Question

The central idea of the collision model is that molecules must collide in order to react. Give two reasons why not all collisions of reactant molecules result in product formation.

Short answer

Not all collisions of reactant molecules result in product formation due to two main reasons: firstly, the collisions might not have sufficient activation energy, which is the minimum energy required for a successful reaction; secondly, the reactant molecules might collide at an incorrect orientation, preventing efficient overlapping of the reacting orbitals and thus, hindering the bonding process.

Step-by-step solution

Understand the Collision Theory

The collision model proposes that for a chemical reaction to occur, the reactant molecules must collide with each other. However, not all collisions will lead to a reaction, and certain conditions must be met for the reaction to proceed.

Reason 1: Insufficient Activation Energy

The reactant molecules need to have a minimum amount of energy for a successful reaction, which is called the activation energy. Activation energy is the energy barrier that needs to be overcome for a reaction to occur. When molecules collide with less than the required activation energy, they may simply bounce off each other and not react.

Reason 2: Incorrect Orientation

The reactant molecules must also collide at the correct orientation (approach angle) for a successful reaction. The collision orientation determines the efficient overlap between the reacting orbitals of the molecules, allowing the bonding process to take place. If the reactant molecules collide at an incorrect orientation, they may bounce off each other and not form a reaction.