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Fluorine forms a great variety of chemical compounds, within which it almost 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.〔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.〕 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). 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 (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. 抄文引用元・出典: フリー百科事典『 ウィキペディア（Wikipedia）』 ■ウィキペディアで「Compounds of fluorine」の詳細全文を読む スポンサード リンク
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