A reaction intermediate is a species that is produced in one elementary step of a reaction mechanism and consumed in a later step, so it never appears in the overall balanced equation. On the AP Chem exam, you identify intermediates by canceling species that show up on both sides of a multistep mechanism.
A reaction intermediate is a chemical species that gets made partway through a reaction mechanism and then used up before the reaction finishes. It shows up as a product of an early elementary step and a reactant in a later step, which means it cancels out when you add the steps together. That's why intermediates never appear in the overall balanced equation, even though they're very real molecules (or ions, or radicals) that exist briefly during the reaction.
Intermediates are usually high-energy and short-lived, but here's the important distinction for AP Chem. An intermediate is an actual species with real bonds that sits in a small energy valley on the reaction coordinate diagram, between two activation energy peaks. It is not the same thing as a transition state, which sits at the top of a peak and is just a fleeting arrangement of atoms mid-bond-breaking. If you can isolate it (even for a microsecond), it's an intermediate. If it exists only at the instant of collision, it's a transition state.
Reaction intermediates live in Unit 5 (Kinetics), especially Topic 5.11 (Catalysis) under learning objective 5.11.A, which asks you to explain how a catalyst changes a reaction mechanism. Catalyzed mechanisms almost always create new intermediates. Per essential knowledge 5.11.A.3, some catalysts work by binding to a reactant to form an intermediate, opening up a new pathway with a lower activation energy. So you can't fully explain catalysis without being able to spot intermediates in a mechanism.
Intermediates also matter for mechanism analysis across all of Unit 5. To check whether a proposed mechanism is valid, you add up the elementary steps and cancel intermediates to see if you recover the overall equation. And because intermediates don't appear in the overall equation, they can't appear in an experimentally measured rate law either, a detail that trips up a lot of people on rate-law questions.
Keep studying AP Chemistry Unit 5
Catalyst (Unit 5)
Catalysts and intermediates are mirror images in a mechanism. A catalyst is consumed first and regenerated later (it appears in the overall mechanism unchanged), while an intermediate is produced first and consumed later. Per 5.11.A.2, the catalyst's net concentration stays constant; the intermediate's net amount is zero.
Transition State (Unit 5)
On an energy diagram, transition states are the peaks and intermediates are the valleys between peaks. A two-step mechanism has two transition states but only one intermediate. Keeping these straight is one of the most common point-savers in Unit 5.
Rate-Determining Step (Unit 5)
The slow step controls the rate, and if an intermediate appears in that slow step, the simple rate law gets complicated because intermediates can't appear in a measurable rate law. AP questions love mechanisms where the slow step comes first precisely to avoid this issue, so notice the step order.
Activation Energy (Unit 5)
Catalysts lower activation energy by replacing one big energy hill with two or more smaller ones (5.11.A.1). The intermediate is what sits in the dip between those smaller hills. New pathway, new intermediate, lower Ea.
Multiple-choice questions hand you a multistep mechanism and ask you to identify the intermediate. The move is mechanical. Find the species produced in one step and consumed in a later step. For example, in a mechanism like Step 1: A + B → C + D (slow), Step 2: C + E → F + G (fast), Step 3: G → H + B (fast), both C and G are intermediates (made then destroyed), while B is the catalyst (consumed then regenerated). Questions on acid-catalyzed mechanisms, like ester hydrolysis or alcohol dehydration, test whether you can spot the step where the catalyst bonds to a reactant to form an intermediate (5.11.A.3).
No released FRQ has leaned on this term verbatim, but free-response mechanism questions routinely ask you to verify that elementary steps sum to the overall equation (which requires canceling intermediates), sketch or interpret reaction coordinate diagrams (intermediates are the valleys), and explain why an intermediate can't appear in the rate law.
An intermediate is a real, isolable species with fully formed bonds that sits in an energy valley between two activation barriers. A transition state is the unstable, highest-energy arrangement of atoms at the very top of a barrier, with bonds partially broken and partially formed, and it can never be isolated. Quick check on an energy diagram. Peaks are transition states, valleys between peaks are intermediates.
A reaction intermediate is produced in one elementary step and consumed in a later step, so it cancels out and never appears in the overall balanced equation.
Intermediates sit in the energy valleys of a reaction coordinate diagram, while transition states sit at the peaks.
An intermediate is the opposite of a catalyst in a mechanism. The intermediate is made first and used up later, while the catalyst is used up first and regenerated later.
Some catalysts speed up reactions by covalently bonding to a reactant to form an intermediate, creating a new pathway with lower activation energy (5.11.A.3).
Intermediates can never appear in an experimentally determined rate law, because their concentrations aren't directly measurable or controllable.
To validate a proposed mechanism, add the elementary steps and cancel the intermediates; the result must match the overall equation.
It's a species formed in one step of a reaction mechanism and consumed in a later step, so it doesn't show up in the overall equation. You find it by spotting whatever appears as a product in an early step and a reactant in a later one.
No. An intermediate is a real species with complete bonds that sits in an energy valley between two activation barriers, while a transition state is the fleeting highest-energy point at the top of a barrier. A two-step mechanism has one intermediate but two transition states.
Order matters. A catalyst is consumed in an early step and regenerated later (reactant first, product later), while an intermediate is the reverse (product first, reactant later). In the mechanism A + B → C + D, then C + E → F + G, then G → H + B, the species C and G are intermediates and B is the catalyst.
No, not in an experimentally measured rate law. Intermediates don't appear in the overall equation and you can't control their concentrations, so AP rate laws are written only in terms of reactants (and sometimes catalysts).
Often, yes. Per essential knowledge 5.11.A.3, some catalysts bind to a reactant to form a new intermediate, which is exactly how the catalyzed pathway achieves a lower activation energy than the uncatalyzed one.