Final answer:
For a chemical reaction to proceed efficiently, it is necessary for the collision to have the proper orientation, which ensures the reactants can form the necessary activated complex, and it must also have enough energy to overcome the activation energy barrier.
Step-by-step explanation:
In the context of chemical reactions, a collision must happen with the proper orientation because it is a critical factor that determines whether the reactants will form the intended products or not. For example, in a simple bimolecular reaction where two reactants collide, the atoms that are to form a new bond need to be aligned correctly, so they can effectively interact with each other. Improper orientation can prevent the reactants from forming the necessary activated complex or transition state needed to result in product formation.
In a two-dimensional collision scenario, proper orientation is also important to prevent undesired outcomes like the rotation of the molecules which can further complicate the collision process. This is analogous to two ice skaters who hook arms and spin, which is not ideal as we want to consider only the direct impact of the collision without any rotation or spinning of point masses.
Moreover, the energy of activation is another crucial aspect; even with the correct orientation, the collision must possess sufficient energy to overcome the reaction's energy barrier. If it doesn't, the reaction won't proceed efficiently. A good illustration of this is the collision between an F ion and H3C-I; if the collision is not effectively oriented to promote C-I bond breaking, the reaction is less likely to proceed or might be ineffective.