Catalysis

A shortcut through the energy barrier β€” speeding a reaction without being used up, and without changing where it ends.

High schoolIntro Gen ChemUni Year 1
⏱️ About 14 min

The catalytic converter under your car takes toxic exhaust gases and turns them harmless in a fraction of a second β€” using a metal that is never consumed. Your own body runs thousands of such tricks every moment, with enzymes. That is catalysis: faster reactions, for free.

πŸ’‘
The big idea: A catalyst speeds a reaction by providing an alternative pathway with a lower activation energy. It is regenerated (not consumed) and it speeds the forward and reverse reactions equally, so it never shifts the equilibrium position.
🎯 By the end, you'll be able to
  • Explain how a catalyst lowers activation energy without being consumed
  • State why a catalyst speeds forward and reverse reactions equally and does not shift equilibrium
  • Distinguish homogeneous, heterogeneous and enzyme (biological) catalysts
  • Connect a lower Ea to a larger rate constant through the Arrhenius equation
πŸ“Ž Helpful to know first

A lower road over the mountain

A catalyst is a substance that speeds up a reaction without being used up in the process. It works not by pushing molecules harder but by opening a different route β€” a mechanism with a lower activation energy. Same reactants, same products, but a smaller hill to climb.

Through the Arrhenius equation, a lower Ea means a bigger fraction of molecules can react at a given temperature, so k rises and the reaction accelerates β€” often by orders of magnitude.

\[ k = A\,e^{-E_a / RT} \;\;\xrightarrow{\;E_a \downarrow\;}\;\; k \uparrow \]
Lowering Ea raises the exponential term, so the rate constant β€” and the rate β€” go up.
πŸ”‘ A catalyst is NOT consumed
A catalyst takes part in the reaction β€” it may bond to a reactant partway through β€” but it is regenerated by the end, emerging unchanged and ready to work again. That is why a tiny amount can process an enormous quantity of reactant, and why a catalyst never appears in the overall balanced equation.
⚠️ A catalyst does NOT shift equilibrium
This is the crucial misconception to avoid. A catalyst lowers the barrier for the forward and the reverse reaction by the same amount, so it speeds both directions equally. It helps the system reach equilibrium sooner, but the equilibrium position β€” and the value of K β€” are unchanged. A catalyst changes kinetics, never thermodynamics: it does not raise the final yield.

It changes the path, not the endpoints

On the energy profile, a catalyst carves a lower peak between the same reactants and products. Because the reactant and product energies are untouched, the reaction's Ξ”H is unchanged β€” only the barrier between them is lower. Faster to get there; same place you arrive.

Three flavours of catalyst

  • Homogeneous β€” in the same phase as the reactants (e.g. an acid dissolved in the reacting solution).
  • Heterogeneous β€” in a different phase, usually a solid surface that gases or liquids react on. The metals in a catalytic converter are heterogeneous catalysts.
  • Enzymes β€” nature's protein catalysts, astonishingly fast and specific; they run the biochemistry of every living cell.
πŸ“ Worked example: A catalyst lowers a reaction's activation energy from 60 kJ/mol to 35 kJ/mol. What happens to the reaction's Ξ”H (enthalpy change)?
  1. A catalyst provides a new pathway with a lower barrier, but it does not change the energies of the reactants or the products.
  2. Ξ”H depends only on the reactant and product energies (their difference), not on the height of the barrier between them.
  3. So Ξ”H is completely unchanged β€” the change in Ξ”H is 0.
βœ“ Ξ”H is unchanged (change = 0). Only the activation energy is lowered.
✏️ Practice: A catalyst lowers the activation energy of a reaction from 80 kJ/mol to 55 kJ/mol. By how many kJ/mol does the reaction's Ξ”H change?
kJ/mol
Solution
  1. A catalyst lowers the barrier but leaves the reactant and product energies alone.
  2. Ξ”H = (product energy) βˆ’ (reactant energy), neither of which the catalyst touches.
  3. So Ξ”H changes by 0 kJ/mol β€” a catalyst speeds the reaction, it does not make it more exothermic or shift equilibrium.

Check your understanding

1. How does a catalyst speed up a reaction?
A catalyst opens an alternative route with a lower Ea, so more collisions succeed and k rises. It does not change temperature, concentration, or the equilibrium position.
2. What happens to a catalyst over the course of the reaction?
A catalyst is not consumed. It may participate mid-reaction but is regenerated by the end, which is why a small amount can catalyse a large amount of reaction.
3. Adding a catalyst to a reaction at equilibrium will…
A catalyst lowers the forward and reverse barriers equally, so equilibrium is reached faster but the position (and K, and the final yield) are unchanged. Rate β‰  yield.
πŸ”Ž Related lessons
Previous: The Arrhenius Equation Collision Theory & Activation Energy Dynamic Equilibrium & K
βœ… Key takeaways
  • A catalyst speeds a reaction by providing a lower-activation-energy pathway.
  • It is not consumed β€” it is regenerated and never appears in the overall equation.
  • It speeds forward and reverse equally, so it does NOT shift equilibrium or change K.
  • It leaves Ξ”H unchanged: catalysts change kinetics, not thermodynamics.
  • Types: homogeneous (same phase), heterogeneous (surface), and enzymes (biological).
➑️ A catalyst gets a system to equilibrium faster without moving where that balance lies. So what DOES move the balance β€” and how do we describe it precisely? That is the story of chemical equilibrium, the next module.
Want to test yourself on this? Try the Chemistry practice test β†’