Heat & Temperature

Intermolecular forces

Forces of attraction between particles, including hydrogen bonding, dipole-dipole interactions, and London dispersion forces.

Solid

A phase of matter with a definite shape and volume, most dense, difficult to compress, and characterized by fixed particle positions within a crystalline lattice.

Amorphous solids

Solids that do not have a regular crystalline structure, examples include glass, plastic, wax, and silly putty.

Liquid

A phase of matter with definite volume but no definite shape, where particles slide past each other and are hard to compress.

Gas

A phase of matter with no definite shape or volume, characterized by low density and little attraction between particles.

Vapor

A gaseous state of a substance that is normally liquid, e.g., water vapor.

Sublimation

The phase transition from solid directly to gas, often occurring in substances with very weak intermolecular forces.

Deposition

The phase transition from gas directly to solid.

Endothermic process

A physical change that absorbs energy, causing a substance to move from solid to liquid or liquid to gas.

Exothermic process

A physical change that releases energy, causing a substance to move from gas to liquid or liquid to solid.

Energy

The capacity to do work or produce heat, manifesting in various forms such as electrical, thermal, atomic, and mechanical.

Kinetic energy

The energy of motion, associated with the movement of particles.

Potential energy

Stored energy, such as the energy held in the bonds between atoms.

Law of Conservation of Energy

The principle stating that energy cannot be created or destroyed, only transformed from one form to another.

Thermal energy

The kinetic energy gained by atoms or molecules as heat is added, measured in Joules (J) or calories (cal).

Law of Thermodynamics

The principle that heat transfers from a higher temperature object to a lower temperature object until thermal equilibrium is reached.

Temperature

A measure of the average kinetic energy of particles in a sample, not a form of energy itself.

Celsius scale

A temperature scale based on the freezing point (0 °C) and boiling point (100 °C) of water.

Kelvin scale

A temperature scale based on kinetic energy, where the lowest temperature possible is absolute zero (0 K).

Absolute zero

The temperature at which particles have minimal kinetic energy and movement, equivalent to 0 Kelvin or -273 °C.

Phase transition

The transformation of a substance from one state of matter to another, such as solid to liquid, liquid to gas, and vice versa.

Phase diagram

A graphical representation showing the phases of a substance at varying temperatures and pressures.

Critical point

The temperature and pressure at which a substance can coexist in both liquid and gas phases.

Boiling point

The temperature at which a substance changes from liquid to gas under a given pressure.

Melting point

The temperature at which a substance changes from solid to liquid.

Heat capacity

The quantity of heat required to change a substance's temperature by one degree Celsius.

Latent heat

The energy absorbed or released during a phase transition without a change in temperature.

Density

The mass of a substance per unit volume, typically expressed in grams per cubic centimeter (g/cm³).

Chemical change

A change that results in the formation of new chemical substances, often indicated by a change in color, temperature, or the production of gas.

Physical change

A change that affects one or more physical properties of a substance without altering its chemical composition.

Equilibrium

A state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in stable concentrations of reactants and products.

Pressure

The force exerted per unit area, commonly affecting the state and phase transitions of matter.

Vapor pressure

The pressure exerted by a vapor in equilibrium with its liquid or solid form at a given temperature.

Saturation point

The point at which a liquid can dissolve no more solute at a given temperature and pressure.

Molarity

A measure of concentration, defined as the number of moles of solute per liter of solution.

Colligative properties

Properties that depend on the number of solute particles in a solution, such as boiling point elevation and freezing point depression.

Solubility

The maximum amount of solute that can dissolve in a given quantity of solvent at a specific temperature.

Hess's Law

The principle stating that the total enthalpy change during a chemical reaction is the same, regardless of the pathway taken.