Chlorine trifluoride: Difference between revisions
m Reverted edits by 71.212.148.165 (talk) to last version by Edgar181 |
Removing unsourced reference to quartz vessels given ClF3 is stated later to react hypergolically with similar materials. The decomposition temperature still needs a citation but it's at least plausible. |
||
| Line 108: | Line 108: | ||
ClF<sub>3</sub> is approximately [[T-shaped (chemistry)|T-shaped]], with one short bond (1.598 [[Ångström|Å]]) and two long bonds (1.698 Å).<ref>{{cite doi|10.1063/1.1698976}}</ref> 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]]. |
ClF<sub>3</sub> is approximately [[T-shaped (chemistry)|T-shaped]], with one short bond (1.598 [[Ångström|Å]]) and two long bonds (1.698 Å).<ref>{{cite doi|10.1063/1.1698976}}</ref> 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 ClF<sub>3</sub> is stable to 180 °C |
Pure ClF<sub>3</sub> is stable to 180 °C; above this temperature it decomposes by a [[Radical (chemistry)|free radical]] mechanism to the elements.{{Citation needed|date=January 2012}} |
||
==Reactions== |
==Reactions== |
||
| Line 138: | Line 138: | ||
}}</ref> |
}}</ref> |
||
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 products of hydrolysis are mainly [[hydrofluoric acid]] and [[hydrochloric acid]], usually released as 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. [[Hydrochloric acid]] is secondary in its danger to living organisms, but is several times more corrosive to most inorganic materials than hydrofluoric acid.{{citation needed|date=September 2012}} |
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 products of hydrolysis are mainly [[hydrofluoric acid]] and [[hydrochloric acid]], usually released as 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. [[Hydrochloric acid]] is secondary in its danger to living organisms, but is several times more corrosive to most inorganic materials than hydrofluoric acid.{{citation needed|date=September 2012}} |
||
==Uses== |
==Uses== |
||
Revision as of 23:31, 27 January 2015
| |||
| Names | |||
|---|---|---|---|
| Systematic IUPAC name
Trifluoro-λ3-chlorane[1] (substitutive) | |||
| Identifiers | |||
3D model (JSmol)
|
|||
| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.029.301 | ||
| EC Number |
| ||
| 1439 | |||
| MeSH | chlorine+trifluoride | ||
PubChem CID
|
|||
| RTECS number |
| ||
| UN number | 1749 | ||
CompTox Dashboard (EPA)
|
|||
| |||
| |||
| Properties | |||
| ClF3 | |||
| Molar mass | 92.45 g·mol−1 | ||
| Appearance | Colourless gas | ||
| Odor | sweet, pungent, irritating[2] | ||
| 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) | ||
| Reacts violently[3] | |||
| Solubility | reacts violently with benzene, toluene, ether, alcohol, acetic acid, selenium tetrafluoride, nitric acid, sulfuric acid, alkali, hexane.[3] Forms shock-sensitive explosive solution in CCl4 | ||
| Vapor pressure | 175 kPa | ||
| Viscosity | 91.82 μPa s | ||
| Structure | |||
| T-shaped | |||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
281.59 J K−1mol−1[4] | ||
Std enthalpy of
formation (ΔfH⦵298) |
−158.87 kJ mol−1[4] | ||
| Hazards | |||
| GHS labelling: | |||
   
| |||
| Danger | |||
| NFPA 704 (fire diamond) | |||
| Related compounds | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |||
Chlorine trifluoride is an interhalogen compound with the formula ClF3. This colourless, 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,[5][6] in nuclear reactor fuel processing,[7] and other industrial operations.[8]
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.[9]
- 3 F2 + Cl2 → 2 ClF3
ClF3 is approximately T-shaped, with one short bond (1.598 Å) and two long bonds (1.698 Å).[10] 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; above this temperature it decomposes by a free radical mechanism to the elements.[citation needed]
Reactions
Reaction with several 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 reacts explosively with water, in which it oxidizes water to give oxygen or in controlled quantities, oxygen difluoride (OF2), as well as hydrogen fluoride and hydrogen chloride. Metal oxides will react to form metal halides and oxygen or oxygen difluoride.
