Chemical Kinetics

Introduction to Chemical Kinetics

Chemical kinetics is the branch of chemistry that studies the rates at which chemical reactions occur and the factors that affect these rates. Understanding reaction rates is essential for predicting how quickly products form and reactants disappear, as well as for controlling industrial and biological processes.

• Rate of a chemical reaction: The speed at which reactants are converted into products.

• Feasibility: Determines if a reaction can occur under given conditions.

• Competition: Reactions may compete with one another for reactants.

• Molecular understanding: Kinetics helps explain how reactions proceed at the molecular level.

Rate of a Chemical Reaction

Defining Reaction Rate

The rate of a chemical reaction quantifies how fast reactants are consumed or products are formed. It is typically expressed as the change in concentration of a species per unit time.

• General reaction:

• Rate expression: For reactants, rate is negative (disappearance); for products, rate is positive (appearance).

• Example: For , the rate can be written for each species.

Calculating Average Rate

The average rate over a time interval is calculated using the change in concentration divided by the change in time.

• Formula:

• Units: or

• Example: If decreases from 0.065 M to 0.051 M in 10 min,

Factors Affecting Reaction Rate

Main Factors

The rate of chemical reactions is primarily controlled by several factors:

• Chemical nature of reactants

• Ability of reactants to come into contact

• Concentration of reactants

• Temperature

• Presence of catalysts

Rate Laws and Reaction Order

The Rate Law

The rate law expresses the relationship between the rate of a reaction and the concentration of reactants. It is determined experimentally.

• General form:

• k: Rate constant (depends on reaction, temperature, catalyst)

• m, n: Reaction order with respect to A and B

• Total order:

Effect of Concentration

Changing the concentration of reactants affects the rate according to the rate law. The value of k and the reaction order must be determined experimentally.

• Units of k: Depend on the overall reaction order.

• Higher k: Faster reaction.

Method of Initial Rates

This method involves measuring the initial rate of reaction for different initial concentrations of reactants to determine the reaction order.

• Compare how the rate changes as the concentration of one reactant is varied while others are held constant.

• Use ratios of rates and concentrations to solve for reaction orders.

• Example: If doubling [A] doubles the rate, the reaction is first order in A.

Sample Table: Initial Rates Data

Additional info: The table above is inferred from the provided data and shows how initial rates change with concentration.

Collision Theory

Principles of Collision Theory

Collision theory explains how chemical reactions occur and why reaction rates differ for different reactions. For a reaction to occur, reactant molecules must collide with sufficient energy and proper orientation.

• Effective collisions: Only collisions with enough energy and correct orientation lead to reaction.

• Factors: Concentration, molecular orientation, kinetic energy.

• Collision frequency: Number of collisions per second, affected by concentration and temperature.

Molecular Orientation and Kinetic Energy

• Steric factor: Probability that molecules are oriented favorably during collision.

• Activation energy (): Minimum energy required for reaction.

• Maxwell-Boltzmann distribution: Describes the distribution of kinetic energies among molecules.

Transition State Theory

Activated Complex

Transition state theory proposes that reactants form a short-lived, unstable activated complex (transition state) before forming products. The energy required to reach this state is the activation energy.

• Activated complex: Cannot be isolated; exists at the peak of the energy barrier.

• Example:

The Arrhenius Equation

Temperature Dependence of Rate Constant

The Arrhenius equation describes how the rate constant (k) varies with temperature and activation energy.

• Equation:

• A: Pre-exponential (frequency) factor

• : Activation energy

• R: Gas constant

• T: Temperature (Kelvin)

• Graph: Plotting vs yields a straight line with slope

Reaction Mechanisms

Elementary Steps and Overall Mechanism

A reaction mechanism is the sequence of elementary steps by which a chemical reaction occurs. Each step has its own rate law and molecularity.

• Elementary process: Single-step reaction

• Overall mechanism: Series of elementary steps

• Rate law for elementary step: Can be written directly from the stoichiometry

• Example: , rate =

Molecularity

Molecularity refers to the number of reactant particles involved in an elementary reaction step.

• Unimolecular: 1 reactant (e.g., )

• Bimolecular: 2 reactants (e.g., )

• Termolecular: 3 reactants (e.g., products)

Rate-Determining Step and Reaction Intermediates

The slowest step in a reaction mechanism is called the rate-determining step. Intermediates are species formed in one step and consumed in another; they do not appear in the overall reaction equation.

• Rate-determining step: Limits the overall reaction rate.

• Intermediates: Temporary species within the mechanism.

Additional info: Some equations and tables have been expanded for clarity and completeness. The notes cover all major aspects of chemical kinetics relevant to a General Chemistry course.