Fluorine: History, Properties, Structure, Obtaining, Risks, Uses

The fluorine is a chemical element with symbol F and 17 leads the group, to which belong the halogens. It is distinguished above the other elements of the periodic table, for being the most reactive and electronegative; It reacts with almost all atoms, so it forms an infinite number of salts and organofluorinated compounds.

Under normal conditions it is a pale yellow gas, which can be confused with yellowish green. In liquid state , shown in the image below, its yellow color intensifies a little more, which disappears completely when it solidifies at its freezing point.

Such is its reactivity, despite the volatile nature of its gas, that it remains trapped in the earth’s crust; especially in the form of the mineral fluorite, known for its violet crystals. Also, its reactivity makes it a potentially dangerous substance; reacts vigorously to everything it touches and burns in flames.

However, many of its by-products can be harmless and even beneficial, depending on their applications. For example, the most popular use of fluoride, added in its ionic or mineral form (such as fluoride salts), is the preparation of fluoride toothpastes, which help protect tooth enamel.

Fluorine has the peculiarity that it can stabilize the high numbers or oxidation states for many other elements. The higher the number of fluorine atoms, the more reactive the compound (unless it is a polymer). Likewise, its effects with molecular matrices will increase; for better or worse.


Use of fluorite

In 1530, the German mineralogist Georgius Agricola discovered that the mineral fluorspar could be used in the purification of metals. Fluorspar is another name for fluorite, a fluorine mineral that consisted of calcium fluoride (CaF 2 ).

The element fluorine had not been discovered by then and the “fluoir” in fluorite came from the Latin word “fluere” which means “to flow”; since, this was precisely what fluorspar or fluorite did with metals: it helped them to leave the sample.

Preparation of hydrofluoric acid

In 1764, Andreas Sigismud Margraff succeeded in preparing hydrofluoric acid, heating fluorite with sulfuric acid. The glass retorts were melted by the action of the acid, so the glass was replaced by metals.

It is also attributed to Carl Scheele in 1771, the preparation of the acid by the same method followed by Margraff. In 1809, the French scientist Andre-Marie Ampere proposed that fluoric or hydrofluoric acid was a compound made up of hydrogen and a new element similar to chlorine.

Scientists tried to isolate fluoride by using hydrofluoric acid for a long time; but its dangerousness made progress in this sense difficult.

Humphry Davy, Joseph Louis Gay-Lussac and Jacques Thénard had severe pain when they inhaled hydrogen fluoride (hydrofluoric acid without water and in gaseous form). Scientists Paulin Louyet and Jerome Nickles died of poisoning under similar circumstances.

Edmond Frémy, a French researcher, tried to create dry hydrofluoric acid to avoid the toxicity of hydrogen fluoride by acidifying potassium bifluoride (KHF 2 ), but during electrolysis there was no conduction of electrical current.


In 1860, the English chemist George Gore attempted the electrolysis of dry hydrofluoric acid and succeeded in isolating a small amount of the fluorine gas. However, an explosion occurred as hydrogen and fluorine violently recombined. Gore attributed the explosion to an oxygen leak.

In 1886, the French chemist Henri Moisson succeeded in isolating fluorine for the first time. Previously, Moisson’s work was interrupted four times by severe hydrogen fluoride poisoning while attempting to isolate the element.

Moisson was a student of Frémy and relied on his experiments to isolate fluorine. Moisson used a mixture of potassium fluoride and hydrofluoric acid in the electrolysis. The resulting solution conducted electricity and fluorine gas collected at the anode; that is, at the positively charged electrode.

Moisson used corrosion resistant equipment, in which the electrodes were made of an alloy of platinum and iridium. In the electrolysis he used a platinum container and cooled the electrolyte solution to a temperature of -23ºF (-31ºC).

Finally, on June 26, 1886, Henri Moissson succeeded in isolating fluorine, work that allowed him to win the Nobel Prize in 1906.

Interest in fluoride

Interest in fluoride research was lost for a time. However, the development of the Manhattan Project for the production of the atomic bomb, drove it again.

