Hydrolysis of Acetylsalicylic Acid (Aspirin)

Estimated Completion Times Pre lab Investigation 1.5 hours Experimentation Overnight reaction period + 1 hr Post Lab Assessment 30 min Hydrolysis of Acetylsalicylic Acid (Aspirin) Note: This experiment requires a 12-hour reaction period. In this experiment you will perform a hydrolysis reaction by mixing aspirin with water and using iron(III) chloride to detect the presence of salicylic acid. You will also compare the purity of freshly prepared aspirin powder to powder exposed to air for 8 to 12 hours. Learning Objectives: Understand why salicylic acid (C7H6O3) is an appropriate starting material for the synthesis of aspirin; Explain how the organic synthesis of aspirin from salicylic acid in a non-aqueous medium is conducted; Ability to identify the reactants, products, solvent and catalyst in the synthesis of aspirin from salicylic acid; Describe the procedure and its limitations in recovering the product from the reaction mixture using solubility criteria; Understand how to test the purity of aspirin by melting point, iron(III) chloride test and analyzing the IR spectrum; Perform a hydrolysis reaction and carryout a chemical test to demonstrate the desired product was obtained. Introduction: Aspirin, which is also known as acetylsalicylic acid (ASA), is an over the counter medicine used to reduce fever, pain, and/or inflammation. In certain populations it is also used to reduce the risk of heart attack. 1 One way to synthesize aspirin is to react salicylic acid with an excess of acetic anhydride (Figure 1). A few drops of a strong acid, such as sulfuric acid can be used as a catalyst in order to speed up the reaction. Acetic acid will be produced as an unwanted byproduct, and unreacted (excess) acetic anhydride will also remain. The excess acetic anhydride can be converted to additional acetic acid by hydrolysis (addition of water). Since acetic acid is very soluble in water it will all be quenched. The aspirin product has a fairly low solubility in cold water, so it will precipitate when ice water is added to the product solution. The synthesis reaction of aspirin is shown below: Figure 1. Synthesis of aspirin Since acetic acid is very soluble in water, it can be easily separated from the aspirin product. The aspirin isolated in this step is the “crude product”. A “purified product” can be obtained through recrystallization of the crude product in hot ethanol. If we were in the laboratory performing the synthesis, the percent yield of the crude product would be determined for this reaction and the purity of the product would be analyzed by three different methods: melting point, ferric chloride test and infrared spectroscopy. In this at-home experiment, instead of performing a synthesis you will perform a hydrolysis reaction by mixing aspirin with water. Iron(III) chloride will be used to detect the presence of salicylic acid. In addition, you will also compare the purity of freshly made aspirin powder to aspirin powder that was exposed to air for 8 to 12 hours. Background - Synthesis of Aspirin: 1. Establishing the Starting Materials and the Reaction The structure of aspirin is shown below in Figure 2. The choice of appropriate starting materials and synthetic route requires knowledge of the availability of substances and the reactions that they undergo, as well as safety and environmental concerns. Figure 2: Reactant (salicylic acid) and product (aspirin) 2 Salicylic acid (Figure 2, above) is easily synthesized and available. The question remains: how to convert the H of the OH group directly connected to the benzene ring to the CH3CO (acetyl) group in aspirin? Organic chemistry provides a general synthetic process for such conversions. In our case, the reaction is between salicylic acid and a compound closely related to acetic acid, acetic anhydride. Acetic anhydride results from the elimination of a molecule of water from two molecules of acetic acid, as shown in Figure 3 below. Acetic anhydride is a common reagent. Figure 3. Production of acetic anhydride We have now established reasonable starting materials and an appropriate reaction for the synthesis of aspirin. The starting materials are salicylic acid and acetic anhydride, both are readily available and inexpensive and are safe to use in a chemistry laboratory equipped with a modern fume hood. Our reaction becomes: Figure 4. Synthesis of acetylsalicylic acid (aspirin) Observe that the reaction also produces a second unwanted product, acetic acid. 2. The Reaction Conditions The conditions under which the reaction should be conducted depend on some of the properties of the reactants and products. Acetic anhydride is a liquid at room temperature and boils at a temperature well above 100°C. It reacts with water to form acetic acid – the reverse of the reaction shown above for its formation. Salicylic acid and aspirin are both solids that melt at temperatures well above 100°C. They are both slightly soluble in cold water (2.2 mg/mL, and 3.3 mg/mL, respectively). In addition, salicylic acid is also soluble in acetic anhydride. From these properties, we deduce that water is an inappropriate solvent in which to carry out this reaction, since water would deactivate the acetic anhydride. Since acetic anhydride is a liquid, we will use this starting material as the solvent as well. Thus it performs two roles here, reactant and solvent. Since the reaction is slow in pure acetic anhydride, we will use a catalyst for the reaction, a strong acid (concentrated sulfuric acid). According to Le Chatelier’s principle (see discussion in Background – 3 Hydrolysis of Aspirin), the presence of an excess of acetic anhydride forces the equilibrium towards the desired product, aspirin. We will also heat the solution to accelerate the approach to equilibrium. 