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Chlorine trifluoride

Chlorine trifluoride
CAS number 7790-91-2
Molecular formula ClF3
Molar mass 92.45 g/mol
Melting point

-76.3 °C

Boiling point

11.75 °C

Solubility in other solvents Hydrolysis
Std enthalpy of
-158.87 kJ/mol
Standard molar
281.59 J.K–1.mol–1
Main hazards Toxic, corrosive, oxidizer.
NFPA 704
Related Compounds
Related compounds ClF5
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Chlorine trifluoride is the chemical compound with the formula ClF3. This colourless, poisonous, corrosive and very 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 industrial cleaning and etching operations primarily in the semiconductor industry[1] [2], nuclear reactor fuel processing[3] and other industrial operations.[4]


Preparation, structure, and properties

Main article: chlorine fluorides

It was first reported by Ruff and Krug who prepared it by fluorination of chlorine, this also produced ClF and the mixture was separated by distillation.[5]

3 F2 + Cl2 → 2 ClF3

ClF3 is approximately T-shaped. This structure is explicable in the context of VSEPR theory, which considers also lone pairs of electrons as occupying two equatorial positions of a hypothetic trigonal bipyramid. The elongated Cl-Faxial bonds are consistent with hypervalent bonding.

Pure ClF3 is stable to 180° in glass vessels, but above this temperature it decomposes by a free radical mechanism to the elements.


ClF3 is a very strong oxidizing and fluorination agent. ClF3 is extremely reactive with most inorganic and organic materials and will initiate the combustion of many materials without an ignition source and these reactions are often violent or in some cases explosive. Several metals give chlorides and fluorides, phosphorus yields PCl3 plus PF5, sulfur SCl2 plus SF4. ClF3 is also violently water reactive in which it hydrolyses to a variety of hazardous chemicals such as hydrofluoric acid. H2S explodes on being mixed with ClF3 at room temperature.

The ability to surpass the oxidizing ability of oxygen leads to corrosivity against oxide-containing materials often thought as incombustible. In an industrial accident, a spill of 900 kg of chlorine trifluoride burned itself through 30 cm of concrete and 90 cm of gravel beneath.[6] Any equipment that comes into contact with chlorine trifluoride must be carefully selected and cleaned, because any contamination can ignite on contact.

Exposure of larger amounts of chlorine trifluoride, as a liquid or as a gas, ignites tissue. The hydrolysis reaction with water is violent and exposure results in a thermal burn. The product of hydrolysis is hydrofluoric acid, which is corrosive to human tissue, absorbs through skin, selectively attacks bone and stimulates pain nerves, and causes a potentially lethal poisoning.

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 from slightly 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 at Falkenhagen which was intended to produce 50 tonnes of N-stoff per month, plus Sarin. However by the time it was captured by the advancing Red Army in 1944, 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.[7]

Rocket propellant

Chlorine trifluoride has been investigated as a high-performance storable oxidizer in rocket propellant systems. Handling concerns, however, prevented its use. 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."[8][9]


  1. ^ Hitoshi Habuka, Takahiro Sukenobu, Hideyuki Koda, Takashi Takeuchi, and Masahiko Aihara (2004). "Silicon Etch Rate Using Chlorine Trifluoride". Journal of the Electrochemical Society 151 (11): G783–G787. doi:10.1149/1.1806391.
  2. ^ United States Patent 5849092 "Process for chlorine trifluoride chamber cleaning"
  3. ^ Board on Environmental Studies and Toxicology, (BEST) (2006). Acute Exposure Guideline Levels for Selected Airborne Chemicals: Volume 5 (citation at the National Academies Press). Washington D.C.: National Academies Press, 40. ISBN 0-309-10358-4. 
  4. ^ United States Patent 6034016 "Method for regenerating halogenated Lewis acid catalysts"
  5. ^ Otto Ruff, H. Krug (1931). "Über ein neues Chlorfluorid-CIF3". Zeitschrift für anorganische und allgemeine Chemie 190 (1): 602–608. doi:10.1002/zaac.19301900127.
  6. ^ Air Products Safetygram.
  7. ^ "Bunker Tours" report on Falkenhagen
  8. ^ Clark, John D. (2001). Ignition!. UMI Books on Demand. ISBN 0-8135-0725-1. 
  9. ^ ClF3/Hydrazine at the Encyclopedia Astronautica.
  • Groehler, Olaf (1989). Der lautlose Tod. Einsatz und Entwicklung deutscher Giftgase von 1914 bis 1945. Reinbek bei Hamburg: Rowohlt. ISBN 3-499-18738-8. 
  • Ebbinghaus, Angelika (1999). Krieg und Wirtschaft: Studien zur deutschen Wirtschaftsgeschichte 1939–1945. Berlin: Metropol, 171–194. ISBN 3-932482-11-5. 
  • 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. 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. doi:10.1063/1.1698975.
  • A. A. Banks and 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: 246.

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

This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Chlorine_trifluoride". A list of authors is available in Wikipedia.
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