why is one lobe of hybrid orbital small and the other lobe large?

why is one lobe of hybrid orbital small and the other lobe large?


1 Answers

879 Points
11 years ago

Dear student,

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 1s orbitals of H and the 2s, 2px, 2py and 2pz orbitals of carbon.

Hybridisation 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 hybridisation for the carbon is sp3.   One sees that the hybridisation 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.  Unhybridised orbitals are responsible for all the p bonding overlaps in a molecule.


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