Chlorine trifluoride

Chlorine trifluoride
Names
Systematic IUPAC name
Trifluoro-λ3-chlorane[1] (substitutive)
Other names
Chlorotrifluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.029.301
EC Number 232-230-4
1439
MeSH chlorine+trifluoride
RTECS number FO2800000
UN number 1749
Properties
ClF3
Molar mass 92.45 g·mol−1
Appearance Colorless gas or greenish-yellow liquid
Odor sweet, pungent, irritating, suffocating[2][3]
Density 4 mg cm−3
Melting point −76.34 °C (−105.41 °F; 196.81 K)
Boiling point 11.75 °C (53.15 °F; 284.90 K) (decomposes @ 180 °C (356 °F; 453 K))
Exothermic hydrolysis[4]
Solubility Reacts with benzene, toluene, ether, alcohol, acetic acid, selenium tetrafluoride, nitric acid, sulfuric acid, alkali, hexane.[4] Soluble in CCl4 but can be explosive in high concentrations.
Vapor pressure 175 kPa
-26.5·10−6 cm3/mol
Viscosity 91.82 μPa s
Structure
T-shaped
Thermochemistry
281.59 J K−1mol−1[5]
−158.87 kJ mol−1[5]
Hazards
Main hazards explosive when exposed to organics, reacts violently with water[3]
Safety data sheet natlex.ilo.ch
GHS pictograms
GHS signal word Danger
NFPA 704
Flash point noncombustible [3]
Lethal dose or concentration (LD, LC):
95 ppm (rat, 4 hr)
178 ppm (mouse, 1 hr)
230 ppm (monkey, 1 hr)
299 ppm (rat, 1 hr)
[6]
US health exposure limits (NIOSH):
PEL (Permissible)
 C 0.1 ppm (0.4 mg/m3)[3]
REL (Recommended)
 C 0.1 ppm (0.4 mg/m3)[3]
IDLH (Immediate danger)
 20 ppm[3]
Related compounds
Related compounds
 Chlorine pentafluoride

Chlorine monofluoride
Bromine trifluoride
Iodine trifluoride

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☑Y verify (what is ☑Y☒N ?)
Infobox references

Chlorine trifluoride is an interhalogen compound with the formula ClF3. This colorless, poisonous, corrosive, and extremely reactive gas condenses to a pale-greenish yellow liquid, the form in which it is most often sold (pressurized at room temperature). The compound is primarily of interest as a component in rocket fuels, in plasmaless cleaning and etching operations in the semiconductor industry,[7][8][9] in nuclear reactor fuel processing,[10] and other industrial operations.[11]

Preparation, structure, and properties

It was first reported in 1930 by Ruff and Krug who prepared it by fluorination of chlorine; this also produced ClF and the mixture was separated by distillation.[12]

3 F2 + Cl2 → 2 ClF3

ClF3 is approximately T-shaped, with one short bond (1.598 Å) and two long bonds (1.698 Å).[13] This structure agrees with the prediction of VSEPR theory, which predicts lone pairs of electrons as occupying two equatorial positions of a hypothetic trigonal bipyramid. The elongated Cl-F axial bonds are consistent with hypervalent bonding.

Pure ClF3 is stable to 180 °C in quartz vessels; above this temperature it decomposes by a free radical mechanism to its constituent elements.

Reactions

Reactions with many metals give chlorides and fluorides; phosphorus yields phosphorus trichloride (PCl3) and phosphorus pentafluoride (PF5); and sulfur yields sulfur dichloride (SCl2) and sulfur tetrafluoride (SF4). ClF3 also violently reacts with water, oxidizing it to give oxygen or, in controlled quantities, oxygen difluoride (OF2), as well as hydrogen fluoride and hydrogen chloride:

ClF3 + 2H2O → 3HF + HCl + O2
ClF3 + H2O → HF + HCl + OF2

It will also convert many metal oxides to metal halides and oxygen or oxygen difluoride.

One of the main uses of ClF3 is to produce uranium hexafluoride, UF6, as part of nuclear fuel processing and reprocessing, by the fluorination of uranium metal:

U + 3 ClF3 → UF6 + 3 ClF

The compound can also dissociate under the scheme:

ClF3 → ClF + F2

Uses

Semiconductor industry

In the semiconductor industry, chlorine trifluoride is used to clean chemical vapour deposition chambers.[14] It has the advantage that it can be used to remove semiconductor material from the chamber walls without the need to dismantle the chamber.[14] Unlike most of the alternative chemicals used in this role, it does not need to be activated by the use of plasma since the heat of the chamber is enough to make it decompose and react with the semiconductor material.[14]

Rocket propellant

Chlorine trifluoride has been investigated as a high-performance storable oxidizer in rocket propellant systems and it may be used as such someday. Handling concerns, however, severely limit its use. John Drury Clark summarized the difficulties:

