Activation Energy and Catalysts
A paraffin candle can sit in a drawer for years surrounded by oxygen and nothing happens, even though burning releases a lot of energy. But bring a match close and it burns on its own until it is used up. Why is that initial "push" needed if the reaction releases energy in the end? And in your body, how do reactions happen at 37 °C that in the lab would need a lot of heating? The answer lies in activation energy and catalysts.
The Energy Barrier
For reactants to turn into products, bonds first have to be broken (or weakened). That costs energy. Afterward, when the new bonds form, energy is released. But the initial "bill" has to be paid before you receive what comes next.
Imagine pushing a cart to the top of a hill so that it can then roll down on its own. The climb is the activation energy; the descent is the energy that is released.
This connects to collision theory: only collisions with enough energy are effective. If $E_a$ is large, few collisions reach it and the reaction is slow; if it is small, the reaction is fast.
The graph shows the energy (in kJ) along the reaction coordinate for an exothermic reaction. From it: a) calculate the activation energy $E_a$ of the forward reaction;…
The graph shows the energy (in kJ) along the reaction coordinate for an exothermic reaction. From it: a) calculate the activation energy $E_a$ of the forward reaction; b) calculate the $\Delta H$ of the reaction; c) explain what would change in the graph if a catalyst were added.
The Energy Profile
An energy profile (or energy diagram) is a graph of the energy of the substances versus the reaction coordinate, which represents the progress from reactants to products. The highest point is the transition state (or activated complex): an unstable grouping where the old bonds are breaking and the new ones are forming.
How to read it:
-
$E_a$ (forward) = energy of the transition state − energy of the reactants.
-
$\Delta H$ = energy of the products − energy of the reactants.
-
If the products end up lower than the reactants, $\Delta H<0$: exothermic reaction. If they end up higher, $\Delta H>0$: endothermic.
What a Catalyst Does
On the profile, the catalyst lowers the peak. Notice what changes and what doesn't:
| Changes with a catalyst | Does not change |
|---|---|
| $E_a$ goes down (forward and reverse) | $\Delta H$ |
| More collisions are effective: the reaction is faster | Energy of reactants and products |
| Other transition states form | Amount of product that can be obtained |
Why $\Delta H$ doesn't change: the energy of the reactants and products is a property of the substances themselves. The catalyst only changes the path, not the starting point or the end point. And since it lowers the barrier in both directions, it speeds up the forward and reverse reactions equally.
Explanation with Particles
At a given temperature, particles have different kinetic energies: many slow ones and a few very fast ones. Only collisions with energy greater than or equal to $E_a$ are effective.
-
Raising the temperature makes more particles exceed $E_a$ and makes them collide more often.
-
A catalyst lowers $E_a$, so, at the same temperature, a larger fraction of collisions already exceeds it.
This explains why refrigeration preserves food (decomposition reactions are slower) and why enzymes allow reactions at 37 °C.
Solving a Complete Problem
Go back to the beginning: the candle doesn't burn on its own because it needs to overcome its $E_a$; the match provides that energy and, since the reaction is exothermic, the energy it releases keeps the combustion going. In the body, enzymes lower $E_a$ and allow fast reactions at 37 °C.