Chemical Kinetics – Collision Theory, Activation Energy & Catalysis

A set of vocabulary flashcards summarizing key terms from the lecture on collision theory, activation energy, temperature and catalyst effects on reaction rates.

Collision Theory

Model stating that reactant particles must collide to react; rate depends on collision frequency and effectiveness.

Effective Collision

A collision with correct molecular orientation and sufficient energy (≥ Ea) that leads to product formation.

Molecular Orientation

The spatial arrangement of colliding molecules that determines whether a collision is effective.

Activation Energy (Ea)

Minimum energy that reacting particles must possess to reach the transition state and form products.

Transition State / Activated Complex

Unstable, high-energy species with partial bonds formed at the peak of the potential-energy diagram.

Potential Energy Diagram

Graph plotting potential energy versus reaction progress, showing reactants, products, Ea and transition state.

Rate Constant (k)

Proportionality factor in the rate law whose value depends on temperature and catalyst but not on reactant concentration.

Arrhenius Equation

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

Frequency Factor (A)

Pre-exponential term in the Arrhenius equation representing collision frequency and orientation probability.

Concentration Effect

Increasing reactant concentration raises collision frequency, giving a higher reaction rate.

Product of Concentrations

Number of possible collisions is proportional to the product (not the sum) of reactant particle numbers.

Collision Frequency

Number of collisions per unit time; increases with higher concentration or temperature.

Kinetic Energy

Energy of motion of particles; rises with temperature, increasing the fraction exceeding Ea.

Temperature Effect on Rate

Higher temperature raises kinetic energy, enlarges the high-energy fraction, increases k and speeds up the reaction.

Maxwell-Boltzmann Distribution Curve

Plot showing distribution of molecular kinetic energies; area beyond Ea represents effective collisions.

Alternative Reaction Pathway

Lower-energy route provided by a catalyst, reducing Ea and increasing reaction rate.

Catalyst

Substance that increases reaction rate by lowering Ea without being consumed in the reaction.

Homogeneous Catalyst

Catalyst present in the same physical phase as the reactants (e.g., all aqueous).

Heterogeneous Catalyst

Catalyst in a different phase than the reactants, often a solid surface interacting with gaseous or liquid reactants.

Rate Law

Mathematical expression linking reaction rate to reactant concentrations, each raised to an experimentally determined power.

Threshold Energy

Energy level equal to Ea that must be reached for a collision to be effective.

Reaction Mechanism

Step-by-step sequence of elementary reactions describing the overall chemical change.

Influence of Catalyst on k

Lower Ea raises the exponential term in Arrhenius equation, giving a larger rate constant k.

Influence of Temperature on k

As T increases, the exponential factor e^(−Ea/RT) becomes larger, increasing k and reaction rate.