How is the fluorine gas color

Name, symbol, atomic number Fluorine, F, 9
Group, period, block17, 2, p
Look pale, greenish-yellowish gas
Mass fraction of the earth's envelope 0,03 %
Atomic mass 18.9984 u
Atomic radius (calculated) 50 (42) pm
Covalent radius 71 pm
Van der Waals radius 147 pm
Electron configuration [He] 2s2 2p5
Electrons per energy level 2, 7
1. Ionization energy 1681.0 kJ / mol
2. Ionization energy 3374.2 kJ / mol
3. Ionization energy 6050.4 kJ / mol
4. Ionization energy 8407.7 kJ / mol
5. Ionization energy 11023 kJ / mol
6. Ionization energy 15164 kJ / mol
7. Ionization energy 17868 kJ / mol
Physical state gaseous
Modifications 1
Crystal structure none (gas)
density 1.696 kg m−3 at 273 K
Mohs hardness none (gas)
magnetism diamagnetic
Melting point 53.53 K (−219.62 ° C)
boiling point 85.03 K (−188.12 ° C)
Molar volume 22,40 · 10−3 m3/ mol
Heat of evaporation 3.2698 kJ / mol
Heat of fusion 0.2552 kJ / mol
Vapor pressure

- Pa

Speed ​​of sound 286 m / s
Specific heat capacity 824 J / (kg K) at 293 K
Electric conductivity 0 S / m
Thermal conductivity 0.0279 W / (m K)
Oxidation states −1
Oxides (basicity) (very acidic)
Normal potential 2.87 V (F + e → F)
Electronegativity 4.0 (Pauling scale)
isotopeNHt1/2ZMZE MeVZP


64.49 sε2,76117O


109.77 mβ +, ε1,65618O

100 %



11.00 sβ-7,02520No


4.158 sβ-5,68421No
NMR properties
safety instructions
Hazardous substance labeling
from RL 67/548 / EEC, Appendix I.
R and S phrases R: 7-26-35
S: (1 / 2-) 9-26-36 / 37 / 39-45
As far as possible and customary, SI units are used.
Unless otherwise noted, the data given apply to standard conditions.

fluorine [ˈFluːoːr] is a chemical element in the periodic table of the elements with the symbol F and the atomic number 9. The poisonous, colorless, pale yellow to yellow-green gas in layers greater than one meter is the most electronegative and reactive of all chemical elements.

Its name derives from lat. fluorine "River" from. Used in the plural (“fluores”) the term “flux” (in metallurgy; see Agricola) denotes and in this sense stood for fluorspar, the most important naturally occurring fluorine mineral.


Fluorine in the form of its calcium salt (fluorspar, CaF2) was described by Georgius Agricola in 1530 and mentioned by him in 1556 as an aid for melting ores.[1] It makes ore smelting and slag more fluid and allows them to flow. In 1670 Schwanhard showed glass etching using acid-treated fluorspar. However, all attempts to produce the free halogen failed - sometimes tragically. It was not until 1886 that Henri Moissan succeeded in using the electrolytic decomposition of a solution of potassium difluoride (KHF2) to produce pure fluorine in liquid hydrogen fluoride (HF) at -55 ° C, which was a coincidence, as the potassium difluoride was only added to the hydrogen fluoride to improve the conductivity (pure hydrogen fluoride cannot be electrolytically decomposed).

Fluorine production took off in the Second World War, on the one hand through the development of the atomic bomb in the USA (Manhattan Project). The isotope enrichment of uranium takes place via gaseous uranium hexafluoride (UF6), which is made with the help of elemental fluorine [2][3]. On the other hand, I. G. Farben in Gottow operated a fluorine electrolysis cell at that time, the product of which was supposedly only used to manufacture a new incendiary agent (chlorine trifluoride) for incendiary bombs [4]. Whether uranium could also be processed with this fluorine production in Germany at that time is controversial [5][6].


