Tetrafluoroammonium

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2D model of the tetrafluoroammonium ion

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 It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine.[1] Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF
4
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 (XF
5
where X is Ge, Sn, or Ti), hexafluorides (XF
6
where X is P, As, Sb, Bi, or Pt), heptafluorides (XF
7
where X is W, U, or Xe), octafluorides (XeF2−
8
),[2] various oxyfluorides (XF
5
O
where X is W or U; FSO
3
, BrF
4
O
), and perchlorate (ClO
4
).[3] Attempts to make the nitrate salt, NF
4
NO
3
, were unsuccessful because of quick fluorination: NF+
4
+ NO
3
NF
3
+ FONO
2
.[4]

Structure[edit]

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[edit]

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]

NF
3
+ F
2
+ SbF
5
NF
4
SbF
6

The hexafluoroarsenate salt was also prepared by a similar reaction with arsenic pentafluoride at 120 °C:[5]

NF
3
+ F
2
+ AsF
5
NF
4
AsF
6

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
4
BF
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
]
2
TiF2−
6
.[7]

Many tetrafluoroammonium salts can be prepared with metathesis reactions.

Reactions[edit]

Tetrafluoroammonium salts are extremely hygroscopic. The NF+
4
ion is readily hydrolysed into nitrogen trifluoride, H
2
F+
, and oxygen gas:

2 NF+
4
+ 2 H
2
O
→ 2 NF
3
+ 2 H
2
F+
+ O
2

Some hydrogen peroxide (H
2
O
2
) is also formed during this process.[5]

Reaction of NF+
4
SbF
6
with alkali metal nitrates yields fluorine nitrate, FONO
2
.[4]

Properties[edit]

Because tetrafluoroammonium salts are destroyed by water, it cannot be used as a solvent. Instead anhydrous hydrogen fluoride or bromine pentafluoride can be used as a to dissolve these salts.[8]

Tetrafluoroammonium salts usually have no colour. However some are coloured due to other metals in them. Red salts include (NF+
4
)
2
CrF2−
6
, (NF+
4
)
2
NiF2−
6
and (NF+
4
)
2
PtF2−
6
. (NF+
4
)
2
MnF2−
6
, NF+
4
UF
7
, NF+
4
UOF
5
and NF+
4
XeF
7
are yellow.[8]

Applications[edit]

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]

See also[edit]

References[edit]

  1. ^ Nikitin, I. V.; Rosolovskii, V. Y. (1985). "Tetrafluoroammonium Salts". Russian Chemical Reviews 54 (5): 426. Bibcode:1985RuCRv..54..426N. doi:10.1070/RC1985v054n05ABEH003068. 
  2. ^ 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. 
  3. ^ 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. ^ a b Hoge, B.; Christe, K. O. (2001). "On the stability of NF+
    4
    NO
    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. ^ a b c d e Sykes, A. G. (1989). Advances in Inorganic Chemistry. Academic Press. ISBN 0-12-023633-8. 
  6. ^ Patnaik, Pradyot (2002). Handbook of inorganic chemicals. McGraw-Hill Professional. ISBN 0-07-049439-8. 
  7. ^ 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. 
  8. ^ a b Sykes, A. G. (1989-07-17). Advances in Inorganic Chemistry. Academic Press. p. 154. ISBN 9780080578828. Retrieved 22 June 2014.