WB01542_8.gif (729 bytes)Activity 1: Synthesis of Aspirin

WB01542_8.gif (729 bytes)Activity 1: Synthesis of Aspirin


Activity 1: Synthesis of Aspirin


In this laboratory activity you will synthesize aspirin, a derivative of salicylic acid. Salicylic acid and its derivatives are antipyretics. Such compounds lower the body temperature of a person with a fever. They have little effect if the body temperature is normal. Salicylates are also mild analgesics that relieve pain associated with headache, neuralgia, and rheumatism. Salicylic acid itself is not used for these purposes because it has an irritating effect on the stomach.

The most common salicylate used in medicine today is aspirin. In this activity, aspirin is prepared from salicylic acid and acetic anhydride.

When ingested, aspirin (acetylsalicylic acid) passes through the stomach largely unchanged. It is hydrolyzed in the intestinal tract liberating the active ingredient, salicylic acid.

In this laboratory synthesis, you will heat a mixture of salicylic acid and acetic anhydride with a trace of sulfuric acid as catalyst. Because aspirin is not very soluble in water, it can be isolated by addition of cold water to the reaction mixture followed by a gravity filtration.


To synthesize aspirin from salicylic acid, test for the purity of the product obtained, and determine the number of aspirin tablets that can be obtained from your preparation.


  1. Wear protective goggles throughout the laboratory activity.
  2. Acetic anhydride is irritating in the liquid and vapor state. Avoid contact with skin and eyes. Rinse with water any body area that comes in contact with acetic anhydride.
  3. Concentrated sulfuric acid is a powerful dehydrating agent; handle it carefully to avoid contact with skin and clothing.
  4. Methanol is toxic; breathing its vapors and skin contact with methanol should be avoided.


  1. Set up a ring stand with a ring, wire gauze, and burner. The top of the burner should be about 12 cm below the ring and wire gauze. Place on the wire gauze a 400-mL beaker half-filled with water. Heat the water to boiling. This is your water bath. As you wait for the water to boil, go to the next step.
  2. Weigh 2.0g salicylic acid on weighing paper, and transfer the solid to a 125-mL Erlenmeyer flask.
  3. Add 5mL acetic anhydride from a 50-mL buret.
  4. Add 5 drops concentrated sulfuric acid. (CAUTION: Sulfuric acid is corrosive.) Stir the mixture.
  5. Heat the flask in the boiling water bath for 10 min.
  6. Remove the flask from the water bath, and carefully add 2mL water from a 10-mL graduated cylinder. Swirl to mix the contents. Allow the flask to stand for 5 min.
  7. Add 40mL water from a graduated cylinder, and stir the solution until crystals begin to form.
  8. Cool the flask in an ice-water mixture in a 400-mL beaker for 10 min to complete the crystallization.
  9. Collect the product by gravity filtration using a 250-mL beaker to collect the filtrate. Wash the product twice with 5mL cold water.
  10. Let the aspirin dry until the next laboratory period.
  11. Transfer the dried aspirin to a weighed piece of filter paper and weigh. Calculate the number of aspirin tablets you have prepared:

  1. Calculate the percent yield of your aspirin from the formula below. The quantity in the denominator (2.5 g) represents the theoretical yield of aspirin based on the moles of salicylic acid and acetic anhydride used in the synthesis.

  1. Determine the qualitative purity of your aspirin in the following manner. Place 1mL methanol in each of three separate test-tubes. Add a few crystals of the following: Test-tube 1, salicylic acid; Test-tube 2, your prepared aspirin; and Test-tube 3, commercial aspirin, crushed. Add 1 drop of 1% iron(III) chloride solution to each test-tube. Observe the color in each tube. What can you conclude?
  2. Extension: Determine the melting point of the aspirin as directed by your instructor.
  3. Place the aspirin in a dry test-tube, label with your name(s), and give it to your teacher.
  4. Thoroughly wash your hands before leaving the laboratory.

