Английская Википедия:Fluorine compounds

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Файл:Fluorine bridging ligands.svg

Fluorine as straight (a) or bent (b) bridging ligands

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. 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). Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. 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.

Difluorine

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. In contrast, the diatomic molecules of the neighboring element oxygen, with two unpaired electrons per molecule, are paramagnetic (attracted to magnets).Шаблон:Sfn

Halogen bond energies (kJ/mol)Шаблон:Sfn

X	XX	HX	BX3	AlX3	CX4
F	159	574	645	582	456
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

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. 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. 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.Шаблон:Sfn

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. Fluorine's electrons cannot exhibit this d character since there are no such d orbitals close in energy to fluorine's valence orbitals. This also helps explain why bonding in F₂ is weaker than in Cl₂.Шаблон:Sfn

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.Шаблон:Sfn

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.Шаблон:Sfn 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).Шаблон:Sfn

Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.Шаблон:Sfn The halogens react readily with fluorine gasШаблон:Sfn as does the heavy noble gas radon. The heavier 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. Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.Шаблон:Sfn Fluorine reacts with ammonia to form nitrogen and hydrogen fluoride .

Chemical characteristics, effects of presence in a molecule

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.

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 Шаблон:Chem and Шаблон:Chem with intermediate oxidation states exist at very low temperatures, decomposing at around 40 K.Шаблон:Sfn (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). Also, the Шаблон:Chem cation and a few related species have been predicted to be stable.

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, and fluorine is also known to form compounds with mendelevium, element 101, rutherfordium, element 104, and seaborgium, element 106.

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.Шаблон: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);Шаблон:Sfn in other words, the trifluoro derative is 33,800 times stronger an acid than acetic. Fluorine is a principal component of the strongest known charge-neutral acid, fluoroantimonic acid (Шаблон:Chem). There is evidence for an even stronger acid called fluoroauric acid (Шаблон:Chem) but it has not proved isolable.

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.Шаблон:Citation needed

Hydrogen fluoride

HF and H2O similarities
graph showing trend-breaking water and HF boiling points: big jogs up versus a trend that is down with lower molecular weight for the other series members.	graph showing humps of melting temperature, most prominent is at HF 50% mole fraction
Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H2O break trends because of hydrogen bonds.	Freezing point of HF/ H2O mixtures: arrows indicate compounds in the solid state.

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.

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.Шаблон:Sfn

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.Шаблон:Sfn 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. At great concentrations, a property called homoconjugation is revealed. HF begins to accept fluoride ions, forming the polyatomic ions (such as bifluoride, Шаблон:Chem) and protons, thus greatly increasing the acidity of the compound. Hydrofluoric acid is also the strongest of the hydrohalic acids in acetic acid and similar solvents. Its hidden acidity potential is also revealed by the fact it protonates acids like hydrochloric, sulfuric, or nitric. Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass (hydrated only).

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.Шаблон:Sfn

Metal fluorides

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.Шаблон:Sfn

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 Шаблон:Sfn liquids,Шаблон:Sfn or gasesШаблон:Sfn 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.

The fluorides of transition metal elements 25–29
white powder in a tube and on a spoon	brown powder in a tube and on a spoon	pink powder in a tube and on a spoon	light green powder in a tube and on a spoon	light blue powder in a tube and on a spoon
Manganese(II) fluoride	Iron(III) fluoride	Cobalt(II) fluoride	Nickel(II) fluoride	Copper(II) fluoride

Low oxidation state metal fluorides

checkerboard-like lattice of small blue and large yellow balls, going in three dimensions so that each ball has 6 nearest neighbors of opposite type

Sodium fluoride: yellow is fluorine, purple is sodium. They are isoelectronic, but fluorine is bigger because its nuclear charge is lower.

The alkali metals form monofluorides. All are soluble and have the sodium chloride (rock salt) structure,Шаблон:Sfn Because the fluoride anion is basic, many alkali metal fluorides form bifluorides with the formula MHF₂. Among other monofluorides, only silver(I) and thallium(I) fluorides are well-characterized. Both are very soluble, unlike the other halides of those metals.

Unlike the monofluorides, the difluorides may be either soluble or insoluble. Several transition metal difluorides, such as those of copper(II) and nickel(II), are soluble. The alkaline earth metals form difluorides that are insoluble. In contrast, the alkaline earth chlorides are readily soluble.

