Fluorine (F): Properties & Uses

 

Fluorine is the lightest member of the halogen family in the periodic table with atomic no. 9 and symbol F. It is the most electronegative element in the periodic table. Therefore, it can react with almost every other element present in the periodic table, except argon, neon, and helium. The existence of fluorine was first detected in the early sixteenth century. The name fluorine comes from the Latin word “fluere,” which means to flow as it was used in its molten state for smelting and welding purposes during that time. Although the people of the sixteenth-century were well aware of the toxic properties of fluorine, they still managed to employ it for commercial applications. Since then, from providing us with a non-stick frypan to keeping the astronauts safe in outer space, scientists have made several uses of fluorine’s extreme reactivity. Many of us also have fluorine to thank for our shining, bright smiles.

Discovery of Fluorine

In 1530, a German meteorologist, Georgius Agricola, found a crystalline material with a remarkable property, which let the ores to melt at low temperatures. That might be the first time in history that humankind became aware of fluorine. This mineral is now known as fluorite, fluorspar, or Calcium fluoride. It took scientists over 200 hundred years to isolate fluorine in its elemental form. During this time, many people got injured, and some even lost their lives to isolate it.

As per records, the first commercial application of fluorine goes back to 1670, when a german glass worker used fluorspar to engrave designs into his glasses. Later in 1771, a German-Swedish chemist, Carl Wilhelm Scheele, found corrosion in one of those glasses. He concluded that there is some specific acid in the fluorspar and named it fluoric acid, now known as hydrofluoric acid (HF). After that, in 1810, André-Marie Ampère found that hydrofluoric acid had an identical similarity to hydrochloric acid. The quest for isolated fluorine got its first potential approach in the nineteenth century when a British scientist, Sir Humphry Davy, conducted electrolysis experiments to test ampere’s theory. During the electrolysis of hydrogen fluoride produced by the distillation of fluorspar, Sir Davy found that a combustible gas was evolving at the anode, and a brown powder was accumulated at the cathode. He postulated that fluorine might have combined with the platinum cathode. However, he was not successful in isolating fluorine. The high reactivity of fluorine posed a major obstacle in choosing a suitable electrolyte to isolate it. In 1886, a French chemist, Ferdinand Frederic Henri Moissan, decided to use dry potassium acid fluoride, {KHF}_{2}, prepared by dissolving perfectly dry potassium fluoride in anhydrous hydrofluoric acid, as an electrolyte. It proved to be a good conductor of electricity. To overcome the extreme reactivity of fluorine with platinum, he used an alloy of platinum and iridium that was reasonably inert to attack by fluorine. Moissan finally succeeded in his quest to isolate fluorine on 26 June 1886, when he observed the formation of a gas at the anode, which burst into flames when tested with silicon. He was awarded the Nobel Prize for this discovery.

Isotopes and Occurrence

Fluorine exists in nineteen different isotopes out of which only one is stable, _{ }^{ 17 }{ F }. Whereas,  _{ }^{ 14-16 }{ F }, _{ }^{ 18-29 }{ F }, and _{ }^{ 31 }{ F } are radioactive isotopes of fluorine. _{ }^{ 18 }{ F }  is the most stable radioisotope of fluorine, with a half-life of  109.77 minutes. Due to its high reactivity, the natural occurrence of fluorine is in compound form. It is the thirteenth most abundant element on earth and 24th among the other elements in the universe. In other words, { 4 }×{ 10 }^{ -5  }% of the universe is made up of fluorine, while 0.054% of the earth constitutes fluorine. Unlike many other elements present in the universe, which depends majorly on stellar nucleosynthesis for their formation and continued survival, fluorine owes its existence to a ghostly particle named neutrinos. However, the relatively less abundance of fluorine from its neighbouring elements nitrogen, oxygen, carbon, and neon in the universe, suggests that it is not created much by the stellar nucleosynthesis processes. The little amount of fluorine formed in a star gets destroyed by the most common elements found in stars, hydrogen, and helium. In a star’s hot interior, hydrogen–nuclei–proton split fluorine into oxygen and helium, while helium-4 nuclei convert it into neon. However, trace amounts of fluorine still manage to escape from the star by stellar winds. There are mainly three proposed factories of fluorine in the universe.