It is, of course, extremely toxic, but that's the least of the problem. It is hypergolic with every known fuel, and so rapidly hypergolic that no ignition delay has ever been measured. It is also hypergolic with such things as cloth, wood, and test engineers, not to mention asbestos, sand, and water—with which it reacts explosively. It can be kept in some of the ordinary structural metals—steel, copper, aluminum, etc.—because of the formation of a thin film of insoluble metal fluoride that protects the bulk of the metal, just as the invisible coat of oxide on aluminum keeps it from burning up in the atmosphere. If, however, this coat is melted or scrubbed off, and has no chance to reform, the operator is confronted with the problem of coping with a metal-fluorine fire. For dealing with this situation, I have always recommended a good pair of running shoes.[2][15][16]

Proposed military applications

Under the code name N-Stoff ("substance N"), chlorine trifluoride was investigated for military applications by the Kaiser Wilhelm Institute in Nazi Germany not long before the start of World War II. Tests were made against mock-ups of the Maginot Line fortifications, and it was found to be an effective combined incendiary weapon and poison gas. From 1938, construction commenced on a partly bunkered, partly subterranean 31.76 km2 munitions factory, the Falkenhagen industrial complex, which was intended to produce 90 tonnes of N-Stoff per month, plus sarin. However, by the time it was captured by the advancing Red Army in 1945, the factory had produced only about 30 to 50 tonnes, at a cost of over 100 German Reichsmark per kilograma. N-Stoff was never used in war.[17]

Hazards

ClF3 is a very strong oxidizing and fluorinating agent. It is extremely reactive with most inorganic and organic materials, such as glass, and will initiate the combustion of many otherwise non-flammable materials without any ignition source. These reactions are often violent, and in some cases explosive. Vessels made from steel, copper, or nickel resist the attack of the material due to formation of a thin layer of insoluble metal fluoride, but molybdenum, tungsten, and titanium form volatile fluorides and are consequently unsuitable. Any equipment that comes into contact with chlorine trifluoride must be scrupulously cleaned and then passivated, because any contamination left may burn through the passivation layer faster than it can re-form. Chlorine trifluoride has also been known to corrode materials otherwise known to be non-corrodible such as iridium, platinum, and gold.

The power to surpass the oxidizing ability of oxygen leads to corrosivity against oxide-containing materials often thought as incombustible. Chlorine trifluoride and gases like it have been reported to ignite sand, asbestos, and other highly fire-retardant materials. It will also ignite the ashes of materials that have already been burned in oxygen. In an industrial accident, a spill of 900 kg of chlorine trifluoride burned through 30 cm of concrete and 90 cm of gravel beneath.[18][16] Other than the use of nitrogen and noble gases, no known fire control/suppression methods are capable of suppressing this oxidation, so the surrounding area must be flooded with nitrogen or helium or simply kept cool until the reaction ceases.[19] The compound reacts with water-based suppressors, and oxidizes even in the absence of atmospheric oxygen, rendering traditional atmosphere-displacement suppressors such as CO2 and halon ineffective. It ignites glass on contact.[20]

Exposure to larger amounts of chlorine trifluoride, as a liquid or as a gas, ignites living tissue. The hydrolysis reaction with water is violent and exposure results in a thermal burn. The products of hydrolysis are mainly hydrofluoric acid and hydrochloric acid, usually released as acidic steam or vapor due to the highly exothermic nature of the reaction. Hydrofluoric acid is corrosive to human tissue, is absorbed through skin, selectively attacks bone, interferes with nerve function, and causes often-fatal fluorine poisoning. Although hydrochloric acid is much less toxic to humans, it is often more corrosive than hydrofluoric acid.

