Fluorine

Fluorine
Atomic number	9
Symbol	F
Group	17 (halogens)
Boiling point	−188.1 °C
Electronegativity	3.98 (Pauling)
Electron configuration	[He] 2s2 2p5
Melting point	−219.7 °C
Main isotopes	19F
Block	p
Phase STP	Gas
Oxidation states	−1, 0
Wikidata	Q650

Fluorine is a pale yellow-green, highly reactive nonmetallic halogen gas with atomic number 9 (symbol F). It is the lightest halogen and is notorious for its extreme chemical reactivity. Fluorine atoms have seven electrons in their outer shell (electron configuration [He] 2s²2p⁵), making them eager to gain one more electron. This gives fluorine the highest electronegativity of any element (about 4.0 on the Pauling scale), so it attracts electrons very strongly in chemical bonds. At standard conditions (0 °C and 1 atm), fluorine is a diatomic gas (F₂). It has a melting point near –219.6 °C and a boiling point around –188.1 °C, so it is a gas in everyday temperatures. The gaseous density of F₂ at STP is about 1.70 grams per liter. Fluorine occurs in nature only in compounds (never free) and almost always with oxidation state –1. Its common oxidation state in compounds is –1 (in F₂ it is 0), and it belongs to Group 17 of the periodic table (period 2, block p) like other halogens.

Atomic Structure and Electron Configuration

An atom of fluorine has 9 protons and 9 electrons. Its electron configuration is [He] 2s²2p⁵. In other words, fluorine’s valence shell (the second shell) contains seven electrons: two in the 2s orbital and five in the 2p orbitals. With seven outer-shell electrons, fluorine is one electron short of a full shell (eight electrons). This makes the neutral fluorine atom extremely eager to gain an extra electron to achieve a stable neon-like configuration. This accounts for fluorine’s very high electron affinity and electronegativity.

Among all elements, fluorine has one of the smallest atomic radii. Its covalent atomic radius (for forming a single bond) is about 60–65 picometers, and its van der Waals radius is roughly 135 picometers. The small size of the fluorine atom and its high effective nuclear charge (nine protons holding few inner electrons) lead to a very high Pauling electronegativity of about 3.98. In periodic trends, fluorine’s electronegativity is the highest of all elements. Its first ionization energy (the energy to remove one electron) is also extremely high – about 1,681 kJ/mol (17.42 eV), third highest after helium and neon. (By comparison, chlorine’s first ionization energy is about 1,250 kJ/mol.) Fluorine’s small size and strong attraction to electrons means it strongly resists losing an electron. Its second ionization energy (removing a second electron) is enormous (about 3,374 kJ/mol). In short, fluorine’s atomic structure – a high nuclear charge and a nearly full valence shell – makes it the runner-up in electron affinity (just behind chlorine) and the champion of electronegativity among reactive elements.

Isotopes and Nuclear Properties

Naturally occurring fluorine consists of only one stable isotope: fluorine-19 (^19F). This isotope has 9 protons and 10 neutrons, a mass of about 18.998 u, and 100% natural abundance. It has a nuclear spin of ½, making ^19F highly useful in nuclear magnetic resonance (NMR) studies. In fact, ^19F NMR is widely used in chemistry and medicine because ^19F is nearly as NMR-sensitive as ^1H and is found in many organic molecules. Because it is the only stable isotope, ^19F has no radioactive decay modes and makes up all of ordinary fluorine.

A handful of radioactive isotopes of fluorine have been synthesized or occur in trace amounts. The most important radioisotope is fluorine-18 (^18F), which has a half-life of about 109.8 minutes and decays by positron emission (β⁺) to oxygen-18. ^18F is produced naturally in the atmosphere (by cosmic-ray reactions) and is also produced in cyclotrons. It is extensively used as a tracer in medical positron emission tomography (PET) imaging – for example in the glucose tracer ^18F-FDG used in cancer diagnostics. Other short-lived isotopes include ^17F (half-life ~64 seconds) and ^20F (half-life ~11 seconds); these decay by positron emission (^17F) or beta decay (^20F) and have no practical uses outside of research. In general, fluorine has no long-lived radioisotopes besides ^18F, and no isotope of fluorine is used for radiometric dating because they decay too quickly.

