Chemical Kinetics and Chemical Equilibrium

Chemical Kinetics Overview

Introduction

  • Definition of Chemical Kinetics: Study of reaction rates and the factors affecting them.

  • Relation to Thermodynamics: Thermodynamics assesses whether reactions occur, while kinetics measures the rate of reactions.

Reaction Rate Basics

  • Definition of Reaction Rate: Change in concentration of reactants/products over time (M/s).

    • Rate A ⇌ B:

      • Rate = -D[A]/Dt (for reactants, decreases concentration).

      • Rate = D[B]/Dt (for products, increases concentration).

  • Example: Tracking concentrations of A and B over time.

Average vs. Instantaneous Rates

  • Average Rate Calculation:

    • Using initial and final concentrations over a time span.

    • Average rate = - D[Br2]/Dt

  • Instantaneous Rate: Rate at a particular moment, represented by the slope of the tangent to the concentration vs time graph.

Rate Law and Expression

  • Rate Law: Defines the relationship between rate of reaction, rate constant (k), and reactant concentrations raised to specific powers.

    • Example: Rate = k [A]^x [B]^y.

    • Reaction order: xth in A, yth in B, and (x+y)th overall.

First-Order Reactions

  • Equation: Rate = -D[A]/Dt = k[A].

  • Characteristics:

    • Doubling concentration of A doubles the rate.

    • Half-life (t½) is constant and can be calculated: t½ = 0.693/k.

  • Decay Example: If [A]0 = 0.88 M to [A] = 0.14 M, use the natural logarithm to determine time.

Second-Order Reactions

  • Equation: Rate = k[A]^2.

  • Half-Life: t½ = 1/(k[A]0).

  • Characteristics: The rate is proportional to the square of the concentration of the reactant.

Zero-Order Reactions

  • Equation: Rate = k.

  • Half-Life: t½ = [A]0/(2k).

  • Characteristics: Rate remains constant regardless of concentration.

Activation Energy and Rate Constant

  • Definition: Minimum energy needed for a reaction to occur.

  • Arrhenius Equation: k = A * exp(-Ea/RT).

    • Where A is the frequency factor, R is the gas constant, and T is the temperature in Kelvin.

  • Graphical Representation: Plot ln(k) vs. 1/T to find Ea from the slope.

Reaction Mechanisms

  • Elementary Steps: Sequence that leads to product formation.

  • Intermediates: Species formed in a reaction mechanism not present in the overall equation.

  • Molecularity: Number of molecules involved in the elementary step (unimolecular, bimolecular, termolecular).

Catalysis

  • Definition: Substance that increases the rate of a reaction without being consumed.

  • Types:

    • Heterogeneous: reactants and catalysts in different phases.

    • Homogeneous: reactants and catalysts in the same phase.

  • Effect on Activation Energy: Catalysts lower activation energy for both directions.

Chemical Equilibrium

  • Definition: State where rates of the forward and reverse reactions are equal.

  • Le Châtelier’s Principle: System adjusts to offset applied stress (changes in concentration, pressure, and temperature).

  • Equilibrium Constant: K = [products]/[reactants] raised to their stoichiometric coefficients.

Summary of Concepts

  • Rate Laws: Determined experimentally and relate to individual components in a chemical equation.

  • Shifting Equilibrium: The addition or removal of reactants/products shifts the reaction to the left or right, affecting concentration and rates.

  • Temperature Effects: Temperature changes can favor either the endothermic or exothermic direction of a reaction, thus affecting K values.