Experimental Overview:
In this experiment, I was making a sample of aspirin and then testing it in order to see how pure the sample of aspirin was. By doing this experiment, I was leaning how to crystalize products, and then used the theoretical yield, along with the percentage yield in order to calculate the amount of aspirin that I had created in the sample. Aspirin is an anti-inflammatory, and analgesic, meaning this medication can reduce inflammation, fever, and pain by blocking the enzymes that promote these issues, and reducing the production of more of these enzymes all over the body.
When making an organic compound, it is important to separate the compound from other compounds to make it as pure as possible, this is called crystallization.
…show more content…
When crystals begin to form it is because the molecules of the compound have weak attractions to one another and form molecules, this leads to very pure crystal products. You can do this by creating a very hot environment and saturating a compound, as it cools the compounds solubility decreases, and the molecules begin to come together to form crystals. In order to get the rest of the impurities out of the sample you simply filter the solvent. In this particular experiment we used salicylic acid as the limiting reactant in the synthesis reaction. So that it could be used up and then the amount of aspirin could be calculated by the amount of salicylic acid that was started in the reaction. So we use the theoretical yield of aspirin can be calculated by how much salicylic acid is used in the reaction. The we look at the percentage yield, also know as the percentage recovery, this is the amount recovered compared to how much was expected by using a calculation. Our final product was expected to have a small amount of unreacted salicylic acid in it. Materials and Methods: As soon as I got to my station I got out the hot plate and turned it on all the way to start warming it up in order to boil water. I then got a 400 mL beaker, and put about 150 mL in it, the set the beaker on the hot plate and waited for the water to boil. While I was watching the water my waiting my partner weighed out exactly 3.0 grams of salicylic acid on some weighing paper and then pored it into a clean and dry 125 mL erlenmeyer flask. I then took that to the fume hood and added 8mL of acetic anhydride, and swirled it under the hood.
Then, I added 8 drops of concentrated phosphoric acid to the mixture. swirling it a few times. Then, I carefully took the flask to the station as I avoided trying to breath the vapors of the acetic anhydride. I put the e-flask into the beaker of water sitting on the hot plate in order to heat it for seven minutes. Once the seven minutes was up, my partner carried the e-flask to the fume hood, and added 3 mL of de-ionized water to the flask. She swirled it for a couple of minutes there. She brought it back tot he station where I gradually added 60 Ml of de-ionozed water to the mixture while my partner stirred the mixture constantly. I was able to see some of the aspirin beginning to form. In order to complete the crystallization process we cooled the flask in an ice-water bath from 4:00 until 4:20. As we waited I began to set up our filtration system. I used a ring stand, right angel clamp, three finger clamp, Buchner funner, filtering flask,rubber tubing, and filter paper in the Buchner funnel. I turned on the aspirator and pored some water over the filtering paper in order to create a good …show more content…
suction. I then pored the flask containing our forming aspirin into the funnel, and rinsed and of the remaining reaction out of the flask and onto the filter with ice-cold water. We then washed the product with 10 Ml of ice-cold water to remove the extra traces of acetic acid from the aspirin. My partner and I noticed that our filtrate was not crystal clear like everyone else's was, in fact it was quite cloudy. So we decided to remove the filtrate sheet with the aspirin crystals that we had already formed and placed it on a piece of paper towel while we placed a new piece of filter paper into the Buchner funnel and then filtered our solution again. this time our filtrate was a little less cloudy, but we didn’t have many crystals collect in the filtrate sheet. So we filtrated it again, using a new filtrate system. This time we got crystal clear filtrate, and a small amount of aspirin crystals. We then used the paper towels to pat dry the aspirin crystals we had collected. I weighed a very small beaker that weighed about 38 grams. Then I transferred the aspirin to the small beaker, making sure most of the extra aspirin was off of the paper towel and placed it in the drying over around 5:45 and left it in there unit about 5:55. We then weighed the beaker again with the aspirin in it this time, and ran into a problem, our beaker weighed more empty then it did with the aspirin in it. we used multiple different scales to make sure this was correct. then, I weighed a bigger beaker, and transferred the aspirin over to that and it ended up correcting the problem. At the very end of our experiment I took a small sample of aspirin and salicic acid in two different test tubes, I the put in 2 mL of FeCl3 in them and watched the reactions. I noticed that the aspirin turned dark orange while the salicylic acid turned dark purple. Results: Moles of salicylic acid used 3.0 g C7H6O3*1 Mol C7H6O3/138.118 g C7H6O3 = 0.022 Mol C7H6O3 Theoretical yield of aspirin in grams 3 g S.
