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 (MF 5O− where M is W or U; FSO− 3, BrF 4O− ), and perchlorate (ClO− 4).[3] Attempts to make the nitrate salt, NF 4NO 3, were unsuccessful because of quick fluorination: NF+ 4 + NO− 3 → NF 3 + FONO 2.[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]
The reaction of nitrogen trifluoride with fluorine and boron trifluoride at 800 °C yields the tetrafluoroborate salt:[6]
NF 3 + F 2 + BF 3 → NF 4BF 4
NF+ 4 salts can also be prepared by fluorination of NF 3 with krypton difluoride (KrF 2) and fluorides of the form MF n, where M is Sb, Nb, Pt, Ti, or B. For example, reaction of NF 3 with KrF 2 and TiF 4 yields [NF+ 4] 2TiF2− 6.[7]
Tetrafluoroammonium salts are extremely hygroscopic. The NF+ 4 ion, when dissolved in water, readily decomposes into NF 3, H 2F+ , and oxygen gas. Some hydrogen peroxide (H 2O 2) is also formed during this process:[5]
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 NF 3–F 2 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]
^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.
^ abHoge, 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.
^ abcdeSykes, A. G. (1989). Advances in Inorganic Chemistry. Academic Press. ISBN0-12-023633-8.
^Patnaik, Pradyot (2002). Handbook of inorganic chemicals. McGraw-Hill Professional. ISBN0-07-049439-8.
^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. ISSN0002-7863.
