Acetylsalicylic Acid LabThis print version free essay Acetylsalicylic Acid Lab.
Autor: reviewessays 03 December 2010
Words: 1364 | Pages: 6
September 15, 2005
September 23, 2005
The objective of this lab experiment is to prepare aspirin (acetylsalicylic acid) by reacting salicylic acid and acetic anhydride.
This purpose of this lab was to prepare aspirin using its basic components: salicylic acid and acetic anhydride. Salicylic acid and acetic anhydride, along with an acid catalyst, react to form acetylsalicylic acid (aspirin) and acetic acid. In this reaction, the hydroxyl group on the benzene ring in salicylic acid reacts with acetic anhydride to form an ester functional group. The acetylsalicylic acid crystallizes as the solution cools and begins to come out of the solution. The acetylsalicylic acid is further purified through recrystallization with ethyl acetate.
Test Tubes Pasteur Pipette
Salicylic acid acetic anhydride
Sulfuric acid 125 mL Erlenmeyer flask
Scale Water Hot water bath
Ice Bath Buchner FÑŒnnel
1% Ferric Chloride
Ethyl acetate Glass Rod
Iodine Aspirin Tablet
Commercial Aspirin tablet
1. 2.0 g of salicylic acid was placed in a 125 mL Erlenmeyer flask.
2. 5.0 mL of acetic anhydride was added to the flask and five drops of concentrated sulfuric acid was also added.
3. The flask was swirled to dissolve the salicylic acid.
4. The flask was placed in a hot water bath for ten minutes.
5. The flask had to be placed in an ice bath briefly and the sides were scratched with a glass rod until crystallization occurred.
6. 50 mL of water was added and the flask was returned to the ice bath for another five minutes. This allowed for more crystallization to occur.
7. The solution was placed in vacuum filtration using a bÑŒchner funnel for five minutes. Distilled water was added to the flask to flush solution off the sides of the flask to be poured into funnel for filtration.
8. The remaining product was allowed to air dry and was weighed. The product weight was 2.88 g.
9. The percentage yield of the product was 11.02 percent.
This experiment was performed to determine if their was any salicylic acid left in the product.
1. Three test tubes were filed with .5 mL of water.
2. Small portions of the product were added to two test tubes, while the third was filled with water to act as the control.
3. One drop of Ferric Chloride solution was added to each tube.
4. The test tubes containing product turned into a tan/peach color while the control remained yellow. The concluded that the product had little to no salicylic acid present.
1. 3 mL of hot ethyl acetate was added to a small Erlenmeyer flask.
2. Product from the previous experiments was added to the flask.
3. The solution was placed in a hot water bath until dissolved and then placed in an ice bath and scraped with a glass rod until crystallized.
4. The solution was vacuum filtered until dry.
5. The ferric chloride purity test was performed on the new product. The same results from Part B occurred, further proving that the previous solution had little to no salicylic acid.
6. The melting point was determined by using a â€œmelting point apparatusâ€(?) The melting point of the product was 131Â°C.
1. The purified product, Ñ˜ an aspirin tablet, and Ñ˜ a commercial aspirin tablet were each placed in separated test tubes.
2. 2 mL of water was added to each tube and each tube was boiled in a hot water bath.
3. A drop of iodine solution was added to each test tube to test the presence of starch. Starch was only present in the regular aspirin tablet, which turned into a violet blue color after having the iodine added to it.
DATE, RESULTS, AND CALCULATIONS
â€¢ 2.0g Salicylic acid(s) + 5.0 mL acetic anhydride(aq) + 5 drops sulfuric acid  2.88g acetylsalicylic acid + acetic acid Percentage yield= 11.02%
Test Tubes w/ ferric chloride Product Added? Result
Tube #1 Yes Solution turned tan/peach color (pure)
Tube #2 Yes Solution turned tan/peach color (pure)
Tube #3 (Control) No Solution turned yellow
Recrystallization was performed using ethyl acetate
â€¢ The melting point of the product was 131Â°C.
