What Happens in the Salicylic Acid and Acetic Anhydride Reaction

Stir a white powder into a colourless liquid, add a few drops of strong acid, warm it for ten minutes, and you have aspirin, plus a lab that smells faintly of vinegar. The salicylic acid and acetic anhydride reaction is one of the most familiar syntheses in chemistry teaching. It’s also a compact lesson in how esters form, why chemists choose some reagents over others, and where the leftover smells and gunk come from.

A safety note before the chemistry: acetic anhydride is corrosive, flammable and a lachrymator (its vapour makes eyes stream), and concentrated sulfuric acid is corrosive, according to PubChem’s hazard data. This reaction belongs in a fume hood with eye protection and gloves, under a supervised lab procedure.

What the salicylic acid and acetic anhydride reaction makes

Salicylic acid is a benzene ring carrying two groups side by side: a carboxylic acid (–COOH) and a phenol, an –OH attached directly to the ring. Acetic anhydride is two acetic acid units joined through a shared oxygen. When the two react, an acetyl group (CH₃CO–) moves from the anhydride onto the phenol oxygen. The products are acetylsalicylic acid, which is aspirin, and one molecule of acetic acid. Odinity’s aspirin synthesis and analysis lab writes it exactly that way: salicylic acid plus acetic anhydride, with phosphoric acid as the catalyst, gives aspirin and acetic acid.

Two names describe the same change. It’s an esterification, because the phenol oxygen ends up as part of an ester. It’s also an acetylation, because what’s added is specifically an acetyl group. The carboxylic acid group stays as it is. A West Virginia University lab handout for this experiment asks students to explain why the phenol –OH, not the carboxylic acid –OH, is the one that gets acetylated. A LibreTexts chemistry page notes that industry runs essentially the same reaction, making aspirin from salicylic acid and acetic anhydride at 90 °C.

The mechanism, step by step

Chemists call this a nucleophilic acyl substitution. An electron-rich atom (the nucleophile) attacks the carbon of a C=O group, and something else leaves. The West Virginia handout lays out the acid-catalysed version in five moves:

1. The acid catalyst protonates acetic anhydride, giving an “activated” carbonyl.
2. A lone pair on salicylic acid’s phenol oxygen attacks that carbonyl carbon.
3. A proton shifts from the attacking oxygen to another oxygen in the intermediate.
4. Acetic acid is pushed out as the leaving group.
5. The protonated ester loses its extra proton to a weak base, giving aspirin and regenerating the acid catalyst.

LibreTexts’ organic chemistry textbook fills in step 2. The nucleophile adds to the carbonyl carbon and forms a short-lived tetrahedral intermediate. When the C=O double bond re-forms, a carboxylate group is expelled. With acetic anhydride, that leaving group ends up as acetic acid.

Why an acid catalyst is used

Protonating the anhydride’s carbonyl oxygen pulls electron density away from the neighbouring carbon. That makes the carbon a better target for the phenol oxygen. The catalyst is handed back at the end, so a few drops are enough. The University of Missouri–St. Louis procedure on LibreTexts uses 5 drops of concentrated sulfuric acid for 2 grams of salicylic acid. The West Virginia version uses 3 drops of 85% phosphoric acid.

Acid isn’t the only option. One student lab report on Odinity, the synthesis of aspirin catalyst study, ran the same reaction four ways and timed how long each mixture took to warm by 4 °C. With sulfuric acid it took 5.5 seconds. The ranking from fastest to slowest was sulfuric acid, boron trifluoride etherate, pyridine, then sodium acetate. The report explains that bases work differently, by pulling a proton off salicylic acid to make a stronger nucleophile. Acids, it argues, help at several steps.

That ranking deserves a caveat, and the report gives one itself. A temperature rise is an indirect way to measure rate, since heats of mixing and side reactions also change the temperature. What it does show is that, with sulfuric acid, the mixture warmed measurably within seconds. The report attributes that mainly to the exothermic acetylation, though mixing and side reactions add heat too.

Illustration of needle-like crystals forming on glass (artist’s impression).

Why acetic anhydride and not acetic acid

On paper, acetic acid could donate the same acetyl group. That’s the classic Fischer esterification: a carboxylic acid plus an alcohol, heated with an acid catalyst. Odinity’s Fischer esterification lab shows the problem. Water forms as a by-product, and as it builds up, the reverse reaction (hydrolysis of the ester) speeds up until the mixture reaches equilibrium. LibreTexts’ textbook gives the reason it’s so reversible: carboxylic acids are about as reactive as esters, so neither direction is strongly favoured.

Phenols make it worse. A LibreTexts page on preparing esters states that phenols react with carboxylic acids “so slowly that the reaction is unusable for preparation purposes.” A Vanderbilt University organic chemistry lecture says the same in fewer words: the Fischer esterification “works poorly” for phenyl esters.

