The fluorine It is a chemical element that is represented by the symbol F and heads group 17, to which halogens belong. It is distinguished above the other elements of the periodic table, for being the most reactive and electronegative; reacts with almost all atoms, so it forms countless 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..
Its reactivity is such, 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. Likewise, 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.
Article index
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 (CaFtwo).
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.
In 1764, Andreas Sigismud Margraff managed to prepare 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 a 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 hydrogen fluoride toxicity by acidifying potassium bifluoride (KHFtwo), but during the electrolysis there was no conduction of the electric 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 research was lost for a time. However, the development of the Manhattan Project for the production of the atomic bomb, boosted 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 fluoride..
In 1986, at a conference about a century on from the isolation of fluorine, the American chemist Karl O. Christe presented a chemical method for the preparation of fluorine by the reaction between KtwoMnF6 and the SbF5.
Fluorine is a pale yellow gas. In liquid state it is bright yellow. Meanwhile, the solid can be opaque (alpha) or transparent (beta).
9.
18,998 u.
-219.67 ºC.
-188.11 ºC.
At room temperature: 1.696 g / L.
At melting point (liquid): 1.505 g / mL.
6.51 kJ / mol.
31 J / (mol K).
At a temperature of 58 K it has a vapor pressure of 986.92 atm.
0.0277 W / (mK)
Diamagnetic
Characteristic pungent and pungent odor, detectable even at 20 ppb.
-1, which corresponds to the fluoride anion, F-.
-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 the OFtwo oxygen has an oxidation number of +2 (Otwo+Ftwo-); in the UF6, Uranium has an oxidation number of +6 (U6+F6-); the same happens with sulfur in SF6 (S6+F6-); and finally there is the AgFtwo, 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..
Fluorine is the most powerful oxidizing element, so no substance is capable of oxidizing it; and for this reason, it is not found 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, XeFtwo. 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, FNO3. Meanwhile, hydrochloric acid reacts vigorously with fluorine to form HF, OFtwo and ClF3.
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] 2stwo 2 P5
The valence bond theory (VTE) states that two fluorine atoms, F, are covalently bonded to each complete their 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 Ftwo (top image), has a single covalent bond, F-F. 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 Ftwo, it is a very toxic and dangerous species.
Because the Ftwo 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 originate an instantaneous dipole that induces another in a neighboring molecule; so that they attract each other slowly and weakly.
The molecule Ftwo it is very small and diffuses into space relatively quickly. In its gaseous phase, it exhibits a pale yellow color (which can be mistaken for a lime green). When the temperature drops to -188 ºC, the dispersion forces become more effective and make the F moleculestwo cohesive enough to define a liquid.
Liquid fluorine (first image) looks even more yellow than its respective gas. In it, the molecules Ftwo they 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 moleculestwo are so close (but without their molecular rotations stopping), it seems that their electrons gain a certain stability and, therefore, their electronic jump is too great for the light to even interact with the crystal.
This cubic crystal corresponds to the β phase (it is not an allotrope because it is still the same Ftwo). 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-Ftwo, the α-Ftwo it is opaque and hard. Perhaps it is because the F moleculestwo they no longer have so much freedom to rotate in their fixed positions on 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-Ftwo it 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 had little interaction with radiation.
It took about fifty years before German scientists (Florian Kraus et al.) Fully deciphered the structure of α-Ftwo more precisely thanks to neutron diffraction techniques.
Fluorine ranks 24th among the most common elements in the Universe. However, on Earth it is mass 13vo element, with a concentration of 950 ppm in the earth's crust, and a concentration of 1.3 ppm in 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 it has been recorded in some cities up to 50 ppb.
Fluorine is obtained mainly from three minerals: fluorite or fluorospar (CaFtwo), fluoroapatite [Ca5(PO4)3F] and cryolite (Na3AlF6).
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.
Silicates and carbonates settle out while fluorite is collected and taken to drying ovens..
Once the fluorite is obtained, it is reacted with sulfuric acid to produce hydrogen fluoride:
CaFtwo + HtwoSW4 => 2 HF + CaSO4
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 ºC.
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 => Htwo + Ftwo
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 the 19F the only stable isotope with 100% abundance. The 18F has a half-life of 109.77 minutes and is the radioactive isotope of fluorine with the longest half-life. The 18F is used as a source of positrons.
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 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 weakening of the bone 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 occurs with small streaks or specks in the tooth enamel. Children under the age of 6 should not use fluoride-containing mouthwashes..
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 begin 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 F molecule is appreciatedtwo, extremely poisonous and dangerous, and the anion F-, which depending on your 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 Catwo+ to form a fluorapatite matrix; more stable and durable than hydroxyapatite. Or at least, this is the proposed mechanism to explain the action of fluoride anion on teeth. It is likely to be more complex and to have a pH-dependent hydroxyapatite-fluorapatite balance.
These anions F- They are available in dental plants in the form of salts; such as: NaF, SnFtwo (the famous stannous fluoride) and NaPOF. However, the concentration of F- It must be low (less than 0.2%), since otherwise it causes negative effects on the body.
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 distrust and controversy, as it has been attributed possible carcinogenic effects.
Gas Ftwo it behaves as a very strong oxidizing agent. This causes many compounds to burn faster 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 the Ftwo or F-, but directly to its electronegative atoms as part of an organic compound. In essence, we speak of a C-F link..
Depending on the structure, polymers or fibers with C-F 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 nafion are examples of these fluorinated polymers.
The reactivity of fluorine makes its use for the synthesis of multiple fluorine compounds, organic or inorganic, 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 by adding fluorine atoms is always on the table..
Very similar happens with herbicides and fungicides. The fluoride in them can increase their action and effectiveness on insect and fungal pests..
Hydrofluoric acid, due to its aggressiveness with 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.
One of the most relevant uses of elemental fluorine is to help enrich uranium as 235U. For this, uranium minerals are dissolved in hydrofluoric acid, producing UF4. This inorganic fluoride then reacts with the Ftwo, to thus transform into UF6 (235PHEW6 Y 238PHEW6).
Subsequently, and by means of a gas centrifugation, the 235PHEW6 separates from 238PHEW6 to later be oxidized and stored as nuclear fuel.
Yet No Comments