3. Isolation of the Product When the reaction has proceeded to its final state, in addition to aspirin, the reaction vessel will contain unreacted acetic anhydride, unreacted salicylic acid (possibly), acetic acid and the catalyst, sulfuric acid. We have noticed that both aspirin and salicylic acid are only slightly soluble in water. After the reaction is complete, we can consider adding water to the reaction vessel. The water will react with the excess acetic anhydride and produce acetic acid. If we use relatively small amounts of water, only small amounts of the aspirin and salicylic acid will dissolve and the bulk will form a precipitate in the water solution. Since the solubility of most substances decreases with decreasing temperature, we use cold water (or ice) to further limit the solubility of aspirin. The acetic acid formed as a product will dissolve in the water and remain in the liquid phase. Therefore, we are able to recover the aspirin by filtering the cold solution. However, the unreacted salicylic acid might also be recovered along with the aspirin. 4. Effectiveness of the Synthesis (Yield) and Purity of the Product The remaining issues are: how much of the final product, aspirin, have we produced and, how pure is the product? From the amounts of the starting materials and observations of the limiting reactant, we can calculate the maximum amount that could be produced – i.e., if the reaction had gone “to completion.” The stoichiometry of the reaction shows that 1 mol of acetic anhydride reacts with 1 mol of salicylic acid to produce 1 mol of aspirin. If we used 1.38 g of salicylic acid (138 g/mol), which is 0.0100 mol and an excess of acetic anhydride, the maximum amount of aspirin we could make is 0.0100 mol or 1.80 g. This maximum amount is called the theoretical yield. If the reaction did not go “to completion”, or if we lose some of the aspirin through solubility in water that we add at the completion of the reaction, we would expect to get less than the theoretical yield. Suppose the actual amount (actual yield) of aspirin produced is 1.50 g. The percent yield (% yield) is defined as: % yield = (actual yield/theoretical yield) x 100 In the current example, the % yield is (1.50g/1.80 g) x 100 = 83.3% Note that in defining the theoretical yield, we calculate the maximum amount of the desired product. In the current exercise, even if the final product were dry, it could contain some unreacted salicylic acid. Therefore, we need a way of determining the purity of the product. In the case of aspirin, a relatively simple qualitative test can detect the presence of salicylic acid in the final product. 4 5. Ferric Chloride Test for Salicylic Acid The addition of ferric chloride to salicylic acid produces a specific color caused by the reaction of salicylic acid with aqueous ferric (Fe(H2O)63+) ion. The oxygen atoms of the acid group –COOH, and the –OH group on salicylic acid together form a complex with Fe(H2O)63+ as shown below in Figure 5. That complex has an intense violet color. In aspirin, the – OH group of salicylic acid has been replaced by an O–COCH3 group, which prevents the second bond from being formed. The resulting complex with aspirin shows only a slight yellow color, not very different from that of Fe(H2O)63+ itself. Figure 5. Results of ferric chloride test - aspirin is present (yellow); salicylic acid is present (purple). Two additional tests of aspirin purity that can be performed are 1) melting point determination and 2) analyzing an infrared absorption spectrum of the product. 6. Melting Point Determination The melting point of a product is useful for determining the identity and purity of a product. For a pure compound, its melting point (at a constant pressure) has a constant value, and these values can be found in a variety of reference books such as the Handbook of Chemistry and Physics. The melting point range is defined as the temperature at which the sample starts to melt (the first appearance of a liquid droplet) to the temperature at which the entire sample has melted. Pure compounds have very narrow melting point ranges, typically 1-2°C. Melting point lowering, a colligative property, is due to the presence of an additional solute or an impurity. The presence of impurities will typically lower the melting point and broaden its range. Keep in mind the absolute error in your thermometer reading when reporting melting points and agreement with the literature