Threshold energy is the minimum total energy that colliding molecules must possess for a reaction to occur, while activation energy is the extra energy that reactant molecules need above their average energy to reach this threshold. The two are related by threshold energy = activation energy + average energy of the reactants. In other words, activation energy is the gap between the average energy of the molecules and the threshold required for an effective collision. The option stating threshold energy equals activation energy minus average energy reverses the relationship. The option that they are unrelated contradicts the very definition of activation energy. The option that threshold energy is always zero would imply every collision is effective, which is false. Understanding this distinction clarifies energy-profile diagrams in NCERT kinetics. This concept also bridges to Chemical Thermodynamics and Some Basic Concepts in Chemistry, so mastering it strengthens performance on linked questions from those topics as well. It is worth emphasising that this is not a special case but a representative example of how threshold energy operates throughout chemical kinetics. Plausibility check: since molecules already carry average thermal energy, only the additional activation energy is needed to climb to the threshold, confirming that the threshold is the sum of the two.