Nitrogen trifluoride

Applications

Nitrogen trifluoride is used in the plasma etching of silicon wafers. Today nitrogen trifluoride is predominantly employed in the cleaning of the PECVD chambers in the high-volume production of liquid-crystal displays and silicon-based thin-film solar cells. In these applications NF₃ is initially broken down in situ by a plasma. The resulting fluorine atoms are the active cleaning agents that attack the polysilicon, silicon nitride and silicon oxide. Nitrogen trifluoride can be used as well with tungsten silicide, and tungsten produced by CVD. NF₃ has been considered as an environmentally preferable substitute for sulfur hexafluoride or perfluorocarbons such as hexafluoroethane. The process utilization of the chemicals applied in plasma processes is typically below 20%. Therefore some of the PFCs and also some of the NF₃ always escape into the atmosphere. Modern gas abatement systems can decrease such emissions.

Elemental fluorine has been introduced as an environmentally friendly replacement for nitrogen trifluoride in the manufacture of flat-panel displays and thin-film solar cells.

Nitrogen trifluoride is also used in hydrogen fluoride and deuterium fluoride lasers, which are types of chemical lasers. It is preferred to fluorine gas due to its convenient handling properties, reflecting its considerable stability.

It is compatible with steel and Monel, as well as several plastics.

Synthesis and reactivity

Nitrogen trifluoride is a rare example of a binary fluoride that can be prepared directly from the elements only at very uncommon conditions, such as electric discharge. After first attempting the synthesis in 1903, Otto Ruff prepared nitrogen trifluoride by the electrolysis of a molten mixture of ammonium fluoride and hydrogen fluoride. It proved to be 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. Today, it is prepared both by direct reaction of ammonia and fluorine and by a variation of Ruff's method. It is supplied in pressurized cylinders.

Reactions

NF₃ 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). This difference arises from the fluorine atoms acting as electron withdrawing groups, attracting essentially all of the lone pair electrons on the nitrogen atom. NF₃ is a potent yet sluggish oxidizer.

It oxidizes hydrogen chloride to chlorine:

2 NF₃ + 6 HCl → 6 HF + N₂ + 3 Cl₂

It converts to tetrafluorohydrazine upon contact with metals, but only at high temperatures:

2 NF₃ + Cu → N₂F₄ + CuF₂

NF₃ reacts with fluorine and antimony pentafluoride to give the tetrafluoroammonium salt:

NF₃ + F₂ + SbF₅ → NF+
4SbF−
6

Greenhouse gas

NF
3 is a greenhouse gas, with a global warming potential (GWP) 17,200 times greater than that of CO
2 when compared over a 100-year period. Its GWP place it 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, although other work suggests a slightly shorter lifetime of 550 years (and a corresponding GWP of 16,800).

Although NF
3 has a high GWP, for a long time its radiative forcing in the Earth's atmosphere has been assumed to be small, spuriously presuming that only small quantities are released into the atmosphere. Industrial applications of NF
3 routinely break it down, while in the past previously used regulated compounds such as SF
6 and PFCs were often released. Research has questioned the previous assumptions. High-volume applications such as DRAM computer memory production, the manufacturing of flat panel displays and the large-scale production of thin-film solar cells.

Since 1992, when less than 100 tons were produced, production has grown to an estimated 4000 tons in 2007 and is projected to increase significantly. 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. 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. 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. 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 overestimated. Instead, the contribution of the nitrogen trifluoride to the CO₂-budget of thin film solar cell production is compensated already within a few months by the CO₂ saving potential of the PV technology.

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.

Safety

Skin contact with NF
3 is not hazardous, and it is a relatively minor irritant to mucous membranes and eyes. It is a pulmonary irritant with a toxicity considerably lower than nitrogen oxides, and overexposure via inhalation causes the conversion of hemoglobin in blood to methemoglobin, which can lead to the condition methemoglobinemia. The National Institute for Occupational Safety and Health (NIOSH) specifies that the concentration that is immediately dangerous to life or health (IDLH value) is 1,000 ppm.