Topic 5.11 Notes – Catalysis

1. What a Catalyst Is

A catalyst is a substance that increases the rate of a reaction by changing the mechanism.

Here’s what that means in practice:

• It provides a pathway with a lower activation energy, $E_a$.
• It is used in one step of the mechanism.
• It is regenerated in a later step.
• Its net concentration stays constant.
• It does not change:
  • $\Delta H$
  • $\Delta G$
  • The equilibrium constant $K$
  • The equilibrium position

So equilibrium is reached faster, but it’s the same equilibrium.

If you add all elementary steps together and the catalyst cancels out, that’s your confirmation it’s truly a catalyst.

The big picture:
Change the mechanism → change the activation energy → change the rate.

2. How Catalysts Change the Reaction Pathway

When a catalyst is added, you’re not just “speeding up” the same reaction. You are introducing a new mechanism.

Lower Activation Energy

On a reaction coordinate diagram:

• Reactant energy stays the same.
• Product energy stays the same.
• The highest peak (activation energy) becomes lower.
• Sometimes the reaction now has multiple peaks because there are new intermediates.

Here’s what that looks like on a reaction coordinate diagram comparing uncatalyzed and catalyzed pathways:

Reaction coordinate diagram: catalyzed vs. uncatalyzed pathways

The key is that the highest barrier is reduced.

Sometimes one big step becomes two smaller steps. Even if there are more steps, the tallest hill is shorter, so the reaction goes faster.

Effective Collisions

For rate to increase, one (or both) must happen:

• More effective collisions
• Lower minimum energy required for reaction

Catalysts help by:

• Holding reactants in the right orientation
• Stabilizing the transition state
• Forming reactive intermediates

On AP questions, if they ask why the rate increases, mention:

• Lower $E_a$
• More effective collisions
• New mechanism

Those phrases matter.

3. What Happens to the Catalyst in a Mechanism

In a catalyzed mechanism, look for this pattern:

• Appears as a reactant early
• Appears as a product later
• Does not appear in the overall balanced equation

Example pattern:

Step 1: $\text{A} + \text{Cat} \rightarrow \text{Intermediate}$
Step 2: $\text{Intermediate} \rightarrow \text{B} + \text{Cat}$

The catalyst is often involved in the rate-determining step, then regenerated afterward.

Students sometimes think “it was used up, so it’s not a catalyst.” It’s allowed to be temporarily consumed. What matters is the net concentration stays constant.

If it never comes back, it wasn’t a catalyst.

4. Types of Catalytic Mechanisms

Different catalysts lower $E_a$ in different ways.

A. Binding Catalysis (including enzymes)

The catalyst binds to the reactant(s) and forms a new intermediate.

Pattern:

1. Catalyst + Reactant → Bound complex
2. Complex → Product + Catalyst

Effects:

• Better orientation
• Stabilized transition state
• Lower activation energy

Enzymes work this way. The enzyme-substrate complex is the new intermediate.

B. Acid-Base Catalysis

Here the catalyst transfers a proton (H⁺).

• A reactant may become protonated (more electrophilic).
• Or deprotonated (more nucleophilic).

This creates:

• A new intermediate
• New elementary steps
• A lower activation energy

Important detail:
The proton donor or acceptor is regenerated at the end.

If you see H⁺ appear in one step and reappear later, that’s acid catalysis.

C. Surface Catalysis (Heterogeneous Catalysis)

This happens on a solid surface.

• Reactants bind to the surface.
• They may form covalent bonds with surface atoms.
• New surface-bound intermediates form.
• New elementary steps involve those bound species.

Common in industrial reactions using metal catalysts.

Why this lowers $E_a$:

• Weakens existing bonds
• Brings reactants close together
• Provides a new bond-breaking pathway

On exams, if they mention “adsorbed to surface” or “bound to metal,” think surface catalysis.