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There is an element so reactive that it was one of the last to be isolated, and chemists were injured trying to obtain it.
Yet once tamed, it became essential to everything from aluminium production and nuclear technology to medicine, batteries, and nonstick coatings.
Its compounds have also been at the center of major debates over public health and environmental contamination.
Learn more about the extraordinary element fluorine on this episode of Everything Everywhere Daily.
I’m guessing that the vast, vast, vast majority of you have no direct experience with elemental fluorine… and that is a good thing. You wouldn’t want to mess around with elemental fluorine because it is the most reactive element on the periodic table.
That said, almost everyone has encountered fluorine when it’s bonded to something else to form a highly stable molecule.
Before we get into the uses of fluorine, we should start with the basics of what it is.
Fluorine has the atomic number 9, meaning there are 9 protons in its nucleus, and the chemical symbol F. It is a member of the halogens, Group 17 of the periodic table, along with chlorine, bromine, iodine, and astatine. At ordinary temperatures, elemental fluorine exists as F?, a pale yellow gas.
It is the most electronegative element and generally regarded as the most chemically reactive element.
The reason for Fluorine’s extraordinary chemistry arises largely from its atomic structure. It has nine electrons, with seven electrons in its outer shell. It needs only one additional electron to complete that shell.
Fluorine therefore has an exceptionally strong tendency to attract electrons from other atoms. Its electronegativity on the Pauling scale is 3.98, the highest of any element.
Because fluorine is so reactive, essentially no free elemental fluorine exists naturally near Earth’s surface. It almost immediately combines with other elements.
Fluorine is relatively uncommon in the universe. It ranks roughly 24th in cosmic abundance, with only about one fluorine atom for every few hundred million hydrogen atoms.
Its scarcity is partly because ordinary stellar fusion does not produce fluorine efficiently. Most cosmic fluorine is thought to form in relatively specialized environments, including certain giant stars and supernovae.
The most important naturally occurring fluorine mineral is fluorite, also known commercially as fluorspar. It consists of 1 calcium atom and 2 fluorine atoms. Other fluorine-containing minerals include fluorapatite, cryolite, and topaz. Fluorite is overwhelmingly the most important mineral commercially mined specifically as a source of fluorine.
Fluorine-based minerals tend to be gorgeous. They can have deep, vibrant colors, and if you are into mineral collecting, you can find them relatively cheaply.
The name fluorine comes from the Latin fluere, meaning “to flow.” Miners and metallurgists discovered long before anyone understood fluorine chemistry that fluorite could be added to ores as a flux, helping materials melt and flow more easily during smelting.
Fluorine has one of the more dramatic discovery stories in chemistry.
By the seventeenth and eighteenth centuries, experimenters knew that fluorite had unusual chemical properties. They discovered that treating fluorite with strong acids produced a substance capable of attacking glass.
The Swedish chemist Carl Wilhelm Scheele investigated fluorite in the 1770s and produced what we now recognize as hydrofluoric acid. Chemists eventually realized that chemicals produced from fluorite probably contained an unknown element analogous to chlorine. In 1812, the French physicist and chemist André-Marie Ampère proposed the name fluorine.
Actually obtaining the element was another matter.
The problem was that whatever apparatus researchers used tended to be attacked by the substance they were trying to produce. Hydrofluoric acid was itself extremely dangerous, and newly liberated fluorine reacted violently with water, glass, metals, organic matter, and many electrode materials.
The breakthrough finally came from the French chemist Henri Moissan.
Moissan dissolved potassium hydrogen fluoride, KHF?, in anhydrous hydrogen fluoride. This produced an electrolyte through which sufficient electrical current could pass. He constructed specialized equipment using materials able to withstand the chemicals and cooled the apparatus significantly to suppress unwanted reactions.
On June 26, 1886, Moissan successfully produced and identified elemental fluorine gas.
Moissan received the 1906 Nobel Prize in Chemistry, in significant part for isolating fluorine. The basic electrochemical principle he developed remains the foundation of industrial fluorine production.
While fluorine is very reactive, once it reacts, it bonds so strongly that the resulting compound tends to be extremely stable.
These seemingly contradictory properties of fluorine both have industrial and commercial uses.
One of the earliest uses of fluorine was in refining aluminum.
Producing aluminum metal requires electrolyzing aluminum oxide. Pure aluminum oxide has an extremely high melting point, making direct electrolysis impractical.
The Hall-Héroult process dissolves alumina in a molten fluoride-based electrolyte, historically associated with cryolite. This dramatically lowers the operating temperature and allows aluminiumum to be produced economically.
One of hydrofluoric acid’s most famous properties is its ability to attack glass. Fluoride chemistry can convert silicon-containing material into fluorosilicate compounds, allowing hydrofluoric acid to etch the surface.
This makes hydrofluoric acid valuable for glass treatment, industrial cleaning, semiconductor processing, and chemical manufacturing.
It also makes hydrofluoric acid unusually hazardous.
Unlike many acids that mainly cause surface burns, hydrofluoric acid can penetrate tissues. Fluoride ions can bind with calcium and magnesium, potentially producing serious toxicity. Significant exposures can therefore become medical emergencies even when the initial burn does not appear catastrophic.
One of fluorine’s most strategically important applications is uranium enrichment.
