Charge-shift bond
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In theoretical chemistry, the charge-shift bond is a proposed new class of chemical bonds that sits alongside the three familiar families of covalent, ionic, and metallic bonds where electrons are shared or transferred respectively.[1][2] 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 electric charge between the bonded atoms is not necessarily a feature of covalent bonds.[3]
An example where charge shift bonding has been used to explain the low electron density found experimentally is in the central bond between the inverted tetrahedral carbon atoms in [1.1.1]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.[2] The charge shift bond (CSB) has also been shown to exist at the cation-anion interface of protic ionic liquids (PILs).[4] The authors have also shown how CSB character in PILs correlates with their physicochemical properties.
The valence bond view of chemical bonding that owes much to the work of Linus 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[5] 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.[5]
| Actual H−H | Actual Cl−Cl | H−Clcov Covalent bond energy H−Cl, arithmetic mean (H−H) and (Cl−Cl) | H−Clact Actual H−Cl | "Ionic contribution" H−Clact – H−Clcov | |
|---|---|---|---|---|---|
| Bond dissociation energy(kcal mol−1) | 103.5 | 57.8 | 80.6 | 102.7 | 22.1 |
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.[2]
Calculated bond energies
The contribution of ionic resonance structures has been termed the charge−shift resonance energy, REcs, and values have been calculated for a number of single bonds, some of which are shown below:[2]
| Covalent contribution kcal mol−1 | REcs kcal mol−1 | % REcs contribution | |
|---|---|---|---|
| H−H | 95.8 | 9.2 | 8.8 |
| Li−Li | 18.2 | 2.8 | 13.1 |
| H3C−CH3 | 63.9 | 27.2 | 30.2 |
| H2N−NH2 | 22.8 | 43.8 | 65.7 |
| HO−OH | –7.1 | 56.9 | 114.3 |
| F−F | –28.4 | 62.2 | 183.9 |
| Cl−Cl | –9.4 | 48.7 | 124.1 |
| H−F | 33.2 | 90.8 | 73.2 |
| H−Cl | 57.1 | 34.9 | 37.9 |
| H3C−Cl | 34.0 | 45.9 | 57.4 |
| H3Si−Cl | 37.0 | 65.1 | 63.8 |
The results show that for homonuclear bonds the charge shift resonance energy can be significant, and for F2 and Cl2 show it is the attractive component whereas the covalent contribution is repulsive. The reduced density along the bond axis density is apparent using ELF, electron localization function, a tool for determining electron density.[2][6]
The bridge bond in a propellane
The bridge bond (inverted bond between the bridgehead atoms which is common to the three cycles) in a substituted [1.1.1]propellane has been examined experimentally.[7] A theoretical study on [1.1.1]propellane has shown that it has a significant REcs stabilisation energy.[8]
