Download AP HYBRIDIZATION

HYBRIDIZATION AND THE LOCALIZED ELECTRON MODEL
The arrangement of valence electrons is represented by the Lewis structure (or structures if
resonance is in play). Once a correct Lewis structure is drawn, the molecular geometry
(shape) can be predicted from the VSEPR model as well as the bond angles.
Consider the typical bonding pattern of carbon (a 1s22s22p2 atom). If you went only with
the idea of atomic orbital overlap (valence bond theory), it would seem that when carbon
bonds there would be only two bonds forming – one from each half filled 2p orbital. But
carbon never bonds twice! If one invokes electron promotion to move an electron from
the 2s to the 2p, you would achieve 4 half filled orbitals (a 2s and three 2p) and now carbon
could bond via orbital overlap 4x. However, when you consider the shape and orientation
of the s and p orbitals you still have a problem. Consider CH 4. Since s and p orbitals have
slightly different energies, the three 2p-orbitals overlapping with the H 1s orbital and one
2s orbital overlapping with the H 1s would seem to imply a significant energy difference in
one of the C-H bonds. However, methane is known to have four identical energy C-H
bonds. Secondly, the 2p orbitals are mutually perpendicular. We might expect three C-H
bonds to be oriented at 90° (the p orbital shape) and one stuck somewhere in an open gap
(the s orbital). Experiments, however, confirm a tetrahedral molecule with bond angles of
109.5°.
Perhaps the 2s and 2p orbitals on an isolated carbon atom may not be the best set of
orbitals for bonding; a new set of atomic orbitals might better serve the carbon atom. To
account for the known structure of methane, it makes sense to assume that the carbon
atom has four equivalent energy atomic orbitals, arranged tetrahedrally – since that’s what
experiments reveal. Such a set of orbitals can be obtained by mathematically combining
the carbon 2s and 2p orbitals. This mixing of the "native" atomic orbitals to form special
orbitals for bonding is called hybridization. These four new orbitals are called sp3 orbitals
because they are formed from one 2s and three 2p orbitals (how clever). We say that the
carbon atom undergoes sp3 hybridization or is sp3 hybridized. In fact, whenever the
electron geometry is tetrahedral, the model assumes sp3 orbitals: the central atom
becomes sp3 hybridized.
Consider ethene C2H4. Draw its Lewis structure:
Earlier we prescribed that a double bond acts as one direction – it doesn’t affect shape.
The ethane molecule thus requires a trigonal planar arrangement with bond angles of 120°.
What orbitals do the carbon atoms in this molecules employ? The 2s and 2p orbitals do not
have the required arrangement so we need a set of hybrid orbitals. The sp 3 orbitals gave us
109.5° angles rather than the required 120° angles. If one combines one 2s and two 2p
orbitals we get sp2 hybridization. In forming the sp2 orbitals, one 2p orbital on carbon has
not been used. This remaining p orbital is oriented perpendicular to the plane of the sp2
orbitals.
In each of the sp2 orbital bonds, the electron pair is shared in an area centered on a line
running between the atoms. This type of bond is called a sigma bond ().
But what of the double bond? The second bond must therefore result from sharing an
electron pair in the space above and below the sigma bond. This type of bond can be
formed using the 2p orbital perpendicular to the sp 2 hybrid orbitals on each carbon atom.
This type of bond is called a pi bond (π). Sigma bonds are formed from orbitals whose
lobes point towards each other and pi bonds result from lateral or side-to-side overlap.
Whenever three directions of electrons surround an atom, a set of sp 2 hybrid orbitals is
required.
If triple bonding is in play as in ethyne, C2H2, the carbon has sp hybridization. The
hybridized electron pairs are found between the nuclei forming a sigma bond. The
unhybridized electrons which occupy two degenerate p orbitals on each C, laterally overlap
creating high electron density above and below the plane forming one pi bond, and high
electron density in front of and in back of the plane forming the second pi bond.