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
- 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.
- 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?
- 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.
- 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?
- 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.
- 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%.
- 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.
- 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
- 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.
- 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?
- 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.
- 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
- 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.
- 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
- 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.
- 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
- 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.
- 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
- 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.
- 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
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