 Liquid phase Chlorine trifluoride discoloured by the presence of chlorine | |||
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| Names | |||
|---|---|---|---|
| Systematic IUPAC name Trifluoro-λ3-chlorane[1](substitutive) | |||
| Other names Chlorotrifluoride | |||
| Identifiers | |||
3D model (JSmol) | |||
| ChEBI | |||
| ChemSpider |
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| ECHA InfoCard | 100.029.301 | ||
| EC Number |
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| 1439 | |||
| MeSH | chlorine+trifluoride | ||
| RTECS number |
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| UNII | |||
| UN number | 1749 | ||
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| 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 incarbon 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] | |||
| 63.9 J K−1 mol−1 | |||
Std molar entropy(S⦵298) | 281.6 J K−1 mol−1 | ||
Std enthalpy of formation(ΔfH⦵298) | −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 theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Chlorine trifluoride is aninterhalogen compound with the formulaClF3. It is a colorless, poisonous, corrosive, and extremely reactivegas 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 andetching operations in thesemiconductor industry,[8][9] innuclear reactor fuel processing,[10] historically as a component inrocket fuels, and various other industrial operations owing to its corrosive nature.[11]
It was first reported in 1930 by Ruff and Krug who prepared it byfluorination ofchlorine; this also producedchlorine monofluoride (ClF) and the mixture was separated bydistillation.[12]
Several hundred tons are produced annually.[13]
Themolecular geometry ofClF3 is approximatelyT-shaped, with one short bond (1.598 Å) and two long bonds (1.698 Å).[14] This structure agrees with the prediction ofVSEPR theory, which predictslone pairs of electrons as occupying two equatorial positions of a hypothetic trigonalbipyramid. The elongated Cl-F axial bonds are consistent withhypervalent bonding.
ClF3 also reacts explosively with water to givehydrogen fluoride andhydrogen chloride, along with oxygen andoxygen difluoride (OF2):[15]
Upon heating, it decomposes:[13]
Reactions with many metals and even metal oxides givefluorides:[15]
ClF3 is used to produceuranium hexafluoride:
Withphosphorus, it yieldsphosphorus trichloride (PCl3) andphosphorus pentafluoride (PF5), whilesulfur yieldssulfur dichloride (SCl2) andsulfur tetrafluoride (SF4).
It reacts withcaesium fluoride to give a salt containing the anionF(ClF3)−3.[16]
It reacts with inorganic oxides to give salts ofchloronium ([ClO2]+).[17]
In thesemiconductor industry, chlorine trifluoride is used to cleanchemical 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 ofplasma since the heat of the chamber is sufficient to make it decompose and react with the semiconductor material.
ClF3 is used for the fluorination of a variety of compounds.[13]
Chlorine trifluoride has been investigated as a high-performance storableoxidizer inrocket propellant systems. Handling concerns, however, severely limit its use. The following passage by rocket scientistJohn 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 ishypergolic 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 mentionasbestos, 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 theformation 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.[18]
Under thecode nameN-Stoff ("substance N"), chlorine trifluoride was investigated for military applications by theKaiser Wilhelm Institute inNazi Germany not long before the start ofWorld War II. Tests were made against mock-ups of theMaginot Line fortifications, and it was found to be an extremely effectiveincendiary weapon andpoison gas. From 1938, construction commenced on a partlybunkered, partly subterranean 14,000 m2 (150,000 sq ft) munitions factory, theFalkenhagen industrial complex, which was intended to produce 90tonnes of N-Stoff per month, in addition tosarin (a deadlynerve agent). However, by the time it was captured by the advancingRed Army in 1945, the factory had produced only about 30 to 50 tonnes, at a cost of over 100German Reichsmarks per kilogram.a N-Stoff was never used in war.[19][20]
ClF3 is a very strongoxidizer. It is extremely reactive with most inorganic and organic materials and will combust many otherwise non-flammable materialswithout any ignition source. These reactions are often violent and in some casesexplosive.Steel,copper, andnickel are not consumed because apassivation layer of metal fluoride will form which prevents further corrosion, butmolybdenum,tungsten, andtitanium are unsuitable as their fluorides are volatile.ClF3 will quickly corrode evennoble metals like iridium, platinum, or gold, oxidizing them to chlorides and fluorides.
Thisoxidizing power, surpassing that of oxygen, causesClF3 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 ofClF3 burned through 30 cm of concrete and 90 cm of gravel beneath.[21][18] There is exactly one known fire control/suppression method capable of dealing withClF3—flooding the fire with nitrogen ornoble gases such asargon. Barring that, the area must simply be kept cool until the reaction ceases.[22] The compound reacts with water-based suppressors and CO2, rendering them counterproductive.[23]
ClF3 causes severechemical and thermal burns. The products of hydrolysis are mainlyhydrofluoric acid andhydrochloric 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.
^a Using data from Economic History Services[24] and The Inflation Calculator[25] 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.
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