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About Fluorine

Fluorine Bohr

Though its primary ore, fluorite, has long been used in smelting to reduce the melting point of metal ores, fluorine was the last of the halogens to be isolated. Hydrofluoric acid was commonly used as a glass etching agent by the eighteenth century, but it was not until 1810 that Sir Humphry Davy proved that this acid was analogous to hydrochloric acid and therefore must contain an element similar to chlorine. However, the methods used to isolate chlorine were not successful in isolating fluorine, instead resulting only in frustration and sometimes the lethal hydrofluoric acid poisoning of many chemists who attempted to do so. Ferdinand Frederic Henry Moissan finally produced the pure element in 1886 with a method still in use today. The name for the element was derived from that of its ore, which itself had been derived from the latin fluere, meaning “to flow”--a reference to the effect the ore had on molten metal. Moissan was awarded the Nobel Prize in Chemistry for his achievement in 1906.

Fluorine’s earliest function, serving as a flux, remains vital in modern industry. Fluorine compounds are used in this capacity in steelmaking, welding, glassmaking, and in producing some forms of ceramics and cements. Additionally, the fluorine-containing compounds sodium hexafluoroaluminate and aluminum trifluoride are absolutely essential fluxes used in the extraction of aluminum from bauxite. Without the use of these compounds, the temperature required for this process would be too high for aluminum metal to be an economically viable industrial material.

Fluorine is also an important industrial tool in the form of hydrofluoric acid. Though considered a weak acid in dilute solution, as hydrogen and fluorine tend to remain bound rather than dissociating into ions under these conditions, hydrofluoric acid is highly corrosive. Its unusual ability to eat away even at glass led to its first use in glass etching. It is also used to etch the tough ceramics used in dental implants, as this improves the ability of the ceramic to be bound to other materials using adhesives. In steelmaking, hydrofluoric acid often serves as a pickling agent, removing oxides and other impurities on the surface of the metal. Hydrofluoric acid is also used in the production of extremely strong acids known as super acids, and in producing intermediate fluorine compounds used in the production of organochlorines and other fluorine products.

Organofluorine compounds are ubiquitous in modern industry. By volume of production, the most common type of organofluorine compounds are refrigerant gases. These include chlorofluorocarbons (CFCs), hydrochlorofluorocarbons (HCFCs), and hydrofluorocarbons (HFCs). These are used primarily for refrigeration and air conditioning, but are also found as aerosol propellants and solvents. CFCs were the first of these to be used, but were replaced by HCFCs when they were found to cause depletion of the ozone layer. Unfortunately HCFCs still cause ozone depletion, but to a lesser degree. HFCs cause little or no ozone depletion, and are therefore preferred to HCFCs when applicable. Another useful class of organofluorine compounds are the perfluorocarbons (PFCs). PFCs may be used as surfactants, or in the production of fluoropolymers such as Teflon. Teflon and related fluoropolymers are perhaps best known for use in non-stick coatings on cookware, but they are additionally used as insulating wire coating, friction reducing coatings on mechanical parts, and as Gore-Tex, a microporous polymer membrane that finds use in rain-repelling clothing, filters, and packing seals.

Additionally, fluorine-carbon bonds are often used in pharmaceuticals and to make organic agrichemicals such as pesticides more chemically stable. In drug design, this is often necessary because the non-fluorinated compound degrades too quickly in the body to have the desired effect. As such, many common pharmaceuticals are fluorinated, including SSRIs, quinolone antibiotics, and many steroids and anesthetics. However, the stability of organofluorine compounds is not always positive--these compounds tend to persist in the environment and to bioaccumulate. As many organofluorine agrichemicals and some common industrial organofluorines are toxic, this is a significant environmental concern.

Fluorine has a few other major uses in medicine. The isotope fluorine-18 is used as a tracer in positron emission tomography (PET). This type of medical imaging detects relative concentrations of a radioisotope that has been chemically bound to a biologically active molecule, most often glucose, to trace the accumulation of said compound in various body tissues. Glucose is used because more metabolically active areas of the body consume more glucose. In the brain, this is used to monitor which areas are most active, while in whole body scans this is typically used to detect cancer cells, which consume glucose at a disproportionately high rate compared to surrounding tissue. Additionally, fluoride compounds are used in toothpaste and added to drinking water, as the fluorine reacts with and hardens tooth enamel, reducing rates of cavities.

Due to the extreme reactivity of fluorine gas, most fluorine chemistry makes use of less-reactive intermediates rather than elemental fluorine, but the gas is produced for a few industrial processes. The vast majority of fluorine gas is used to produce either uranium or sulfur hexafluoride. Sulfur hexafluoride is used widely as a gaseous dielectric medium in high-voltage circuit breakers and other electrical equipment, often replacing old devices that contained harmful PCBs, while uranium hexafluoride is used in the production of nuclear fuels. The remainder of fluorine gas produced is used in the production of nitrogen fluoride, several metal fluorides, and fluorinating agents for use in organic synthesis. The gas nitrogen trifluoride is used in the cleaning of chambers used in the production of many electronics, in chemical lasers, and in plasma etching of silicon wafers. Rhenium hexafluoride and tungsten hexafluoride are used as precursor materials in chemical vapor deposition. Halogen fluorides such as chlorine and bromine trifluoride and iodine pentafluoride, as well as sulfur tetrafluoride, are common fluorinating agents used in the production of industrial fluorocarbons and fluorinated pharmaceuticals.

