Chemistry
Chemistry 9: 1-2
WebLecture
Valence Bond Theory
Outline
Electrons Between Atoms
We've talked at length about orbitals in isolated atoms, and how precisely the wave nature of the electron and the electrical forces — the mutual repulsion between the electrons and their attraction to the positive nucleus — define where electrons can exist. Each element's orbitals are unique, resulting in a "fingerprint" emission and absorption spectral pattern. Now we step back to get the larger picture when multiple electrical fields overlap in molecules.
Orbitals in Molecules
It shouldn't be much of a surprise that when atoms move into molecules, the orbitals shift, affected by the overlapping fields of the repulsive electrons and the multiple attractions of several nearby positive nuclei, creating new patterns and new spectra.
To account for these variations from the orbitals of isolated atoms, Linus Pauling proposed the valence bond theory, in which the isolated orbitals become hybrid orbitals in molecules. We look at how electrons are shared between two atoms in each bond of the molecule (localized orbitals), and determine the possible hybrid orbitals. The theory accounts for many observations of molecular structures, but not all.
A second proposed theory, the molecular orbital theory, was developed by Robert Mulliken. This theory looks at the electrons as though they were a common resource to be shared shared among all the bonds of the atoms in the molecule (delocalized orbitals). This theory accounts for some of the molecules the VSEPR theory cannot explain...but not for some that VSEPR does explain! So for now, we need both theories.
Hybrid Orbitals
Where orbitals overlap, electrons are attracted to two nuclei — and the nuclei are attracted to the electrons, forming the bond. The greater the amount of overlap, the stronger the bond. As the nuclei get closer, the overlap increases, but there is a point at which the mutual positive charge of the nuclei causes them to repel each other.
A sigma bond can be created in three ways:
- Two s orbitals overlap
- An s and a p orbital that points toward the center of the s orbital overlap
- Two p orbitals aligned on the same axis, pointing toward one another, overlap.
Consider the case of carbon, with 2 s and 2 p electrons. The electrons first rearrange themselves to half-fill the orbitals available at the valence level, so we get one s and three p orbitals containing one electron each. The electrons then shift so they are evenly spaced at the same level between s and p. Now we have four hybrid orbitals, capable of forming exactly the same kind of bond. We call such hybridization an sp3 hybridization: 4 new and equal orbitals built from two s and two p orbitals. Because these bonds are the same length and equally spaced around the carbon atom, the resulting shape of a CX4 molecule is tetrahedral.
WARNING: Don't be misled by the notation. Hybrid orbitals are not the same thing as electron configuration, even though the notation is similar.
Several rules appear to govern the formation of hybrids and the resulting characteristics of the molecule, including its shape.
- Rule 1: The hybridization can only occur in a half-filled orbital. A filled orbital has a lone pair.
- Rule 2: The number of hybrid orbitals must equal the number of valence orbitals.
- Rule 3: The total number of hybrid orbitals and lone pairs determines the shape of the molecule.
There can only be one s and three p orbitals in a valence shell, so atoms of the AB5 or higher form must utilize d orbitals. This gives us the following possible hybrids:
| Hybrid set | Shape |
|---|---|
| sp | linear |
| sp2 | planar-trigonal |
| sp3 | tetrahedral |
| sp3d | trigonal-bipyramidal |
| sp3d2 | octahedral |
Multiple Bonds
Multiple bonds form only when parallel p orbitals overlap. For example, assume that you have two atoms sharing a sigma bond along the X-axis, forming a molecule that looks like a dumbbell. If each atom also has a half-filled, un-hybridized p orbital along the Y-axis, these orbitals can overlap as well, forming a pi bond. An additional pi bond could form along the Z-axis, creating a triple bond, if an un-hybridized electron exists on each atom after formation of the first pi bond.
Rotation can occur around a single sigma bond, but not around a double bond (sigma plus pi). This causes a higher level of stability in double bonds, and allows for the formation of isomers, molecules that have the same molecular formula and the same bonds, but different spatial organization.
Another form of multiple bond is the ring, which we see in benzene (CH6) and in saccharides when they are in water solutions. Benzene in particular is interesting become some of the electrons are delocalized across the entire ring, forming a resonance bond situation.
In most situations involving sigma bonds, the VSEPR theory is adequate. It is only when we need to account for pi bonds that we need a theory that allows for electrons to exist in delocalized orbits — the molecular orbital theory.
Practice with the Concepts
Discussion Questions
- Why do we need two theories to account for the same type of phenomena in different situations. Which is "right"?
- How could we apply VSEPR theory to organic molecules like starches, which do not contain central atoms and terminal atoms, but rather form long chains?
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