The American company Dupont developed, between the years 1930 and 1940, fluorinated products such as chlorofluorocarbons (Freon-12), used as refrigerants; and polytetrafluoroethylene plastic, better known by the name Teflon. This produced an increase in the production and consumption of fluorine.

In 1986, at a conference to mark a century of the isolation of fluorine, the American chemist Karl O. Christe presented a chemical method for the preparation of fluorine by the reaction between K 2 MnF 6 and SbF 5 .

Physical and chemical properties


Fluorine is a pale yellow gas. In liquid state it is bright yellow. Meanwhile, the solid can be opaque (alpha) or transparent (beta).

Atomic number (Z)


Atomic weight

18,998 u.

Melting point

-219.67 ° C.

Boiling point

-188.11 ° C.


At room temperature: 1.696 g / L.

At melting point (liquid): 1.505 g / mL.

Heat of vaporization

6.51 kJ / mol.

Molar caloric capacity

31 J / (mol K).

Vapor pressure

At a temperature of 58 K it has a vapor pressure of 986.92 atm.

Thermal conductivity

0.0277 W / (m K)

Magnetic order



Characteristic pungent and pungent odor, detectable even at 20 ppb.

Oxidation numbers

-1, which corresponds to the fluoride anion, F .

Ionization energy

-First: 1,681 kJ / mol

-Second: 3,374 kJ / mol

-Third: 6.147 KJ / mol


3.98 on the Pauling scale.

It is the chemical element with the highest electronegativities; that is, it has a high affinity for the electrons of the atoms with which it bonds. Because of this, fluorine atoms generate large dipole moments in specific regions of a molecule.

Its electronegativity also has another effect: the atoms bound to it lose so much electron density that they begin to acquire a positive charge; this is, a positive oxidation number. The more fluorine atoms there are in a compound, the central atom will have a more positive oxidation number.

For example, in OF 2 oxygen has an oxidation number of +2 (O 2+ F 2 ); in UF 6 , uranium has an oxidation number of +6 (U 6+ F 6 ); the same happens with sulfur in SF 6 (S 6+ F 6 ); and finally there is AgF 2 , where silver even has an oxidation number of +2, rare for it.

Therefore, the elements manage to participate with their most positive oxidation numbers when they form compounds with fluorine.

Oxidizing agent

Fluorine is the most powerful oxidizing element, so no substance is capable of oxidizing it; and for this reason, it is not free in nature.


Fluorine is capable of combining with all other elements except helium, neon, and argon. It also does not attack mild steel or copper at normal temperatures. Reacts violently with organic materials such as rubber, wood, and fabric.

Fluorine can react with the noble gas xenon to form the strong oxidant xenon difluoride, XeF 2 . It also reacts with hydrogen to form a halide, hydrogen fluoride, HF. In turn, hydrogen fluoride dissolves in water to produce the famous hydrofluoric acid (as glass).

The acidity of the acidic acids, classified in increasing order is:

HF <HCl <HBr <HI

Nitric acid reacts with fluorine to form fluorine nitrate, FNO 3 . Meanwhile, hydrochloric acid reacts vigorously with fluorine to form HF, OF 2 and ClF 3 .

Structure and electronic configuration

Diatomic molecule

The fluorine atom in its ground state has seven valence electrons, which are in the 2s and 2p orbitals according to the electronic configuration:

[He] 2s 2 2p 5

The valence bond theory (TEV) states that two fluorine atoms, F, are covalently bonded to each complete its valence octet.

This happens quickly because it takes just one electron to become isoelectronic to the neon noble gas; and its atoms are very small, with a very strong effective nuclear charge that easily demands electrons from the environment.

The molecule F 2 (upper image), has a single covalent bond, FF. Despite its stability compared to free F atoms, it is a highly reactive molecule; homonuclear, apolar, and eager for electrons. That is why fluorine, like F 2 , is a very toxic and dangerous species.

Because F 2 is apolar, its interactions depend on its molecular mass and the London scattering forces. At some point, the electronic cloud around both F atoms must deform and give rise to an instantaneous dipole that induces another in a neighboring molecule; so that they attract each other slowly and weakly.

