Antibonding

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Antibonding (or anti-bonding) is a type of chemical bonding. An antibonding orbital is a form of molecular orbital (MO) that is located outside the region of two distinct nuclei. The overlap of the constituent atomic orbitals is said to be 'out of phase' and as such the electrons present in each antibonding orbital are repulsive and act to destabilize the molecule as a whole. (See Electron Phases)

Antibonding molecular orbits (MOs) are normally higher in energy than bonding MOs. They are occupied by two electrons at a time and (in the case of hydrogen), each atom can contribute only one electron, therefore only the s (bonding) MO is occupied and the H₂ molecule is more stable than two separate H atoms.

A molecular orbital becomes antibonding as there is actually less electron density between the two nuclei than there would be if there was no bonding interaction at all. When an MO changes sign (from positive to negative) between two atoms, it is said to be antibonding with respect to those atoms. Antibonding orbitals are often labelled with an asterisk (*) on molecular orbital diagrams.

In molecules with several atoms, such as benzene, a particular MO may be bonding with respect to some adjacent pairs of atoms and antibonding with respect to other pairs. If the bonding interactions outnumber the antibonding interactions, the MO is said to be "bonding," while if the antibonding interactions outnumber the bonding interactions, the MO is said to be "antibonding". Since each carbon atom contributes only one electron to the π-system of benzene, there are six π-electrons and therefore only the three lowest-energy MOs (the bonding ones) are filled.

Another particular feature of antibonding is that the antibonding orbital is more antibonding than the bonding orbital is bonding. This leads to the conclusion that the energy of both MOs are raised by the presence of nucleus-nucleus repulsion.

Antibonding orbitals are also important for explaining chemical reactions in terms of molecular orbital theory. Roald Hoffmann and Kenichi Fukui shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative MO explanations for chemical reactions.

[edit] References

• Atkins, P.W. (2002). Atkins Physical Chemistry. 7th ed. Oxford. ISBN 0-19-879285-9
• Orchin, M. Jaffe, H.H. (1967) The Importance of Antibonding Orbitals. Houghton Mifflin. ISBN B0006BPT5O
• The 1981 Nobel Prize in Chemistry

v • d • e Chemical bonds "Strong" Covalent bonds & Antibonding Sigma bonds: 3c-2e (bent bond) · 3c-4e (Hydrogen bond, Dihydrogen bond, Agostic interaction) · 4c-2e Pi bonds: π backbonding · Conjugation · Hyperconjugation · Aromaticity · Metal aromaticity Delta bond: Quadruple bond · Quintuple bond · Sextuple bond Coordinate covalent bond · Hapticity Ionic bonds Cation-pi interaction · Salt bridge Metallic bonds - "Weak" Hydrogen bond Dihydrogen bond · Dihydrogen complex · Low-barrier hydrogen bond · Symmetric hydrogen bond · Hydrophile Other noncovalent van der Waals force · Mechanical bond · Halogen bond · Aurophilicity · Intercalation · Stacking · Entropic force · Chemical polarity other Disulfide bond · Peptide bond · Phosphodiester bond Note: the weakest strong bonds are not necessarily stronger than the strongest weak bonds