Chem 1A Gas Laws and Kinetic Molecular Theory

These flashcards cover key concepts from the Chem 1A lecture on gas laws and the kinetic molecular theory, including definitions and relationships relevant to the behavior of gases.

Gas

Gases assume the volume and shape of their containers and are the most compressible state of matter.

Pressure (P)

The force exerted by the gas per unit area.

Volume (V)

The amount of space a gas occupies.

Temperature (T)

A measure of the average kinetic energy of gas particles.

Boyle's Law

Describes the relationship between pressure and volume; P1V1 = P2V2, assuming temperature and amount of gas are constant.

Charles's Law

Describes the relationship between volume and temperature; V1/T1 = V2/T2, assuming pressure and amount of gas are constant.

Avogadro's Law

States that the volume of a gas is directly proportional to the number of moles of the gas at constant temperature and pressure.

Ideal Gas Law

PV = nRT; relates pressure, volume, number of moles, gas constant, and temperature.

Dalton's Law of Partial Pressures

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

Mole Fraction (c)

The ratio of the moles of one component to the total moles of all components in a mixture, cA = nA/(nA+nB).

Kinetic Molecular Theory

A model that describes the behavior of gases in terms of particles in constant motion, with no attractive or repulsive forces between them.

Diffusion

The spreading of gas molecules throughout a container until uniformly distributed.

Effusion

The process by which gas molecules pass through a tiny opening into a vacuum or another container.

Graham's Law of Effusion

The rate of effusion of a gas is inversely proportional to the square root of its molar mass.

Root Mean Square Velocity

A measure of the average speed of gas particles, calculated using the equation u_rms = √(3RT/M).