1.1 Fundamentals of chemical kinetics

Chemical kinetics studies how fast reactions happen and what affects their speed. It's all about tracking how quickly reactants turn into products over time. This knowledge helps scientists control reactions in various fields, from making medicines to developing new materials.

Reaction rates depend on things like how much of each chemical is present and how easily they can overcome energy barriers. Factors like temperature, concentration, and catalysts can speed up or slow down reactions. Understanding these influences helps predict and control chemical processes.

Introduction to Chemical Kinetics

Definition of chemical kinetics

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• Studies the rates of chemical reactions and the factors that influence them
• Investigates how quickly reactants are consumed and products are formed over time
• Provides insights into the mechanisms and pathways of chemical reactions
• Crucial for understanding and controlling reactions in various fields (chemical engineering, pharmaceuticals, materials science)

Reaction rate and concentration relationships

• Reaction rate quantifies the change in concentration of reactants or products per unit time
• Mathematically expressed as $rate = -\frac{d[A]}{dt} = \frac{d[B]}{dt}$, where [A] and [B] are concentrations of reactant and product
• Rate law describes the relationship between reaction rate and reactant concentrations: $rate = k[A]^m[B]^n$
  • k is the rate constant, m and n are the orders of the reaction with respect to reactants A and B
• Increasing reactant concentrations generally increases the reaction rate by providing more molecules for collisions
• Product accumulation can sometimes slow down the reaction rate through product inhibition

Factors Influencing Reaction Rates

Role of activation energy

• Activation energy (Ea) is the minimum energy required for reactants to initiate a chemical reaction
• Reactant molecules must collide with sufficient energy to overcome the activation energy barrier
• Higher activation energies lead to slower reaction rates, while lower activation energies result in faster rates
• Catalysts can lower the activation energy and increase the reaction rate without being consumed
• Arrhenius equation relates the rate constant (k) to the activation energy: $k = Ae^{-Ea/RT}$
  • A is the pre-exponential factor, R is the gas constant, and T is the absolute temperature
  • Exponential relationship between rate constant and activation energy

Elementary vs complex reactions

• Elementary reactions occur in a single molecular event with no intermediates
  • Involve a single transition state
  • Molecularity determines the order of elementary reactions
    1. Unimolecular reactions: first-order (radioactive decay)
    2. Bimolecular reactions: second-order (dimerization)
    3. Termolecular reactions: third-order (rare, ozone formation)
• Complex reactions consist of multiple elementary steps with intermediates
  • Overall reaction order may not be easily determined from the balanced equation
  • Rate-determining step (slowest step) governs the overall rate of the complex reaction
  • Steady-state approximation and rate-limiting step simplifications are used to analyze complex reactions
  • Examples: enzymatic reactions, combustion of hydrocarbons