|Molar mass||108.999 g/mol|
|Molecular shape||trigonal bipyramidal|
|Dipole moment||0 D|
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Nitrogen pentafluoride is a theoretical compound of nitrogen and fluorine that is hypothesized to exist based on the existence of the pentafluorides of the atoms below nitrogen in the periodic table, such as phosphorus pentafluoride. Theoretical models of the nitrogen pentafluoride (NF5) molecule are either a trigonal bipyramidal covalently bound molecule with symmetry group D3h, or NF4+F−, which would be an ionic solid. Also, a related compound NH4+F− is known.
In 1966 W. E. Tolberg first synthesized a five valent nitrogen compound of nitrogen and fluorine when tetrafluoroammonium compounds, tetrafluoroammonium heaxafluoroantimonate NF4SbF6 and tetrafluoroammonium heaxafluoroarsenate NF4AsF6 were made. In 1971 C. T. Goetschel announced the preparation of NF4BF4 and also produced a white solid assumed to be tetrafluoroammonium fluoride (NF4+F−). This was made by treating nitrogen trifluoride and fluorine with 3MeV electron radiation at 77K. It decomposed above 143K back into those ingredients. Theoretical studies also show the ionic compound is very likely to decompose to nitrogen trifluoride and fluorine gas.
Karl O. Christe synthesised bis(tetrafluoroammonium) hexafluoronickelate (NF4)2NiF6. He also prepared compounds with manganese, a fluorouranate, a perchlorate, a fluorosulfate and N2F3+ salts. Christe attempted to make NF4F by metathesis of NF4SbF6 with CsF in HF solvent at 20 °C. However, a variant, tetrafluoroammonium bifluoride (NF4HF2·nHF), was produced. At room temperature it was a milky liquid, but when cooled turned pasty. At −45 °C it had the form of a white solid. When reheated it frothed, giving off F2, HF and NF3 as gases. This has CAS number 71485-49-9.
I. J. Solomon believed that nitrogen pentafluoride was produced by the thermal decomposition of NF4AsF6, but experimental results were not reproduced.
For a NF5 molecule to form, five fluorine atoms have to be arranged around a nitrogen atom. There is insufficient space to do this in the most compact way, so that bond lengths are forced to be longer. Calculations show that the NF5 molecule is thermodynamically favourably inclined to form NF4 and F radicals with energy 8.5 kcal mol−1 and a transition barrier around 20±4 kcal mol−1.
- Lewars, Errol G. (03/11/2008). Modeling marvels: computational anticipation of novel molecules. Springer. pp. 53–67. Check date values in:
- Goetschel, C. T.; V. A. Campanile; R. M. Curtis; K. R. Loos; C. D. Wagner; J. N. Wilson (July 1972). Inorganic Chemistry 11 (7): 1696–1701. doi:10.1021/ic50113a051. Missing or empty
- Christe, Karl O; William W. Wilson (December 1992). "Nitrogen pentafluoride: covalent NF5 versus ionic NF4+F− and studies on the instability of the latter". Journal of the American Chemical Society 114 (25): 9934–9936. doi:10.1021/ja00051a027.
- Christe, Karl O. (23 May 1980). "Research Studies in NF4+ Salts". Rockwell. Retrieved 23 February 2012.
- Christe, Karl O. (23 May 1980). "Research Studies in NF4+ Salts". Rockwell. pp. G7–G8. Retrieved 23 February 2012.
- Tetrafluoroammonium bifluoride
- Christe, Karl O.; William W. Wilson; Gary J. Schrobilgen; Raman V. Chirakal; George A. Olah (March 1998). "On the existence of pentacoordinated nitrogen". Inorganic Chemistry 27 (5): 789–790. doi:10.1021/ic00278a009.
- Holger F. Bettinger, Paul v. R. Schleyer, and Henry F. Schaefer III (27 October 1998). "NF5 — Viable or Not?". Journal of the American Chemical Society 120 (44): 11439–11448. doi:10.1021/ja9813921.