Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist.^[2] Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding (a weaker bridging link to certain nonmetals).^[3] Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds.^{[note 1]} For many elements (but not all) the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others (elements in certain groups) the highest oxidation states of oxides and fluorides are always equal.^[4]

While an individual fluorine atom has one unpaired electron, molecular fluorine (F₂) has all the electrons paired. This makes it diamagnetic (slightly repelled by magnets) with the magnetic susceptibility of −1.2×10⁻⁴ (SI), which is close to theoretical predictions.^[5] In contrast, the diatomic molecules of the neighboring element oxygen, with two unpaired electrons per molecule, are paramagnetic (attracted to magnets).^[6]

The fluorine–fluorine bond of the difluorine molecule is relatively weak when compared to the bonds of heavier dihalogen molecules. The bond energy is significantly weaker than those of Cl₂ or Br₂ molecules and similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines.^[8] The covalent radius of fluorine of about 71 picometers found in F₂ molecules is significantly larger than that in other compounds because of this weak bonding between the two fluorine atoms.^[9] This is a result of the relatively large electron and internuclear repulsions, combined with a relatively small overlap of bonding orbitals arising due to the small size of the atoms.^[10]

The F₂ molecule is commonly described as having exactly one bond (in other words, a bond order of 1) provided by one p electron per atom, as are other halogen X₂ molecules. However, the heavier halogens' p electron orbitals partly mix with those of d orbitals, which results in an increased effective bond order; for example, chlorine has a bond order of 1.12.^[11] Fluorine's electrons cannot exhibit this d character since there are no such d orbitals close in energy to fluorine's valence orbitals.^[11] This also helps explain why bonding in F₂ is weaker than in Cl₂.^[10]

Reactivity

Reactions with elemental fluorine are often sudden or explosive. Many substances that are generally regarded as unreactive, such as powdered steel, glass fragments, and asbestos fibers, are readily consumed by cold fluorine gas. Wood and even water burn with flames when subjected to a jet of fluorine, without the need for a spark.^[12][13]

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External videos
Cold gas impinging on several substances causes bright flames: video by the University of Nottingham.Extra footage. Fluorine reacting with caesium, video by the Royal Institution. (Both videos filmed at a fluorine laboratory of the University of Leicester.)

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Reactions of elemental fluorine with metals require diverse conditions that depend on the metal. Often, the metal (such as aluminium, iron, or copper) must be powdered because many metals passivate by forming protective layers of the metal fluoride that resist further fluoridation.^[7] The alkali metals can react with fluorine explosively, while the alkaline earth metals react not quite as aggressively. The noble metals ruthenium, rhodium, palladium, platinum, and gold react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F).^[14]

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.^[15] The halogens react readily with fluorine gas^[16] as does the heavy noble gas radon.^[17] The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions, while argon will undergo chemical transformations only with hydrogen fluoride.^[18] Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.^[19] Fluorine reacts with ammonia to form nitrogen and hydrogen fluoride .

Fluorine's chemistry is dominated by its strong tendency to gain an electron. It is the most electronegative element and elemental fluorine is a strong oxidant. The removal of an electron from a fluorine atom requires so much energy that no known reagents are known to oxidize fluorine to any positive oxidation state.^[20]

Therefore, fluorine's only common oxidation state is −1. It differs from this value in elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0, and a few polyatomic ions: the very unstable anions F⁻
₂ and F⁻
₃ with intermediate oxidation states exist at very low temperatures, decomposing at around 40 K.^[21] (In very rare circumstance, fluorine can exist in zero oxidation state other than elemental form - namely, in AuF₇ and in cluster of SF₆⁺ with helium atoms^[22]). Also, the F⁺
₄ cation and a few related species have been predicted to be stable.^[23]

Fluorine forms compounds with all elements except neon and helium. In particular, it forms binary compounds, named fluorides, with all said elements except argon. All of the elements up to einsteinium, element 99, have been checked except for astatine and francium,^[24] and fluorine is also known to form compounds with mendelevium, element 101,^[25] rutherfordium, element 104,^[26] and seaborgium, element 106.^[27]

As a result of its small size and high negative charge density, the fluoride anion is the "hardest" base (i.e., of low polarizability). Because of this, fluorides in real salt crystals often have higher effective charges than oxides of the same metal, even though oxygen's formal charge is twice as great as fluorine's.^{[citation needed]}

