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Experiment 18 Discussion
Organic Chemistry Lab II (CHEM201401)
Boston College
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In this experiment we synthesized esters. We first synthesized cholesteryl benzoate through the reaction of an alcohol (cholesterol) and an acid chloride (benzoyl chloride). When the two were reacted with each other in the presence of a base, pyridine, a cloudy white precipitate formed within the solution, indicating that the esterification had occurred. The precipitate was then separated then recrystallized. The percent yield for this reaction was 25%. The melting range of the product was 146–147°C which is very close to the standard melting range of cholesteryl benzoate, 149–150°C, which indicates that the product was mostly pure, but that there were still some impurities present. I also found the melting point of my purified product mixed with cholesterol which was 125–147°C. This is a very large range which is explained by the presence of the cholesterol. Due to its chirality, cholesteryl benzoate is a liquid crystal meaning that the molecules are ordered to a certain extent in the liquid phase. The crystals were heated and cooled, and when they were cooled, under the light they looked mainly blue. This is because the crystal structure can change through the spacing between the layers when there is a temperature change and thus the reflected wavelength can change. Next, we synthesized acetylsalicylic acid (or aspirin) through a reaction between an alcohol (salicylic acid) and an acid anhydride (acetic anhydride). When the two were reacted with each other in the presence of an acid (phosphoric acid), the solution turned from white to yellow, and all of the solid dissolved when heated. When water was added, all of the yellow color disappeared leaving a white and cloudy mixture, indicating that the esterification into acetylsalicylic acid had occurred. The product was separated and then recrystallized.
The percent yield of the product was 34%, and the melting range was 135–138°C, which is very close to the actual melting range of 138–140°C, which indicates that the product was mostly pure but that there was still some impurities present. I also found the melting range of commercial aspirin tablet which was 133–139°C. The low starting melting point can be explained by some substance used as a binder that is acting as an impurity. I performed solubility tests on my purified acetylsalicylic acid product and the aspirin tablet. I found that my product was insoluble in toluene, hot water, after being rubbed with a glass rod, and in hydrochloric acid, and it was soluble in sodium hydroxide. The commercial aspirin was insoluble in toluene and in hot water, which matches my results. Ferric chloride tests were also performed on salicylic acid, my crude acetylsalicylic acid product, and my purified acetylsalicylic acid product. These tests are used to test for the presence of a phenol group which is a benzene ring with an -OH group directly attached. If a solution turns purple upon the addition of ferric chloride, there are phenols present. As salicylic acid contains a phenol, when it was tested with ferric chloride, it turned purple. When my crude product was tested with ferric chloride, it also turned purple, indicating that there was still unreacted salicylic acid present because acetylsalicylic acid shouldn’t react as it doesn’t contain a phenol. When I tested my purified acetylsalicylic acid product, it also turned purple. This indicated that I hadn’t completely purified out all of the salicylic acid, which is supported by the melting range being slightly off and the low yield. In test tube 4, I first reacted my purified acetylsalicylic acid product with hydrochloric acid and heated it, then I added ferric chloride and the solution turned dark purple. This happened because HCl is a strong acid that leads to the decomposition of acetylsalicylic acid back into salicylic acid and acetic acid, and as salicylic acid contains a phenol group, it reacted with the ferric chloride and became dark purple.
Experiment 18 Discussion
Course: Organic Chemistry Lab II (CHEM201401)
University: Boston College
