What is activation energy?

When reactant molecules collide, they don't automatically turn into products. They must collide with enough force to reach an unstable, high-energy arrangement called the transition state. The activation energy is the exact difference in energy between the reactants and this transition state. If the molecules hit each other gently, they just bounce off without reacting. They need to collide with an energy equal to or greater than $E_a$ to successfully react.

The concept of activation energy explains why temperature affects reaction rates. Heating a substance gives its molecules more kinetic energy. With more energy, a larger fraction of the molecules can overcome the activation energy barrier when they collide, making the reaction happen faster. If you want to speed up a reaction without adding heat, you can use a catalyst. A catalyst provides an entirely different pathway for the reaction—one that has a lower activation energy. It is like finding a tunnel through the hill instead of pushing the boulder over the top.

A common mistake is confusing activation energy ($E_a$) with the net energy change of the reaction ($\Delta H$). Activation energy is the hill you must climb to start the reaction. The net energy change is the difference in elevation between where you started (reactants) and where you finished (products). A reaction can be highly exothermic (releasing a lot of energy, meaning a negative $\Delta H$), but if its activation energy is very high, it won't happen without a massive initial spark.