Allotropes and Typical Compounds

Elemental fluorine has no allotropes at room temperature beyond its diatomic gas form, F₂. (In contrast, chlorine, bromine, and iodine also occur as diatomic molecules under standard conditions.) When cooled to very low temperatures, solid fluorine does exhibit two crystalline phases. Solid F₂ has been found to form an α-phase and a β-phase at cryogenic temperatures (below –160 °C). The α-F₂ form is monoclinic (space group C2/c) and β-F₂ is cubic (Pm3n), but both contain only F₂ molecules arranged in a lattice. These solid structures were difficult to determine because fluorine is so reactive it can attack sample containers. Overall, there are no polymeric or multi-atom allotropes of pure fluorine under ordinary conditions – it exists as discrete F₂ molecules (much like oxygen as O₂).

Fluorine forms a rich variety of chemical compounds, mainly owing to the fluoride anion (F⁻) and the strong F–X bond. Some of the most characteristic compounds and classes are:

• Hydrogen fluoride (HF): A diatomic gas that dissolves in water to give hydrofluoric acid. HF is unique among hydrogen halides. It is a weak acid compared to HCl or HBr because its H–F bond is very strong; instead of fully dissociating, HF molecules form extensive hydrogen-bonded networks in liquid or solution. (HF has a boiling point far above analogous hydrogen halides due to H-bonding.) Because HF is both a weak acid and a weak base, it uniquely often exists as polymeric ions like H₂F²⁻ (bifluoride). HF and its aqueous solutions aggressively etch glass and react with many metals.

• Fluoride salts and oxides: Fluorine commonly occurs in nature as fluoride anions bound to metals. Calcium fluoride (CaF₂, “fluorspar” or fluorite) is the chief ore. Calcium and sodium produce common salts (e.g. CaF₂, NaF, K₂SiF₆), which are ionic solids. Working up fluorite with sulfuric acid yields HF, which is the main route to many fluorine chemicals. Metal fluorides (e.g. MgF₂, AlF₃, and uranium hexafluoride UF₆) often have high melting points and strong ionic/covalent bonds. Many metal-fluorine compounds are very stable; for example, aluminum oxide (Al₂O₃) is not easily fluoridized, whereas sodium chloride is easily converted to sodium fluoride by fluorine.

• Interhalogen and oxide compounds: Fluorine forms reactive interhalogen molecules, such as ClF, ClF₃, BrF₃, BrF₅, and IF₇. These typically involve one fluorine and one other halogen, and they are often strong fluorinating or oxidizing agents themselves. Fluorine also bonds to oxygen in a few molecules: for example, dioxygen difluoride (O₂F₂) and oxygen difluoride (OF₂) – in these compounds fluorine is –1 and oxygen is in a positive oxidation state. Oxygen difluoride is a toxic, corrosive gas. In general, fluorine tends to oxidize oxygen rather than the reverse (oxygen can have a +2 state in OF₂).

• Noble gas fluorides: Unreactive noble gases actually form stable compounds with fluorine under extreme conditions. Xenon fluorides (XeF₂, XeF₄, XeF₆) and krypton difluoride (KrF₂) are well-known. Xenon tetrafluoride and hexafluoride are important for producing reactive fluoride complexes (e.g. XeF₆ reacts with fluoride acceptors to generate XeF₅⁺). Argon fluorides have been observed only at very low temperatures or in discharges.

• Organic fluorides (fluorocarbons): Fluorine forms very strong bonds to carbon. Perfluorinated organic compounds, especially polytetrafluoroethylene (PTFE) and related fluoropolymers, are famous. PTFE (Teflon) has the repeating unit –CF₂– and is extremely chemically inert, having a very high melting point (~327 °C). Other fluoropolymers like polyvinylidene fluoride (−(CH₂CF₂)n−) and Nafion (a sulfonated fluoropolymer) are important in plastics and membranes. Many organic solvents and refrigerants contain fluorine (e.g. chlorofluorocarbons CCl₂F₂, hydrochlorofluorocarbons CH₂ClF, or hydrofluorocarbons like CH₃CF₃). These compounds tend to be quite stable, nonflammable, and have applications in industries.

