kinetic molecular theory

Meaning: The theory explains gas behavior from particle motion. An ideal gas is assumed to have particles in constant, random, straight-line motion; particle volume negligible compared with the container; no attractions or repulsions between particles; elastic collisions, so no kinetic energy is lost; and average kinetic energy proportional to absolute temperature. Pressure comes from collisions with the walls. Raising T makes collisions harder and more frequent. Particles have a spread of speeds (a Maxwell–Boltzmann distribution), so statements refer to averages.

Equation: KEavg = ½mv² = (3/2)kBT, and urms = √(3RT/M), where m is particle mass (kg), v is speed (m/s), kB is the Boltzmann constant (J/K), T is absolute temperature (K), R is the gas constant (8.314 J/(mol·K)), and M is molar mass (kg/mol).

On the MCAT: Passages compare gases at the same T, asking which has the greater average speed (the lighter one) or effusion rate (Graham's law, rate ∝ 1/√M). Maxwell–Boltzmann graphs show that higher T shifts the peak right and flattens the curve. Questions also ask why real gases deviate from ideal behavior: at high pressure particle volume matters, and at low temperature attractions matter.

Common confusion: Same temperature means the same average kinetic energy, not the same speed, because mass differs. Also, T must be in kelvins: going from 20 °C to 40 °C does not double the average kinetic energy.