Kinetics
Reaction Rate
How fast reactants become products — controlled by concentration, temperature, catalyst
Reaction rate is the change in concentration of reactant or product per unit time. Rate = -Δ[A]/Δt for reactants; +Δ[P]/Δt for products. Rate law: rate = k[A]^m[B]^n where k is rate constant, m and n are reaction orders (determined experimentally). Affected by: concentration, temperature (Arrhenius equation), catalyst, surface area (heterogeneous), light. Different from equilibrium — equilibrium is thermodynamic; rate is kinetic. Catalysts increase rate without changing equilibrium.
- RateΔ[concentration]/Δt
- Rate lawrate = k[A]^m[B]^n
- Reaction orderSum of exponents (m + n)
- Rate constant kTemperature-dependent (Arrhenius)
- Activation energyMinimum energy for reaction
- CatalysisIncreases rate; doesn't change equilibrium
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Why rate matters
- Industrial. Optimize productivity.
- Drug delivery. Controlled release.
- Food preservation. Slow degradation.
- Catalysis. Accelerate reactions.
- Atmospheric chemistry. Ozone destruction, smog.
- Biology. Enzyme kinetics.
- Materials. Polymerization, corrosion.
Common misconceptions
- Rate = stoichiometric coefficients. Order from experiment.
- Rate constant is constant. Depends on T.
- All reactions follow simple laws. Many complex (multistep, autocatalytic).
- Catalyst changes equilibrium. No — only rate.
- Faster reactions are more thermodynamically favored. Different concept.
- Rate doesn't depend on temperature much. Strong dependence (exponential).
Frequently asked questions
What is reaction rate?
Speed of reaction. Rate = change in concentration over time. 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. Stoichiometry adjusts so all give same rate. Units: M/s (concentration per time).
What's a rate law?
Rate = k[A]^m[B]^n. k = rate constant (T-dependent). m, n = reaction orders for A and B (determined experimentally, not from stoichiometry). Total order = m + n. Most common: 1st order (rate ∝ concentration) or 2nd order. Zero order: rate = k (independent of concentration).
How is rate law determined?
Method of initial rates. Run reaction with different starting concentrations; measure initial rates. Compare ratios. If doubling [A] doubles rate: 1st order in A. If doubling quadruples rate: 2nd order. Calculate exponents from rate ratios. Modern: integrated rate laws and computer fitting.
How does temperature affect rate?
Strongly. Arrhenius equation: k = A × exp(-Ea/RT). Higher T → higher k → faster rate. Rule of thumb: 10°C increase doubles rate (Ea ≈ 50 kJ/mol). Smaller activation energy → less T sensitivity. Used in: kinetics studies, food preservation (refrigeration), chemical processing.
What's activation energy?
Minimum energy required for reaction to proceed. Reactant bonds must break partially before products form. Higher Ea → harder reaction; slower at given T. Reactants need kinetic energy ≥ Ea. Catalysts lower Ea — alternative pathway. Diagram: reaction profile shows energy peak (transition state) at Ea above reactants.
How does a catalyst work?
Provides alternative reaction pathway with lower activation energy. Catalyst not consumed. Speeds up both forward and reverse reactions equally — doesn't change K. Examples: enzymes (proteins, very specific), metal catalysts (Pt for hydrogenation), acid catalysts (H₂SO₄ in esterification). Industrial: Haber process, catalytic converters.
What about heterogeneous reactions?
Reactants in different phases (solid/liquid, solid/gas). Rate depends on surface area. Examples: combustion (O₂ + fuel surface), catalytic converters (gas + Pt surface). Larger surface = faster rate (catalysts often porous). Diffusion can be rate-limiting in heterogeneous systems.