Chapter 16: Kinetics - Rates and Mechanisms of Chemical Reactions

Vocabulary-style flashcards covering the fundamental concepts of chemical kinetics, rate laws, collision theory, and catalysis based on the Chapter 16 lecture transcript.

Chemical kinetics

The study of the speeds of chemical processes and the specific variables that influence those speeds.

Reaction rate

The manner in which the concentration of a reactant or product changes over a specific duration of time.

Rate Law

A mathematical expression, such as $\text{Rate} = k[A]^m[B]^n$, that relates the speed of a reaction to the concentrations of its reactants.

Rate constant ($k$)

A temperature-dependent proportionality constant in the rate law that is unique to a specific chemical reaction.

Reaction order

Determined experimentally, these exponents in the rate law indicate the degree to which the concentration of a specific reactant affects the overall rate.

Zero-order reaction

A reaction where the rate is independent of the reactant concentration, resulting in a constant rate over time.

First-order reaction

A reaction where the rate is directly proportional to the concentration of a single reactant.

Second-order reaction

A reaction where the rate is proportional to the square of a reactant concentration or the product of two reactant concentrations.

Half-life ($t_{1/2}$)

The time required for the concentration of a reactant to reach half of its initial value.

Collision Theory

The principle stating that for a reaction to occur, particles must collide with sufficient energy and the correct spatial orientation.

Activation energy ($E_a$)

The minimum energy threshold that colliding molecules must surpass to successfully transform into products.

Transition state

An unstable, high-energy arrangement of atoms representing the peak of the potential energy barrier during a chemical change.

Arrhenius equation

The equation $k = Ae^{-E_a/RT}$ that describes the relationship between the rate constant, temperature, and activation energy.

Frequency factor ($A$)

The product of the collision frequency ($Z$) and the orientation probability factor ($p$) in the Arrhenius equation.

Reaction mechanism

The chronological sequence of single molecular events, or elementary steps, that result in the overall chemical transformation.

Elementary step

A single reaction event within a larger mechanism whose rate law can be directly inferred from its balanced stoichiometry.

Molecularity

The specific number of reactant particles involved in an individual elementary step, such as unimolecular or bimolecular.

Rate-determining step

The slowest elementary step in a reaction mechanism, which governs the overall reaction rate.

Reaction intermediate

A chemical species that is generated in one step of a mechanism and subsequently consumed in a later step.

Catalyst

A substance that accelerates a reaction by offering an alternative pathway with a lower activation energy, without being consumed in the process.

Homogeneous catalyst

A catalyst that occupies the same physical phase as the reacting substances.

Heterogeneous catalyst

A catalyst that exists in a different physical phase than the reactants, such as a solid metal surface in a gas reaction.

Enzyme

A complex biological protein that functions as a highly specific catalyst for biochemical reactions.

Active site

The specific region of an enzyme where substrate molecules bind and undergo a chemical reaction.

Induced-fit model

A theory of enzyme action where the enzyme's active site undergoes a slight shape change to bind more effectively with its substrate.