Collision Theory of Chemical Reactions
Molecular Collisions & Reaction Rates | Chemical Kinetics
1. Basic Principles
According to Collision Theory (developed by Max Trautz and William Lewis):
- Collision Frequency ($Z$): The number of collisions per second per unit volume of the reaction mixture.
- Effective Collisions: Collisions that actually lead to the formation of products.
2. Criteria for Effective Collisions
For a collision to be effective, it must overcome two main barriers:
A. Energy Barrier (Activation Energy)
The colliding molecules must possess a minimum amount of energy called Threshold Energy ($E_T$).
If molecules have lower energy, they just bounce off. The extra energy required by reactant molecules to reach the threshold is the Activation Energy ($E_a$).
$$ E_T = E_{reactants} + E_a $$The fraction of molecules having energy $\ge E_a$ is given by the Boltzmann factor: $e^{-E_a/RT}$.
B. Orientation Barrier (Steric Factor)
Molecules must collide with the Proper Orientation to facilitate the breaking of old bonds and formation of new bonds.
Example: In the reaction of $CH_3Br + OH^- \rightarrow CH_3OH + Br^-$, the $OH^-$ ion must attack the Carbon atom from the side opposite to $Br$.
3. Rate Expression
Combining the factors of Collision Frequency ($Z_{AB}$) and Activation Energy:
$$ \text{Rate} = Z_{AB} \times e^{-E_a/RT} $$However, to account for the Orientation Factor (Probability Factor, $P$), the equation is modified:
Comparing this with the Arrhenius Equation ($k = A e^{-E_a/RT}$), we see that the Arrhenius Factor ($A$) is related to collision frequency and orientation:
$$ A = P \times Z_{AB} $$4. Limitations of Collision Theory
- It assumes atoms/molecules to be Hard Spheres.
- It ignores the structural aspect of molecules (though the P factor attempts to correct this).
- It does not fully explain reactions in solution or complex chain reactions.
Practice Quiz
Test your understanding of Reaction Dynamics.
No comments:
Post a Comment