Many of the difluorides adopt the fluorite structure, named after calcium fluoride (and also adopted by several metal dioxides such as CeO₂, UO₂, ThO₂, etc.), which surrounds each metal cation with 8 fluorides. Some difluorides adopt the rutile structure, named after a form of titanium dioxide and adopted by several other metal dioxides also. The structure is tetragonal and puts metal atoms in octahedral coordination.

Beryllium difluoride is different from the other difluorides. In general, beryllium has a tendency to bond covalently, much more so than the other alkaline earths and its fluoride is partially covalent (although still more ionic than its other halides). BeF₂ has many similarities to SiO₂ (quartz) a mostly covalently bonded network solid. BeF₂ has tetrahedrally coordinated metal and forms glasses (is difficult to crystallize). When crystalline, beryllium fluoride has the same room temperature crystal structure as quartz and shares many higher temperatures structures also.

Beryllium difluoride is very soluble in water, unlike the other alkaline earths. (Although they are strongly ionic, they do not dissolve because of the especially strong lattice energy of the fluorite structure.) However, BeF₂ has much lower electrical conductivity when in solution or when molten than would be expected if it were ionic.Шаблон:SfnШаблон:Sfn

Order and disorder in difluorides
A parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers	cube of 8 yellow atoms with white ones at the holes of the yellow structure
The fluorite structure, adopted by many difluorides	Disordered structure of beryllium fluoride glass (sketch, two dimensions)

Many metals form trifluorides, such as iron, bismuth, the rare-earth elements, and the metals in the aluminium and scandium columns of the periodic table. The trifluorides of many rare earths, as well as bismuth, have the YF₃ structure. Trifluorides of plutonium, samarium (at high temperature), and lanthanum adopt the LaF₃ structure. Iron and gallium trifluorides have the FeF₃ structure, which is similar to rhenium trioxide. Only ScF₃ is cubic (ReO₃) at ambient temperature; this material also has the unusual property of negative thermal expansion, meaning it shrinks on heating, over a quite broad temperature range.

Файл:AuF3structure.jpg

The spiral chains of gold trifluoride

Gold trifluoride adopts a structure of linked –AuF₄– squares that align in a helix (spiral chain). In contrast to gold's distinctly ionic trifluoride, its trichloride and tribromide are volatile dimeric molecules.Шаблон:Sfn Aluminium trifluoride is a high melting point solid which is a monomer in the gas phase, while its other trihalides are low-melting, volatile molecules or linear polymeric chains that form dimers as gases phase.Шаблон:Sfn No trifluoride is soluble in water, but several are soluble in other solvents.

The tetrafluorides show a mixture of ionic and covalent bonding. Zirconium, hafnium, plus many of the actinides form tetrafluorides with an ionic structure that puts the metal cation in an 8-coordinate square antiprism. Melting points are around 1000 °C.Шаблон:Sfn

Titanium and tin tetrafluorides are polymeric, with melting points below 400 °C. (In contrast, their tetrachlorides are molecular and liquids at room temperature.) Vanadium tetrafluoride has a similar structure to tin's and disproportionates at 100–120 °C to the trifluoride and the pentafluoride.

The tetrafluorides of iridium, platinum, palladium, and rhodium all share the same structure which was not known until 1975. They have octahedrally coordinated metal atoms with four of the fluorines shared and two unshared. The melting points, where known, are below 300 °C.

Manganese tetrafluoride is an unstable solid that decomposes even at room temperature. Only one of the two allotropes, α-MnF₄, is understood. In this compound, manganese forms –MnF₆– octahedra which share bridging fluorines to make –Mn₄F₂₀– rings which are then further connected three dimensionally.

Some tetrafluoride structures
4 unit cells shown of a crystal structure that is skewed (not cubic) and has 8 bonds to a fluiride, for each metal (fluorides are 2 coordinate). Hard to resolve though even if you could see the picture.	Файл:SnF4structure.jpg	4 tilted octahedra linked in a square
ZrF4, common tetrafluoride structure	The polymeric structure of SnF4	The –Mn4F20– ring of MnF4

High oxidation state metal fluorides

Metal penta- and higher fluorides are all molecular and hence at least somewhat volatile.

