Experiment 19: Synthesis of Aspirin and Oil of Wintergreen Discussion
The main purpose of this lab is to allow students to have the opportunity to observe the synthesis of various organic compounds, namely aspirin and the oil of wintergreen. This is done through utilization of the processes of esterification between an alcohol and an acid. After obtaining the synthesized products, purification techniques were used in order to create pure aspirin. The concepts of recrystallization, standardization of titrants, melting point tests, and back titration are all utilized throughout this lab in order to determine and obtain data associated with the purity and crudeness of both aspirin as well as oil of wintergreen.
Utilizing 2.545 grams of salicylic acid and mixing it with 5mL of acetic anhydride, 6.051 grams of aspirin (acetylsalicylic acid) was obtained. The synthesis of this reaction obtained a percent yield of 98.19%. The theoretical yield for aspirin synthesis is 6.15grams and the actual yield was 6.051 grams. In order to test the purity of the product, a ferric chloride (FeCly) solution is added. This ferric chloride solution tests for the presence of phenols, which is composed of families with an OH group, and if that phenol group is present, a magenta color is observed. Since aspirin doesn’t have a hydroxyl group, no color will form because no salicylic acid is present. Therefore, when a magenta color is observed it is primarily due to the impurities of salicylic acid. In order to further purify the product, a method referred to as recrystallization is utilized. The percent yield for recrystallized aspirin recovered is 13.75%. This yield is significantly lower than part one primarily because the process of recrystallization is attempting to remove the unreacted salicylic acid impurity from the product, thus we obtained a lower amount of yield. Additionally, if then aspirin wasn’t completely dry before weighing it, then the water will add to the mass of the crude product and make the % yield higher than if a lower yield should have been expected. Another source of error could be associated with how the lack of purification of our recrystallized product was most likely due to leaving our solution in the ice bath for a long time, until most of the solution was recrystallized. Because of this, some of the salicylic acid probably was able to recrystallize along with the aspirin product, no longer separating the two.
The melting range point observed for the crude aspirin is 22 degrees Celsius, while as the melting point observed for the pure aspirin is 14 degrees Celsius. The pure aspirin’s melting point is 125 degrees Celsius and has a 7.4 % error when it is compared to the known melting point of 135 degrees Celsius. The crude aspirin’s melting point is 102 degrees Celsius and has a 24.44% error when it is compared to the known melting point of 135 degrees Celsius. The difference between the melting point ranges is primarily associated with the intuition that if a sample is not completely pure, the melting point is lower and the temperature range associated with the melting point is larger. Therefore this confirms that the pure sample had a lot less impurities than the crude sample simply because the pure sample was close to the known melting point, and the crude sample was a lot farther from the known melting point meaning it had a lot of impurities.
In part 6, KHP is titrated with a solution of NaOH to determine the exact concentration of the NaOH. The average concentration observed for the standardized NaOH is 0.096755 M and the average deviation is 0.004626, therefore further demonstrating that the 3 samples differed by less than the required one percent. NaOH solution is standardized in order to determine its exact concentration. The standardized NaOH is effectively used to determine later on, the quantity of aspirin in the sample. Specifically, the standardized molarity of the NaOH(average molarity) was used to come up with mmoles of NaOH titrated and to help come up with the mmoles of NaOH in hydrolysis.
Through the addition of 1.003 grams of salicylic acid and 5mL. of methanol and 3 drops of concentrated sulfuric acid. The resulting solution after being heated and then cooled still contains unreacted salicylic acid in the solution. We added a drop of 1% ferric chloride to the solution and it turned purple/magenta. This occurred because the hydroxyl group of the salicylic acid attaches to the aromatic ring in the ferric chloride which produces the colored complex (purple/magenta color). This showed us that there was still unreacted salicylic acid in the solution.
In part 7, we analyzed the purity of the aspirin by doing a back titration of the aspirin This is used to determine the amount of acetylsalicylic acid in the sample. To do this we used 2 samples of crude aspirin and 2 samples of pure aspirin. Their mass was respectively, crude: 0.475g, 0.551g and pure: 0.512g, 0.466g. After the back titration of the 4 samples, we were able to calculate percent aspirin in the 4 samples. The average mass of the aspirin was 0.219g. The average yield of aspirin in the crude sample is 0.0765 g, while as in the pure sample the average yield of aspirin is 0.36185g. The percent yield for aspirin in the first crude sample is 3.74% and in the second crude sample it is 24.52%. The percent vield for aspirin in the first pure sample is 61.58% and for the second pure sample it is 87.74%. Thus proving that our pure samples were had a lot less impurities than the crude samples. A source of error could be associated with how the stopper was not continuously placed down all the way in the Erlenmeyer flask containing the 0.1M NaOH.
