In regard to collision theory, why do molecules have to have the correct orientation in order to react? Hypothetically speaking, if a reagent in a bimolecular reaction has enough energy to break the bond of the other reagent but is orientated in the wrong direction, why does the bond not break? Does the reagent's orientation affect the activation energy of the reaction somehow? The point I'm trying to make is if you have enough energy to break a bond, why would molecular orientation matter when the same amount of energy is being used to distort bonds regardless of orientation?
Suppose we have a simple reaction of a diatomic and a single atom $\ce{AB + C -&gt;A +BC}$. So the reaction is breaking the bond between $\ce{A}$ and $\ce{B}$, which then bonds to $\ce{C}$. So we might expect the transition state to look like $\ce{A--B--C}$. This suggests the reaction is highly orientation dependent. 

Looking at the process as a function of the angle formed by ABC we would expect that a 0 degree angle would vary likely lead to a reaction, since everything is lined up as we would expect, while a 180 degree angle would have almost no chance of forming the desired product as it would aligned more like this $\ce{C--A--B}$. While its not a perfectly linear relationship, we should expect that the reactivity will decrease monotonically as it goes from 0 to 180 degrees.

For more complex structures, the orientation becomes even more important. A large molecule may only have a very small reactive site, and if the other reactant collides with a completely different section of the molecule, there is virtually no chance the desired reaction will occur.