Examples of antibonding in the following topics:
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- In molecular orbital theory, bond order is also defined as the difference, divided by two, between the number of bonding and antibonding electrons; this often, but not always, yields the same result.
- In the second diagram, one of the bonding electrons in H2 is "promoted" by adding energy and placing it in the antibonding level.
- However, removing an electron from the antibonding level produces the molecule He2+, which is stable in the gas phase with a bond order of 0.5.
- If the molecule He2 were to exist, the 4s electrons would have to fully occupy both the bonding and antibonding levels, giving a bond order of zero.
- By adding energy to an electon and pushing it to the antibonding orbital, this H2 molecule's bond order is zero, effectively showing a broken bond.
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- Bonding and antibonding orbitals are illustrated in MO diagrams, and are useful for predicting the strength and existence of chemical bonds.
- For a π-bond, corresponding bonding and antibonding orbitals would not have such symmetry around the bond axis, and are designated π and π* respectively.
- The electrons in the bonding MOs are called bonding electrons, and any electrons in the antibonding orbital are called antibonding electrons.
- The presence of a filled antibonding orbital, after fulfilling the conditions above, indicates that the bond in this case does not exist.
- Notice the two electrons occupying the antibonding orbital, which explains why the He2 molecule does not exist.
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- As in SN2, the leaving group (LG) is "pushed" away by electrons that access the C-LG antibonding orbital.
- That way, the electrons from the C-H bond can easily fall into the antibonding C-LG orbital, which is found 180 degrees from the C-LG bonding orbital.
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- In the following diagram, two 1s atomic orbitals combine to give a sigma (σ) bonding (low energy) molecular orbital and a second higher energy MO referred to as an antibonding orbital.
- The 1s and 2s atomic orbitals do not provide any overall bonding, since orbital overlap is minimal, and the resulting sigma bonding and antibonding components would cancel.
- In both these cases three 2p atomic orbitals combine to form a sigma and two pi-molecular orbitals, each as a bonding and antibonding pair.
- The overall bonding order depends on the number of antibonding orbitals that are occupied.
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- Initial bonding with a nucleophile is believed to involve the empty antibonding π*-orbital.
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- When two double bonds are conjugated, the four p-atomic orbitals combine to generate four pi-molecular orbitals (two are bonding and two are antibonding).
- In a similar manner, the three double bonds of a conjugated triene create six pi-molecular orbitals, half bonding and half antibonding.
- The energetically most favorable π __> π* excitation occurs from the highest energy bonding pi-orbital (HOMO) to the lowest energy antibonding pi-orbital (LUMO).
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- These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond, in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.
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- The remaining six molecular orbitals are antibonding, and are empty.
- The symmetries of the appropriate reactant and product orbitals were matched to determine whether the transformation could proceed without a symmetry imposed conversion of bonding reactant orbitals to antibonding product orbitals.
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- Typically they either have low-charge (Na+), electrons in d orbitals that are antibonding with respect to the ligands (Zn2+), or lack covalency (Ln3+, where Ln is any lanthanide).
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- Orbital interactions that produce bonding or antibonding orbitals in heteronuclear diatomics occur if there is sufficient overlap between atomic orbitals, as determined by their symmetries and similarity in orbital energies.