Elemental fluorine does not occur naturally in nature due to its high reactivity. It therefore occurs under natural conditions almost exclusively in the form of fluorides and various fluorine complex salts (e.g. sodium hexafluoridoaluminate = cryolite); however, a few organisms can also produce organofluorine compounds. The leaves of the South African plant Gifblaar contain the deadly fluoroacetic acid. In the form of its salts, fluorides and fluorido complexes, it is widespread and, for example, also in many waters (0.1–1.5 mg / l F) contain. Fluorspar (CaF) is mainly used for the production of fluorine and fluorochemicals2), which was also mined in many places in Germany in the past. When hydrofluoric acid is dissolved in sulfuric acid, hydrofluoric acid (HF) is formed, which is then electrochemically converted to F.2 can be split. Not inconsiderable amounts of fluorosilicic acid arise in the production of phosphoric acid. Fluorosilicic acid (and the sodium hexafluoridosilicate produced from it with the help of soda) is used directly for water fluoridation in many places in the USA. The natural cryolite deposits in Greenland have been exploited since the 1960s.


Elemental fluorine can be prepared from fluorides in particular by electrochemical means. On an industrial scale, it is produced by the electrolysis of low-melting fluorides - for example KF * xHF - with carbon electrodes in iron or Monel® cells. In the almost continuously carried out industrial electrolysis process, the complexly bound hydrogen fluoride (HF) is converted into hydrogen (H.2) and fluorine (F2) disassembled. The resulting HF loss is compensated for by continuously feeding gaseous HF into the melt.

The raw fluorine that leaves the electrolysis cell is more or less strong with HF, oxygen (O2), Tetrafluoromethane (CF4) and perfluorinated hydrocarbons - primarily from the reaction of the electrode material with the fluorine formed - contaminated and can be cleaned if necessary. The cleaning is done by freezing out (HF and volatile metal fluorides), absorption (HF) and low-temperature distillation (removal of the PFHCs).

The "pure fluorine" obtained in this way mostly still contains traces of HF and is more or less free of O2, Nitrogen (N.2) and CF4. Undiluted, pure fluorine is rarely on the market (problematic handling!). However, the much safer fluorine-inert gas mixtures with a fluorine content of up to about 20%, which are typically delivered to the end user in pressurized gas cylinders and are used on an industrial scale, for example in car tank production, are commercially available.

The electrochemical production of fluorine is not a trivial process. A reliable process can only be guaranteed with sufficient technical knowledge, which not only takes into account the extremely aggressive nature of the process media, but also the difficult electrochemistry of the electrolysis process.

In addition to the traditional production of fluorine in large-scale plants, decentralized fluorine production with systems of small and medium-sized systems has recently emerged.

Fluorine can be produced chemically by adding K2MnF6 with SbF5 at about 150 ° C the unstable MnF4 is released, which in turn is shown in F2 and MnF3 disintegrates [7].



Elemental fluorine is "canary yellow" in its liquefied state[8]In its pure form it is pale yellow in gaseous form [9]. In the liquid state it has a density of 1.51 g / cm at the boiling point3 and a density of 1.639 g / cm at -207 ° C3. It is one of the strongest oxidants that are stable at room temperature. It is the most electronegative element. Under normal conditions it is in the form of F2Molecules. Fluorine spontaneously forms compounds with many other elements. Fluorine even reacts with the noble gases xenon, radon and krypton under special conditions. In contrast to all other halogens, fluorine reacts explosively with hydrogen even as a solid at −200 ° C without light activation. The reaction leads to the formation of hydrogen fluoride.

Many other substances also react vigorously with fluorine, including many hydrogen compounds such as water (H.2O), ammonia (NH3), Monosilane (SiH4), Propane (C.3H8), organic solvents, etc. For example, fluorine turns water into oxygen (O2) and hydrogen fluoride (HF) split. In addition, small amounts of ozone O are produced3 and HOF. The driving force behind all of these reactions is the exothermic formation of hydrogen fluoride.