As I’ve covered in previous episodes, natural uranium consists primarily of uranium-238 with only about 0.7 percent uranium-235. Nuclear reactors and nuclear weapons may require uranium with a higher proportion of U-235.
Uranium metal itself is unsuitable for many isotope-separation methods. But uranium can be converted to uranium hexafluoride.
Uranium hexafluoride has the extremely useful property of becoming gaseous at relatively modest temperatures. The gaseous molecules containing U-235 are every so slightly lighter than those containing U-238.
This allows the isotopes to be separated using technologies such as gaseous diffusion historically and, today, primarily gas centrifuges.
Demand from the Manhattan Project and the nuclear industry played an important role in expanding industrial fluorine production during and after World War II.
One of the best-known fluorinated materials is polytetrafluoroethylene, or PTFE, but you probably know it better by the brand name, Teflon.
Its repeating structure contains carbon atoms surrounded by fluorine atoms. The carbon-fluorine bonds are exceptionally strong, and the fluorine effectively shields the carbon backbone.
As a result, the material has low friction, resists many chemicals, tolerates high temperatures, and does not readily react with other materials.
PTFE and related fluoropolymers are used far beyond cookware. They appear in chemical processing equipment, electrical insulation, aerospace systems, seals, tubing, medical devices, laboratory equipment, semiconductor plants, and countless specialized applications.
PTFE and Teflon will be the subject of a future episode, as their origin and how they work are really interesting.
Closely related to PTFE, yet very different, are PFAS, or polyfluoroalkyl substances.
They are a large family of synthetic chemicals valued for resisting heat, water, oil, and chemical breakdown. They have been used in products such as firefighting foams, stain-resistant fabrics, food packaging, industrial coatings, and some nonstick applications.
The controversy is that many PFAS break down extremely slowly, so they can persist in water, soil, wildlife, and the human body for years, earning the nickname “forever chemicals.”
Their stability is the problem. Certain PFAS have been linked to health concerns including immune effects, developmental problems, higher cholesterol, and increased risk of some cancers.
This too will be the subject of a future episode.
Fluorine also transformed refrigeration during the twentieth century. Early refrigeration systems sometimes used dangerous substances such as ammonia, sulfur dioxide, or methyl chloride.
Chemists developed chlorofluorocarbons, or CFCs, which appeared nearly ideal. They were stable, relatively nonflammable, and useful as refrigerants, aerosol propellants, foam-blowing agents, and solvents.
…but their chemical stability became an environmental problem.
CFC molecules survive long enough to reach the stratosphere, where ultraviolet radiation breaks them apart. Importantly, it is primarily the chlorine, rather than fluorine, released from CFCs that destroys stratospheric ozone.
When CFC’s were phased out, the replacements still relied on fluorine. HCFCs, or hydrochlorofluorocarbons, replaced CFCs, but they still contained chlorine, so they have also been phased out.
HFCs, or hydrofluorocarbons, replaced them; they contain hydrogen, fluorine, and carbon but no chlorine, so they do not significantly deplete ozone. However, many HFCs are powerful greenhouse gases, so they too are being phased down.
Fluorine plays an essential role in powering modern consumer electronics, as today’s lithium-ion batteries regularly rely on fluorine chemistry.
Lithium hexafluorophosphate serves as one of the most widely used electrolyte salts. When dissolved in organic solvents, this compound facilitates the transport of lithium ions between battery electrodes.
Fluorinated electrolyte additives and fluorinated electrode materials are also becoming increasingly important in battery research because chemists can use fluorine to alter voltage stability, conductivity, surface chemistry, and resistance to degradation.
While fluorine has many other uses, you probably thought of one immediately when you saw the episode title: dental products.
Fluoride, which is just a negative fluorine ion, can become incorporated into or associated with tooth mineral, making enamel more resistant to acid attack. It can also promote remineralization of early tooth decay.
Fluoride can be found in almost all toothpastes, and Many communities also adjust fluoride levels in drinking water.
The current U.S. Public Health Service recommendation is approximately 0.7 milligrams per liter, or 0.7 parts per million. The Centers for Disease Control states that fluoridation at recommended levels reduces tooth decay and cites an approximate 25 percent reduction in cavities among children and adults.
Water fluoridation began in the United States in 1945, after observations showed that communities with certain naturally occurring fluoride levels had lower rates of dental cavities.
For decades, major dental and public-health organizations have strongly supported fluoridation because of its effects on tooth decay.
However, the decision to fluoridate water has been controversial, as Fluoride is toxic at high levels. Excessive fluoride exposure while children’s permanent teeth are developing can cause dental fluorosis, in which enamel develops white streaks.
Of all the uses for fluorine I’ve covered, the benefits and drawbacks both stem from the same thing: fluorine forms incredibly strong chemical bonds that do not break easily.
Fluorine’s reactivity allows it to etch silicon, process metals, produce uranium hexafluoride, and modify surfaces.
The strength of fluorinated compounds gives us chemically resistant plastics, durable coatings, powerful pharmaceuticals, stable refrigerants,, and batteries.
In nature, it can form beautiful minerals and gemstones.
But that same chemical stability created CFCs that survived long enough to reach the ozone layer, HFCs that remain powerful greenhouse gases, and PFAS compounds that can persist in the environment for decades or longer.
Fluorine is not inherently a technological blessing or an environmental curse. It is an unusually powerful chemical tool that requires careful consideration of how it is used.