Chemical Kinetics – Core Vocabulary

A comprehensive set of vocabulary flashcards covering definitions and key terms from the lecture notes on chemical kinetics, including reaction types, rate concepts, factors influencing rate, kinetic parameters, and theoretical models.

Chemical kinetics

Branch of chemistry that studies reaction rates and the factors affecting them.

Kinesis (origin)

Greek word meaning movement; root of the term kinetics.

Instantaneous reaction

A very fast reaction that occurs immediately upon contact of reactants; rate cannot be measured.

Very slow reaction

A reaction that proceeds so slowly (e.g., rusting of iron) that rate is difficult to observe in short time frames.

Moderately slow reaction

Reaction whose speed is measurable under laboratory conditions (e.g., decomposition of H₂O₂).

Rate of reaction

Change in concentration of reactants or products per unit time.

Average rate

Overall change in concentration over a long time interval divided by that time.

Instantaneous rate

Reaction rate at a specific moment; equals the slope of concentration-time curve at that point.

Rate unit (solution)

Commonly mol L⁻¹ s⁻¹ (or mol L⁻¹ min⁻¹) for reactions in solution.

Rate unit (gas)

Often expressed in bar s⁻¹, atm s⁻¹, or corresponding per-minute units for gaseous reactions.

Concentration effect

Increasing reactant concentration raises rate by providing more molecules for effective collisions.

Temperature effect

Raising temperature generally doubles rate for each 10 °C increase.

Nature of reactants

Intrinsic bond strengths and structures govern how easily reactants transform, affecting rate.

Catalyst

Substance that increases reaction rate without being consumed, typically by lowering activation energy.

Pressure effect

For gases, higher pressure (higher concentration) accelerates reaction rate.

Surface area effect

Smaller particle size or powdered form increases surface area, enhancing reaction rate.

Radiation effect

Photons can initiate or speed photochemical reactions (e.g., H₂ + Cl₂ → 2 HCl under light).

Rate law / rate expression

Experimentally determined equation relating rate to molar concentrations of reactants, each raised to a power.

Rate constant (k)

Proportionality factor in the rate law; equals reaction rate when each reactant concentration is unity.

Order of reaction

Sum of the powers of concentration terms in the rate law; may be 0, 1, 2, etc., or fractional.

Zero-order reaction

Rate is independent of reactant concentrations; common in catalysed surface or photochemical processes.

First-order reaction

Rate is directly proportional to the concentration of one reactant.

Second-order reaction

Rate depends on either the square of one reactant’s concentration or the product of two concentrations.

Third-order reaction

Rate proportional to the product of three concentration terms (e.g., 2 NO + O₂ → 2 NO₂).

Fractional order reaction

Reaction whose overall order is a non-integer (e.g., 3⁄2 for H₂ + Br₂ → 2 HBr).

Elementary reaction

Reaction that occurs in a single molecular step.

Complex reaction

Overall process comprising two or more elementary steps.

Molecularity

Number of species that collide simultaneously in an elementary step; always a positive integer.

Integrated rate equation

Mathematical expression (after integration) that links concentration with time and rate constant.

Half-life (t₁⁄₂)

Time required for reactant concentration to decrease to half its initial value.

Pseudo first-order reaction

Multireactant process that behaves as first order because one reactant is in large excess (e.g., ester hydrolysis in water).

Arrhenius equation

k = A e^(−Ea/RT); relates rate constant to temperature and activation energy.

Activation energy (Ea)

Minimum energy that colliding molecules must possess to form activated complex and react.

Threshold energy

Total minimum energy (reactants’ energy + Ea) required for an effective collision.

Maxwell–Boltzmann distribution

Curve showing fraction of molecules versus kinetic energy at a given temperature.

Collision frequency (Z)

Number of molecular collisions per second per unit volume.

Effective collision

Collision with proper orientation and sufficient energy to form products.

Probability (steric) factor (P)

Term in collision theory accounting for orientation requirements in formation of products.

Promoter (catalysis)

Substance that enhances catalyst activity.

Poison (catalysis)

Substance that diminishes or destroys catalyst activity.

Inhibitor

Substance that decreases reaction rate without necessarily deactivating the catalyst.