POTASSIUM FLUORIDE (KF): STRUCTURE, PROPERTIES AND USES - CHEMISTRY - 2025

The potassium fluoride is an inorganic halide which is a salt formed between the metal and the halogen. Its chemical formula is KF, which means that for every K ⁺ cation there is a counterpart F ^-. As can be seen, the interactions are electrostatic, and as a consequence there are no covalent KF bonds.

This salt is characterized by its extreme solubility in water, which is why it forms hydrates, absorbs moisture and is deliquescent. Therefore, it is very easy to prepare aqueous solutions of it, which serve as a source of fluoride anions for all those syntheses where it is desired to incorporate it into some structure.

Potassium fluoride. Source: Gabriel Bolívar

Shown above are the K ⁺ cation (purple sphere) and the F ^- anion (bluish sphere). Both ions interact attracting each other by their +1 and -1 charges.

Although KF is not as dangerous as HF, the fact that it has "total freedom" to the anion F ^-, makes it a toxic salt. That is why their solutions have been used as insecticides.

KI is produced by reacting potassium carbonate with hydrofluoric acid, producing potassium bifluoride (KHF ₂); which by thermal decomposition ends up originating potassium fluoride.

Structure of potassium fluoride

Source: Cubic arrangement or gem salt for potassium fluoride. Benjah-bmm27, from Wikimedia Commons

The upper image shows the structure of potassium fluoride. The purple spheres, as in the first image, represent the K ⁺ cations; while the yellowish spheres represent the F ^- anions.

Note that the arrangement is cubic and corresponds to a rock-salt-like structure, very similar to that of sodium chloride. All the spheres are surrounded by six neighbors, which make up a KF ₆ or FK ₆ octahedron; that is, each K ⁺ is surrounded by six F ^-, and the same happens vice versa.

It was mentioned earlier that KF is hygroscopic and therefore absorbs moisture from the environment. Thus, the arrangement shown would correspond to the anhydrous form (without water) and not to its hydrates; which absorb so much water that they even become solubilized and "melt" (deliquescence).

Hydrates

The crystalline structures of hydrates become less simple. Why? Because now the water molecules intervene directly in the arrangements and interact with the K ⁺ and F ^- ions. Some of the most stable hydrates are KF · 2H ₂ O and KF · 4H ₂ O.

In both hydrates the octahedra just mentioned are deformed by the water molecules. This is mainly due to hydrogen bonds between F ^- and H ₂ O (F ^- -HOH). Crystallographic studies have determined that despite this the two ions continue to have the same number of neighbors.

As a result of all this, the original cubic structure for anhydrous potassium fluoride is transformed into a monoclinic and even rhombohedral arrangement.

Anhydrous ones share the deliquescent property, so their white crystals if left in contact with a cold mist would turn watery in a short time.

Properties

Molecular weight

58.097 g / mol.

Physical appearance (color)

White cubic crystals or deliquescent white crystalline powder.

Taste

Sharp salty taste.

Boiling point

2,741 ° F to 760 mmHg (1502 ° C). In the liquid state, it becomes a conductor of electricity, although the F ^- anions may not collaborate to the same degree as K ⁺.

Melting point

1,576 ° F; 858 ° C; 1131 K (anhydrous KF). This is indicative of its strong ionic bonds.

Solubility

Soluble in HF, but insoluble in alcohol. This shows that the hydrogen bonds between fluoride and alcohols, F ^- -HOR, do not favor the solvation process against the dissolution of their crystalline lattice.

Water solubility

Anhydrous 92 g / 100 ml (18 ° C); 102 g / 100 ml (25 ° C); dihydrate 349.3 g / 100 ml (18 ° C). That is, as KF is hydrated, it becomes more soluble in water.

Density

2.48 g / cm ³.

Vapor pressure

100 kPa (750 mmHg) at 1,499 ° C.

Decomposition

When heated to decomposition it emits a toxic smoke of potassium oxide and hydrogen fluoride.

Corrosive action

A watery solution corrodes glass and porcelain.

Flashpoint

It is not a flammable substance

Experimental refractive index (ηD)

1,363.

Stability

Stable if protected from moisture, otherwise the solid will dissolve. Incompatible with strong acids and bases.

Applications

Adjust pH

Aqueous solutions of potassium fluoride are used in industrial applications and processes; for example, KF solutions allow the pH to be adjusted in manufacturing in textile processing facilities and in laundries (close to a value of 7).

Fluorine source

Potassium fluoride is after hydrogen fluoride, the main source of obtaining fluorine. This element is used in nuclear plants and in the production of inorganic and organic compounds, some with uses such as its incorporation in toothpastes.

Synthesis of fluorocarbons

Potassium fluoride can be used in the synthesis of fluorocarbon or fluorocarbon from chlorocarbon, using the Finkeistein reaction. Ethylene glycol and dimethyl sulfoxide are used as solvents in this reaction.

Fluorination

As it is a source of fluorine where it lies dissolved in water, complex fluorides can be synthesized from its solutions; that is, an F ^- is incorporated into the structures. An example is in the following chemical equation:

MnBr ₂ (ac) + 3KF (ac) => KMnF ₃ (s) + 2KBr (ac)

The mixed KMnF ₃ fluoride then precipitates. Thus, F ^- could be added to make it part of a complex metal salt. Besides manganese, fluorides of other metals can be precipitated: KCoF ₃, KFeF ₃, KNiF ₃, KCuF ₃ and KZnF ₃.

Likewise, fluorine can be covalently incorporated into an aromatic ring, synthesizing organofluorides.

Various

KF is used as an intermediate or raw material for the synthesis of compounds that are mainly used in agrochemical or pesticide products.

In addition, it is used as a fluxing agent for welding and glass etching; that is, its aqueous solution eats away at the surface of the glass and, on a mold, prints the desired finish.

References

1. Chemical Book. (2017). Potassium fluoride. Recovered from: chemicalbook.com
2. PubChem. (2019). Potassium fluoride. Recovered from: pubchem.ncbi.nlm.nih.gov
3. TH Anderson and EC Lincafelte. (1951). The structure of potassium fluoride dihydrate. Acta Cryst. 4, 181.
4. Royal Society of Chemistry. (2015). Potassium fluoride. ChemSpider. Recovered from: chemspider.com
5. Maquimex. (sf). Potassium fluoride. Recovered from: maquimex.com