lab report - organic chem

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Question Description

  • do not plagiarize
  • I have started the question that required calculations. USE THE DATA PROVIDED TO DISCUSS RESULTS.
  • Feel free to ask questions
  • Make sure all questions are fully answered
  • If you use any outside source, please list it below*
  • I have started question #1 (not sure if its right)
  • Read the lab before answering all questions (additional information for answer question can be found there)
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o do not plagiarize o I have started the question that required calculations. USE THE DATA PROVIDED TO DISCUSS RESULTS. o Feel free to ask questions o Make sure all questions are fully answered o If you use any outside source, please list it below* o I have started question #1 (not sure if its right) o Read the lab before answering all questions 1. What is the unknown painkiller in the pill? Explain your evidence for its identity, and how you were able to rule out the other possible unknown. (5 points) - DMS melting point report for phenacetin: 134˚ - DMS melting point for acetanilide: 114.3˚ In the lab we learned that both phenacetin and acetanilide are compounds reported with melting points. The reported melting point for phenacetin is 134 degree Celsius, whereas the melting point for acetanilide is 114.3 degree Celsius. Based on the reported melting point of these compounds, it was able to compare it with the unknown’s melting point to determine which of the two compounds had been substituted for acetaminophen. According to the melting points obtained in the lab, the unknown painkiller in the pill is phenacetin because the melting point of the recrystallized sample was 136˚ and for the crude, it was 135.8˚ which is close to the reported melting point for phenacetin. 2a. Based on the sum of your weights of your separated compounds, calculate the total percent recovery of the painkiller. Use the crude weight of the unknown for this calculation. Discuss your results – how well were you able to collect each component during the separation steps? Initial mass of pain killer = 1.501g o o o o Sucrose mass = 0.146 g Aspirin mass = 0.3977 g Unknown (crude) mass = 0.54 g Total mass collected = 1.0837 g (sucrose mass + aspirin mass + unknown mass)/initial mass of painkiller x 100% (0.146 +0.3977 + 0.54)/ 1.501 x 100% = 72.19% Total percent recovery = 72.19% The separation of the unknown painkiller didn’t go as expected. The 72.19% recovery is a low amount of solid recovered. It needs to be taken in consideration that after the filtration, and the solution exposed to heat 2b. Calculate the percent composition of your painkiller. Again, use the crude weight of the unknown for this calculation. Does your percent composition match the label? You may assume a range of 8-12% sucrose, 35-45% aspirin, and 45-55% unknown is a match to the label. Discuss your results – explain any experimental errors that could have led to deviations of your values from the label values. o Total mass collected = 1.0837 g Each component mass/total mass collected x 100% Sucrose 0.146/1.0837 x 100% = 13.5% Aspirin 0.3977/1.0837 x 100% = 36.7% Unknown (crude) 0.54/1.0837 x 100% = 49.8% 13.5 + 36.7 + 49.8 = 100% 3a. Calculate a percent recovery for the recrystallization of the unknown component of the painkiller. Discuss your results – what experimental factors would cause a low percent recovery of your sample? (5 points) Recrystallized/crude x 100% 0.3855/0.54 x 100% = 71.38% 3b. Based on your crude and recrystallized melting points, was your recrystallization a success? Did you see any improvement to your melting point after recrystallization? Crude Melting point : 135.8˚ Recrystallized melting point : 136˚ 4. While acetanilide and phenacetin are not very acidic, acetaminophen (like aspirin) has an acidic hydroxyl group. Based on this information, what problem would you encounter during the separation steps if the unknown component were acetaminophen rather than acetanilide or phenacetin? 5. Equation 1 shows the balanced reaction equation for the deprotonation of aspirin by sodium hydroxide. What is the balanced reaction equation when sodium acetylsalicylate is treated with concentrated hydrochloric acid? 6. Assuming the reaction from question #5 is spontaneous under standard conditions, label the stronger acid, stronger base, weaker acid, and weaker base. INFORMATION YOU WILL NEED TO ANSWER QUESTIONS CAN BE FOUND HERE* READ CAREFULLY Experiment - SEPARATION AND IDENTIFICATION OF THE COMPONENTS OF A PAINKILLER Analgesic drugs reduce pain and antipyretic drugs reduce fever; some drugs, such as aspirin, do both. Most of the common over the counter analgesic/antipyretic drug preparations contain aspirin, acetaminophen, or combinations of these substances with other ingredients. You can see from the chemical structures of acetaminophen, phenacetin, and acetanilide (Figure 1) that they are structurally related. Figure 1. Chemical structures of analgesic drugs As is often the case for scientific discoveries, the pain-killing and fever reducing effects of acetanilide were discovered by chance. In 1886, Arnold Cahn and Paul Hepp were trying to treat patients with intestinal worms. They tried everything in their stock room. One patient was treated with a substance labeled as naphthalene. The chemical did not affect the worms, but it did effectively treat the patient’s fever and pain. Since the substance did not