Liquid and solid

The F 2 molecule is very small and diffuses in space relatively quickly. In its gaseous phase, it exhibits a pale yellow color (which can be confused with a lime green). When the temperature drops to -188 ° C, the dispersion forces become more effective, causing the F 2 molecules to coalesce enough to define a liquid.

Liquid fluorine (first image) looks even more yellow than its respective gas. In it, the F 2 molecules are closer and interact with light to a greater degree. Interestingly, once the distorted cubic fluorine crystal is formed at -220 ° C, the color fades and remains as a transparent solid.

Now that the F 2 molecules are so close together (but without their molecular rotations stopping), it seems that their electrons gain a certain stability and, therefore, their electronic jump is too great for light to even interact with the crystal.

Crystalline phases

This cubic crystal corresponds to the β phase (it is not an allotrope because it remains the same F 2 ). When the temperature drops even further, down to -228 ºC, the solid fluorine undergoes a phase transition; the cubic crystal becomes a monoclinic one, the α phase:

Unlike β-F 2 , α-F 2 is opaque and hard. Perhaps it is because the F 2 molecules no longer have as much freedom to rotate in their fixed positions in monoclinic crystals; where they interact to a greater degree with light, but without exciting their electrons (which would superficially explain their opacity).

The crystal structure of α-F 2 was difficult to study by conventional X-ray diffraction methods. This is because the transition from the β to the α phase is highly exothermic; so the crystal practically exploded, at the same time that it interacted little with radiation.

It took about fifty years before German scientists (Florian Kraus et al.) Fully deciphered the structure of α-F 2 with greater precision thanks to neutron diffraction techniques.

Where to find and obtaining

Fluorine ranks 24th of the most common elements in the Universe. However, in the earth mass is 13 vo element, with a concentration of 950 ppm in the crust, and a concentration of 1.3 ppm in the seawater.

Soils have a fluoride concentration between 150 and 400 ppm, and in some soils the concentration can reach 1,000 ppm. In atmospheric air it is present in a concentration of 0.6 ppb; but up to 50 ppb has been recorded in some cities.

Fluorine is mainly obtained from three minerals: fluorite or fluorospar (CaF 2 ), fluoroapatite [Ca 5 (PO 4 ) 3 F] and cryolite (Na 3 AlF 6 ).

Fluorite Processing

After collecting the rocks with the mineral fluorite, they are subjected to a primary and secondary crushing. With secondary crushing very small rock fragments are obtained.

The rock fragments are then taken to a ball mill for reduction to powder. Water and reagents are added to form a paste, which is placed in a flotation tank. Air is injected under pressure to form bubbles, and thus the fluorite ends up floating on the aqueous surface.

The silicates and carbonates settle out while the fluorite is collected and taken to the drying ovens.

Once the fluorite is obtained, it is reacted with sulfuric acid to produce hydrogen fluoride:

CaF 2       + H 2 SO 4      => 2 HF + CaSO 4

Electrolysis of hydrogen fluoride

In the production of fluorine, the method used by Moisson in 1886 is followed, with some modifications.

An electrolysis is made of a mixture of molten potassium fluoride and hydrofluoric acid, with a molar ratio of 1: 2.0 to 1: 2.2. The temperature of the molten salt is 70-130 ° C.

The cathode consists of a Monel alloy or steel, and the anode is degraphite carbon. The fluorine production process during electrolysis can be outlined as follows:

2HF => H 2     + F 2

Water is used to cool the electrolysis chamber, but the temperature must be above the melting point of the electrolyte to avoid solidification. Hydrogen produced in electrolysis is collected at the cathode, while fluorine at the anode.


Fluorine has 18 isotopes, with 19 F being the only stable isotope with 100% abundance. The 18 F has a half life of 109.77 minutes and is the radioactive isotope of fluorine with the longer half – life. The 18 F is used as a source of positrons.

Biological role

There is no known metabolic activity of fluorine in mammals or higher plants. However, some plants and marine sponges synthesize monofluoroacetate, a poisonous compound, which they use as a protection to prevent its destruction.


