reaction intermediate
When you fold a paper airplane, there are half-folded shapes that exist only briefly — not a flat sheet anymore, not yet a plane. They are real objects you could photograph mid-fold, but they are nobody's goal; they appear and then vanish into the next fold. A reaction intermediate is chemistry's version of such an in-between species: a real molecule or fragment that is made in one step of a reaction and used up in a later step.
More precisely, a reaction intermediate is a species formed during a reaction mechanism that does not appear in the overall balanced equation, because it is produced by one elementary step and consumed by another. Unlike the starting materials and final products, it is usually present only in tiny amounts and for a short time. Common examples include free radicals, carbocations, and small fragments like single atoms. Because they are genuine molecules, intermediates can sometimes be detected by fast spectroscopy or trapped by clever experiments.
Why it matters: spotting an intermediate is strong evidence for a proposed mechanism, since the overall equation gives no hint that it exists. The key point of confusion to avoid is the difference between an intermediate and a transition state: an intermediate sits in a small valley on the energy landscape and has a real, if fleeting, existence, whereas a transition state sits at the very top of a barrier and cannot be isolated at all.
In the two-step decomposition of ozone, the first step splits an O3 into an O2 and a lone oxygen atom. That free O atom is the intermediate: it is born in step one and immediately consumed in step two, so it never appears in the overall equation 2 O3 → 3 O2.
Made in one step, used up in the next — so it is absent from the overall equation.
Do not confuse an intermediate with a catalyst: a catalyst is added at the start and recovered at the end, whereas an intermediate is created inside the reaction and consumed inside it.