AP Chemistry Unit 5 Kinetics: Learning Reaction Mechanisms and Deriving Rate Laws

Elementary reaction (elementary step)

A single molecular event that occurs exactly as written, where the reactant particles in that step collide/rearrange in one move to form the step’s products.

Reaction mechanism

A proposed sequence of elementary steps that adds up to the overall chemical reaction and explains the experimentally observed rate law.

Molecularity

The number of reacting particles participating in a single elementary step (defined only for elementary steps, not overall reactions).

Unimolecular step

An elementary step involving one reactant particle that breaks apart or rearranges; typical rate law form is rate = k[A].

Bimolecular step

An elementary step involving collision of two reactant particles; typical rate law form is rate = k[A][B] (or k[A]^2 if the reactants are identical).

Termolecular step

An elementary step involving three particles colliding simultaneously; rare because three-body collisions are unlikely.

Elementary-step rate law rule

For an elementary step, the rate law exponents match the stoichiometric coefficients of the reactants in that step.

Overall reaction (in kinetics context)

The net balanced equation obtained by summing mechanism steps; its coefficients generally cannot be used to write the rate law.

Collision model (conceptual link)

Idea that reaction rate depends on how often the required particles meet; e.g., bimolecular frequency scales with [A][B].

Transition state

A high-energy arrangement of atoms along the reaction pathway for an elementary step (represented by a peak on an energy diagram).

Activation energy (Ea)

The energy barrier that must be overcome for a particular elementary step to occur.

Potential energy diagram (multi-step)

Energy vs. reaction progress plot where each elementary step corresponds to a peak (“hump”); multi-step mechanisms show multiple humps.

Rate-determining step (RDS)

The slowest elementary step in a mechanism that acts as a bottleneck and often controls the overall reaction rate.

Intermediate

A species produced in one mechanism step and consumed in a later step; it cancels out and does not appear in the overall reaction.

Catalyst (in a mechanism)

A species consumed in an early step and regenerated in a later step; it appears in the mechanism but not in the net overall reaction.

Mechanism validity tests

A mechanism must (1) sum to the overall balanced equation and (2) produce a derived rate law consistent with the experimental rate law.

Experimental rate law

Rate expression determined from experiments (e.g., rate = k[A]^m[B]^n), where orders m and n are found empirically and may be non-integers.

Reaction order (with respect to a reactant)

The exponent on a reactant concentration in the experimental rate law, indicating how the rate depends on that reactant’s concentration.

Rate constant (k)

Proportionality constant in a rate law; its units depend on the overall reaction order.

Fast pre-equilibrium

Mechanism situation where an early reversible step rapidly reaches equilibrium, allowing an equilibrium relationship to replace an intermediate concentration.

Equilibrium substitution for an intermediate

Using K = [I]/([A][B]) (from a fast equilibrium A + B ⇌ I) to write [I] = K[A][B] and eliminate [I] from the rate law.

Observed rate constant (k_obs)

An effective constant that combines multiple constants from a mechanism (e.g., k_obs = kK) after substituting for intermediates.

Steady-state approximation

Method assuming an intermediate’s concentration stays approximately constant: d[I]/dt ≈ 0, meaning formation rate ≈ consumption rate (not that [I]=0).

Steady-state equation (intermediate balance)

Expression setting intermediate formation terms minus consumption terms ≈ 0, built from the elementary-step rate laws for every step that creates or destroys the intermediate.

Limiting-case behavior (steady-state result)

Interpretation that a complex steady-state rate law can simplify under certain conditions (e.g., depending on which term dominates the denominator), changing apparent orders with concentration.