As a part of a molecule, it is a part with great inductive effect. In the latter case, it significantly increases the acidity of a molecule: the anion formed after giving the proton off becomes stable as a result. Consider acetic acid and its mono-, di-, and trifluoroacetic derivatives and their pK_a values (4.74, 2.66, 1.24, and 0.23^{[note 2]});^[28] in other words, the trifluoro derative is 33,800 times stronger an acid than acetic.^[29] Fluorine is a principal component of the strongest known charge-neutral acid, fluoroantimonic acid (H
₂FSbF
₆).^[30] There is evidence for an even stronger acid called fluoroauric acid (H
₂FAuF
₆) but it has not proved isolable.^[31]

In a molecule that is composed of a central atoms and fluorines attached to it, the intermolecular bonding is not very strong. Moreover, the dense negative balls that fluorines are repel each other. Therefore, the intermolecular bonding strength falls further down, a result of which is the low melting point of high fluorides.^{[citation needed]}

Fluorine combines with hydrogen to make a compound (HF) called hydrogen fluoride or, especially in the context of water solutions, hydrofluoric acid. The H-F bond type is one of the few capable of hydrogen bonding (creating extra clustering associations with similar molecules). This influences various peculiar aspects of hydrogen fluoride's properties. In some ways the substance behaves more like water, also very prone to hydrogen bonding, than one of the other hydrogen halides, such as HCl.^[32][33][34]

Hydrogen bonding amongst HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase. Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C and −35 °C (−120 °F and –30 °F). HF is miscible with water (will dissolve in any proportion), while the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water also form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C (−40 °F), which is 44 degrees Celsius (79 degrees Fahrenheit) above the melting point of pure HF.^[35]

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant (pK_a) equal to 3.19.^[36] HF's weakness as an aqueous acid is paradoxical considering how polar the HF bond is, much more so than the bond in HCl, HBr, or HI. The explanation for the behavior is complicated, having to do with various cluster-forming tendencies of HF, water, and fluoride ion, as well as thermodynamic issues.^{[note 3]} At great concentrations, a property called homoconjugation is revealed. HF begins to accept fluoride ions, forming the polyatomic ions (such as bifluoride, HF⁻
₂) and protons, thus greatly increasing the acidity of the compound.^[38] Hydrofluoric acid is also the strongest of the hydrohalic acids in acetic acid and similar solvents.^[39] Its hidden acidity potential is also revealed by the fact it protonates acids like hydrochloric, sulfuric, or nitric.^[40] Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass (hydrated only).^[38]

Dry hydrogen fluoride dissolves low-valent metal fluorides readily. Several molecular fluorides also dissolve in HF. Many proteins and carbohydrates can be dissolved in dry HF and can be recovered from it. Most non-fluoride inorganic chemicals react with HF rather than dissolving.^[41]

Metal fluorides are rather dissimilar from other metal halides, adopting distinctive structures. In many respects, metal fluorides are more similar to oxides, often having similar bonding and crystal structures.^[42]

Owing to its high electronegativity, fluorine stabilizes metals in higher oxidation states with high M:halide ratios. Numerous charge-neutral penta- and hexafluorides are known, whereas analogous chlorides and bromides are rarer. The molecular binary fluorides are often volatile, either as solids ^[43] liquids,^[44] or gases^[45] at room temperature.

The solubility of fluorides varies greatly but tends to decrease as the charge on the metal ion increases. Dissolved fluorides produce basic solutions.^[46]

The nonmetal binary fluorides are volatile compounds. They show a great difference between period 2 and other fluorides. For instance, period 2 elements elements fluorides never exceed the octet in their atoms. (Boron is an exception due to its specific position in the periodic table.) Lower-period elements, however, may form hypervalent molecules, such as phosphorus pentafluoride or sulfur hexafluoride.^[78] The reactivity of such species varies greatly—sulfur hexafluoride is inert, while chlorine trifluoride is extremely reactive—but there are some trends based on periodic table locations.

