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Compounds

[edit]
Fluorine as straight (a) or bent (b) bridging ligands.[1]

Fluorine's only common oxidation state is −1.[note 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.[4] Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding.[5] Fluorine has a rich chemistry including inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds.[note 2]

Inorganic

[edit]

Hydrogen fluoride

[edit]
Main articles: Hydrogen fluoride and Hydrofluoric acid
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. 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 hydrogen–fluorine 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 other hydrogen halides, such as HCl.[6][7][8]

Hydrogen bonding between 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 (−120 °F) and −35 °C (−30 °F). HF is fully miscible with water (dissolves 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 °C (79 °F) above the melting point of pure HF.[9]

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant (pKa) of 3.19.[10] 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.[12] Hydrofluoric acid is also the strongest of the hydrohalic acids in acetic acid and similar solvents.[13] Its hidden acidity potential is also revealed by the fact it protonates acids like hydrochloric, sulfuric, or nitric.[14] Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass (hydrated only).[12]

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 recovered from it. Most non-fluoride inorganic chemicals react with HF rather than dissolving.[15]

Metal fluorides

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Main article: Structural chemistry of the metal fluorides

Metal fluorides have similarities with other metal halides but are more ionic. In many respects, metal fluorides differ from other metal halides (chlorides, bromides, iodides), which are very similar to each other. Instead, fluorides are more similar to oxides, often having similar bonding and crystal structures.[16]

The solubility of fluorides varies greatly but tends to decrease as the charge on the metal ion increases. Dissolved fluorides produce basic solutions. (F- is a weak base because HF is a weak acid.)[17] Like hydroxides, fluorides may be viewed as basic, amphoteric, and acidic, with the acidity property generally increasing with the oxidation state of the metal;[18] a fluoride is a much weaker base than a hydroxide, though.[19]

Alkali metals fluorides are very ionic and soluble (and are similar to a few oxides such as CaO.[20]), and give off the fluorides easily.[21] Other ionic metals fluorides show clear difference than their chlorides, being more ionic in nature. Thallium[22] and silver[23] monofluorides are soluble, while the chlorides are not (however, alkaline earth metals' are opposite in that respect[23]); aluminium[24] and gold[25] form ionic trifluorides while the trichlorides are covalent.

However, as the oxidation state of the metal grows, fluorides begin to show covalent character: Titanium,[26] niobium,[27] and vanadium[28] tetrafluorides are polymeric, and some transition metals tetrafluorides show unusual structures intermediate between ionic and monomeric covalent (notably iridium tetrafluoride featuring both shared and unshared fluorines[29]). They normally have low melting points or decompose easily. Despite the metal's relatively low oxidation state, beryllium fluoride also tends to bond covalently, so it shows a significant covalent character. BeF2 is similar to SiO2 (quartz) in many respects: both are mostly covalently bonded network solids, commonly form glasses rather than crystallize, share the same room temperature crystal structure, etc. It is very soluble in water, unlike the other alkaline earths.[23][30][31][32][33]

Metal penta- and higher fluorides are all covalently bonded and volatile. This behavior contrasts with the corresponding oxides. Oxygen is a weaker oxidant and inherently more likely to form covalent bonds, but it only forms molecules with five metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide, osmium tetroxide, and iridium tetroxide[34]). Fluorine forms molecules with fifteen metals because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can.

Higher fluorides are normally discrete molecules. This behavior contrasts with the corresponding oxides. Oxygen is a weaker oxidant and inherently more likely to form covalent bonds, but it only forms molecules with five metals (manganese heptoxide, technetium heptoxide, ruthenium tetroxide, osmium tetroxide, and iridium tetroxide[34]). Fluorine forms molecules with fifteen metals because its small size and single charge as an ion allows surrounding metal atoms with more fluorines than oxygen can. These compounds are highly reactive, acting like acids. For example, platinum hexafluoride was the first compound to oxidize molecular oxygen[35] and xenon.[36]

Major fluoride types
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 ?????? Ball and stick drawing showing central rhenium with a fluorine directly above and below and then an equatorial belt of 5 surrounding fluorines.
Sodium fluoride, ionic (NaCl-like) Bismuth pentafluoride, polymeric Rhenium heptafluoride, discrete molecule

Nonmetal fluorides

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The nonmetal binary fluorides are volatile compounds. Period 2 elements fluorides differ greatly from other fluorides. For instance, they always follow the octet rule. (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.[37] 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.

