antibonding orbital

Examples of antibonding orbital in the following topics:

• Bonding and Antibonding Molecular Orbitals

  • 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.

• Bond Order

  • 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.
  • The 1s electrons do not take part in the bonding, but the 2s electrons fill the bonding orbital.
  • Without the 1s electrons participating in bonding, the p electrons completely fill the bonding orbital; this leaves the antibonding orbital empty and gives a bond order of one, indicating a stable molecule (in this case, in the gas phase).
  • 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.

• Atomic and Molecular Orbitals

  • 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.
  • A mixing of the 2s-orbital with two of the 2p orbitals gives three sp2 hybrid orbitals, leaving one of the p-orbitals unused.
  • 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.

• Elimination

  • 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.

• Addition to Carbonyl Double Bonds

  • In the following diagram the essential molecular orbitals are drawn to the left of the arrow.
  • Initial bonding with a nucleophile is believed to involve the empty antibonding π*-orbital.

• Theoretical Models for Pericyclic Reactions

  • A molecular orbital diagram of ethene is created by combining the twelve atomic orbitals associated with four hydrogen atoms and two sp2 hybridized carbons to give twelve molecular orbitals.
  • Six of these molecular orbitals (five sigma & one pi-orbital) are bonding, and are occupied by the twelve available valence shell electrons.
  • The remaining six molecular orbitals are antibonding, and are empty.
  • The π-orbital on the left has one nodal plane (colored light blue), and the π*-orbital on the right has a second nodal plane (colored yellow).
  • 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.

• The Phase of Orbitals

  • When constructing molecular orbitals, the phase of the two orbitals coming together creates bonding and anti-bonding orbitals.
  • One orbital, based on in-phase mixing of the orbitals, will be lower in energy and termed bonding.
  • This molecular orbital is called the bonding orbital and its energy is lower than that of the original atomic orbitals.
  • For a π-bond, corresponding bonding and antibonding orbitals would not have such symmetry around the bond axis and would be designated π and π*, respectively.
  • P-orbital overlap is less than head-on overlap between two s orbitals in a σ-bond due to orbital orientation.

• Heteronuclear Diatomic Molecules

  • In carbon monoxide (CO), the oxygen 2s orbital is much lower in energy than the carbon 2s orbital, so the degree of mixing is low.
  • In hydrogen fluoride (HF), the hydrogen 1s orbital can mix with the fluorine 2pz orbital to form a sigma bond because experimentally, the energy of 1s of hydrogen is comparable with 2p of fluorine.
  • While MOs for homonuclear diatomic molecules contain equal contributions from each interacting atomic orbital, MOs for heteronuclear diatomics contain different atomic orbital contributions.
  • 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.
  • In hydrogen fluoride, HF, symmetry allows for overlap between the H 1s and F 2s orbitals, but the difference in energy between the two atomic orbitals prevents them from interacting to create a molecular orbital.

• The Importance of Conjugation

  • To understand why conjugation should cause bathochromic shifts in the absorption maxima of chromophores, we need to look at the relative energy levels of the pi-orbitals.
  • 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).
  • Increased conjugation brings the HOMO and LUMO orbitals closer together.

• Reactions of Coordination Compounds

  • There are also organic ligands such as alkenes whose pi (π) bonds can coordinate to empty metal orbitals.
  • 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).
  • Complexes that have unfilled or half-filled orbitals often show the capability to react with substrates.
  • These substrates need an empty orbital to be able to react with a metal center.
  • Metals with half-filled orbitals have a tendency to react with such substrates.