• Special compounds: Some fluorine compounds are notable for technological reasons. For example, uranium hexafluoride (UF₆) is a volatile solid used to separate uranium isotopes by gas diffusion or gas centrifuge (because UF₆ forms a gas near room temperature and can be centrifuged). Sulfur hexafluoride (SF₆) is an extremely inert, high-electronegativity gas used as a high-voltage insulator. Boron trifluoride (BF₃) and hexafluoroantimonic acid (HSbF₆) are superacids (very strong acids).

In summary, fluorine’s compounds are usually characterized by very strong bonds to fluorine. Fluoride (F⁻) is a very small, strongly basic anion that forms strong ionic bonds. Covalent F–X bonds (to carbon, oxygen, metals, nonmetals) are very polar and strong. Fluorine does not exhibit positive oxidation states in stable compounds (unlike chlorine, which can go +1, +3, etc). Instead, it usually takes –1 (with the only other oxidation state being 0 in F₂ itself).

Physical Properties

Fluorine is a gas at standard temperature and pressure. It appears as a pale yellow-green gas with a distinctive, pungent odor. (Some describe it as a bleach-like smell.) In liquid form (below –188 °C) it is a pale golden color. The liquid density at the boiling point is about 1.505 grams per cubic centimeter (i.e. fluorine becomes quite dense as a liquid). The limits are low: fluorine melts at 53.48 K (–219.67 °C) and boils at 85.03 K (–188.11 °C). Its triple point is at 53.52 K and 0.252 kPa.

Solid fluorine (below 53 K) consists of F₂ molecules in a molecular crystal. As mentioned, there are two solid phases, α-F₂ and β-F₂, depending on temperature (and pressure). Under normal pressure, the transition between α and β occurs around 45 K. Both solids are composed of discrete F₂ molecules held in a lattice. The F–F bond length in F₂ gas or solid is about 142 pm (1.42 Å).

Thermally, fluorine gas has low heat capacity and thermal conductivity typical of diatomic gases. Its thermal conductivity at STP is about 25×10^–3 W/m·K. Electrically, fluorine gas is an excellent insulator (no conduction by free carriers) and is diamagnetic (all electrons in F₂ are paired). Under electric fields or UV light, fluorine becomes highly ionized and reactive, but in ordinary conditions it has no appreciable conductivity.

Fluorine exhibits characteristic spectroscopic lines in the ultraviolet. Atomic fluorine lines occur in the extreme UV (for example around 600 Å for F II transitions). Most spectroscopy involving fluorine focuses on molecules (HF, etc.) or on ^19F NMR.

Chemical Reactivity and Trends

Fluorine’s chemical reactivity is legendary – it is the most electronegative and one of the strongest oxidizing agents of all elements. Its standard reduction potential (F₂ + 2 e⁻ → 2 F⁻) is about +2.87 V, higher than for any other element. In practice, fluorine will oxidize nearly every other element, converting it to a fluoride. For example, fluorine reacts explosively with hydrogen gas to form HF. It reacts with water to produce hydrofluoric acid and oxygen gas (2 F₂ + 2 H₂O → 4 HF + O₂). It will fluorinate metals and nonmetals alike; for instance, fluorine oxidizes most metals to their highest oxidation state and often forms stable metal fluorides. Even noble metals like gold or platinum can be converted to fluorides (e.g. PtF₆ is known). Many organic compounds are set on fire by fluorine – contact with most hydrocarbons causes vigorous combustion, often with a bright flame.

Key aspects of fluorine’s chemical behavior include:

• Oxidizing power: Fluorine can oxidize substances that other oxidizers cannot. It can convert sulfides to sulfates, chlorides to chlorates, and even burn glass (SiO₂) if in contact with silicon compounds. It is able to oxidize oxygen under extreme conditions to form compounds like O₂F₂ (difluorine oxide) – in effect oxygen becomes the less electronegative partner. In practical terms, fluorine gas can ignite nearly any reducing agent on contact.