As the early chemists who worked with fluorine quickly learned, its unique properties present significant difficulties and dangers. Fluorine-fluorine bonds are relatively weak, while the bonds between it and other elements are usually very strong, due to fluorine’s high electronegativity; this combination makes fluorine gas extraordinarily reactive. It is corrosive to substances usually considered profoundly inert, including glass, and is therefore stored only in containers made from metals that acquire a passivating metal-fluorine compound coating upon contact with the gas--typically nickel or nickel-alloys. Hydrofluoric acid causes both severe burns and poisoning upon exposure, the latter occurring as it easily diffuses through skin and the fluorine quickly binds with essential ions. However, despite the risks and difficulties inherent to working with fluorine and some of its compounds, it remains a powerful chemical tool, and component of vital materials.

Fluorine is produced directly from the calcium fluoride mineral fluorite, sometimes also known as fluorspar. Lower grade deposits, termed “metspar”, are typically used directly in flux applications, while higher grade deposits known as “acid spar” are treated with sulfuric acid to produce hydrofluoric acid. This acid is then used in the production of virtually all other fluorine compounds. For the few applications in which elemental fluorine is required, it is still prepared via the original method devised in the nineteenth century: the electrolysis of a mixture of hydrogen fluoride and potassium fluoride, the latter being added to allow the mixture to conduct electricity. Additionally, fluorosilicic acid is a byproduct of phosphoric acid production, and this compound is increasingly being recovered as a secondary fluorine source.


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Fluorine Properties

Fluorine Element SymbolFluorine is a Block P, Group 17, Period 2 element. Its electron configuration is [He]2s22p5. The fluorine atom has a covalent radius of 64 pm and its Van der Waals radius is 135 pm. In its elemental form, CAS 7782-41-4, fluorine gas has a pale yellow appearance. Fluorine was discovered by André-Marie Ampère in 1810. Fluorine Bohr ModelIt was first isolated by Henri Moissan in 1886. Fluorine information, including technical data, properties, and other useful facts are specified below. Scientific facts such as the atomic structure, ionization energy, abundance on Earth, conductivity and thermal properties are included.

Symbol: F
Atomic Number: 9
Atomic Weight: 19
Element Category: halogen
Group, Period, Block: 17 (halogens), 2, p
Color: pale yellow
Other Names: Fluor, Fluoro
Melting Point: -219.67°C, -363.406°F, 53.48 K
Boiling Point: -188.11°C, -306.598°F, 85.04 K
Density: (0 °C, 101.325 kPa 1.696 3 g/L
Liquid Density @ Melting Point: 1.505 g·cm3
Density @ 20°C: 0.001696 g/cm3
Density of Solid: 1700 kg·m3
Specific Heat: N/A
Superconductivity Temperature: N/A
Triple Point: 53.48 K, 90 kPa
Critical Point: 144.41 K, 5.1724 MPa
Heat of Fusion (kJ·mol-1): 0.26 (per mol F atoms)
Heat of Vaporization (kJ·mol-1): 3.27 (per mole F atoms)
Heat of Atomization (kJ·mol-1): 76.9
Thermal Conductivity: 0.02591 W·m-1·K-1
Thermal Expansion: N/A
Electrical Resistivity: N/A
Tensile Strength: N/A
Molar Heat Capacity: (Cv) (21.1 °C) 610 J·mol-1·K-1, (Cp) (21.1 °C) 825 J·mol-1·K-1
Young's Modulus: N/A
Shear Modulus: N/A
Bulk Modulus: N/A
Poisson Ratio: N/A
Mohs Hardness: N/A
Vickers Hardness: N/A
Brinell Hardness: N/A
Speed of Sound: N/A
Pauling Electronegativity: 3.98
Sanderson Electronegativity: 4
Allred Rochow Electronegativity: 4.1
Mulliken-Jaffe Electronegativity: 3.91 (14.3% s orbital)
Allen Electronegativity: 4.193
Pauling Electropositivity: 0.02
Reflectivity (%): N/A
Refractive Index: 1.000195
Electrons: 9
Protons: 9
Neutrons: 10
Electron Configuration: [He]2s22p5
Atomic Radius: N/A
Atomic Radius,
non-bonded (Å):
Covalent Radius: 64 pm
Covalent Radius (Å): 0.6
Van der Waals Radius: 135 pm
Oxidation States: -1 (oxidizes oxygen)
Phase: gas
Crystal Structure: monoclinic
Magnetic Ordering: diamagnetic
Electron Affinity (kJ·mol-1) 328.147
1st Ionization Energy: 1,681 kJ·mol-1
2nd Ionization Energy: 3,374 kJ·mol-1
3rd Ionization Energy: 6,147 kJ·mol-1
CAS Number: 7782-41-4
EC Number: N/A
MDL Number: N/A
Beilstein Number: N/A
SMILES Identifier: [F]
InChI Identifier: InChI=1S/F
PubChem CID: 5360525
ChemSpider ID: 4514530
Earth - Total: 13.5 ppm
Mercury - Total: 2.2 ppm 
Venus - Total: 15 ppm
Earth - Seawater (Oceans), ppb by weight: 1300
Earth - Seawater (Oceans), ppb by atoms: 420
Earth -  Crust (Crustal Rocks), ppb by weight: 540000
Earth -  Crust (Crustal Rocks), ppb by atoms: 590000
Sun - Total, ppb by weight: 500
Sun - Total, ppb by atoms: 30
Stream, ppb by weight: 100
Stream, ppb by atoms: 5
Meterorite (Carbonaceous), ppb by weight: 89000
Meterorite (Carbonaceous), ppb by atoms: 96000
Typical Human Body, ppb by weight: N/A
Typical Human Body, ppb by atom: N/A
Universe, ppb by weight: N/A
Universe, ppb by atom: N/A
Discovered By: André-Marie Ampère
Discovery Date: 1810
First Isolation: Carl Wilhelm Scheele (1774)