- ClF3 + 2H2O → 3HF + HCl + O2
- ClF3 + H2O → HF + HCl + OF2
The main use 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
Dissociates under the scheme:
- ClF3 ↔ ClF + F2
Hazards
ClF3 is a very strong oxidizing and fluorinating agent. It is extremely reactive with most inorganic and organic materials, including glass and teflon, 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.
The ability to surpass the oxidizing ability of oxygen leads to extreme 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. In an industrial accident, a spill of 900 kg of chlorine trifluoride burned through 30 cm of concrete and 90 cm of gravel beneath.[11] Fire control/suppression is incapable of suppressing this oxidation, therefore the surrounding area is kept cool until the reaction ceases.[12] The compound reacts violently with water-based suppressors, and oxidizes in the absence of atmospheric oxygen, rendering atmosphere-displacement suppressors such as CO2 and halon completely ineffective. It ignites glass on contact.[13]
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 products of hydrolysis are mainly hydrofluoric acid and hydrochloric acid, usually released as 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. Hydrochloric acid is secondary in its danger to living organisms, but is several times more corrosive to most inorganic materials than hydrofluoric acid.[citation needed]
Uses
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, the Falkenhagen industrial complex, 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 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.[14]
Semiconductor industry
In the semiconductor industry, chlorine trifluoride is used to clean chemical vapour deposition chambers.[15] It has the advantage that it can be used to remove semiconductor material from the chamber walls without having to dismantle the chamber.[15] 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.[15]
Rocket propellant
Chlorine trifluoride has been investigated as a high-performance storable oxidizer in rocket propellant systems. Handling concerns, however, prevented 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 which 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][16][17]
References
- ^ "Chlorine trifluoride - Compound Summary". PubChem Compound. USA: National Center for Biotechnology Information. 16 September 2004. Identification and Related Records. Retrieved 9 October 2011.
- ^ a b ClF3/Hydrazine at the Encyclopedia Astronautica.
- ^ a b Chlorine fluoride (ClF3) at Guidechem Chemical Network
- ^ a b "Chlorine trifluoride". NIST Chemistry WebBook. USA: National Institute of Standards and Technology. Gas phase thermochemistry data. Retrieved 9 October 2011.
- ^ 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.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ United States Patent 5849092 "Process for chlorine trifluoride chamber cleaning"
- ^ 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. p. 40. ISBN 0-309-10358-4.
{{cite book}}: External link in|title= - ^ United States Patent 6034016 "Method for regenerating halogenated Lewis acid catalysts"
- ^ 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.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1063/1.1698976, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1063/1.1698976instead. - ^ Air Products Safetygram, http://web.archive.org/web/20060318221608/http://www.airproducts.com/nr/rdonlyres/8479ed55-2170-4651-a3d4-223b2957a9f3/0/safetygram39.pdf
- ^ "Chlorine Trifluoride Handling Manual". Canoga Park, CA: Rocketdyne. September 1961. p. 24. Retrieved 2012-09-19.
- ^ Pradyot Patnaik (2007). A comprehensive guide to the hazardous properties of chemical substances (3rd ed.). Wiley-Interscience. p. 478. ISBN 0-471-71458-5.
- ^ "Bunker Tours" report on Falkenhagen
- ^ a b c "In Situ Cleaning of CVD Chambers". Semiconductor International. 6/1/1999.
{{cite news}}: Check date values in:|date=(help) - ^ Clark, John D. (2001). Ignition!. UMI Books on Demand. ISBN 0-8135-0725-1.
- ^ Clark, John D. (1972). Ignition! An Informal History of Liquid Rocket Propellants. Rutgers University Press. p. 214. ISBN 0-8135-0725-1.
- Notes
- 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. pp. 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.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - 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 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 (5): 246. Bibcode:1951JChEd..28..246L. doi:10.1021/ed028p246.
{{cite journal}}: CS1 maint: multiple names: authors list (link)
^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.
External links
- National Pollutant Inventory - Fluoride and compounds fact sheet
- NIST Standard Reference Database
- CDC - NIOSH Pocket Guide to Chemical Hazards - Chlorine Trifluoride
- WebElements
- Sand Won't Save You This Time, blog post by Derek Lowe on the hazards of handling ClF3