Q.According to collision theory, why do not all molecular collisions between reactant molecules result in a chemical reaction? State the two conditions a collision must satisfy in order to be effective.

[!TLDR] Only collisions with (1) enough energy to exceed the activation energy and (2) the correct mutual orientation are effective. [!ANSWER] Most collisions fail because the colliding molecules either lack sufficient energy to overcome the activation-energy barrier, or are no …

Collision theory recognises that the observed rate of most reactions is far smaller than the total frequency of molecular collisions calculated from kinetic theory, so not every collision can be leading to reaction. Two separate requirements must both be satisfied for a collision to be effective (reaction-producing). First, the colliding molecules must jointly possess kinetic energy at least equal to the activation energy $E_a$ -- since molecular energies are spread over the Maxwell-Boltzmann distribution, only a fraction of collisions involve molecules energetic enough to clear this barrier at any given temperature. Second, even sufficiently energetic molecules must collide with the correct relative orientation, so that the specific atoms and bonds that must interact are actually brought together -- a collision between the wrong parts of two molecules simply results in an unproductive bounce, however much kinetic en …

State both requirements for an effective collision (energy threshold and orientation) and …

Giving only the energy condition and omitting the orientation requirement is an incomplete answer -- …