According to collision theory, what must happen for a reaction to occur?

For a reaction to occur according to collision theory, particles must collide with the correct orientation and sufficient energy. This principle is foundational in understanding how chemical reactions take place at the molecular level.

The concept of collision theory emphasizes that merely colliding is not enough for a reaction; the collisions must also meet specific requirements. First, the kinetic energy of the colliding particles must be high enough to overcome the activation energy barrier of the reaction. This energy is necessary to break bonds in the reactants so that new bonds can form in the products. Second, the orientation of the molecules during the collision plays a crucial role. Molecules need to align in such a way that the reactive parts of the molecules are positioned correctly to facilitate bond formation.

While high speeds of molecules can contribute to the likelihood of effective collisions, it is the combination of energy and orientation that truly determines whether or not a reaction will proceed. The presence of a catalyst may increase the reaction rate by lowering the activation energy, but it is not a requirement for the occurrence of a reaction as defined by collision theory. Lastly, particles must interact to some extent for reactions to be possible; avoiding interaction would prevent any reaction from occurring. Thus, focusing on correct orientation and energy is key to understanding why this