values. If one were to report a melting point for ice as: −2.0 to −1.0°C ± 2°C, this would agree with the literature value of 0.0 °C and indicate that the ice was pure. Literature melting points: Salicylic acid 159 °C; Aspirin = 136 °C. 7. Infrared Spectroscopy The infrared absorption spectra of aspirin and salicylic acid each contain features unique to these compounds. The product, aspirin, has a carbonyl (−C=O) group and a methyl (−CH3) group that are absent in the starting material. The starting material has an –OH group attached to the benzene ring that is absent in the product. Within the limits of detectability of the infrared technique, we 5 can verify the qualitative composition of the product by comparing the spectrum of the product with the spectra of pure aspirin and salicylic acid. Background – Hydrolysis of Aspirin: 1. Starting Materials Salicylic acid can also react with acetic acid to form acetylsalicylic acid (aspirin) via a condensation reaction as shown in Figure 6. A condensation reaction is when two organic molecules (salicylic acid and acetic acid) combine to produce one organic molecule (aspirin) and release a water molecule. Figure 6. Condensation reaction for the synthesis of aspirin. Look carefully at Figure 6 and take note of the double arrows. Double arrows tell you that the reaction is reversible. Once the reaction starts it will proceed until equilibrium is reached. Once equilibrium is reached, the rate of the forward and reverse reactions as well as the concentrations of the reactants and products stays constant. The reverse reaction is our hydrolysis reaction. Hydrolysis simply means “to break using water”. In this experiment hydrolysis (being broken down) refers to converting a larger organic molecule (aspirin) into smaller organic molecules (salicylic acid and acetic acid) by the addition of water. 2. Le Châtelier’s Principle According to Le Châtelier’s principle, when a change is imposed on a chemical system, the system will act to reduce the total effect of the change. In the case of our experiment with respect to acetylsalicylic acid, when water (a product) is added to the system, the system compensates by producing more of the reactants. Therefore, adding water to acetylsalicylic acid yields salicylic acid and acetic acid. In this experiment, you will use the concept of Le Châtelier’s principle to carry out the hydrolysis of acetylsalicylic acid (aspirin). You will add water to aspirin powder with the desired outcome of favoring the reverse reaction (producing salicylic acid and acetic acid). A simple test can be 6 performed to determine if the hydrolysis reaction proceeded successfully (iron(III) chloride test). As discussed in section 5 above, FeCl3 is a reagent that can be used to test for the presence of (OH) (phenols). In solution, the hydrated iron(III) ions bond with salicylic acid molecules to produce iron complex molecules with a violet color (Figure 5). When FeCl3 is mixed with acetylsalicylic acid (aspirin), the solution remains yellow with no purple color (Figure 5). Since FeCl3 does not react with aspirin, it is an appropriate test to investigate the purity of aspirin. Precautions: Ferric Chloride Solution: Corrosive Aspirin: Pulmonary irritant – don’t inhale dust. Procedure: Note: This experiment requires a 12-hour reaction period. Part A: Hydrolysis reaction 1. Collect all items needed including aspirin tablets, permanent marker, test tubes, well plate, FeCl3 pipet, spatula, empty long thin-stem pipet, test tube cleaning brush, 2 sheets of paper, metal spoon, distilled water, digital camera/smart phone and scissors. 2. Place your safety goggles securely on your face and put on gloves. 3. Write your name and date on a sheet of paper. 4. Use the permanent marker to label one test tube “A” and the other test tube “B.” Place test tube A and test tube B in the 24 well plate as shown in Figure 7. Figure 7. Test tube A and B and FeCl3 pipet in well plate 7 5. Pulverize both aspirin tablets into a fine powder using a metal spoon (or provided spatula) on the other sheet of paper as shown in Figure 8. Figure 8. Pulverizing aspirin with spoon 6. Use the spatula to transfer about 1/3rd of the powder from one aspirin tablet to test tube B as shown in Figure 9 below. The amount does not have to be exact. It should look similar to the amount in Figure 9 (~0.5 cm3). Figure 9. Transfer approximately 0.5 cm3 fine aspirin powder to test tube B 8 7. Add 20 mL of distilled water to the 25 mL graduated cylinder as shown in Figure 10. Look carefully at the graduation units. Is each graduation 1 mL or 0.5 mL? Figure 10. 25 mL graduated cylinder with 20 mL of distilled water. 8. Use the long thin-stem pipet to transfer 5 mL of distilled water from the graduated cylinder to test tube A and test tube B. 9. Gently swirl the mixture in test tube B so all of the aspirin powder is dissolved in the water. Use the metal spatula to stir the mixture to fully dissolve the powder if needed. 