• Bonding patterns: Because fluorine is so small and electronegative, it forms very polar covalent bonds. In organic chemistry, carbon–fluorine bonds are among the strongest single bonds (bond energy ~485 kJ/mol), making organofluorine compounds highly stable. In inorganic compounds, fluorine nearly always carries a –1 formal charge (with the rare exception of F₂ or F radicals). As a result, compounds like OF₂ have oxygen in a positive oxidation state (+2) and F at –1. The inability of fluorine to have a positive oxidation state means it never forms species like F⁺ in normal chemistry – fluorine will not release an electron to form cations, it only forms anions or covalent bonds.

• Acid–base behavior: Fluorine greatly influences acidity. Hydrofluoric acid (HF) in water is a weak acid relative to HCl or HBr because the H–F bond is very strong. Its pKa is about 3.2, much higher (weaker) than HCl or HBr (pKa ~ –7 and –8). HF forms strong hydrogen-bonded chains in liquid, giving it a high boiling point (19.5 °C) for a hydrogen halide. Conversely, the fluoride ion (F⁻) is a strong base and a powerful hydrogen-bond acceptor; it readily forms HF again by capturing a proton. In aqueous solution, F⁻ will often pull protons from water to re-form HF (making F⁻ solutions somewhat basic). Many fluoride complexes exist (for instance, the bifluoride ion HF₂⁻ is common in HF solutions).

• Complexation and coordination: Fluoride ions bind strongly to many metal centers. For example, aluminum forms the complex ion AlF₄⁻, and silicon forms SiF₄. Hard metal ions like Al³⁺, Be²⁺, and rare-earth cations often coordinate with fluoride in complex ions (due to fluoride’s small, highly charged nature). This “hard–hard” Lewis acid-base interaction is exploited in chemical processes: AlF₃ is used to catalyze alkylation, and SbF₅ plus HF forms “fluoroantimonic acid,” one of the strongest known superacids.

• Corrosion and passivation: Fluorine gas is highly corrosive and will attack many materials, including glass and metals. Most solids are destroyed or burned by F₂. Some metals (aluminum, stainless steel) can develop a thin fluoride layer that slows further reaction, but this passive film can be broken easily. Only a few materials stand up to fluorine, such as nickel, copper-nickel alloys, or Teflon-lined containers. Fluorine also reacts with water; any exposure of F₂ to humidity produces HF (which itself is corrosive). In fact, a key safety rule is that fluorine cannot be stored in glass or metallic containers – it is typically handled as a cryogenic liquid or gas in specially coated metal or plastic systems.

• Reactivity series: In the context of halogens, fluorine is at the top. It will oxidize other halide ions (in theory F₂ can oxidize Cl⁻ to ClF, although fluoride is so reactive it usually oxidizes water first). It ranks above oxygen on the Pauling scale. Down the group, chlorine, bromine, and iodine are progressively less electronegative and less reactive. Compared to other group 17 elements, fluorine’s electron affinity is actually slightly less than chlorine’s (due to its small size causing electron-electron repulsion), but its net effect is that it is the most aggressive electron-seeker. In summary, any chemical trend in halogens is largely dominated by fluorine’s extremity: it has the smallest size, the highest charge density, and the highest tendency to attract electrons.

In practical terms, almost no substance can resist hot or concentrated fluorine. Water, hydrogen, rocks, and organic matter all burn or react vigorously. This makes elemental fluorine difficult to handle but extremely useful for quick, powerful chemical transformations.

Occurrence and Production

Fluorine is relatively abundant in the Earth’s crust but is found only in compounds, not as a free element. It ranks around the 13th most common element in the crust, at roughly 0.054% by weight. Its cosmic abundance is modest; fluorine is not produced in large amounts by stellar nucleosynthesis, so it is much rarer in the universe than lighter elements. On Earth, fluorine’s most common natural forms are mineral fluorides. The chief fluorine-bearing minerals include:

• Fluorite (fluorspar, CaF₂): The most important source of fluorine. Fluorite is a common gangue mineral in ore veins. It appears as colorless or colored crystals. When treated with sulfuric acid (H₂SO₄), CaF₂ yields hydrogen fluoride (HF) gas, which is the starting point for other fluorine compounds.

