Fluorine compounds


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. 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 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 the highest oxidation states of oxides and fluorides are always equal.

Difluorine

While an individual fluorine atom has one unpaired electron, molecular fluorine has all the electrons paired. This makes it diamagnetic with the magnetic susceptibility of −1.2×10−4, which is close to theoretical predictions. In contrast, the diatomic molecules of the neighboring element oxygen, with two unpaired electrons per molecule, are paramagnetic.
XXXHXBX3AlX3CX4
F159574645582456
Cl243428444427327
Br193363368360272
I151294272285239

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.
The F2 molecule is commonly described as having exactly one bond 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.

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.
Reactions of elemental fluorine with metals require diverse conditions that depend on the metal. Often, the metal must be powdered because many metals passivate by forming protective layers of the metal fluoride that resist further fluoridation. 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.
Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals. The halogens react readily with fluorine gas as does the heavy noble gas radon. 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. Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.
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 and with intermediate oxidation states exist at very low temperatures, decomposing at around 40 K.. Also, the 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.
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 pKa values ; 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. There is evidence for an even stronger acid called fluoroauric acid but it has not proved isolable.

Hydrogen fluoride

Fluorine combines with hydrogen to make a compound 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. 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. HF is miscible with water, 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, which is 44 degrees Celsius above the melting point of pure HF.
Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in water solution, with acid dissociation constant equal to 3.19. 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 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.
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.

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.
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 liquids, or gases 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.

Low oxidation state metal fluorides

The alkali metals form monofluorides. All are soluble and have the sodium chloride structure, Because the fluoride anion is basic, many alkali metal fluorides form bifluorides with the formula MHF2. Among other monofluorides, only silver and thallium 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 and nickel, 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, 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. BeF2 has many similarities to SiO2 a mostly covalently bonded network solid. BeF2 has tetrahedrally coordinated metal and forms glasses. 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. However, BeF2 has much lower electrical conductivity when in solution or when molten than would be expected if it were ionic.
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, and lanthanum adopt the LaF3 structure. Iron and gallium trifluorides have the FeF3 structure, which is similar to rhenium trioxide. Only ScF3 is cubic 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.
Gold trifluoride adopts a structure of linked –AuF4– squares that align in a helix. In contrast to gold's distinctly ionic trifluoride, its trichloride and tribromide are volatile dimeric molecules. 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. 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.
Titanium and tin tetrafluorides are polymeric, with melting points below 400 °C. 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.