Collision Theory & Activation Energy

Molecules have to hit hard enough — and line up right. That double requirement explains almost everything about reaction speed.

High schoolIntro Gen ChemUni Year 1
⏱️ About 20 min

Fill a room with reactant molecules and they collide billions of times a second — yet most reactions crawl. If every collision reacted, everything would react instantly. So the puzzle isn't why reactions happen; it's why they mostly DON'T. The answer is a barrier called activation energy.

💡
The big idea: For a collision to react, two things must both be true: the molecules must collide with at least the activation energy (Ea), and they must be oriented correctly. Most collisions fail one test or the other — which is why reactions are far slower than the raw collision rate would suggest.
🎯 By the end, you'll be able to
  • State the two requirements for a successful (reactive) collision
  • Define activation energy Ea as the minimum energy barrier a reaction must clear
  • Read a reaction energy profile, including forward and reverse activation energies
  • Explain why heating speeds reactions mainly by shifting the energy distribution above Ea
📎 Helpful to know first

Reactions happen when molecules collide — usefully

Collision theory pictures a reaction as molecules bumping into one another. But most bumps do nothing. For a collision to actually make product, it must satisfy two conditions at once:

  • Enough energy — the colliding molecules must carry at least the activation energy, so old bonds can break as new ones form.
  • Correct orientation — the molecules must be lined up so the right atoms meet. A glancing hit on the wrong side does nothing, however hard.

Fail either test and the molecules simply bounce apart unchanged.

🔑 Activation energy (Ea)
Activation energy is the minimum energy a collision must supply to reach the transition state — the strained, half-broken arrangement at the top of the barrier. A high Ea means a tall barrier and a slow reaction; a low Ea means an easy barrier and a fast one. Ea is always positive: even 'downhill' exothermic reactions have a hill to climb first.

The reaction energy profile

Plot energy against the progress of the reaction and you get a hill. Reactants sit on the left, products on the right, and the peak in between is the transition state. The climb from reactants to the peak is the forward activation energy; the climb from products up to the same peak is the reverse activation energy.

\[ E_{a,\text{reverse}} = E_{a,\text{forward}} - \Delta H_{\text{rxn}} \]
The peak is the same height for both directions, so the two activation energies differ by the reaction's overall enthalpy change ΔH.

Why heat speeds reactions so much

At any temperature the molecules have a spread of energies — the Maxwell–Boltzmann distribution. Only the molecules out in the high-energy tail, with energy ≥ Ea, can react. Raising the temperature shifts that distribution and, crucially, fattens the tail: a much larger fraction of molecules now clears the barrier.

⚠️ The key subtlety about temperature
It is tempting to say 'heat makes molecules collide more often, so it is faster.' Higher temperature does raise the collision frequency a little — but that is a minor effect. The dominant reason heating speeds a reaction is that it sharply increases the fraction of collisions with energy ≥ Ea (the Boltzmann tail). A modest 10 °C rise can roughly double the rate — far more than the tiny change in how often molecules meet.
📝 Worked example: A forward reaction has activation energy Ea(forward) = 50 kJ/mol and releases energy: ΔH = −30 kJ/mol (exothermic). What is the reverse activation energy?
  1. The transition-state peak is at the same height for both directions.
  2. Ea(reverse) = Ea(forward) − ΔH = 50 − (−30).
  3. = 50 + 30 = 80 kJ/mol. The reverse barrier is taller, which makes sense: the reverse reaction is uphill (endothermic).
✓ Ea(reverse) = 80 kJ/mol.
✏️ Practice: A reaction has Ea(forward) = 45 kJ/mol and ΔH = +20 kJ/mol (endothermic). What is the reverse activation energy in kJ/mol? (Ea,rev = Ea,fwd − ΔH.)
kJ/mol
Solution
  1. Ea(reverse) = Ea(forward) − ΔH.
  2. = 45 − (+20).
  3. = 25 kJ/mol. The reverse barrier is lower because the forward reaction was uphill.

Check your understanding

1. Raising the temperature speeds a reaction MAINLY because it…
Heating shifts the Maxwell–Boltzmann distribution so a much larger fraction of molecules exceed Ea. The rise in collision frequency is real but minor; temperature does not change Ea itself.
2. Two molecules collide with plenty of energy but the wrong way round. The collision…
A successful collision needs BOTH sufficient energy AND correct orientation. Energy alone is not enough if the molecules are poorly aligned.
3. For an exothermic reaction, how do the forward and reverse activation energies compare?
Both climb to the same peak. For an exothermic reaction the products sit lower than the reactants, so the reverse climb (and thus reverse Ea) is larger.
🔎 Related lessons
Previous: Integrated Rate Laws & Half-Life Next: The Arrhenius Equation Catalysis
✅ Key takeaways
  • Collision theory: a reaction needs collisions, but most collisions do not react.
  • A successful collision needs BOTH energy ≥ Ea AND correct orientation.
  • Activation energy Ea is the barrier to the transition state; higher Ea → slower reaction.
  • On an energy profile, Ea(reverse) = Ea(forward) − ΔH.
  • Heating speeds reactions mainly by enlarging the fraction of collisions above Ea, not by more frequent collisions.
➡️ We now have the ingredients — a barrier Ea, an orientation requirement, and a temperature-sensitive energy distribution. The Arrhenius equation packages all of them into one formula that predicts exactly how k depends on Ea and temperature.
Want to test yourself on this? Try the Chemistry practice test →