Molecular Orbital theory is primarily used to explain the bonding in molecules that cannot be defined by Valence bond Theory. These are molecules that typically involve some type of resonance. Resonance implies that a bond is neither single nor dual but some hybrid of the two. Valence bond theory only explains the bonding of solitary or twin or triple bonds. The does not provide an explanation for resonance bonding.

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Molecular orbital concept does describe resonance.

The rules of Molecular orbital Theory:

First principle: The number of molecular orbitals created is always equal to the number of atomic orbitals carried by the atoms that have actually combined. second principle: Bonding molecule orbitals are reduced in power that the parental orbitals, and also the antibonding orbitals are higher in energy. third principle: electron of the molecule room assigned to orbitals from lowest to successively higher energy fourth principle: atomic orbitals integrate to form molecular orbitals most efficiently when the atom orbitals room of comparable energy.

OK. For this reason what execute those values actually mean???

Principle 1: example - Hydrogen ( H2 ) each hydrogen atom has a solitary valence orbital, this gift the 1s orbital. Two molecular orbitalsmay be developed by the constructive and also destructive overlap the these two atomic orbitals. Therefore if you have two 1s atomic orbitals you can only make two molecular orbitals indigenous them. This is the an initial Principle.


According come MO Theory, the two molecular orbitals that kind are referred to as s (sigma = bonding) and also s* (sigma star = antibonding). In the situation of H2 both the the valence electron that kind the bond between the hydrogens to fill the bonding or s orbital.

Principle 2 & 3: This communication of atom orbitals, which offers rise come the molecular orbitals, may likewise be represented in the type of an orbital (electron) power diagram which reflects the family member energies that the orbitals. In the specific case of hydrogen every of the secluded atoms has one electron in its 1s orbital and when the atoms combine to form H2 the two electrons may be accommodated (with the contrary spins) in the bonding molecule orbital, as shown below. The second principle describes why electron would desire to fill molecular orbitals in the very first place. Together you should recognize by now, stability originates from lowering power needs. Think around it. Don"t girlfriend feel better when your energy need is lowered? If not, I would be happy to increase your homework? Anyway...because the bonding molecular orbitals carry out a reduced energy, more stable state because that the electrons, they fill these orbitals first. This additionally explains the 3rd principle statement as well.


Principle 4: If you keep in mind in H2 we linked two 1s orbitals to kind a single lower energy s molecule orbital. The 4th principle says that steady molecular orbitals are easiest to type when constructed out of atomic orbitals of comparable energies. This way that 1s orbitals should combine with 1s orbitals and 2p orbitals should integrate with 2p orbitals etc. To form the most stable molecule orbitals.

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Filling up the molecular Orbitals:

We begin with:


We then type Molecular Orbitals:


Which fill from shortest to highest power as follows:


Bond Order

The shortcut order because that a molecule deserve to be figured out as follows: bond order = ½ (bonding electron − antibonding electrons). Therefore, the H2 molecule has a link order the ½ (2 − 0) = 1. In other words, there is a single bond connecting the two H atom in the H2 molecule. In the case of He2, ~ above the other hand, the link order is ½ (2 − 2) = 0. This means that He2 is not a steady molecule.

Let"s try an example: N2

Each Nitrogen atom has 7 electrons and an electron construction of 1s2 2s2 2p3 . This returns a total of 14 electron to job-related with. How many molecular orbitals have the right to form? Which molecular orbitals will form? i beg your pardon orbitals space filled first? What is the shortcut order because that N2?


Something come notice:

1) because dinitrogen completely fills the s1s and s*1s orbitals they cancel each other out and also it becomes more obvious as to why the is the valence electrons the actually manage bonding. As you continue from duration to duration in the periodic table this tendency of cancellation among the main point electrons continues.