• Cryolite (Na₃AlF₆): A rare mineral that was historically mined in Greenland. Cryolite is used as a flux in aluminum smelting (it dissolves Al₂O₃ and lowers its melting point). Cryolite itself is a double fluoride of sodium and aluminum.

• Apatite (Ca₅(PO₄)₃F or Ca₅(PO₄)₃(Cl, F) depending on type): A family of phosphate minerals, some of which contain fluoride (fluorapatite). Apatite is the main source of phosphate rock and contains a small percentage of fluorine. Phosphate rock processing can release HF.

• Villiaumite (NaF): A rare mineral (sodium fluoride) found in some pegmatites. It provides fluoride but is not a major source historically.

Fluorine does not occur in significant amounts in water; seawater contains only about 1.3 ppm fluorine, mostly as F⁻. The elemental gas F₂ is never found naturally (it reacts away immediately). Most fluorine in the environment is bound as fluoride in soil, minerals, or biological matter (teeth and bones of animals have fluorapatite).

The world production of fluorine compounds is large, driven by demand for hydrofluoric acid (HF) and related chemicals. HF is produced on the order of millions of tons per year globally, and it serves as an intermediate for fluorine chemistry. Industrial fluorine (F₂ gas) is manufactured on a smaller scale (thousands of tons per year). The primary method to make F₂ is electrolysis of anhydrous hydrogen fluoride. In the modern “Moissan” process, HF is dissolved in potassium fluoride (KF) to form potassium bifluoride (KHF₂) and then electrolyzed:

<code>2 HF + 2 KF (molten) → 2 KHF₂ (electrolysis) → F₂ (gas at anode) + H₂ (gas at cathode).
</code>

Nickel or copper electrodes and cell materials resistant to fluorine are used. The result is pure F₂ gas and hydrogen. (Hydrogen is usually vented or burned.)

Most industrial fluoride comes from sourcing fluorite and converting it to HF. Major fluorite-producing countries are China (about 60–70% of world production), Mexico, Mongolia, South Africa, and Vietnam, among others. Thus, China and Mexico dominate the supply of raw fluorine-containing ore. The largest chemical producers of fluorine gases and HF include companies in the United States, Europe, Japan, and China. For example, DuPont (USA) and Bayer (Germany) historically were major producers of F₂ gas. Today many chemical firms manufacture specialty fluorine products (plans often use imported HF or F₂), while China has expanded rapid mining and HF production capacity.

Fluorine from uranium: In the nuclear fuel industry, fluorine is involved indirectly. Uranium is converted to uranium hexafluoride (UF₆) gas for isotope separation during enrichment. This UF₆ is not a direct mining product but is made in plants: uranium ore (U₃O₈) is treated with fluorine-containing reagents to make UF₆. Thus fluorine is also indirectly mined in any nation with uranium fuel processes. Major UF₆ production occurs in countries with nuclear programs (USA, Russia, France, China, etc.).

Fluorine has no significant “biological” sources or sinks beyond being locked in minerals or bone. In space, fluorine is produced in certain stellar environments (some red giants) and by cosmic-ray interactions (for example, atmospheric argon can yield small amounts of ^18F). The principal terrestrial fluorine cycle is the rock cycle: weathering of minerals slowly releases fluoride into soil and water; organisms deposit it in bones and teeth; and mining/industrial processes redistribute it.

Applications and Technology

Fluorine’s unique chemical properties have led to many important applications across technology and industry. Below are some key uses and products that involve fluorine:

• Fluoropolymers (nonstick plastics): Perhaps the most famous is PTFE (polytetrafluoroethylene), discovered accidentally by DuPont chemist Roy Plunkett in 1938. PTFE, known by the brand name Teflon, is used to coat nonstick cookware, to make chemical-resistant seals, gaskets, and hoses, and as an insulator in aerospace wiring. Its C–F bonds make it extremely stable to heat and chemicals. Other fluoropolymers include PVDF (polyvinylidene fluoride) for wire coatings and Nafion (a sulfonated fluoropolymer) used in fuel cells and electrolysis membranes.

