Chemistry
Why Do Some Reactions Require a High Activation Energy?
Quick fact
Activation energy is the minimum energy colliding molecules need to break bonds and form new ones – some reactions need a lot because their bonds are very strong or their orientation is highly specific.
Why this is interesting
Why does a spark ignite a forest fire but leave a rock unchanged? The answer lies in the energy barrier that some reactions must overcome.
Read the full explanation
Understanding Why Do Some Reactions Require a High Activation Energy?
For a chemical reaction to happen, reactant molecules must collide with enough energy and the correct orientation to break existing bonds and create new ones. The activation energy is the energy needed to reach the transition state – the highest-energy, most unstable arrangement along the reaction pathway. In reactions with high activation energy, the reactants have very strong bonds (like the triple bond in nitrogen gas) or require a very precise, strained geometry to form the transition state. This means only a tiny fraction of collisions have sufficient energy, making the reaction slow without added heat or a catalyst.
A deeper explanation
The height of the activation energy barrier is determined by the electronic structure of the reactants. Stronger bonds (higher bond dissociation energy) require more energy to break. Additionally, the transition state itself is inherently unstable – if it involves significant angle strain, charge separation (e.g., formation of a carbocation), or violation of octet rules, the energy penalty is large. This is why reactions like the combustion of diamond (C → CO₂) are thermodynamically favorable but kinetically hindered: the strong C–C bonds in diamond have a high activation barrier. Understanding this mechanism explains why catalysts work by providing an alternative pathway with a lower-energy transition state, and why some reactions need high temperatures to proceed.