Chapter 8 – Gases, Liquids, and Solids (Vocabulary Review)

A comprehensive set of vocabulary flashcards covering key terms, laws, and concepts from Chapter 8 on gases, liquids, and solids.

State of Matter

A physical form in which a substance exists—solid, liquid, or gas—determined by the balance between kinetic energy and intermolecular forces.

Phase Change

The conversion of a substance from one state of matter to another (e.g., melting, freezing, vaporization).

Endothermic Process

A phase change that absorbs heat (ΔH positive), such as melting, vaporization, or sublimation.

Exothermic Process

A phase change that releases heat (ΔH negative), such as freezing, condensation, or deposition.

Intermolecular Forces (IMFs)

Attractive forces acting between molecules or discrete atoms that influence physical properties; also called van der Waals forces.

London Dispersion Forces

Weak temporary attractions present in all molecules, arising from momentary electron polarization; strength increases with molecular size and surface area.

Dipole–Dipole Forces

Attractions between the positive end of one polar molecule and the negative end of another polar molecule.

Hydrogen Bonding

A strong dipole–dipole interaction between a hydrogen bonded to O, N, or F and a nearby O, N, or F atom.

Kinetic-Molecular Theory of Gases

A set of assumptions describing an ideal gas: random motion, negligible volume, elastic collisions, and kinetic energy proportional to Kelvin temperature.

Ideal Gas

A hypothetical gas that exactly obeys all postulates of the kinetic-molecular theory and the gas laws under all conditions.

Pressure (P)

Force per unit area exerted by gas particles colliding with a surface.

Atmospheric Pressure

The pressure exerted by the weight of the atmosphere at Earth’s surface, equal to 1 atm or 760 mmHg at sea level.

Pascal (Pa)

SI unit of pressure equal to one newton per square meter; 1 atm = 101,325 Pa.

Millimeter of Mercury (mmHg)

A common pressure unit based on the height of a mercury column in a barometer; 1 mmHg = 1 torr.

Barometer

An instrument that measures atmospheric pressure, often using a column of mercury.

Boyle’s Law

For a fixed amount of gas at constant temperature, pressure is inversely proportional to volume (P₁V₁ = P₂V₂).

Charles’s Law

For a fixed amount of gas at constant pressure, volume is directly proportional to Kelvin temperature (V₁/T₁ = V₂/T₂).

Gay-Lussac’s Law

For a fixed amount of gas at constant volume, pressure is directly proportional to Kelvin temperature (P₁/T₁ = P₂/T₂).

Combined Gas Law

Relates pressure, volume, and temperature for a fixed amount of gas (P₁V₁/T₁ = P₂V₂/T₂).

Avogadro’s Law

At constant temperature and pressure, volume is directly proportional to moles of gas (V₁/n₁ = V₂/n₂).

Standard Temperature and Pressure (STP)

Reference conditions for gases: 0 °C (273 K) and 1 atm (760 mmHg).

Standard Molar Volume

The volume occupied by one mole of an ideal gas at STP, equal to 22.4 L.

Ideal Gas Law

The equation PV = nRT that relates pressure, volume, temperature, and moles of an ideal gas.

Gas Constant (R)

Proportionality constant in the ideal gas law; common values are 0.0821 L·atm / mol·K and 62.4 L·mmHg / mol·K.

Partial Pressure

The pressure exerted by a single gas in a mixture, as if it were alone in the container.

Dalton’s Law of Partial Pressures

The total pressure of a gas mixture equals the sum of the partial pressures of each component (Ptotal = Σ Pi).

Vapor

Gas-phase molecules that are in equilibrium with the liquid phase of the same substance.

Vapor Pressure

The partial pressure of vapor molecules above a liquid at equilibrium; increases with temperature and weaker IMFs.

Normal Boiling Point

The temperature at which a liquid’s vapor pressure equals exactly 1 atm.

Viscosity

A liquid’s resistance to flow; increases with stronger intermolecular forces.

Surface Tension

Energy required to increase the surface area of a liquid due to unequal IMFs at the surface.

Crystalline Solid

A solid whose particles are arranged in an orderly, repeating three-dimensional lattice.

Amorphous Solid

A solid lacking long-range order; particles are arranged irregularly (e.g., glass, plastic).

Ionic Solid

A crystalline solid composed of cations and anions held together by ionic bonds (e.g., NaCl).

Molecular Solid

A crystalline solid whose particles are molecules held together by intermolecular forces (e.g., ice).

Covalent Network Solid

A solid where atoms are linked by covalent bonds in a giant, continuous network (e.g., diamond).

Vaporization (Evaporation)

Endothermic phase change from liquid to gas.

Condensation

Exothermic phase change from gas to liquid.

Melting (Fusion)

Endothermic phase change from solid to liquid.

Freezing

Exothermic phase change from liquid to solid.

Sublimation

Endothermic phase change directly from solid to gas.

Deposition

Exothermic phase change directly from gas to solid.

Elastic Collision

A collision in which gas particles rebound without loss of total kinetic energy.

Specific Heat

The quantity of heat needed to raise the temperature of 1 g of a substance by 1 °C.

Kinetic Energy (of Gas Particles)

Energy of motion proportional to the Kelvin temperature of the gas.