acid*1 mol s. Acid/ 138.1g s acid*1 mol aspirin/1 mol s. acid * 180.2 g aspirin/1 mol aspirin = 3.9145 aspirin
mass of aspirin recovered
59.628g - 56.230 g= 3.398
percentage yield of aspirin:
3.398 g/3.9145 g *100= 86.8%
Conclusion:
Looking at my results, I can see that I only obtained 86.6% of aspirin. this is not the highest percentage that it could be. I probably lost some aspirin when I had to re-filtered the solution multiple times, or when I had to weigh a new beaker and maybe not all of it got transferred over to the new beaker, or maybe I didn’t get all of the aspirin crystals off of the paper towel and into the beaker. Regardless, of where the aspirin sample was lost, I obtained a 86.6% of aspirin from this
experiment.
The experiment was not a success, there was percent yield of 1,423%. With a percent yield that is relatively high at 1,423% did not conclude a successful experiment, because impurities added to the mass of the actual product. There were many errors in this lab due to the product being transferred on numerous occasions as well, as spillage and splattering of the solution. Overall, learning how to take one product and chemically create something else as well as how working with others effectively turned out to be a
The purpose of this experiment was to learn and preform an acid-base extraction technique to separate organic compounds successfully and obtaining amounts of each component in the mixture. In this experiment, the separation will be done by separatory funnel preforming on two liquids that are immiscible from two layers when added together. The individual components of Phensuprin (Acetylsalicylic acid, Acetanilide, and Sucrose as a filler) was separated based upon their solubility and reactivity, and the amount of each component in the mixture was obtained. Also, the purity of each component will be determined by the melting point of the component.
Another thing that must be kept constant is who much of the Alka-Seltzer tablet you use and what the surface area is. This has to be kept constant since maybe the crushed Alka-Seltzer tablet will dissolve faster because of its surface area, if you use different surface area your data will then again be
2-ethyl-1,3-hexanediol. The molecular weight of this compound is 146.2g/mol. It is converted into 2-ethyl-1-hydroxyhexan-3-one. This compounds molecular weight is 144.2g/mol. This gives a theoretical yield of .63 grams. My actual yield was .42 grams. Therefore, my percent yield was 67%. This was one of my highest yields yet. I felt that this was a good yield because part of this experiment is an equilibrium reaction. Hypochlorite must be used in excess to push the reaction to the right. Also, there were better ways to do this experiment where higher yields could have been produced. For example PCC could have been used. However, because of its toxic properties, its use is restricted. The purpose of this experiment was to determine which of the 3 compounds was formed from the starting material. The third compound was the oxidation of both alcohols. This could not have been my product because of the results of my IR. I had a broad large absorption is the range of 3200 to 3500 wavenumbers. This indicates the presence of an alcohol. If my compound had been fully oxidized then there would be no such alcohol present. Also, because of my IR, I know that my compound was one of the other 2 compounds because of the strong sharp absorption at 1705 wavenumbers. This indicates the presence of a carbonyl. Also, my 2,4-DNP test was positive. Therefore I had to prove which of the two compounds my final product was. The first was the oxidation of the primary alcohol, forming an aldehyde and a secondary alcohol. This could not have been my product because the Tollen’s test. My test was negative indicating no such aldehyde. Also, the textbook states that aldehydes show 2 characteristic absorption’s in the range of 2720-2820 wavenumbers. No such absorption’s were present in my sample. Therefore my final product was the oxidation of the secondary alcohol. My final product had a primary alcohol and a secondary ketone
