According to collision theory, for a chemical reaction to occur, colliding molecules must possess

1. NCERT §3.6 (Collision Theory) gives two conditions for a successful (productive) collision. 2. ENERGY: the kinetic energy of the colliding molecules must equal or exceed the activation energy $E_a$ (the threshold) so that bonds can break. 3. ORIENTATION: the molecules must collide along an axis that allows the relevant bonds to form/break — this is captured by the steric factor $p$ in the Arrhenius pre-exponential $A$. 4. So a successful collision needs BOTH: sufficient energy AND correct geometry. Most collisions fail one or both tests, which is why measured rate constants are much smaller than the collision frequency. 5. Option A reverses the energy requirement. Option C is irrelevant (it's the SUM of energies relative to $E_a$ that matters, not individual values). Option D is plainly wrong — high energy enables, not hinders, reaction. _Source: NCERT Class 12 Chemistry Part 1, Ch 3, §3.6 (Collision Theory of Chemical Reactions), p. 22–23._