Gibbs free energy (ΔG) is crucial for predicting chemical equilibrium because it tells us whether a reaction will proceed or has reached balance. Here’s a explanation:
1. Relationship between ΔG and equilibrium
- ΔG < 0 → reaction is spontaneous, moves toward products.
- ΔG > 0 → reaction is non-spontaneous, moves toward reactants.
- ΔG = 0 → reaction is at equilibrium, no net change occurs.
2. Standard free energy change (ΔG°) and equilibrium constant (K)
- ΔG° is related to the equilibrium constant K of a reaction:
- Large negative ΔG° → K is large → products are favored.
- Large positive ΔG° → K is small → reactants are favored.
- This allows chemists to predict the position of equilibrium without running the experiment.
3. Adjusting conditions
- ΔG also depends on temperature, pressure, and concentration:
- Changing these can shift ΔG toward negative → reaction moves toward products.
- At equilibrium, ΔG = 0 → system has minimized its free energy.
4. Example
- Ammonia synthesis (Haber process):
- At certain temperature and pressure, ΔG < 0 → reaction favors ammonia formation.
- At higher temperature, ΔG becomes less negative → equilibrium shifts back toward reactants.
Summary
- Gibbs free energy predicts reaction spontaneity.
- When ΔG reaches zero, the reaction is at equilibrium.
- By calculating ΔG, chemists and engineers can determine the direction of reaction and optimize conditions for maximum yield.