The structure is tetragonal and puts metal atoms in octahedral coordination. Properties depend on the number and identity of the halogen atoms. [11] This also helps explain why bonding in F2 is weaker than in Cl2. [85] Nitrogen is not known to form a pentafluoride, although the tetrafluoroammonium cation (NF+4) features nitrogen in the formal oxidation state of +5. [84] The New Latin name fluorum gave the element its current symbol F; Fl was used in early papers. [39] Despite its weakness, hydrofluoric acid is very corrosive, even attacking glass (hydrated only). ),[125] 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. Excision or amputation of affected parts may be required. The main HFC is R-134a[63] with a new type of molecule HFO-1234yf, a Hydrofluoroolefin (HFO) coming to prominence owing to its global warming potential of less than 1% that of HFC-134a. American sheep and cattle collars may use 1080 against predators like coyotes. Gases, such as sulfur dioxide and hydrogen sulfide readily combine with fluorine. At room temperature, tungsten hexafluoride is a gas. Dwell time in the body varies greatly by species, with half-lives of days in rodents, and years in humans. The noble gases are generally non-reactive because they have filled electronic shells. Hungarian-American chemist George Olah received the 1994 Nobel Prize in chemistry for investigating such reactions.[98]. Using calcium chloride – a common laboratory reagent – in lieu of calcium gluconate is contraindicated, and may lead to severe complications. [99][100][note 11], Alkali metals form ionic and highly soluble monofluorides; these have the cubic arrangement of sodium chloride and analogous chlorides. The fluoride ion from dissolved fluoride salts inhibits dental cavities, and so finds use in toothpaste and water fluoridation. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. For example, fluorine gas consists of F 2 molecules in which the fluorine atoms are bound together by sharing a pair of electrons, one contributed by each atom. [131], Fluoroelastomers, like other elastomers (artificial rubbers), consist of disordered polymer chains connected in three dimensions. [200] Other fluoropolymers, including fluorinated ethylene propylene, mimic PTFE's properties and can substitute for it; they are more moldable, but also more costly and have lower thermal stability. It is isoelectronic with O2 and, unusually, like BF, has a higher bond order than single-bonded fluorine.[2][87]. Xenon can be brought to react with fluorine, F 2 [3]: Xe (g) + 2 F 2 (g) XeF 4 (s), mix gasses at 400 °C then cool to -78 °C Fluorine's chemistry is dominated by its strong tendency to gain an electron. [17][18][note 1], The bond energy of difluorine is much lower than that of either Cl2 or Br2 and similar to the easily cleaved peroxide bond; this, along with high electronegativity, accounts for fluorine's easy dissociation, high reactivity, and strong bonds to non-fluorine atoms. The whole world has admired the great experimental skill with which you have studied that savage beast among the elements. 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). [264] Insoluble calcium fluoride formation triggers strong pain[265] and burns larger than 160 cm2 (25 in2) can cause serious systemic toxicity. Note that aside from the molecular one, other forms of the highest fluorides are not considered to be possible. Neon, on the other hand, already has a full outer energy level. Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals. Chlorofluorocarbon restrictions lowered this to 3.6 million tons in 1994; production has since been increasing. [15] The halogens react readily with fluorine gas[16] as does the heavy noble gas radon. [58][61] Fluorite, also known as fluorspar, (CaF2), abundant worldwide, is the main source of fluoride, and hence fluorine. [117][139][140], Chalcogens have diverse fluorides: unstable difluorides have been reported for oxygen (the only known compound with oxygen in an oxidation state of +2), sulfur, and selenium; tetrafluorides and hexafluorides exist for sulfur, selenium, and tellurium. [63], As with other iron alloys, around 3 kg (6.5 lb) metspar is added to each metric ton of steel; the fluoride ions lower its melting point and viscosity. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. [11] Fluorine's electrons cannot exhibit this d character since there are no such d orbitals close in energy to fluorine's valence orbitals. [163] Polytetrafluoroethylene (PTFE), the simplest fluoropolymer and perfluoro analogue of polyethylene with structural unit –CF2–, demonstrates this change as expected, but its very high melting point makes it difficult to mold. [30] Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the very strong triple bond in elemental nitrogen;[31] ammonia may react explosively. [129][note 14] However, it can attack glass, something the other acids cannot do. Perfluoroalkyl acids (PFAAs), which are sparingly water-soluble owing to their acidic functional groups, are noted persistent organic pollutants;[289] perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are most often researched.