Data Analysis and Concept Development

  1. Can the purity of the prepared aspirin be determined by the color test with the iron(III) chloride?
  2. Can the purity of the prepared aspirin be determined by the melting point?
  3. In Step 12, the quantity 2.5g represents the theoretical yield of aspirin. Show the calculation leading to a theoretical yield of 2.5g of aspirin. [You must first determine the limiting reactant from the following data: Molar mass of salicylic acid = 138 g; molar mass of acetic anhydride = 102 g; density of acetic anhydride 1.08 g/mL.]

Implications and Applications

  1. What steps do aspirin manufacturers have to take to make sure that theirproduct is fit for human ingestion?
  2. What other types of analgesics are on the market? How does their action compare with aspirin and why are these alternative drugs necessary?



Activity 1: Synthesis of Aspirin

Major Chemical Concept

The synthesis of aspirin is known in organic chemistry as an esterification reaction. This is a substitution reaction in which an alcohol (the –OH group in salicylic acid) reacts with acetic anhydride to form an ester, aspirin.



Honor chemistry students

Expected Student Background

None expected. The techniques illustrated here (synthesis, melting point determination) are not likely to have been previously encountered.


The activity will take one 50-min period for the synthesis, an overnight drying step, and a second period for calculations, color tests, and melting point determination.


Read the Safety Considerations in the Student Version. The acetic anhydride,methanol and concentrated sulfuric acid should be handled with care. Dispensing the anhydride from a buret minimizes spills and possible body contact. Concentrated sulfuric acid should be dispensed from a small dropping bottle. The aspirin produced in this activity is not pure. It is often contaminated with salicylic acid, acetic acid, and/or sulfuric acid. Students should not taste their aspirin.

Materials (For 24 students working in pairs)



Advance Preparation

Fill the 50-mL buret with acetic anhydride. Prepare 1% iron(III) chloride solution in dropping bottles. Set out container of salicylic acid. Use of vacuum filtration, if available, will facilitate the filtration. Purchase commercial aspirin tablets, or obtain samples from the school health office. Set up mineral oil baths for melting point determination as shown in Possible Extensions.

Pre-Laboratory Discussion

  1. Briefly explain the reaction type (substitution, esterification) and synthetic goal.
  2. Caution students about reagent hazards.
  3. Review the use of a buret for delivering a given volume of liquid.
  4. Review technique of stirring versus swirling.
  5. Review theoretical yield calculation.
  6. Explain the concept in the iron(III) chloride test. Many phenols form colored coordination compounds with iron(III) ion, in which six molecules of a monohydric phenol are combined with one atom of iron in the form of a complex anion. Salicylic acid contains a phenolic –OH group and gives a positive test (turns purple) with iron(III) chloride. Complete reaction of salicylic acid with acetic anhydride substitutes an ester for the phenolic group. Therefore, aspirin uncontaminated with the salicylic acid starting material will give a negative test with Fe 3+ . What can students conclude if their aspirin preparation turns purple? [The preparation still has salicylic acid starting material present.]
  7. Explain how to obtain a melting point range (temperature at which solid first melts to the temperature at which solid is melted).

Teacher-Student Interaction

Walk around room during laboratory exercise emphasizing proper technique. Talk with students to ascertain if they understand activity. Help with yield calculations and melting point determination.

Anticipated Student Results

Typically students obtain a 60-70% yield. The product is often contaminated with salicylic acid. The result of the iron(III) chloride test is: salicylic acid (dark purple), prepared aspirin (dark purple), and commercial aspirin (light salmon). The melting point range is broad, i.e., there is a large temperature range between first melting and complete melting of the sample. Typical values are between 128-137 C. Recrystallization of the aspirin would remove unreacted salicylic acid and narrow the melting point range.


Answers to Data Analysis and Concept Development

  1. Yes. The prepared aspirin should not give a positive color test with the iron(III) chloride because salicylic acid, which gives the positive test, is absent.
  2. Yes. The sharpness of the melting point is one of the best methods of determining purity of an organic solid. A broad melting point indicates the presence of impurities. In this case, the impurity is probably a trace of unreacted salicylic acid.

