Monday, May 23, 2011

Full Report: Synthesis of Aspirin

(This is posted to help my fellow college students, but please, use this only as a guide and don't copy paste because it's bad.)

Organic synthesis is the process where a desired organic compound is constructed or prepared from commercially available materials. The objective of organic synthesis is to design the simplest synthetic routes to a molecule.
Acetylsalicylic acid, also known as aspirin, is one of the most widely used medications to reduce fever and is also used as a pain killer. It is an acetyl derivative of salicylic acid. It is a white, crystalline, weakly acidic substance which melts at 135°C.
Aspirin is synthesized through the reaction of salicylic acid with acetyl anhydride which causes a chemical reaction that turns salicylic acid's hydroxyl group into an acetyl group, (R—OH R—OCOCH3).
For the reaction to take place, an inorganic acid such as phosphoric acid is used as a catalyst.

At the end of the exercise, the student should be able to:
1.     be introduced to the concept of organic synthesis;
2.     synthesize acetylsalicylic acid from salicylic acid by nucleophilic acyl substitution; and
3.     differentiate acetylsalicylic acid from salicylic acid by simple chemical tests.

V. SAMPLE CALCULATIONS (For parts III, IV and V, please refer to your own data. Be resourceful in this one. LOL!)

The first part of the experiment was the preparation of Acetylsalicylic Acid (Aspirin). A white, milky mixture was obtained when salicylic acid, acetic anhydride and phosphoric acid (a catalyst) were mixed. The mechanism of the reaction is:

This shows that the oxygen in salicylic acid attacks one of the carbons in acetic anhydride. Also, the mechanism shows how acetic acid was separated from the acetylsalicylic acid.

In the first part of the experiment, heating of the mixture was done and a clear yellow liquid was obtained (Table 2). Heating was employed so that salicylic acid would melt and react with acetic anhydride. On the other hand, water was added after heating (not at the start of the experiment). This is to prevent the reaction of acetic anhydride with water at the start of the experiment, if this had happened, no aspirin could have formed. In this manner, acetic anhydride was decomposed after the formation of aspirin.

After the adding 40mL ice-cold water, cooling to room temperature and placing in an ice bath, the liquid became whitish/cloudy with white precipitates. This addition of cold water is very important in purification and isolation of the crystals from the liquid since aspirin is insoluble in cold water. Purification is needed to eliminate any salicylic acid and acetic anhydride that did not react, as well as the acetic acid product and phosphoric acid. In this part, purification is not yet complete (it was continued on the recrystallization part). Isolation was done through suction filtration, white, sugar-like crystals were obtained.

The crude/impure product was then weighed and it weighed 4.40g. This is quite far from the theoretical yield because it still contains impurities. This data was used to calculate the percent recovery on the latter part of the exercise.

The second part of the experiment was recrystallization. This is the second part of the purification process. Here, 95% ethanol was added dropwise to the crystals until dissolved and after this, distilled water was added dropwise until cloudy/until recrystallization. Ethanol was used to dissolve aspirin along with the impurities such as salicylic acid and others. Cold water, on the other hand, is used to recrystallize only aspirin, thus, leaving all the impurities behind. Since aspirin is an ester, it should not be recrytallized from hot water since esters hydrolyses in hot water. After cooling in an ice bath (which further facilitates recrystallization and purification), the mixture was then suction filtered.

The weight of the recovered sample was 2.85g. The calculations for percent yield was shown in Table 6. The percent yield was 109.25%, meaning there was a slight error. Perhaps, the sample was not weighed properly or it was weighed when still wet.  On the other hand, the calculated percent recovery was 64.77%. Certainly, another error occurred. This could be due to handling problems in suction filtration or drying, etcetera.

As for the melting point data, the range of the crude sample was 120-124˚C and the range of the purified sample was 122-124˚C. (Actually, there could have been an error here since I wasn’t able to observe it. A classmate just told me the MP range of purified sample, too bad, I forgot to tell my groupmates!). Comparing the results to the literature value of 135˚C, both the purified and crude had a precise value BUT since the purified sample has a narrower range, it is logically more comparable to the literature value (hehe!).

In Table 8, the differentiation of synthesized acetylsalicylic acid from commercially available aspirin was accounted for. The test used in this part was Iodine test, which is a test for the presence of starch (since iodine can form a black complex with starch).After dissolving synthesized aspirin in 2mL water and 1mL iodine solution, a mixture of red-orange liquid and white precipitates was obtained while when commercially available aspirin was dissolved in 2mL water and 1mL iodine solution, a black precipitate in a dark brown to black solution was formed. This shows that commercially available aspirin contains starch.

Other tests that were performed were summarized in Table 7. Since salicylic acid has a phenol group, it gave a positive result to FeCl3 Test and KMnO4 Test, both of which react with phenol. Acetic anhydride gave a positive result to water solubility test to form acetic acid. The recrystallized aspirin, an ester, did not give any positive result to the tests since esters do not react with FeCl3 Test, KMnO4 Test and Tollen’s Test. Small esters are actually fairly soluble in water but solubility falls with chain length and hydrophobic parts. Since aspirin has a hydrophobic aromatic ring, it did not dissolve in water. Having these results, the recrystallized sample was then identified (or assumed) as acetylsalicylic acid.

Aspirin was prepared from the reaction of salicylic acid and acetic anhydride. Phosphoric acid was used as a catalyst. Upon addition of cold water, acetic acid was formed and thus eliminated. Other impurities like salicylic acid were removed upon the process of recrystallization.

The melting point range of the purified and crude samples were compared to the literature value and it showed that the purified sample is logically “near” to the literature value because of its narrow range.

The recrystallized product was differentiated from commercial aspirin through iodine test and it showed that the commercial aspirin contains starch. Other tests such as water solubility test, FeCl3 Test, KMnO4 Test and Tollen’s Test differentiated the starting materials, salicylic acid and acetic anhydride, from aspirin.

Rainsford, K. D. (2004). Aspirin and Related Drugs. USA: Taylor & Francis Inc, no page (e-book).

Whitten, K. W., R. E. Davis, M. L. Peck, G. G. Stanley (2007). Chemistry. 8th ed. USA: Thomson Brooks/Cole, p. 947.


  1. This comment has been removed by the author.

  2. Thanks for this! Through your references, I answered my own papers. (Though it was a bit hard to find that book from Whitten)

  3. i think the mechanism is quite wrong esp. on the second product where the electrons on O-H bond move to +O without a pushing force... like H could've abstracted first by a base before moving the electrons to O to neutralize its charge...

    and how about making the acetyl anhydride more electrophilic first by protonating some O?

  4. Thank you so much for your discussion and introduction. Wouldn't totally agree with the mechanism. Thank you.

  5. Thank you so much for this! It helped me in making our laboratory report!

  6. which of the two reactants was in excess

  7. which of the two reactants was in excess