Slide 1: Introduction

  • Definition of Chemical Kinetics
  • Importance of studying Chemical Kinetics
  • What is Half-life in Chemical Kinetics?
  • Types of reactions in Chemical Kinetics
  • Rate of Reaction

Slide 2: Rate Expression

  • Rate Expression Definition
  • Equation for the rate of reaction:
    • Rate = k[A]^m[B]^n
    • Where k is the rate constant and m, n are the order of reaction with respect to A and B respectively
  • Example: Rate = k[H2]^2[O2]

Slide 3: Reaction Order

  • Definition of Reaction Order
  • Reaction Order for:
    • Zero Order
    • First Order
    • Second Order
    • Mixed Order
  • Example equations for each reaction order

Slide 4: Integrated Rate Laws

  • Definition of Integrated Rate Laws
  • Integrated Rate Laws for:
    • Zero Order
    • First Order
    • Second Order
  • Deriving equations and solving for concentration

Slide 5: Half-life

  • Introduction to Half-life
  • Formula for Half-life:
    • Zero Order: t1/2 = [A]0 / 2k
    • First Order: t1/2 = 0.693 / k
    • Second Order: t1/2 = 1 / k[A]0
  • Calculation of Half-life

Slide 6: Reaction Mechanisms

  • Definition of Reaction Mechanism
  • Elementary Reactions and Rate-Determining Step
  • Arrhenius Equation for Activation Energy
  • Rate Laws for Elementary Reactions

Slide 7: Catalysts

  • Introduction to Catalysts
  • Definition of Catalyst
  • Homogeneous Catalysts vs Heterogeneous Catalysts
  • Role of Catalysts in Chemical Kinetics
  • Examples of Common Catalysts

Slide 8: Temperature and Reaction Rate

  • Effect of Temperature on Reaction Rate
  • Activation Energy and its Significance
  • Arrhenius Equation: k = A * e^(-Ea/RT)
  • Graphical representation of the Arrhenius Equation

Slide 9: Concentration and Reaction Rate

  • Effect of Concentration on Reaction Rate
  • Collision Theory and Activation Energy
  • Rate law and Order of Reaction
  • Examples demonstrating the effect of concentration on reaction rate

