Chemical Kinetics

Introduction to Chemical Kinetics

• Definition: A branch of physical chemistry that studies the rate of chemical reactions.
• Focus Areas: Rate and mechanism of chemical reactions.

Classification of Chemical Reactions

• Types of Reactions Based on Kinetics:
  1. Fast Reactions:
     • Complete in a fraction of a second.
     • Examples:
       • Explosive reactions
       • Neutralization reactions (e.g., NaOH + HCl → NaCl + H2O)
       • Precipitation reactions (e.g., AgNO3 + NaCl → AgCl + NaNO3)
  2. Intermediate Reactions:
     • Complete in a measurable time period.
     • Examples:
       • Hydrolysis of esters
       • Decomposition of hydrogen peroxide (H2O2 → H2O + O2)
  3. Slow Reactions:
     • Take a long time to complete.
     • Examples:
       • Rusting of iron
       • Aging of food, humans, and erosion of soil

Rate of Reaction

Definitions

• Rate of Reaction: Describes how reactants decrease and products increase over time.

Rate in Terms of Amount

• Rate = Change in amount of reactant/product per unit time.
• Units: mol/time (e.g., mol/s, mol/min).

Rate in Terms of Concentration

• Rate = Change in concentration of reactant/product per unit time.
• Units: mol/L/time (e.g., mol/L/s).

Expressions of Rate of Reaction

• Rate of Disappearance of Reactants:
  • Formula: Rate = -Δ[A]/Δt
• Rate of Formation of Products:
  • Formula: Rate = Δ[B]/Δt

Stoichiometric Relationships

• In a balanced reaction:
  • 1 mole of A leads to 2 moles of B:
    • Rate of disappearance of A = 1/2 Rate of formation of B.

Equivalent Rate

• General Reaction: A + B → C + D
• Rate expressions involve coefficients of reactants and products in the rate of reaction.
• Order of reaction is derived from the sum of the powers in the rate law.

Average Rate vs. Instantaneous Rate

• Average Rate: Rate over a specific time interval.
  • Formula: Average Rate = -Δ[A]/Δt or Δ[B]/Δt.
• Instantaneous Rate: Rate at a specific moment.
  • Formula: Limit as Δt → 0 of -Δ[A]/Δt or Δ[B]/Δt.

Factors Affecting the Rate of Reaction

1. Nature of Reactants:
   • Ionic reactants react faster than covalent ones.
2. Concentration:
   • Increased concentration results in higher reaction rates due to more collisions.
3. Surface Area:
   • Larger surface area increases reaction rates (e.g., powdered substances vs. solids).
4. Catalysts:
   • Substances that increase reaction rates without being consumed.
5. Temperature:
   • Higher temperatures increase kinetic energy, leading to faster reactions.
6. Light/Radiation:
   • Affects photochemical reactions; increasing light intensity increases the reaction rate.

Rate Law Equation

• Rate Law: Mathematical relationship showing how the rate of reaction depends on concentration.
• General form: Rate = k[A]^m[B]^n, where k is the rate constant.
  • m and n are the orders of the reaction with respect to A and B respectively.

Rate Constant and Units

• Rate Constant (k): Defines the relationship between rate and concentration at unit values.
• Units of Rate Constant:
  • Zero-order: mol/L/time
  • First-order: 1/time
  • Second-order: L/(mol*time)
  • Third-order: L^2/(mol^2*time)

Order of Reaction

• Definition: Sum of the powers of concentration terms in the rate law.
• Can be affected by physical conditions (temperature, pressure, etc.).
• Values can be 0, positive, negative, or fractional.

Molecularity

• Definition: Number of reacting species that simultaneously collide to form products.
• Theoretical parameter that is always a positive integer.
• Types:
  • Unimolecular: 1 species
  • Bimolecular: 2 species
  • Termolecular: 3 species (rare).

Differences Between Order and Molecularity

• Order: Experimental parameter based on reaction concentration.
• Molecularity: Theoretical parameter based on the molecularity of the reaction, always a positive integer.
• Order can vary with conditions while molecularity is invariant.