collisions

Meaning: Collision theory says a reaction needs reactant particles to meet, yet most collisions do not give products. A collision is effective when the particles carry kinetic energy at least equal to Ea and approach with an orientation that lets bonds break and form. Rate rises with higher concentration (or gas pressure), larger solid surface area, and higher temperature. A catalyst lowers Ea, so a larger fraction of collisions is effective at the same temperature.

Equation: k = Ae^(−Ea/RT), where k is the rate constant (units depend on reaction order), A is the frequency factor reflecting collision frequency and orientation (same units as k), Ea is activation energy (J/mol), R = 8.314 J/(mol·K), and T is temperature (K).

On the MCAT: Passages give rates at different concentrations or temperatures, or Maxwell–Boltzmann curves, and ask you to predict the change. Warmer gas shifts the curve right, so more molecules exceed Ea. Elementary steps with three reactants are unlikely because three-body collisions are rare. In gas-law questions, collisions with the container walls explain pressure in kinetic molecular theory.

Common confusion: Students assume more collisions means a proportionally faster reaction. Raising temperature increases collision frequency only slightly; the large rate increase comes from the much larger fraction of collisions with energy ≥ Ea. A catalyst does not work by adding collisions; it lowers Ea.