• Type 2 Supernovae.

• Wolf-Rayet Stars.

• Asymptotic Gaint Branch (AGB) stars of intermediate-mass range.

When a massive star goes supernova (i.e., it explodes), it creates around { 10 }^{ 58 } high-energy neutrinos. Although neutrinos aren’t harmful, trillions of them pass through us each second, in a supernovae explosion, these neutrinos travel at super-high speeds and have very high energy, which can be harmful. In other words, if our sun goes supernovae explosion, the entire human population will be wiped out by the neutrino storm alone. These massive stars have an enormous amount of neon, _{ }^{ 20 }{ Ne } before they explode. During the explosion, these highspeed neutrinos would either knock off a proton to make fluorine-19 or remove a neutron to make radioactive neon-19, which decays into fluorine-19. In Wolf-Rayet stars (blue stars with mass 40 times that of the sun), the nucleosynthesis for fluorine production involves helium burning. Whereas in AGB stars, it involves both hydrogens to helium conversion and helium burning to produce fluorine. Ironically, helium and hydrogen are also responsible for its destruction in the Wolf-Rayet stars. However, in such massive stars, fluorine can be rescued by the solar wind that can carry it through the hydrogen envelope to outer space. The same phenomenon also helps fluorine to come out of AGB stars, except in this case, the process is carried out by convection.

In earth’s atmosphere, fluorine never occurs in its elemental form due to its high reactivity. Industrial dependence for fluorine is fulfilled mainly by three minerals.

• Fluorite ({ CaF }_{ 2 }), also known as Fluorspar, Calcium flouride.

• Fluorapatite ({ Ca }_{ 5 }({ PO }_{ 4 })_{ 3 }{ F }).

• Cryolite ({ Na }_{ 3 }{ AlF }_{ 6 }).

Among these, fluorapatite is the most abundant source of fluorine in earth’s crust. Moreover, the oceanic presence of fluorine is very rare due to the insolubility of its compounds and salts.

Properties of Fluorine

Physical Properties

Fluorine is a pale yellow colour diatomic gas at room temperature. It has an atomic no. 9 and atomic mass of 18.99u. It changes its state from gas to liquid when condensed at the temperature of -188.11°C (boiling point), and from liquid to solid at -219.67°C (melting point). Fluorine in its gaseous form has a density of 0.001553 gram per cubic cm (three times denser than air) and has a pungent odour like other halogen gases. If we decrease the temperature beyond the melting point, at -220°C, fluorine will form a transparent, disordered cubic crystal structure named β-fluorine. Lowering the temperature further, up to −228°C, it will form an opaque crystal structure named α-fluorine. Fluorine is the smallest ion with an ionic radius of 136 pm. Fluorine is poorly soluble in liquid hydrogen fluoride, with a solubility of { 2.5 }×{ 10 }^{ -3  }g per 100 g of HF at –70°C and { -4 }×{ 10 }^{ -3  }g at –20°C. In liquid form, it is freely soluble in liquid oxygen and ozone.

Chemical Properties

The ground state electronic configuration of fluorine is {1s}^{ 2 }{ 2s }^{ 2 }{ 2p }^{ 5 }. The reason behind the highest electronegativity of fluorine is its electronic configuration. The outermost p orbital has only five electrons, and the optimal configuration for p orbital is six. Therefore, it is one electron less than stable configuration. This lack of electron can easily be filled in the case of fluorine, as there are not many shells in between the valance shell and nucleus. In simple terms, the atomic radius of fluorine is so small that the outer electrons are inefficient to provide any effective electronic shielding. Hence, it has the highest electronegativity. The strong nuclear pull due to high effective nuclear charge can also be the reason for its third-highest ionization energy, after helium and neon.

