General Chemistry II - Chemical Kinetics

CHEM 115.3: General Chemistry II Study Notes

Instructor Information

  • Instructor: Dr. Pearson W. K. Ahiahonu

  • Email: pearson.ahiahonu@usask.ca

  • Textbook: Chemistry, 2020, 6th Edition, by Gilbert et al.

Chapter 13: Chemical Kinetics

Definition of Chemical Kinetics

  • Chemical kinetics refers to the study of the rate of change of concentrations of substances involved in chemical reactions.

Reaction Rate

  • Reaction rate is defined as how rapidly a reaction occurs.

  • This rate is related to the rates of change in concentration of reactants and products over time.

Factors Affecting Reaction Rates

  • The reaction rate can be influenced by several factors:

    • Physical state of reactants

    • Concentration of reactants

    • Temperature

    • Presence of a catalyst

Reaction Rate and Stoichiometry

  • The relative rates of consumption of reactants and formation of products are based on the stoichiometry of the reaction.

  • For a general reaction of the form: aA + bB \rightarrow cC + dD

    • Where A and B are reactants while C and D are products, with a, b, c, and d being the stoichiometric coefficients.

Average Rate

  • Average rate is defined as the change in concentration of a reactant or product over a specified time interval.

Example Reaction

  • For the reaction: N2(g) + O2(g) \rightarrow 2 \text{NO}(g)

    • The reaction rate at a given instant can be graphically determined as the tangential slope of a concentration versus time plot.

Instantaneous Reaction Rates

  • To determine the instantaneous rate for a reaction at a certain time, use the following:

Sample Calculation

  • Given:
    \Delta N \Delta O \Delta NO\text{ Rate} = = = \frac{\Delta}{2 \Delta \Delta}

  • Time (t) is utilized in the calculation.

Worked Example 1

  • Balanced Chemical Equation:
    \text{H}2\text{O}2(aq) + 3 \text{I}^-(aq) + 2 \text{H}^+(aq) \rightarrow \text{I}3^-(aq) + 2 \text{H}2\text{O}(l)

  • For the first 10.0 seconds of the reaction:

    • The concentration of I⁻ dropped from 1.000 M to 0.868 M.

    1. (a) Calculate the average rate of reaction in this time interval.

    2. (b) Predict the rate of change in the concentration of H⁺ (i.e., \frac{\Delta[H^+]}{\Delta t}) during this time interval.

    3. (c) Predict the rate of formation of I₃⁻ (\frac{\Delta[I_3^-]}{\Delta t}) during this time interval.

Result

  • \text{Rate} = + \frac{\Delta[I_3^-]}{\Delta t} = 4.40 \times 10^{-3} \text{M/s}

The Rate Law: The Effect of Reactant Concentration on Reaction Rate

  • Rate Law: A mathematical expression that defines the experimentally determined relationship between reactant concentrations and the rate of the reaction.

  • The rate law of a reaction expresses how the rate of the reaction correlates to the concentration of reactants.

Definitions

  • Reaction Order: An experimentally determined number that indicates the dependence of the reaction rate on the concentration of a reactant.

    • Overall Order of Reaction: The sum of all exponents of concentration terms in the rate law.

  • Rate Constant (k): Proportionality constant that connects the rate of a reaction to the concentration of reactants.

    • The general form of the rate law is:
      \text{Rate} = k[A]^m[B]^n

    • Where k is the rate constant, m and n are the reaction orders with respect to reactants A and B, determined experimentally.

Example Reaction

  • For the reaction: 2 \text{NO}(g) + \text{O}2(g) \rightarrow 2 \text{NO}2(g)

    • The rate law can be expressed as:
      \text{Rate} = k[O_2]^m[NO]^n

    • Values for m and n are determined by measuring the initial rate under varying experimental conditions.

Reaction Orders and Their Effects

  • First Order Reaction: When the reactant concentration doubles, the rate also doubles.

  • Second Order Reaction: For a doubling of the reactant concentration, the initial rate increases four-fold.

  • For a first-order reaction like A \rightarrow \text{products}

    • The rate law is:
      \text{Rate} = k[A]^n

    • If n = 0, becomes a zero-order reaction with rate independent of A's concentration.

Exercise 1

  • Reaction:
    \text{CHCl}3(g) + \text{Cl}2(g) \rightarrow \text{CCl}_4(g) + \text{HCl}(g)

  • Tasks: (a) Determine the rate law and (b) determine the rate constant k for this reaction.

Exercise 2

  • Use the following data to determine the rate law for the reaction:
    A + B + C \rightarrow D + E