collision theory

Chemical Kinetics

The study of reaction rates, factors affecting rates, and the step-by-step molecular mechanisms of reactions.

Reaction Rate

The change in concentration of a reactant or product per unit time (Δ[C] / Δt).

Units for Reaction Rate

M/s or mol/(L·s).

Experimental Methods to Track Reaction Rate

Spectrophotometry (color intensity), gas pressure/volume tracking, pH change, electrical conductivity, or mass changes.

Average Rate vs. Instantaneous Rate vs. Initial Rate

Average: over a time interval; Instantaneous: at a specific moment (slope of tangent line); Initial: instantaneous rate at t = 0.

Stoichiometric Rate Relationship

For aA + bB -> cC + dD: Rate = -(1/a)(Δ[A]/Δt) = -(1/b)(Δ[B]/Δt) = +(1/c)(Δ[C]/Δt) = +(1/d)(Δ[D]/Δt)

Collision Theory

Reactions occur when particles collide with kinetic energy >= Ea and correct spatial orientation.

Activation Energy (Ea)

The minimum kinetic energy colliding molecules must possess to react and reach the transition state.

Activated Complex (Transition State)

An unstable, high-energy atomic arrangement formed at the peak of the potential energy barrier.

Five Primary Factors Affecting Reaction Rates

1. Reactant nature, 2. Concentration / Pressure, 3. Surface area, 4. Temperature, 5. Catalyst addition

Maxwell-Boltzmann Energy Distribution and Temperature

Increasing T shifts the energy curve right and flattens it, increasing the fraction of particles with E >= Ea.

Rule of Thumb for Temperature and Rate

For many reactions near room temperature, a 10°C rise doubles the reaction rate.

Differential Rate Law Formula

Rate = k[A]^m [B]^n, where k is the rate constant and m, n are reaction orders.

Method of Initial Rates

Comparing initial rates across experiments where reactant concentrations are varied one at a time to isolate orders m and n.

General Formula for Rate Constant (k) Units

Units of k = M^(1 - n) · t^-1 (where n is the overall reaction order).

Units of k by Reaction Order

0th Order: M·s^-1 | 1st Order: s^-1 | 2nd Order: M^-1·s^-1

Reaction Profile: ΔH (Enthalpy Change)

ΔH = E_products - E_reactants. Negative for exothermic, positive for endothermic.

Activation Energy of Forward vs. Reverse Reactions

Ea,reverse = Ea,forward - ΔH

Elementary Reaction Step

A single molecular event step in a mechanism whose rate law exponents match its stoichiometric coefficients directly.

Molecularity of Elementary Steps

Unimolecular (1 particle), Bimolecular (2 particles), Termolecular (3 particles - rare).

Reaction Intermediate vs. Catalyst

Intermediate: Produced then consumed (absent at start and end). Catalyst: Consumed then regenerated (present at start and end).

Rate-Determining Step (RDS)

The slowest elementary step in a reaction mechanism; dictates the overall reaction rate and rate law.

Mechanisms with Fast Initial Equilibrium Step

Set forward rate = reverse rate of the fast step to solve for an intermediate in terms of reactants, then substitute into the RDS rate law.

Homogeneous vs. Heterogeneous Catalysts

Homogeneous: Same phase as reactants. Heterogeneous: Different phase (usually solid surface binding gas/liquid reactants).

Enzyme Catalysis Key Terms

Enzyme (biological catalyst), Substrate (reactant molecule), Active Site (specific binding pocket forming the enzyme-substrate complex).

How Catalysts Change Reaction Diagrams

Lower peak height (Ea), often split into multiple lower activation peaks; ΔH and equilibrium position remain unchanged.

Inhibitors

Substances that bind to catalysts or enzymes to decrease or block their catalytic activity.

Explain why the rates have different signs.

Negative sign: Oxygen is a reactant, so its concentration decreases over time, making it negative.

Positive sign for: Nitrogen dioxide is a product, so its concentration increases over time,