Aspirin Synthesis Your team of scientists has been hired by the drug research division of Santa Monica Pharmaceuticals, a new start-up company formed by some SM alumni. You have been asked to evaluate a simple organic synthesis for Aspirin and to determine if the synthesis merits further funding by Santa Monica Pharmaceuticals. You will have one week to test the synthesis, evaluate the purity of the product and make your recommendations. In many ways, the science of organic synthesis is like finding the solution to a puzzle. The goal of organic synthesis is always to prepare a target molecule, perhaps chosen because of its intrinsic scientific importance or, perhaps, because of its economic importance. The constraints imposed upon the synthesis are what makes it challenging to the synthetic chemist. For example, often a route is chosen because the starting materials and reagents are available inexpensively or in sufficient quantities. ften working against this constraint is the fact that, for efficiency, a synthetic route should involve the least number of steps possible, because each chemical and physical step reduces the overall efficiency. The choice of the most efficient route is also affected by other constraints. Sometimes (although not in this lab!), efficient synthetic routes involve highly reactive or toxic species, even if the product has none of those properties. Such a route might improve efficiency, but it would also increase the complexity of the synthesis because of the additional safety precautions that would need to be implemented. Finally, the end use of the product often dictates the acceptable level of purity of the material. Typically, chemicals destined for use as pharmaceuticals or in the electronics industry are needed in extremely pure form. Purification strategies can often be as complex as the syntheses themselves. In the first part of the lab your team will make aspirin, purify your product and calculate the percent yield. In the second part of the lab you will perform quality control experiments on your aspirin, similar to the kinds of tests that are required by the FDA. Figure 1 Acetylsalicylic acid (aspirin), 9 8 4 Aspirin (acetylsalicylic acid) is a versatile drug that is consumed in huge quantities worldwide. It is a non-steroidal anti-inflammatory drug (NSAID) with a wide range of physiological effects. At very low doses, aspirin is used to treat and prevent heart attacks and blood clots. At higher doses
it is used as an analgesic to reduce pain and as an antipyretic to reduce fever. At very high doses, it is an effective anti-inflammatory agent used to treat rheumatic fever, gout and rheumatoid arthritis. It is also an anticoagulant, it dissolves corns and calluses, and it provokes loss of uric acid (a toxin) but promotes retention of fluids in the kidneys. It kills bacteria and induces peptic ulcers. The exact mechanisms of its pharmacological actions are still under study. In many plants, salicylates can induce flowering. In the synthesis you will be evaluating, you will start with salicylic acid and make aspirin using acetic anhydride according to the following reaction: acetic anhydride + + salicylic acid ( 7 6 3 ) acetic acid acetylsalicylic acid (aspirin, 9 8 4 ) Note that it is the - group of salicylic acid that reacts with acetic anhydride to form an esterlike product. The carboxylic acid group of salicylic acid remains unchanged. Phosphoric acid will be used as a catalyst in this experiment. aving synthesized aspirin, you will purify it by crystallization and filtration. Although there will be some loss of product, good experimental technique will minimize the losses. If your aspirin were going to be used pharmaceutically it would require even further purification. The main impurity in the crystallized aspirin will be salicylic acid, which will co-precipitate with the aspirin if the procedure is done too quickly. The first method you will use to determine the purity is the melting point of your product. We will use commercial Mel-temp apparatuses to determine the temperature at which your synthesized crystals melt. If the aspirin is pure, it will melt sharply at the literature value. If it is impure, it will be lower than the literature value by an amount that is roughly proportional to the amount of impurity present. A more quantitative method of determining the purity of your aspirin is to use absorption spectroscopy (the same method we used in the manganese lab earlier this semester you may want to review this material now). In this method you will react a sample of your purified product with Fe 3+ (aq), introduced as Fe(N 3 ) 3. This will form an intensely purple-colored Fe 3+ - salicylate complex with any remaining salicylic acid impurities, but will not complex with the aspirin. You will use absorption spectroscopy to measure the amount of the complex formed and determine the amount of impurity of your aspirin.