Excessive consumption of fluoride has been associated with bone fluorosis in adults and dental fluorosis in children, as well as with alterations in kidney function. For this reason, The United States Public Health Service (PHS) suggested that the concentration of fluoride in drinking water should not be greater than 0.7 mg / L.

Meanwhile, The Us Enviromental Protection Agency (EPA) established that the concentration of fluoride in drinking water should not be greater than 4mg / L, in order to avoid skeletal fluorosis, in which fluoride accumulates in the bones. This can lead to bone weakening and fractures.

Fluoride has been associated with damage to the parathyroid gland, with a decrease in calcium in bone structures and high concentrations of calcium in plasma.

Among the alterations attributed to excess fluoride are the following: dental fluorosis, skeletal fluorosis, and damage to the parathyroid gland.

Dental fluorosis

Dental fluorosis occurs with small streaks or specks in the tooth enamel. Children under the age of 6 should not use mouthwashes that contain fluoride.

Skeletal fluorosis

In skeletal fluorosis, pain and damage to the bones, as well as the joints, can be diagnosed. The bone can harden and lose elasticity, increasing the risk of fractures.



We start with the section on the uses for fluoride with the one best known: that of serving as a component of many toothpastes. This is not the only use where the contrast between its extremely poisonous and dangerous molecule F 2 and the anion F is appreciated , which depending on its environment can be beneficial (although sometimes not).

When we eat food, especially sweets, bacteria break it down by increasing the acidity of our saliva. Then there comes a point where the pH is acidic enough to degrade and demineralize tooth enamel; hydroxyapatite breaks down.

However, in this process the F ions interact with the Ca 2+ to form a fluorapatite matrix; more stable and durable than hydroxyapatite. Or at least, this is the proposed mechanism to explain the action of the fluoride anion on teeth. It is likely to be more complex and to have a pH-dependent hydroxyapatite-fluorapatite balance.

These F anions are available in dental teeth in the form of salts; such as: NaF, SnF 2 (the famous stannous fluoride) and NaPOF. However, the concentration of F must be low (less than 0.2%), as otherwise it causes negative effects on the body.

Water fluoridation

Much like toothpaste, fluoride salts have been added to drinking water sources to combat cavities in those who drink it. The concentration should still be much lower (0.7 ppm). However, this practice is often the subject of mistrust and controversy, as it has been attributed possible carcinogenic effects.

Oxidizing agent

The F 2 gas behaves as a very strong oxidizing agent. This causes many compounds to burn more quickly than when exposed to oxygen and a heat source. That is why it has been used in rocket fuel mixtures, in which it can even replace ozone.


In many uses, the contributions of fluorine are not due to F 2 or F , but directly to their electronegative atoms as part of an organic compound. In essence, we are talking about a CF link.

Depending on the structure, polymers or fibers with CF bonds are usually hydrophobic, so they do not get wet or resist the attack of hydrofluoric acid; Or better yet, they can be excellent electrical insulators, and useful materials from which objects such as pipes and gaskets are made. Teflon and naphion are examples of these fluorinated polymers.


Fluorine’s reactivity makes its use for the synthesis of multiple inorganic or organic fluorine compounds questionable. In organics, specifically those with pharmacological effects, replacing one of their heteroatoms with F atoms increases (positively or negatively) their action on their biological target.

That is why in the pharmaceutical industry the modification of some drugs is always on the table by adding fluorine atoms.

Very similar happens with herbicides and fungicides. The fluoride in them can increase their action and effectiveness on insect and fungal pests.

Glass engraving

Hydrofluoric acid, due to its aggressiveness towards glass and ceramics, has been used to engrave thin and delicate pieces of these materials; usually destined for the manufacture of microcomponents of computers, or for electric bulbs.

Uranium enrichment

One of the most relevant uses of elemental fluorine is to help enrich uranium as 235 U. For this, uranium minerals are dissolved in hydrofluoric acid, producing UF 4 . This inorganic fluoride then reacts with F 2 , thus transforming into UF 6 ( 235 UF 6 and 238 UF 6 ).

Subsequently, and by gas centrifugation, the 235 UF 6 is separated from the 238 UF 6 to later be oxidized and stored as nuclear fuel.


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