Kinetics
What to know
Kinetics explains how fast reactions proceed. Key concepts: collision theory (molecules must collide with sufficient energy and correct orientation); activation energy (the energy barrier); and the effect of temperature, concentration, surface area, and catalysts on rate. A catalyst lowers activation energy without being consumed and does not change the thermodynamics of the reaction.
Rate laws express rate = k[A]^m[B]^n, where exponents are determined experimentally, not from coefficients. Reaction order is the sum of exponents. Zero-order reactions have constant rate; first-order reactions have a constant half-life. The Arrhenius equation relates rate constant to temperature and activation energy. Energy diagrams show reactants, products, activation energy, and the transition state; exothermic reactions release energy, endothermic absorb it. Reaction mechanisms must sum to the overall reaction, and the slowest step (rate-determining step) controls rate.
Traps
- Reading rate-law exponents directly from balanced-equation coefficients (only valid for elementary steps).
- Confusing thermodynamics (spontaneity) with kinetics (speed)-a spontaneous reaction can be slow.
- Assuming a catalyst changes the equilibrium constant or yield.
- Misreading an energy diagram's activation energy as the net energy change.
Worked example
Doubling [A] doubles the rate; doubling [B] quadruples it. Then order in A is 1, in B is 2, overall third order: rate = k[A][B]². Rule of thumb: determine exponents from how rate responds to concentration changes, never from coefficients.