Catalysis

Core Concept

A catalyst is a substance that speeds up a reaction by providing an alternative reaction pathway with a lower activation energy, while being chemically unchanged at the end. Because Ea appears in the exponent of the Arrhenius equation, lowering Ea dramatically increases the fraction of effective collisions and therefore the rate. Crucially, a catalyst lowers the barrier for both the forward and reverse reactions equally, so it does not change the thermodynamics (ΔG°) or the equilibrium constant (K).

Key Ideas

• Catalysts work by lowering Ea, not by raising temperature or concentration.
• Homogeneous catalysis: catalyst in the same phase as reactants (often dissolved).
• Heterogeneous catalysis: catalyst in a different phase, usually a solid surface for gas/liquid reactants.
• Enzymes are highly specific biological catalysts that stabilize the transition state.
• Catalysts are regenerated — they are not consumed in the net reaction.

Equations and Variables

The connection to the Arrhenius equation:

k = A e^-E_a/RT

Lowering Ea (smaller barrier) increases e^(−Ea/RT) and thus k. The key thermodynamic constraint:

ΔG^° = -RTlnK

Since a catalyst changes the pathway but not the free-energy difference between reactants and products, ΔG° — and therefore K — are unaffected.

How It Works

A catalyst opens a new route with a lower energy hill. For a solid metal catalyst (heterogeneous), reactant molecules adsorb onto the surface, where they are held in favorable orientations and weakened by bonding to the surface, effectively lowering the barrier to reaction; products then desorb. For a homogeneous catalyst, the catalyst enters an intermediate compound with the reactants and is regenerated later.

The Haber process (N₂ + 3 H₂ → 2 NH₃) uses an iron catalyst to break the very strong N≡N bond more easily; catalytic converters use platinum/rhodium/palladium surfaces to oxidize CO and unburned hydrocarbons and reduce NOₓ. Enzymes, such as those in the liver, accelerate specific biological reactions millions of times by binding substrates in a precise active-site geometry that stabilizes the transition state.

Because the catalyst lowers Ea equally in both directions, it reaches equilibrium faster but arrives at the same equilibrium mixture. This is why a catalyst cannot make a thermodynamically unfavorable reaction favorable — it can only get you to equilibrium sooner.

Worked Example

A reaction has Ea = 80 kJ/mol without a catalyst and Ea = 50 kJ/mol with one. Compare the Boltzmann factors at 300 K to estimate how much faster the catalyzed path is.

Uncatalyzed:

E_aRT = 80,000(8.314)(300) = 32.1   ⇒  e^-32.1 = 1.1 × 10^-14

Catalyzed:

E_aRT = 50,000(8.314)(300) = 20.0   ⇒  e^-20.0 = 2.1 × 10^-9

The fraction of productive collisions rises by about 2.1 × 10⁻⁹ / 1.1 × 10⁻¹⁴ ≈ 190,000-fold — a dramatic speedup from lowering Ea by only 30 kJ/mol. (In practice, the orientation factor and A may also change, but the exponential dominates.) The equilibrium constant, meanwhile, is unchanged because ΔG° is the same with or without the catalyst.

How it works

A catalyst opens a new route with a lower energy hill. For a solid metal catalyst (heterogeneous), reactant molecules adsorb onto the surface, where they are held in favorable orientations and weakened by bonding to the surface, effectively lowering the barrier to reaction; products then desorb. For a homogeneous catalyst, the catalyst enters an intermediate compound with the reactants and is regenerated later.

The Haber process (N₂ + 3 H₂ → 2 NH₃) uses an iron catalyst to break the very strong N≡N bond more easily; catalytic converters use platinum/rhodium/palladium surfaces to oxidize CO and unburned hydrocarbons and reduce NOₓ. Enzymes, such as those in the liver, accelerate specific biological reactions millions of times by binding substrates in a precise active-site geometry that stabilizes the transition state.

Because the catalyst lowers Ea equally in both directions, it reaches equilibrium faster but arrives at the same equilibrium mixture. This is why a catalyst cannot make a thermodynamically unfavorable reaction favorable — it can only get you to equilibrium sooner.

Common confusions

• Thinking a catalyst changes K or ΔG°. It does not — it only lowers Ea and raises the rate.
• Confusing a catalyst with a reactant. A catalyst is regenerated and appears in neither the overall equation nor the equilibrium constant expression.
• Assuming "faster" means "more product." At equilibrium, the yield is the same; the catalyst just gets you there sooner.
• Forgetting heterogeneous surface effects. Solid catalysts work by adsorption, orientation, and surface bonding, not by dissolving into the reaction.

Quick review

1. How does a catalyst increase a reaction's rate?
2. What is the difference between homogeneous and heterogeneous catalysis?
3. Give one industrial and one biological example of catalysis.
4. Why does a catalyst not change the equilibrium constant?
5. In the Arrhenius equation, which quantity does a catalyst alter?