Vanadium, niobium, and tantalum form pentafluorides as their highest charge-neutral fluoride. Vanadium pentafluoride is the only non-volatile high-charged metal fluoride, with vanadium being centers of –VF₆– octahedra. The niobium and tantalum pentafluorides, have the same tetrahedra in their structures, with the difference being the formation of the tetra- (rather than poly-) meric molecules.Шаблон:Sfn

Bismuth's highest fluoride is a volatile penta species that is a powerful fluorinating agent. In the solid state, it is polymeric, consisting of linear chains of octahedra, sharing axial fluorides. In combination with alkali metals, pentavalent bismuth can form hexafluorobismuthate, [BiF₆]⁻, upon reaction with a fluoride donor, either strong (such as NaFШаблон:Sfn) or not (such as XeF₄).

Many metals that form hexafluorides also can form pentafluorides. For instance, uranium, which has a well-known hexafluoride, also forms two different pentafluoride structures. The room-temperature (alpha) form has the same linear chain structure as bismuth pentafluoride. As a molecular (gas) species, UF₅ has a square pyramidal structure.

The structure of bismuth (or alpha-uranium) pentafluoride
Шаблон:Center	Шаблон:Center
Structure of a (XF5)n chain; X=Bi, U.	Packing of chains

The metals that make well-characterized hexafluorides include nine metals in the center of the periodic table (molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, and platinum) along with elements 92–94: uranium, neptunium, and plutonium. At room temperature, tungsten hexafluoride is a gas. Molybdenum hexafluoride and rhenium hexafluoride are liquids. The rest are volatile solids.

test-tube like glass container, with pointy sealed end, filled with white crystals

Uranium hexafluoride in sealed glass

Metal hexafluorides are oxidants because of their tendency to release fluorines: for example, platinum hexafluoride was the first compound to oxidize molecular oxygen and xenon. Polonium also forms a hexafluoride, but it is understudied.Шаблон:Sfn

Rhenium is the only metal known to bond with seven fluorides, which is the record for number of charged ligands for a charge-neutral metal compound. Rhenium heptafluoride adopts a pentagonal bipyramid molecular geometry. Calculations shows that the currently unknown but perhaps possible iridium heptafluorideШаблон:Sfn (report of synthesis is being prepared), technetium heptafluoride, and osmium heptafluorideШаблон:Sfn will also have this structure.

Osmium octafluoride was first reported in 1913, but in 1958 that compound was shown to be actually osmium hexafluoride. A 1993 theoretical study predicted very weak bonds in osmium octafluoride and said that it would be difficult to ever detect experimentally. The study predicted that, if made, OsF₈ would have Os–F bonds of two different lengths.

ball and stick drawing, shows central rhenium with a fluorine directly above and below and then an equatorial belt of 5 surrounding fluorines.

Rhenium heptafluoride

Nonmetal fluorides

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. 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.Шаблон:Sfn 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.

Файл:BF resonance lewis structures.svg

Lewis dot diagram structures show three formal alternatives for describing bonding in boron monofluoride.

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

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, PF₃ hydrolyzes very slowly in moist air,Шаблон:Sfn while AsF₃ completely hydrolyzes. 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. The pentafluorides of phosphorusШаблон:Sfn and arsenic are much more reactive than their trifluorides; antimony pentafluoride is such a strong acid that it holds the title of the strongest Lewis acid. Nitrogen is not known to form a pentafluoride, although the tetrafluoroammonium cation (Шаблон:Chem) features nitrogen in the formal oxidation state of +5. 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.

Файл:Sulfur-tetrafluoride-3D-balls.png

SF₄: The unusual see-saw shape is predicted by VSEPR theory.

