Localized Bonding Model for Carbon Dioxide
The isosurface viewer below illustrates the Valence Bond Theory description of bonding in the carbon dioxide molecule. The small red spheres on the right and left indicate the positions of the oxygen nuclei. The gray sphere marks the position of the carbon nucleus.
Examine each of the atomic orbitals. Identify orbitals for lone pairs 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 all three nuclei lie along the z axis. The vertical axis is the x axis.
- On which atom are the lone pairs located? Which orbitals account for the lone pairs of electrons?
- Which atomic orbitals overlap to form a C=O σ bond?
- Which atomic orbitals overlap to form a C=O π bond?
- How is electron density distributed differently in a C=O σ bond compared with a C=O π bond?
- Compare the two C=O π bonds. How are they the same? How are they different?
- In the viewer, the C=O π bond formed from the 2px atomic orbitals is located between the carbon and the oxygen on the left. The C=O π bond to the oxygen on the right uses the 2py atomic orbitals. Could this arrangement be reversed, with the πx orbital to the right and the πy to the left? How does Valence Bond Theory deal with this issue?
sp (1) sp (2) 2px 2py
Oxygen (1) Atomic Orbitals
sp2 (z) sp2 (front) sp2 (back) 2px
Oxygen (2) Atomic Orbitals
sp2 (z) sp2 (up) sp2 (down) 2py
Carbon Dioxide Orbitals
C-O σ orbitals: 1 2
C=O π orbitals: πx πy
Oxygen (2) Lone Pairs:
VB_CO2.html version 3.0
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