Understanding Reaction Kinetics

• Kinetics involves the study of reaction rates, how reaction rates change under different conditions, and the molecular events during a reaction.
• Reaction rate signifies the speed at which reactants turn into products during a chemical reaction.

Factors Affecting Reaction Rates

• Reactant Concentration: Higher reactant concentrations usually increase reaction rates.
• Higher concentrations increase reaction rates due to a greater number of reactant molecules available to react.
• Temperature: Reaction rates generally increase with temperature.
• Increased temperature causes molecules to possess more kinetic energy, leading to more frequent and energetic collisions.
• Physical State: Reactants in the same phase react more readily.
• Reactants in the same phase have increased contact compared to those in different phases.
• Surface Area: Increased surface area leads to higher reaction rates in reactions involving solids.
• Presence of a Catalyst: Catalysts speed up reactions without being consumed.
• Catalysts provide an alternative pathway with lower activation energy.
• Light: Light can influence reaction rates, especially in photochemical reactions.

Rate Laws

• Rate laws express the relationship between the rate of reaction and reactant concentrations.
• For the reaction aA + bB → cC + dD, the rate law typically is: rate = k[A]^m[B]^n.
  • k is the rate constant, which depends on temperature and is specific to each reaction.
  • [A] and [B] represent the concentrations of reactants A and B.
  • m and n are the reaction orders with respect to A and B, and they must be determined experimentally.
• Overall reaction order is the sum of the individual orders (m + n).

Determining Rate Laws

• The method of initial rates involves measuring the initial rate of a reaction using varying initial concentrations of reactants, then comparing how the rate changes with concentration.
• Integrated rate laws relate the concentration of reactants to time, allowing the determination of reaction order and the rate constant from concentration-time data.

Reaction Order

• Zero Order: Rate is independent of reactant concentration; rate = k.
• First Order: Rate is directly proportional to reactant concentration; rate = k[A].
• Second Order: Rate is proportional to the square of reactant concentration: rate = k[A]^2, or to the product of the concentrations of two reactants: rate = k[A][B].
• Pseudo-Order: Occurs if one or more reactants are present in large excess such that their concentration remains essentially constant.

Integrated Rate Laws

• Integrated rate laws illustrate how reactant concentration changes over time.

Zero-Order Reactions

• Rate Law: rate = k
• Integrated Rate Law: [A]t = -kt + [A]0
• Half-life: t1/2 = [A]0 / 2k

First-Order Reactions:

• Rate Law: rate = k[A]
• Integrated Rate Law: ln[A]t = -kt + ln[A]0 or [A]t = [A]0 * e^(-kt)
• Half-life: t1/2 = 0.693 / k

Second-Order Reactions

• Rate Law: rate = k[A]^2
• Integrated Rate Law: 1/[A]t = kt + 1/[A]0
• Half-life: t1/2 = 1 / (k[A]0)

Collision Theory

• Collision theory states that reactant molecules must collide with sufficient energy and proper orientation for a reaction to occur.
• Activation Energy (Ea): Minimum energy for a collision to result in a chemical reaction.
• Orientation Factor (p): Fraction of collisions with effective orientations.
• Arrhenius equation: k = A * e^(-Ea/RT).
  • Describes the relationship between the rate constant, activation energy, and temperature.
  • A is the frequency factor relating to collision frequency and molecular orientation.
  • R is the gas constant (8.314 J/(mol·K)).
  • T is the absolute temperature in Kelvin.

Transition State Theory

• Transition state theory (or activated-complex theory) focuses on the formation of a high-energy transition state (or activated complex) during a reaction.
• The transition state represents the point of maximum energy along the reaction pathway.
• The difference in energy between reactants and the transition state represents the activation energy.

Reaction Mechanisms

• A reaction mechanism details the step-by-step sequence of elementary reactions composing the overall chemical reaction.
• Elementary Reaction: A single-step reaction that cannot be broken down further.
• Molecularity: The number of molecules participating as reactants in an elementary reaction (unimolecular, bimolecular, termolecular).
• Rate-Determining Step: The slowest step in a reaction mechanism, determining the reaction's overall rate.
• Intermediates: Species formed in one step and consumed in a subsequent step; they do not appear in the overall balanced equation.
• Catalysts: Substances increasing reaction rate without being consumed, providing an alternative pathway with lower activation energy.

Catalysis

• Homogeneous Catalysis: The catalyst is in the same phase as the reactants.
• Heterogeneous Catalysis: The catalyst is in a different phase from the reactants, often involving reactant adsorption onto the surface of a solid catalyst.
• Enzymes: Biological catalysts, typically proteins, that demonstrate high specificity and efficiency.