Liquids and Intermolecular Forces

This set of vocabulary flashcards covers states of matter, various types of intermolecular forces, liquid properties like viscosity and surface tension, and the energetics and equations associated with phase changes.

Condensed States

The states of matter, specifically solids and liquids, where molecules are brought more closely together by stronger intermolecular forces compared to gases.

Crystalline Solids

Solids in which atoms are arranged in a well-ordered 3D array.

Amorphous Solids

Solids that exhibit no long-range order in their atomic arrangement.

Intermolecular Forces

Forces that determine the states of matter based on small charges over long distances; they are much weaker than chemical bonds.

Dispersion Forces (London Dispersion Forces)

The only intermolecular force present between all molecules and atoms, resulting from instantaneous or temporary dipoles created by electron distribution.

Polarizability

The ease with which an electron cloud can be distorted; higher polarizability leads to stronger dispersion forces.

Dipole-Dipole Interaction

An intermolecular force present in molecules with permanent dipoles where the positive end of one dipole attracts the negative end of another.

Miscibility

The ability of two liquids to mix without separating into distinct layers.

Hydrogen Bonding

A strong type of dipole-dipole interaction occurring in polar molecules with $H$ bonded to small, highly electronegative atoms such as $F$, $O$, or $N$.

Ion-Dipole Interaction

The strongest type of intermolecular force, occurring when an ionic compound is mixed with a polar compound.

Surface Tension

The energy required to increase the surface area of a liquid by a unit amount, caused by the tendency of liquids to minimize their surface area.

Viscosity

A liquid's resistance to flow; it increases with stronger intermolecular forces and molar mass, and usually decreases as temperature increases.

Poise (P)

The unit of measurement for viscosity, defined as $g \text{ cm}^{-1} \text{ s}^{-1}$.

Cohesion

The attraction between molecules of the same substance for each other.

Adhesion

The attraction between molecules of a liquid and a surface.

Capillary Action

The ability of liquids to flow against gravity in narrow tubes, driven by the balance of cohesive and adhesive forces.

Vaporization

The endothermic process of a substance converting from a liquid state to a gaseous state.

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

The amount of heat required to convert one mole of liquid into a gas at its boiling point.

Dynamic Equilibrium

A state in a closed system where the rate of condensation and the rate of vaporization are equal.

Vapor Pressure

The pressure exertion by a gas when it is in dynamic equilibrium with its liquid state.

Boiling Point

The temperature at which a liquid's vapor pressure is equal to the external pressure.

Normal Boiling Point

The temperature at which a liquid's vapor pressure is exactly $1 \text{ atm}$.Scale.

Clausius-Clapeyron Equation

The mathematical relationship used to relate vapor pressure and temperature: $\ln P = \frac{-\Delta H_{vap}}{R}(\frac{1}{T}) + \ln C$.

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

The endothermic amount of heat required to turn one mole of a solid into a liquid.

Sublimation

The process where molecules break free from a solid and go directly into the gas phase.

Heat (Enthalpy) of Sublimation ($\Delta H_{sub}$)

The total heat required to turn one mole of solid into a gas, calculated as $\Delta H_{sub} = \Delta H_{fus} + \Delta H_{vap}$.

Critical Point

The point on a phase diagram where the gas and liquid phases combine into a supercritical fluid.

Critical Temperature

The specific temperature at which a supercritical fluid is formed and the distinct liquid and gas phases cease to exist.