Unit 5: Kinetics — Unlocking Reaction Mechanisms

Elementary Reaction

Describes a single molecular event in a chemical reaction, such as a distinct collision between two particles.

Molecularity

Refers to the number of reactant particles involved in a single elementary step.

Unimolecular Reaction

A reaction where a single molecule rearranges or decomposes.

Bimolecular Reaction

A reaction involving two particles colliding.

Termolecular Reaction

A reaction involving three particles colliding simultaneously, which is extremely rare.

Rate Law for Unimolecular Reaction

Rate = k[A] when one reactant reacts.

Rate Law for Bimolecular Reaction

Rate = k[A][B] or Rate = k[A]^2 when two reactants collide.

Rate Law for Termolecular Reaction

Rate = k[A]^2[B] for three particles colliding.

Overall Reaction

A reaction that combines all steps in a reaction mechanism to show the net equation.

Reaction Mechanism

The proposed series of elementary steps that leads from reactants to products.

Intermediate

A species that is produced in an early step and consumed in a later step; not present in the final products.

Catalyst

Substance that increases the reaction rate without being consumed; is regenerated after being used.

Rate-Determining Step (RDS)

The slowest step in a multi-step reaction, which determines the overall rate.

Fast Equilibrium

A state where a fast step reaches equilibrium quickly, influencing the rate law.

Steady-State Approximation

Assumes that the concentration of a reactive intermediate remains constant during the major part of the reaction.

Formation Rate of an Intermediate

Rate = k_1[A][B] for the formation of an intermediate I from reactants A and B.

Consumption Rate of an Intermediate

Rate of consumption of an intermediate I is the sum of rates of processes consuming it.

Equilibrium Condition

A relationship that defines equal rates of forward and reverse reactions.

Substitution in Rate Laws

The process of substituting intermediates out of a rate law to express it in terms of reactants.

Common Mistake: Intermediates vs Catalysts

Confusing these two; catalysts are regenerated, intermediates are consumed.

Common Mistake: Coefficients in Overall Rate Laws

Applying coefficients from an overall reaction to rate laws without justification.

Common Mistake: Leaving Intermediates in Rate Laws

Not substituting intermediates when writing the final rate law expression.

Activation Energy (Ea)

The energy barrier that must be overcome for a reaction to proceed.

Equilibrium Rate Expression

The rate expression derived from an equilibrium condition for a fast step.

Reaction Intermediary

A transient species formed during the conversion of reactants to products.

Formal Rate Law

Specifies how the rate of a reaction is affected by the concentration of reactants.

Elementary Step

A single step in a reaction mechanism characterized by a specific molecularity.