Think of activation energy like pushing a heavy boulder up a steep hill. You have to put in a lot of hard work to reach the very top. Once you push it over the peak, it rolls down the other side easily.
In chemistry, molecules must crash into each other with enough force to break their existing bonds. If they bump together too softly, they just bounce off each other with no reaction. The Swedish chemist Svante Arrhenius first explained this important concept in 1889.
He created the Arrhenius equation to link activation energy directly to the reaction rate. This equation shows why adding heat makes chemical reactions happen much faster. A higher temperature means the molecules are moving around with much more kinetic energy.
Because they are moving faster, many more molecules can successfully climb the energy hill. Students sometimes think that a catalyst lowers the original energy hill. Instead, a catalyst provides a completely different pathway with a lower energy requirement. It is like digging a tunnel through the hill instead of climbing over it.