• Refrigerants and blowing agents: Fluorine-containing refrigerants were once widely used. Early refrigerants like R-12 (CCl₂F₂) and R-22 (CHClF₂) are chlorofluorocarbons (CFCs) or hydrofluorocarbons (HCFCs). Though many are now phased out (due to ozone depletion and global warming concerns), fluorinated gases like HFC-134a (CH₂FCF₃) and hydrofluoroolefins (HFOs) have replaced them in many uses. Pentafluoroethane, tetrafluoroethane, and related fluorocarbons are used in air conditioners and aerosol propellants. These fluorocarbons are stable, nonflammable, and excellent at absorbing infrared (so they have high global warming potential).

• Insulating gas (SF₆): Sulfur hexafluoride (SF₆) is a fluorine compound used as an electrical insulator in high-voltage equipment. SF₆ is extremely inert, has a very high dielectric strength, and remains gaseous over a wide temperature range. It has largely replaced air in switchgear and circuit breakers. (SF₆ does have one of the highest global warming potentials of any substance, so its use is now being re-evaluated in some systems.)

• Chemical intermediates (HF, fluorides): Hydrogen fluoride (HF) is itself a major industrial chemical. It is used to etch glass and silicon (HF dissolves SiO₂), to refine and purify gasoline (alkylation catalysts), and to produce specialty chemicals. Many metal fluorides are used in aluminum production (cryolite) and in steelmaking (fluorspar as a flux). Fluoride compounds like cryolite lower melting points of oxides in metallurgy.

• Nuclear fuel cycle: An important application is uranium processing. To prepare enriched fuel, uranium dioxide or yellowcake is converted to uranium hexafluoride (UF₆) via reaction with fluorine chemicals. This volatile UF₆ gas is then enriched (by gas diffusion or centrifuges) to increase the concentration of fissile ^235U. After enrichment, UF₆ is converted back to oxide fuel. Thus fluorine (as UF₆) plays a critical role in nuclear energy and weapons technology. (Note: UF₆ is also extremely reactive and toxic; it reacts with moisture to form HF.)

• Battery electrolytes: Lithium-ion batteries use fluorine-containing salts in their electrolytes. For example, lithium hexafluorophosphate (LiPF₆) is a common electrolyte salt. Fluorinated salts help to form stable solid-electrolyte interfaces and increase battery performance.

• Pharmaceuticals and agrochemicals: Many modern drugs and pesticides contain fluorine atoms. Incorporating F can improve a molecule’s stability and bioactivity. It is estimated that over 20% of pharmaceuticals on the market contain fluorine. Examples include Fluoxetine (Prozac) and many heart and antiviral drugs. Fluoroalkanes and related compounds are widely used as solvents and oil additives because C–F bonds resist degradation.

• Etching and cleaning: Hydrofluoric acid (and related fluoride treatments) are used to etch glass, clean silicon wafers, and remove metal oxides. HF vapor is used in semiconductor manufacturing to strip photoresist and silicon dioxide layers.

• Specialty applications: Fluorine lasers (e.g. F₂ laser at 157 nm) have been developed for microfabrication. Radioisotope ^18F is used in medical imaging (the glucose analog ^18F-FDG is used in PET scanners to detect tumors). Certain rocket fuels have been proposed using fluorine or fluorine compounds (e.g. liquid F₂ mixed with liquid hydrogen) due to fluorine’s high energy release, though safety issues have prevented widespread use.

In summary, fluorine sees use whenever highly stable chemicals, strong electronegativity, or high oxidation are needed. From plastics and refrigerants to nuclear fuel and pharmaceuticals, fluorine makes many modern technologies possible.

Biology, Environment, and Safety

Fluorine’s interactions with biology and the environment are mixed and require caution. In small amounts, fluorine (as fluoride, F⁻) can be beneficial. Many organisms incorporate fluoride into bones and teeth as fluorapatite (Ca₅(PO₄)₃F), which is harder than pure calcium phosphate. Fluoride ions at low concentrations (around 0.7–1.2 ppm in drinking water) help prevent tooth decay by stabilizing enamel. However, excess fluoride can cause health problems. Chronic overexposure to fluoride (often via drinking water or industrial exposure) leads to dental and skeletal fluorosis: staining and pitting of teeth and increased bone density that makes bones brittle. Extremely high fluoride intake (more than tens of ppm) can be toxic to soft tissues and organs.

