Nitrogen trifluoride
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2008-07-15T12:35:58Z
Plumbago
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{{Chembox new
| Name = Nitrogen trifluoride
| ImageFile = Nitrogen-trifluoride-2D-dimensions.png
<!-- | ImageSize = 120px -->
| ImageName = Nitrogen trifluoride
| ImageFile1 = Nitrogen-trifluoride-3D-vdW.png
<!-- | ImageSize1 = 120px -->
| ImageName1 = Nitrogen trifluoride
| IUPACName = Nitrogen trifluoride
| OtherNames = Nitrogen fluoride<br /> Trifluoramine<br /> Trifluorammonia<br />
| Section1 = {{Chembox Identifiers
| CASNo = 7783-54-2
}}
| Section2 = {{Chembox Properties
| Formula = NF<sub>3</sub>
| MolarMass = 71.0019 g/mol
| Appearance = colorless gas
| Density = 3.003 kg/m<sup>3</sup> (1.013 bar and 15 °C) gas<br />1540 kg/m<sup>3</sup> (1.013 bar at boiling point) liquid
| Solubility = 0.021 vol/vol (20 °C and 1 bar)
| MeltingPt = −206.8 °C (66.35 K)
| BoilingPt = −129.1 °C (144.05 K) (1.013 bar)
| pKb =
}}
| Section3 = {{Chembox Structure
| MolShape = trigonal pyramidal
| Dipole = 0.234 D
}}<!--
| [[Bond angle]]
| 102.3°
-->
| Section7 = {{Chembox Hazards
| ExternalMSDS = [http://encyclopedia.airliquide.com/sds/en/091_AL_EN.pdf Air Liquide MSDS]
| EUClass = not listed
| NFPA-H = 2 | Other=OX
| NFPA-F =
| NFPA-R =
| FlashPt = non-flammable
}}
}}
'''Nitrogen trifluoride''' is the [[inorganic compound]] with the [[chemical formula|formula]] NF<sub>3</sub>. This [[nitrogen]]-[[fluorine]] compound is a colorless, toxic, odourless, nonflammable gas. It finds increasing use as an etchant in [[microelectronics]].
==Applications==
Nitrogen trifluoride is used in the [[plasma etching]] of silicon wafers. In this application NF<sub>3</sub> is broken down into nitrogen and fluorine gases ''in situ,'' and the resulting fluorine [[radical (chemistry)|radicals]] are the active cleaning agents that attack the [[polysilicon]]. It can be used as well with [[silicon nitride]], [[tungsten silicide]], and [[tungsten]] produced by [[chemical vapor deposition|CVD]]. In these applications, NF<sub>3</sub> was originally chosen in about 2000 as an environmentally preferable substitute for [[perfluorocarbon]]s such as [[hexafluoroethane]] and [[sulfur hexafluoride]].<ref name=Reich>{{cite journal
| title = Environmentally friendly wafer production: {{chem|NF|3}} remote microwave plasma for chamber cleaning
| author = H. Reichardt , A. Frenzel and K. Schober
| doi = 10.1016/S0167-9317(00)00505-0
| journal = Microelectronic Engineering
| volume = 56
| issues = 1-2
| year = 2001
| pages 73-76}}</ref>
About two-thirds of the PFCs escaped into the atmosphere; NF<sub>3</sub> was less likely to escape into the air.
Nitrogen trifluoride is also used in [[hydrogen fluoride laser|hydrogen fluoride and deuterium fluoride lasers]], which are types of [[chemical laser]]s. It is preferred to fluorine gas due to its convenient handling properties, reflecting its considerable stability.
In 2008, about three-quarters of the chemical is now used to manufacture computer microchips; the rest is used to make [[Liquid crystal display|LCD panels]].
It is compatible with steel and [[Monel|Monel(TM)]], as well as several plastics.
==Synthesis and reactivity==
NF<sub>3</sub> is a rare example of a binary fluoride that cannot be prepared directly from the elements (i.e., N<sub>2</sub> does not react with F<sub>2</sub>). Almost all other elements in the periodic table react directly, often violently, with fluorine.
After first attempting the synthesis in [[1903]], [[Otto Ruff]] prepared nitrogen trifluoride 25 years later by the electrolysis of a molten mixture of [[ammonium fluoride]] and [[hydrogen fluoride]].<ref>{{cite journal
| title = Das Stickstoff-3-fluorid
| author = [[Otto Ruff]], Joseph Fischer, Fritz Luft
| journal = Zeitschrift für anorganische und allgemeine Chemie
| year = 1928
| volume = 172
| issue = 1
| pages = 417–425
| doi = 10.1002/zaac.19281720132
}}</ref> It proved to be far less reactive than [[nitrogen trichloride]]. Today, it is prepared both by direct reaction of ammonia and fluorine and by a variation of Ruff's method.<ref name=Kirk>Philip 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</ref>
It is supplied in pressurized cylinders.
