Aspirin (Experiment #1)

William Butler and Fernando Marquez

I. Introduction

Acetylsalicyclic acid, or Aspirin, belongs to a group of medications called anaglesics. These medications can be used to reduce fever, swelling, and relieve pain. Felix Hoffman (1893) is created for discovering the synthesis of aspirin while working for the Bayer, the German chemical company. Today, 80 million aspirin tablets are used daily and is found in over 100 nonprescription drugs.1

The goal of this experiment was to synthesize Aspirin by using H2SO4 (sulfuric acid) in an acid catalyzed acetylation reaction. Essentially, sulfuric acid speeds up the rate of the reaction without creating unwanted products. This allows salicyclic acid and acetic anhydride to create acetyl salicyclic acid and acetic acid, as seen in figure 1:

Figure 1: Synthesis of Aspirin via Acid Catalyzed Acetylation Reaction2

To help purify this reaction, the crude product can be recrystallized. This process involves re-dissolving the crude product. When placed in a second solvent, the molecules can interact more selectively and can be refiltered.

The purity of this process can be tested through two different methods; the Ferric Chloride Test and by measuring the melting point of the final product. The phenol test serves as a quality control for the amount of phenol (found in the starting material) that results in the product. The ferric chloride test for phenols is assessed qualitatively. Once the test is conducted, a color change from orange to purple is a positive test for the phenol group. If this is the case, then the product is not pure and the crude product must be recrystallized, as described above.

The second test for purity involves a Mel-Temp Apparatus which is a machine that can be used to measure the melting point of a solid. The ideal melting point of pure acetyl salicyclic acid is 136 Co. By using the Mel-Temp Apparatus, the actual melting point can be

compared to the ideal melting point. A compound is considered pure when it is within 2oC of the pure substance. Impurities with the crystal lattice will disrupt this trend and cause a melting point that is lower than the idea.

II. Results and Discussion

Limiting Reagent: Salicyclic Acid (2g)

Excess Reagent: Acetic Anhydride (5 ml)

Salicyclic Acid: 2 grams (1 mol/138.12 g) = 0.014 moles

Acetic Anhydride: 5 ml (1.08 g/1 ml) = 5.4 grams (1 mole/102.09 grams) = 0.05 moles

Because .014 mol < 0.05 mol, salicyclic acid is present in lesser amounts and is the limiting reagent.

Theoretical Yield of Acetyl Salicyclic Acid:

0.014 mol C7H6O3 (1 mol C9H8O4 /1 mol C7H6O3) = 0.014 mol C9H8O4

0.014 mol C9H8O4 = (180.16 g C9H8O4 / 1 mol C9H8O4) = 2.52 g Salicyclic Acid

Crude Product Actual Yield:

Percent yield = actual/theoretical x 100%

= 2.126 g/2.612 g = 81.38%

No Recrystallization necessary.

Melting Point of Pure Acetyl Salicyclic Acid: 136oC

Melting Point of Product: 115oC

Results from Ferric Chloride Test

Color After the Addition of Ferric ChlorideSalicyclic Acid PurpleCrude Product Orange/ClearRecrustalized Product N/A

Although the crude product yield was only moderately high, the phenol test did not require recrystallization. This provided evidence that our product was mostly pure but there may have been functional errors in measurements which deviated from the ideal yield. The Ferric Chloride test for Phenols did not render a purple hue which would have been indicative of a phenol group which was not desired. The phenol test was mostly clear which means that the phenols were only found in the reactants, as expected.

The expected melting point was 136oC while the experimental melting point was 115oC. To be considered “pure”, the melting point would have to be within 20C of the ideal melting point. This would mean that our experimental melting point would have to be 134oC or higher which it was not. This suggests impurity with the sample. These impurities are often attributed to weak binding between the molecules. It is likely that the sample was wet which would account for a relatively pure crude product with a lower melting point.

As a result of the ferric chloride test it can be determined that our product was relatively pure. This is because the product did not possess purple character. Ambiguity can be found in the relatively moderate yield and the low melting point. These results suggest that both random and systematic errors were involved. A lower than expected yield could be due to measuring errors while the lower melting point could be rationalized systematically as the Mel-Temp requires a dry sample. Still, this experiment is considered a success given the purity of our product and the amount of information learned.

III. Experimental

2.0 grams of salicyclic acid, 5 ml of acetic anhydride, and 5 drops of sulfuric acid were added to a 125 ml Erlenmeyer flask and swirled until dissolved. The solution was then heated with a water bath and kept at 50oC for 10 minutes. The product was then cooled to room temperature for crystallization. The crystals were then collected by vacuum filtration with 50 ml H2O. The crystals were rinsed with water and vacuum dried. The crude product was collected and weighed. This was used to calculate the percent yield. The ferric chloride test was then performed. This test was used to determine the presence of unwanted phenols in the crude product. A color change from orange to purple is indicative of phenols. 10 drops of FeCl3 was added to saliclycic acid which served as a control and the crude product. The salicyclic acid was purple as expected and the product was orange. The data for this experiment can be found above in the Results and Discussion.

IV. References

1) Kluger, E. Introduction to Organic Lab Techniques. CHM 2211L, 2011. Print.2) University of California, Davis.

http://chemwiki.ucdavis.edu/@api/deki/files/2782/image063.png?revision=1 (access May 27, 2014).