Intermolecular forces (IMFs) are the attractive forces that operate between individual molecules, and they are distinct from the much stronger intramolecular forces, which are the covalent bonds holding atoms within a molecule together. These external forces are electrostatic in nature, arising from the attraction between opposite partial charges on neighboring molecules. The strength of a substance’s intermolecular forces directly determines many of its bulk physical properties, such as its melting point, boiling point, and viscosity. Stronger attractions require significantly more energy to overcome, leading to higher temperatures required for a phase change from liquid to gas.
The Three Categories of Intermolecular Forces
The weakest attractive forces are the London Dispersion Forces (LDFs), which are present in all molecules. LDFs arise from the constant movement of electrons, which temporarily creates an uneven charge distribution, resulting in an instantaneous dipole moment. This momentary imbalance induces a corresponding dipole in a nearby molecule, leading to a weak, transient attraction. LDF strength increases with molecular size, as larger molecules have more electrons and diffuse electron clouds that are easier to distort (polarizability).
Molecules that possess a permanent separation of charge, meaning they are polar, experience an additional type of attraction called dipole-dipole forces. These forces occur when the partially positive end of one polar molecule is electrostatically attracted to the partially negative end of an adjacent polar molecule. This attraction is stronger and more consistent than London Dispersion Forces because the dipole moment is permanent, not temporary. The net strength of the dipole-dipole force is proportional to the magnitude of the molecule’s overall polarity.
A specific and strong form of dipole-dipole interaction is known as Hydrogen Bonding. This attraction requires a hydrogen atom to be covalently bonded directly to one of three highly electronegative atoms: nitrogen (N), oxygen (O), or fluorine (F). The large electronegativity difference in these \(\text{H-N}\), \(\text{H-O}\), or \(\text{H-F}\) bonds creates an extremely polarized bond, leaving the hydrogen atom with a large partial positive charge. This partially positive hydrogen is then strongly attracted to a lone pair of electrons on a nitrogen, oxygen, or fluorine atom of a neighboring molecule.
Determining Molecular Shape and Polarity
To identify the intermolecular forces present in ammonia (\(\text{NH}_3\)), its molecular shape and polarity must first be determined. The structure is predicted using Valence Shell Electron Pair Repulsion (VSEPR) theory. The central nitrogen atom has five valence electrons, forming three single covalent bonds with hydrogen atoms and possessing one lone pair of electrons.
These four electron regions (three bonding pairs and one lone pair) arrange themselves in a tetrahedral geometry to minimize repulsion. The lone pair occupies more space than the bonding pairs, exerting a greater repulsive force that distorts the geometry.
As a result of this distortion, the \(\text{NH}_3\) molecule adopts a trigonal pyramidal molecular shape, with the nitrogen atom at the apex and the three hydrogen atoms forming the base. Nitrogen is significantly more electronegative than hydrogen, meaning it pulls the shared electrons closer to itself. This unequal sharing creates individual bond dipoles directed toward the nitrogen atom. Because the molecular shape is asymmetrical, the three bond dipoles do not cancel each other out in three-dimensional space. The result is a significant net dipole moment for the entire molecule, with the negative pole concentrated near the nitrogen atom and the lone pair. Therefore, ammonia is confirmed to be a polar molecule.
Identifying the Forces Present in Ammonia
Since ammonia (\(\text{NH}_3\)) is a molecule, the universal London Dispersion Force (LDF) is present, contributing to the overall attraction between molecules. Because ammonia is confirmed to be a polar molecule with a permanent net dipole moment, it also experiences dipole-dipole interactions. The partially negative nitrogen atom of one molecule is attracted to the partially positive hydrogen atoms of a neighboring molecule. This force is substantially stronger than the LDFs.
The most powerful intermolecular force in ammonia is Hydrogen Bonding. Ammonia meets the structural requirement because each molecule contains hydrogen atoms directly bonded to nitrogen. The highly polarized \(\text{N-H}\) bond allows the hydrogen atom of one molecule to form a strong electrostatic attraction with the lone pair on the nitrogen atom of an adjacent molecule. The relative strength is ordered: Hydrogen Bonding is the most dominant, followed by Dipole-Dipole interactions, and finally, the weakest, London Dispersion Forces. The dominance of hydrogen bonding largely dictates the physical behavior of ammonia.
Physical Consequences of Ammonia’s Intermolecular Forces
The strong network of hydrogen bonds in liquid ammonia leads to an abnormally high boiling point compared to its heavier Group 15 counterparts, such as phosphine (\(\text{PH}_3\)). The expected trend is a steady increase in boiling points down a group due to increasing LDF strength, but ammonia breaks this trend. Ammonia boils at \(-33.3^{\circ}\text{C}\), significantly higher than phosphine’s \(-87.7^{\circ}\text{C}\), despite phosphine having a greater molecular weight. This difference highlights the strength of the hydrogen bonds in \(\text{NH}_3\), which require large amounts of energy to break.
The strong IMFs also explain ammonia’s high solubility in water. Water molecules also form hydrogen bonds, and ammonia can form extensive hydrogen bonds with water, acting as both a donor and an acceptor. This ability to mix readily allows for its dissolution, following the principle that “like dissolves like.”