===Reactions===
NF<sub>3</sub> 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).<ref>Thomas M. Klapötke “Nitrogen–fluorine compounds” Journal of Fluorine Chemistry Volume 127, 2006, pp. 679-687. {{doi|10.1016/j.jfluchem.2006.03.001}}</ref> NF<sub>3</sub> is a potent yet sluggish oxidizer. It converts to [[tetrafluorohydrazine]] upon contact with metals, but only at high temperatures:
:2 NF<sub>3</sub> + Cu → N<sub>2</sub>F<sub>4</sub> + CuF<sub>2</sub>
NF<sub>3</sub> reacts with fluorine and antimony pentafluoride to give the tetrafluoroammonium salt:
: NF<sub>3</sub> + F<sub>2</sub> + SbF<sub>5</sub> → NF<sub>4</sub><sup>+</sup>SbF<sub>6</sub><sup>-</sup>
==Greenhouse gas==
{{chem|NF|3}} is a potent [[greenhouse gas]], with a [[global warming potential]] (GWP) 17,200 times greater than that of [[carbon dioxide|{{chem|CO|2}}]] when compared over a 100 year period.<ref name="WG1Full">{{Citation | title=Climate Change 2007: The Physical Sciences Basis | url=http://www.ipcc.ch/pdf/assessment-report/ar4/wg1/ar4-wg1-chapter2.pdf | publisher=[[IPCC]] | accessdate=2008-07-03}}</ref><ref name="CAT.INIST">{{cite journal | last=Robson | first=J.I. | coauthors=Gohar, L.K., Hurley, M.D., Shine, K.P. and Wallington, T. | year=2006 | title=Revised IR spectrum, radiative efficiency and global warming potential of nitrogen trifluoride | journal=[[Geophysical Research Letters|Geophys. Res. Lett.]] | volume=33 | url=http://cat.inist.fr/?aModele=afficheN&cpsidt=17893800 | doi=10.1029/2006GL026210}}</ref> Its GWP would place it second only to [[Sulfur hexafluoride|{{chem|SF|6}}]] in the group of [[Kyoto_Protocol|Kyoto-recognised]] greenhouse gases, although {{chem|NF|3}} is not currently included in that grouping. It has an estimated [[Greenhouse gas#Atmospheric lifetime|atmospheric lifetime]] of 740 years,<ref name="WG1Full"/> although recent work suggests a slightly shorter lifetime of 550 years (and a GWP of 16,800).<ref name=prather08/>
Although {{chem|NF|3}} has a high GWP, its [[radiative forcing]] in the [[Earth's atmosphere]] is small, as it is only released into the atmosphere in small quantities. Industrial applications involving {{chem|NF|3}} routinely break it down as it is used, whereas the regulated compounds {{chem|SF|6}} and [[Perfluorocarbon|PFCs]] are typically released.<ref name=prather08/><ref name=tsai08>{{cite journal | last=Tsai | first= W.-T. | year=2008 | title=Environmental and health risk analysis of nitrogen trifluoride ({{chem|NF|3}}), a toxic and potent greenhouse gas | journal=J. Hazard. Mat. | volume= | doi=10.1016/j.jhazmat.2008.02.023}}</ref>
==Production volume==
Since [[1992]], when less than 100 tons were produced, production grew to an estimated 4000 tons in [[2007]] and is projected to increase significantly.<ref name=prather08>{{cite journal | last=Prather | first=M.J. | coauthors=Hsu, J. | year=2008 | title={{chem|NF|3}}, the greenhouse gas missing from Kyoto | journal=[[Geophysical Research Letters|Geophys. Res. Lett.]] | volume=35 | doi=10.1029/2008GL034542 | url=http://www.agu.org/journals/gl/gl0812/2008GL034542/}}</ref>
World production of NF<sub>3</sub> is expected to reach 8,000 tons a year by 2010. It is not currently known how much is ultimately released into the atmosphere; estimates are either less or more than 2%, but there is not good independent data about releases, nor measurements of atmospheric concentration.<ref>{{cite news | url = http://www.latimes.com/news/nationworld/nation/la-na-climate8-2008jul08,0,7460950.story | title = A climate threat from flat TVs, microchips | date = 2008-07-08 | author = M. Roosevelt}}</ref>
==Safety==
Skin contact with {{chem|NF|3}} is not hazardous, and it is a relatively minor irritant to [[mucous membrane]]s and eyes. It is a pulmonary irritant with a [[toxicity]] comparable with [[nitrogen oxides]], and overexposure via inhalation causes the conversion of [[hemoglobin]] in blood to [[methemoglobin]], which can lead to the condition [[methemoglobinemia]].<ref name=gasworld08>{{cite web | url=http://www.gasworld.com/news.php?a=2896 | title=Nitrogen trifluoride - Cleaning up in electronic applications | last=Malik | first=Yogender | publisher=[[Gasworld]] | date=2008-07-03 | accessdate=2008-07-15}}</ref>
==References==
{{reflist}}
==External links==
*[http://www.npi.gov.au/database/substance-info/profiles/44.html National Pollutant Inventory - Fluoride and compounds fact sheet]
[[Category:Nitrogen compounds]]
[[Category:Fluorides]]
[[Category:Nonmetal halides]]
[[Category:Greenhouse gases]]
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[[de:Stickstoff(III)-fluorid]]
[[ru:Трифторид азота]]