Book Cover Organic Chemistry 4e Carey
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Chapter 1: Chemical Bonding



Summary | Atoms, Electrons and Orbitals | Bond types: Ionic, Covalent, Polar Covalent | Lewis Structures and Formal Charge | The Shapes of Molecules | Hybridization | Self Assessment | Quiz |


Hybridization

Chapter 1: Chemical Bonding

 

Covalent bonds are formed when atomic orbitals overlap.  There are two types of orbital overlap that an organic chemist needs to be familiar with.  Sigma, s, overlap occurs when there is one bonding interaction that results from the overlap of two orbitals.  Pi, p, overlap occurs when two bonding interactions result from the overlap of orbitals.

The organic chemist also needs to realise how these orbital overlaps relate to the type of bonding that is occuring between atoms:
single bond      s overlap
double bond     s and p overlaps
triple bond       s and two p overlaps
If one tries to correlate the overlap of atomic orbitals to the shape of a molecule, however, the expected geometry does not correspond to a maximum orbital overlap.   Take a look at methane, CH4 . VSEPR predicts a tetrahedral geometry about the carbon atom but this is not achieved when one considers a maximum orbital overlap between four H atoms - each with a 1s orbital and the 2s, 2px, 2py and 2pz orbitals of carbon.
Hybridization is a solution to this problem.  It is the imaginary mixing of the 2s, 2px, 2py and 2pz atomic orbitals of carbon to form a new set of 'hybrid' orbitals that orient themselves in the desired VSEPR geometry.  The hybrid orbitals are equivalent to one another making all orbital overlaps equivalent, therefore, all C-H bonding interactions equivalent.
Hybrid orbitals are named by considering the type and number of atomic orbitals from which they arose.  For CH4 then the hybridization for the carbon is sp3.   One sees that the hybridization of an atom can be determined very quickly by considering the number of electron groups about an atom.  Hybrid orbitals are responsible for all the s bonding overlaps in a molecule.  Unhybridized orbitals are responsible for all the p bonding overlaps in a molecule.
 
 

You should be able to predict the hybridization of all non-hydrogen and non-terminal atoms in a molecule and draw the bonding interactions in that molecule based on the hybridisation model.  Try some questions.

If you are still having difficulty with hybridization, then you will need to review your concept of hybridization.

 

Going deeper into Hybridization:

Consider a molecule of carbon monoxide, CO.  Start by visualising the imaginary mixing of the valence atomic orbitals of carbon.  To do this refer to the valence orbital energy level diagram of carbon, the number of electrons is determined by taking into account any formal charges.  In CO, carbon has a formal charge of one so that the electron configuration is 2s22p3, rather than 2s22p2.

Next the number of groups about carbon for which hybridization must occur is considered.  It has two electron groups, one bonding the other a lone pair.
The number of electron groups is saying that two atomic orbitals must hybridize.  Orbitals are always hybridized from lowest to highest energy, therefore, one s and one p orbital will hybridize.  The energy level diagram will change so that the two hybridized orbitals will become equal in energy, which will be an average of the energy for the atomic orbitals from which they arose, and two unhybridized p orbitals remain.
The energy level diagram for the oxygen is determined in the same way but the electron count will take into consideration the 1+ formal charge.
Recall that CO has a triple bond, which should be the result of one s and two p overlaps.  The hybrid orbitals are involved in forming the s overlap of the triple bond and in housing the lone pairs of electrons.  The unhybridized p orbitals are involved in two p overlaps, which will be located perpendicular to one another.

 

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