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. 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. Oxygen's highest fluoride is oxygen difluoride, but fluorine can theoretically (as of 2012) oxidize it to a uniquely high oxidation state of +4 in the fluorocation: [[trifluorooxonium|Шаблон:Chem]]. In addition, several chalcogen fluorides occur which have more than one chalcogen (O₂F₂,Шаблон:Sfn S₂F₁₀,Шаблон:Sfn 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.Шаблон:Sfn While chlorine and bromine heptafluorides are not known, the corresponding cations Шаблон:Chem and Шаблон:Chem, extremely strong oxidizers, are.Шаблон:Sfn Astatine is not well-studied, and although there is a report of a non-volatile astatine monofluoride, its existence is debated. 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.Шаблон:Sfn

Superacids

Several important inorganic acids contain fluorine. They are generally very strong because of the high electronegativity of fluorine. One such acid, fluoroantimonic acid (HSbF₆), is the strongest charge-neutral acid known. The dispersion of the charge on the anion affects the acidity of the solvated proton (in form of Шаблон:Chem): The compound has an extremely low pK_a of −28 and is 10 quadrillion (10¹⁶) times stronger than pure sulfuric acid. Fluoroantimonic acid is so strong that it protonates otherwise inert compounds like hydrocarbons. Hungarian-American chemist George Olah received the 1994 Nobel Prize in chemistry for investigating such reactions.

Noble gas compounds

The noble gases are generally non-reactive because they have filled electronic shells. Until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett found that platinum hexafluoride reacts with xenon. He called the compound he prepared xenon hexafluoroplatinate, but since then the product has been revealed to be mixture, perhaps monofluoroxenyl(II) pentafluoroplatinate, [XeF]⁺[PtF₅]⁻, monofluoroxenyl(II) undecafluorodiplatinate, [XeF]⁺[Pt₂F₁₁]⁻, and trifluorodixenyl(II) hexafluoroplatinate, [Xe₂F₃]⁺[PtF₆]⁻.Шаблон:Sfn Bartlett's fluorination of xenon has been highly praised. Later in 1962, xenon was found to react directly with fluorine to form the di- and tetrafluorides. Since then, other noble gas fluorides have been reported.

black-and-white photo showing icy-looking crystals in a dish

Xenon tetrafluoride crystals

The binary compounds xenon include xenon difluoride, xenon tetrafluoride, and xenon hexafluoride.Шаблон:Sfn Xenon forms several oxyfluorides, such as xenon oxydifluoride, XeOF₂, by hydrolysis of xenon tetrafluoride.Шаблон:Sfn Its lighter neighbor, krypton also forms well-characterized compounds, e.g., krypton difluoride. Krypton tetrafluoride was reported in 1963, but was subsequently shown to be a mistaken identification; the compound seems to be very hard to synthesize now (although even the hexafluoride may exist).

In accordance with the periodic trends, radon is more reactive toward fluorine. Radon difluoride has been claimed. The lighter noble gases (helium through argon) do not form stable binary fluorides.

Highest oxidation states: fluorine versus oxygen

Ruthenium's highest fluoride and oxide
A molecule diagram, with a blue sphere being connected with a stick to 6 yellow-green ones to form an octahedron	A molecule diagram, with a blue sphere being connected with two sticks to 4 red ones to form a tetrahedron
Ruthenium hexafluoride: Six fluorines fit around the ruthenium but only make a +6 oxidation state.	Ruthenium tetroxide: Four oxygens fit around the ruthenium, making a +8 oxidation state.

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 highest oxidation states that are otherwise hard to achieve. For example, no binary oxide is known for krypton, but krypton difluoride is well-studied.Шаблон:Sfn 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, while ruthenium tetroxide has even found an industrial use.) 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.Шаблон:Sfn However, octafluoride anions, such as the octafluoroiodate (Шаблон:Chem), octafluorozirconate (Шаблон:Chem), and octafluoroxenate (Шаблон:Chem) 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, occurred at cryogenic temperatures and the compound decomposes at the temperatures of solid nitrogen. More unstable still, a rare cobalt(V) species, the Шаблон:Chem cation, has only been observed in gas phase (with no interactions with other atoms, thus no shown stability in any chemical environment). The reason why such unstable species exist is complicated, yet can be summarized as follows on the example of the hypothesized [[nitrogen pentafluoride|Шаблон:Chem]] molecule: According to the modern calculations, five fluorine atoms and one nitrogen atoms can theoretically arrange themselves in different ways, such as Шаблон:Chem and Шаблон:Chem, Шаблон:Chem and Шаблон:Chem, Шаблон:Chem, etc. The Шаблон:Chem+Шаблон:Chem system is of the smallest energy (most stable). However, if a Шаблон:Chem 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).Шаблон:Sfn

Fluorine oxoacids

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 (Шаблон:Chem). 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 (Шаблон:Chem), can form certain stable compounds, like lithium hypofluorite (Шаблон:Chem) and trifluoromethyl hypofluorite (Шаблон:Chem).

