10.1 Intermolecular Forces

Molecules attract each other through different types of forces, and these forces vary widely in strength. Understanding intermolecular forces is essential for explaining why substances have different boiling points, why some liquids flow easily while others don't, and why water behaves so unusually compared to similar molecules.

Types of Intermolecular Forces

Explain the three main types of intermolecular forces: dispersion forces, dipole-dipole attractions, and hydrogen bonding

Every molecule experiences at least one type of intermolecular force. The three main types build on each other in strength, and recognizing which ones are present in a substance is the key skill for this section.

Dispersion forces (London dispersion forces) are the weakest intermolecular force, but they're present in all molecules, including nonpolar ones like $\ce{H2}$, $\ce{Cl2}$, and $\ce{CH4}$. They arise because electrons are constantly moving around in a molecule. At any given instant, the electrons might be unevenly distributed, creating a temporary dipole. That temporary dipole can then induce a dipole in a neighboring molecule, producing a brief attractive force. Individually these attractions are tiny, but they add up. Dispersion forces get stronger as molecular size and surface area increase, because larger molecules have more electrons and more opportunities for these temporary dipoles to form.

Dipole-dipole attractions occur between polar molecules that have permanent dipoles, such as $\ce{HCl}$, $\ce{SO2}$, and $\ce{CHCl3}$. The partially positive end of one molecule is attracted to the partially negative end of another. These are stronger than dispersion forces alone, and their strength depends on the size of the dipole moment. A larger difference in electronegativity between bonded atoms means a bigger dipole moment and a stronger attraction.

Hydrogen bonding is a special, extra-strong type of dipole-dipole attraction. It occurs when hydrogen is bonded directly to nitrogen, oxygen, or fluorine (N, O, or F). These three atoms are small and highly electronegative, so they pull electron density strongly away from hydrogen. Because hydrogen is so small, the resulting positive charge is concentrated in a tiny area, allowing it to get very close to the lone pair on a neighboring N, O, or F atom. This close approach creates a surprisingly strong electrostatic attraction. Common examples include $\ce{H2O}$, $\ce{NH3}$, $\ce{HF}$, alcohols, and carboxylic acids.

Prediction of Intermolecular Forces

Predict the intermolecular forces present in substances based on molecular structure

To figure out which intermolecular forces a substance has, work through these steps:

1. Start with dispersion forces. Every molecule has them, so they're always present.
2. Check if the molecule is polar. If the molecule has a net dipole (asymmetric shape + polar bonds), add dipole-dipole attractions.
3. Check for hydrogen bonding. If the molecule has H bonded directly to N, O, or F, add hydrogen bonding.

Here's how that plays out for different categories:

• Nonpolar molecules have only dispersion forces, since there are no permanent dipoles. Examples: $\ce{H2}$, $\ce{Cl2}$, $\ce{CH4}$, and noble gases like $\ce{Ar}$.
• Polar molecules without H bonded to N, O, or F have dispersion forces and dipole-dipole attractions. Examples: $\ce{HCl}$, $\ce{SO2}$, $\ce{CHCl3}$.
• Polar molecules with H bonded to N, O, or F have all three: dispersion forces, dipole-dipole attractions, and hydrogen bonding. Examples: $\ce{H2O}$, $\ce{NH3}$, $\ce{HF}$, alcohols, amines, and carboxylic acids.

A common mistake is saying a molecule "only has hydrogen bonding." Remember, the forces are cumulative. If hydrogen bonding is present, dispersion and dipole-dipole forces are also at work.

Impact on Physical Properties

Connect intermolecular forces to physical properties, especially melting and boiling points

The general rule is straightforward: stronger intermolecular forces mean more energy is needed to pull molecules apart, which leads to higher melting and boiling points.

• Substances with only dispersion forces (noble gases, nonpolar hydrocarbons) tend to have the lowest boiling points. Methane ($\ce{CH4}$) boils at $-161\,°\text{C}$.
• Substances with dipole-dipole attractions boil at higher temperatures than similar-sized nonpolar molecules. Chloroform ($\ce{CHCl3}$), which has both dispersion and dipole-dipole forces, boils at $61\,°\text{C}$, far above methane.
• Substances with hydrogen bonding have the highest boiling points among molecules of similar size. Water ($\ce{H2O}$) boils at $100\,°\text{C}$, while hydrogen sulfide ($\ce{H2S}$), which is heavier but only has dipole-dipole attractions, boils at $-60\,°\text{C}$. That 160°C difference shows just how powerful hydrogen bonding is.

Intermolecular forces also affect other physical properties:

• Viscosity: Stronger forces mean higher viscosity (greater resistance to flow). Honey is more viscous than water partly because of extensive hydrogen bonding in its sugar molecules.
• Surface tension: Stronger forces mean higher surface tension (the liquid surface resists being broken). Water's high surface tension is why small insects can walk on it.
• Vapor pressure: Stronger forces mean lower vapor pressure. Molecules are held more tightly in the liquid phase, so fewer escape into the gas phase at a given temperature.

Effects of Intermolecular Forces on Liquid Properties

Several liquid behaviors trace directly back to intermolecular forces:

Cohesion is the attraction between molecules of the same substance. Liquids with stronger intermolecular forces have stronger cohesion, which is why water forms droplets rather than spreading into a thin film.

Adhesion is the attraction between molecules of different substances. When water climbs up the side of a glass, that's adhesion between water molecules and the glass surface.

Surface tension results from cohesion. Molecules at the surface of a liquid are only pulled inward and sideways by their neighbors (there are no liquid molecules above them). This net inward pull causes the surface to behave like a stretched elastic sheet, shrinking to the smallest possible area.

Capillary action is what happens when adhesion and cohesion work together. In a narrow tube, if the adhesive forces between the liquid and the tube walls are stronger than the cohesive forces within the liquid, the liquid climbs up the tube against gravity. This is how water moves upward through thin plant roots and how a paper towel soaks up a spill.