| Discovery date | 1886 |
| Discovered by | Henri Moissan |
| Origin of the name | The name is derived form the Latin 'fluere', meaning to flow |
| Allotropes | F2 |
F
Fluorine
9
18.998
| Group | 17 | Melting point | −219.67°C, −363.41°F, 53.48 K |
| Period | 2 | Boiling point | −188.11°C, −306.6°F, 85.04 K |
| Block | p | Density (g cm−3) | 0.001553 |
| Atomic number | 9 | Relative atomic mass | 18.998 |
| State at 20°C | Gas | Key isotopes | 19F |
| Electron configuration | [He] 2s22p5 | CAS number | 7782-41-4 |
| ChemSpider ID | 4514530 | ChemSpider is a free chemical structure database | |
Image explanation
The image reflects the highly reactive nature of the element.
Appearance
A very pale yellow-green, dangerously reactive gas. It is the most reactive of all the elements and quickly attacks all metals. Steel wool bursts into flames when exposed to fluorine.
Uses
There was no commercial production of fluorine until the Second World War, when the development of the atom bomb, and other nuclear energy projects, made it necessary to produce large quantities. Before this, fluorine salts, known as fluorides, were for a long time used in welding and for frosting glass.
The element is used to make uranium hexafluoride, needed by the nuclear power industry to separate uranium isotopes. It is also used to make sulfur hexafluoride, the insulating gas for high-power electricity transformers.
In fact, fluorine is used in many fluorochemicals, including solvents and high-temperature plastics, such as Teflon (poly(tetrafluoroethene), PTFE). Teflon is well known for its non-stick properties and is used in frying pans. It is also used for cable insulation, for plumber’s tape and as the basis of Gore-Tex® (used in waterproof shoes and clothing).
Hydrofluoric acid is used for etching the glass of light bulbs and in similar applications.
CFCs (chloro-fluoro-carbons) were once used as aerosol propellants, refrigerants and for ‘blowing’ expanded polystyrene. However, their inertness meant that, once in the atmosphere, they diffused into the stratosphere and destroyed the Earth’s ozone layer. They are now banned.
Biological role
Fluoride is an essential ion for animals, strengthening teeth and bones. It is added to drinking water in some areas. The presence of fluorides below 2 parts per million in drinking water is believed to prevent dental cavities. However, above this concentration it may cause children’s tooth enamel to become mottled. Fluoride is also added to toothpaste.
The average human body contains about 3 milligrams of fluoride. Too much fluoride is toxic. Elemental fluorine is highly toxic.
Natural abundance
The most common fluorine minerals are fluorite, fluorspar and cryolite, but it is also rather widely distributed in other minerals. It is the 13th most common element in the Earth’s crust.
Fluorine is made by the electrolysis of a solution of potassium hydrogendifluoride (KHF2) in anhydrous hydrofluoric acid.
The early chemists were aware that metal fluorides contained an unidentified element similar to chlorine, but they could not isolate it. (The French scientist, André Ampère coined the name fluorine in 1812.) Even the great Humphry Davy was unable to produce the element, and he became ill by trying to isolate it from hydrofluoric acid.
The British chemist George Gore in 1869 passed an electric current through liquid HF but found that the gas which was liberated reacted violently with his apparatus. He thought it was fluorine but was unable to collect it and prove it. Then in 1886 the French chemist Henri Moissan obtained it by the electrolysis of potassium bifluoride (KHF2) dissolved in liquid HF.
| Atomic radius, non-bonded (Å) | 1.47 | Covalent radius (Å) | 0.60 |
| Electron affinity (kJ mol−1) | 328.165 |
Electronegativity (Pauling scale) |
3.98 |
|
Ionisation energies (kJ mol−1) |
1st
1681.045
2nd
3374.17
3rd
6050.441
4th
8407.713
5th
11022.755
6th
15164.128
7th
17867.734
8th
92038.447
|
||
|
| Common oxidation states | -1 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 19F | 18.998 | 100 | - | - | |
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Specific heat capacity (J kg−1 K−1) |
824 | Young's modulus (GPa) | Unknown | |||||||||||
| Shear modulus (GPa) | Unknown | Bulk modulus (GPa) | Unknown | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Fluorine Podcast |
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Transcript :
Chemistry in its element: fluorine (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith This week, a strong acid it's not, but deadly it definitely is. Kira J. Weissman The 37-year old technician spilled only a few hundred milliliters or so in his lap during a routine palaeontology experiment. He took the normal precaution in such situations, quickly dowsing himself with water from a laboratory hose, and even plunged into a nearby swimming pool while the paramedics were en route. But a week later, doctors removed a leg, and a week after that, he was dead. The culprit: hydrofluoric acid (colloquially known as HF), and the unfortunate man was not its first victim. Unlike its close relatives, hydrochloric and hydrobromic acid, HF is a weak acid. This, coupled with its small molecular size, allows it to penetrate the skin and migrate rapidly towards the deeper tissue layers. Once past the epidermis, HF starts to dissociate, unleashing the highly-reactive fluoride ion. Free fluoride binds tightly to both calcium and magnesium, forming insoluble salts which precipitate into the surrounding tissues. Robbed of their co-factors, critical metabolic enzymes can no longer function, cells begin to die, tissues to liquefy and bone to corrode away. And if calcium loss is rapid enough, muscles such as the heart stop working. Burns with concentrated HF involving as little as 2.5% of the body surface area - the size of the