Boron trifluoride is a planar molecule. It has only six electrons around the central boron atom (and thus an incomplete octet), but it readily accepts a Lewis base, forming adducts with lone-pair-containing molecules or ions such as ammonia or another fluoride ion which can donate two more electrons to complete the octet.^[79] Boron monofluoride is an unusual molecule with higher than single B-F bond. The bond order has been described as 1.4 (intermediate between a single and double bond). It is isoelectronic with N₂.^[80]

Silicon tetrafluoride, similar to carbon tetrafluoride and germanium tetrafluoride, adopts a molecular tetrahedral structure.^[81] SiF₄ is stable against heating or electric spark, but reacts with water (even moist air), metals, and alkalis, thus demonstrating weak acidic character.^[82] Reactions with organomagnesium compounds, alcohols, amines, and ammonia yield adduction compounds.^[82] Fluorosilicic acid, a derivative of SiF₄, is a strong acid in aqueous solution (the anhydrous form does not exist).^[83]

Pnictogens (nitrogen's periodic table column) show very similar trends in reactivity and acidity of the highest fluorides (pentafluorides) and most common ones (trifluorides), with the said property increasing down the group: NF₃ is stable against hydrolysis,^[84] PF₃ hydrolyzes very slowly in moist air,^[85] while AsF₃ completely hydrolyzes.^[84] SbF₃ hydrolyzes only partially because of the increasing ionic character of the bond to fluorine. The compounds are weak Lewis bases, with NF₃ again being an exception.^[84] The pentafluorides of phosphorus^[85] and arsenic^[86] are much more reactive than their trifluorides; antimony pentafluoride is such a strong acid that it holds the title of the strongest Lewis acid.^[86] Nitrogen is not known to form a pentafluoride, although the tetrafluoroammonium cation (NF⁺
₄) features nitrogen in the formal oxidation state of +5.^[87] Nitrogen monofluoride is a metastable species that has been observed in laser studies. It is isoelectronic with O₂ and, unusually, like BF, has a higher bond order than single-bonded fluorine.^[2][88]

The chalcogens (oxygen's periodic table column) are somewhat similar: The tetrafluorides of S, Se, and Te hydrolyze and are Lewis acidic. Sulfur and selenium tetrafluorides are molecular while TeF₄ is a polymer.^[89] The hexafluorides are the result of direct fluorination of the elements (compare: other hexahalides of these elements do not even exist). They increase in reactivity with atomic number: SF₆ is extremely inert, SeF₆ is less noble (for example, reacts with ammonia at 200 °C (400 °F)), and TeF₆ easily hydrolyzes to give an oxoacid.^[89] Oxygen's highest fluoride is oxygen difluoride,^[89] but fluorine can theoretically (as of 2012) oxidize it to a uniquely high oxidation state of +4 in the fluorocation: OF⁺
₃.^[90] In addition, several chalcogen fluorides occur which have more than one chalcogen (O₂F₂,^[91] S₂F₁₀,^[92] etc.).

The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF₃, and XF₅. Of the neutral +7 species, only iodine heptafluoride is known.^[93] While chlorine and bromine heptafluorides are not known, the corresponding cations ClF⁺
₆ and BrF⁺
₆, extremely strong oxidizers, are.^[94] Astatine is not well-studied, and although there is a report of a non-volatile astatine monofluoride,^[95] its existence is debated.^[96] Many of the halogen fluorides are powerful fluorinators. Chlorine trifluoride is particularly noteworthy—readily fluorinating asbestos and refractory oxides—and may be even more reactive than chlorine pentafluoride. Used industrially, ClF₃ requires special precautions similar to those for fluorine gas because of its corrosiveness and hazards to humans.^[97][98]

Elements frequently have their highest oxidation state in the form of a binary fluoride. Several elements show their highest oxidation state only in a few compounds, one of which is the fluoride; and some elements' highest known oxidation state is seen exclusively in a fluoride.

For groups 1–5, 13–16 (except nitrogen), the highest oxidation states of oxides and fluorides are always equal. Differences are only seen in chromium, groups 7–10, copper, mercury, and the noble gases. Fluorination allows some elements to achieve relatively low^{[note 4]} highest oxidation states that are otherwise hard to achieve. For example, no binary oxide is known for krypton, but krypton difluoride is well-studied.^[106] At the same time, for some other elements, certain very high oxidation states are known only for the oxygen-based species, not the fluorine-based ones. For the previously mentioned volatile oxides, there are no corresponding hepta- or octafluorides except for rhenium. (For example, ruthenium octafluoride is unlikely to be ever synthesized,^[107] while ruthenium tetroxide has even found an industrial use.^[108]) The main problem that prevents fluorine from forming the highest states in covalent hepta- and octafluorides is that it is hard to attach such a large number of ligands around a single atom; the number of ligands is halved in analogous oxides.^{[109][note 5]} However, octafluoride anions, such as the octafluoroiodate (IF⁻
₈), octafluorozirconate (ZrF⁴⁻
₈), and octafluoroxenate (XeF²⁻
₈) anions are well-known.

