Why do some collisions not lead to product formation?

Some collisions do not lead to product formation because they lack the necessary activation energy or correct orientation.

In a chemical reaction, not all collisions between reactant molecules result in product formation. This is primarily due to two key factors: the activation energy and the orientation of the colliding molecules.

The activation energy is the minimum amount of energy that reactant molecules must possess for a successful collision to occur, leading to the formation of products. If the colliding molecules do not have this minimum energy, the collision will not result in a reaction, but instead, the molecules will simply bounce off each other and remain as reactants. This concept is part of the Collision Theory, which states that for a reaction to occur, particles must collide with sufficient energy and in the correct orientation.

The orientation of the colliding molecules is also crucial. Even if the molecules possess the required activation energy, they must collide in a manner that allows the breaking and forming of bonds to produce the desired products. If the molecules collide in an incorrect orientation, the reaction will not occur, and no products will be formed.

For instance, consider a reaction between two molecules where a bond must be formed between specific atoms in each molecule. If the molecules collide in such a way that these specific atoms do not come into contact, the reaction will not occur, regardless of the energy of the collision.

In summary, the formation of products in a chemical reaction is not guaranteed with every collision. Both the activation energy and the correct molecular orientation are necessary for a successful reaction. Without these, the collision will not lead to product formation, explaining why some collisions do not result in a reaction.