Molecular structure determines how a molecule interacts with its environment and other substances. The specific three-dimensional arrangement of atoms dictates a molecule’s properties, including its polarity, reactivity, and biological function. To understand these properties, chemists rely on models that predict the spatial orientation of atoms within a molecule. This structural organization follows predictable rules governing how electrons arrange themselves around a central atom.
The VSEPR Theory Framework
The foundation for predicting molecular shape rests on the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory operates on the principle that the negatively charged electron groups surrounding a central atom repel one another. To minimize this electrostatic repulsion, these electron groups organize themselves to be as far apart in three-dimensional space as possible. This arrangement dictates the overall electronic structure of the molecule.
An electron group, or electron domain, is defined as any region of high electron density around the central atom. This includes a single bond, a double bond, a triple bond, or a non-bonding lone pair of electrons; crucially, a multiple bond counts as only one electron domain. The total number of these domains around the central atom is the primary factor that determines the molecule’s basic geometry.
Determining Electron Geometry by Electron Domains
Electron geometry specifically describes the spatial arrangement of all electron domains—both bonding pairs and lone pairs—around the central atom. The overall shape is determined by the total count of these domains, which are treated equally under the VSEPR theory. The resulting shape is named for the geometric figure created by connecting the positions of these electron domains.
The simplest structure is the linear geometry, which results from a central atom having only two electron domains (e.g., \(\text{BeCl}_2\)). These two domains position themselves at a \(180^\circ\) angle from each other to maximize separation. Increasing the count to three domains results in a trigonal planar geometry, where the domains are spaced \(120^\circ\) apart on a single plane (e.g., \(\text{BF}_3\)).
A molecule with four electron domains (e.g., \(\text{CH}_4\)) adopts a tetrahedral geometry, with bond angles of approximately \(109.5^\circ\). Five electron domains lead to a trigonal bipyramidal structure (e.g., \(\text{PCl}_5\)), featuring three equatorial positions and two axial positions. Finally, six electron domains arrange themselves in an octahedral geometry, forming \(90^\circ\) angles (e.g., \(\text{SF}_6\)). This electron geometry provides the foundational shape, regardless of whether the domains are involved in bonding or exist as lone pairs.
Electron Geometry Versus Molecular Geometry
While electron geometry considers all electron domains, molecular geometry describes the observable shape of the molecule, which is defined only by the arrangement of the atoms themselves. This distinction is necessary because lone pairs of electrons occupy space and influence the overall electronic arrangement, but they are not considered when naming the final molecular shape. The fundamental electron geometry and the molecular geometry are identical only when the central atom has no lone pairs.
The difference becomes apparent when lone pairs are present on the central atom, as they exert a greater repulsive force than bonding pairs. This increased repulsion compresses the angles between the bonding atoms, distorting the final molecular shape. For example, both methane (\(\text{CH}_4\)) and ammonia (\(\text{NH}_3\)) have four electron domains, giving both a tetrahedral electron geometry. However, the nitrogen atom in ammonia has one lone pair, causing the molecular geometry to be trigonal pyramidal. Similarly, the water molecule (\(\text{H}_2\text{O}\)) also has a tetrahedral electron geometry (two bonding pairs and two lone pairs), but the two lone pairs on the oxygen atom cause the molecular shape to be bent. Understanding the electron geometry is the first step, but recognizing the influence of the lone pairs is necessary to accurately predict the molecule’s final molecular geometry.