â€¢ Ferric chloride purity test results were also the same.
Product + water After Iodine is added.
Acetylsalicylic acid No color change
Ñ˜ aspirin Violet blue color change
Ñ˜ commercial aspirin No color change
In Part A, aspirin was prepared by the reaction between salicylic acid and acetic anhydride. This reaction is possible because of the two factors: The hydroxyl group on the benzene ring in salicylic acid reacting with the acetic anhydride to form an ester function group, which is an acid derivative with an alkyl group replacing the acid proton, and the presence of an acid catalyst, sulfuric acid. When the reaction is complete, acetylsalicylic acid and acetic acid are the products. The most likely impurity is salicylic acid, which could be due to incomplete reaction or hydrolysis.
In Part B, the Ferric Chloride purity test was performed on the product. As said before, salicylic acid could still be present in the product due to a incomplete reaction or hydrolysis since salicylic acid is a phenol, meaning it processes a hydroxyl group on the benzene ring. Phenols form a highly colored complex with ferric chloride and because aspirin is not a phenol, salicylic acid is easily detected. After performing the test, it was discovered that our product had no salicylic acid and was for the most part pure.
In Part C, recrystallization, the process of the re-crystallizing a solid in order to obtain further purity, was performed on the acetylsalicylic product obtained in Part A. Hot ethyl acetate was used as a solvent this time because water decomposes aspirin when heated. After the solution was crystallized, the ferric chloride purity test was performed again, only to yield the same results as before, which prove the experiment was properly performed. A sample of the product was placed inside of a melting point apparatus to determine the melting point and compare it to the melting point of pure aspirin. Our result was 131Â°C, only a couple of degrees short of the normal melting point range 135-136Â°C.
In Part D, the product, a normal aspirin tablet, and a commercial aspirin tablet were each tested to determine if starch is present. Small samples of each were boiled in water and a drop of iodine was placed in each tube. If starch is present, the solution will become a violet-blue color. The only solution to become violet-blue was the normal, store brand aspirin tablet. This is because aspirin tablets are pressed together with small amounts of binding material, such as starch, methylcellulose and microcrystalline cellulose. These are normally the inactive ingredients.
This lab shows how compounds containing hydrocarbons can be combined with the help of an acid catalyst. This lab also shows that crystallization can occur when a compound with a high melting point begins to get insoluble at room temperature.
1. What is the purpose of the concentrated sulfuric acid used in the first step? The concentrated sulfuric acid was used as a catalyst to aid in the reaction.
2. What would happen if the sulfuric acid were left out? Without the presence of the H+ cation, the acetic anhydride would not have split and acetic acid would have formed.
3. If you used 5.0 g of salicylic acid and excess acetic anhydride in the preceding synthesis of aspirin, what would be the theoretical yield of acetylsalicylic acid in moles? In grams?
4. What is the equation for the decomposition reaction that can occur with aspirin? The equation is: acetylsalicylic acid +water + heat -> salicylic acid + acetyle anhydride.
5. Most aspirin tablets contain five grains of acetylsalicylic acid. How many milligrams is this? 325 milligrams.
6. A student performed the reaction in this experiment using a water bath at 90Â°C instead of 50Â°C. The final product was tested for the presence of phenols with ferric chloride. This test was negative; however, the melting point of the dry product was 122-125Â°C. Explain these results as completely as possible.
7. If the aspirin crystals were not compleely dried before the melting point was determined, what effect would this have on the observed melting point? The melting point would have been lower because the product would still be partially liquid, making it easier to melt.
A good follow up experiment to this lab would be to synthesize other over-the-counter drugs, such as ibuprofen and acetaminophen.
Pavia, L Donald et al. (2005). Introduction to Organic Laboratory Techniques, A small scale approach. 2nd edition (57-61 & 662-667).