Names | |
---|---|
IUPAC name
Nitrogen trifluoride
| |
Other names
Nitrogen fluoride
Trifluoramine Trifluorammonia | |
Identifiers | |
3D model (
JSmol)
|
|
ChEBI | |
ChemSpider | |
ECHA InfoCard | 100.029.097 |
EC Number |
|
1551 | |
PubChem
CID
|
|
RTECS number |
|
UNII | |
UN number | 2451 |
CompTox Dashboard (
EPA)
|
|
| |
| |
Properties | |
NF3 | |
Molar mass | 71.00 g/mol |
Appearance | colorless gas |
Odor | moldy |
Density | 3.003 kg/m3 (1 atm, 15 °C) 1.885 g/cm3 (liquid at b.p.) |
Melting point | −207.15 °C (−340.87 °F; 66.00 K) |
Boiling point | −129.06 °C (−200.31 °F; 144.09 K) |
0.021 g/100 mL | |
Vapor pressure | 44.0 atm [1](−38.5 °F or −39.2 °C or 234.0 K) [a] |
Refractive index (nD)
|
1.0004 |
Structure | |
trigonal pyramidal | |
0.234 D | |
Thermochemistry | |
Heat capacity (C)
|
53.26 J/(mol·K) |
Std molar
entropy (S⦵298) |
260.3 J/(mol·K) |
Std enthalpy of
formation (ΔfH⦵298) |
−31.4 kcal/mol
[2] −109 kJ/mol [3] |
Gibbs free energy (ΔfG⦵)
|
−84.4 kJ/mol |
Hazards | |
GHS labelling: | |
H270, H280, H332 | |
P220, P244, P260, P304+P340, P315, P370+P376, P403 | |
NFPA 704 (fire diamond) | |
Flash point | Non-flammable |
Lethal dose or concentration (LD, LC): | |
LC50 (
median concentration)
|
2000 ppm (mouse, 4
h) 9600 ppm (dog, 1 h) 7500 ppm (monkey, 1 h) 6700 ppm (rat, 1 h) 7500 ppm (mouse, 1 h) [5] |
NIOSH (US health exposure limits): | |
PEL (Permissible)
|
TWA 10 ppm (29 mg/m3) [4] |
REL (Recommended)
|
TWA 10 ppm (29 mg/m3) [4] |
IDLH (Immediate danger)
|
1000 ppm [4] |
Safety data sheet (SDS) | AirLiquide |
Related compounds | |
Other
anions
|
nitrogen trichloride nitrogen tribromide nitrogen triiodide ammonia |
Other
cations
|
phosphorus trifluoride arsenic trifluoride antimony trifluoride bismuth trifluoride |
Related binary fluoro-
azanes
|
tetrafluorohydrazine |
Related compounds
|
dinitrogen difluoride |
Except where otherwise noted, data are given for materials in their
standard state (at 25 °C [77 °F], 100 kPa).
|
Nitrogen trifluoride (NF
3) is an
inorganic, colorless, non-
flammable,
toxic gas with a slightly musty odor. It finds increasing use within the manufacturing of
flat-panel displays,
photovoltaics,
LEDs and other
microelectronics.
[6] Nitrogen trifluoride is also an extremely strong and long-lived
greenhouse gas. Its atmospheric burden exceeded 2
parts per trillion during 2019 and has doubled every five years since the late 20th century.
[7]
[8]
Nitrogen trifluoride did not exist in significant quantities on Earth prior to its synthesis by humans. It is a rare example of a binary fluoride that can be prepared directly from the elements only at very uncommon conditions, such as an electric discharge. [9] 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. [10] 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. It is prepared in modern times both by direct reaction of ammonia and fluorine and by a variation of Ruff's method. [11] 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 difference arises from the fluorine atoms acting as electron-withdrawing groups, attracting essentially all of the lone pair electrons on the nitrogen atom.
Similar to dioxygen, NF3 is a potent yet sluggish oxidizer. [11] It oxidizes hydrogen chloride to chlorine:[ citation needed]
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. [11]
Above 200-300 °C, NF3 radicalizes to nitrogen difluoride and free fluorine radicals. In the presence of metals to remove the fluorine radicals, the mixture cools to give tetrafluorohydrazine:
NF3 reacts with fluorine and antimony pentafluoride to give the tetrafluoroammonium salt: [11]
Mixtures of NF3 and B2H6 are explosive even at cryogenic temperatures, reacting to produce nitrogen gas, boron trifluoride, and hydrofluoric acid. [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
3
dissociates 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.
[14]
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.
[15] 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.
[16]
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. [17]
Nitrogen trifluoride is also used in hydrogen fluoride and deuterium fluoride lasers, which are types of chemical lasers. There it is also preferred to fluorine gas due to its more convenient handling properties
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.
[18]
[19]
[20] 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,
[18] although other work suggests a slightly shorter lifetime of 550 years (and a corresponding GWP of 16,800).
[21]
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 use NF
3.
[21]
[22]
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.
[21] 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]
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.
[29] 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.
