Valence Bond Theory: Hydrogen Cyanide

In Valence Bond Theory (sometimes called the Localized Bonding Model), a σ bond betweeen two atoms is viewed as arising from direct overlap of an orbital from each atom. The orbitals used to describe the bond should be oriented directly at the other atom.

Consider the hydrogen cyanide molecule: HCN. The central atom is carbon. The Lewis structure for HCN involves a H-C σ bond, a C-N σ bond, a lone-pairs of electrons on the nitrogen, and two CN π bonds. In terms of the VSEPR Model, the carbon and the nitrogen each have two electron groups, which means the valence electrons for both the carbon and the nitrogen are sp hybrid orbitals.

The isosurface viewer below illustrates the Valence Bond Theory description of bonding in the hydrogen cyanide molecule. The small blue sphere on the left indicates the position of the nitrogen nucleus. The gray sphere marks the position of the carbon nucleus. The white sphere marks the position of the hydrogen nucleus.

Examine each of the atomic orbitals. Identify the orbital for the lone pair of electrons. Identify pairs of orbitals that overlap to form a chemical bond. Compare pairs of overlapping atomic orbitals with the corresponding bonding orbitals for the molecule.

Note that orbital positions are based upon the original view, in which the three nuclei lie along the z axis. The vertical axis is the x axis.

Questions

1. Why are the nitrogen atomic orbitals smaller than the carbon atomic orbitals?
2. There is one lone pair in this molecule. Which orbital accounts for the lone pair?
3. Both the carbon and nitrogen have 2p_x and 2p_y orbitals that are not part of a hybridization scheme. What is the role of the 2p_x and 2p_y orbitals in bonding in the HCN molecule?
4. The HCN molecule contains two σ bonds and two π bonds. Locate each of these bonds. What is the difference between a σ bond and a π bond?
5. Is bonding stronger between the C and N or between the C and H? Explain.