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

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.[38] Boron monofluoride is an unstable molecule with an unusual (higher than single) bond to fluorine. The bond order has been described as 1.4, which is a superposition of the three possible options: 1, 2, and 3 (see the image). It is isoelectronic with N2.[39]

Silicon tetrafluoride, similar to carbon tetrafluoride and germanium tetrafluoride, adopts a molecular tetrahedral structure.[40] SiF4 is stable against heating or electric spark, but reacts with water (even moist air), metals, and alkalis, thus demonstrating weak acidic character.[41] Reactions with organomagnesium compounds, alcohols, amines, and ammonia yield adduction compounds.[41] Fluorosilicic acid, a derivative of SiF4, is a strong acid in aqueous solution (the anhydrous form does not exist).[42]

Pnictogens (nitrogen's periodic table column) fluorides, both the highest (pentafluorides) and most common (trifluorides), become more reactive and acidic as the pnictogen becomes heavier: NF3 is stable against hydrolysis,[43] PF3 hydrolyzes very slowly in moist air,[44] while AsF3 completely hydrolyzes.[43] 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.[43] The pentafluorides of phosphorus[44] and arsenic[45] are much more reactive than their trifluorides; antimony pentafluoride is the strongest Lewis acid of all charge-neutral compounds.[45] Nitrogen is not known to form a pentafluoride, although the tetrafluoroammonium cation (NF+
4) features nitrogen in the formal oxidation state of +5.[46] Nitrogen monofluoride is a metastable species that has been observed in laser studies. It is isoelectronic with O2 and like BF, unusually, has a higher bond order than single-bonded fluorine.[4][47]

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

The chalcogens (oxygen's periodic table column) are somewhat similar: The tetafluorides are thermally unstable and hydrolyze, and are also ready to use their lone pair to form adducts to other (acidic) fluorides. Sulfur and selenium tetrafluorides are molecular while TeF4 is a polymer.[48] The hexafluorides are the product of direct fluorination of the elements (compare: other hexahalides of the elements do not even exist). They increase in reactivity with atomic number: SF6 is extremely inert, SeF6 is less noble, and TeF6 easily hydrolyzes to give an oxoacid.[48] Oxygen's highest fluoride is oxygen difluoride,[48] but fluorine can theoretically (as of 2013) oxidize it to a uniquely high oxidation state of +4 in the fluorocation: OF+
3.[49]

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.[50] While chlorine and bromine heptafluorides are not known, the corresponding cations ClF+
6 and BrF+
6, extremely strong oxidizers, are.[51] Astatine is not well-studied, and although there is a report of a non-volatile astatine monofluoride,[52] its existence is debated.[53] 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.[54][55]

Superacids

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Main article: Superacid

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.[56] The dispersion of the charge on the anion affects the acidity of the solvated proton (in form of H
2F+
): The compound has an extremely low pKa of −28 and is 10 quadrillion (1016) times stronger than pure sulfuric acid.[56] Fluoroantimonic and related acids are so strong that they protonate otherwise inert compounds like hydrocarbons; an early superacid was given the nickname (and eventual tradename) of "Magic acid" after a 1966 demonstration in which this fluorinated acid dissolved a paraffin wax candle.[57] Hungarian-American chemist George Olah received the 1994 Nobel Prize in chemistry for investigating such reactions.[58]

Noble gas compounds

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Main article: Noble gas compound