See also

References

  1. ↑ "Chlorine trifluoride - Compound Summary". PubChem Compound. USA: National Center for Biotechnology Information. 16 September 2004. Identification and Related Records. Retrieved 9 October 2011.
  2. 1 2 ClF3/Hydrazine Archived 2007-02-02 at the Wayback Machine. at the Encyclopedia Astronautica.
  3. 1 2 3 4 5 6 "NIOSH Pocket Guide to Chemical Hazards #0117". National Institute for Occupational Safety and Health (NIOSH).
  4. 1 2 Chlorine fluoride (ClF3) Archived 2013-10-29 at the Wayback Machine. at Guidechem Chemical Network
  5. 1 2 "Chlorine trifluoride". NIST Chemistry WebBook. USA: National Institute of Standards and Technology. Gas phase thermochemistry data. Retrieved 9 October 2011.
  6. ↑ "Chlorine trifluoride". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  7. ↑ Hitoshi Habuka; Takahiro Sukenobu; Hideyuki Koda; Takashi Takeuchi; Masahiko Aihara (2004). "Silicon Etch Rate Using Chlorine Trifluoride". Journal of the Electrochemical Society. 151 (11): G783–G787. doi:10.1149/1.1806391. (author ResearchGate link)
  8. ↑ United States Patent 5849092 "Process for chlorine trifluoride chamber cleaning" Archived 2007-09-26 at the Wayback Machine.
  9. ↑ Habuka, Hitoshi (2012). "Etching of Silicon Carbide Using Chlorine Trifluoride Gas". Physics and Technology of Silicon Carbide Devices. doi:10.5772/50387. ISBN 978-953-51-0917-4.
  10. ↑ Board on Environmental Studies and Toxicology, (BEST) (2006). Acute Exposure Guideline Levels for Selected Airborne Chemicals: Volume 5. Washington D.C.: National Academies Press. p. 40. ISBN 978-0-309-10358-9. (available from National Academies Press)
  11. ↑ United States Patent 6034016 "Method for regenerating halogenated Lewis acid catalysts" Archived 2007-09-26 at the Wayback Machine.
  12. ↑ Otto Ruff, H. Krug (1930). "Über ein neues Chlorfluorid-CIF3". Zeitschrift für anorganische und allgemeine Chemie. 190 (1): 270–276. doi:10.1002/zaac.19301900127.
  13. ↑ Smith, D. F. (1953). "The Microwave Spectrum and Structure of Chlorine Trifluoride". The Journal of Chemical Physics. 21 (4): 609–614. Bibcode:1953JChPh..21..609S. doi:10.1063/1.1698976.
  14. 1 2 3 "In Situ Cleaning of CVD Chambers". Semiconductor International. June 1, 1999.
  15. ↑ Clark, John D. (2001). Ignition!. UMI Books on Demand. ISBN 978-0-8135-0725-5.
  16. 1 2 Clark, John D. (1972). Ignition! An Informal History of Liquid Rocket Propellants. Rutgers University Press. p. 214. ISBN 978-0-8135-0725-5.
  17. ↑ "Germany 2004". www.bunkertours.co.uk.
  18. ↑ Air Products Safetygram, https://web.archive.org/web/20060318221608/http://www.airproducts.com/nr/rdonlyres/8479ed55-2170-4651-a3d4-223b2957a9f3/0/safetygram39.pdf
  19. ↑ "Chlorine Trifluoride Handling Manual". Canoga Park, CA: Rocketdyne. September 1961. p. 24. Retrieved 2012-09-19.
  20. ↑ Pradyot Patnaik (2007). A comprehensive guide to the hazardous properties of chemical substances (3rd ed.). Wiley-Interscience. p. 478. ISBN 978-0-471-71458-3.
Notes

  • Groehler, Olaf (1989). Der lautlose Tod. Einsatz und Entwicklung deutscher Giftgase von 1914 bis 1945. Reinbek bei Hamburg: Rowohlt. ISBN 978-3-499-18738-4.
  • Ebbinghaus, Angelika (1999). Krieg und Wirtschaft: Studien zur deutschen Wirtschaftsgeschichte 1939–1945. Berlin: Metropol. pp. 171–194. ISBN 978-3-932482-11-3.
  • Harold Simmons Booth, John Turner Pinkston, , Jr. (1947). "The Halogen Fluorides". Chemical Reviews. 41 (3): 421–439. doi:10.1021/cr60130a001.
  • Yu D Shishkov; A A Opalovskii (1960). "Physicochemical Properties of Chlorine Trifluoride". Russian Chemical Reviews. 29 (6): 357–364. Bibcode:1960RuCRv..29..357S. doi:10.1070/RC1960v029n06ABEH001237.
  • Robinson D. Burbank; Frank N. Bensey (1953). "The Structures of the Interhalogen Compounds. I. Chlorine Trifluoride at −120 °C". The Journal of Chemical Physics. 21 (4): 602–608. Bibcode:1953JChPh..21..602B. doi:10.1063/1.1698975.
  • A. A. Banks; A. J. Rudge (1950). "The determination of the liquid density of chlorine trifluoride". Journal of the Chemical Society: 191–193. doi:10.1039/JR9500000191.
  • Lowdermilk, F. R.; Danehower, R. G.; Miller, H. C. (1951). "Pilot plant study of fluorine and its derivatives". Journal of Chemical Education. 28 (5): 246. Bibcode:1951JChEd..28..246L. doi:10.1021/ed028p246.

^a Using data from Economic History Services and The Inflation Calculator, we can calculate that 100 Reichsmark in 1941 is approximately equivalent to US$540 in 2006. Reichsmark exchange rate values from 1942 to 1944 are fragmentary.

This article is issued from Wikipedia. The text is licensed under Creative Commons - Attribution - Sharealike. Additional terms may apply for the media files.