AP Chemistry: Factors Affecting Reaction Rate (Concentration, Temperature, Catalyst)

AP Chemistry – Factors Affecting Reaction Rate (Concentration, Temperature, Catalyst)

AP Chemistry: Factors Affecting Reaction Rate (Concentration, Temperature, Catalyst) – Exam-Ready Study Guide

What This Is

Reaction rate measures how fast reactants turn into products. The AP exam tests how concentration, temperature, and catalysts change this speed. These factors are crucial for designing industrial processes (e.g., Haber-Bosch for ammonia synthesis) and understanding everyday reactions (e.g., why food spoils faster in heat). Mastering this topic helps you predict reaction behavior, optimize lab experiments, and ace kinetics questions on the exam.

Key Terms & Concepts

• Reaction Rate: How fast reactants disappear or products form, measured in M/s (molarity per second).

• Example: If [A] decreases from 0.5 M to 0.3 M in 10 s, rate = (0.5 – 0.3)/10 = 0.02 M/s.

• Collision Theory: Reactions occur when particles collide with enough energy (activation energy, Eₐ) and proper orientation.

• Key idea: More collisions = faster rate.

• Concentration Effect: Increasing reactant concentration (or pressure for gases) increases rate because more particles collide per second.

• Exception: Zero-order reactions (rate doesn’t depend on concentration).

• Rate Law: Rate = k[A]ᵐ[B]ⁿ, where:

• k = rate constant (temperature-dependent),
• [A], [B] = reactant concentrations,

• m, n = reaction orders (determined experimentally, not from coefficients).

• Temperature Effect: Higher temperature = faster rate because:

• Particles move faster → more collisions.

• More particles exceed Eₐ (see Maxwell-Boltzmann distribution).

• Arrhenius Equation: k = Ae^(-Eₐ/RT), where:

• k = rate constant,
• A = frequency factor (collision frequency),
• Eₐ = activation energy (J/mol),
• R = 8.314 J/mol·K,
• T = temperature (K).

• AP shortcut: If T increases, k increases exponentially.

• Catalyst: A substance that lowers Eₐ (by providing an alternate pathway) without being consumed.

• Example: Enzymes (e.g., catalase breaks down H₂O₂) or catalytic converters in cars (convert CO to CO₂).

• Homogeneous vs. Heterogeneous Catalysts:

• Homogeneous: Same phase as reactants (e.g., H⁺ in esterification).

• Heterogeneous: Different phase (e.g., solid Pt in H₂ + O₂ → H₂O).

• Reaction Mechanism: A step-by-step sequence of elementary reactions. The slowest step (rate-determining step) controls the rate law.

• Example: For 2 NO + O₂ → 2 NO₂, the mechanism might be:

  
  
  1. NO + NO → N₂O₂ (fast),
  2. N₂O₂ + O₂ → 2 NO₂ (slow, rate-determining).
     → Rate law = k[NO]²[O₂].

• Activation Energy (Eₐ): Minimum energy needed for a collision to result in a reaction.

• Visual: The "hill" in a reaction energy diagram.

• Maxwell-Boltzmann Distribution: Shows how particle energies are distributed at a given temperature. Higher T shifts the curve right, increasing the fraction of particles with E ≥ Eₐ.

Step-by-Step: Solving Rate Problems

1. Determine the Rate Law from Data
2. Compare experiments where one reactant’s concentration changes while others stay constant.

3. Example: If doubling [A] quadruples the rate, the order for A is 2 (rate ∝ [A]²).

4. Calculate the Rate Constant (k)

5. Plug known values (rate, concentrations, orders) into the rate law and solve for k.

6. Units of k: Depend on overall order (e.g., 1st order: s⁻¹; 2nd order: M⁻¹s⁻¹).

7. Predict Rate Changes

8. Concentration: If [A] triples in a 1st-order reaction, rate triples.

9. Temperature: Use the Arrhenius equation or the rule of thumb: Rate doubles for every 10°C rise (approximate).

10. Analyze Catalysts

11. Identify if a catalyst is present (lowers Eₐ in the energy diagram).

12. AP trick: Catalysts do not change ΔH or equilibrium position.

13. Sketch Energy Diagrams

14. Label Eₐ (with and without catalyst), ΔH, and transition state.
15. Example: Exothermic reaction: products lower than reactants.

