Nitrogen monofluoride

Today we are going to explore Nitrogen monofluoride, a topic that has captured the attention of people of all ages and walks of life. Since its emergence, Nitrogen monofluoride has generated great interest due to its impact on our society and our daily lives. In this article, we are going to dive into the history of Nitrogen monofluoride, explore its implications in today's world, and reflect on its future. Whether you are an expert on the topic or just curious to learn more about it, this article will provide you with a complete and insightful overview of Nitrogen monofluoride. Join us on this fascinating journey!

Nitrogen monofluoride
Names
Other names
Fluoroimidogen
Identifiers
3D model (JSmol)
  • InChI=1S/FN/c1-2
    Key: CMUBZTZNXGBJMQ-UHFFFAOYSA-N
  • =
Properties
FN
Molar mass 33.005 g·mol−1
Related compounds
Dioxygen, nitroxyl anion
Except where otherwise noted, data are given for materials in their standard state (at 25 °C , 100 kPa).

Nitrogen monofluoride (fluoroimidogen) is a metastable species that has been observed in laser studies. It is isoelectronic with O2. Like boron monofluoride, it is an instance of the rare multiply-bonded fluorine atom.[1][2] It is unstable with respect to its formal dimer, dinitrogen difluoride, as well as to its elements, nitrogen and fluorine.

Nitrogen monofluoride is produced when radical species (H, O, N, CH3) abstracts a fluorine atom from nitrogen difluoride (NF2). Stoichiometrically, the reaction is extremely efficient, regenerating a radical for long-lasting chain propagation. However, radical impurities in the end product also catalyze that product's decomposition. Azide decomposition offers a less-efficient but more pure technique: fluorine azide (which can be formed in situ via reaction of atomic fluorine with hydrazoic acid) decomposes upon shock into NF and N2.[3][4]

Many NF-producing reactions give the product in an excited state with characteristic chemiluminescence at 870 and 875 nm (infrared), or at 525–530 nm (green). They have thus been investigated for development as a chemical laser.[4][5]

References

  1. ^ Davis, Steven J.; Rawlins, Wilson T.; Piper, Lawrence G. (Feb 1989). "Rate coefficient for the H + NF(a1Δ) reaction" (PDF). The Journal of Physical Chemistry. 93 (3). American Chemical Society: 1078–1082. doi:10.1021/j100340a013. ISSN 0022-3654 – via MetastableStates.com.
  2. ^ Harbison, G. S. (2002). "The Electric Dipole Polarity of the Ground and Low-lying Metastable Excited States of NF". Journal of the American Chemical Society. 124 (3): 366–367. doi:10.1021/ja0159261. PMID 11792193.
  3. ^ Gmelin-lnstitut für Anorganische Chemie der Max-Planck-Gesellschaft zur Förderung der Wissenschaften (2013). Gmelin Handbook of Inorganic Chemistry: F Fluorine: Compounds with Oxygen and Nitrogen. Springer Science & Business Media. pp. 263–271. ISBN 9783662063392.
  4. ^ a b Avizonis, Petras V. (2012). "Chemically Pumped Electronic Transition Lasers". In Onorato, Michele (ed.). Gas Flow and Chemical Lasers. Plenum Press. pp. 1–19. doi:10.1007/978-1-4615-7067-7_1. ISBN 978-1-4615-7067-7.
  5. ^ Kenner, Rex D.; Ogryzlo, Elmer A. (1985). "Chemiluminescence in Gas Phase Reactions; 4. NF(a1Δ) (870, 875 nm) and (b1Σ+) (525–530 nm)". In Burr, John G. (ed.). Chemi- and Bioluminescence. Chemical and Biochemical Analysis. Vol. 16. Dekker. pp. 84–87. ISBN 0-8247-7277-6.