharmacological Reviews. 50 (1): 89–105. PMID 9549759.
  19. ^ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 439. ISBN 0123526515.
  20. ^ Byrns, A. C.; Rollefson, G. K. (1934). "The formation of chlorine heptoxideon illuminations of mixtures of chlorine and ozone". Journal of the American Chemical Society. 56 (5): 1250–51. doi:10.1021/ja01320a506.
  21. ^ Emeléus, Harry Julius; Sharpe, A. G. (1978). Advances in inorganic chemistry and radiochemistry. Vol. 21. Academic Press. p. 248. ISBN 0120236214.{{cite book}}: CS1 maint: multiple names: authors list (link)
  22. ^ Alan Isaacs, John Daintith, Elizabeth Martin, ed. (1984). Concise Science Dictionary. Oxford University Press. p. 356. ISBN 0192115936.{{cite book}}: CS1 maint: multiple names: editors list (link)
  23. ^ Gundersen, G.; Hedberg, K.; Huston, J. L. (1970). "Molecular Structure of Xenon Tetroxide, XeO4". J. Chem. Phys. 52: 812–15. Bibcode:1970JChPh..52..812G. doi:10.1063/1.1673060.
  24. ^ Seppelt, Konrad (1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research. 12 (6): 211–16. doi:10.1021/ar50138a004.
  25. ^ Chandra, Sulekh (2004). Comprehensive Inorganic Chemistry. New Age International. p. 308. ISBN 8122415121.
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