Kinetics and Fundamental Equilibrium Concepts Flashcards

Vocabulary practice flashcards covering reaction kinetics, rate laws, order of reaction, integrated rate laws, collision theory, activation energy, reaction mechanisms, catalysts, and fundamental chemical equilibrium concepts.

Kinetics

The branch of chemistry concerned with the study of reaction rates and the factors that affect them.

Reaction Rate

The change in concentration of a reactant or product per unit of time, typically expressed in units of $M/s$.

Average Reaction Rate

The rate at which a reaction proceeds over a specified time period, calculated using initial and final concentrations or as the slope of a secant line on a concentration-versus-time graph.

Instantaneous Reaction Rate

The reaction rate at a specific moment in time, determined graphically from the slope of a line tangent to the concentration-versus-time curve at that exact moment.

Initial Reaction Rate

The instantaneous reaction rate evaluated at time zero ($t = 0$), immediately upon mixing the reactants.

Rate Law

An expression showing the mathematical relationship between the rate of a reaction, the rate constant ($k$), and the reactant concentrations raised to specific powers.

Rate Constant ($k$)

A proportionality constant in a rate law that is specific to a given reaction at a particular temperature and accounts for collision frequency and orientation effectiveness.

Reaction Order

The exponent to which a reactant concentration is raised in a rate law, representing the sensitivity of the reaction rate to changes in that reactant's concentration.

Differential Rate Law

A rate law expression that explicitly shows how the speed of a reaction depends directly on reactant concentrations.

Integrated Rate Law

An equation derived from a differential rate law that expresses the concentration of reactants or products as a function of time.

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

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

Zero-Order Reaction

A reaction whose rate is completely independent of the concentration of the reactant, expressed as $\text{Rate} = k$.

First-Order Reaction

A reaction whose rate is directly proportional to the concentration of a single reactant, with an integrated rate law of $\text{ln}[A]_t = -kt + \text{ln}[A]_0$.

Second-Order Reaction

A reaction whose rate depends on the concentration of a reactant raised to the second power, with an integrated rate law of $\frac{1}{[A]_t} = kt + \frac{1}{[A]_0}$.

Collision Theory

A theory stating that chemical reactions occur as a result of collisions between reactant particles, provided the collisions possess adequate kinetic energy and proper molecular orientation.

Activation Energy ($E_a$)

The minimum kinetic energy required during a molecular collision for reactant molecules to break existing bonds and form products.

Arrhenius Equation

An equation relating the rate constant ($k$) to activation energy ($E_a$), absolute temperature ($T$), and a frequency factor ($A$): $k = A \times e^{-\frac{E_a}{RT}}$.

Frequency Factor ($A$)

A constant term in the Arrhenius equation that reflects the collision frequency and the probability that collisions occur with favorable molecular orientation.

Reaction Mechanism

The sequence of individual elementary steps by which an overall chemical reaction takes place at the molecular level.

Elementary Step

A single molecular event or step in a reaction mechanism that represents an actual collision or rearrangement and cannot be broken down into simpler reactions.

Intermediate

A chemical species that is produced in an early elementary step of a mechanism and consumed in a later elementary step, preventing it from appearing in the overall chemical equation.

Molecularity

The number of reactant molecules, atoms, or ions involved in a single elementary step (e.g., unimolecular, bimolecular, or termolecular).

Rate-Determining Step

The slowest elementary step in a reaction mechanism sequence, which limits and dictates the overall reaction rate.

Catalyst

A substance that increases the rate of a chemical reaction by providing an alternate reaction pathway with a lower activation energy, without being consumed in the overall process.

Homogeneous Catalyst

A catalyst that exists in the same physical phase as the reacting species.

Heterogeneous Catalyst

A catalyst that exists in a different physical phase from that of the reacting species, most commonly a solid catalyst interacting with gas or liquid reactants.

Chemical Equilibrium

A dynamic state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, causing reactant and product concentrations to remain constant over time.

Equilibrium Constant ($K_c$)

The numerical ratio of product molar concentrations to reactant molar concentrations at equilibrium, each raised to the power of its stoichiometric coefficient in the balanced equation.

Reaction Quotient ($Q$)

A calculated ratio of product concentrations to reactant concentrations structured like the equilibrium expression, evaluated at any given point in time to determine the system's direction toward equilibrium.

Le Châtelier's Principle

The principle stating that if a dynamic equilibrium system is subjected to a disturbance or stress, the system will shift its equilibrium position to counteract the effect of that stress.

Exothermic Reaction

A chemical reaction that releases heat energy to its surroundings, characterized by a negative enthalpy change ($\text{Δ}H < 0$).

Endothermic Reaction

A chemical reaction that absorbs heat energy from its surroundings, characterized by a positive enthalpy change ($\text{Δ}H > 0$).

Successful Collision

A collision between reactant particles that occurs with both proper orientation and kinetic energy equal to or greater than the activation energy, resulting in product formation.

Multi-Step Energy Profile

A reaction progress diagram illustrating potential energy changes across multiple elementary steps, where peaks represent transition states and valleys represent reaction intermediates.