The Valence Shell Electron Pair Repulsion (VSEPR) theory is a straightforward model used in chemistry to predict the three-dimensional shape, or geometry, of molecules. This theory focuses specifically on the arrangement of electron groups surrounding a central atom within a chemical structure. It provides a simple, systematic way to understand why molecules adopt a particular spatial orientation, which significantly influences a substance’s physical and chemical characteristics.
The Core Principle of VSEPR
The fundamental premise of VSEPR theory is that all pairs of valence electrons around a central atom possess a negative charge and will naturally repel one another. This repulsion forces the electron groups to arrange themselves in three-dimensional space to achieve the maximum possible separation. By maximizing the distance between these negatively charged regions, the molecule minimizes its overall energy, leading to a stable configuration.
The strength of this repulsive force is not uniform across all electron groups. Bonding pairs are shared between two nuclei, while lone pairs—non-bonding electrons localized on the central atom—are only attracted to one nucleus, causing them to occupy more space. This establishes a hierarchy of repulsion: lone pair-lone pair repulsion is the strongest, followed by lone pair-bonding pair, and bonding pair-bonding pair repulsion is the weakest. This variation is essential for explaining subtle distortions in a molecule’s final shape and its bond angles.
Defining Electron Domains and Geometry
Applying the VSEPR model begins with identifying and counting the electron domains, or electron groups, around the central atom. An electron domain is defined as any region of high electron density, which can be a single bond, a multiple bond (double or triple), or a lone pair of electrons. A multiple bond is counted as a single electron domain because all electrons in that bond are confined to one localized region. The total number of these domains dictates the electron geometry, which is the overall spatial arrangement of all electron groups—both bonding and non-bonding—around the central atom.
The electron geometry is determined solely by the count of electron domains, leading to a set of five fundamental arrangements. These basic geometries represent the lowest-energy state for that specific number of electron groups.
- Two domains adopt a linear geometry with a \(180^{\circ}\) angle.
- Three domains arrange themselves in a trigonal planar geometry, forming \(120^{\circ}\) angles.
- Four domains adopt a tetrahedral geometry, where all angles are \(109.5^{\circ}\).
- Five domains create a trigonal bipyramidal arrangement.
- Six domains result in an octahedral electron geometry.
Predicting Molecular Shape and Bond Angles
While electron geometry describes the placement of all electron domains, the final molecular shape is defined only by the positions of the atomic nuclei. Lone pairs are not included in the description of the shape itself, and their presence on the central atom is what causes the molecular shape to differ from the electron geometry. For instance, both methane (\(\text{CH}_4\)) and ammonia (\(\text{NH}_3\)) have four electron domains, giving them both a tetrahedral electron geometry.
Methane has four bonding pairs and no lone pairs, so its molecular shape is also tetrahedral, with ideal bond angles of \(109.5^{\circ}\). Ammonia, however, has three bonding pairs and one lone pair, resulting in a trigonal pyramidal molecular geometry. Because the lone pair exerts a greater repulsive force than the bonding pairs, it pushes the three hydrogen atoms closer together, reducing the H-N-H bond angle from the ideal \(109.5^{\circ}\) to approximately \(107^{\circ}\).
The effect is even more pronounced in water (\(\text{H}_2\text{O}\)), which has two bonding pairs and two lone pairs. The two lone pairs create an even stronger repulsion, compressing the H-O-H bond angle further to about \(104.5^{\circ}\), resulting in a bent or angular molecular shape. The final molecular shape is a practical consequence of the lone pairs occupying space and dictating the spatial positions of the atoms that are actually visible.