Slide 10: Factors Affecting Reaction Rate

  • Effect of Surface Area on Reaction Rate
  • Effect of Pressure on Reaction Rate
  • Effect of Agitation on Reaction Rate
  • Effect of Catalysts on Reaction Rate
  • Examples illustrating the influence of these factors on reaction rate
  1. Chemical Kinetics - Half-life for 2nd order reaction
  • Definition of Half-life for 2nd order reaction
  • Formula for Half-life for 2nd order reaction:
    • t1/2 = 1 / (k[A]0)
    • Where [A]0 is the initial concentration of reactant A
  • Example: Consider the reaction 2A → B with rate constant k. If the initial concentration of A is 0.1 M, calculate the half-life of the reaction.
  1. Collision Theory
  • Explanation of Collision Theory
  • Factors influencing the Reaction Rate according to Collision Theory:
    • Collision Frequency
    • Orientation of Colliding Molecules
    • Energy of Collision (Activation Energy)
  • The role of Activation Energy in the Reaction Rate
  • Example: How does an increase in temperature affect the activation energy and the overall reaction rate?
  1. Arrhenius Equation
  • Introduction to the Arrhenius Equation
  • Equation:
    • k = A * e^(-Ea/RT)
    • Where k is the rate constant, A is the pre-exponential factor, Ea is the activation energy, R is the gas constant, and T is the temperature in Kelvin.
  • Importance of the Arrhenius Equation in Chemical Kinetics
  • Example calculation using the Arrhenius Equation to determine the rate constant at a given temperature.
  1. Rate-Determining Step
  • Definition of Rate-Determining Step
  • Identification of the Slowest Step in a Reaction Mechanism
  • Role of Rate-Determining Step in the Overall Reaction Rate
  • Example: Consider the reaction A + B → C + D. If the mechanism consists of the steps A + B → X and X → C + D, which step is the rate-determining step?
  1. Zero-Order Reactions
  • Explanation of Zero-Order Reactions
  • Characteristics of Zero-Order Reactions
  • Rate Law Equation: Rate = k
  • Graphical Representation of Zero-Order Reactions
  • Example: The decomposition of nitrogen pentoxide, 2N2O5 → 4NO2 + O2, follows a zero-order reaction. If the initial concentration of N2O5 is 0.4 M, calculate the concentration after 20 minutes.
  1. First-Order Reactions
  • Explanation of First-Order Reactions
  • Characteristics of First-Order Reactions
  • Rate Law Equation: Rate = k[A]
  • Graphical Representation of First-Order Reactions
  • Example: The decomposition of ethyl chloride, C2H5Cl → C2H4 + HCl, follows a first-order reaction. If the rate constant is 0.03 s^-1, calculate the time it takes for the concentration of C2H5Cl to decrease by 75%.
  1. Second-Order Reactions
  • Explanation of Second-Order Reactions
  • Characteristics of Second-Order Reactions
  • Rate Law Equation: Rate = k[A]^2
  • Graphical Representation of Second-Order Reactions
  • Example: The reaction between nitric oxide and hydrogen, 2NO + 2H2 → N2O + 2H2O, follows a second-order reaction with a rate constant of 0.05 M^-1s^-1. If the initial concentration of NO is 0.1 M, calculate the time it takes for the concentration of NO to decrease to 0.02 M.
  1. Determining Reaction Order from Experimental Data
  • Analysis of Concentration vs Time Data
  • Method for determining reaction order:
    • Zero Order: Plot [A] vs Time
    • First Order: Plot ln[A] vs Time
    • Second Order: Plot 1/[A] vs Time
  • Example: Given the following concentration vs time data for a reaction, determine the reaction order. [A]0 = 0.2 M.
    • Time (s): 0, 10, 20, 30
    • [A] (M): 0.2, 0.1, 0.05, 0.025
  1. Temperature Dependence of Reaction Rate Constant
  • Explanation of the Temperature Dependence of Rate Constant
  • Effect of Temperature on Average Kinetic Energy of Molecules
  • Collision Frequency and Activation Energy
  • Relationship between Temperature and Rate Constant: Arrhenius Equation
  • Example: The rate constant of a reaction is 0.02 s^-1 at 25°C and the activation energy is 50 kJ/mol. Calculate the rate constant at 50°C.
  1. Factors Affecting Reaction Rate: Concentration
  • Influence of Reactant Concentration on Reaction Rate
  • Effect of Increasing Reactant Concentration on Collision Frequency
  • Rate Law Equation and Relation to Concentration
  • Example: The rate of a reaction is directly proportional to the square of the concentration of reactant A. If the concentration of A is doubled, how does it affect the reaction rate?
  1. Chemical Equilibrium
  • Definition of Chemical Equilibrium
  • Forward and Reverse Reactions
  • Equilibrium Constant (K) and its Expression
  • Relationship between K and Reaction Quotient (Q)
  • Le Chatelier’s Principle and its Application to Equilibrium
  • Example: N2(g) + 3H2(g) ⇌ 2NH3(g). If the equilibrium constant (Kc) is 4.0 x 10^2, calculate the value of Kc for the reverse reaction.
  1. Factors Affecting Equilibrium Position
  • Effect of Concentration on Equilibrium Position
  • Effect of Temperature on Equilibrium Position
  • Effect of Pressure (for gaseous reactions) on Equilibrium Position
  • Effect of Catalysts on Equilibrium Position
  • Examples illustrating the influence of these factors on the equilibrium position
  1. Solubility Equilibrium
  • Definition of Solubility Equilibrium
  • Solubility Product (Ksp) and its Expression
  • Common Ion Effect and its Impact on Solubility
  • Calculating Solubility using Ksp
  • Example: Determine the solubility of lead(II) iodide (PbI2) in a solution with a Ksp of 2.5 x 10^-8.
  1. Acid-Base Equilibria
  • Definition of Acid-Base Equilibria
  • Acid Dissociation Constant (Ka) and its Expression
  • Base Dissociation Constant (Kb) and its Expression
  • pH and pOH Calculations
  • Examples involving acid-base equilibria and pH calculations
  1. Henderson-Hasselbalch Equation
  • Introduction to the Henderson-Hasselbalch Equation
  • Equation: pH = pKa + log ([A-]/[HA])
  • Calculations using the Henderson-Hasselbalch Equation
  • Applications of the Henderson-Hasselbalch Equation
  • Example: Calculate the pH of a solution containing 0.1 M acetic acid (CH3COOH) and 0.05 M sodium acetate (CH3COONa). The pKa of acetic acid is 4.76.
  1. Buffer Solutions
  • Definition of Buffer Solutions
  • Composition of a Buffer Solution
  • Buffer Capacity and pH Range
  • Preparation of Buffer Solutions
  • Examples of Buffer Solutions in Biological Systems
  1. Acid-Base Titrations
  • Introduction to Acid-Base Titrations
  • Equivalence Point and Endpoint in Titration
  • Indicator Selection and Use in Acid-Base Titrations
  • Calculations involving Acid-Base Titrations
  • Example: A 25 mL sample of HCl solution of unknown concentration is titrated with 0.1 M NaOH. The volume of NaOH required to reach the equivalence point is 20 mL. Calculate the concentration of HCl.
  1. Redox Reactions
  • Definition of Redox Reactions
  • Oxidation and Reduction Half-Reactions
  • Balancing Redox Equations
  • Oxidizing and Reducing Agents
  • Examples of Redox Reactions in Everyday Life
  1. Electrochemical Cells
  • Introduction to Electrochemical Cells
  • Half-Cells, Anode, and Cathode
  • Cell Potential (Ecell) and its Calculation
  • Nernst Equation and its Application
  • Example: Calculate the cell potential of a standard hydrogen electrode (SHE) when the concentration of H+ ions is 0.5 M at 25°C.
  1. Standard Electrode Potentials
  • Definition of Standard Electrode Potentials
  • Reduction Potential (E°) and its Significance
  • Standard Hydrogen Electrode (SHE) as Reference
  • Tabulated Standard Electrode Potentials
  • Examples using Standard Electrode Potentials in Determining Cell Potential