Health, Safety & Transportation Information for Fluorine

Fluorine is a pale yellow gas with a pungent odor. It is commonly shipped as a cryogenic liquid. It is toxic by inhalation and skin absorption. Contact with skin in lower than lethal concentrations causes chemical burns. It reacts with water to form hydrofluoric acid and oxygen. It is corrosive to most common materials. It reacts with most combustible materials to the point that ignition occurs. Under prolonged exposure to fire or intense heat the containers may violently rupture and rocket. Safety data for Fluorine and its compounds can vary widely depending on the form. For potential hazard information, toxicity, and road, sea and air transportation limitations, such as DOT Hazard Class, DOT Number, EU Number, NFPA Health rating and RTECS Class, please see the specific material or compound referenced in the Products tab.

Safety Data
Material Safety Data Sheet MSDS
Signal Word Danger
Hazard Statements H280-H270-H330-H314-EUH071
Hazard Codes T+,C
Risk Codes 7-26-35
Safety Precautions 9-26-36/37/39-45
RTECS Number LM6475000
Transport Information UN 1045/9192
WGK Germany 3
Globally Harmonized System of
Classification and Labelling (GHS)
Skull and Crossbones-Acute Toxicity  Oxidizing Liquid - Oxidizing Gas Corrosion - Corrosive to metals Gas Cylinder - Gases Under Pressure

Fluorine Isotopes

Fluorine has one stable isotope: 19F.

Nuclide Isotopic Mass Half-Life Mode of Decay Nuclear Spin Magnetic Moment Binding Energy (MeV) Natural Abundance
(% by atom)
14F 14.03506(43)# N/A p to 13O 2-# N/A 70.4 -
15F 15.01801(14) 410(60)E-24 s [1.0(2) ] p to 14O (1/2+) N/A 94.31 -
16F 16.011466(9) 11(6)E-21 s [40(20) keV] p to 15O 0- N/A 108.91 -
17F 17.00209524(27) 64.49(16) s EC to 17O 5/2+ 4.722 125.38 -
18F 18.0009380(6) 109.771(20) min EC to 18O 1+ N/A 135.32 -
19F 18.99840322(7) STABLE - 1/2+ 2.628867 145.26 100
20F 19.99998132(8) 11.163(8) s ß- to 20Ne 2+ 2.094 152.41 -
21F 20.9999490(19) 4.158(20) s ß- to 21Ne; ß- + n to 20Ne 5/2+ N/A 160.49 -
22F 22.002999(13) 4.23(4) s ß- to 22Ne; ß- + n to 21Ne 4+,(3+) N/A 165.77 -
23F 23.00357(9) 2.23(14) s ß- to 23Ne; ß- + n to 22Ne (3/2,5/2)+ N/A 172.92 -
24F 24.00812(8) 400(50) ms ß- to 24Ne; ß- + n to 23Ne (1,2,3)+ N/A 176.34 -
25F 25.01210(11) 50(6) ms ß- to 25Ne; ß- + n to 24Ne (5/2+)# N/A 180.69 -
26F 26.01962(18) 9.6(8) ms ß- to 26Ne; ß- + n to 25Ne 1+ N/A 182.25 -
27F 27.02676(40) 4.9(2) ms ß- to 27Ne; ß- + n to 26Ne 5/2+# N/A 183.8 -
28F 28.03567(55)# <40 ns ß- to 28Ne; ß- + n to 27Ne N/A N/A 183.5 -
29F 29.04326(62)# 2.6(3) ms ß- to 29Ne; ß- + n to 28Ne 5/2+# N/A 184.12 -
30F 30.05250(64)# <260 ns Unknown N/A N/A 183.82 -
31F 31.06043(64)# 1# ms [>260 ns] Unknown 5/2+# N/A 184.44 -