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Credit Sprint's Guide to Chemistry Unit 5: Kinetics

AP® weight 7–9%CLEP® weight 3–5%

Kinetics is about how fast reactions go. Collision theory, activation energy, rate laws and mechanisms explain what controls speed, and the integrated rate laws let you predict concentrations over time.

Reaction energy diagramReaction progressEnergy reactantsproductstransition state Ea ΔH

What you need to know

  • Rate = k[A]ᵐ[B]ⁿ; orders are found from experiments (initial rates), never from the balanced equation.
  • The rate constant k depends on temperature and catalyst, not concentration.
  • Units of k: zero order M/s, first order s⁻¹, second order M⁻¹s⁻¹.
  • Integrated laws: zero [A] = [A]₀ − kt; first ln[A] = ln[A]₀ − kt; second 1/[A] = 1/[A]₀ + kt. The linear plot identifies the order.
  • First-order half-life t½ = 0.693/k is constant. Zero-order half-life shrinks as [A] falls; second-order half-life grows.
  • Collision theory: collisions need enough energy (Ea) and the right orientation. Higher temperature, higher concentration and more surface area increase the rate.
  • A catalyst lowers Ea by offering a new pathway; it does not change ΔH or K.
  • Mechanisms: the slowest elementary step sets the rate. Intermediates are made and used; catalysts are used and regenerated.
  • Reaction coordinate diagrams show Ea, ΔH, the transition state and intermediates (valleys between peaks).

Equations and rules

rate = k[A]ᵐ[B]ⁿ
ln[A] = ln[A]₀ − kt
t½ = 0.693/k (first order)
ln k = −Ea/RT + ln A

Exam traps

Watch forOrders come from data, not coefficients (except for elementary steps).
Watch forA catalyst changes speed, not the equilibrium position or ΔH.
Watch forIntermediates cancel out of the overall equation; catalysts appear on both sides of steps.
Watch forln[A] vs t is linear only for first-order reactions.

Key terms

  • Rate law: An equation relating rate to reactant concentrations: rate = k[A]ᵐ[B]ⁿ.
  • Reaction order: The exponent of a concentration in the rate law, found from experiments.
  • First-order half-life: t½ = 0.693/k; constant, independent of concentration.
  • Activation energy: The minimum energy that colliding particles need to react.
  • Catalyst: A substance that speeds a reaction by lowering Ea without being consumed.
  • Intermediate: A species made in one step of a mechanism and used up in a later step.
  • Rate-determining step: The slowest elementary step; it sets the rate law.
  • Collision theory: Reactions need collisions with enough energy and the correct orientation.
  • Arrhenius equation: k = Ae^(−Ea/RT); k increases with temperature.

Practice questions

Five of the 50 Unit 5 questions in Credit Sprint. Every choice is explained in the app.

1. When [A] is doubled from 0.10 M to 0.20 M, the initial rate of a reaction increases from 2.0 × 10⁻³ M/s to 8.0 × 10⁻³ M/s. What is the order with respect to A?
  1. Second order
  2. First order
  3. Zero order
  4. Third order
Show answer

Second order. Doubling [A] quadruples the rate: rate ∝ [A]².

2. A first-order reaction has k = 0.0231 s⁻¹. What is its half-life?
  1. 15.0 s
  2. 30.0 s
  3. 0.0231 s
  4. 43.3 s
Show answer

30.0 s. t½ = 0.693 ÷ 0.0231 s⁻¹ = 30.0 s.

3. For a zero-order reaction, the half-life
  1. increases as the reaction proceeds
  2. is constant
  3. decreases as the reaction proceeds because the concentration decreases
  4. is independent of k
Show answer

decreases as the reaction proceeds because the concentration decreases. t½ = [A]₀ ÷ 2k, so it depends on the initial amount.

4. On a reaction energy diagram, the activation energy of the forward reaction is represented by
  1. the energy of the products
  2. the difference between reactants and products
  3. the difference between the reactant energy and the top of the barrier
  4. the difference between products and the top of the barrier
Show answer

the difference between the reactant energy and the top of the barrier. Ea is the energy needed to reach the transition state.

5. In the previous mechanism, what is the overall reaction?
  1. NO₂ → NO + CO₂
  2. NO₂ + CO → NO + CO₂
  3. 2 NO₂ + CO → NO₃ + CO₂
  4. NO₃ + NO → NO₂ + NO₂
Show answer

NO₂ + CO → NO + CO₂. NO₃ cancels as an intermediate and one NO₂ cancels as it is regenerated.