The history of fluorine isolation is marked by decades of dangerous and sometimes fatal attempts. Fluorine compounds, particularly the mineral fluorite (calcium fluoride, CaF₂), have been known since the 16th century. In 1810, the French chemist André-Marie Ampère (1775-1836) suggested the existence of a new element analogous to chlorine in hydrofluoric acid. For over 70 years, many chemists attempted to isolate fluorine, but its extreme reactivity made the task extremely perilous. Several researchers were seriously poisoned or died during these attempts, including the Knox brothers, Thomas and George, in Ireland.
It was not until 1886 that the French chemist Henri Moissan (1852-1907) finally succeeded in isolating fluorine gas by electrolysis of a mixture of hydrofluoric acid and potassium bifluoride in a platinum-iridium apparatus cooled. This feat earned him the Nobel Prize in Chemistry in 1906. The name fluorine derives from the Latin fluere (to flow), referring to the use of fluorite as a flux in metallurgy to lower the melting point of ores.
Fluorine (symbol F, atomic number 9) is a halogen of group 17 in the periodic table, consisting of nine protons, usually ten neutrons (for the stable isotope), and nine electrons. The only natural stable isotope is fluorine-19 \(\,^{19}\mathrm{F}\) (100% natural abundance).
At room temperature, fluorine exists as a pale yellow diatomic gas (F₂), with a pungent and acrid odor, extremely toxic and corrosive. Fluorine is the most electronegative chemical element of all (Pauling electronegativity: 3.98), meaning it attracts electrons more strongly than any other element. This property makes fluorine the most powerful oxidant and the most aggressive reactive known. F₂ gas has a density of about 1.696 g/L at standard temperature and pressure.
The temperature at which the liquid and solid states can coexist (melting point): 53.48 K (−219.67 °C). The temperature at which it transitions from liquid to gas (boiling point): 85.03 K (−188.12 °C). Liquid fluorine has a characteristic bright yellow color.
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Fluorine-18 — \(\,^{18}\mathrm{F}\,\) | 9 | 9 | 18.000938 u | Unnatural | 109.77 minutes | Radioactive β\(^+\) decay to \(\,^{18}\mathrm{O}\); widely used in PET scans (FDG labeled with fluorine-18). |
| Fluorine-19 — \(\,^{19}\mathrm{F}\,\) | 9 | 10 | 18.998403 u | 100 % | Stable | Only stable isotope of fluorine; present in all natural and artificial fluorinated compounds. |
| Fluorine-20 — \(\,^{20}\mathrm{F}\,\) | 9 | 11 | 19.999981 u | Unnatural | 11.00 s | Radioactive β\(^-\) decay to \(\,^{20}\mathrm{Ne}\); artificially produced in accelerators. |
| Fluorine-21 — \(\,^{21}\mathrm{F}\,\) | 9 | 12 | 20.999949 u | Unnatural | 4.158 s | Radioactive β\(^-\); used in nuclear research. |
| Fluorine-17 — \(\,^{17}\mathrm{F}\,\) | 9 | 8 | 17.002095 u | Unnatural | 64.49 s | Radioactive β\(^+\); positron emitter used in medical imaging. |
| Other isotopes — \(\,^{14}\mathrm{F}-\,^{16}\mathrm{F},\,^{22}\mathrm{F}-\,^{31}\mathrm{F}\) | 9 | 5-7, 13-22 | — (resonances) | Unnatural | \(10^{-22}\) — 5 s | Very unstable states observed in nuclear physics; decay by particle emission or β radioactivity. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Fluorine has 9 electrons distributed across two electron shells. Its full electronic configuration is: 1s² 2s² 2p⁵, or simplified as: [He] 2s² 2p⁵. This configuration can also be written as: K(2) L(7).
K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 7 electrons distributed as 2s² 2p⁵. The 2s orbitals are complete, while the 2p orbitals contain only 5 out of 6 possible electrons. Thus, only 1 electron is missing to achieve the stable neon configuration with 8 electrons (octet).
The 7 electrons in the outer shell (2s² 2p⁵) are the valence electrons of fluorine. This configuration explains its chemical properties:
By gaining 1 electron, fluorine forms the F⁻ ion (oxidation state -1), its unique and systematic oxidation state in all its compounds, thus adopting the stable neon configuration [Ne].
Fluorine cannot exhibit any positive oxidation state because it is the most electronegative element of all chemical elements (electronegativity of 4.0 on the Pauling scale).
The oxidation state 0 corresponds to difluorine F₂, its natural molecular form, where two fluorine atoms share a pair of electrons.
