Valence Bond Theory
Valence Bond Theory
The valence bond (VB) theory of bonding was mainly developed by Walter Heitler and Fritz London in 1927, and later modified by Linus Pauling to take bond direction into account. The VB approach concentrates on forming bonds in localized orbitals between pairs of atoms, and hence retains the simple idea of Lewis structures and electron pairs. The wave function for the bonding electrons is described as the overlap of atomic orbitals. Thus, in the H2 molecule, the spherical s orbitals of the two H atoms simply overlap, so that the electrons in the bond sense the nuclei of both atoms. This method also works well for simple molecules like H2O, CH4, and NH3. First the appropriate hybrid orbitals are constructed on the central atom to give the correct molecular geometry (e.g., four sp3 tetrahedral orbitals on the C atom in CH4). The bonding picture is then constructed by simple overlap of the atomic orbitals, (e.g., one 1s orbital of a H atom with each sp3 lobe of the C atom in CH4 to give four C–H bonds).
Bonding in BF3
The electron configuration of the boron atom is 2s22p1, with one unpaired electron. This electron is excited to the higher energy configuration 2s12p2, with three unpaired electrons. These three orbitals are now hybridised to give three equivalent sp2 hybrid orbitals, coplanar, and lying 120° apart, each containing one electron. These hybrid orbitals then overlap the half-filled p orbitals of the three fluorine atoms, thus forming three coplanar B–F electron-pair bonds. (See Figure 1.)
All seems well, except that the boron atom does not have an octet: One perpendicular 2p orbital remains empty. This orbital can accept electron density from the F atoms to a maximum of one electron pair. On average, each F atom donates one-third of an electron pair to the empty p orbital on boron. One model for BF3 is a resonance hybrid of three structures, each having one double bond and two single bonds. The B–F bond is said to have a bond order of 1⅓.
The valence bond approach is especially useful in organic chemistry where so many molecules are built of tetrahedral C atoms, sp3 hybridised. The concept of hybrids is intuitively very satisfying because they fit visually with our perceived picture of the shape of a molecule with its directed bonds between pairs of atoms. Unfortunately, the VB approach is not satisfactory for species like CO3=, NO3−, and benzene because the VB picture does not reflect the known chemical structure. A new concept of resonance hybrids must be introduced, and CO3= must now be represented by a combination of three Lewis-octet structures. Worse still, the VB approach cannot easily give a satisfactory bonding picture for either of the important molecules O2 or CO.
In cases where the VB approach does not work well, the molecular orbital (MO) method is often more successful. The situation is best summarized by using the strengths of the VB approach where they are appropriate, as in CH4, and using the MO approach where it is best suited, as in O2 and benzene. After all, each approach is an approximation, incomplete and imperfect.
see also Bonding; Lewis Structures; Molecular Orbital Theory.
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de Kock, Roger L., and Gray, Harry B. (1980). Chemical Structure and Bonding. Menlo Park, CA: Benjamin/Cummings Publishing Co.
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Pauling, Linus (1960). The Nature of the Chemical Bond. Ithaca, NY: Cornell University Press.