Valence Bond Theory

A bond according to Valence Bond Theory results from the in-phase overlap of two atomic valence orbitals.  The orbitals that overlap can be s, p, or sp^x hybrid orbitals.

As an example consider two hydrogen atoms H₁ and H₂ each having the same phase.  When they are separated rather far (maybe 4 angstroms or more) there is little or no interaction between them (right side of diagram).  As the hydrogen atoms become closer together there is an attractive force between the electrons and nuclei.  If the atoms are pushed too close together (left side of the diagram) then the nuclei repel one another and the energy increases.  At the minimum on the diagram, the attractive and repulsive forces are equal or balanced. This is the equilibrium bond length (r_e) for a hydrogen atom. We would say the H-H σ bond is made up of a combination of an H1s orbital on H₁ and an H1s orbital on H₂.

When we describe the structures of alkanes, akenes, alkynes, and other functional groups you will need to be able to identify the valence orbitals that overlap to make the various bonds.  You should note that sigma (σ) bonds are a result of the in-phase overlap of s-like orbitals (1s, 2s or sp^x hybrid orbitals), while pi (π) bonds are the result of the overlap of p-type orbitals.  It's easy to remember, s for sigma(σ) and p for pi(π).

Every covalent bond in a given molecule has a characteristic length and strength. In general, the length of a typical carbon-carbon single bond in an organic molecule is about 150 pm, while carbon-carbon double bonds are about 130 pm, carbon-oxygen double bonds are about 120 pm, and carbon-hydrogen bonds are in the range of 100 to 110 pm. The strength of covalent bonds in organic molecules ranges from about 234 kJ/mol for a carbon-iodine bond (in thyroid hormone, for example), about 410 kJ/mole for a typical carbon-hydrogen bond, and up to over 800 kJ/mole for a carbon-carbon triple bond.

Typical Bond Energies and Lengths