Valence Bond (VB) Theory Definition

Valence bond (VB) theory is a chemical bonding theory that explains the chemical bonding between two atoms. Like molecular orbital (MO) theory, it explains bonding using principles of quantum mechanics. According to valence bond theory, bonding is caused by the overlap of half-filled atomic orbitals. The two atoms share each other's unpaired electron to form a filled orbital to form a hybrid orbital and bond together. Sigma and pi bonds are part of valence bond theory.

Theory

Valence bond theory predicts covalent bond formation between atoms when they have half-filled valence atomic orbitals, each containing a single unpaired electron. These atomic orbitals overlap, so electrons have the highest probability of being within the bond region. Both atoms then share the single unpaired electrons to form weakly coupled orbitals.

The two atomic orbitals do not need to be the same as each other. For example, sigma and pi bonds may overlap. Sigma bonds form when the two shared electrons have orbitals that overlap head-to-head. In contrast, pi bonds form when the orbitals overlap but are parallel to each other.

Sigma bonds form between electrons of two s-orbitals because the orbital shape is spherical. Single bonds contain one sigma bond. Double bonds contain a sigma bond and a pi bond. Triple bonds contain a sigma bond and two pi bonds. When chemical bonds form between atoms, the atomic orbitals may be hybrids of sigma and pi bonds.

The theory helps explain bond formation in cases where a Lewis structure can't describe real behavior. In this case, several valence bond structures may be used to describe a single Lewis stricture.

History

Valence bond theory draws from Lewis structures. G.N. Lewis proposed these structures in 1916, based on the idea that two shared bonding electrons formed chemical bonds. Quantum mechanics was applied to describe bonding properties in the Heitler-London theory of 1927. This theory described chemical bond formation between hydrogen atoms in the H2 molecule using Schrödinger's wave equation to merge the wavefunctions of the two hydrogen atoms. In 1928, Linus Pauling combined Lewis's pair bonding idea with the Heitler-London theory to propose valence bond theory. Valence bond theory was developed to describe resonance and orbital hybridization. In 1931, Pauling published a paper on valence bond theory entitled, "On the Nature of the Chemical Bond." The first computer programs used to describe chemical bonding used molecular orbital theory, but since the 1980s, principles of valence bond theory have become programmable. Today, the modern versions of these theories are competitive with each other in terms of accurately describing real behavior.

Uses

Valence bond theory can often explain how covalent bonds form. The diatomic fluorine molecule, F₂, is an example. Fluorine atoms form single covalent bonds with each other. The F-F bond results from overlapping p_z orbitals, which each contain a single unpaired electron. A similar situation occurs in hydrogen, H₂, but the bond lengths and strength are different between H₂ and F₂ molecules. A covalent bond forms between hydrogen and fluorine in hydrofluoric acid, HF. This bond forms from the overlap of the hydrogen 1s orbital and the fluorine 2p_z orbital, which each have an unpaired electron. In HF, both the hydrogen and fluorine atoms share these electrons in a covalent bond.

Sources

• Cooper, David L.; Gerratt, Joseph; Raimondi, Mario (1986). "The electronic structure of the benzene molecule." Nature. 323 (6090): 699. doi:10.1038/323699a0
• Messmer, Richard P.; Schultz, Peter A. (1987). "The electronic structure of the benzene molecule." Nature. 329 (6139): 492. doi:10.1038/329492a0
• Murrell, J.N.; Kettle, S.F.A.; Tedder, J.M. (1985). The Chemical Bond (2nd ed.). John Wiley & Sons. ISBN 0-471-90759-6.
• Pauling, Linus (1987). "Electronic structure of the benzene molecule." Nature. 325 (6103): 396. doi:10.1038/325396d0
• Shaik, Sason S.; Phillipe C. Hiberty (2008). A Chemist's Guide to Valence Bond Theory. New Jersey: Wiley-Interscience. ISBN 978-0-470-03735-5.