CHEM LECTURE (02/07/25)

Homework and Upcoming Assessments

• Homework is due next Thursday (Alex assignment).

• No current quizzes to worry about.

• Next quiz will be on kinetics and is scheduled for next week.

Recap of Last Class

• Focus on second order reactions involving iodine atoms.

• Discussed the relationship between concentration and half-life.

  • Half-life increases as concentration decreases.

  • Lower concentrations result in a slower rate of reaction due to fewer particle collisions.

Understanding Reaction Orders and Half-Life Equations

• Emphasized the importance of knowing the equations for different reaction orders:

  • Zero Order: Rate is constant and does not depend on concentration. Example: uranium decomposition.

  • First Order: Rate depends on the concentration of one reactant.

  • Second Order: Rate depends on the concentrations of two reactants.

    • Can be different reactants or the same (e.g., I + I -> I2).

• For calculations:

  • Identify the order of the reaction.

  • Determine necessary calculations (rate law, concentrations, half-lives, etc.).

Kinetics and Temperature Effects

• Reaction rates are affected by temperature changes.

• Need to consider that experiments may be conducted at varying temperatures.

• Example illustrated an experiment during different thermal conditions leading to discrepancies in results.

• Important Equation: ln(k) = -Ea/R(1/T) + constant, where:

  • k = rate constant

  • Ea = activation energy

  • R = gas constant (temperature in Kelvin)

  • T = absolute temperature in Kelvin

Importance of Activation Energy and Units

• When calculating activation energy, ensure:

  • Use consistent units (Ea in kJ, and R in J).

  • Activation energy generally results in a positive value, indicating an energy barrier in reactions.

• Ensure accurate plotting of data to linearize the equation for calculations.

Reaction Mechanisms

• The overall reaction may differ from the reaction mechanism.

• Elementary Steps: These are the individual reactions that combine to form the overall reaction.

  • Intermediates: Substances formed and consumed in a reaction that do not appear in the overall equation.

  • Catalysts: Substances that facilitate the reaction and appear in the overall equation.

• Example using NO and O2 to form NO2 reveals underlying reactions and intermediates.

Rate Determining Steps

• Discussed the concept of the rate determining step being the slowest step in a reaction mechanism.

  • The overall rate can only occur as quickly as the rate of the slowest step.

• Elementary Reactions Recap:

  • Unimolecular: One molecule reaction (first order).

  • Bimolecular: Two molecules react (second order).

  • Termolecular: Rare; involves three molecules reacting simultaneously.

Proposed Mechanisms and Examples

• When evaluating reactions:

  • Overall reaction must match the sum of elementary steps.

  • Rate law of the overall reaction should align with the rate law derived from the rate determining step.

• Example of hydrogen peroxide decomposition illustrates the mechanism and identification of intermediates and catalysts.

• Energy diagrams depict activation energies and which step is slower based on energy barriers.

Key Takeaways for Quizzes

• Be prepared to identify intermediates and catalysts in reactions.

• Understand how temperature affects reaction rates.

• Practice calculations involving activation energy and rates using provided equations.

• Familiarize yourself with reaction mechanisms and rate determining steps.