Chlorine trifluoride

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Chlorine trifluoride
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Template:Longitem 1439
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MeSH chlorine+trifluoride
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UN number 1749
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Template:Longitem Template:Chembox Elements/molecular formula
Molar mass Template:Chem molar mass
Appearance Colorless gas or greenish-yellow liquid
Odor Sweet, pungent, irritating, suffocating[2][3]
Density 3.779 g/L[4]
Melting point Template:Chembox CalcTemperatures
Boiling point Template:Chembox CalcTemperatures
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
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Viscosity 91.82 μPa s
Template:Longitem T-shaped molecular geometry
Template:Longitem 63.9 J K−1 mol−1
Template:Longitem 281.6 J K−1 mol−1
Template:Longitem −163.2 kJ mol−1
NFPA 704 (fire diamond) Template:NFPA 704 diamond
Flash point Template:Chembox CalcTemperatures
Template:Longitem Chlorine pentafluoride
Chlorine monofluoride
Bromine trifluoride
Iodine trifluoride

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Chlorine trifluoride is an interhalogen compound with the formula Template:Chem2. It is a colorless, poisonous, corrosive, and extremely reactive gas that condenses to a pale-greenish yellow liquid, the form in which it is most often sold (pressurized at room temperature). It is notable for its extreme oxidation properties. The compound is primarily of interest in plasmaless cleaning and etching operations in the semiconductor industry,[7][8] in nuclear reactor fuel processing,[9] historically as a component in rocket fuels, and various other industrial operations owing to its corrosive nature.[10]

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.[11]

Template:Chem2

Several hundred tons are produced annually.[12]

The molecular geometry of Template:Chem2 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.

Reactions

Template:Chem2 also reacts explosively with water to give hydrogen fluoride and hydrogen chloride, along with oxygen and oxygen difluoride (Template:Chem2):[14]

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Template:Chem2

Upon heating, it decomposes:[12]

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Reactions with many metals and even metal oxides give fluorides:[14]

Template:Chem2
Template:Chem2

Template:Chem2 is used to produce uranium hexafluoride:

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With phosphorus, it yields phosphorus trichloride (Template:Chem2) and phosphorus pentafluoride (Template:Chem2), while sulfur yields sulfur dichloride (Template:Chem2) and sulfur tetrafluoride (Template:Chem2).

It reacts with caesium fluoride to give a salt containing the anion Template:Chem2.[15]

It reacts with inorganic oxides to give salts of chloronium ([ClO2]+).[16]

Uses

Semiconductor industry

In the semiconductor industry, chlorine trifluoride is used to clean chemical vapour deposition chambers. It can be used to remove semiconductor material from the chamber walls without the need to dismantle the chamber. 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 sufficient to make it decompose and react with the semiconductor material.

Fluorination reagent

Template:Chem2 is used for the fluorination of a variety of compounds.[12]

Military applications

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.[17]

Chlorine pentafluoride (Template:Chem2) 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.

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 extremely effective incendiary weapon and poison gas. From 1938, construction commenced on a partly bunkered, partly subterranean Script error: No such module "convert". munitions factory, the Falkenhagen industrial complex, which was intended to produce 90 tonnes of N-Stoff per month, in addition to sarin (a deadly nerve agent). 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 Reichsmarks per kilogram.a N-Stoff was never used in war.[18]Template:R/superscript[19]

Hazards

Template:Chem2 is a very strong oxidizer. It is extremely reactive with most inorganic and organic materials and will combust many otherwise non-flammable materials without any ignition source. These reactions are often violent and in some cases explosive. Steel, copper, and nickel are not consumed because a passivation layer of metal fluoride will form which prevents further corrosion, but molybdenum, tungsten, and titanium are unsuitable as their fluorides are volatile. Template:Chem2 will quickly corrode even noble metals like iridium, platinum, or gold, oxidizing them to chlorides and fluorides.

This oxidizing power, surpassing that of oxygen, causes Template:Chem2 to react vigorously with many other materials often thought of as incombustible and refractory. It ignites sand, asbestos, glass, and even ashes of substances that have already burned in oxygen. In one particular industrial accident, a spill of 900 kg of Template:Chem2 burned through 30 cm of concrete and 90 cm of gravel beneath.[20][17] There is exactly one known fire control/suppression method capable of dealing with Template:Chem2Template:Px2Template:MdashTemplate:Px2flooding the fire with nitrogen or noble gases such as argon. Barring that, the area must simply be kept cool until the reaction ceases.[21] The compound reacts with water-based suppressors and CO2, rendering them counterproductive.[22]

Template:Chem2 causes severe chemical and thermal 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, and these substances present hazards of their own.

See also

Explanatory notes

<templatestyles src="Citation/styles.css"/>^a Using data from Economic History Services[23] and The Inflation Calculator[24] 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

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  6. ↑ Template:IDLH
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  8. ↑ Xi, Ming et al. (1997) U.S. patent 5849092 "Process for chlorine trifluoride chamber cleaning"
  9. ↑ Script error: No such module "citation/CS1". (available from National Academies Press Script error: No such module "webarchive".Template:Open access)
  10. ↑ Boyce, C. Bradford and Belter, Randolph K. (1998) U.S. patent 6034016 "Method for regenerating halogenated Lewis acid catalysts"
  11. ↑ Script error: No such module "Citation/CS1".
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  20. ↑ Safetygram. Air Products
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Further reading

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External links

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