Lewis structures illustrate the bonding arrangement of atoms within a molecule. These two-dimensional drawings serve as the starting point for predicting a molecule’s three-dimensional shape, which determines how it interacts with neighboring molecules. The attractive forces existing between individual molecules, known as intermolecular forces (IMFs), dictate a substance’s physical behavior, such as its boiling point or solubility. Identifying these forces requires a sequential analysis that begins with the Lewis structure.
Step One: Translating Lewis Structures into Molecular Geometry
The first step involves converting the Lewis structure into a three-dimensional shape using the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR is based on the premise that all regions of electron density around a central atom arrange themselves to minimize electrostatic repulsion. These electron groups include single bonds, multiple bonds, and lone pairs.
Counting the total number of electron groups around the central atom establishes the electron-group geometry. For example, four electron groups (such as in methane (\(\text{CH}_4\))) arrange into a tetrahedral geometry. If a central atom has three electron groups, the arrangement is trigonal planar.
The actual molecular geometry, defined only by the atomic nuclei, is determined by the ratio of lone pairs versus bonding pairs. A lone pair occupies more space than a bonding pair, exerting a greater repulsive force that compresses the angles between bonded atoms. For instance, a molecule with four total electron groups but one lone pair, like ammonia (\(\text{NH}_3\)), is shaped as a trigonal pyramid.
Step Two: Assessing Molecular Polarity
Determining the three-dimensional geometry is a prerequisite for the next stage, which is to assess whether the molecule possesses a net dipole moment, making it polar or nonpolar. This assessment requires a two-part analysis involving the polarity of the individual bonds and the overall molecular symmetry. A bond is considered polar if the two atoms sharing the electrons have a significant difference in electronegativity, causing the electron density to be pulled toward the more electronegative atom.
This unequal sharing creates a bond dipole, represented as a vector pointing toward the atom with the greater electron density, which establishes regions of partial negative (\(\delta-\)) and partial positive (\(\delta+\)) charge. Once all bond dipoles are identified, the molecular geometry from Step One is used to determine if these individual vectors cancel each other out in three-dimensional space.
If the molecular geometry is symmetrical and all surrounding atoms are identical, the bond dipoles will perfectly cancel, resulting in a nonpolar molecule with a net dipole moment of zero, even if the individual bonds are polar. Carbon dioxide (\(\text{CO}_2\)) is an example, where the two polar carbon-oxygen bonds pull equally in opposite directions along a linear axis. Conversely, if the geometry is asymmetrical or if the central atom has lone pairs that distort the charge distribution, the vectors do not cancel, classifying the molecule as polar.
Step Three: Identifying the Intermolecular Forces
The determination of molecular polarity provides the necessary information to classify the types of intermolecular forces (IMFs) acting between molecules. Every molecule, whether polar or nonpolar, experiences London Dispersion Forces (LDFs), which are the weakest of the three primary IMFs. These forces arise from the continuous, temporary fluctuations in electron distribution that create instantaneous, short-lived dipoles that induce dipoles in neighboring molecules. LDFs are the only IMFs present in nonpolar molecules, and their strength increases with the molecule’s size and mass because larger molecules have more electrons and are more easily polarized.
In addition to LDFs, molecules identified as polar in Step Two also experience Dipole-Dipole forces. These forces are the electrostatic attractions between the permanent partial positive end of one polar molecule and the permanent partial negative end of an adjacent polar molecule. Since these forces involve permanent dipoles rather than instantaneous ones, they are typically stronger than LDFs between molecules of comparable size.
A strong type of dipole-dipole attraction is known as Hydrogen Bonding. Hydrogen bonding occurs only when a hydrogen atom is covalently bonded directly to one of the three most electronegative atoms: nitrogen (N), oxygen (O), or fluorine (F). This specific arrangement creates a strong partial positive charge on the hydrogen atom, which is then attracted to a lone pair on a nearby N, O, or F atom of an adjacent molecule.
Connecting IMFs to Observable Properties
The successful identification and classification of a molecule’s intermolecular forces explain many of its bulk physical properties. The relative strength of the IMFs dictates how tightly molecules are held together in the condensed phases (liquid and solid). Stronger intermolecular forces require a greater input of energy to overcome the attraction between molecules, leading to a higher boiling point and melting point for the substance.
For example, water’s high boiling point compared to similar-sized molecules is a direct consequence of its ability to form extensive hydrogen bonds. Similarly, the principle of “like dissolves like” is explained by IMFs. Polar substances dissolve readily in polar solvents because they can form strong dipole-dipole attractions with the solvent molecules. Nonpolar substances, which rely only on LDFs, dissolve best in nonpolar solvents, forming comparable weak LDFs with the solvent.