Fluorous acid (Шаблон:Chem), fluoric acid (Шаблон:Chem), and perfluoric acid (Шаблон:Chem) are determined to be much, much more unstable than hypofluorous acid and have yet to be discovered.

Organofluorine compounds

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

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).Шаблон:Sfn "Solid fluorine carriers", compounds that can release fluorine upon heating, notably cobalt trifluoride,Шаблон:Sfn 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.Шаблон:Sfn 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).Шаблон:Sfn In contrast, the fluoropolymers are formed by polymerizing free radicals; other techniques used for hydrocarbon polymers do not work in that way with fluorine.Шаблон:Sfn

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.

Some important fluorocarbons

Файл:Fluoromethane-3D-vdW.png	Файл:1,1,1,2-tetrafluoroethane-3D-vdW.png	Файл:Dichlorodifluoromethane-3D-vdW.png	Файл:TfOH-3D-vdW.png	Файл:PFOS-anion-3D-vdW.png	Файл:PTFE-3D-vdW.png
Fluoromethane (methyl fluoride), used in semiconductor processing	1,1,1,2-Tetrafluoroethane (R-134a), a HFC	Dichlorodifluoromethane (R-12 refrigerant), a CFC	Trifluoromethanesulfonic acid (triflic acid), a superacid	Perfluorooctanesulfonic acid, a fluorosurfactant (the anion is depicted)	A section of polytetrafluoroethylene (Teflon) polymer

Small molecules

Monofluoroalkanes (alkanes with one hydrogen replaced with fluorine) may be chemically and thermally unstable, yet are soluble in many solvents; but as more fluorines are in instead of hydrogens, the stability increases, while melting and boiling points, and solubility decrease. While the densities and viscosities are increased, the dielectric constants, surface tensions, and refractive indices fall.Шаблон:Sfn

Partially fluorinated alkanes are the hydrofluorocarbons (HFCs). Substituting other halogens in combination with fluorine gives rise to chlorofluorocarbons (CFCs) or bromofluorocarbons (BFCs) and the like (if some hydrogen is retained, HCFCs and the like). Properties depend on the number and identity of the halogen atoms. In general, the boiling points are even more elevated by combination of halogen atoms because the varying size and charge of different halogens allows more intermolecular attractions. As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they do not have carbon–hydrogen bonds to react and released halides quench flames.

two six-membered rings joined on a side, all F substituted

Perfluorodecalin (the trans isomer is shown), a perfluorocarbon that is a liquid at room temperature. It boils at a lower temperature than its hydrocarbon analog, decalin.

When all hydrogens are replaced with fluorine to achieve perfluoroalkanes, a great difference is revealed. Such compounds are extremely stable, and only sodium in liquid ammonia attacks them at standard conditions. They are also very insoluble, with few organic solvents capable of dissolving them.Шаблон:Sfn

However, if a perfluorocarbon contains double or triple bonds (perfluoroalkenes or -alkynes), a very reactive towards ligand accepting result, even less stable than corresponding hydrocarbons.Шаблон:Sfn Difluoroacetylene, which decomposes even under liquid nitrogen temperatures, is a notable example. If such a molecule is asymmetric, then the more fluorinated carbon is attacked, as it holds positive charge caused by the C–F bonds and is shielded weaklyШаблон:Sfn (similarly to that how unsaturated hydrocarbons attacked by HF add hydrogen to the more hydrogen-rich atom per Markovnikov's ruleШаблон:Sfn).

Perfluorinated compounds, as opposed to perfluorocarbons, is the term used for molecules that would be perfluorocarbons—only carbon and fluorine atoms—except for having an extra functional group (even though another definition exists). They share most of perfluorocarbon properties (inertness, stability, non-wettingness and insolubility in water and oils, slipperiness, etc.), but may differ because of the functional group properties, although the perfluorocarbon tail differ the group-specific properties as compared to those of hydrocarbon-tailed compounds.

