Aspirin Synthesis in a Teaching Laboratory

Aspirin, or acetylsalicylic acid, is a useful compound for introducing organic synthesis, stoichiometry, crystallisation and analytical chemistry. In a supervised teaching laboratory, students can prepare a small sample by acetylating salicylic acid with acetic anhydride, then isolate the product as crystals.

This experiment is suitable for senior secondary or undergraduate chemistry classes when conducted under an approved risk assessment. The prepared material is for laboratory analysis only and must never be swallowed or substituted for pharmacy-grade aspirin purchased in Australia.

Choosing Scale and Safety

A practical microscale preparation begins with 2.00 g of salicylic acid, approximately 0.0145 mol. The reaction also requires about 5 mL of acetic anhydride and a few drops of an acid catalyst, such as concentrated phosphoric acid. Acetic anhydride is corrosive, lachrymatory and moisture-sensitive, so all transfers should take place in a functioning fume cupboard.

Wear a buttoned laboratory coat, splash-resistant safety glasses and suitable chemical-resistant gloves. Keep the reaction vessel away from the edge of the bench, and use a water bath rather than a naked flame. Australian schools and universities should follow their local chemical risk assessment, Safety Data Sheet requirements and relevant Safe Work Australia guidance. Students in laboratories in Sydney, Melbourne, Brisbane or elsewhere should also follow the specific procedures of their institution.

Reaction Chemistry

The synthesis is an esterification reaction. The phenolic hydroxyl group of salicylic acid reacts with acetic anhydride, producing acetylsalicylic acid and acetic acid as the main by-product. Phosphoric acid helps activate the acetic anhydride without becoming part of the final aspirin molecule.

The balanced reaction can be represented as:

C₇H₆O₃ + C₄H₆O₃ → C₉H₈O₄ + C₂H₄O₂

Salicylic acid is the limiting reagent when acetic anhydride is used in excess. Its molar mass is 138.12 g mol⁻¹, while aspirin has a molar mass of 180.16 g mol⁻¹. From 0.0145 mol of salicylic acid, the theoretical yield is approximately 2.61 g of aspirin.

Experimental Procedure

Place 2.00 g of finely powdered salicylic acid into a dry 100 mL conical flask. Add 5 mL of acetic anhydride carefully, followed by about five drops of concentrated phosphoric acid. Swirl gently to wet the solid, then place the flask in a water bath maintained at approximately 70–80 °C for 10–15 minutes. Do not seal the vessel, and avoid breathing vapour from the flask.

Remove the flask with tongs and allow it to cool briefly. Add approximately 20 mL of distilled water slowly down the inside wall of the flask. This destroys remaining acetic anhydride and encourages crude aspirin to precipitate. The mixture may warm and fizz slightly, so add water gradually and keep the flask pointed away from people.

Cool the suspension to room temperature and then place it in an ice bath for several minutes. Scrape the inside of the flask gently with a glass rod if crystals are slow to appear. A cold mixture generally produces a better solid recovery because aspirin has limited solubility in cold water.

Isolation and Recrystallisation

Collect the crude crystals using a Büchner funnel, filter paper and gentle vacuum. Rinse the crystals with a small volume of ice-cold distilled water to remove soluble acetic acid and unreacted reagents. Avoid washing with too much water because some aspirin will dissolve and the measured yield will decrease.

For purification, transfer the damp solid to a small flask and add the minimum volume of warm ethanol needed to dissolve it. Add warm water gradually until the solution is just cloudy, then warm carefully until it becomes clear. Allow the solution to cool slowly before placing it in an ice bath. Filter the purified crystals and leave them to dry to constant mass.

Ethanol is flammable, so it must be heated with a hot-water bath or approved heating equipment rather than a Bunsen burner. In an Australian teaching laboratory, solvent use and storage should match the institution’s fire-safety procedures, particularly in busy practical classes.

Checking Product and Yield

Record the mass of dry aspirin and calculate the percentage yield using:

Percentage yield = actual yield ÷ theoretical yield × 100

A pure sample of aspirin commonly melts at about 135–136 °C, although the exact result depends on calibration, heating rate and sample quality. A lower or broader melting range can indicate residual salicylic acid, trapped solvent or another impurity. Comparing the result with a reliable reference value is more informative than treating a single temperature as proof of identity.

An iron(III) chloride test can provide a simple qualitative check for unreacted salicylic acid. Salicylic acid gives a purple-coloured complex with iron(III) ions, whereas properly purified aspirin gives little or no purple colour under the same conditions. This test uses a corrosive reagent and should be carried out only with suitable supervision and waste controls.

Waste and Scientific Interpretation

Liquid residues containing acetic acid, phosphoric acid or dissolved aspirin should be placed in the labelled aqueous organic-acid waste container specified by the laboratory. Used ethanol, contaminated filter paper and excess reagents must be handled according to the local waste plan. Do not pour concentrated chemical residues into a sink simply because the final mixture appears dilute.

The experiment demonstrates why purification affects both yield and quality. Using too much washing water may reduce recovery, while insufficient washing leaves impurities in the crystals. Students can compare crude and recrystallised material through appearance, melting range, mass and, where available, infrared spectroscopy or thin-layer chromatography.

Aspirin is widely recognised in Australian pharmacies and supermarket medicine aisles, including products sold through major pharmacy chains and independent chemists. That commercial product has been manufactured, tested and labelled for therapeutic use; a student preparation has not. The practical value of the experiment lies in understanding reaction control, isolation and analysis, not in producing a medicine for personal use.

Accurate weighing, controlled heating, careful cooling and complete drying are the foundations of a reliable result. Keeping a clear record of reagent masses, temperatures, observations, yield and melting behaviour turns a simple aspirin preparation into a rigorous lesson in organic laboratory practice.