The noble gases are generally non-reactive because they have full electronic shells, which are extremely stable. Until the 1960s, no chemical bond with a noble gas was known. In 1962, Neil Bartlett conducted the reaction of platinum hexafluoride and xenon. He called the compound he prepared xenon hexafluoroplatinate, but since then the product has been revealed to be mixture of different chemicals. Bartlett probably synthesized a mixture of monofluoroxenyl(II) pentafluoroplatinate, [XeF]+[PtF5]–, monofluoroxenyl(II) undecafluorodiplatinate, [XeF]+[Pt2F11]–, and trifluorodixenyl(II) hexafluoroplatinate, [Xe2F3]+[PtF6]–.[59] Bartlett's fluorination of xenon has been called one of the ten most beautiful experiments in the history of chemistry.[60] Later in 1962, xenon was reported to react directly with fluorine to form the di- and tetrafluorides. Since then, chemists have made extensive efforts to form other noble gas fluorides.

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

Having started the noble gases–fluorine compounds class, xenon has the most well-studied such compounds. Its binary compounds include xenon difluoride, xenon tetrafluoride, and xenon hexafluoride, as well as their derivatives.[61] Xenon forms several oxyfluorides, such as xenon oxydifluoride, XeOF2, by reaction of xenon tetrafluoride with water.[62] Its lighter homolog, krypton, is the only other one to form a well-established compound: krypton difluoride. Krypton tetrafluoride was reported in 1963,[63] but was subsequently shown to be a mistaken identification; the compound seems to be very hard to synthesize now (however, even the hexafluoride may theoretically exist).[64]

In strong accordance with the periodic trends, radon is notably more reactive to oxidizers, including fluorine. It has been shown to readily react with fluorine to form a solid compound, which is generally thought to be radon difluoride. The exact composition is uncertain as the compound is prone to decomposition. Calculations indicate that radon difluoride may be ionic, unlike all other binary noble gas fluorides.[52]

The lighter noble gases (helium through argon) do not form stable binary fluorides. Argon forms no binary fluoride, but reacts in extreme conditions with hydrogen fluoride to form argon fluorohydride; it is the only "stable"[note 4] argon compound. Helium and neon do not form any stable chemical compounds at all. Helium forms helium fluorohydride; it has been shown to be unstable in gas phase, but it may be stable under enormous pressure.[65] Neon, the least reactive element,[note 5] is not now expected to form a stable compound capable of synthesis.

Ununoctium, the last currently known group 18 element, is predicted to form ununoctium difluoride, UuoF
2, and ununoctium tetrafluoride, UuoF
4, which is likely to have the tetrahedral molecular geometry.[68] However, only a few atoms of ununoctium have been synthesized,[69] and its chemical properties have not been examined.

Highest oxidation states: fluorine versus oxygen

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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, 10, 13–16, the highest oxidation states of oxides and fluorides are always equal. Differences are only seen in chromium, groups 7–9, copper, mercury, and the noble gases. Fluorination allows elements to achieve relatively low[note 6] oxidation states that are, however, hard to achieve. For example, no binary oxide is known for krypton, but krypton difluoride is well-studied.[70] At the same time, very high oxidation states are known for oxygen-based species only. For the previously mentioned volatile oxides, there are no corresponding hepta- or octafluorides. (For example, ruthenium octafluoride is unlikely to be ever synthesized,[71] while ruthenium tetroxide has even found an industrial use.[72]) 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.[73][note 7] 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 of theoretical interest: mercury tetrafluoride is the first compound to achieve an oxidation state above +2 for a group 12 element, breaking the filled 5d-shell. This again shows the significance of the relativistic effects on the heavy elements and fueled the debate over whether mercury, cadmium, and zinc are transition metals.[75] They, however, may be uncommon to everyday life, or even industrial usage. The synthesis of the same mercury tetrafluoride occurred at cryogenic temperatures and the compound decomposes at the temperatures of solid nitrogen.[76] More unstable still, the only 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).[71] The reason why such unstable species exist is basically that such an unstable molecule, once synthesized, cannot directly decay to the low energy state, but only through a transitional state. Normally, the transitional state is higher in energy than the original unstable one, and cannot be reached if the molecule is not given enough energy to offset the difference (low temperatures). However, synthesis of such molecules is still a great problem (in some cases not yet solved).[77]

