Overview
This lecture explains collision theory, including its definition, significance of orientation and temperature, and why reactions often occur in steps.
Collision Theory Basics
- Collision theory states that atoms or molecules must collide to react.
- Successful reactions require collisions with sufficient energy and correct orientation.
Importance of Orientation
- Correct orientation means atoms must be aligned properly during collision for a reaction to occur.
- Example: Only specific atoms on different molecules must collide for a reaction; wrong alignments don't react.
Role of Temperature and Concentration
- Increasing temperature causes molecules to move faster, raising collision chance and energy.
- Higher concentration means more molecules in solution, increasing collision likelihood.
- Both higher temperature and concentration speed up reaction rates.
Reaction Steps
- Simultaneous collision of many molecules is rare due to low probability of correct orientation and energy.
- Reactions often proceed in steps, allowing fewer molecules to interact at each stage.
- Reaction steps form intermediates that can react further, making the process more feasible.
Key Terms & Definitions
- Collision Theory — The idea that atoms/molecules must collide with enough energy and the correct orientation to react.
- Orientation — The specific alignment required between reacting atoms for a reaction to occur.
- Concentration — The amount of reactant molecules in a solution.
- Reaction Steps — Stages in a reaction where smaller groups of molecules combine, often forming intermediates.
Action Items / Next Steps
- Review textbook sections on collision theory and reaction mechanisms.
- Practice identifying reaction steps and intermediates for sample reactions.