Chemistry:Tetrafluoroammonium
 2D model of the tetrafluoroammonium ion
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| IUPAC name
Tetrafluoroammonium
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3D model (JSmol)
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| ChemSpider | |
| 2028 | |
PubChem CID
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| Properties | |
| F4N+ | |
| Molar mass | 90.000 g·mol−1 |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
The tetrafluoroammonium cation (also known as perfluoroammonium) is a positively charged polyatomic ion with chemical formula NF+4. It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine.[1] Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF4, trifluoramine oxide ONF3 and the tetrafluoroborate BF−4 anion.
The tetrafluoroammonium ion forms salts with a large variety of fluorine-bearing anions. These include the bifluoride anion (HF−2), tetrafluorobromate (BrF−4), metal pentafluorides (MF−5 where M is Ge, Sn, or Ti), hexafluorides (MF−6 where M is P, As, Sb, Bi, or Pt), heptafluorides (MF−7 where M is W, U, or Xe), octafluorides (XeF2−8),[2] various oxyfluorides (MF5O− where M is W or U; FSO−3, BrF4O−), and perchlorate (ClO−4).[3] Attempts to make the nitrate salt, NF4NO3, were unsuccessful because of quick fluorination: NF+4 + NO−3 → NF3 + FONO2.[4]
Structure
The geometry of the tetrafluoroammonium ion is tetrahedral, with an estimated nitrogen-fluorine bond length of 124 pm. All fluorine atoms are in equivalent positions.[5]
Synthesis
Tetrafluoroammonium salts are prepared by oxidising nitrogen trifluoride with fluorine in the presence of a strong Lewis acid which acts as a fluoride ion acceptor. The original synthesis by Tolberg, Rewick, Stringham, and Hill in 1966 employs antimony pentafluoride as the Lewis acid:[5]
- NF3 + F2 + SbF5 → NF4SbF6
The hexafluoroarsenate salt was also prepared by a similar reaction with arsenic pentafluoride at 120 °C:[5]
- NF3 + F2 + AsF5 → NF4AsF6
The reaction of nitrogen trifluoride with fluorine and boron trifluoride at 800 °C yields the tetrafluoroborate salt:[6]
- NF3 + F2 + BF3 → NF4BF4
NF+4 salts can also be prepared by fluorination of NF3 with krypton difluoride (KrF2) and fluorides of the form MFn, where M is Sb, Nb, Pt, Ti, or B. For example, reaction of NF3 with KrF2 and TiF4 yields [NF+4]2TiF2−6.[7]
Many tetrafluoroammonium salts can be prepared with metathesis reactions.
Reactions
Tetrafluoroammonium salts are extremely hygroscopic. The NF+4 ion, when dissolved in water, readily decomposes into NF3, H2F+, and oxygen gas. Some hydrogen peroxide (H2O2) is also formed during this process:[5]
- NF+4 + H2O → NF3 + H2F+ + 1⁄2 O2
- NF+4 + 2 H2O → NF3 + H2F+ + H2O2
Reaction of NF+4SbF−6 with alkali metal nitrates yields fluorine nitrate, FONO2.[4]
Properties
Because NF+4 salts are destroyed by water, water cannot be used as a solvent. Instead, bromine trifluoride, bromine pentafluoride, iodine pentafluoride, or anhydrous hydrogen fluoride can be used.[8]
Tetrafluoroammonium salts usually have no colour. However, some are coloured due to other elements in them. (NF+4)2CrF2−6, (NF+4)2NiF2−6 and (NF+4)2PtF2−6 have a red colour, while (NF+4)2MnF2−6, NF+4UF−7, NF+4UOF−5 and NF+4XeF−7 are yellow.[8]
Applications
NF+4 salts are important for solid propellant NF3–F2 gas generators. They are also used as reagents for electrophilic fluorination of aromatic compounds in organic chemistry.[5] As fluorinating agents, they are also strong enough to react with methane.[9]
See also
References
- ↑ Nikitin, I. V.; Rosolovskii, V. Y. (1985). "Tetrafluoroammonium Salts". Russian Chemical Reviews 54 (5): 426. doi:10.1070/RC1985v054n05ABEH003068. Bibcode: 1985RuCRv..54..426N.
- ↑ Christe, K. O.; Wilson, W. W. (1982). "Perfluoroammonium and alkali-metal salts of the heptafluoroxenon(VI) and octafluoroxenon(VI) anions". Inorganic Chemistry 21 (12): 4113–4117. doi:10.1021/ic00142a001.
- ↑ Christe, K. O.; Wilson, W. W. (1986). "Synthesis and characterization of tetrafluoroammonium(1+) tetrafluorobromate(1-) and tetrafluoroammonium(1+) tetrafluorooxobromate(1-)". Inorganic Chemistry 25 (11): 1904–1906. doi:10.1021/ic00231a038.
- ↑ 4.0 4.1 Hoge, B.; Christe, K. O. (2001). "On the stability of NF+4NO−3 and a new synthesis of fluorine nitrate". Journal of Fluorine Chemistry 110 (2): 87–88. doi:10.1016/S0022-1139(01)00415-8.
- ↑ 5.0 5.1 5.2 5.3 5.4 Sykes, A. G. (1989). Advances in Inorganic Chemistry. Academic Press. ISBN 0-12-023633-8.
- ↑ Patnaik, Pradyot (2002). Handbook of inorganic chemicals. McGraw-Hill Professional. ISBN 0-07-049439-8.
- ↑ 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.
- ↑ 8.0 8.1 Sykes, A. G. (1989-07-17). Advances in Inorganic Chemistry. Academic Press. p. 154. ISBN 9780080578828. https://books.google.com/books?id=qzN5pnPwuaoC&pg=PA154. Retrieved 22 June 2014.
- ↑ Olah, George A.; Hartz, Nikolai; Rasul, Golam; Wang, Qi; Prakash, G. K. Surya; Casanova, Joseph; Christe, Karl O. (1994-06-01). "Electrophilic Fluorination of Methane with "F+" Equivalent N2F+ and NF4+ Salts". Journal of the American Chemical Society 116 (13): 5671–5673. doi:10.1021/ja00092a018. ISSN 0002-7863.
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