Liquids and Intermolecular Forces

This set of vocabulary flashcards covers the fundamental concepts of intermolecular forces, liquid properties, phase transitions, and phase diagrams as discussed in General Chemistry chapter 11.

Intermolecular Forces (IMFs)

The attractive forces between particles that tend to draw them together, which are significantly weaker than intramolecular attractions (atomic bonds).

London dispersion forces (LDFs)

The weakest intermolecular forces arising from the constant, random motion of electrons that momentarily create instantaneous dipoles and induce dipoles in neighboring particles.

Polarizability

The ease with which an electron cloud is distorted; atoms or molecules with more electrons are more polarizable (or "squishier"), leading to stronger London dispersion forces.

Dipole-dipole interactions

Attractions between the permanent dipoles of polar molecules that orient themselves to maximize attractive forces and minimize repulsive forces.

Hydrogen bonding

A particularly strong form of dipole–dipole interaction occurring when hydrogen is covalently bonded to a highly electronegative atom ($N$, $O$, or $F$) and is attracted to a lone pair on a neighboring $N$, $O$, or $F$ atom.

Hydrogen bond donor

A hydrogen atom that is covalently bonded to nitrogen, oxygen, or fluorine.

Hydrogen bond acceptor

The lone pair on a neighboring $N$, $O$, or $F$ atom to which a hydrogen bond donor is attracted.

Surface tension

A property of a liquid surface that enables it to resist external forces, caused by cohesive intermolecular forces exerting a net inward pull on surface molecules.

Cohesive forces

Intermolecular forces that act among identical molecules within a liquid.

Adhesive forces

Intermolecular forces of attraction between liquid molecules and the surface of a surrounding material.

Capillary action

The ability of a liquid to rise against gravity within a narrow space due to the combined effects of cohesive and adhesive forces.

Viscosity

A measure of a fluid's resistance to flow, reflecting the internal friction between adjacent layers moving past one another.

Poise ($mPa \cdot s$)

The unit of measurement for viscosity.

Endothermic Transitions

Processes that require an input of energy ($+\Delta H$) to separate particles by weakening intermolecular forces, such as melting, vaporization, and sublimation.

Exothermic Transitions

Processes that release energy ($-\Delta H$) as particles slow down and intermolecular attractions become stronger, including freezing, condensation, and deposition.

Sublimation

A phase transition where a solid transitions directly to a gas without passing through the liquid phase.

Deposition

A phase transition where a gas transitions directly to a solid.

Sensible heat

Heat that produces a measurable change in the temperature of a substance within a single phase.

Latent heat

Heat absorbed or released during a phase change that occurs without a change in temperature.

Latent Heat of Fusion ($\Delta H_{fus}$)

The heat required to convert a solid to a liquid at its melting point.

Latent Heat of Vaporization ($\Delta H_{vap}$)

The heat required to convert a liquid to a gas at its boiling point.

Vapor pressure

The pressure exerted by a vapor in equilibrium with its liquid in a closed system at a given temperature.

Escape energy ($E_0$)

The threshold of kinetic energy required for particles at the surface of a liquid to overcome intermolecular attractions and enter the vapor phase.

Dynamic equilibrium

A state where the rates of evaporation and condensation are equal, resulting in constant amounts of liquid and vapor.

Volatile

A description for liquids with weak intermolecular forces that vaporize readily and have high vapor pressures.

Boiling point

The temperature at which a liquid's vapor pressure equals the external atmospheric pressure.

Normal boiling point

The boiling point of a substance at a standard pressure of $1\,atm$.

Phase diagram

A graph showing the physical state of a substance under different conditions of temperature (x-axis) and pressure (y-axis).

Triple Point

The unique temperature and pressure at which the solid, liquid, and gas phases of a substance all exist simultaneously in equilibrium.

Critical Point

The endpoint of the liquid–gas boundary beyond which the distinction between liquid and gas disappears.

Supercritical fluid

A substance existing beyond the critical point that possesses gas-like mobility and liquid-like density.