A. Abstract

Aspirin is the common name for the compound acetylsalicylic acid, widely used as a fever reducer and as a pain killer. The first part of the experiment aims to synthesize aspirin from the reaction of salicylic acid with acetic anhydride with the aid of phosphoric acid as a catalyst. The second part of the experiment aims to assess the purity of aspirin through the determination of its melting point and comparing it to the accepted value. An oil bath was used in the determination of the melting point. Results showed 79.56% yield of impure aspirin. It was also determined that its melting point range is about 125˚C - 130˚C while that of the pure aspirin is about 130˚C – 134.5˚C which is closer to the accepted value, 136˚C. The latter also has a narrower range which shows that it really is purer than the crude one, although not pure enough when compared with commercial aspirin.

We will write a custom essay sample on

Aspirin Synthesis specifically for you

for only $13.90/page

Order Now

B. Introduction

Acetylsalicylic acid or more commonly known as aspirin is one of the most widely used medications to reduce fever (an antipyretic), to reduce pain (an analgesic), and to reduce swelling, soreness, and redness (an anti-inflammatory agent). It is a white, crystalline substance which melts at a temperature of about 136˚C.

Aspirin is synthesized through the reaction of salicylic acid with an excess of acetyl anhydride which causes a chemical reaction that turns salicylic acid's hydroxyl group into an acetyl group. A small amount of strong acid such as Phosphoric acid is used as a catalyst which speeds up the reaction.

The reaction for the synthesis of acetylsalicylic acid or aspirin is shown in Figure 1.

Figure 1.

Since acetic acid is very soluble in water, it is easily separated from the aspirin product. The aspirin isolated in this step is the crude product. This crude product is the one desired for the first part of the experiment. A purified product can be obtained through recrystallization of the crude product in hot ethanol.

Recrystallization is the primary method for purifying solid organic compounds. Compounds obtained from natural sources or from reaction mixtures almost always contain impurities. The impurities may include some combination of insoluble, soluble, and colored impurities. To obtain a pure compound these impurities must be removed. Each is removed in a separate step in the recrystallization procedure.8 The recrystallized aspirin, which in theory is purer, will be used in the second part of the experiment along with the crude one.

Since the melting point of a solid can be easily determined even with small amounts of the compound, it’s a physical property that is often used to identify and characterize a compound. The purity of a compound can also be assessed through the determination of its melting point. An impure compound generally melts at temperatures lower than that of the accepted or true value of the melting point of the compound. It also has a wider range of melting point than that of a pure compound. Purer compounds usually have a melting point range of 1-2˚C only. If the supposed pure aspirin melts at a wide range, this indicates that further purification is needed or the compound is not aspirin at all.

C. Review of Related Literature

This section presents citations of literature and related studies and experiments regarding this topic. Aspirin is an effective analgesic (pain reliever), antipyretic (fever reducer) and anti-inflammatory agent and is one of the most widely used non-prescription drugs. The use of aspirin had its origin in the 18th century, when it was found that an extract from the bark of willow trees which contains salicylic acid was useful in reducing pain and fever. In folk medicine, willow bark teas were used as headache remedies and other tonics.4 Although salicylic acid was effective at reducing pain and fever, it also had some unpleasant side effects. It is irritating to the lining of the mouth, esophagus, and stomach, and can cause hemorrhaging of the stomach lining.

Acetylsalicylic acid, a weaker acid than salicylic acid, was found to have the medicinal properties of salicylic acid without having the objectionable taste or producing the stomach problems. The acetyl group effectively masks the acidity of the drug during its ingestion and after it passes into the small intestine, it is converted back to salicylic acid where it can enter the bloodstream and do its pain relieving action.2 Aspirin can be synthesized in the laboratory through a simple reaction between salicylic acid and acetic anhydride. This type of reaction actually has a specific name, esterification. An esterification reaction usually combines an organic alcohol and an organic acid.

Organic functional groups are important because they help to determine the behavior of the compounds to which they are attached. For example, most acids have a sour taste (like citric acid in lemons) and many esters have odors (like intergreen or vanilla). During the reaction process, a molecule of water splits off and the remaining carboxylic acid and alcohol fragments become attached producing an ester. In addition, a small amount of phosphoric acid will be added to help the reaction occur at a reasonable rate.3 Nowadays, salicylic acid is administered in the form of aspirin which is less irritating to the stomach than salicylic acid.

The most common method of purifying solid organic compounds is by recrystallization.6 Recrystallization, also known as fractional crystallization, is a procedure for purifying an impure compound in a solvent. The method of purification is based on the principle that the solubility of most solids increases with increased temperature. This means that as temperature increases, the amount of solute that can be dissolved in a solvent increases.7 Most organic compounds have a sharp melting point, which can be measured accurately to within 1oC or better. Furthermore, the measurement is easily made with a small quantity of material (a few small crystals) using a simple apparatus. Melting point determinations are routinely made on solid organic compounds, and extensive compilations of melting points are available in reference books. One use of the melting point is to establish that a preparative or isolation procedure has led to an expected product.