Fluorine is known to be a special element in the periodic table as it is one of the few elements that can lure the noble gas elements, such as xenon, krypton, and radon, to make compounds with it. Although fluorine is the most vicious element of the halogen family, its reactivity makes it a useful element also. The reactivity of fluorine owes everything to its oxidizing nature, which frequently leads to a sudden or explosive reaction. Due to its smaller atomic structure, the interatomic repulsion among the valance shell electrons of difluorine is very high, as a result, the bond energy of { F }{ 2 } is much lower than that of { Cl }_{ 2 } and { Br }_{ 2 }. In simple words, it is easier to break a bond between { F }_{ 2 } molecule in comparison to other halogen-halogen molecules. This property accounts for fluorine’s high reactivity and strong bonds with non-fluoride elements.

Compounds of Fluorine

Fluorine’s high reactivity, electronegativity, and high electron affinity account for its extensive chemistry in the periodic table. Fluorine takes part in compound formation with almost all other elements except helium and neon. In most of its compounds, fluorine adopts the oxidation state of -1 and either forms polar covalent bonds or ionic bonds. Fluorine is only one electron less from an optimal electronic configuration, so in most of the reactions, it acts as a reducing agent and carries a net negative charge. For instance, hydrogen fluoride is classified among the most stable diatomic molecule due to its hydrogen bonding, with dissociation energy of 560kJ. In an aqueous solution, hydrofluoric acid is a weak acid, with an activity comparable to that of formic acid. It is characterized both by its water-like physical structure due to hydrogen bonding (regardless of its phases – solid, liquid, gas) and by its high reactivity and toxicity. For industry-centric manufacturing, fluorspar or calcium difluoride is treated with sulphuric acid to form hydrogen fluoride.

{ CaF }_{ 2 } + { H }_{ 2 }{ SO }_{ 4 } →  { 2 }{ HF } + { Ca }{ SO }_{ 4 }

Fluorine makes ionic bonds in case of alkali metals and alkaline earth metals. Nevertheless, alkali metals form highly soluble monofluorides, whereas alkaline earth metals form insoluble difluorides, except for { Be }{ F }_ { 2 }. In the case of beryllium fluoride, the hydration energy released during its dissolution overcomes the lattice energy, and therefore, makes it exceptionally soluble in water.

 

Fluorine shows variable chemistry throughout the periodic table. For instance, it can form trifluorides, tetra fluorides, pentafluorides, hexafluorides, and even heptafluorides, with several d-block elements. In the case of p-block elements, fluorine mostly make covalent bonds with other elements. In chalcogens, fluorine compounds of sulphur and oxygen have special characteristics. Oxygen forms an unstable difluoride, in which it has an exceptional oxidation state of +2, while sulphur forms a chemically inert hexafluoride with fluorine ({ S }{ F }_{ 6 }). In its own family, fluorine can form mono, tri, and pentafluorides with chlorine, bromine, and iodine. However, there is only one possible heptafluoride in the halogen family ({ I }{ F }_{ 7 }). Fluorine tends to form hypervalent molecules with heavier p-block elements. Conceivably, the most interesting fact about fluorine compounds is noble gas fluorides. Fluorine was the first element in history to form a compound with noble gases.

Fluorine forms the strongest bond with carbon when we talk about organic compounds. There is no known evidence that organofluorines exist naturally. However, they can be prepared artificially in labs. Fluorine provides extra stability to organic compounds by replacing hydrogen and increasing their melting point, and therefore, helps in forming fluoropolymers that have high resistance to solvents, acids, and bases.

Uses of Fluorine

The high reactivity of fluorine has made it one of the most employable elements of the periodic table. We daily come across several household items that contain fluorine. From welding purposes to make non-stick pans, humankind has faced several challenges while working with fluorine. Regardless of its toxicity, fluorine has provided exciting alchemy for the betterment of humanity. Let’s discuss its role in various sectors.

