Nitrogen trifluoride can be prepared from the elements in the presence of an electric discharge.[10] In 1903, Otto Ruff prepared nitrogen trifluoride by the electrolysis of a molten mixture of ammonium fluoride and hydrogen fluoride.[11] It is far less reactive than the other nitrogen trihalides nitrogen trichloride, nitrogen tribromide, and nitrogen triiodide, all of which are explosive. Alone among the nitrogen trihalides it has a negative enthalpy of formation. It is prepared in modern times both by direct reaction of ammonia and fluorine and by a variation of Ruff's method.[6] It is supplied in pressurized cylinders.
NF 3 is slightly soluble in water without undergoing chemical reaction. It is nonbasic with a low dipole moment of 0.2340 D. By contrast, ammonia is basic and highly polar (1.47 D).[12] This contrast reflects the differing electronegativities of H vs F.
Similar to dioxygen, NF3 is a potent yet sluggish oxidizer.[6] It oxidizes hydrogen chloride to chlorine:[citation needed]
2 NF3 + 6 HCl → 6 HF + N2 + 3 Cl2
However, it only attacks (explosively) organic compounds at high temperatures. Consequently it is compatible under standard conditions with several plastics, as well as steel and Monel.[6]
Above 200-300 °C, NF3 reacts with metals, carbon, and other reagents to give tetrafluorohydrazine:[13]
Nitrogen trifluoride is primarily used to remove silicon and silicon-compounds during the manufacturing of semiconductor devices such as LCD displays, some thin-film solar cells, and other microelectronics. In these applications NF 3 is initially broken down within a plasma. The resulting fluorine radicals are the active agents that attack polysilicon, silicon nitride and silicon oxide. They can be used as well to remove tungsten silicide, tungsten, and certain other metals. In addition to serving as an etchant in device fabrication, NF 3 is also widely used to clean PECVD chambers.
NF 3dissociates more readily within a low-pressure discharge in comparison to perfluorinated compounds (PFCs) and sulfur hexafluoride (SF 6). The greater abundance of negatively-charged free radicals thus generated can yield higher silicon removal rates, and provide other process benefits such as less residual contamination and a lower net charge stress on the device being fabricated. As a somewhat more thoroughly consumed etching and cleaning agent, NF3 has also been promoted as an environmentally preferable substitute for SF 6 or PFCs such as hexafluoroethane.[17]
The utilization efficiency of the chemicals applied in plasma processes varies widely between equipment and applications. A sizeable fraction of the reactants are wasted into the exhaust stream and can ultimately be emitted into Earth's atmosphere. Modern abatement systems can substantially decrease atmospheric emissions.[18]NF 3 has not been subject to significant use restrictions. The annual reporting of NF 3 production, consumption, and waste emissions by large manufacturers has been required in many industrialized countries as a response to the observed atmospheric growth and the international Kyoto Protocol.[19]
Highly toxic fluorine gas (F2, diatomic fluorine) is a climate neutral replacement for nitrogen trifluoride in some manufacturing applications. It requires more stringent handling and safety precautions, especially to protect manufacturing personnel.[20]
Growth in atmospheric concentration of NF3 since the 1990s is shown in right graph, along with a subset of similar man-made gases. Note the log scale.[21]
The GWP of NF 3 is second only to SF 6 in the group of Kyoto-recognised greenhouse gases, and NF 3 was included in that grouping with effect from 2013 and the commencement of the second commitment period of the Kyoto Protocol. It has an estimated atmospheric lifetime of 740 years,[7] although other work suggests a slightly shorter lifetime of 550 years (and a corresponding GWP of 16,800).[15]
Nitrogen trifluoride concentration at several latitudes since 2015.[22]NF3 measured by the Advanced Global Atmospheric Gases Experiment (AGAGE) in the lower atmosphere (troposphere) at stations around the world. Abundances are given as pollution free monthly mean mole fractions in parts-per-trillion.
Since 1992, when less than 100 tons were produced, production grew to an estimated 4000 tons in 2007 and is projected to increase significantly.[15] World production of NF3 is expected to reach 8000 tons a year by 2010. By far the world's largest producer of NF 3 is the US industrial gas and chemical company Air Products & Chemicals. An estimated 2% of produced NF 3 is released into the atmosphere.[23][24] Robson projected that the maximum atmospheric concentration is less than 0.16 parts per trillion (ppt) by volume, which will provide less than 0.001 Wm−2 of IR forcing.[25]
The mean global tropospheric concentration of NF3 has risen from about 0.02 ppt (parts per trillion, dry air mole fraction) in 1980, to 0.86 ppt in 2011, with a rate of increase of 0.095 ppt yr−1, or about 11% per year, and an interhemispheric gradient that is consistent with emissions occurring overwhelmingly in the Northern Hemisphere, as expected. This rise rate in 2011 corresponds to about 1200 metric tons/y NF3 emissions globally, or about 10% of the NF3 global production estimates. This is a significantly higher percentage than has been estimated by industry, and thus strengthens the case for inventorying NF3 production and for regulating its emissions.[26]
One study co-authored by industry representatives suggests that the contribution of the NF3 emissions to the overall greenhouse gas budget of thin-film Si-solar cell manufacturing is clear.[27]
The UNFCCC, within the context of the Kyoto Protocol, decided to include nitrogen trifluoride in the second Kyoto Protocol compliance period, which begins in 2012 and ends in either 2017 or 2020. Following suit, the WBCSD/WRI GHG Protocol is amending all of its standards (corporate, product and Scope 3) to also cover NF3.[28]
^ abcdePhilip B. Henderson, Andrew J. Woytek "Fluorine Compounds, Inorganic, Nitrogen" in Kirk‑Othmer Encyclopedia of Chemical Technology, 1994, John Wiley & Sons, NY. doi:10.1002/0471238961.1409201808051404.a01 Article Online Posting Date: December 4, 2000
^Ruff, John K. (1967). "Derivatives of Nitrogen Fluorides". Chemical Reviews. 67 (6): 665–680. doi:10.1021/cr60250a004.
^Parry, Robert W., and Thomas C. Bissot. "The Preparation and Properties of Phosphorus Trifluoride-Borane and Phosphorus Trifluoride-Borane-d31." Journal of the American Chemical Society 78, no. 8 (1956): 1524-1527.
^Tsai, W.-T. (2008). "Environmental and health risk analysis of nitrogen trifluoride (NF 3), a toxic and potent greenhouse gas". J. Hazard. Mater. 159 (2–3): 257–63. doi:10.1016/j.jhazmat.2008.02.023. PMID18378075.
^H. Reichardt, A. Frenzel and K. Schober (2001). "Environmentally friendly wafer production: NF 3 remote microwave plasma for chamber cleaning". Microelectronic Engineering. 56 (1–2): 73–76. doi:10.1016/S0167-9317(00)00505-0.
^J. Oshinowo; A. Riva; M Pittroff; T. Schwarze; R. Wieland (2009). "Etch performance of Ar/N2/F2 for CVD/ALD chamber clean". Solid State Technology. 52 (2): 20–24.
