AP* Chemistry: Chemical Kinetics*

Introduction to Chemical Kinetics

• Chemical Kinetics: Study of the speed/rate of a reaction under various conditions.
• Spontaneity vs. Speed: A spontaneous reaction does not imply a rapid reaction (e.g., diamond to graphite).
• Mechanism: Sequence of events controlling speed and outcome of the reaction.

Factors Affecting Reaction Rates

1. Nature of Reactants:
   • Physical state affects reaction speed (e.g., gasoline in different states).
   • Chemical identity: Usually ions of opposite charge react rapidly.
   • Bond strength: Stronger bonds (higher bond energies) slow down reactions.
2. Concentration of Reactants:
   • More molecules lead to more collisions.
3. Temperature:
   • Increased temperature increases successful collisions, thus increasing reaction rate.
   • Rule of thumb: 10°C increase doubles the reaction rate.
4. Catalysts:
   • Catalysts accelerate reactions without being consumed.
   • Lower activation energy, thus increasing reaction rate.
   • Example: MnO2 speeds up decomposition of H2O2.
5. Surface Area:
   • More exposed surface area increases reaction rate (e.g., coal dust vs. charcoal).
6. Inert Gas:
   • Adding an inert gas has no effect on reaction rate.

Collision Theory

• Particles Must Collide: Only two particles can collide at once.
• Proper Orientation: Colliding molecules must be properly oriented.
• Activation Energy: Collision energy must overcome electron repulsion and transform energy to allow bond formation.

Chemical Reaction Rates

• Rate Expression: Rate = change in concentration of a species/time interval.
• Instantaneous Rate: Slope of a tangent line on a concentration-time graph.

Relative Rates

• Focus on the disappearance of reactants or the appearance of products.
• Respect stoichiometry and algebraic signs in rate calculations.

Differential Rate Law

• Reversible Reactions: Forward and backward reactions affect equilibrium.
• Rate Law: Relation between reaction rate and concentrations of reactants.
• Rate Constant (k): Temperature-dependent, determined experimentally.

Order of Reaction

• Zero-Order: Concentration change doesn't affect rate.
• First-Order: Rate proportional to reactant concentration.
• Second-Order: Rate quadruples when reactant concentration doubles.
• Fractional Orders: Rare but possible.

Determining Reaction Orders

• Use experimental data to determine exponents in rate laws for reactants.
• Example: Adjust concentrations and observe rate changes to deduce order.

Integrated Rate Laws

• Relation between concentration and time for reactions.
• Use graphical methods to determine reaction order.

Reaction Mechanisms

• Elementary Steps: Steps in a reaction mechanism showing molecularity.
• Rate Expressions: Derived from elementary steps, not overall stoichiometry.
• Rate-Determining Step: Slowest step controlling overall reaction speed.

Catalysis

• Catalysts: Lower activation energy, provide alternative pathways.
• Types:
  • Heterogeneous: Different phase than reactants.
  • Homogeneous: Same phase as reactants.
• Applications: Catalytic converters, hydrogenation of oils, etc.

Exercises and Examples

• Various exercises to practice calculating reaction rates, rate laws, and understanding mechanisms based on given data and scenarios.