Industrial Applications

The worldwide demand for the elemental fluorine is increasing on the back of the various advantages it has over other elements. Commercial grade calcium fluoride is used to make over 17000 tonnes of fluorine every year. The price of elemental fluorine is $5 to $8 per kg. However, the main challenge with fluorine is not its production but its transportation and storage. Due to such problems, the price become much higher when it is supplied in cylinders. To tackle these problems, many industries employ in situ production under vertical integration, in which units for elemental fluorine production are integrated into { SF }_{ 6 } or { UF }_{ 6 } units.

1. Optics

The ancient known source of fluorine, Calcium fluoride, has a very remarkable property. Calcium fluoride is transparent not only to visible light but also to the ultraviolet and infrared wavelengths of light falling in the region (about 0.15 µm to 9 µm) and exhibits an extremely low change in refractive index. This property of calcium fluoride made it employable in the manufacture of semiconductor steppers, a lithographic device used for making integrated circuits, as the refractive index of { Ca }{ F }_{ 2 } shows non-linearity to the wavelength of 157 nm, even at high power densities. However, in the early years of the 21st century, the stepper market for calcium fluoride collapsed, and many large manufacturing facilities were been closed. Nowadays, newer glasses and computer-aided designs are used for this purpose.

Flourite has unique optical transparency at shorter wavelengths (about the order of few nanometers). Well known lense companies, such as Nikon, Olympus, Carl Zeiss, and Leica, make use of fluorite to make objective lenses. Their transparency to ultraviolet light enables them to be used for fluorescence microscopy and in the production of ultraviolet images.

2. Fluorescence

Another source of fluorine, Fluorapatite ({ Ca }_{ 5 }({ PO }_{ 4 })_{ 3 }{ F }), can also be produced synthetically. When doped with manganese-II and antimony-V, and when irradiated by the resonated mercury radiations, synthetic fluorapatite shows fluorescence, and therefore, it is used to make fluorescent tubes. The well-known Warm white, White, and Daylight tubes, generally referred to as “halophosphors,” are based on this property.

These tubes can generally be seen at food markets or art studios to provide special lighting effects.

Fluorapatite is also used as a gemstone.

3. Metal Industry

Fluorine compounds can easily be found in several metal industries. For instance, alumina (aluminium oxide, Al2O3) can be reduced to metallic aluminium by electrolysis when fused with a flux consisting of sodium fluoroaluminate ({Na}_{ 3 }{ Al }{ F }_{ 6 }), usually called cryolite. The metallurgical grade calcium fluoride (<97% CaF2, also called metspar) is used for iron and steel casting and steel making to lower the melting point and increase the fluidity of the slag.

4. Refrigerant 

Fluorine Refrigerants (also known as synthetic refrigerants, fluorocarbon refrigerants, and halocarbon refrigerants) are a wide group of products most commonly known as CFCs, HCFCs, HFCs. CFCs are thermodynamically efficient and effective across a wide range of applications & operating temperatures. However, environmental concerns relating to ozone depletion due to the homolytic cleavage of the carbon-chlorine bonds.

6. Fluorosurfactants

Fluorosurfactants are the fluorine compound that lower the surface tension (or interfacial tension) between two liquids, between a gas and a liquid, or between a liquid and a solid. The word surfactant is a shorthand notation of “surface-active-agents.” Fluorosurfactants have a wide range of applications. They can act as detergents, wetting agents, emulsifiers, foaming agents, or dispersants. The most common use of fluorosurfactants is in paints. However, due to health concerns, fluorine-based surfactants are facing a downfall in the market due to their replacement by other eco-friendly substitutes, such as hydrocarbon-based surfactants.

