A catalyst is a substance that increases the rate of a chemical reaction by providing an alternate reaction pathway with a lower activation energy; it is consumed in an early step of the mechanism but regenerated later, so its net concentration stays constant (AP Chem 5.11.A.1-2).
A catalyst speeds up a reaction without getting used up overall. The key word is overall. In a real mechanism, the catalyst is usually consumed in an early elementary step (often the rate-determining one) and then regenerated in a later step, so it walks out of the reaction unchanged (EK 5.11.A.2). That's why catalysts appear in the elementary steps of a mechanism but cancel out of the overall balanced equation.
How does it actually make things faster? Per the CED, a catalyst must do one (or both) of two things: increase the number of effective collisions, or provide a new reaction pathway with a lower activation energy than the uncatalyzed route (EK 5.11.A.1). Picture activation energy as a hill between reactants and products. A catalyst doesn't push the reaction harder up the hill; it tunnels through a shorter one. Some catalysts, like enzymes, do this by binding to reactants and holding them in a favorable orientation. Crucially, a catalyst changes the speed of reaching equilibrium, never the position of equilibrium, and never the thermodynamics (ΔG, ΔH, K all stay the same).
Catalysts are the centerpiece of Topic 5.11 (Catalysis) and learning objective AP Chem 5.11.A, which asks you to explain how a catalyst changes the reaction mechanism. But the term shows up across three units. In Unit 5, it's part of what controls reaction rates (AP Chem 5.1.A) and a named component of mechanisms (AP Chem 5.7.A.1), so you have to spot catalysts in a list of elementary steps. In Unit 9, catalysts are the answer to a classic AP question type built on AP Chem 9.4.A: a reaction is thermodynamically favored (ΔG < 0) but doesn't happen at a measurable rate because activation energy is too high. The fix is kinetic (add a catalyst), not thermodynamic. In Unit 7, the trap runs the other way: a catalyst is NOT a stress under Le Châtelier's principle (AP Chem 7.9.A), so adding one never shifts equilibrium.
Keep studying AP Chemistry Unit 5
Activation Energy (Unit 5)
Activation energy is the lever a catalyst pulls. The catalyzed pathway has a lower Ea, which means more molecular collisions have enough energy to react. On an energy diagram, the catalyzed curve has a lower peak (or several smaller peaks for a multistep mechanism), but the same starting and ending energies. ΔH doesn't budge.
Elementary Step (Unit 5)
On the exam, you identify a catalyst from a mechanism by tracking species across elementary steps. A catalyst is consumed first and regenerated later. An intermediate does the opposite, produced first and consumed later. Same cancellation trick, mirror-image order.
Thermodynamic and Kinetic Control (Unit 9)
Hydrogen peroxide decomposition is thermodynamically favored (ΔG° < 0), yet a bottle of H₂O₂ sits stable for months because the activation energy is high. The reaction is under kinetic control. Add a catalyst and it fizzes immediately. This Unit 5 plus Unit 9 crossover is one of the most-tested catalyst scenarios.
Le Châtelier's Principle (Unit 7)
Adding a catalyst is the classic 'no effect' answer choice. A catalyst speeds up the forward and reverse reactions equally, so the system reaches equilibrium faster but lands at exactly the same K and the same equilibrium concentrations.
Catalysts show up in three predictable ways. First, mechanism questions hand you a list of elementary steps and ask you to identify the catalyst (consumed then regenerated) versus the intermediate (produced then consumed), or to explain mechanistically why the catalyzed path is faster. Second, kinetic-control questions describe a thermodynamically favorable reaction that runs too slowly, like H₂O₂ stored for months without decomposing, and the correct answer involves lowering activation energy with a catalyst rather than changing ΔG. Third, equilibrium questions test whether you know a catalyst does NOT shift equilibrium or change K. Catalysts also appear as context in released FRQs, like the platinum catalyst in the 2022 methanol decomposition problem and the iron hand-warmer reaction in 2021 Q4, where you're expected to reason about rates and thermodynamics around the catalyzed system. The verbs to be ready for are identify (from a mechanism) and explain (in terms of activation energy or effective collisions).
Both catalysts and intermediates appear in elementary steps but cancel out of the overall equation, which is exactly why the AP exam loves making you tell them apart. The difference is order of appearance. A catalyst goes in first (consumed in an early step) and comes back out (regenerated later), so it's present before, during, and after the reaction. An intermediate is produced by one step and consumed by a later one, so it only exists while the reaction is running (EK 5.7.A.3). Quick check on any mechanism: consumed-then-made means catalyst, made-then-consumed means intermediate.
A catalyst speeds up a reaction by providing a lower-activation-energy pathway or increasing effective collisions, not by changing the thermodynamics.
In a mechanism, a catalyst is consumed in an early step and regenerated in a later step, so its net concentration is constant; an intermediate is the reverse (made first, consumed later).
A catalyst speeds up the forward and reverse reactions equally, so it changes how fast equilibrium is reached but never shifts the equilibrium position or changes K.
A catalyst does not change ΔG, ΔH, or the energy of reactants and products; on an energy diagram, only the peak (Ea) gets lower.
If a reaction is thermodynamically favored (ΔG < 0) but happens too slowly to measure, it's under kinetic control, and adding a catalyst is the way to make it usable.
A catalyst is a substance that increases a reaction's rate by lowering the activation energy or increasing effective collisions, while its net concentration stays constant. It's consumed in an early step of the mechanism and regenerated in a later one (EK 5.11.A.1-2).
No. A catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached faster but the equilibrium concentrations and K are unchanged. 'Add a catalyst' is the classic no-shift answer in Le Châtelier questions (Topic 7.9).
A catalyst is consumed in an early elementary step and regenerated later, so it exists before and after the reaction. An intermediate is produced by one step and consumed by a later one, so it only exists while the reaction is happening. Track the order of appearance in the mechanism to tell them apart.
No. A catalyst only lowers the activation energy, the hill between reactants and products. The energies of reactants and products themselves don't move, so ΔH, ΔG, and thermodynamic favorability are all unchanged.
Because the decomposition has a high activation energy, the reaction is under kinetic control and proceeds at an immeasurably slow rate even though ΔG° < 0 (EK 9.4.A.2). Add a catalyst and the decomposition to water and oxygen happens fast.