Unit 5.2 Rate Laws and Reaction Rates

  • Chemical Reactions and Equilibrium

    • Chemical reactions are reversible.
    • Products can convert back into reactants.
    • Equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction.

  • Focus on Rates Before Equilibrium

    • Study reaction rates just after reactants are mixed (initial rate method).
    • At time zero, only reactants are present (no products).

  • Understanding Rate Law

    • A rate law relates the rate of a reaction to the concentration of reactants.
    • General form: \text{Rate} = k [A]^m [B]^n
      • k: rate constant
      • [A], [B]: concentrations of reactants
      • m, n: rate orders (exponents that need to be determined).
    • Note: Products are not included in the rate law.

  • Types of Rate Laws

    • Differential Rate Law: Used to express how the rate of reaction depends on the concentrations of reactants. (Common in AP questions)
    • Integrated Rate Law: This will be discussed in later lessons.

  • Initial Rate Method

    • To determine the order of reactants based on experimental data.

  • Example of Determining Rate Law

    • Given a balanced equation and experimental data:
    • Compare two experiments (e.g., Experiments 2 and 3) and focus on one reactant (e.g., ammonium).
    • Set up equations based on initial rates and concentrations from those experiments.
    • Use the ratio of rates in the format:
      \frac{ \text{Rate}3 }{ \text{Rate}2 } = \frac{ k [A]3^x [B]3^y }{ k [A]2^x [B]2^y }
    • Cancel out the constant k.
    • Solve for exponents.

  • Finding Exponents

    • For ammonium:
    • After simplifying, when comparing experiment rates, find that x = 1 (first order).
    • For nitrite:
    • Again, compare the relevant experimental rates where nitrite varies.
    • Solve to find y = 2 (second order).

  • Overall Reaction Order

    • Sum of exponents gives overall reaction order.
    • Example: 1 + 2 = 3 (third order).
    • The higher the order, the faster reactants are consumed.

  • Determining Rate Constant (k)

    • Use any set of experimental data and the established rate law.
    • Example: For nitrogen monoxide (NO) squared and chlorine (Cl) to the first power.
    • The units for the rate constant vary based on the overall order of reaction.

  • Units for Rate Constant (k)

    • Depends on the reaction order:
    • Overall reaction order of n means k has units of \text{M}^{- (n-1)} \text{s}^{-1}
    • Practice applies units when deriving k from the rate law.

  • Final Review Activity

    • Conduct further practice on finding rate laws and overall reaction orders using new reaction data provided.

  • Conclusion

    • Review concepts of chemical kinetics and the calculation of rate laws in preparation for exams.
    • Emphasis on understanding rather than memorization to enhance problem-solving skills.