Answers to Implications and Applications

  1. Purification by recrystallization and testing for impurities by melting point determination and chromatography are necessary. Strict regulatory codes apply.
  2. Other analgesics may be anti-inflammatory or antipyretic, but only aspirin is a "triple action" drug (fulfills all three functions). Aspirin is sometimes irritating to the digestive tract lining, so alternatives must be available for individuals affected by aspirin in this way.

Post-Laboratory Discussion

Review the experimental techniques, the experimental results and their significance.

Possible Extensions

  1. Determine the purity of your aspirin by obtaining its melting point. Use the set-up shown in Figure 1. Crush about 50-100 mg of dry aspirin with a spatula against the walls of a 50-mL beaker. Thrust the open end of a melting point capillary tube into the powdered aspirin several times. Work the plug of solid material down to the tube’s sealed end by dropping it 3-4 times into the open end of the 50-mL beaker used for crushing. The capillary tube should contain the amount of packed solid as shown in Figure 1. Attach the capillary tube to the thermometer as shown in Figure 1. The rubber band can be a small section of rubber tubing of suitable diameter. Place the thermometer and tube in the 250-mL beaker oil bath. Heat the oil bath so that the temperature rises about 1-2 C/min. Record as the melting point range the temperature at which the sample begins to melt and the temperature at which the sample is completely melted. The melting point of aspirin is about 135 C. Discard the used melting point tubes in a waste glass container.


  1. Synthesize methyl salicylate (oil of wintergreen) in a medium-sized test-tube from 0.5g salicylic acid and 1mL methanol. Stir the mixture, and add 3 drops of concentrated sulfuric acid. Heat the mixture in a water bath for 5 min. Remove the test-tube, cool it in a test-tube rack, and then add the contents to a 100-mL beaker containing 4-5 ice cubes. Waft the odor toward thenose. What commercial products contain oil of wintergreen? [Wintergreen Life-Savers, Ben-Gay, etc.]

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  1. Use a Spectronic 20 to determine the amount of salicylic acid in the laboratory-synthesized aspirin and in commercial aspirin. To prepare an aspirin solution, dissolve 0.05g aspirin in 50mL water. Prepare an iron(III) nitrate solution by dissolving 2g iron(III) nitrate nonahydrate [Fe(NO 3 ) 3 . 9 H 2 O] in 50mL water and adding 0.5mL concentrated nitric acid. To 1mL aspirin solution add 5mL iron(III) nitrate solution. Shake well for 5min (a violet color should appear if salicylic acid is present). Measure the absorbance at 535–540 nm by placing the violet solution in the cell of a Spectronic 20 instrumentafter standardizing with distilled H 2 O. NOTE: Old Spectronic 20’s may still have their wavelength scales calibrated in mm (millimicrons) but the accepted term for this unit is nanometers(nm). The absorbance is proportional to the concentration of salicylic acid in the sample. Compare your laboratory aspirin with several brands of commercial aspirin.[The laboratory synthesized aspirin should have an absorbance >0.5 at 540 nm. The commercial aspirin has an absorbance of <0.2 at 540 nm. The commercial aspirin solution does not change to a violet color when mixed with the iron(III) nitrate solution.]

Assessing Laboratory Learning

  1. What are the most likely impurities present in the sample of aspirin you prepared? [Unreacted salicylic acid, traces of acetic acid.]
  2. If aspirin sits for long periods of time, the odor of vinegar can be noticed in the container. A hydrolysis reaction occurs very slowly. Explain, using a chemical equation,thesourceof thevinegar,andthe reasonforthehydrolysis. [Slow reaction with moisture in air can produce small amounts of vinegar.]

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  1. A white powder was tested by a police chemist with iron(III) chloride solution. A purple color is seen. What conclusion was drawn by the police chemist?[The white powder contained a phenolic compound, possibly salicylic acid.]
  2. Aspirin tablets are sold as containing 5grains of aspirin. If 1 grain equals 65mg, how many milligrams of aspirin does each tablet contain? [325mg]