[30]
{{
cite journal}}
: CS1 maint: multiple names: authors list (
link)
{{
cite journal}}
: Cite journal requires |journal=
(
help)
{{
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: Cite journal requires |journal=
(
help)
Names | |
---|---|
IUPAC name
Nitrogen trifluoride
| |
Other names
Nitrogen fluoride
Trifluoramine Trifluorammonia | |
Identifiers | |
3D model (
JSmol)
|
|
ChEBI | |
ChemSpider | |
ECHA InfoCard | 100.029.097 |
EC Number |
|
1551 | |
PubChem
CID
|
|
RTECS number |
|
UNII | |
UN number | 2451 |
CompTox Dashboard (
EPA)
|
|
| |
| |
Properties | |
NF3 | |
Molar mass | 71.00 g/mol |
Appearance | colorless gas |
Odor | moldy |
Density | 3.003 kg/m3 (1 atm, 15 °C) 1.885 g/cm3 (liquid at b.p.) |
Melting point | −207.15 °C (−340.87 °F; 66.00 K) |
Boiling point | −129.06 °C (−200.31 °F; 144.09 K) |
0.021 g/100 mL | |
Vapor pressure | 44.0 atm [1](−38.5 °F or −39.2 °C or 234.0 K) [a] |
Refractive index (nD)
|
1.0004 |
Structure | |
trigonal pyramidal | |
0.234 D | |
Thermochemistry | |
Heat capacity (C)
|
53.26 J/(mol·K) |
Std molar
entropy (S⦵298) |
260.3 J/(mol·K) |
Std enthalpy of
formation (ΔfH⦵298) |
−31.4 kcal/mol
[2] −109 kJ/mol [3] |
Gibbs free energy (ΔfG⦵)
|
−84.4 kJ/mol |
Hazards | |
GHS labelling: | |
H270, H280, H332 | |
P220, P244, P260, P304+P340, P315, P370+P376, P403 | |
NFPA 704 (fire diamond) | |
Flash point | Non-flammable |
Lethal dose or concentration (LD, LC): | |
LC50 (
median concentration)
|
2000 ppm (mouse, 4
h) 9600 ppm (dog, 1 h) 7500 ppm (monkey, 1 h) 6700 ppm (rat, 1 h) 7500 ppm (mouse, 1 h) [5] |
NIOSH (US health exposure limits): | |
PEL (Permissible)
|
TWA 10 ppm (29 mg/m3) [4] |
REL (Recommended)
|
TWA 10 ppm (29 mg/m3) [4] |
IDLH (Immediate danger)
|
1000 ppm [4] |
Safety data sheet (SDS) | AirLiquide |
Related compounds | |
Other
anions
|
nitrogen trichloride nitrogen tribromide nitrogen triiodide ammonia |
Other
cations
|
phosphorus trifluoride arsenic trifluoride antimony trifluoride bismuth trifluoride |
Related binary fluoro-
azanes
|
tetrafluorohydrazine |
Related compounds
|
dinitrogen difluoride |
Except where otherwise noted, data are given for materials in their
standard state (at 25 °C [77 °F], 100 kPa).
|
Nitrogen trifluoride (NF
3) is an
inorganic, colorless, non-
flammable,
toxic gas with a slightly musty odor. It finds increasing use within the manufacturing of
flat-panel displays,
photovoltaics,
LEDs and other
microelectronics.
[6] Nitrogen trifluoride is also an extremely strong and long-lived
greenhouse gas. Its atmospheric burden exceeded 2
parts per trillion during 2019 and has doubled every five years since the late 20th century.
[7]
[8]
Nitrogen trifluoride did not exist in significant quantities on Earth prior to its synthesis by humans. It is a rare example of a binary fluoride that can be prepared directly from the elements only at very uncommon conditions, such as an electric discharge. [9] 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. [10] 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. It is prepared in modern times both by direct reaction of ammonia and fluorine and by a variation of Ruff's method. [11] 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 difference arises from the fluorine atoms acting as electron-withdrawing groups, attracting essentially all of the lone pair electrons on the nitrogen atom.
Similar to dioxygen, NF3 is a potent yet sluggish oxidizer. [11] It oxidizes hydrogen chloride to chlorine:[ citation needed]
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. [11]
Above 200-300 °C, NF3 radicalizes to nitrogen difluoride and free fluorine radicals. In the presence of metals to remove the fluorine radicals, the mixture cools to give tetrafluorohydrazine:
NF3 reacts with fluorine and antimony pentafluoride to give the tetrafluoroammonium salt: [11]
Mixtures of NF3 and B2H6 are explosive even at cryogenic temperatures, reacting to produce nitrogen gas, boron trifluoride, and hydrofluoric acid. [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
3
dissociates 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.
[14]
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.
[15] 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.
[16]
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. [17]
Nitrogen trifluoride is also used in hydrogen fluoride and deuterium fluoride lasers, which are types of chemical lasers. There it is also preferred to fluorine gas due to its more convenient handling properties
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.
[18]
[19]
[20] 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,
[18] although other work suggests a slightly shorter lifetime of 550 years (and a corresponding GWP of 16,800).
[21]
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 use NF
3.
[21]
[22]
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.
[21] 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]
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.
[29] 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.
[30]
{{
cite journal}}
: CS1 maint: multiple names: authors list (
link)
{{
cite journal}}
: Cite journal requires |journal=
(
help)
{{
cite journal}}
: Cite journal requires |journal=
(
help)