have the traditional moth ball smell, it was sent to Hepp’s cousin for analysis. Tests indicated that the new drug was acetanilide. Acetanilide was soon marketed under the name antifebrin. It is a pain reliever and fever reducer, but it has a nasty side effect. Some patients who took antifebrin developed methemoglobinemia, a condition in which hemoglobin molecules are altered in a way that reduces their ability to transport oxygen. Acetanilide is considered too toxic for medicinal use today, but its discovery led to the development of safer and more effective analgesic and antipyretic drugs. Just a few months after the discovery of antifebrin, Carl Duisberg had to get rid of 50 tons of paminophenol, a by-product of dye manufacturing. Duisberg knew that many compounds with a hydroxyl (-OH) group attached directly to a benzene ring were toxic, so he decided to hide the hydroxyl group by converting the alcohol to an ether. Converting the primary amine (-NH2) to the amide generated phenacetin (Figure 2). Phenacetin was used for years as an effective analgesic/antipyretic drug. Only recently was it banned in the United States since it can cause kidney damage, hemolytic anemia, or even cancer in some patients. If you look carefully at the synthesis of phenacetin, you might notice that if Duisberg had not masked the hydroxyl group, he would have synthesized acetaminophen, which has proved to be safer than either acetanilide or phenacetin as an analgesic. Organic chemistry lab techniques can be used to solve a variety of everyday “real” problems. You will be asked to solve such a problem in this experiment. We have discovered an over-thecounter painkiller that looks as if it might be a counterfeit. The label says the product is manufactured in the United States, but a savvy consumer noticed minor discrepancies in the labeling and noticed that the pills in the bottle seemed to crumble more easily than the previously purchased bottle of the same painkiller. These discrepancies caused her to contact us. We have been charged with determining if the pills in the bottle are what the bottle claims them to be. The label indicates that each tablet contains 200 mg of aspirin, 250 mg of acetaminophen, and 50 mg of sucrose. The sucrose is simply added to make the pills taste better. The other components are painkillers. Initial analysis does indicate that the tablets contain sucrose and aspirin, but the third component is still a mystery. The third component of the mixture is probably very similar in structure to acetaminophen, being either acetanilide or phenacetin. Both acetanilide and phenacetin are effective painkillers, so a substitution would be unnoticeable to the consumer, but they are both banned in the United States because of their toxicity. If either of the banned substances is present in the over-the-counter product, then we will need to have the product removed from the shelves and turn the case over to the FDA. It is your job to find out what percentages of aspirin, sucrose, and the unknown painkiller component are present in the tablets, and to determine the true identity of the third component of the painkiller. You will be separating the three components of the painkiller by making use of their solubilities and acid-base properties. The painkiller contains sucrose, aspirin, and an unknown third component. You will find these solubility characteristics to be of utmost importance: 1. Sucrose is soluble in water but insoluble in the organic solvent dichloromethane (also called methylene chloride, with chemical formula CH2Cl2). 2. Aspirin is soluble in dichloromethane but relatively insoluble in water. Sodium hydroxide, a strong base, converts aspirin to a salt that is insoluble in dichloromethane but soluble in water. 3. Acetanilide and phenacetin, like aspirin, are soluble in dichloromethane and insoluble in water. They are not converted to salts by sodium hydroxide. Dichloromethane will dissolve aspirin and the unknown component, but not sucrose. The sucrose can then be removed from the solution of the other two components by gravity filtration To separate the aspirin from the unknown component, you will take advantage of the acid/base properties of aspirin and the fact that the solubility of aspirin changes when it is deprotonated (Equation 1). Equation 1. Formation of sodium acetylsalicylate from aspirin A dichloromethane solution of aspirin and the unknown component can be placed in a separatory funnel and extracted with an aqueous solution of sodium hydroxide. The salt of aspirin will move to the water layer. The unknown component will stay in the dichloromethane. The layers can then be separated from one another. A separatory funnel is shown in Figure 3. The funnel is fitted with a stopcock and a glass stopper. The separatory funnel is top heavy, especially when filled with liquid, and should be supported in an iron ring of proper size attached to a ring stand. Before adding any liquid to the separatory funnel, make sure that the stopcock is closed. To perform a separation, add the two immiscible liquids for the extraction into the separatory funnel, never filling the separatory funnel more than about 3⁄4 full to allow room for mixing. Place the stopper in the separatory funnel and, holding the stopper firmly in place with the index finger of one hand, remove the separatory funnel from the iron ring (Figure 4). Invert the funnel, pointing the stem up away from you and away from everybody else, and carefully open the stopcock to vent any