Collision Theory – Chemical Kinetics | Class 12 Notes
Date: August 7, 2025
📌 Definition
Collision Theory is a model that explains how chemical reactions occur and why reaction rates differ for different reactions. It states that:
🧪 Basic Assumptions of Collision Theory
- Molecules are always in random motion and collide frequently.
- Only a small fraction of the total collisions result in product formation.
- Colliding molecules must possess energy equal to or greater than the Activation Energy (Ea).
- Collisions must occur with the proper orientation.
⚙️ Rate of Reaction (Based on Collision Theory)
The rate of a chemical reaction according to collision theory is:
- Z: Collision frequency
- P: Steric factor (orientation probability)
- e-Ea/RT: Fraction of effective collisions
🔍 Important Terms
| Term | Meaning |
|---|---|
| Activation Energy (Ea) | Minimum energy needed for the reaction to occur |
| Effective Collision | Collision with correct orientation and sufficient energy |
| Steric Factor (P) | Fraction of molecules with correct orientation |
| Collision Frequency (Z) | Number of collisions per second |
📊 Maxwell-Boltzmann Distribution
This distribution shows the kinetic energies of molecules. A larger area under the curve beyond the activation energy represents a greater number of successful collisions.
- As temperature increases, the curve flattens and shifts to the right.
- More molecules attain energy ≥ Ea, increasing the reaction rate.
⚡ Temperature Effect on Rate
Increasing temperature raises the average kinetic energy of molecules, resulting in:
- Higher fraction of molecules with energy ≥ Ea
- Increased number of effective collisions
- Faster reaction rates
💡 Why All Collisions Do Not Lead to Reaction?
- Insufficient energy (energy < Ea)
- Incorrect orientation of colliding molecules
Hence, only a small portion of all collisions result in the formation of products.
🧪 Examples
Fast Reaction:
H2 + Cl2 → 2HCl (Low Ea, high Z, fast rate)Slow Reaction:
C + O2 → CO2 (High Ea, slow rate)
🧠 Transition State Theory vs Collision Theory
| Feature | Collision Theory | Transition State Theory |
|---|---|---|
| Focus | Energy & orientation of collisions | Formation of transition complex |
| Rate Prediction | Better for simple molecules | Better for complex reactions |
| Energy Barrier | Must overcome Ea by collision | Energy to form activated complex |
📗 Sample Numerical Problem
Q: The rate constant of a reaction increases from 2×10-3 to 8×10-3 mol-1L s-1 when temperature rises from 300K to 310K. Calculate the activation energy (Ea).
Use Arrhenius equation:
ln(k2/k1) = Ea/R × (T2-T1)/(T1×T2)
🔬 Applications of Collision Theory
- Understanding how temperature and catalysts affect rate
- Designing efficient chemical reactions in industries
- Explaining why gaseous reactions proceed faster
📌 Limitations
- Fails for complex, multi-step reactions
- Assumes spherical molecules
- Does not explain catalysis well
- Steric factor is often unknown
📝 Summary
- Collisions must have enough energy and correct orientation
- Only effective collisions lead to product formation
- Rate ∝ Z × P × e-Ea/RT
- More temperature ⇒ faster reaction
No comments:
Post a Comment