Experimental Your team will need to split into two groups. ne group should be responsible for synthesizing the aspirin while the other group prepares the standard solutions and plots the curve you will use to analyze your product s purity using absorption spectrometry. Part I: Preparation of aspirin Fill a large 600-mL beaker with about 450-500 ml of water and begin heating this over a Bunsen burner using a ring stand and wire gauze. You will need the water to be around 75. Fill your squirt bottle with deionized water and place it inside a large beaker of ice water to chill. Now weigh approximately 2.0 grams (record the exact mass) of salicylic acid into a 50 ml Erlenmeyer flask. Measure out 5.0 ml of acetic anhydride in your graduated cylinder and pour it into the flask in such a way as to wash any crystals on the walls down into the bottom. Add five drops of 85% phosphoric acid to serve as a catalyst. aution: both the acetic anhydride and the 85% phosphoric acid can cause chemical burns. Be sure to handle these substances with caution and wash your hands after use. nce the temperature of the water bath reaches 75, clamp the flask in place inside the water bath and heat for at least 15 minutes. Stir occasionally. Maintain the temperature of the water at 75 throughout this process. At the end of this time cautiously add about 2 ml of deionized water to the flask to decompose any excess acetic anhydride. aution: There may be some hot acetic acid vapors evolved as a result of this decomposition. When the liquid has stopped giving off vapors, remove the flask from the water bath and add 20 ml more of water. Let the flask cool on the bench for about 5 minutes, during which time you should see crystals of aspirin beginning to form. After 5 minutes, place the flask in a beaker of ice water to hasten the crystallization and increase the yield of products. Allow the flask to cool in the ice water for an additional 5 10 minutes. If crystals are slow to form it may help to gently scratch the inside of the flask with a glass stirring rod.
ollect the aspirin crystals by filtering the cold liquid through a Buchner funnel using suction. Be sure to weigh the dry filter paper first. Use the chilled water from your squirt bottle to wash any remaining crystals from your flask into the funnel while filtering. Rinse the crystals several times using the chilled water from your squirt bottle to help remove any impurities. Finally, draw air through the filter for several minutes to help hasten drying of the crystals. Place the filter paper under a heat lamp to dry. Allow the crystals to dry for at least 30 minutes. Weigh the product and then heat for an additional 10 minutes. The crystals are dry when subsequent weighings remain the same (i.e. all the water has been evaporated). The dry crystals should have a light flaky appearance. lumpy crystals suggest that they are still wet and need further drying. It may help speed drying to break up clumps of crystals using your stirring rod. Part II: Testing the purity nce the crystals are completely dry you may test the purity using the Mel-Temp apparatus. Your instructor will demonstrate the proper use of this device. You will also need to determine the purity of your crystals using absorption spectrometry. For this you will need to make a set of five or more standard salicylic acid solutions in the range of 10-3 to 10-4 M. You will need to prepare these standards using dilutions in the same way the standards were prepared in the Manganese lab (refer to the information below, the Manganese lab on-line handout and the on-line prelab for this experiment for how to prepare these solutions). In this case the samples will be purple and so we will measure all the data at a wavelength of 525 nm. For this experiment we will use 50.00 ml volumetric flasks. btain 2 from the stockroom. Use a few ml of acetone (not water) to rinse the flasks if necessary. Salicylic acid is not soluble in water, but salicylic acid dissolved in acetone is. Thus, to make your stock solution you should first calculate how much salicylic acid you will need, weigh this amount directly into 50.00 ml volumetric flask, and then completely dissolve the salicylic acid in 10 ml of acetone solution before adding acetone to the mark. To make up the diluted standards, pipette your calculated quantity of stock solution into a second 50.00 ml volumetric flask. Add 10.00 ml of acetone using a 10-mL pipette to be sure the salicylic acid remains in solution. Then add 10.00 ml of 0.025 M Fe(N 3 ) 3 solution using the pipette, and fill to the mark with deionized water. Measure the absorbance of these standards (what should you use as a blank?) and create a Beer s Law plot.
When testing the purity of your aspirin product you will need to weigh out approximately 0.10 grams (accurately record the exact mass) of your aspirin crystals into a clean 50-mL volumetric flask. Then pipette 10.00 ml of acetone into the flask and completely dissolve the solid product to be sure that everything goes into solution. Next add 10.00 ml of 0.025 M Fe(N 3 ) 3 solution and fill to the mark with deionized water. Measure the absorbance of this solution. If the absorbance is greater than 1.2 absorption units, you will need to repeat this procedure using a smaller mass of your product. You can now use your Beer s law plot to determine the molarity of the salicylic acid impurity in this solution, and from this the percent mass of salicylic acid impurity in your product. The percent purity of your aspirin can be determined by subtracting the percentage of impurities from one hundred. Lab Report The report for this experiment is an abbreviated lab report and need only include: Data, Results, References, Discussion and onclusions section. No introduction or method is required. The report must be type written. A short 2-5 page report is acceptable. Be sure to include your Beer s law plot and a table detailing how you prepared your standards. Discuss the yield and purity of your and include a discussion of error analysis. Finally make your recommendations to the company regarding the practicality of using this synthesis to make aspirin and any further testing they may wish to do in this regards.