Hybridization

This is the idea that atoms involved in bonding rearrange their orbitals shapes and energies in order to produce orbitals that can overlap successfully with suitable orbitals on adjacent atoms and be used in bonding.

Carbon, for example, forms four bonds with hydrogen to make the molecule CH₄. Studies show that the methane molecule is symmetrical and the hydrogen atoms are arranged in a tetrahedral fashion around the central carbon atom.

However, the atomic orbitals in the outer shell of carbon have two electrons in a 2s orbital and 2 electrons in two 2p orbitals. There are three problems here:

• The 2s orbital is spherical and can't overlap with adjacent atoms.
• The 2s orbital is full. It can't share an electron from another overlapping orbital.
• There are only two singly occupied orbitals available for overlap and shaing electrons suggesting that carbon should only form two bonds and that these bonds would be at 90º to one another.

Carbon atoms overcome these problems by reforming the orbitals to give four degenerate orbitals, each of which is singly occupied. This is called Hybridization.

It must be stressed that Hybridization is simply an model to explain the difference between the shapes and orientation of the orbitals found in atoms and molecules. It gives us a process that explains how one situation can change into another.

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Hybridization in double and triple bonds

Double bonds

Hybridization can also explain the formation and shape of molecules that contain double and triple bonds.

Using carbon as the example, in the case of a double bond the central atom needs to only bond to three other atoms. To do this it needs three degenerate orbitals which it obtains by Hybridization of the three valence orbitals that have the lowest energy, i.e. the 2s and two of the 2p orbitals.

Triple bonds

Atoms that are triple bonded can only bond to two other atoms. This means that two degenerate orbitals are needed. These are provided by Hybridization of the two orbitals from the valence shell that have the lowest energy, i.e. the 2s and onle of the 2p orbitals.

This Hybridization scheme produces two 'sp' orbitals at 180º to one another (linear) that can be used for direct orbital overlap (sigma bonding) and leaves two unchanged 'p' orbitals that can be used for lateral overlap with suitable orbitals on adjacent atoms (pi bonding).

It must be stressed that Hybridization occurs in all molecules, not just carbon-containing molecules. It is a simple matter to spot sp³, sp² and sp Hybridization by the number and type of bonds on an atom.

Carbon, for example, forms four bonds. If all four bonds are single then the Hybridization is sp³. If there is one double bond then it is using sp² Hybridization and if there is a triple bond then the Hybridization is sp.

Nitrogen forms only three bonds with other atoms and in nitrogen containing compounds one double bond is indicative of sp² Hybridization, just like in carbon. In the ammonia molecule the four electron pairs are arranged in a tetrahedral orientation using sp³ Hybridization.

In the hydrogen cyanide molecule the nitrogen is attached to the carbon atom by a triple bond. Both atoms are sp hybridised.

In summary:

• If an atom has only single bonds then it is using sp³ Hybridization.
• if an atom is attached with one double bond then it is sp² hybridised
• If an atom is attached by a triple bond then it is sp hybridised.

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