Unit 9: States of Matter - Vocabulary Flashcards

Vocabulary flashcards for Unit 9: States of Matter, covering key terms and definitions from the lecture notes.

Energy

The ability to do work; it manifests in various forms and is conserved (neither created nor destroyed) in any process.

Potential Energy

Energy due to the composition or position of an object (PE = mgh).

Kinetic Energy

Energy of motion (KE = ½ mv2).

Kinetic Molecular Theory (KMT)

Explains the behavior of matter at a particle level, stating that particles are in constant, random motion, do not attract or repel each other, have elastic collisions, and kinetic energy depends on temperature.

Heat (Thermal) Energy

A form of energy that flows between two samples of matter due to their difference in temperature, flowing from hot to cold, and measured in calories or Joules.

System

A specific reaction or process being studied.

Surroundings

Everything else in the universe outside of the system.

Endothermic Process

A process where the system gains energy from the surroundings (heat is transferred from the surroundings to the system).

Exothermic Process

A process where the system loses energy to the surroundings (heat is transferred from the system to the surroundings).

Phase

States of substances that coexist as physically distinct parts of a mixture (e.g., ice water).

Melting

The process of a solid changing to a liquid.

Sublimation

The process of a solid changing directly to a gas.

Freezing

The process of a liquid changing to a solid.

Vaporization

The process of a liquid changing to a gas.

Condensation

The process of a gas changing to a liquid.

Deposition

The process of a gas changing directly to a solid.

Normal Freezing Point

The temperature at which the solid and liquid phases are in equilibrium at 1 atm pressure.

Boiling

Conversion of a liquid to a vapor within the liquid as well as at its surface.

Boiling Point

The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

Normal Boiling Point

The boiling point at normal atmospheric pressure (1 atm, 760 torr, or 101.3 kPa).

Evaporation

Occurs only at the surface of a liquid.

Phase Diagram

A graph of pressure versus temperature that shows in which phase a substance exists under different conditions.

Triple Point

The temperature and pressure at which three phases of a substance can coexist.

Critical Point

The pressure and temperature above which a substance cannot exist as a liquid.

Critical Temperature (Tc)

The temperature above which a substance cannot exist in the liquid state.

Critical Pressure (Pc)

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

Temperature

A measure of the average kinetic energy of the particles in a sample of matter.

Joule

SI unit of heat as well as all other forms of energy.

Specific Heat

The amount of heat required to raise the temperature of one gram of a substance one degree Celsius.

Calorimeter

An insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

Volume

The space matter occupies (measured in Liters (L) or milliliters (mL)).

Pressure

The force over a given area (measured in Atmospheres (atm), Kilopascals (kPa), torr, or mmHg).

STP (Standard Temperature and Pressure)

Temperature at 273 K and Pressure at 1 atm; used to compare gases.

Boyle’s Law

Explains the effect pressure has on volume; Temperature stays constant; inverse relationship (↑P↓V).

Charles’ Law

Explains the effect temperature has on volume; Pressure stays constant; direct relationship (↑T↑V).

Gay-Lussac’s Law

Explains the effect temperature has on pressure; Volume stays constant; direct relationship (↑T↑P).

Avogadro’s Principle

One mole (6.02 x 1023 particles) of any gas at STP occupies a volume of 22.4 L.

Combined Gas Law

Integration of the gas laws into a single equation.

Ideal Gas Law

Relates pressure, temperature, volume, and number of moles using the universal gas constant R.

Real Gas

A gas that does not behave completely according to the assumptions of the kinetic-molecular theory.

Dalton’s Law of Partial Pressures

The total pressure of a gas mixture is the sum of the partial pressures of the component gases.

Mole Fraction

Ratio of the number of moles of a certain component of a mixture to the total number of moles in the mixture.

Diffusion

Spontaneous mixing of the particles of two substances caused by their random motion.

Effusion

Process by which gas particles pass through a tiny opening.

Graham’s Law

Explains the effect mass has on velocity and is derived from the KE equation (1/2mv2).

Density

The ratio of an object's mass and volume (D = m / V).