7. Agrochemicals

During World War II, fluoro-DDT or “Gix” was used for the control of insects of medical concerns. More recently, fluoroacetamide and analogues have been used as systemic insecticides, and a large variety of other fluorinated organic compounds have shown insecticidal activity. Sulfuryl fluoride has recently been marketed as a fumigant for household and structural pests. Cryolite was first registered as a pesticide in 1957. It is used on many fruits, vegetables, and ornamental crops to protect against leaf-eating pests.

Medical Applications

1. Pharmaceuticals

When it comes to the pharmaceutical industry, one-fifth of the drugs contain some amount of fluoride in them. A fluorine atom in a bioactive material may simulate a hydrogen atom, and although this does not prevent metabolic processes from occurring, the end products may be ineffective or toxic. Some organofluorides are highly lipophilic; they increase the lipid solubility of pharmaceuticals, and thus, accelerates their absorption and transport in a living organism. Some of the commonly known examples of fluorine-containing drugs are 5-fluorouracil, flunitrazepam (Rohypnol), fluoxetine (Prozac), paroxetine (Paxil), ciprofloxacin (Cipro), mefloquine, and fluconazole.

2. Inhaler Propellent

Fluorocarbons are also used as a propellant for metered-dose inhalers used to administer some asthma medications. The current generation of propellant consists of hydrofluoroalkanes (HFA), which has replaced CFC-propellant-based inhalers. CFC inhalers were banned in 2008 as part of the Montreal Protocol because of environmental concerns of ozone layer depletion.

3. Positron Emission Tomography

Positron Emission Tomography (PET) is a specialized imaging method that can quantify functions of internal organs and biochemical and physiologic responses. PET offers unique and important information about disease processes and is useful for making the diagnosis and prognosis of the disease and analyzing treatment._{ }^{ 18 }{ F } is the most commonly used radionuclide compound for PET scanning. It is generally known as 2-deoxy-2-_{ }^{ 18 }{ F }-fluoro-β-D-glucose (¹⁸F-FDG).

4. Dental Care

 

Dental caries can be prevented with the help of fluorine therapy. The regular use of fluorides, such as in toothpaste and drinking water, is extremely effective in preventing dental caries. However, it is often advised to use these fluorides more frequently in low concentrations rather than using them in high concentrations in low frequency. Fluoride ions (F–) replace hydroxyl groups (OH–) in the formation of the apatite crystal lattice, resulting in a stronger, fluoridated tooth mineral (fluorapatite). Since fluorapatite is less soluble than hydroxyapatite even under acidic conditions, it is also more resistant to subsequent demineralization when acid is produced by the bacteria.

5. Oxygen Carrier

The concept of land mammals breathing liquids was first introduced into the medical literature as an experiment. In a pioneering study, it was shown that mice with fluid-filled lungs could survive up to 4 hours when submerged at pressures up to 160 mmHg. Then, in the early 1960s, it was discovered that fluorocarbon, a substance first produced during World War 2 as part of the Manhattan Project, is an excellent carrier of oxygen. In fact, one such perfluorocarbon Oxcyte has been through initial clinical trials.

Miscellaneous Applications

1. The Manhattan Project (role of fluorine in the quest of the atomic bomb)

Even after the successful isolation of fluorine in 1866, it was not until the Second World War when scientists could easily research fluorides due to their high reactivity. The first large-scale production of fluorine was carried out for the atomic bomb Manhattan Project, which needed { UF }_{ 6 } as a gaseous carrier of uranium to separate the _{ }^{235}{ U }  and _{ }^{238}{ U } isotopes of uranium. Uranium tetra-fluoride (UF4), which was prepared from uranium dioxide (UO2) and hydrogen fluoride (HF), was converted to UF6 by reaction with fluorine.

2. Teflon

Teflon is a common brand name given to a fluoropolymer of tetrafluoroethylene, Polytetrafluoroethylene (PTFE). It has a wide range of applications from household items to aerospace engineering. For instance, the most common application of Teflon can be seen in a non-stick pan, as it is hydrophobic.