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 |
| ||
| 1439 | |||
| MeSH | chlorine+trifluoride | ||
PubChem CID
|
|||
RTECS number
|
| ||
| UNII | |||
| UN number | 1749 | ||
CompTox Dashboard (EPA)
|
|||
| |||
| |||
| Properties | |||
| ClF3 | |||
| Molar mass | 92.45 g·mol−1 | ||
| Appearance | Colorless gas or greenish-yellow liquid | ||
| Odor | Sweet, pungent, irritating, suffocating[2][3] | ||
| Density | 3.779 g/L[4] | ||
| Melting point | −76.34 °C (−105.41 °F; 196.81 K)[4] | ||
| Boiling point | 11.75 °C (53.15 °F; 284.90 K)[4] (decomposes at 180 °C, 356 °F, 453 K) | ||
| Reacts with water[1] | |||
| Solubility | Soluble in carbon tetrachloride but explosive in high concentrations. Reacts with hydrogen-containing compounds e.g. hydrogen, methane, benzene, ether, ammonia.[1] | ||
| Vapor pressure | 175 kPa | ||
| −26.5×10−6 cm3/mol[5] | |||
| Viscosity | 91.82 μPa s | ||
| Structure | |||
| T-shaped molecular geometry | |||
| Thermochemistry[6] | |||
Heat capacity (C)
|
63.9 J K−1 mol−1 | ||
Std molar
entropy (S⦵298) |
281.6 J K−1 mol−1 | ||
Std enthalpy of (ΔfH⦵298)formation |
−163.2 kJ mol−1 | ||
Gibbs free energy (ΔfG⦵)
|
−123.0 kJ mol−1 | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards
|
Very toxic, very corrosive, powerful oxidizer, violent hydrolysis[3] | ||
| GHS labelling: | |||
   
| |||
| Danger | |||
| NFPA 704 (fire diamond) | |||
| Flash point | Noncombustible[3] | ||
| Lethal dose or concentration (LD, LC): | |||
LC50 (median concentration)
|
95 ppm (rat, 4 hr) 178 ppm (mouse, 1 hr) 230 ppm (monkey, 1 hr) 299 ppm (rat, 1 hr) [7] | ||
| NIOSH (US health exposure limits): | |||
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] | ||
| Safety data sheet (SDS) | [1] | ||
| 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).
| |||
Chlorine trifluoride is an
Preparation, structure, and properties
It was first reported in 1930 by Ruff and Krug who prepared it by fluorination of chlorine; this also produced Chlorine monofluoride (ClF) and the mixture was separated by distillation.[12]
- 3 F2 + Cl2 → 2 ClF3
The
Pure ClF3 is stable to 180 °C (356 °F) in quartz vessels; above this temperature, it decomposes by a free radical mechanism to its constituent elements.[clarification needed][citation needed]
Reactions
Reactions with many metals give chlorides and fluorides. With phosphorus, it yields phosphorus trichloride (PCl3) and phosphorus pentafluoride (PF5), while sulfur yields sulfur dichloride (SCl2) and sulfur tetrafluoride (SF4). ClF3 also reacts with water to give hydrogen fluoride and hydrogen chloride, along with oxygen and oxygen difluoride (OF2):
- ClF3 + H2O → HF + HCl + OF2
- ClF3 + 2H2O → 3HF + HCl + O2
It will also convert many
It occurs as a ligand in the complex CsF(ClF3)3.[14]
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
Rocket propellant
Chlorine trifluoride has been investigated as a high-performance storable oxidizer in rocket propellant systems. Handling concerns, however, severely limit its use. The following passage by rocket scientist John D. Clark is widely quoted in descriptions of the substance's extremely hazardous nature:
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 aluminium 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.[16]
Chlorine pentafluoride (ClF5) has also been investigated as a potential rocket oxidizer. It offered improved specific impulse over chlorine trifluoride, but with all of the same difficulties in handling. Neither compound has been used in any operational rocket propulsion system.
Proposed military applications
Under the
Hazards
ClF3 is a very strong
This
Exposure to larger amounts of ClF3, as a liquid or as a gas, ignites living tissue, resulting in severe chemical and thermal burns. ClF3 reacts violently with water and exposure to the reaction also results in burns. The products of hydrolysis are mainly hydrofluoric acid and hydrochloric acid, which are usually released as steam or vapor due to the highly exothermic nature of the reaction.
See also
Explanatory notes
^a Using data from Economic History Services[22] and The Inflation Calculator[23] it can be calculated that the sum of 100 Reichsmarks in 1941 is approximately equivalent to US$4,652.50 in 2021. Reichsmark exchange rate values from 1942 to 1944 are fragmentary.
References
- ^ a b c "Chlorine trifluoride". PubChem Compound. National Center for Biotechnology Information. 4 July 2023. Retrieved 8 July 2023.
- ^ ClF3/Hydrazine Archived 2007-02-02 at the Wayback Machine at the Encyclopedia Astronautica.
- ^ a b c d e f NIOSH Pocket Guide to Chemical Hazards. "#0117". National Institute for Occupational Safety and Health (NIOSH).
- ^ ISBN 978-1-4398-5511-9.
- ISBN 978-1-4398-5511-9.
- ISBN 978-1-4398-5511-9.
- ^ "Chlorine trifluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- from the original on 2022-01-25. Retrieved 2017-04-11.
- ^ Xi, Ming et al. (1997) U.S. patent 5,849,092 "Process for chlorine trifluoride chamber cleaning"
- ISBN 978-0-309-10358-9. (available from National Academies Press Archived 2014-11-07 at the Wayback Machine)
- ^ Boyce, C. Bradford and Belter, Randolph K. (1998) U.S. patent 6,034,016 "Method for regenerating halogenated Lewis acid catalysts"
- .
- .
- PMID 32608053.
- ^ a b c "In Situ Cleaning of CVD Chambers". Semiconductor International. June 1, 1999.[permanent dead link]
- ^ ISBN 978-0-8135-0725-5.
- doi:10.1038/438427a.
- ^ "Germany 2004". www.bunkertours.co.uk. Archived from the original on 2006-06-13. Retrieved 2006-06-13.
- ^ Safetygram. Air Products
- ^ "Chlorine Trifluoride Handling Manual". Canoga Park, CA: Rocketdyne. September 1961. p. 24. Archived from the original on 2013-04-08. Retrieved 2012-09-19.
- ISBN 978-0-471-71458-3.
- ^ Officer, Lawrence H. (2002), Exchange Rate Between the United States Dollar and Forty Other Countries, 1913–1999, EH.net (Economic History Services), archived from the original on 15 June 2006, retrieved 7 July 2023
- ^ "The Inflation Calculator". S. Morgan Friedman's 'Webpage': Ceci N'est Pas Une Homepage. Retrieved 7 July 2023.
Further reading
- 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.
- Booth, Harold Simmons; Pinkston, John Turner Jr. (1947). "The Halogen Fluorides". PMID 18895518.
- Yu D Shishkov; A A Opalovskii (1960). "Physicochemical Properties of Chlorine Trifluoride". S2CID 250863587.
- Robinson D. Burbank; Frank N. Bensey (1953). "The Structures of the Interhalogen Compounds. I. Chlorine Trifluoride at −120 °C". .
- A. A. Banks; A. J. Rudge (1950). "The determination of the liquid density of chlorine trifluoride". .
- Lowdermilk, F. R.; Danehower, R. G.; Miller, H. C. (1951). "Pilot plant study of fluorine and its derivatives". .