The solvent should be easily removed from the purified product, not react with the target substances, and should only dissolve the target substance near it’s boiling point, but none at freezing. A successful recrystallization uses minimum amount of solvent, and cools the solution slowly, if done to fast, many impurities will be left in the crystals. Using the correct solvent, in this case ice water and ethyl acetate, the impurities in the compound can be dissolved to obtain just the pure compound. A mixed solvent was used to control the solubility of the product. The product is soluble in ethanol an insoluble in water. Adding water reduced solubility and saturates the solution and then the crystals
neutralize 35ml of our base. Once we weighed out the KHP we then dissolved it
Aspirin contains the substance acetylsalicylic acid (ASA), which can relieve inflammation, fever, pain, and known as a “blood thinner”. Aspirin was not officially trademarked until March 6, 1899 when the Imperial Office of Berlin made it official. It has been used for the last 110 years, but its natural form, salicylic acid has been around for thousands by Egyptians, Greeks, and Romans. Aspirin is available in over 80 countries and known as the best non-prescription drug. The most common use of aspirin is to cure headaches and use it as a pain reliever, but aspirin is known to prevent heart attack and strokes. It was first proposed in 1940, but wasn’t confirmed until 1970 when doctors would recommend taking aspirin daily [1].
In light of the findings of the study, the pharmacokinetic parameters of this drug would v...
(10) We learned how to extract and partially “clean” DNA from fruit flies, Inversion polymorphism associated with the ebony phenotype.
Analysis of Aspirin Tablets Aim --- To discover the percentage of acetylsalicylic acid in a sample of aspirin tablets. ----------------------------------------------------------------- In order to do this, the amount of moles that react with the sodium hydroxide must be known. This is achieved by using the method of back titration.
The experiment will demonstrate how antacids act inside the body by representing gastric acids with lemon juice. Gastric acid ranges from one to three on the pH scale (UCSB Science Line), while lemon juice ranges from two to three. (Helmenstine) Although the two acids have different pH levels, the rate of neutralization should remain approximately the same because the Antacid’s Neutralization Capacity (ANC) would remain the same no matter the acid type. (Ogbru pg 3,
In this experiment, [Co(NH3)5ONO]Cl¬2 was synthesized with a yield of 1.4314 g. It was then used to obtain UV-Vis Spectroscopy data with other prepared cobalt complexes including [Co(NH3)5(H2O)]Cl3, [Co(NH3)5(Cl)]Cl2 , Co(NH3)5(NO2)]Cl2 and [Co(NH3)6]Cl3. Each compound was a different color. Color, by definition, represents the wavelengths of UV light that a particle reflects. UV-Vis spectroscopy measures the amount of UV light absorbed. The easy way to determine wavelength of absorption from the color of the solution was the use of a color wheel like in Figure 1. The wavelengths of the color opposite of the solution’s color in the color wheel were the expected wavelengths of absorption. Co(NH3)5ONO]Cl¬2 was an reddish-orange color so its wavelength
·Aspirin (salicylic acid acetate) is an anti-inflammatory (decreases swelling and inflammation), anti-pyretic (fever reducing), and anti-platelet (decreases platelets in the body to thin blood). Many heart treatment patients take an aspirin a day to prevent blood clotting. However, if aspirin is taken in large quantities over long periods of time, it may cause gastric ulcers or other internal damage. The molecular formula of aspirin is C9H8O4. Some examples of Aspirin are Bayer, Ecotrin, and Aspergum.
Drop 1 aspirin tablet into the water, measure the time required for the tablet to fully dissolve.
The process of crystallization, especially crystallization from solution, is used as a separation and purification technique and in the production of pharmaceutical solid forms3,4. Knowledge of the crystal structure allows the crystal engineer to know and manipulate the chemistry of the crystal in order to optimize exact characteristics performance 3.