Intermolecular forces (IMFs) are the attractive forces that exist between individual molecules, holding them together in a liquid or solid state. These forces are fundamentally different from the strong covalent or ionic bonds that hold atoms together within a single molecule. The strength of these attractions dictates how much energy is required to separate the molecules. This strength determines the physical state of a substance at a given temperature, explaining why substances exist as gases, liquids, or solids.
The Mechanism of London Dispersion Forces
London Dispersion Forces (LDFs) are a type of attraction that occurs between all molecules, regardless of their polarity, making them a universal force. This attraction arises from the momentary, uneven distribution of electrons within a molecule’s electron cloud. Electrons are constantly in motion, and at any instant, they may cluster on one side of a molecule, creating a temporary, instantaneous dipole.
This temporary partial charge on one molecule influences a neighbor, causing its electron cloud to distort and form an induced dipole. The resulting weak electrostatic attraction between these transient dipoles is the London Dispersion Force. The strength of LDFs is directly related to a molecule’s size and mass. Larger molecules have more electrons and more dispersed electron clouds, making them easier to distort (polarizability), which increases the force.
The Specificity and Strength of Hydrogen Bonding
Hydrogen bonding is a particularly strong form of a dipole-dipole interaction. This interaction is highly specific, requiring a hydrogen atom to be covalently bonded to a small, highly electronegative atom, most commonly Nitrogen (N), Oxygen (O), or Fluorine (F). The large difference in electronegativity pulls the electron density away from the hydrogen atom, leaving it with a substantial partial positive charge.
This highly exposed, partially positive hydrogen atom is then strongly attracted to a lone pair of electrons on a nearby highly electronegative atom (N, O, or F) in an adjacent molecule. This creates an attraction significantly stronger than a regular dipole-dipole force due to the small size of the atoms involved and the concentrated charge separation. Hydrogen bond strengths typically range from approximately 4 kJ/mol to 50 kJ/mol.
Comparing Relative Strengths and Determining Dominance
When comparing individual interactions, hydrogen bonding is substantially stronger than London Dispersion Forces. LDFs are generally considered the weakest of all intermolecular forces on a per-interaction basis. The bond energy of a single hydrogen bond is far greater than that of a single London dispersion interaction. This difference explains why small molecules like water, which can hydrogen bond, have properties vastly different from similarly sized nonpolar molecules like methane.
However, the cumulative effect of LDFs can sometimes overcome the strength of isolated hydrogen bonds. LDFs increase with molecular size, so in very large molecules, such as long hydrocarbon chains or polymers, the sheer number of transient attractions adds up. For example, a large nonpolar oil molecule may have thousands of individual LDFs along its length. The combined effect of these forces can be stronger than the few hydrogen bonds that a small molecule like ethanol can form.
How Intermolecular Forces Determine Physical Properties
The differences in the strength of these intermolecular forces are directly responsible for the varying physical properties observed in different substances. Stronger IMFs require more thermal energy to overcome, which translates directly to higher boiling and melting points. Water, for instance, has an unusually high boiling point of 100 °C because its extensive network of hydrogen bonds must be broken to transition into a gas.
In contrast, nonpolar substances like methane, which are only held together by weak LDFs, have very low boiling points and exist as gases at room temperature. Stronger attractions also lead to increased viscosity and higher surface tension. In biological systems, the strength and directionality of hydrogen bonds are responsible for the structural integrity of DNA, holding the two strands of the double helix together.