12 Intermolecular Forces and States of Matter

Particles, attractions, and physical behavior

Intermolecular forces are attractions between separate particles, such as molecules, atoms, or ions. They are different from chemical bonds within a particle. A water molecule, for example, can remain chemically intact while attractions between separate water molecules influence whether the sample is liquid or gas.

The state of a substance reflects a balance: attractions keep particles near one another, while thermal motion tends to spread them apart. Increasing temperature generally increases particles’ average kinetic energy; cooling generally decreases it. The kinds of attractions present, along with particle size, shape, and packing, help explain differences in physical properties.

Four important types of intermolecular force

Several kinds of intermolecular forces help explain how particles attract one another:

• London dispersion forces arise from temporary shifts in electron distribution. They occur between all particles, including individual atoms and nonpolar molecules. More electrons and more easily distorted electron clouds tend to strengthen these forces. For example, dispersion forces increase across the halogens from $\mathrm{F_2}$ to $\mathrm{I_2}$; their room-temperature phases range from gases through liquid $\mathrm{Br_2}$ to solid $\mathrm{I_2}$.

• Dipole–dipole attractions occur between polar molecules. The partially positive region of one molecule attracts the partially negative region of another, and molecular orientation affects the attraction.

• Hydrogen bonding is a particularly strong, directional dipole–dipole attraction. It occurs when hydrogen bonded to nitrogen, oxygen, or fluorine is attracted to a nitrogen, oxygen, or fluorine atom with a lone pair on a neighboring particle. Hydrogen bonding between water molecules contributes to water’s comparatively high boiling point.

• Ion–dipole attractions occur between ions and polar molecules. For example, water molecules interact with dissolved $\mathrm{Na^+}$ and $\mathrm{Cl^-}$ ions.

A particle or substance can experience more than one kind of attraction. Polar molecules, for instance, also have dispersion forces. When comparing substances, consider the kinds of forces present as well as molecular size, shape, and how closely particles can approach.

Connecting attractions to physical properties

Stronger attractions generally make it harder for particles to separate. For comparable substances, this often means higher melting and boiling points, lower vapor pressure, and greater viscosity and surface tension. These are useful trends, not absolute rules: molecular size, shape, and packing can also affect measured properties.

Vapor pressure is the pressure of vapor in equilibrium with its liquid in a closed container. Some surface molecules escape into the gas even below the boiling point. Stronger attractions make escape less likely, so they generally lower vapor pressure at a given temperature.

Boiling occurs when a liquid’s vapor pressure equals the external pressure. A liquid with stronger attractions generally must reach a higher temperature to boil at a given external pressure. Lower external pressure lowers the boiling temperature; higher external pressure raises it.

Viscosity describes resistance to flow. Stronger attractions can make neighboring liquid molecules harder to move past one another, increasing viscosity. Surface tension reflects the extra energy needed to increase a liquid’s surface area; attractions between surface molecules contribute to it.

A useful comparison strategy is to identify which attractions each substance can have, then consider particle size and shape. Avoid predicting a measured property from the name of an intermolecular force alone.

Phases, heating, and changes of state

In a solid, particles are close together and generally vibrate around relatively fixed positions. In a liquid, particles remain close but can move past one another. In a gas, particles are far apart and move freely.

A phase change occurs when conditions shift the balance between particle motion and attractions. During melting and vaporization, heat is absorbed; during freezing and condensation, heat is released. Sublimation changes a solid directly into a gas and absorbs heat, while deposition changes a gas directly into a solid and releases heat.

At a substance’s melting or boiling temperature at fixed pressure, added heat changes the phase rather than raising the temperature. The energy rearranges or separates particles against their attractions. Freezing and condensation release energy as particles form closer, more strongly interacting arrangements. These physical changes normally leave the particles’ chemical identities unchanged: vaporizing water does not break the $\mathrm{O-H}$ bonds within its molecules.

Evaporation and boiling are both forms of vaporization, but they are not the same. Evaporation occurs at a liquid’s surface and can happen below its boiling point. Boiling occurs throughout the liquid when its vapor pressure matches the surrounding pressure. Thus, liquid in an open container can evaporate without boiling.

Reading phase diagrams

A phase diagram shows which phase is stable at different temperatures and pressures. Boundaries mark conditions where two phases coexist in equilibrium. The triple point is where solid, liquid, and gas coexist. At the critical point, the liquid–gas boundary ends; above the critical temperature and pressure, there is no distinct boundary between liquid and gas.

Pressure can change which phase is favored, so a substance’s melting or boiling temperature is not always independent of its surroundings. To interpret a phase diagram, locate the temperature and pressure of interest, identify the region containing that point, and use the boundaries to determine where phases coexist or a transition occurs.

Takeaway: Intermolecular attractions influence how readily particles separate, while temperature and pressure affect their motion and the conditions under which phases are stable. Phase changes rearrange particles without normally changing their chemical identities.