Chemical reactions occur when reactants collide. For what reasons may a collision fail to produce a chemical reaction?

Collision theory tells us that in order for chemicals to react, three things must occur. The reactants must collide, the collision must have sufficient energy, and the orientation of reactants must be conducive for the reaction to proceed.
The question presumes that a collision occurs. The collision could fail to produce a reaction if either of the remaining conditions is not met, that is, if there is insufficient energy or incorrect orientation of reactants.
What is "sufficient energy" for a chemical reaction? A chemical reaction generally requires the breaking of existing chemical bonds and the formation of new ones. Breaking bonds requires energy, and forming bonds releases energy. Bond breaking must begin before bond formation, so energy must be supplied to initiate a reaction. A certain minimum energy input is required to form an "activated complex," which can continue to form products with no further energy needed. This minimum required energy is called the "activation energy." It is supplied by the kinetic energy of the collision and is related to the relative speed of the colliding reactants.
Thus only collisions having a certain minimum relative speed will have sufficient energy to produce a reaction. At a given temperature there will be a distribution of reactant speeds, and some proportion of collisions will have sufficient energy to react. (Raising the temperature increases the average speed reactants move.)
What is the "correct orientation" for a chemical reaction? The discussion of formation of an activated complex provides the explanation. Energy input causes bonds to begin to break, but unless atoms are in all of the right positions for new bonds to form, reaction will not occur. Completely breaking chemical bonds without simultaneous formation of new bonds would require far more than the activation energy. Collisions in which the atoms expected to bond remain far apart thus do not have the correct orientation to react.
In summary, a collision between potential reactants will not result in a chemical reaction if either there is insufficient energy, meaning the kinetic energy of the collision is less than the activation energy, or the orientation is not correct for reaction to occur.

A reaction will occur if the reactants are at or above the activation energy. This energy level is different for every set of chemicals and is related to temperature. Hotter reactants will react faster than colder ones for this reason. Reactions will still occur slowly at even extremely low energy states when energy from one cold molecule transfers to another cold molecule making it just warm enough to react.
Reactions can also occur when a catalyst is introduced to the reaction. Catalysts allow the reaction to occur at a lower energy state.This means a reaction will occur faster and at a lower temperature when exposed to a catalyst.
Finally, in reactions involving very large reactants, you must take into account activation sites or places on the molecules where the reactions actually occur. This is usually the case when dealing with large proteins or enzymes and can be ignored when dealing with smaller molecules like salts. The result is that a reaction will not occur on a larger molecule if the reactants are not properly aligned.