Intermolecular forces (IMFs) are the attractive forces that exist between neighboring molecules. They are fundamentally different from the much stronger intramolecular forces, such as covalent or ionic bonds, which hold atoms together within a single molecule. These weak attractions govern how molecules interact, determining how “sticky” they are. The strength of these forces dictates a substance’s physical state—solid, liquid, or gas—at a given temperature and pressure.
Forces Driven by Permanent Polarity
Dipole-dipole interactions occur between molecules that possess a permanent separation of charge, known as a permanent dipole moment. This permanent polarity arises when two atoms in a bond share electrons unequally due to a difference in their electronegativity. The resulting molecule has a partially negative end (\(\delta-\)) and a partially positive end (\(\delta+\)). These molecules naturally align themselves so that the positive end of one molecule is attracted to the negative end of a neighboring molecule. For example, in hydrogen chloride (\(\text{HCl}\)), the more electronegative chlorine atom pulls electron density toward itself, creating a permanent dipole moment that facilitates attraction between adjacent \(\text{HCl}\) molecules.
Hydrogen bonding is a particularly strong type of dipole-dipole interaction, often categorized separately due to its unique strength. This force only occurs when a hydrogen atom is covalently bonded to a highly electronegative atom: nitrogen (\(\text{N}\)), oxygen (\(\text{O}\)), or fluorine (\(\text{F}\)). This specific bonding arrangement leaves the hydrogen nucleus virtually unshielded, creating a much stronger partial positive charge than in typical polar bonds.
This highly positive hydrogen is then strongly attracted to a lone pair of electrons on a nearby \(\text{N}\), \(\text{O}\), or \(\text{F}\) atom of an adjacent molecule. Water (\(\text{H}_2\text{O}\)) provides the most common example, where the partial positive hydrogen of one molecule is attracted to the partial negative oxygen of another. Because each water molecule can participate in up to four hydrogen bonds, this extensive network of strong forces gives water many unique properties, such as its high boiling point.
Forces Driven by Instantaneous Polarity
London Dispersion Forces (LDFs), sometimes referred to as Van der Waals forces, exist in all substances, though they are the only intermolecular forces present in non-polar molecules. These forces arise from the constant, random motion of electrons within a molecule’s electron cloud. At any given moment, the electrons may be distributed unevenly, creating a temporary, short-lived charge separation called an instantaneous dipole.
This fleeting dipole in one molecule can then influence the electron distribution in a nearby molecule, distorting its cloud and inducing a corresponding temporary dipole in the neighbor. The momentary attraction between these two induced dipoles constitutes the London Dispersion Force. The strength of LDFs is directly related to a molecule’s polarizability, which is the ease with which its electron cloud can be distorted.
Larger molecules with more electrons are generally more polarizable because their electrons are farther from the nucleus and less tightly held, leading to stronger LDFs. Molecular shape also plays a significant role. Elongated molecules, such as n-pentane, have a larger surface area for contact, allowing for greater interaction between neighboring molecules and stronger LDFs. Conversely, a more compact, spherical molecule like neopentane has a smaller contact area, resulting in weaker dispersion forces and a lower boiling point.
Interactions Involving Charged Particles
Ion-dipole forces represent a very strong type of intermolecular attraction that occurs when an ion interacts with a neutral, polar molecule. These forces are most often observed when an ionic compound dissolves in a polar solvent, such as when table salt (\(\text{NaCl}\)) is added to water. The fully charged ion, like the positive sodium cation (\(\text{Na}^{+}\)), strongly attracts the partially negative end of the polar water molecule.
Simultaneously, the negative chloride anion (\(\text{Cl}^{-}\)) attracts the partially positive end of the water molecule. This attraction causes the water molecules to surround and pull the individual ions away from the solid crystal lattice structure. This process, called hydration when water is the solvent, is driven by the energy released from forming these numerous ion-dipole attractions, which overcomes the strong forces holding the ions together in the solid.
How Intermolecular Forces Determine Physical States
The collective strength of a substance’s intermolecular forces directly determines its physical properties. Stronger attractions require more energy to overcome, resulting in higher melting points and higher boiling points. For a liquid to transition into a gas, the molecules must gain enough thermal energy to escape the attractive pull of their neighbors. Strong IMFs also result in higher viscosity, which is a liquid’s resistance to flow, because molecules are held together more tightly.
The relative strength of the forces discussed follows a specific order. For example, the high boiling point of water (100 °C) compared to other similar-sized molecules is due to the powerful effect of hydrogen bonding.
Relative Strength of IMFs
- Ion-Dipole forces (Strongest)
- Hydrogen Bonding
- Dipole-Dipole interactions
- London Dispersion Forces (Weakest)
Intermolecular forces also govern solubility, often summarized by the rule “like dissolves like.” Polar solutes (which use Dipole-Dipole or Hydrogen Bonding) will readily dissolve in polar solvents like water because the new solute-solvent interactions are similar in type and strength to the original solvent attractions. Conversely, non-polar substances, which rely only on LDFs, will not dissolve in water because the weak LDFs cannot compete with the water’s strong hydrogen bonds, causing the substances to remain separate.