Organic

[edit]
Main article: Organofluorine chemistry

Carbon–fluorine chemical bond of the organofluorine compounds is the strongest bond in organic chemistry.[78] Along with the low polarizability of the molecules, these are the most important factors contributing to the great stability of the organofluorines.[79]

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)[80]. Hydrogen fluoride or "solid fluorine carriers", compounds that can release fluorine upon heating, notably cobalt trifluoride[81], may be used instead. After the reaction, the molecular size is not changed significantly, as the elements have very similar van der Waals radii.[79] Direct fluorination becomes even less important when it comes to organohalogens or unsaturated compounds reactions, or when a prefluorocarbon is desired (then HF-based electrolysis is typically used).[82] In contrast, the fluoropolymers are formed by polymerizing free radicals; other techniques used for hydrocarbon polymers do not work in that way with fluorine.[83]

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.[84]

Some important fluorocarbons
Fluoromethane (methyl
fluoride), used in
semiconductor processing 		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

[edit]

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.[85]

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.[86] As with fluorocarbons, chlorofluorocarbons and bromofluorocarbons are not flammable: they do not have carbon–hydrogen bonds to react and released halides quench flames.[86]

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.[85]

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.[87] Difluoroacetylene, which decomposes even under liquid nitrogen temperatures,[88] 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 weakly[87] (similarly to that how unsaturated hydrocarbons attacked by HF add hydrogen to the more hydrogen-rich atom per Markovnikov's rule[89]).

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[90]). They share most of perfluorocarbon properties (inertness, stability, non-wettingness and insolubility in water and oils, slipperiness, etc.),[91] 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.

Perfluoroalkanic acids organic acids may also be seen as perfluorinated compounds, with perfluoroalkyl joined to a carboxyl group. As the fluorines are added to the acid, its strength grows: consider acetic acid and its mono-, di-, and trifluoroacetic derivatives and their pKa values (4.74, 2.66, 1.24, and 0.23).[92] This happens because with fluorines, the anion formed after giving the proton off becomes stable, an effect caused by fluorine's great inductive effect. Because of this, the trifluoro derivative is 33,800 times stronger an acid than acetic.[93] Similarly, the acidity is greatly increased for other perfluorocarboxyl acids, as well as the amines (which are not acids but become less basic if fluorinated).[79]

The perfluoroalkanesulfonic acids are also very notable for their acidity. The sulfonic acid derivative, trifluoromethanesulfonic acid, is comparable in strength to perchloric acid.[94] 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 (see above).[94]

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. Industrially, such compounds are treated as perfluorinated. 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.[91]

Polymers

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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.[83]

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.[95] The compound, however, lacks an ability to transform upon melting, which is not a problem for various PTFE derivatives, namely FEP (fluorinated ethylene propylene, structurally similar to PTFE but has some fluorines replaced with the –CF3 groups) 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).[96]

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.[97] PCTFE (polychlorotrifluoroethylene, structural unit: –CF2CFCl–) is another important compound. It differs from PTFE by having a quarter of fluorines replaced with chlorines, yet this difference brings even greater hardness, creep resistance, and moisture persistence.[97]

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.[98]

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.[99]

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 fluoroelastomers. VDF/HFP is a copolymer system of vinylidene fluoride and (at least 20%) hexafluoropropylene. TFE/propylene is another copolymer system with better chemical resistance to some solvents. TFE/PMVE (perfluoromethylvinyl ether) is a copolymer system which creates a perfluorinated fluoroelastomer.[100]

See also

[edit]
  • Caesium—the least electronegative element, fluorine's "opposite"

