Lecture on Chemical Kinetics

Presented by: Roshni from Learn Hub

Key Aspects of Chemical Reactions

1. Feasibility of the Reaction (Thermodynamics)
   • \( ext{ΔG} < 0\): Spontaneous
   • \( ext{ΔG} > 0\): Non-spontaneous
   • \( ext{ΔG} = 0\): Equilibrium
2. Extent to Which Reaction Happens (Equilibrium)
   • Extent of reactants converting to products or vice versa.
3. Speed of the Reaction (Chemical Kinetics)
   • Fast vs. slow reactions.
   • Example: Diamond to graphite (super slow), mixing ink in water (super fast).

Factors Affecting Reaction Rates

1. Nature and State of Reactants
   • Example: Sodium reacts explosively with water, calcium reacts moderately.
   • Solid-state reactions are slower compared to those in the liquid or gaseous state.
2. Concentration of Reactants
   • Higher concentration usually increases the rate of reaction.
3. Temperature
   • Higher temperatures generally increase reaction rates.
4. Catalysts
   • Increase reaction rate without being consumed.

Rate of Reaction Expressions

• Macroscopic Level: Focus on amounts of reactants/products.
• Microscopic Level (Molecular Level): Focus on molecule behavior, energy, collisions.

Average and Instantaneous Rates

• Average Rate: \(- Δ[R] / Δt\)
• Instantaneous Rate: \(- d[R] / dt\) at a specific time.

Rate Law

• Differential Rate Law: \( ext{Rate} = k[R]^n \)
• Integrated Rate Law
  • Zero-order: \( [R] = [R]_0 - kt \)
  • First-order: \( ext{ln}([R]_0 / [R]) = kt \)
  • Half-life (t½): For zero-order, \( t_{1/2} = [R]0/(2k) \); for first-order, \( t{1/2} = 0.693 / k \)

Order of Reaction

• Definition: Sum of the powers of concentration terms in the rate law.
• Types: Zero-order (not dependent on [R]), first-order, second-order, fractional order.

Collision Theory

• Rate depends on collision frequency and effectiveness of collisions.
• Effective collisions require sufficient energy and proper orientation.

Activation Energy

• Minimum energy needed for a reaction.
• Graphical representation includes energy barrier (hill) indicating activation energy.

Arrhenius Equation

• Expression: \( k = Ae^{-E_a/RT} \)
  • k: Rate constant
  • A: Frequency factor (pre-exponential factor)
  • Ea: Activation energy
  • T: Temperature (K)
  • R: Gas constant (8.314 J/mol·K)
• Graphical form: \( ext{ln}(k) = -E_a/(RT) + ext{ln}(A) \)

Catalysts

• Lower the activation energy, thereby increasing the rate of reaction.
• Types: Heterogeneous (different phase), homogeneous (same phase).

Example Reactions

1. Zero Order
   • Decomposition of \( ext{N}_2 ext{O}_5\)
   • Rate: \( - d[ ext{N}_2 ext{O}_5] / dt = k[ ext{N}_2 ext{O}_5]^0 \)
2. First Order
   • Decomposition of hydrogen peroxide ( ext{H}_2 ext{O}_2)
   • Rate: \( - d[ ext{H}_2 ext{O}_2] / dt = k[ ext{H}_2 ext{O}_2] \)
3. Second Order
   • Reaction between \( 2 ext{HI} ightarrow ext{H}_2 + ext{I}_2 \)
   • Rate: \( - d[ ext{HI}] / dt = k[ ext{HI}]^2 \)

Problem-Solving Methods

1. Formula Substitution: Use rate laws to solve for rate constants, concentrations, etc.
2. Graphical Method: Plot graphs to determine reaction order, rate constants.
3. Half-life Method: Use half-life formulas for zero and first-order reactions.

Summary

• Chemical Kinetics involves the study and determination of reaction rates and the factors affecting them.
• Essential tools include understanding differential and integrated rate laws, collision theory, and the role of catalysts.

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