5.11 Catalysis

AP Chem 5.11 Catalysis Summary

A catalyst speeds up a reaction by giving it a new pathway with lower activation energy or by increasing the number of effective collisions. It changes the reaction mechanism but is not used up overall: it gets consumed in one step and regenerated in a later step, so its net concentration stays the same.

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Why This Matters for the AP Chemistry Exam

Catalysis ties together the big kinetics ideas from this unit: activation energy, collision requirements, reaction mechanisms, and energy profiles. You should be able to explain why a catalyst increases reaction rate by connecting what happens at the particle level to what you observe at the macroscopic level. That means describing how a catalyst lowers activation energy or improves collision effectiveness, identifying a catalyst inside a mechanism, and showing how a catalyst changes a reaction energy profile compared to the uncatalyzed path.

You may need to read or draw an energy profile, compare catalyzed and uncatalyzed mechanisms, and justify a claim about how the catalyst affects the rate. Catalysis also sets up later units, since a catalyst speeds up how fast equilibrium is reached without changing the position of equilibrium or the overall energy change of the reaction.

Key Takeaways

• A catalyst increases rate by lowering activation energy and/or increasing the number of effective collisions.
• A catalyst is consumed in one step of a mechanism and regenerated in a later step, so its net concentration stays constant.
• A catalyst changes the mechanism, often by creating a new reaction intermediate, but it does not change the overall reaction or its overall energy change.
• On an energy profile, a catalyzed path has a lower peak (lower activation energy) than the uncatalyzed path.
• Enzyme, acid-base, and surface catalysis all work by creating new, lower-energy intermediates that the original reaction did not have.

What Is a Catalyst?

A catalyst is a species that speeds up a reaction without being used up overall. It shows up as a reactant in one elementary step and is regenerated as a product in a later step, so its net concentration does not change. A catalyst does not change the overall balanced reaction or the overall energy change. Instead, it changes the path the reaction takes.

You will often see a catalyst written above the reaction arrow rather than as part of the reactants or products. For example, the decomposition of hydrogen peroxide into water and oxygen can be catalyzed by iodide ion:

2 H2O2 -> 2 H2O + O2 (catalyzed by I-)

This reaction is usually done with potassium iodide, but the iodide ion is the actual catalyst. Hydrogen peroxide decomposes slowly on its own, but with the catalyst the reaction speeds up dramatically. This is the basis of the "elephant's toothpaste" demonstration, which is an application of catalysis, not required AP content.

How a Catalyst Changes the Energy Profile

A catalyst works by giving the reaction a new, lower-energy path to follow. Activation energy is the energy difference between the reactants and the transition state (the peak of the curve). A catalyst lowers that peak, so more collisions have enough energy to react.

There are two ways to think about how this happens:

• The catalyst provides a reaction path with a lower activation energy than the original path.
• The catalyst increases the number of effective collisions, for example by holding reactants in a better orientation.

A catalyst can also turn one high-activation-energy step into a multistep path made of lower-activation-energy steps. On an energy profile, the catalyzed path has a lower maximum than the uncatalyzed path, even though both start and end at the same energy levels. The catalyst does not change the energy of the reactants or products, so the overall energy change of the reaction stays the same.

How a Catalyst Changes the Mechanism

A catalyst changes the steps a reaction goes through. Compare a catalyzed and an uncatalyzed mechanism for the same overall reaction and you will usually find that the catalyzed version follows a different set of elementary steps, often with a new reaction intermediate that did not exist in the uncatalyzed path.

For a catalyst to increase the rate, its presence must increase the number of effective collisions and/or provide a path with lower activation energy. One common way catalysts do this is by binding to the reactants. When a catalyst binds a reactant, the reactant can be oriented more favorably or react with a lower activation energy. This often creates a new intermediate in which the catalyst is bound to the reactant. Many enzymes work exactly this way.

Keep this rule in mind: even though a catalyst is frequently consumed in the rate-determining step, the net catalyst concentration stays constant because it is regenerated in a later step.

