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Charge-shift bond
The charge-shift bond has been proposed as a new class of chemical bond that sits alongside the two familiar families of covalent and ionic bonds where electrons are shared or transferred respectively. The charge shift bond derives its stability from the resonance of ionic forms rather than the covalent sharing of electrons which are often depicted as having electron density between the bonded atoms. A feature of the charge shift bond is that the predicted electron density between the bonded atoms is low. It has long been known from experiment that the accumulation of electronic charge between the bonded atoms is not necessarily a feature of covalent bonds. An example where charge shift bonding has been used to explain the low electron density found experimentally is in the inverted C-C bond in ()propellanes. Theoretical calculations on a range of molecules have indicated that a charge shift bond is present, a striking example being fluorine, F2, which is normally described as having a typical covalent bond.〔
==Valence bond description==
The valence bond view of chemical bonding that owes much to the work of Pauling is familiar to many, if not all, chemists. The basis of Pauling's description of the chemical bond is that an electron pair bond involves the mixing, resonance, of one covalent and two ionic structures. In bonds between two atoms of the same element, homonuclear bonds, Pauling assumed that the ionic structures make no appreciable contribution to the overall bonding. This assumption followed on from published calculations for the hydrogen molecule in 1933 by Weinbaum and by James and Coolidge.〔The Nature of the Chemical bond, L. Pauling, 1940, 2d edition, pp. 49-59, Oxford University Press〕 that showed that the contribution of ionic forms amounted to only a small percentage of the H-H bond energy. For heteronuclear bonds, A-X, Pauling estimated the covalent contribution to the bond dissociation energy as being the mean of the bond dissociation energies of homonuclear A-A and X-X bonds. The difference between the mean and the observed bond energy was assumed to be due to the ionic contribution. The calculation for HCl is shown below.〔
The ionic contribution to the overall bond dissociation energy was attributed to the difference in electronegativity between the A and X, and these differences were the starting point for Pauling's calculation of the individual electronegativities of the elements. The proponents of charge shift bond bonding re-examined the validity of Pauling's assumption that ionic forms make no appreciable contribution to the overall bond dissociation energies of homonuclear bonds. What they found using modern valence bond methods was that in some cases the contribution of ionic forms was significant, the most striking example being F2, fluorine, where their calculations indicate that the bond energy of the F-F bond is due wholly to the ionic contribution.〔
抄文引用元・出典: フリー百科事典『 ウィキペディア（Wikipedia）』
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