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Orbitals for selected molecules

Posted by RAHMA CAHYANINGRUM pada Maret 6, 2010

1. Saturated molecules

These are molecules in which all valence electrons are involved in the formation of single bonds. There are no non-bonded lone pairs. These molecules are generally less reactive than either electron-rich or electron-deficient species, with all occupied orbitals having relatively low energies.

a. Methane:

The valence molecular orbitals of methane are delocalized over the entire nuclear skeleton – that is, it is not easy to assign any one orbital to a particular C-H bond. It is possible to see how complex the orbital structure becomes with the increase in energy. Methane has four valence molecular orbitals (bonding), consisting of one orbital with one nodal plane (lowest occupied) and three degenerate (equal energy) orbitals that do have a nodal plane.
For the energy diagram and pictorial view of the orbitals – please see below:

b. Ethane:

The ethane molecule has fourteen valence electrons occupying seven bonding molecular orbitals. As can be seen from the energy diagram – four of the molecular orbitals occur as degenerate pairs. Like in methane – the molecular orbitals of ethane show increasing nodal structure with increasing orbital energy.
For the energy diagram and pictorial view of the orbitals – please see below:

2. Molecules with double bonds

In molecules where the number of bonding electron pairs exceeds the number of unions between atoms, the extra electrons occupy higher energy molecular orbitals than the orbitals found in molecules where the number of bonding electron pairs equals the number of unions between atoms. These are double bonds, and the orbitals have a nodal plane containig the atoms sharing these p-type orbitals.

a. Ethene:

The simplest alkene is ethene. Its chemistry is dominated by two “frontier orbitals“, that is the Highest Occupied Molecular Orbital (HOMO) and the Lowest Unoccupied Molecular Orbital (LUMO). For the ethene orbital energy diagram these are shown as pCC for the HOMO, and p*CC for the LUMO.
An important property of the ethene molecule, and alkenes in general is the existence of a high barrier to rotation about the C=C which tends to hold the molecule flat.
For the energy diagram and pictorial view of the orbitals – please see below:

3. Molecules with triple bonds

a. Ethyne:

For the energy diagram and pictorial view of the orbitals – please see below:

4. Molecules with electron lone pairs

a. HF

A simple diatomic molecule is Hydrogen fluoride. There are eight valence electrons which occupy four molecular orbitals. The two highest energy MO’s are degenerate, are p-type and have no electron density associated with the hydrogen atom, ie. they are Non-Bonding Orbitals (NBO) and in Lewis Theory are represented as two “Lone Pairs“. Another important difference between Hydrogen Fluoride and previous molecules is that the electron density is not equally distributed about the molecule. There is a much greater electron density around the fluorine atom. This is because fluorine is an exremely electronegative element, and in each bonding molecular orbital, fluorine will take a greater share of the electron density.

For the energy diagram and pictorial view of the orbitals – please see below:

b. Water

In the water molecule the highest occupied orbital, (1b1) is non-bonding and highly localized on the oxygen atom, similar to the non-bonding orbitals of hydrogen fluoride. The next lowest orbital (2a1) can be thought of as a non-bonding orbital, as it has a lobe pointing away from the two hydrogens. From the lower energy bonding orbitals, it is possible to see that oxygen also takes more than its “fair share” of the total electron density.

c. Ammonia

Ammonia has two pairs of degenerate orbitals, one bonding and one antibonding, and like hydrogen fluoride and water has a non-bonding orbital (2a1). This highest occupied orbital has a lobe pointing away from the three hydrogens, and corresponds to a lone pair orbital localized upon the nitrogen, whereas the three lowest energy MO’s lead to the description of the three N-H bonds of the Lewis structure. The lone pair is relatively high in energy, and is responsible for the well known Lewis base properties of ammonia.



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