The electronic configuration of fluorine, with 7 electrons in its valence shell, classifies it among the halogens and makes it the most reactive element in the periodic table. This structure gives it exceptional characteristic properties: maximum chemical reactivity (fluorine reacts with practically all elements, including the heaviest noble gases and even water), the highest electronegativity of all elements (unmatched ability to attract electrons), and the most powerful oxidizing power known. Fluorine exclusively forms the fluoride ion F⁻ by capturing an electron to complete its octet. Its small atomic size and strong effective nuclear charge explain its exceptional avidity for electrons. Difluorine F₂ is an extremely corrosive and dangerous pale yellow-green gas that violently attacks almost all materials. Despite its extreme reactivity, fluorine and its compounds have important applications: sodium fluoride (NaF) is added to drinking water and toothpaste to prevent tooth decay, fluorinated compounds are used as refrigerants (although CFCs are banned), polytetrafluoroethylene (PTFE, Teflon) is a highly resistant non-stick polymer, and hydrofluoric acid HF is used in glass etching and metallurgy.
Fluorine has seven valence electrons and needs only one electron to complete its outer shell. This configuration, combined with its record electronegativity, makes fluorine an extraordinarily aggressive oxidant that reacts spontaneously with almost all chemical elements, including some noble gases (xenon, krypton, radon) under appropriate conditions. Fluorine can even oxidize oxygen to form oxygen difluoride (OF₂), a compound where oxygen is in an unusual positive oxidation state.
Fluorine reacts violently with most organic and inorganic substances, often with spontaneous ignition. Water reacts explosively with fluorine to produce hydrofluoric acid (HF), oxygen, and ozone. Metals ignite on contact with fluorine gas, forming metal fluorides. Even ordinary glass is attacked by fluorine, requiring the use of special containers made of passivated metals (nickel, copper, stainless steel) whose surface is covered with a thin protective layer of fluoride.
Fluorine forms the fluoride ion (F⁻) in its ionic compounds and extremely strong covalent bonds in its covalent compounds. The C-F (carbon-fluorine) bond is one of the strongest and most stable chemical bonds in organic chemistry, giving fluorocarbon compounds (such as Teflon) exceptional chemical and thermal stability. Hydrofluoric acid (HF) is a weak acid in aqueous solution but extremely corrosive as it can dissolve glass and penetrate deeply into biological tissues.
Despite its toxicity and reactivity in elemental form, the fluoride ion (F⁻) at low concentrations plays a beneficial role by strengthening tooth enamel by converting it into fluorapatite, which is more resistant to acid attacks. This is why fluoride is added to toothpaste and drinking water in many countries to prevent tooth decay.
Fluorine is a relatively rare element in the universe, with a cosmic abundance about 400 times lower than that of oxygen. This rarity contrasts with fluorine's position in the periodic table between oxygen (very abundant) and neon (moderately abundant), creating what is sometimes called the "fluorine deficit" in the cosmos.
Unlike most other light elements, the astrophysical origin of fluorine has long been mysterious. Fluorine cannot be produced efficiently by Big Bang primordial nucleosynthesis or by the usual fusion reactions in stars. Recent research suggests that fluorine is mainly produced by two processes:
Nucleosynthesis in AGB stars (asymptotic giant branch stars, 2-8 solar masses) appears to be the main source. In these evolved stars, fluorine is produced by neutron capture on nitrogen-14 and oxygen-18, followed by nuclear reactions involving protons. These stars then disperse fluorine into the interstellar medium via their powerful stellar winds and matter ejections.
Neutrinos produced during supernovae can also contribute to fluorine production. When a massive star explodes as a supernova, the intense neutrino flux can induce nuclear reactions (nu process) that convert neon-20 into fluorine-19 and sodium-23. This contribution remains debated but could explain part of the abundance of fluorine in the universe.
Fluorine has also been detected in the atmospheres of some carbon-rich evolved stars and in a few planetary nebulae. Variations in fluorine abundance in different stellar populations allow astronomers to constrain models of nucleosynthesis and galactic chemical evolution.
In the solar system, fluorine is mainly found as fluoride in terrestrial minerals (fluorite, apatite) and in some meteorites. The Earth contains about 0.06% fluorine in its crust, mainly in minerals such as fluorite (CaF₂), apatite (Ca₅(PO₄)₃F), and topaz (Al₂SiO₄(F,OH)₂).
The detection of fluorine in interstellar space is difficult because gaseous fluorine (F₂) and simple fluorinated compounds are rare. Hydrofluoric acid (HF) has been detected in some molecular clouds and circumstellar envelopes, providing information on the chemistry of fluorine in space.
N.B.:
The "fluorine paradox" remarkably illustrates the duality of this extraordinary element. In its elemental form (F₂), fluorine is one of the most dangerous chemicals ever handled: toxic, corrosive, reactive with almost everything, and responsible for fatal accidents throughout its history. Yet, in the form of the fluoride ion (F⁻) at low concentrations, it becomes beneficial for human dental health. Similarly, synthetic organofluorine compounds are among the most chemically stable and inert substances ever created (Teflon, Gore-Tex), in complete contrast to the reactivity of elemental fluorine. This spectacular transformation of chemical properties between the free element and its compounds is more pronounced for fluorine than for any other element. Fluorine thus embodies a fundamental lesson in chemistry: the properties of an element in its free state can be radically different from those of its compounds, and the toxicity or danger of a substance depends entirely on its chemical form and concentration.
