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, in nuclear reactor fuel processing, and other industrial operations.
|Systematic IUPAC name|
3D model (JSmol)
|Molar mass||92.45 g·mol−1|
|Appearance||Colorless gas or greenish-yellow liquid|
|Odor||sweet, pungent, irritating, suffocating|
|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))|
|Solubility||Reacts with benzene, toluene, ether, alcohol, acetic acid, selenium tetrafluoride, nitric acid, sulfuric acid, alkali, hexane. Soluble in CCl4 but can be explosive in high concentrations.|
|Vapor pressure||175 kPa|
|Viscosity||91.82 μPa s|
Heat capacity (C)
|63.9 J K−1mol−1|
|281.6 J K−1mol−1|
Std enthalpy of
|−163.2 kJ mol−1|
Gibbs free energy (ΔfG˚)
|−123.0 kJ mol−1|
|Main hazards||explosive when exposed to organics, reacts violently with water|
|Safety data sheet||natlex.ilo.ch|
|GHS Signal word||Danger|
|NFPA 704 (fire diamond)|
|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)
|NIOSH (US health exposure limits):|
|C 0.1 ppm (0.4 mg/m3)|
|C 0.1 ppm (0.4 mg/m3)|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Preparation, structure, and properties
- 3 F2 + Cl2 → 2 ClF3
ClF3 is approximately T-shaped, with one short bond (1.598 Å) and two long bonds (1.698 Å). 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 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
In the semiconductor industry, chlorine trifluoride is used to clean chemical vapour deposition chambers. It has the advantage that 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 enough to make it decompose and react with the semiconductor material.
Chlorine trifluoride has been investigated as a high-performance storable oxidizer in rocket propellant systems. 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.
The discovery of chlorine pentafluoride rendered ClF3 obsolete as an oxidizer.
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 14,000 m2 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.
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 are not consumed by ClF3 because a thin layer of insoluble metal fluoride will form, but molybdenum, tungsten, and titanium form volatile fluorides and are consequently unsuitable. Any equipment that comes into contact with chlorine trifluoride must be meticulously 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 fact that its oxidizing ability surpasses oxygen's 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. There is exactly one known fire control/suppression method capable of dealing with chlorine trifluoride - the use of nitrogen and noble gases: the surrounding area must be flooded with nitrogen or helium. Barring that, the area must simply be kept cool until the reaction ceases. 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.
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.
^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.
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