9 Elements of Chemical Kinetics and Catalysis - Vocabulary Flashcards

Vocabulary flashcards covering key concepts from chemical kinetics and catalysis to aid exam preparation.

Chemical kinetics

The branch of chemistry that studies rates of chemical reactions and the factors that influence these rates.

Thermodynamics vs. kinetics

Thermodynamics tells whether a reaction is feasible; kinetics tells how fast and by what path the reaction proceeds.

Time scale

Different reactions occur on different time scales, from seconds to years (e.g., s, 10^9 s, 10^15 s, etc.).

Reactants

Substances consumed in a chemical reaction.

Products

Substances formed in a chemical reaction.

Reaction rate

The change in concentration of a reactant or product per unit time (units: mol/L·s).

Average rate

Change in concentration over a specific time interval: Δ[Product]/Δt.

Instantaneous rate

Rate at a specific moment, obtained from the slope of the concentration vs. time graph.

Initial rate

Rate of the reaction at the very start (t = 0).

Reaction mechanism

A step-by-step sequence of elementary reactions that leads to the overall reaction.

Catalyst

A substance that speeds up a reaction without being consumed in the overall process.

Homogeneous catalysis

Catalysis where the catalyst is in the same phase as the reactants.

Heterogeneous catalysis

Catalysis where the catalyst is in a different phase than the reactants (e.g., solid catalyst in gas).

Enzyme catalysis

Biological catalysis by enzymes that speed up metabolic reactions.

Autocatalysis

A process in which the product acts as a catalyst for the reaction.

Activation energy (Ea)

The minimum energy required for a collision to lead to a reaction; the energy barrier.

Activated complex / Transition state

A short-lived, high-energy arrangement of atoms at the peak of the energy barrier during a reaction.

Energy profile diagram

A graph of potential energy vs. reaction coordinate showing Ea and relative energies of reactants and products.

Collision theory

A model where reactions occur when reactants collide with sufficient energy and proper orientation.

Effective collision

A collision with enough energy and correct orientation to form products.

Ineffective collision

A collision lacking sufficient energy or proper orientation, leading to no reaction.

Arrhenius equation

Relates rate constant to temperature: k = A e^(-Ea/RT).

Frequency factor (A)

A factor in the Arrhenius equation representing collision frequency and proper orientation.

Rate constant (k)

Proportionality constant in a rate law; its value depends on temperature and Ea.

Rate law

An equation that relates the rate to the concentrations of reactants, e.g., rate = k[A]^m[B]^n.

Reaction order

The sum of the exponents (m + n) in the rate law; indicates how rate responds to concentration changes.

Zero-order reaction

Rate is independent of reactant concentration; rate = k.

First-order reaction

Rate is proportional to the concentration of one reactant; rate = k[A].

Second-order reaction

Rate depends on the square of one reactant or on two different reactants; e.g., rate = k[A]^2 or rate = k[A][B].

Molecularity

The number of molecules involved in an elementary step (unimolecular, bimolecular, termolecular).

Elementary step

A single-step event with a defined molecularity that contributes to the mechanism.

Unimolecular

An elementary step involving one molecule (e.g., A → products).

Bimolecular

An elementary step involving two molecules (e.g., A + B → products).

Termolecular

An elementary step involving three molecules (rare in practice).

Intermediates

Species formed in one step and consumed in a later step; not present in the overall equation.

Transition state

The high-energy, unstable arrangement at the peak of the energy barrier during a reaction.

Rate-determining step

The slowest step in a multistep mechanism that controls the overall rate.

Steric factor

A factor accounting for the orientation/geometry of colliding molecules affecting the probability of reaction.

Beer-Lambert Law

A = εlc, relating absorbance to concentration, where ε is the molar extinction coefficient, l is path length, and c is concentration.

Energetic profile diagram

Graphical representation of energy changes in a reaction, showing reactants, products, and Ea.

Transition State Theory (TST)

Theory describing the formation of the transition state and its role in determining rate constants.

Catalysis applications

Using catalysts to optimize industrial processes, energy efficiency, and environmental impacts.