AP Chemistry Unit 5 Notes

Overview of Kinetics

• Kinetics studies the rates of chemical reactions.

• Goals:

  • Determine reaction rates.

  • Understand factors affecting these rates.

Factors Affecting Reaction Rates

1. Physical State of Reactants

   • Reactants in the same physical state are more likely to react.

   • Different phases limit interaction to surface contact.

2. Concentration of Reactants

   • Higher concentrations increase reaction frequency.

3. Temperature

   • Higher temperatures enhance kinetic energy, increasing frequency and energy of collisions.

4. Presence of a Catalyst

   • Catalysts provide a pathway for effective collisions but inert gases do not impact rates or equilibrium.

Collision Theory

• For a reaction to occur:

  • Energy: Collisions must have enough energy to overcome attractive forces.

  • Orientation: Collisions must be properly oriented.

• Only a small percentage of collisions lead to successful reactions.

• Increasing concentration, surface area, or temperature increases collision frequency.

Reaction Rates

• Reaction rate is the change in concentration over time:
  $ext{Rate} = rac{ riangle [ ext{reactants}]}{ riangle t}$ or $ext{Rate} = rac{ riangle [ ext{products}]}{ riangle t}$

• Delta ($ riangle$) indicates the difference between initial and final concentration or time.

• Rates:

  • Products: Positive (+)

  • Reactants: Negative (−)

Graphical Representation of Rates

• Graphing concentration vs time shows how rates change.

• The slope of the tangent line at any point represents the reaction rate.

• Rates decrease over time due to reduced reactant concentration.

Stoichiometry and Reaction Rates

• Stoichiometric ratios define the relationship between product formation and reactant consumption.

  • Example: For the reaction $A <br>ightarrow B$, the rate of disappearance of A matches the rate of appearance of B (1:1 ratio).

• For $C <br>ightarrow 2D$, the rate of disappearance of C is half the rate of appearance of D (1:2 ratio).

Differential Rate Law

• Relationship between rate and concentration: $ext{Rate} = k[ ext{reactant}<em>1]^m[ ext{reactant}</em>2]^n$ where:

  • $k$ is the rate constant,

  • $m$ and $n$ indicate the order of reaction with respect to each reactant.

Orders of Reaction

• 0th order: changes in concentration have no effect on the rate.

• 1st order: changes in concentration produce proportional changes in rate.

• 2nd order: changes in concentration produce squared changes in rate.

Determining Orders of Reaction

• To find orders, carry out rate trials while holding one reactant constant.

• Use isolated trial data to discern the orders systematically.

Integrated Rate Law

• For zero, first, and second order, the relationships between concentration and time are as follows:

  1. Zero order:
     $[ ext{A}] - [ ext{A}_0] = -kt$

  2. First order:
     $ext{ln}[ ext{A}] - ext{ln}[ ext{A}_0] = -kt$

  3. Second order:
     $rac{1}{[ ext{A}]} - rac{1}{[ ext{A}_0]} = kt$

• Graphs of these relations determine reaction orders based on linearity.

Mechanisms in Kinetics

• Mechanisms are sequences of elementary steps that describe how reactants transform into products.

• To validate a mechanism:

  1. The sum of elementary steps must equal the overall reaction equation.

  2. The mechanism must align with the experimentally determined rate law.

• Rate Determining Step: The slowest step dictates the reaction rate.

Catalysts

• Catalysts enhance reaction rates by lowering activation energy without being consumed.

• Provide pathways for more effective collisions.

• Types:

  1. Homogenous: Same phase as reactants.

  2. Heterogeneous: Different phase from reactants (e.g., platinum in catalytic converters).

  3. Biological (Enzymes): Increase biological reaction rates.

  4. Acid/Base Catalysts: Alter reaction pathways through protonation/deprotonation reactions.

Example of catalyst role in ozone layer depletion:
$ext{CCl}<em>2 ext{F}</em>2(g) <br>ightarrow ext{CClF}<em>2(g) + ext{Cl}(g)$ $ext{Cl}(g) + ext{O}</em>3(g) <br>ightarrow ext{ClO}(g) + ext{O}<em>2(g)$ $ext{O}</em>3(g) + ext{ClO}(g) <br>ightarrow ext{Cl}(g) + 2 ext{O}_2(g)$

• Here, Cl is a catalyst, not consumed in the overall reaction but participates in intermediates.