Elemental fluorine (F₂ gas) and hydrogen fluoride (HF) are extremely hazardous to living organisms. F₂ is a corrosive oxidizer that attacks moisture (including tissues) on contact. Inhaling F₂ gas or HF fumes causes severe chemical burns in the respiratory tract, leading to lung edema (fluid build-up) and often death from asphyxiation. Skin contact with HF is especially dangerous because it penetrates tissues and dissolves into CaF₂ in the body, causing deep burns and disrupting calcium metabolism (potentially leading to cardiac arrest). Protective equipment (special gas masks and plastic or rubber suits) is absolutely required when handling fluorine. The Occupational Safety and Health Administration (OSHA) and NIOSH set very low exposure limits: the recommended exposure limit (REL) and permissible exposure level (PEL) for fluorine gas is around 0.1 ppm (0.2 mg/m³) as a time-weighted average. Concentrations above 25 ppm are immediately dangerous to life (IDLH). Even dilute aqueous HF (a few percent) is hazardous; for example, a 1% HF solution can cause severe eye or skin burns if not washed off quickly.

Fluoride ions in the environment also have complex impacts. Industrial releases of HF or fluoride dust (from phosphate fertilizer manufacturing, aluminum smelting, or brick and glass plants) can lead to local contamination of water and vegetation. Fluoride pollution can damage crops and cause illness in livestock near such plants. Inhaling dust ejected from such plants has been linked to respiratory problems in humans. Worldwide, natural fluoride levels in groundwater vary greatly; in some arid regions, fluoride is geologically high and can exceed safe drinking levels (causing endemic fluorosis).

Another environmental concern involves fluorinated organic compounds, especially the class of PFAS (per- and polyfluoroalkyl substances). These include long-chain fluorocarbons once used in nonstick coatings, firefighting foams, and many consumer products. PFAS are extremely resistant to breakdown (“forever chemicals”) and have been found in water, soil, and organisms globally. While PFAS are usually inert, they raise potential health concerns if ingested in large amounts.

Because of fluorine’s extreme reactivity, almost no natural cycle can handle it safely. Free fluorine or HF is neutralized quickly in nature by reacting with rocks (forming fluoride salts) or water. This means that fluorine’s primary environmental form is the fluoride anion, which is relatively inert. Still, industries that generate fluoride must treat waste scrubbing and emissions carefully.

Biologically, no cell uses elemental fluorine, and fluoride (F⁻) has no known essential metabolic role (beyond strengthening bones/teeth). High doses of fluoride are toxic to most organisms, as fluoride ions can inhibit many enzymes (by binding to metal centers) and can cause skeletal damage. Fish and aquatic life can be sensitive to fluoride levels above a few ppm. Regulatory agencies do monitor and limit fluoride in drinking water, food, and air emissions for this reason.

In short, fluorine and its compounds must be handled with extreme care. Industrial hygiene procedures for fluorine include: strict containment (often handling F₂ as a cryogenic liquid in passivated containers), specialized PPE, continuous HF scrubbers on exhaust, and emergency protocols for leaks (do not use water on pure F₂ leaks, since water reacts dangerously). First aid for HF exposure is unique: calcium gluconate gel is applied to skin burns to bind fluoride ions and prevent systemic toxicity. Because of these hazards, fluorine-related work is limited to trained professionals in well-controlled settings.

History and Etymology

Fluorine’s history is one of caution and persistence. Its name comes from the mineral fluorspar (calcite of calcium fluoride). “Fluorite” was named from the Latin fluere, meaning “to flow,” because it was used as a flux to lower the melting point of ores in metal smelting. Chemists recognized fluorite’s significance early on. In 1670, Georgius Agricola noted that fluorspar produced a corrosive acid (hydrofluoric acid) when treated with sulfuric acid. However, isolating elemental fluorine proved extremely difficult because it was so reactive.

