Valence Bond Theory

In Valence Bond Theory, a bond forms when two atomic valence orbitals overlap in-phase, creating a region of increased electron density between the bonded nuclei. The overlapping orbitals can be s, p, or sp^x hybrid orbitals.

Consider the example of two hydrogen atoms, H1 and H2, each with a 1s orbital in the same phase. When the atoms are far apart (more than 4 angstroms), their orbitals barely interact, and there is little or no attractive force between them, resulting in a high energy state (right side of the diagram). As the atoms move closer together, the electrons from each atom begin to feel the attractive pull of the opposing nucleus, which stabilizes the system and lowers the energy. However, if the atoms come too close, the repulsion between the positively charged nuclei causes the energy to increase sharply (left side of the diagram).

The equilibrium bond length (r_e) is the point where the attractive and repulsive forces balance perfectly, resulting in the lowest energy configuration. At this distance, the H-H sigma (σ) bond forms, created by the overlap of the 1s orbital from H1 and the 1s orbital from 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