The highest oxidation states may be uncommon to everyday life, or even industrial usage. For example, the synthesis of mercury tetrafluoride, the first compound to achieve an oxidation state above +2 for a group 12 element, breaking the filled 5d-shell, again showing the significance of the relativistic effects on the heavy elements, and fueling the debate over whether mercury, cadmium, and zinc are transition metals,^[110] occurred at cryogenic temperatures and the compound decomposes at the temperatures of solid nitrogen.^[111] More unstable still, a rare cobalt(V) species, the CoF⁺
₄ cation, has only been observed in gas phase (with no interactions with other atoms, thus no shown stability in any chemical environment).^[107] The reason why such unstable species exist is complicated, yet can be summarized as follows on the example of the hypothesized NF
₅ molecule: According to the modern calculations, five fluorine atoms and one nitrogen atoms can theoretically arrange themselves in different ways, such as NF
₃ and F
₂, NF^•
₄ and F
_•, NF
₅, etc. The NF
₃+F
₂ system is of the smallest energy (most stable). However, if a NF
₅ molecule was synthesized, it would have to go through a high-energy transitional state, from which it could decay into two molecules. But since the transitional state is higher in energy than the hexatomic molecule, the energy difference would need to be added to reach the transitional state and thus allow the decay. This energy is called reaction activation barrier. (The second decay mode has an analogical position.) Thus if little energy was added (low temperatures), then the compound could exist; however, the synthesis is a serious problem (not yet solved).^[112]

Fluorine is very difficult to form oxoacids with, due to the fact that it will take a lot of energy per mole to join each additional oxygen atom to fluorine owing to its exceptional electronegativity.

The only known oxoacid of fluorine is hypofluorous acid (HOF). It appears as a white-solid below temperatures of -117°C, and a pale-yellow liquid above that temperature. HOF is itself only stable until 0 °C and rapidly decomposes above temperatures above that, being explosive at room temperature. Despite being often found as a short-lasting intermediate in the oxidation of water by fluorine, HOF can be still be stably isolated as a solid and the only hypohalous acid known to be able to do so. Its conjugate base, the hypofluorite anion (OF⁻
), can form certain stable compounds, like lithium hypofluorite (LiOF) and trifluoromethyl hypofluorite (CF
₃OF).

Fluorous acid (HFO
₂), fluoric acid (HFO
₃), and perfluoric acid (HFO
₄) are determined to be much, much more unstable than hypofluorous acid and have yet to be discovered.

The carbon–fluorine chemical bond of the organofluorine compounds is the strongest bond in organic chemistry.^[113] Along with the low polarizability of the molecules, these are the most important factors contributing to the great stability of the organofluorines.^[114]

The carbon–fluorine bond of the smaller molecules is formed in three principal ways: Fluorine replaces a halogen or hydrogen, or adds across a multiple bond. The direct reaction of hydrocarbons with fluorine gas can be dangerously reactive, so the temperature may need to be lowered even to −150 °C (−240 °F).^[115] "Solid fluorine carriers", compounds that can release fluorine upon heating, notably cobalt trifluoride,^[116] may be used instead, or hydrogen fluoride. After the reaction, the molecular size is not changed significantly, as the elements have very similar van der Waals radii.^[114] Direct fluorination becomes even less important when it comes to organohalogens or unsaturated compounds reactions, or when a perfluorocarbon is desired (then HF-based electrolysis is typically used).^[117] In contrast, the fluoropolymers are formed by polymerizing free radicals; other techniques used for hydrocarbon polymers do not work in that way with fluorine.^[118]

The range of organofluorine compounds is diverse, reflecting the inherent complexity of organic chemistry. A vast number of small molecules exist with varying amounts of fluorine substitution, as well as many polymers—research into particular areas is driven by the commercial value of applications.^[119]

• Fluorochemical industry