Chemistry Paper on Laboratory Report: Aspirin Synthesis Package
One of the most important roles associated with organic chemistry is synthesis of chemical products. Organic chemists can synthesize several products both in the laboratory and in the industrial scale. The process of synthesis appears simple when performed yet involves a series of complex reactions. Any synthesis has to involve formation of new products and by products through transformation of certain raw materials. In each of these, some bonds are broken while others are created. The creation of new bonds results in the production of energy in most reactions while the breaking of bonds consumes energy in most of the reactions. Since breaking of bonds precedes the creation of new bonds in any reaction, it is crucial for the energy required for bond breaking to be made available through heating and/ or through provision of reaction energies in other ways. The amount of energy required for bond breaking is dependent on the strength of the bonds to be broken. However, in all cases, the minimum energy requirement for bond breaking in all reactions is referred to as the activation energy.
The rate of reaction is represented by the reaction rate constant which is dependent on the difference between the energy requirement EA and the available energy RT. In the reaction for the synthesis of aspirin, the activation energy is provided through heating in a water bath. The objective of the experiment carried out was to synthesize and purify acetylsalicylic acid (aspirin) from salicylic acid using acetic anhydride and phosphoric acid as a catalyst. After synthesis, the required product will be isolated through crystallization, its purity confirmed through IR and quantified through yield calculations. It is expected that based on the mole ratios in the reaction, the yield from pure reactants should be more than 50% if the reaction conditions are favorable. The diagram below represents the intended reaction.
Figure 1: Reaction Equation (Source: Arizona State University 3)
The laboratory exercise carried out involved three key stages. The first stage was the synthesis of acetylsalicylic acid; this was followed by separation of the product through crystallization and finally purification of the product.
- Salicylic acid
- Acetic anhydride
- Phosphoric acid
Synthesis of acetylsalicylic acid
Approximately 25mg of salicylic acid was measured into a test tube and the exact amount recorded. 0.5 ml of acetic anhydride was then added to the test tube containing the acid. To the mixture, a drop of 85% phosphoric acid was then added to the mixture. The test tube with the three chemicals was then placed in a beaker containing water at 40 – 45⁰C and the mixture stirred to a solution. Heating was continued in the water bath for approximately 15 minutes for the reaction to reach completion. The diagram below shows the set up for aspirin synthesis.
Figure 2: Synthesis of Acetylsalicylic acid (Source: Arizona State University 6)
Isolation of the Product
The test tube was removed from the water bath while still warm and 1.5ml of de-ionized water added to it instantly. The test tube was then shaken for a few minutes until the mixture was clear then cooled to room temperature. At this point, some crystals had begun to form. The test tube was then cooled further in an ice bath while at the same time, a Hirsch funnel was being set up. The test tube was removed from the ice bath as soon as the crystals were sufficient and the mixture had not begun to warm again. The mixture was then filtered to isolate the product from the mixture. The diagram below shows how this was done.
Figure 3: Isolation of the products from the reaction mixture (Arizona State University 6)
Purification of the product
The impure aspirin was transferred to a clean Craig tube. After this, ethanol was boiled in a small test tube and the boiled ethanol added drop wise to the Craig tube containing the impure aspirin. The amount of ethanol added was just enough to dissolve the aspirin. During the process, the Craig tube was kept in the beaker to maintain the right temperatures. A white inner plug was then placed in the Craig tube containing the impure solution of aspirin and the set up placed in a 25ml Erlenmeyer flask and let to cool for between 5 and 10 minutes. At this point, some crystals of pure benzoic acid had begun to form. The Craig tube with the inner plug was transferred to an ice bath and placed firmly to prevent it from falling when ice melts. This enhanced crystallization of the pure aspirin. This went for 15 minutes and then the crystals were separated from the ethanol through centrifugation. To do this effectively, the inner plug was tightly wrapped with a copper wire. The Craig tube was then placed into a plastic centrifuge tube. This was done while ensuring that the inner plug went in first. The remaining copper wire was wrapped around the plastic centrifuge tube to hold it in place. After separation of the crystals, the set up was disassembled and the pure aspirin scrapped off onto a weighed watch glass. The crystals were dried completely by passing air over them before being weighed with the watch glass (Patnaik 1.16). The weight of the aspirin produced was equal to the difference between the final weight measured and the initial weight of the watch glass.
The salicylic acid used in the reaction has a molar mass of 138.121g/mol. The weight of the acid at the start was 0.254g. As such, the number of moles of acid at start is equivalent to:
The end product of the reaction i.e. acetylsalicylic acid is of a molar mass of 180.157g/mol. The obtained aspirin was 0.065g after drying. The number of moles of aspirin produced was thus equal to:
Based on the reaction equation shown in figure 1, the mole ratio for the reaction is 1:1, reactants to products. From this, the percentage yield was calculated through the equation below given by Vogel (774):
For a 100 percent yield, the number of moles of aspirin produced could have been equal to the number of moles of acid used in the reaction. This implies that 100% yield could only be achieved with a product weight equal to:
Based on the results obtained, it can be said that the exercise was successful although not to a large extent. The calculated yield was approximately 19% based on percentage mole calculations. According to the reaction mechanism, it is expected that one mole of salicylic acid should produce one mole of acetylsalicylic acid. The reaction mechanism is as shown in the diagram below. 100% yield is signified by the production an equal number of moles of aspirin to the number of moles of the acid used. The recorded low yield can be attributed to various factors such as presence of impurities in the acid used, incomplete reactions in the synthesis process and incomplete separation in the isolation and purification phases. To enhance productivity, recovery procedures are often used on the extracted liquor following isolation and purification phases.
Figure 3: Reaction Mechanism
During the reaction, the phosphoric acid was added as a catalyst to lower the activation energy associated with the reactants in the process. The heating was performed to provide thermal energy since the reaction in progress was an endothermic reaction. In such reactions, the activation energy required for bond breaking is higher than the energy of reaction hence additional energy was required for the process. In the endothermic reaction, the activation energy is constant for the specific bonds to be broken while the energy of reaction varies depending on the number of bonds broken and the number and strength of bonds created in the process.
In the first phase of the experiment, water was added to isolate the aspirin from the other products since aspirin is hydrophobic and would be suspended. This was contrary to the second phase of purification where alcohol, a protic solvent which acts through the loss of a hydrogen atom, was used (Mark 317). Ethanol dissolves acetylsalicylic acid since it has the potential for gaining a proton hence is the most suitable solvent for purification purposes.
The objective of the experiment was to synthesize acetylsalicylic acid (aspirin). The reaction is endothermic and required heating to complete the synthesis. Moreover, phosphoric acid was also used as a catalyst to speed up the reaction by lowering the activation energy. Although it was expected that the reaction would have high yield of aspirin, the results showed a yield of less than 20%. The probable reasons for this include low purity of reactants, incomplete reactions during synthesis and incomplete separation during the isolation processes. Such reactions can be improved in future through confirming the purity of reactants prior to use and using more efficient separation strategies for the isolation stages.
Arizona State University. CHM 237: Organic Chemistry Lab 1. Arizona State University, 2016.
Mark, Loudon. Organic Chemistry 4th Ed. New York: Oxford University Press. 2002.
Patnaik, Pradyot. Dean’s Analytical Chemistry Handbook (2nd ed). New York: McGraw Hill Companies, 2004.
Vogel, Arthur. A Text- Book of Practical Organic Chemistry including qualitative Organic Analysis (3rd ed). London: Longman Group of Companies, 1956.