The perfluoroalkanesulfonic acids are also very notable for their acidity. The sulfonic acid derivative, trifluoromethanesulfonic acid, is comparable in strength to perchloric acid.Шаблон:Sfn These compounds lower surface energy; for this reason, they, especially perfluorooctanesulfonic acid (PFOS, formerly the active component in brand "Scotchgard") have found industrial use as surfactants.Шаблон:Sfn

If a perfluorinated compound has a fluorinated tail, but also a few non-fluorinated carbons (typically two) near the functional group, it is called a fluorotelomer (such molecules are commercially treated as perfluorinated), but such molecules are more of industrial value than chemical. The chain end may similarly be attached to different functional groups (via the hydrogenized terminal carbon), such as hydroxyl resulting in fluorotelomer alcohols, sulfonate resulting in fluorotelomer sulfonates, etc.

Polymers

Fluoropolymers are similar in many regards with smaller molecules; adding fluorine to a polymer affects the properties in the same manner as in small molecules (increasing chemical stability, melting point, reducing flammability, solubility, etc.). Each fluoropolymer has own characteristic properties, though.Шаблон:Sfn

The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit:–CF₂–. PTFE has a backbone of carbons single bonded in a long chain, with all side bonds to fluorines. It contains no hydrogens and can be thought of as the perfluoro analog of polyethylene (structural unit: –CH₂–). PTFE has high chemical and thermal stability, as expected for a perfluorocarbon, much stronger than polyethylene. Its resistance to van der Waals forces makes PTFE the only known surface to which a gecko cannot stick. The compound, however, lacks an ability to transform upon melting, which is not a problem for various PTFE derivatives, namely FEP (fluorinated ethylene propylene, with some fluorines replaced with the–CF₃ group) or PFA (perfluoroalkoxy, some fluorines replaced with –OCF₃). They share most properties with PTFE, but there are still differences, namely maximum usage temperature (highest for the non-flexible PTFE).Шаблон:Sfn

There are other fluoroplastics other than perfluorinated. Polyvinylidene fluoride (PVDF, structural unit: –CF₂CH₂–), is an analog of PTFE with half the fluorines. PVF (polyvinyl fluoride, structural unit: –CH₂CHF–) contains one one-fourth the fluorines of PTFE. Despite this, it still has many properties of more fluorinated compounds.Шаблон:Sfn PCTFE (polychlorotrifluoroethylene, structural unit: –CF₂CFCl–) is another important compound. It differs from PTFE by having a quarter of fluorine replaced with chlorines, yet this difference brings even greater hardness, creep resistance, and moisture persistence.Шаблон:Sfn

Mild fluorination of polyethylene gives does not make all of the plastic lose its hydrogens for fluorine; only a thin layer (0.01 mm at maximum) is then affected. This is somewhat similar to metal passivation: the bulk properties are not affected, but the surface properties are, most notably, a greater vapor barrier. Therefore, they are a cheaper alternative to the perfluoro plastics if only surface is important.

Skeletal chemical formula

The complex unit structure of Nafion polymer

Nafion is a structurally complicated polymer. It has a PTFE-like backbone, but also contains side chains of perfluoro ether that end in sulfonic acid(–SO₂OH) groups. It also possesses great chemical stability, while exact properties vary with morphology. However, because of the difficult chemical structure, it is also relatively easily converted to an ionomer (shows conductivity) by adding cations like Na⁺ or by converting into the sulfonic acid rather than the given sulfonyl fluoride. The conductivity is due to that the main carbon chain separates from the side chains, thus forming polar and non-polar regions. This form is also very hydroscopic.

Fluoroelastomers, like other elastomers (artificial rubbers), consist of disordered polymer chains connected in three dimensions. The main challenges in making fluorelastomers are cross-linking (reacting the unreactive polymers), as well as removing the HF formed during curing. There are three main families of fluoroelasters. VDF/HFP is a copolymer system of vinylidene fluoride and (at least 20%) hexafluoropropylene. TFE/propylene is another copylymer system with better chemical resistance to some solvents. TFE/PMVE (perfluoromethylvinyl ether) is a copolymer system which creates a perfluorinated fluoroelastomer.

See also

• Fluorochemical industry

Indexed references

• Шаблон:Cite report
• Шаблон:Cite book
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• Шаблон:Cite book
• Шаблон:Cite book
• Шаблон:Cite book
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• Шаблон:Cite book
• Шаблон:Cite book
• Шаблон:Cite book