sole of the foot, for example - have been fatal. HF has a long history of destructive behaviour, claiming the lives of several chemists in the 1800s, including the Belgian Paulin Louyet, and the Frenchman Jérôme Nicklès. These brave scientists were battling to be the first to isolate elemental fluorine (F2) from its various compounds, using electrolysis. However, it was Nicklès' countrymen, Henri Moissan, who succeeded in 1886. To achieve this feat, Moissan not only had to contend with HF - the preferred electrolyte in such experiments - but fluorine itself, a violently reactive gas. His key innovation was to construct an apparatus out of platinum, one of the few metals capable of resisting attack, while cooling the electrolytic solution down to -50 °C to limit corrosion. Moissan's feat earned him the 1906 Nobel Prize in chemistry, but the celebration was short-lived. Another victim of fluorine's toxic effects, he died only two months later. Yet Moissan's method lived on, and is used today to produce multi-ton quantities of fluorine from its ore fluorspar. Ironically, while elemental fluorine is decidedly bad for your health, fluorine atoms turns up in some 20% of all pharmaceuticals. The top-selling anti-depressant Prozac, the cholesterol-lowering drug Lipitor, and the antibacterial Cipro, all have fluorine to thank for their success. How is this possible? Because the flip side of fluorine's extreme reactivity is the strength of the bonds it forms with other atoms, notably including carbon. This property makes organofluorine compounds some of the most stable and inert substances known to man. Fluorine's special status also stems from the 'fluorine factor', the ability of this little atom to fine-tune the chemical properties of an entire molecule. For example, replacing hydrogen with fluorine can protect drugs from degradation by metabolic enzymes, extending their active lifetimes inside the body. Or the introduced fluorine can alter a molecule's shape so that it binds better to its target protein. Such precise chemical tinkering can now be carried out in pharmaceutical labs using an array of safe, commercially-available fluorinating agents, or the tricky transformations can simply be out-sourced to someone else. Most of us also have fluorine to thank for our beaming smiles. The cavity-fighting agents in toothpaste are inorganic fluorides such as sodium fluoride and sodium monofluorophosphate. Fluoride not only decreases the amount of enamel-dissolving acid produced by plaque bacteria, but aids in the tooth rebuilding process, insinuating itself into the enamel to form an even harder surface which resists future attack. And the list of medical applications doesn't stop there. Being put to sleep is a little bit less worrisome thanks to fluorinated anaesthetics such as isoflurane and desflurane, which replaced flammable and explosive alternatives such as diethyl ether and chloroform. Fluorocarbons are also one of the leading candidates in development as artificial blood, as oxygen is more soluble in these materials than most other solvents. And radioactive fluorine (18F rather than the naturally-occurring 19F) is a key ingredient in positron emission tomography (or PET), a whole-body imaging technique that allows cancerous tumours to be discovered before they spread. Fluorochemicals are also a mainstay of industry. One of the most famous is the polymer polytetrafluoroethylene, better known as Teflon, which holds the title of world's most slippery solid. Highly thermostable and water proof, it's used as a coating for pots and pans, in baking sprays, and to repel stains on furniture and carpets. Heating and stretching transforms Teflon into Gore-tex, the porous membrane of sportswear fame. Gore-tex's pores are small enough to keep water droplets out, while allowing water vapour (that is, sweat) to escape. So you can run on a rainy day, and still stay dry. Fluorine plays another important role in keeping you cool, as air-conditioning and household refrigeration units run on energy-efficient fluorocarbon fluids. And fluorine's uses are not limited to earth. When astronauts jet off into space they put their trust in fluoroelastomers, a type of fluorinated rubber. Fashioned into O-rings and other sealing devices, these materials ensure that aircraft remain leak-free even under extreme conditions of heat and cold. And when accidents do happen, space travellers can rely on fluorocarbon-based fire extinguishers to put the flames out. Fluorine has long been known as the 'tiger of chemistry'. And while the element certainly retains its wild side, we can reasonably claim to have tamed it. As only a handful of naturally-occurring organofluorine compounds have ever been discovered, some might argue that we now make better use of fluorine than even Nature herself. Chris Smith So Teflon is acknowledged as the world's most slippery thing and I bet there are one or two politicians knocking around who are thanking fluorine for that. Thank you also to Kira Weismann from Zaarland University in Germany. Next week.ouch Steve Mylon I cannot imagine that this is all someone would be saying if they were unfortunate enough to be stricken with the disease of the same name. The ouch-ouch disease. The disease results from excessive cadmium poisoning and was first reported in a small town about 200 miles north west of Tokyo. Rice grown in cadmium contaminated soils had more than 10 times the cadmium content than normal rice. The ouch-ouch-ness of this disease resulted from weak and brittle bones subject to collapse due to high porosity. Chris Smith And you can find out about the ouch-ouch factor with Steve Mylon when he uncovers the story of cadmium on next week's Chemistry in Its Element. I'm Chris Smith, thank you for listening and goodbye. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
