Activation energy is the minimum energy that colliding particles must have for the collision to lead to a reaction. On an energy profile it is the vertical distance from the reactant level up to the peak.
It follows on from interpreting an energy profile and sits in the same module, energy changes and bonds.
What is activation energy actually describing?
Particles are always colliding, but most collisions do nothing. To react, the colliding particles need enough energy to start breaking the bonds in the reactants. The required minimum is the activation energy.
If the particles have less than that energy, they bounce apart with the same chemical identity. If they have at least that much, the collision can succeed. This is why the same reactants can sit together without reacting until they are given an initial energy input, even when the overall reaction is exothermic.
How do you mark and use it on a diagram?
- Draw or find the reactant level on the left.
- Find the peak of the curve.
- Draw a vertical arrow from the reactant level up to the peak and label it activation energy, Ea.
- Draw a different arrow, from reactants to products, for the overall energy change ΔH.
- Say what each arrow means in words: one is a barrier, one is the net result.
Worked example
The numbers are invented for practice. A fictional reaction has reactants at 100 kJ/mol, a peak at 250 kJ/mol and products at 40 kJ/mol. With a catalyst, the peak drops to 190 kJ/mol.
Step 1, activation energy without a catalyst: 250 − 100 = 150 kJ/mol.
Step 2, energy change: 40 − 100 = −60 kJ/mol, so the reaction is exothermic.
Step 3, activation energy with the catalyst: 190 − 100 = 90 kJ/mol. The barrier is lower by 150 − 90 = 60 kJ/mol.
Step 4, energy change with the catalyst: the reactants and products are at the same levels, so ΔH is still −60 kJ/mol.
Step 5, write the explanation: “The catalyst provides an alternative pathway with a lower activation energy, so more collisions have enough energy to react. The overall energy change is the same because the energy levels of the reactants and products have not changed.”
Check: the catalyst changes the peak only. The two flat lines and the ΔH arrow stay where they were.
The mistake to watch for
The typical slip is to confuse the barrier with the overall energy change.
Mistaken answer: “The catalyst makes the reaction more exothermic because the peak is lower.”
The student linked the lower peak to the amount of energy released. They are different quantities.
The correction is to keep the two arrows separate. The peak controls how hard it is to get started; the final levels control how much energy is released overall. A catalyst touches the first and not the second.
For the collision-theory link to reaction speed, see describing a catalyst’s role without claiming it is used up in the rates module.
Check yourself
1. Reactants are at 60 kJ/mol, the peak is at 200 kJ/mol and products are at 10 kJ/mol. Find the activation energy and ΔH.
Show answer
Ea = 200 − 60 = 140 kJ/mol. ΔH = 10 − 60 = −50 kJ/mol (exothermic).
2. A catalyst lowers the peak in question 1 to 150 kJ/mol. State the new activation energy and the new ΔH.
Show answer
New Ea = 150 − 60 = 90 kJ/mol. ΔH stays −50 kJ/mol, because the reactant and product levels have not changed.
3. Explain in one sentence why some exothermic reactions need a starting input of energy.
Show answer
Bond breaking in the reactants needs energy before the new bonds form, so the particles must reach the activation energy even though more energy is released overall.
Where this leads next
The next lesson puts numbers on the bond breaking and making behind the peak: estimate an energy change using supplied bond energies. You can also revisit the collision ideas in the rates module.
Students often describe the catalyst correctly and still draw the profile wrongly, or the other way round. A teacher in online one-to-one Chemistry tuition can see which half of the link is unsteady and work on that.