Acid-Base Catalysis

In acid-base catalysis, a reactant or intermediate gains or loses a proton (H+). This changes the mechanism by introducing a new intermediate and new elementary reactions involving that intermediate.

What to know:

• Proton transfer creates a new chemical species. Gaining H+ is protonation; losing H+ is deprotonation.
• The new intermediate participates in elementary steps that the uncatalyzed reaction does not have.
• The new intermediate is more reactive, which lowers the overall activation energy.
• The acid or base is regenerated, so the catalyst is not used up overall.

Acid-catalyzed reactions in organic and biochemical systems often follow this pattern. The protonated intermediate reacts more easily than the original neutral molecule, which is what speeds things up. Any specific organic mechanism shown here is an example application, not a required AP mechanism to memorize.

Surface Catalysis

In surface catalysis, a reactant or intermediate binds to, or forms a covalent bond with, a solid surface. This is more than just sticking to the surface: the reactant can form an actual chemical bond with surface atoms, creating a new bound intermediate.

What to know:

• Reactants can form covalent bonds with surface atoms, not just adsorb loosely.
• These surface-bound species are new reaction intermediates with different reactivity.
• The mechanism includes new elementary steps involving those bound intermediates.
• The surface atoms are temporarily part of the reacting system and are freed again at the end.

A common application is ammonia synthesis on an iron catalyst, where nitrogen and hydrogen bind to the metal surface, react as surface-bound intermediates, and release the product. These bound intermediates have different activation energies than the free gas molecules, which lets the reaction proceed under more reasonable conditions. The specific industrial process is an example application, not required AP content.

How to Use This on the AP Chemistry Exam

Free Response

• When asked to explain how a catalyst speeds up a reaction, state that it lowers activation energy and/or increases effective collisions, then connect that to a larger fraction of collisions having enough energy to react.
• If asked to justify a claim, link the particle-level change (new intermediate, better orientation, lower barrier) to the macroscopic result (faster rate). Both levels should appear in a strong answer.
• Say clearly that the overall reaction and overall energy change do not change, only the path.

Problem Solving

• To identify a catalyst in a mechanism, look for a species that appears as a reactant in an early step and is regenerated as a product in a later step. Its net concentration is unchanged.
• To tell a catalyst from an intermediate: a catalyst is added at the start and comes back out; an intermediate is made during the reaction and used up before the end.

Energy Profiles

• A catalyzed path is drawn with a lower peak than the uncatalyzed path.
• The reactant and product energy levels stay the same, so the overall energy change does not move.
• If a catalyst creates a multistep path, the profile may show more than one smaller hump instead of one tall hump.

Common Trap

Do not say a catalyst lowers the energy of the reactants or products, and do not say it changes how much product forms at equilibrium. It only changes the activation barrier and the path.

Common Misconceptions

• "A catalyst is used up in the reaction." It is consumed in one step but regenerated in a later step, so its net amount does not change.
• "A catalyst changes the overall energy of the reaction." It lowers activation energy only. The energy of the reactants and products, and the overall energy change, stay the same.
• "A catalyst makes more product form." It speeds up how fast the reaction reaches its endpoint. It does not change the overall reaction or the equilibrium position.
• "A catalyst and an intermediate are the same thing." A catalyst is present at the start and end; an intermediate is formed and then consumed within the mechanism.
• "Surface catalysis is just molecules sticking to a surface." Reactants can form real covalent bonds with the surface, creating new bound intermediates and new elementary steps.
• "Lowering activation energy means the reaction releases more energy." Activation energy is the barrier to reach the transition state, not the overall energy released.

Related AP Chemistry Guides

• Unit 5 Overview: Kinetics
• 5.1 Reaction Rates
• 5.2 Introduction to Rate Law
• 5.3 Concentration Changes Over Time
• 5.4 Elementary Reactions
• 5.5 Collision Model