Hosea

The fluorine is a chemical element of the family of the Halogène S of symbol F and of Atomic number 9.

With the Normal conditions of temperature and pressure the fluorine is present in the form of Difluor F₂ Gaz Diatomique yellow pale and toxic.

It is the chemical element most reactive. It has strongest electronegativity.

It causes burns in contact with the skin.

History

The fluorine (of Latin fluere meaning flow or melting ) is described by Georgius Agricola in 1530 in its form of Fluorine like a substance used to promote the fusion of metals or minerals.

In 1670, Georges Gore had observed that the fluorite (CaF₂, calcium fluoride known formerly under the name of emerald of Bohemian), treated by an acid, could engrave glass; Schwandhard used this property. Carl Scheele as well as other later researchers such as Davy Humphry, Gay-Lussac, Antoine Lavoisier, and the baron Louis Jacques Thénard did all of the experiments with hydrofluoric acid (hydrofluoric acid solution HF in water). Some of these experiments finished in tragedy because of the dangerosity of this product.

This element could not be insulated during many years, because, hardly separated, it immediately attacks the remainders of its compound. Mendeleïev placed it in its table in 1869, but it is only in 1886 that Henri Moissan, after 74 years of uninterrupted efforts, managed to prepare it by electrolysis fluoride of potassium in hydrofluoric acid, with electrodes in iridic Platine , under a tension of 50 volts. The pure difluor appeared with the positive pole and the Dihydrogène with cathode. Henri Moissan due to lead this experiment to low temperature, because the hydrofluoric acid (HF) boils with 19°C. This discovery to him been worth the Nobel Prize of Chemistry in 1906.

The first industrial production of difluor took place during the manufacture of the atomic bomb, within the framework of the Projet Manhattan at the time of the Second world war, where the Uranium hexafluoride UF₆, which is a volatile molecular compound, was used to separate the various isotopes from the Uranium by gas diffusion. This process besides is always implemented at the time of the manufacture of nuclear fuel used in the current nuclear plants.

Properties

The pure difluor is a corrosive pale yellow gas: it is an oxidizing powerful. It is most reactive and most electronegative of all the elements and forms compounds with the majority of the other elements, including the Rare gas S Xénon and Radon.

Even under basic conditions temperature and without light, the difluor reacts explosivement with the Dihydrogène, even in lower part of -250°C when the fluorine is solid and liquid hydrogen! In a gas difluor jet, the Glass, the metals, the Water and other substances burn with a luminous flame. The fluorine has such an affinity for the majority of the elements, in particular for the Silicium (If), that it can neither be prepared nor to be preserved in containers of glass (made up mainly of SiO₂ silica)

In solution, the fluorine forms Ion S F⁻ fluoride.

Preparation

The Difluor, F₂, is still produced today industrially thanks to the process of electrolysis introduces by Henri Moissan in 1886. The electrolytic bath consists of a mixture KF-2HF melted to 90-100°C approximately. Anhydrous HF is not conducting because little dissociated and it is the addition of KHF₂ which allows ionic conduction by a complex mechanism.

During the reaction of electrolysis, the difluor is produced on a carbon anode according to:

2 HF - → F (G) + 2 HF + 2nd -

With cathode (out of metal), Dihydrogène is produced:

4 HF + 2nd - → H (G) + 2 HF -

In the cell of electrolysis, the potential applied is included/understood between 8 and 10 V, and the density of current is about 12 has DM -2 . The output while running is good (95%), but the total energetic efficiency is only of 30%.

In 1986, at the time of the hundredth birthday of discovered electrochemical preparation of fluorine, Karl Christe discovered an original and purely chemical method of preparation while making react to 150 °C hydrofluoric acid anhydrous HF with K₂MnF₆ and SbF₅. The reaction is:

K₂MnF₆ + 2 SbF₅ → 2 KSbF₆ + MnF₃ + F₂

This process is anecdotic because not exploitable industrially.

Use

The difluor is too reactive for a direct use in a pure state. Its many chemical compounds has a multitude of applications on the other hand. Some examples:

the fluorine enters the composition of Plastics to low coefficient of friction the such Teflon (or Polytétrafluoréthylène (PTFE)) ;
many fluorinated Gases, such as for example the freon, is used as a Refrigerant in the systems of refrigeration and air conditioning. The Chlorofluorocarbone S (CFC) were however banished of these applications because of their probable contribution to the hole of the Couche of ozone and to the Greenhouse effect;
the hydrofluoric Acid (HF) with the single property to be able to dissolve all inorganic oxides almost. It is in fact used to attack glass, to eliminate oxides from surface of the Silicium in the industry of the Semi-conducteur S, like catalyst of the reactions of Alkylation of isobutane and butylene in the Raffinage of oil and to eliminate from the oxidized impurities of the Stainless steel;
In the Cycle of nuclear fuel, the Uranium hexafluoride UF₆ is used to separate the various isotopes from the Uranium by gas diffusion. It is produced in a Four with flame by chemical reaction of the difluor with the Tétrafluorure of uranium UF₄;
the Cryolite, a Mineral composed of fluoride of Sodium and Aluminum (NaAlF), is used as electrolyte in the production of Aluminum by electrolysis;
the monoatomic fluorine is used for “plasma ashing” in the manufacture of the Semi-conducteur S;
With others composed, the fluorine is used in the manufacture of more than one hundred of commercial fluorinated compounds, such as figure high temperature;
the fluoride of sodium was used like Insecticide, particularly against the Cafard S or as deratisor;
Of the Fluorure S is added to the Dentifrice S, or sometimes in certain sources of water or certain food to fight against the dental caries. This use is disputed by various associations and lobbies of consumers in the world.

Toxicity

The fluorine can be toxic for the Man:

< 1 mg/jour: protect from the decay
2 mg/jour: risk dental Fluorose
10 to 40 mg/jour: fluorose of the skeleton
20 to 80 mg/jour: fluorose ankylosante
100 mg/jour : delay of growth
125 mg/jour: renal deterioration
200 to 500 mg/jour: lethal amount

In France, the fluoridation of water is prohibited.

See too

Ion fluoride

External bonds

/Environmental Health Criteria for Fluorides (EHC 227) published in 2002 by the WHO (Summarized and conclusions in French)
Effects on health of fluorides - a summary for non-specialists of the report/ratio of WHO by GreenFacts

Document on the dangers of the flor

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