Acetic anhydride gets around both problems. It’s much more reactive than acetic acid, because its leaving group (acetate) departs far more easily than the –OH of an acid. And its co-product is acetic acid rather than water, so no water builds up during the reaction to push the ester back toward hydrolysis. Labs also add plenty of it. The Missouri–St. Louis recipe uses about 0.05 mole of anhydride for 0.015 mole of salicylic acid, and the West Virginia one uses three equivalents. The Odinity catalyst report adds a practical reason for the excess: acetic anhydride is easily hydrolysed by stray water and lost.

Why not something even more reactive, like acetyl chloride? LibreTexts lists three reasons industry prefers the anhydride. It’s cheaper. It’s safer, being less corrosive and slower to react with water. And it doesn’t release corrosive, poisonous hydrogen chloride fumes. The trade-off is waste: an anhydride throws away half of itself as the leaving group. LibreTexts’ textbook notes that’s why chemists mostly use cheap, readily available anhydrides such as acetic anhydride.

Side products and where the vinegar smell comes from

Acetic acid, twice over. Every molecule of aspirin formed releases one molecule of acetic acid. At the end, water is added to destroy the leftover anhydride. Each anhydride molecule reacts with water to give two molecules of acetic acid. LibreTexts notes this hydrolysis is slow at room temperature and faster with gentle warming, which is why the Odinity catalyst lab heated its quenched mixture for 5 minutes.

That acetic acid is the smell. PubChem, the National Institutes of Health chemical database, describes acetic acid’s odour as pungent, sour and vinegar-like. Acetic anhydride smells of vinegar too, and a LibreTexts page explains why: it reacts with water vapour in the air, and with the moisture in your nose, to make acetic acid. In the Odinity analysis lab, students noted a vinegar-like smell when water met the anhydride and identified it as acetic acid. Pure aspirin itself is odourless, according to the Merck Index entry quoted in the Missouri–St. Louis procedure. That same entry adds that in moist air aspirin is gradually hydrolysed back into salicylic and acetic acids.

Gummy solids. Salicylic acid carries both an –OH and a –COOH group, so in principle it can link to itself and form a polyester. The Missouri–St. Louis procedure says such a polymer forms and filters it off after dissolving the crude product in sodium bicarbonate. In the Odinity catalyst lab, some tubes turned into a gummy mass that the student took to be polymer and blamed on heating for too long.

A more precise culprit. A 2022 study by Cheng and colleagues at Tianjin University of Science and Technology collected that by-product and identified it by thin-layer chromatography, IR, NMR and melting point. It wasn’t a salicylic acid polymer but acetylsalicylic anhydride. In their runs, starting from 2 g of salicylic acid, the amount grew with reaction time at 85 °C, from 0.01 g after 5 minutes to 0.18 g after 15. It also grew with temperature over a fixed 15 minutes, from 0.01 g at 50 °C to 0.18 g at 85 °C. The salicylic acid was already used up within the first 5 minutes, so the authors propose the extra heating lets acetic anhydride join two aspirin molecules together. Older analytical work by Reepmeyer and Kirchhoefer, published in 1979, separated four related impurities of aspirin by chromatography: salicylic acid, salicylsalicylic acid, acetylsalicylsalicylic acid and acetylsalicylic anhydride. They isolated the last three from aspirin tablets, so acetylsalicylic anhydride was already a known impurity.

Unreacted salicylic acid. Any starting material left behind still has its free phenol group. That’s what the classic iron(III) chloride colour test picks up.

Crude product is usually cleaned up by recrystallization, the same purification Odinity’s benzoic acid melting point explainer walks through. Odinity’s acetylsalicylic acid synthesis lab shows a microscale run with phosphoric acid at 50 °C, with a 59.2% yield.

What is settled and what is still debated

The core chemistry is settled. Acetic anhydride transfers an acetyl group to salicylic acid’s phenol oxygen through nucleophilic acyl substitution. An acid catalyst speeds this up by activating the carbonyl, and acetic acid is the co-product behind the vinegar smell. The choice of anhydride over acetic acid follows from basic reactivity: phenols barely react with carboxylic acids, and Fischer esterification stalls at equilibrium.

Some details are rules of thumb. The excess of anhydride, the catalyst drops and the heating time vary between manuals, and one student report’s catalyst ranking by temperature rise isn’t a precise kinetic result.

The by-product is the main unsettled point. Many lab texts describe it as a salicylic acid polymer. The 2022 Tianjin analysis, using a standard textbook recipe, found acetylsalicylic anhydride and no detectable polymer, with more forming the longer and hotter the mixture was heated. That’s a single study of one sulfuric-acid recipe, and we found no independent replication yet. Whether the same is true for other recipes, such as phosphoric-acid or microscale versions, hasn’t been tested. Even a classic experiment taught at many universities has details worth checking.