Chemical Kinetics: Reaction Rates and Rate Laws

Reaction Rates and Rates Laws

• Reaction rate is the speed at which reactants are consumed or products are formed.
• It's typically expressed as the change in concentration of a reactant or product over time.
• Rates are typically positive, signifying the consumption of reactants or formation of products.
• Rate depends on reactant concentrations and reaction temperature.
• The rate law describes the relationship between reaction rate and reactant concentrations.
• It's experimentally determined; not derived from the balanced chemical equation.
• The rate law format is: rate = k[A]^m[B]^n, where k is the rate constant, [A] and [B] are reactant concentrations, and m and n are the reaction orders.
• The overall reaction order is the sum of the individual orders (m + n).
• The rate constant (k) is unique to a reaction and temperature-dependent.
• A reactant's order is experimentally determined.

Determining Rate Laws

• Initial rates method: Measure the initial reaction rate at various initial reactant concentrations.
• This method is valuable for determining individual reactant orders by comparing rate changes with concentration changes.
• Integrated rate laws: Provide equations for calculating reactant or product concentrations at any point during the reaction.
• Different reaction orders yield distinct integrated rate law equations.

Reaction Mechanisms

• A reaction mechanism is a series of elementary steps detailing the overall reaction.
• Elementary steps represent molecular-level processes; the slowest step is the rate-determining step.
• The overall reaction rate law is determined by the rate law of the slowest elementary step.
• The sum of elementary steps equals the overall balanced equation.
• Mechanisms may involve intermediate species, which are neither reactants nor products.

Factors Affecting Reaction Rates

• Temperature: Increased temperature boosts reactant kinetic energy, leading to more frequent and forceful collisions, accelerating the reaction.
• Concentration: Higher reactant concentration increases the number of molecules per unit volume, enhancing collisions and reaction speed.
• Surface area: For heterogeneous reactions (different phases), larger surface area of a solid reactant increases the reaction area, thus accelerating the reaction.
• Catalysts: Catalysts accelerate reactions by providing alternate pathways with lower activation energies. Catalysts are not consumed in the reaction.
• Inhibitors: Inhibitors slow reactions by interfering with the reaction mechanism.

Activation Energy (Ea)

• Activation energy (Ea) is the minimum energy needed for a reaction to occur.
• It represents the energy hurdle between reactants and products.
• Catalysts alter reaction mechanisms, reducing Ea and boosting reaction rates.
• The Arrhenius Equation links the rate constant (k), activation energy (Ea), temperature (T), the ideal gas constant (R), and the frequency factor (A).

First-Order Reactions

• First-order reaction rates are directly proportional to the concentration of one reactant.
• The integrated rate law for a first-order reaction shows a linear relationship between the natural logarithm of reactant concentration and time.
• The half-life of a first-order reaction is independent of the initial reactant concentration.

Second-Order Reactions

• Second-order reactions have rates proportional to the square of one reactant's concentration or the product of two reactants' concentrations.
• More complex integrated rate law equations exist for second-order reactions than for first-order reactions.
• The half-life of a second-order reaction varies with the initial reactant concentration.

Zero-Order Reactions

• Zero-order reaction rates are unaffected by reactant concentrations.
• The integrated rate law for a zero-order reaction shows a linear relationship between reactant concentration and time.
• The half-life of a zero-order reaction changes with the initial reactant concentration.