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Chemists invoke the concept of hybridization to explain this discrepancy. Under this concept, to accommodate the geometry of molecules, atomic orbitals modify themselves to become hybrid orbitals of the suitable geometry. For instance, to attain a tetrahedral arrangement the carbon undergoes sp3 hybridization: the 2s orbital and three 2p orbitals become four sp3 hybrid orbitals. Note that the total number of orbitals is conserved, but the orientation and energy of the orbitals have changed.

Figure %: Hybridization of s and p orbitals into sp3 hybrid orbitals.

Orbital View of Multiple Bonding

Single covalent bonds can be explained by the VB model as the result of a head-on overlap between atomic orbitals. In some cases, these may be hybrid orbitals. Such head-on overlap is a sigma bond (σ bond), so-called because of the bond's cylindrical symmetry. Only one σ bond can exist between two given atoms. How, then, are double and triple bonds formed?

The answer lies in the fact that p-orbitals are capable of overlapping sideways in what is called a pi bond (Π bond). Π bonds are weaker than σ bonds because sideways overlap is not as effective as head-on overlap. For example, a C-C σ bond has a typical bond energy of 80 kcal/mol, but the C-C Π bond energy is usually around 60 kcal/mol.

Consider ethylene, which has a C=C double bond. Each carbon has three bonds in the VSEPR scheme, so each carbon has a trigonal planar geometry. To accommodate this geometry each carbon undergoes sp2 hybridization. The 2s orbital and two of the 2p orbitals hybridize to form three sp2 hybrid orbitals. The last p-orbital of each carbon atom remains unhybridized. These unhybridized p-orbitals overlap with one another to form the necessary Π-bond.

Figure %: Hybridization into sp2 orbitals and Π bond formation in ethylene.

Triple bonds are formed in a similar process. In acetylene, the C-C triple bond is actually one σ bond and two Π bonds. Each carbon undergoes sp hybridization. The two unhybridized p-orbitals on each carbon form two orthogonal Π-bonds.

Figure %: Hybridization into sp orbitals and bonding in acetylene.

The Valence Bond Model: Conclusions

The valence bond model provides a simple and useful framework through which we mayunderstand covalent bonding. However it has several drawbacks. First, when using this model it is difficult to say anything about the energies of electrons. A more serious drawback of the VB model is its assumption that electrons are localized to specific atoms. In fact electrons are commonly delocalized to several atoms, as described by resonance structures. The Molecular Orbital model, while more complex, addresses both of these issues.

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