10. Use scissors to carefully remove the tip of the iron(III) chloride (FeCl3) pipet according to the directions below: a. Make sure goggles and gloves are always worn during the duration of the experiment. b. While holding the pipet away from you, carefully snip open the tip of the pipet using scissors. Be sure to cut in a straight line, not on a diagonal. Be careful as iron(III) chloride may escape the pipet while opening! c. Place the pipet bulb-side-down in the well plate until it is called for in the procedure (see Figure 7). 9 d. Wash your scissors and dry with paper towels. 11. Add one drop of FeCl3 to test tube A and test tube B. Gently swirl both test tubes to thoroughly mix the solutions. 12. Put the test tubes back in the well plate and observe. What is the color of each solution? Is salicylic acid (SA) present in each solution? Record your observations and conclusion for the presence of SA in Data Table 1. 13. Use your camera or smartphone to photograph the two solutions, test tube A and test tube B (Photo 1 and Photo 2 respectively) with the sheet of paper displaying your name and date. Be sure to include your photos with your notebook submission on blackboard. 14. Put the well plate containing test tube A and test tube B in a safe and secure location for 8 to 12 hours while the hydrolysis reaction takes place. 15. Expose the remaining aspirin powder to air for 8 to 12 hours by keeping it in a protected location, away from others including children and pets. 16. Wash the spoon or spatula with soap and water. 17. Put aside all other materials to be used in Part 2. Part B: Identifying Salicylic Acid in Solution 18. Once 8 to 12 hours has passed, gather the aspirin powder and well plate and return to your work area. 19. Place your safety goggles securely on your face and put on gloves. 20. Gather all the other items used in Part A. 21. Observe test tube A and B. What is the color of each solution? Is salicylic acid (SA) present in each solution? Record your observations and conclusion for the presence of SA in Data Table 2. 22. Use a camera or smartphone to photograph the two solutions, test tube A and test tube B (Photo 3 and Photo 4 respectively) with the sheet of paper displaying your name and date. Be sure to include your photos with your notebook submission on blackboard. 10 23. Dispose of the contents and clean test tube B. Be sure to use the test tube cleaning brush and plenty of soap and water. 24. Rinse test tube B with distilled water and shake to dry. 25. Relabel the cleaned and dry test tube “C” using the permanent marker. 26. Use the spatula to transfer approximately 0.5 cm3 (about 5 mg) of the aspirin powder that was exposed to air for 8 to 12 hours to test tube C. 27. Use the long thin-stem pipet to transfer 5 mL of distilled water from the graduated cylinder to test tube C. 28. Gently swirl the mixture in test tube C so all of the aspirin powder is dissolved in the water. Use the metal spatula to stir the mixture to fully dissolve the powder if needed. 29. Add one drop of FeCl3 to test tube C. Gently swirl the test tubes to thoroughly mix the solution. 30. What is the color of each solution? Is salicylic acid (SA) present in each solution? Record your observations and conclusion for the presence of SA in Data Table 3. 31. Use your camera or smartphone to photograph the two solutions, test tube A and test tube C (Photo 5 and Photo 6 respectively) with the sheet of paper displaying your name and date. Be sure to include your photos with your notebook submission on blackboard. Cleanup Properly dispose of the remaining solutions and chemicals according to local laws and regulations. Clean the test tubes and well plate with soap and water using the test tube cleaning brush. Put all clean, dry equipment back in the lab kit for future use. 11 Data Tables: You should draw out data tables before you start the experiment and fill it out during the experimentation step. Be sure to leave space for your notes! Data Table 1: Test tube A (H2O + FeCl3) and test tube B (H2O + FeCl3 + aspirin), time = 0 (beginning of experiment). Sample Description Salicylic Acid Present Test tube A Test tube B Data Table 2: Test tube A (H2O + FeCl3) and test tube B (H2O + FeCl3 + aspirin), time = 8-12 hours. Sample Description Salicylic Acid Present Test tube A Test tube B Data Table 3: Test tube A (H2O + FeCl3) and test tube C (H2O + FeCl3 + aspirin) exposed to air for 8-12 hours. Sample Description Salicylic Acid Present Test tube A Test tube C Discussion: 1. Define equilibrium. How is equilibrium denoted in a written chemical equation? 2. Referring to Figure 6, draw the molecules involved in the synthesis of aspirin in your notebook. Observe the reaction you drew. Is the forward reaction hydrolysis or condensation? Is the reverse reaction hydrolysis or condensation? Describe what occurs in each type of reaction with respect to acetylsalicylic acid. 3. What was the purpose of test tube A in this experiment? 12 4. State Le Châtelier’s principle in your own words. Explain how Le Châtelier’s principle was used in this experiment. Compare and reference Data Table 1 and Data Table 2 in your answer. 5. How did exposing the aspirin powder to air for 8 to 12 hours affect its composition? Compare and reference Data Table 1 and Data Table 3. 13