Gibbs free energy (ΔG) helps explain why some reactions need catalysts, but it’s important to note:
- ΔG determines whether a reaction is spontaneous (can happen on its own).
- Catalysts do not change ΔG — they cannot make a non-spontaneous reaction spontaneous.
- Catalysts only speed up the reaction by lowering the activation energy, the energy barrier that must be overcome for reactants to turn into products.
How it works:
- Spontaneous but slow reactions
- Some reactions have a negative ΔG (thermodynamically favorable) but happen very slowly because the reactants need a lot of energy to start reacting.
- Example: Conversion of diamond to graphite
- ΔG is negative → graphite is more stable.
- But it’s extremely slow at room temperature because of a high activation energy.
- A catalyst can speed it up without changing ΔG.
- Industrial processes
- In the Haber process, forming ammonia is spontaneous under certain conditions (ΔG < 0), but it is slow.
- Iron catalyst lowers the activation energy → reaction reaches equilibrium faster.
Key points:
- ΔG tells us if a reaction can happen.
- Activation energy determines how fast it happens.
- Catalysts help reactions occur faster but do not affect spontaneity (ΔG).
In short: Gibbs free energy explains whether a reaction is possible, and catalysts help it happen more quickly by lowering the energy barrier.