Manganese(IV) fluoride

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Manganese(IV) fluoride

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Template:Longitem	MnF4
Molar mass	130.93 g mol−1
Appearance	blue solid
Density	3.61 g cm−3 (calc.)
Melting point	Template:Chembox CalcTemperatures
Template:Longitem	tetragonal, tI80
Template:Longitem	I41/a (No. 88)
Template:Longitem	a = 1263 pm, c = 604.9 pm
Template:Longitem	Manganese(II) fluoride Manganese(III) fluoride

Manganese tetrafluoride, MnF₄, is the highest fluoride of manganese. It is a powerful oxidizing agent and is used as a means of purifying elemental fluorine.

Preparation

Manganese tetrafluoride was first unequivocally prepared in 1961 by the reaction of manganese(II) fluoride (or other Mn^II compounds) with a stream of fluorine gas at 550 °C: the MnF₄ sublimes into the gas stream and condenses onto a cold finger. This is still the commonest method of preparation, although the sublimation can be avoided by operating at increased fluorine pressure (4.5–6 bar at 180–320 °C) and mechanically agitating the powder to avoid sintering of the grains. The reaction can also be carried out starting from manganese powder in a fluidized bed.

Other preparations of MnF₄ include the fluorination of MnF₂ with krypton difluoride, or with F₂ in liquid hydrogen fluoride solution under ultraviolet light. Manganese tetrafluoride has also been prepared (but not isolated) in an acid–base reaction between antimony pentafluoride and K₂MnF₆ as part of a chemical synthesis of elemental fluorine.

K₂MnF₆ + 2 SbF₅ → MnF₄ + 2 KSbF₆

Chemistry

Decomposition

Manganese tetrafluoride is in equilibrium with manganese(III) fluoride and elemental fluorine:

MnF₄ ⇌ MnF₃ + Template:Sfrac F₂

Decomposition is favoured by increasing temperature, and disfavoured by the presence of fluorine gas, but the exact parameters of the equilibrium are unclear, with some sources saying that MnF₄ will decompose slowly at room temperature, others placing a practical lower temperature limit of 70 °C, and another claiming that MnF₄ is essentially stable up to 320 °C. The equilibrium pressure of fluorine above MnF₄ at room temperature has been estimated at 10⁻⁴ Pa (10⁻⁹ bar), and the enthalpy change of reaction at +44(8) kJ mol⁻¹.

Other reactions

Manganese tetrafluoride reacts violently with water and even with sodium-dried petroleum ether. It immediately decomposes on contact with moist air.

Reaction with alkali metal fluorides or concentrated hydrofluoric acid gives the yellow hexafluoromanganate(IV) anion [MnF₆]²⁻.

MnF₄ reacts with XeF₂ to form Lewis acid-base adducts: 3XeF₂∙2MnF₄, XeF₂∙MnF₄, and XeF₂∙2MnF₄. A tetrameric F-bridged ring with XeF₂ molecules coordinated to metal atoms, observed in the crystal structure of XeF₂∙MnF₄ (XeMnF₆) adduct, could serve as a structural model for the currently unknown structure of the first noble-gas compound, XePtF₆.

Applications

The main application of manganese tetrafluoride is in the purification of elemental fluorine. Fluorine gas is produced by electrolysis of anhydrous hydrogen fluoride (with a small amount of potassium fluoride added as a support electrolyte) in a Moissan cell. The technical product is contaminated with HF, much of which can be removed by passing the gas over solid KF, but also with oxygen (from traces of water) and possibly heavy-metal fluorides such as arsenic pentafluoride (from contamination of the HF). These contaminants are particularly problematic for the semiconductor industry, which uses high-purity fluorine for etching silicon wafers. Further impurities, such as iron, nickel, gallium and tungsten compounds, can be introduced if unreacted fluorine is recycled.

The technical-grade fluorine is purified by reacting it with MnF₃ to form manganese tetrafluoride. As this stage, any heavy metals present will form involatile complex fluorides, while the HF and O₂ are unreactive. Once the MnF₃ has been converted, the excess gas is vented for recycling, carrying the remaining gaseous impurities with it. The MnF₄ is then heated to 380 °C to release fluorine at purities of up to 99.95%, reforming MnF₃, which can be reused. By placing two reactors in parallel, the purification process can be made continuous, with one reactor taking in technical fluorine while the other delivers high-grade fluorine. Alternatively, the manganese tetrafluoride can be isolated and transported to where the fluorine is needed, at lower cost and greater safety than pressurized fluorine gas.

Fluoromanganate(IV) complexes

The yellow hexafluoromanganate(2−) of alkali metal and alkaline earth metal cations have been known since 1899, and can be prepared by the fluorination of MnF₂ in the presence of the fluoride of the appropriate cation. They are much more stable than manganese tetrafluoride. Potassium hexafluoromanganate(IV), K₂MnF₆, can also be prepared by the controlled reduction of potassium permanganate in 50% aqueous hydrofluoric acid.

2 KMnO₄ + 2 KF + 10 HF + 3 H₂O₂ → 2 K₂MnF₆ + 8 H₂O + 3 O₂

The pentafluoromanganate(1−) salts of potassium, rubidium and caesium, MMnF₅, can be prepared by fluorination of MMnF₃ or by the reaction of [MnF₄(py)(H₂O)] with MF. The lemon-yellow heptafluoromanganate(3−) salts of the same metals, M₃MnF₇, have also been prepared.

When potassium hexafluoromanganate is doped into potassium fluorosilicate it forms a narrow band red phosphor.

Notes and references

Further reading

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