Common Mistakes

• Mistake: Assuming the rate law matches the balanced equation coefficients.

• Correction: Rate laws are experimentally determined. Coefficients only match for elementary steps (rare on the AP exam).

• Mistake: Forgetting that catalysts do not change equilibrium or ΔH.

• Correction: Catalysts speed up both forward and reverse reactions equally, so equilibrium is unchanged.

• Mistake: Confusing rate (M/s) with rate constant (k).

• Correction: Rate depends on concentration; k is constant at a given temperature.

• Mistake: Misapplying the Arrhenius equation (e.g., using °C instead of K).

• Correction: Always convert temperature to Kelvin (K = °C + 273).

• Mistake: Ignoring units for k.

• Correction: Units of k must cancel to give M/s for the rate. Example:
  • 1st order: s⁻¹,
  • 2nd order: M⁻¹s⁻¹.

AP Exam Insights

1. FRQ Hot Topics:
2. Experimental design: Propose how to determine the rate law (e.g., "Describe an experiment to find the order with respect to A").
3. Energy diagrams: Sketch and label Eₐ, ΔH, and the effect of a catalyst.

4. Arrhenius calculations: Given Eₐ and k at one temperature, find k at another temperature.

5. Multiple-Choice Traps:

6. Catalysts vs. intermediates: Catalysts are consumed then regenerated; intermediates are produced then consumed.
7. Temperature vs. concentration: Both increase rate, but temperature also changes k (concentration does not).

8. Zero-order reactions: Rate is constant (e.g., enzyme-catalyzed reactions at high substrate concentration).

9. Tricky Distinctions:

10. Rate vs. rate constant: Rate changes with concentration; k changes with temperature.
11. Homogeneous vs. heterogeneous catalysts: Phase matters for mechanism (e.g., surface adsorption in heterogeneous catalysts).

Quick Check Questions

1. Multiple Choice:
   For the reaction 2 NO + O₂ → 2 NO₂, the rate law is rate = k[NO]²[O₂]. If [NO] is doubled and [O₂] is halved, the rate will:
   a) Stay the same
   b) Double
   c) Quadruple
   d) Increase by a factor of 8
   Answer: b) Double. (Rate = k[2NO]²[0.5O₂] = k·4[NO]²·0.5[O₂] = 2·k[NO]²[O₂].)

2. Short FRQ:
   The decomposition of H₂O₂ (2 H₂O₂ → 2 H₂O + O₂) is catalyzed by MnO₂. Sketch a reaction energy diagram for the uncatalyzed and catalyzed reactions. Label Eₐ and ΔH.
   Answer: Two curves (uncatalyzed higher Eₐ, catalyzed lower Eₐ); same ΔH (catalysts don’t change ΔH).

3. Multiple Choice:
   Which of the following will not increase the rate of a reaction?
   a) Adding a catalyst
   b) Increasing temperature
   c) Increasing the concentration of a reactant in a zero-order reaction
   d) Increasing the surface area of a solid reactant
   Answer: c) Increasing concentration in a zero-order reaction. (Rate = k, independent of concentration.)

Last-Minute Cram Sheet

1. Rate = Δ[ ]/Δt (M/s); rate law = k[A]ᵐ[B]ⁿ (orders from data, not coefficients).
2. ⚠️ Zero-order: Rate = k (concentration doesn’t matter).
3. Temperature ↑ → rate ↑ (more collisions + more particles exceed Eₐ).
4. Arrhenius equation: k = Ae^(-Eₐ/RT) (use Kelvin!).
5. Catalysts lower Eₐ but don’t change ΔH or equilibrium.
6. Homogeneous catalyst: Same phase (e.g., H⁺ in solution).
7. Heterogeneous catalyst: Different phase (e.g., solid Pt in gas reaction).
8. Rate-determining step = slowest step in mechanism → determines rate law.
9. Maxwell-Boltzmann: Higher T shifts curve right → more particles with E ≥ Eₐ.
10. ⚠️ Units of k: 1st order = s⁻¹; 2nd order = M⁻¹s⁻¹.