Chapter 10: States of Matter Overview

Kinetic-Molecular Theory

A theory based on the idea that particles of matter are always in motion, used to explain the properties of solids, liquids, and gases in terms of particle energy and forces between them.

Ideal Gas

A hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory.

Elastic Collision

A collision between gas particles or between particles and container walls in which there is no net loss of total kinetic energy.

Kinetic Energy

Energy of motion, possessed by gas particles due to their continuous, rapid, random motion. Formula: KE = 1/2mv².

Expansion (of gases)

Gases do not have a definite shape or volume and completely fill any container in which they are enclosed, as gas particles move rapidly in all directions without significant attraction.

Fluidity (of gases)

The ability of gases to flow, like liquids, due to insignificant attractive forces between gas particles, allowing them to glide easily past one another.

Low Density (of gases)

Gaseous substances at atmospheric pressure have a density about 1/1000 the density of the same substance in liquid or solid state, because the particles are much farther apart.

Compressibility (of gases)

The property where gas particles, initially far apart, are crowded closer together during compression.

Diffusion

The spontaneous mixing of the particles of two substances caused by their random motion, where gases spread out and mix with one another even without being stirred.

Effusion

A process by which gas particles pass through a tiny opening. Note: Molecules of low mass effuse faster than molecules of high mass, and rates of effusion are directly proportional to particle velocities.

Real Gas

A gas that does not behave completely according to the assumptions of the kinetic-molecular theory. Deviations: Real gases deviate from ideal behavior at very high pressures and low temperatures, and if their molecules are more polar.

Liquid

A form of matter that has a definite volume and takes the shape of its container.

Intermolecular Forces (IMF)

Attractive forces between liquid particles, including London dispersion forces, dipole-dipole forces, and hydrogen bonding.

London Dispersion Forces

Weak forces present between non-polar molecules, resulting from temporary shifts in the density of electrons in electron clouds.

Dipole-Dipole Forces

Attractions between oppositely charged regions of polar molecules.

Hydrogen Bonds

Occur between molecules containing a hydrogen atom bonded to a small, highly electronegative atom (N, O, F) with at least one lone electron pair.

Fluids (Liquids)

A substance that can flow and therefore take the shape of its container.

Relative Incompressibility (of liquids)

Liquids are much less compressible than gases because their particles are more closely packed.

Surface Tension

A force that tends to pull adjacent parts of a liquid's surface together, decreasing surface area to the smallest possible size. Stronger forces of attraction lead to higher surface tension.

Capillary Action

The attraction of the surface of a liquid to the surface of a solid, which tends to pull liquid molecules upward against gravity (e.g., forming a meniscus).

Vaporization

The process by which a liquid or solid changes to a gas.

Evaporation

The process by which particles escape from the surface of a nonboiling liquid and enter the gas state, occurring because liquid particles have different kinetic energies.

Boiling

The change of a liquid to bubbles of vapor that appear throughout the liquid.

Freezing (Solidification)

The physical change of a liquid to a solid by removal of energy as heat.

Crystalline Solids

Solids that consist of crystals, where particles are arranged in an orderly, geometric, repeating pattern.

Amorphous Solid

A solid where the particles are arranged randomly, sometimes classified as supercooled liquids. They have no definite melting point (e.g., glass and plastics).

Melting

The physical change of a solid to a liquid by the addition of energy as heat. At the melting point, kinetic energies overcome attractive forces holding particles together.

Crystal Structure

The total three-dimensional arrangement of particles of a crystal.

Lattice

A coordinate system used to represent the arrangement of particles in a crystal.

Unit Cell

The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice.

Ionic Crystals

Consist of positive and negative ions arranged in a regular pattern (e.g., NaCl, KBr, CaCO3). They are hard, brittle, have high melting points, and are good insulators.

Covalent Network Crystals

Each atom is covalently bonded to its nearest neighboring atoms, with bonding extending throughout a large network (e.g., diamond, quartz). They are very hard, brittle, have very high melting points, and are usually nonconductors or semiconductors.

Metallic Crystals

Consist of metal cations surrounded by a sea of delocalized valence electrons, explaining their high electrical conductivity (e.g., Fe, Au, Ag).

Covalent Molecular Crystals

Consist of covalently bonded molecules held together by intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding). They have low melting points, are easily vaporized, relatively soft, and good insulators (e.g., table sugar).

Phase

Any part of a system that has uniform composition and properties.

Condensation

The process by which a gas changes to a liquid.

Vapor

A gas in contact with its liquid or solid phase.

Equilibrium

A dynamic condition in which two opposing changes occur at equal rates in a closed system.

Equilibrium Vapor Pressure of a Liquid

The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. This pressure increases with increasing temperature.

Volatile Liquids

Liquids that evaporate readily due to relatively weak forces of attraction between their particles (e.g., ether).

Nonvolatile Liquids

Liquids that do not evaporate readily due to relatively strong attractive forces between their particles (e.g., molten ionic compounds).

Boiling Point

The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure. The normal boiling point is at 1 atm (760 torr or 101.3 kPa).

Molar Enthalpy of Vaporization (∆Hv)

The amount of energy as heat needed to vaporize one mole of liquid at its boiling point at constant pressure. Its magnitude measures the attraction between liquid particles.

Melting Point

The constant temperature at which a pure crystalline solid changes to a liquid, where solid and liquid are in equilibrium and melting and freezing proceed at equal rates.

Molar Enthalpy of Fusion (∆Hf)

The amount of energy as heat required to melt one mole of solid at its melting point. Its magnitude depends on the attraction between solid particles.

Sublimation

The change of state from a solid directly to a gas.

Deposition

The reverse process of sublimation; the change of state from a gas directly to a solid.

Phase Diagram

A graph of pressure versus temperature that shows the conditions under which the phases of a substance exist.

Triple Point

The temperature and pressure conditions at which the solid, liquid, and vapor of a substance can coexist at equilibrium.

Critical Point

Indicates the critical temperature and critical pressure.

Critical Temperature (tc)

The temperature above which a substance cannot exist in the liquid state, regardless of how much pressure is applied.

Critical Pressure (Pc)

The lowest pressure at which a substance can exist as a liquid at the critical temperature.

Supercritical Fluids

Substances that exist above a critical point of pressure and temperature where no separate liquid and gas phases exist.