In 1810, the French chemist André-Marie Ampère predicted the existence of an element analogous to chlorine but more reactive, based on electrochemical ideas of the time. Humphry Davy in England became interested. Between 1811–1813, Davy tried to electrolyze aqueous HF (a change that would release F₂ at the anode), but he was unable to obtain the gas or identify it. Instead, his platinum electrodes corroded, and he detected hydrogen at the cathode. Davy proposed the name “fluorine” in 1812 for the hypothetical element, after the Latin fluorum (fluorspar). Other chemists continued the quest: in 1831 the Irish Knox brothers used fluorspar vessels and strong reagents hoping to isolate fluorine; none succeeded. French chemist Edmond Frémy in 1855 used molten calcium fluoride as an electrolyte and observed colorless gas evolving (likely fluorine), but he also failed to collect it in pure form.

The breakthrough came in 1886 with the work of French chemist Henri Moissan. Moissan designed a specialized cell with platinum-iridium anodes and cathodes, cooled with liquid sulfur dioxide. He electrolyzed potassium hydrogen fluoride (KHF₂) in hydrogen fluoride, carefully avoiding moisture. After many trials, on June 26, 1886, Moissan announced he had produced pure fluorine gas for the first time. He described it as a faintly yellow gas with a disagreeable odor, which attacked most materials. Moissan’s method (electrolysis of KHF₂ in anhydrous HF) remains the standard approach today. For this achievement, Moissan was awarded the Nobel Prize in Chemistry in 1906.

Earlier, other scientists had also laid groundwork. Notably, French chemist Louis-Joseph Gay-Lussac and his assistant Louiss Jacques Thénard in 1811 recognized a new acid (HF) by distilling fluorite with sulfuric acid. Another Frenchman, René Just Haüy, studied fluorite’s crystal structure. The Swiss chemist Alfred Werner later developed coordination chemistry partly using fluorides. The symbol “F” was chosen (by August Kekulé, or at least it became conventional) to stand for fluorine in the periodic table devised by the mid-19th century chemists.

Industrial milestones included: in the early 1900s, fluorine chemistry began on a larger scale. The Manhattan Project in the 1940s used fluorine (as UF₆) to produce enriched uranium for the first atomic bombs. In 1938 (USA), Roy Plunkett’s accidental discovery of PTFE launched the age of fluoropolymers. During World War II, fluorine compounds were used in refrigerants and other war efforts. After the war, widespread use of CFC refrigerants and Teflon grew.

Etymologically, the word “fluorine” (from Latin fluorum + the suffix “-ine” for halogens) reflects its origin from fluorspar. In many languages, the element’s name is similar (e.g. French “fluor”, German “Fluor”, etc.). Its symbol F stands simply for its English name. The discovery of fluorine spanned more than a century of international effort and was fraught with danger: Moissan himself later died from long-term exposure to hydrofluoric acid. Fluorine’s history exemplifies the challenge of isolating very reactive elements.

In summary, humanity learned of fluorine long ago, but isolating it safely required advances in materials and technique. Its naming honors the mineral flux which first led to its chemistry. Since the late 19th century, fluorine has gone from a laboratory curiosity to a vital element in modern industry, albeit always watched under strict safety measures.

Data Table

Property	Value
Symbol	F
Atomic Number (Z)	9
Atomic Weight	18.9984
Electron Configuration	[He] 2s² 2p⁵
Common Oxidation State(s)	–1
Group / Period / Block	17 (Halogen) / 2 / p
Valence Electrons	7 (2 in 2s, 5 in 2p)
Phase at STP	Gas (diatomic F₂)
Electronegativity (Pauling)	3.98 (highest of all elements)
First Ionization Energy	1,681 kJ/mol (≈17.42 eV)
Atomic Radius (covalent)	≈60–65 pm
van der Waals Radius	≈135 pm
Density (F₂ gas at STP)	1.696 g/L
Density (liquid, at boiling point)	1.505 g/cm³
Melting Point	53.48 K (–219.67 °C)
Boiling Point	85.03 K (–188.11 °C)
Color/Odor	Pale yellow-green gas; pungent odor
Magnetism	Diamagnetic
Major Stable Isotope	^19F (100% natural abundance; spin 1/2)
Notable Radioisotope	^18F (half-life ≈110 min; used in PET imaging)
Crystal Structure (solid α-F₂)	Monoclinic (C2/c); 4 F₂ molecules/unit cell
Crystal Structure (solid β-F₂)	Cubic (Pm3n)