A very pure substance has a very sharp melting point. Further purification will not change the melting point. Less pure substances melt over a range of temperatures that is below the actual melting point of pure material. Thus the sharpness of a melting point is a useful criterion of purity. When a melting point is determined, it is therefore important that the melting range be recorded. The bottom of the melting range is the temperature at which the first signs of liquid can be seen. The top of the melting range is the temperature at which the last of the solid fuses, i.e. turns into liquid. The compound is generally regarded as pure enough for most purposes if the melting range is no greater than 2oC. A wide melting range signals the need for further purification. The purity of the aspirin produced in the experiment was tested using this same method.

D. Methodology

A. I. Preparation of Reaction Aspirin Weigh 3.0 grams of salicylic acid and place in a 250 mL Erlenmeyer flask. Add 6.0 mL of acetic anhydride. (Do this in the hood.) Carefully add 5-10 drops of phosphoric acid and swirl t mix everything thoroughly. Heat the mixture for about ten minutes in a warm water bath (70 ˚C - 80 ˚C).

After the heating process, add 20 drops of distilled water. Then add another 20 mL of distilled water. Cool this in an ice bath. Scratch the walls of the flask with a stirring rod to induce crystallization. Filter the solid aspirin using a pre-weighed filter paper. Wash the crystals with 2-3 mL of chilled water. Dry the solid and the filter paper.

A. II. Determination of Yield Pre-weigh a watch glass. Place the filter paper with the solid aspirin on it and weigh. Obtain the weight of the weight of the aspirin. This impure aspirin still contains traces of water and salicylic acid. Place a small amount of this impure aspirin in a beaker, cover with tissue and set aside for further purification.

A. III. Recrystallization of the Aspirin Transfer the remaining impure aspirin into a 100 mL beaker. Add 10 mL of 95% ethanol. Warm the mixture in a water bath to dissolve the crystals. (Note: Do not boil!) If not all crystals dissolve, add 5 mL more of the ethanol and continue to warm the mixture. When all crystals have dissolved, add 20 mL of warm water. Cover the beaker with a watch glass and allow the mixture to cool slowly. Crystals of aspirin should form. Continue cooling the solution in an ice bath. Filter the pure aspirin using a pre-weighed filter paper again. Dry the crystals and place them in a beaker. Cover it with tissue and allow these crystals to dry further.

B. I. Filling a Capillary Tube Fill the capillary tube by pressing the open end in a mass of the crystals of the substance. Drop the tube, open end up, down a glass tubing about 1 cm in diameter onto a hard flat surface. This will allow the crystals to go down the bottom of the capillary tube. The sample should be tightly packed to a depth of 2-3 mm.

B. II. Oil Baths Pour oil, enough that the height of the sample will be fully immersed, into a beaker. Add some boiling chips and heat.

B. III. Determination of Approximate Melting Point Place a pinch of the sample onto the bulb of the thermometer. Heat the sample until it completely melts. B. IV. Actual Melting Point Range Prepare 2-4 capillary tubes and seal one end of each with a blue flame. Pack the capillary tube with the sample aspirin. 2 capillary tubes will be for the impure sample and the other two will be for the pure. Attach these to the thermometer bulb. Immerse the samples into the oil bath. Stir the oil occasionally to ensure uniform distribution of heat. Watch the solid in the capillary tube carefully. The temperature at which the sample just starts to melt and the temperature at which all the sample has completely turned to liquid must be recorded.

E. Discussion of Data and Results

F. Conclusion and Recommendation The reaction of salicylic acid and acetic anhydride produces acetylsalicylic acid and a by-product acetic acid. As expected of organic compounds, the percent yield was below 100%. Excess reactant acetic anhydride and by-product acetic acid was washed from the produced aspirin. It is inevitable that impurities will be present in the products of reactions. This product requires purification before it can be used for purposes other than experimentation. Recrystallization is a very useful method in the purification of a substance.

This “purified” substance or other substances for that matter can also be tested of its purity through the determination of its melting point. The melting point of the crude aspirin and the pure aspirin were compared to the accepted value. Despite the difference in the melting point range, it was found that the melting point of the pure aspirin is logically nearer to the true value. With this result, it can be said that the “purified” aspirin still contains some impurities and still needs to undergo further purification if we wish to extend the aim of this experiment.

In addition, based on our experiences with this experiment, it is recommended that carefulness is important for reactants and the catalyst used in this reaction can cause irritation to any part of the body that these came in contact with. Proper packing of the capillary tube must also be observed for the air that will be in between the particles of the sample will cause the reading on the thermometer to be higher than true temperature. Careful observation of the changes that happen during the reaction process and the melting point determination must also be done. This is important for the accurate reading and recording of data for the experiment.