Valence Bond Theory (VBT) is a foundational model in chemistry developed in the early 20th century to explain how atoms connect to form molecules. VBT describes a covalent bond as the result of two atomic orbitals on different atoms overlapping in space. The theory maintains a view of bonding as a localized phenomenon, meaning the electron pair forming the bond is shared primarily between the two specific atoms involved.
The Fundamental Concept of Orbital Overlap
The core of Valence Bond Theory is the principle that a covalent bond forms when two half-filled atomic orbitals, one from each atom, physically overlap. This overlap allows the two electrons, which must have opposite spins, to be shared between the two atomic nuclei in the region of overlap. The greater the extent of this overlap, the stronger and more stable the resulting bond will be.
This mechanism gives rise to two distinct types of covalent bonds, differentiated by how the orbitals align in space. A sigma (\(\sigma\)) bond forms from the head-to-head or end-to-end overlap of atomic orbitals, such as two \(s\) orbitals, an \(s\) and a \(p\) orbital, or two \(p\) orbitals along the internuclear axis. This direct overlap results in the electron density being concentrated symmetrically along the line connecting the two nuclei, which permits free rotation around the bond axis.
The second type is the pi (\(\pi\)) bond, which forms from the side-to-side overlap of unhybridized \(p\) orbitals. This lateral overlap results in the electron density being concentrated in two lobes, one above and one below the plane of the \(\sigma\) bond axis. Pi bonds are weaker than sigma bonds because the side-to-side overlap is less efficient than the head-to-head overlap. A double bond consists of one \(\sigma\) bond and one \(\pi\) bond, while a triple bond contains one \(\sigma\) bond and two \(\pi\) bonds, and the presence of a \(\pi\) bond restricts rotation around the bond axis.
Explaining Molecular Geometry through Hybridization
Simple orbital overlap alone fails to explain the observed three-dimensional shapes of many molecules. For example, the carbon atom in methane (\(\text{CH}_4\)) should theoretically only form two bonds using its two unpaired \(p\) electrons, yet it forms four equivalent bonds with a tetrahedral geometry. To resolve this conflict, VBT introduces the concept of hybridization, which is the mathematical mixing of standard atomic orbitals on a single atom.
Hybridization combines an atom’s valence \(s\) and \(p\) orbitals to create a new set of equivalent, directional hybrid orbitals. For carbon in methane, one \(s\) and all three \(p\) orbitals mix to form four equivalent \(\text{sp}^3\) hybrid orbitals, which naturally arrange themselves in a tetrahedral shape. This mixing maximizes the orbital overlap and leads to stronger bonds.
The specific type of hybridization determines the molecular geometry and the number of bonds an atom can form. When an atom mixes one \(s\) and two \(p\) orbitals, it forms three \(\text{sp}^2\) hybrid orbitals arranged in a trigonal planar geometry, leaving one unhybridized \(p\) orbital available to form a \(\pi\) bond. If one \(s\) and one \(p\) orbital mix, two \(\text{sp}\) hybrid orbitals are formed in a linear arrangement, leaving two unhybridized \(p\) orbitals for forming two \(\pi\) bonds, as seen in a triple bond.
The Limitations of Valence Bond Theory
Despite its success in explaining molecular geometry, Valence Bond Theory encounters conceptual difficulties when describing certain chemical phenomena. A significant failure lies in its inability to easily account for electron delocalization, particularly in molecules like benzene (\(\text{C}_6\text{H}_6\)). VBT assumes bonds are localized between two specific atoms, while the bonding electrons in benzene are spread out over the entire ring structure.
To address delocalization, VBT relies on the concept of resonance, which requires depicting the molecule as a blend of multiple theoretical structures. This need for multiple contributing structures highlights the model’s struggle to natively represent electrons shared by more than two atoms. VBT also often fails to predict or explain the magnetic properties of certain simple molecules.
For example, the theory predicts that molecular oxygen (\(\text{O}_2\)) should be diamagnetic (all electrons paired). However, experiments show that oxygen is paramagnetic, meaning it possesses unpaired electrons and is attracted to a magnetic field. Since VBT’s core premise is that bonds form by pairing up electrons, it lacks an internal mechanism to explain these unpaired electrons in the molecule’s ground state.
How VBT Compares to Molecular Orbital Theory
Valence Bond Theory is often contrasted with Molecular Orbital Theory (MOT), an alternative model for describing chemical bonding. The central difference is how each theory treats electrons: VBT views bonding as localized between two atoms using atomic or hybrid orbitals. In contrast, MOT treats the entire molecule as a single entity where atomic orbitals combine to form new molecular orbitals that are delocalized over all the atoms.
VBT’s strength lies in its simplicity and its intuitive explanation of molecular shape through hybridization, making it the preferred model for predicting geometry in introductory chemistry. Conversely, MOT is superior in its ability to accurately explain phenomena that VBT struggles with, such as electron delocalization and magnetic properties. For instance, MOT successfully predicts the paramagnetism of oxygen because its molecular orbitals naturally allow for unpaired electrons. The two theories are not mutually exclusive but rather complementary, offering different perspectives on the nature of the chemical bond.