He hydrofluoric acid (HF) is an aqueous solution in which hydrogen fluoride is dissolved. This acid is obtained mainly from the reaction of the concentrated sulfuric acid with the mineral fluorite (CaF 2 ). The mineral is degraded by the action of the acid and the remaining water dissolves the gases of hydrogen fluoride.
From this same acidic water, the pure product, that is, the hydrogen fluoride anhydride, can be distilled. Depending on the quantities of the dissolved gas, different concentrations are obtained and, therefore, several hydrofluoric acid products are available in the market.
At a concentration of less than 40%, it has a crystalline appearance indistinguishable from water, but at higher concentrations it emits white vapors of hydrogen fluoride. Hydrofluoric acid is known as one of the most aggressive and dangerous chemicals.
It is capable of"eating"almost any material with which it has contact: from glasses, ceramics and metals, to rocks and concretes. In what container is it then stored? In plastic bottles, synthetic polymers inert to their action.
Index
- 1 Formula
- 2 Structure
- 3 Properties
- 3.1 Reactivity
- 4 Applications
- 5 References
Formula
The formula of hydrogen fluoride is HF, but that of hydrofluoric acid is represented in an aqueous medium, HF (ac), to differentiate from the first.
Thus, hydrofluoric acid can be considered as the hydrate of hydrogen fluoride, and this results in its anhydride.
Structure
All acid in water has the ability to generate ions in an equilibrium reaction. In the case of hydrofluoric acid, it is estimated that in a solution there is the ion pair H 3 OR + and F - .
The anion F - probably forms a very strong hydrogen bridge with one of the cation's hydrogens (F-H-O + See the man page trans(1) for more information. 2 ). This explains why hydrofluoric acid is a weak Bronsted acid (proton donor, H + ), despite its high and dangerous reactivity; that is, in water it does not release so many H + compared to other acids (HCl, HBr or HI).
However, in concentrated hydrofluoric acid the interactions between the hydrogen fluoride molecules are effective enough to allow them to escape in the gas phase.
That is, within the water they can interact as if they were in the liquid anhydride, forming hydrogen bridges between them. These hydrogen bonds can be assimilated as almost linear chains (H-F-H-F-H-F-...) surrounded by water.
In the upper image, the non-shared pair of electrons oriented in the opposite direction of the bond (H-F:) interacts with another HF molecule to assemble the chain.
Properties
Because hydrofluoric acid is an aqueous solution, its properties depend on the concentration of the anhydride dissolved in water. The HF is very soluble in water and is hygroscopic, being able to produce a variety of solutions: from very concentrated (smoky and with yellow tones) to very diluted.
As its concentration decreases, HF (ac) adopts properties more similar to pure water than those of anhydride. However, hydrogen bonds H-F-H are stronger than those in water, H 2 OH OH.
Both coexist in harmony in the solutions, raising the boiling points (up to 105ºC). Likewise, the densities increase as more HF anhydride is dissolved. Of rest, all the solutions of HF (ac) have strong and irritating odors and are colorless.
Reactivity
So, what is the corrosive behavior of hydrofluoric acid? The answer lies in the H-F bond and the ability of the fluorine atom to form very stable covalent bonds.
Being fluorine a very small and electronegative atom, it is a powerful Lewis acid. That is, it is separated from hydrogen to bind to species that offer more electrons at a low energy cost. For example, these species can be metals, such as the silicon present in the glasses.
SiO 2 + 4 HF → SiF 4 (g) + 2 H 2 OR
SiO 2 + 6 HF → H 2 SiF 6 + 2 H 2 OR
If the dissociation energy of the H-F bond is high (574 kJ / mol), why does it break in the reactions? The answer has kinetic, structural and energetic nuances. In general, the less reactive the resulting product, the more favored is its formation.
What happens with the F - in water? In concentrated hydrofluoric acid solutions, another HF molecule can form a hydrogen bond with the F - of the pair [H 3 OR + F - ]
This results in the generation of difluoride ion [FHF] - , which is extraordinarily acidic. That is why all physical contact with this is extremely harmful. The slightest exposure can trigger an infinity of damage to the body.
There are many safety rules and protocols for proper handling, and thus prevent potential accidents to those who operate with this acid.
Applications
It is a compound with numerous applications in the industry, in research and in the work of consumers.
- Hydrofluoric acid generates organic derivatives that are involved in the process of purification of aluminum.
- Used in the separation of isotopes from uranium, as in the case of uranium hexafluoride (UF) 6 ). Likewise, it is used in the extraction, processing and refining of metals, rocks and oils, also used for the inhibition of growth and mold removal.
- The corrosive properties of the acid have been used to carve and burn crystals, especially frosted ones, by means of the etching technique.
- It is used in the manufacture of silicon semiconductors, with multiple uses in the development of computing and information technology, responsible for human development.
- It is used in the automotive industry as a cleaner, being used as a mold remover in ceramics.
- In addition to serving as an intermediary in some chemical reactions, hydrofluoric acid is used in some ion exchangers that are involved in the purification of metals and more complex substances.
- Participates in the processing of petroleum and its derivatives, which has allowed the obtaining of solvents for use in the manufacture of products for the cleaning and elimination of fats.
- It is used in the generation of agents for plating and surface treatment.
- Consumers use numerous products in which hydrofluoric acid has participated in its elaboration; for example, some needed for car care, cleaning products for furniture, electrical and electronic components, and fuels, among other products.
References
- PubChem. (2018). Hydrofluoric Acid. Retrieved on April 3, 2018, from: pubchem.ncbi.nlm.nih.gov.
- Kat Day. (April 16, 2013). The acid that really does eat through everything. Retrieved on April 3, 2018, from: chronicleflask.com
- Wikipedia. (March 28, 2018). Hydrofluoric acid. Retrieved on April 3, 2018, from: en.wikipedia.org.
- Shiver & Atkins. (2008). Inorganic chemistry. (Fourth edition, pages 129, 207-249, 349, 407). Mc Graw Hill.
- Hydrofluoric Acid. Musc. Medical University of South Carolina. Retrieved on April 3, 2018, from: academicdepartments.musc.edu