Notes

[edit]
118 squares with numbers and letters in them, mostly colored gray and green, with a few numbers and words outside the boxes
Periodic table colored to show how elements are treated in this article: dark gray elements are metals, green ones are nonmetals, blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray ones are those with indeterminate properties.
  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 temperatues, decomposing at around 40 K.[2] Also, the F+
    4     cation and a few related species have been predicted to be stable.[3]
  2. ^ In this article, metalloids are not treated separately from metals and nonmetals, but among elements they are closer to. For example, germanium is treated as a metal, and silicon as a nonmetal. Antimony is included for comparison among nonmetals, even though it is closer to metals chemically than to nonmetals. The noble gases are treated separately from nonmetals; hydrogen is discussed in the Hydrogen fluoride section and carbon in the Organic compounds section. P-block period 7 elements have not been studied and thus are not included. This is illustrated by the image to the right: the dark gray elements are metals, the green ones are nonmetals, the light blue ones are the noble gases, the purple one is hydrogen, the yellow one is carbon, and the light gray elements have unknown properties.
  3. ^ For more detailed explanation, see [11]
  4. ^ In the context of the noble gas compounds, "stable" means not decaying chemically over time under certain (any count) conditions; other ("meta-stable") compounds are not treated in this overview article.
  5. ^ Even though helium and neon have no known stable compounds, calculations show that helium may be called slightly less unreactive than neon: As mentioned, its fluorohydride may, unlike neon's, be stable under specific conditions. A stable cation of helium (HeH+), the strongest acid of all species,[66] yet existent) has been known since 1925, while no neon cations are. Additionally, clathrates are known for every noble gas but neon.[67]
  6. ^ There is no general line where oxidation states are "relatively low" or "relatively high", they rely on specific elements (and defined only for elements that have highest oxides and fluorides are in different oxidation states); in general, +7 and +8 are high, while +4 and below are low. States +5 and +6 rely on element properties, like atomic radius; for a small nitrogen atom, +5 is "high" here, but for larger palladium and platinum +6 is still "low".
  7. ^ Note that aside from the molecular one, other forms of the highest fluorides are not considered to be possible. Such compounds are extremely unlikely to be completely ionic because of large sixth, seventh, and eighth oxidation energies of all elements, which make the reactions that may produce such compounds highly unfavorable (even though the possible osmium octafluoride may be mostly ionic[74]).

References

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  1. ^ Calderazzo, Fausto (2010). "Halide-bridged polymers of divalent metals with donor ligands – Structures and properties". Coordination Chemistry Reviews. 254 (5–6): 537–554. doi:10.1016/j.ccr.2009.08.007.
  2. ^ Wiberg, Wiberg & Holleman 2001, p. 422.
  3. ^ Schlöder, T.; Riedel, S. (2012). "Investigation of Heterodimeric and Homodimeric Radical Cations of the Series: [F2O2]+, [F2Cl2]+, [Cl2O2]+, [F4]+, and [Cl4]+". RSC Advances. 2 (3). Royal Society of Chemistry: 876–881. doi:10.1039/C1RA00804H.
  4. ^ a b Harbison, G. S. (2002). "The Electric Dipole Polarity of the Ground and Low-lying Metastable Excited States of NF". Journal of the American Chemical Society. 124 (3): 366–367. doi:10.1021/ja0159261. PMID 11792193.
  5. ^ Smart, Bruce E.; Tatlow, J. C. (1994). Organofluorine chemistry: Principles and commercial applications. Springer. p. 515. ISBN 0306446103.{{cite book}}: CS1 maint: multiple names: authors list (link)
  6. ^ Pauling, Linus A. (1960). The nature of the chemical bond and the structure of molecules and crystals: An introduction to modern structural chemistry. Cornell University Press. pp. 454–464. ISBN 978-0-8014-0333-0.
  7. ^ Atkins, Peter (2008). Chemical principles: The quest for insight. W. H. Freeman & Co. pp. 184–185. ISBN 978-1-4292-0965-6. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
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Bibliography

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