gases. You may hear a hissing sound as the gases are released through the stopcock. Close the stopcock and gently shake the two liquids together to mix them for several minutes. Since the reaction between aspirin (in the dichloromethane layer) and the sodium hydroxide (in the aqueous layer) only occurs where the two liquids meet, the mixing step is key to the extraction process, and should not be rushed. It is common for pressure to build up in the separatory funnel during mixing, especially when using a volatile organic solvent, so it is necessary periodically to stop mixing and open the stopcock to vent any pressure that has built up. After mixing is complete, close the stopcock and place the separatory funnel in the iron ring, remove the stopper and allow the layers to settle and separate; a clean interface should form between the two layers. Note: if bubbles have formed at the interface, these can be broken up by using a glass rod or spatula to disturb the interface. Place an Erlenmeyer flask or a beaker under the funnel, slowly open the stopcock and drain the lower layer into the beaker or flask. As the interface between the two solvents approaches the stopcock, slow the rate of draining by adjusting the stopcock. If you are going to perform another extraction on the top layer, leave it in the separatory funnel and add more of the next extracting solvent to the separatory funnel. If, however, you are going to perform another extraction on the lower layer, pour the remaining upper layer out the top of the separatory funnel into a separate flask and return the lower layer to the separatory funnel for the next extraction. A common problem that you may face in doing extractions is trying to determine which one of the two layers in the separatory funnel is the aqueous layer and which one is the organic layer. The heavier layer (i.e., the more dense liquid) is the lower layer, of course, but some organic liquids (e.g., benzene, diethyl ether, ethyl acetate) are lighter than water and some (e.g., chloroform, dichloromethane) are heavier than water, so, depending on the solvents being used, the organic phase might be the upper or the lower layer in the separatory funnel. If you are unsure about which layer is which, carry out a simple test: mix a small sample of each layer with a few drops of water in a test tube and see if the two liquids are miscible; the aqueous solution will dissolve the water drops but the organic solution will be immiscible with the added water. The GOLDEN RULE of extraction is to never throw away a layer until the end of the experiment or until you are absolutely certain that you no longer need it. The most common mistake made during extraction is throwing the wrong layer down the drain or in the waste bottle, which means loss of material and starting experiment over from the beginning. Once the aspirin and the unknown component have been separated from each other using the separatory funnel, they must be recovered. One of the simplest ways to recover the aspirin is to add acid, thereby converting the sodium acetylsalicylate back to aspirin and changing the solubility so that the aspirin precipitates from the aqueous solution. Cooling the aqueous layer in ice will help with precipitation. The aspirin can be isolated by vacuum filtration. The unknown component can be isolated by drying the organic solution, removing the solution from the drying agent, and evaporating the solvent. What remains in the flask is the unknown component of the painkiller. To determine the identity of the unknown painkiller, you will find its melting point. The melting point of a pure solid organic compound is one of its characteristic physical properties along with molecular weight, boiling point, refractive index, and density. A pure solid will melt reproducibly over a narrow range of temperatures, typically 1-2 °. Note that a melting “point” is actually a misnomer, and melting points of compounds are usually recorded as a range. Melting points are determined for three reasons. If the compound is known one, the melting point will help to characterize the sample in hand. If the compound is new, then the melting point is recorded in order to allow future characterization by others. Finally, the range of melting point is indicative of the purity of the compound; an impure compound will melt over a wide range of temperatures with lower than recorded melting range. The process of determining this melting point range is done on a truly micro scale using less than 1 mg of the material. The apparatus is simple, consisting of a thermometer, a capillary tube to hold the sample, and a heating bath. The DigiMelt and Mel-Temp apparatuses (Figure 5) consist of an electrically heated aluminum blocks that accommodates three capillaries. The sample is illuminated through a viewing lens. As the sample is heated to its melting point, two temperatures are recorded. The first temperature in the range is recorded when the sample just begins to barely melt. The second temperature in the range is recorded when the entire sample has become liquid. The accuracy of the melting point depends on a on the accuracy of the thermometer and on a rate of heating. At the melting point the temperature rise should not be greater than 1 °C per minute. This may seem extraordinary slow, but it is necessary in order for heat from the bath to be transferred equally to the sample and to the thermometer, and to obtain an accurate reading. Both acetanilide and phenacetin are known compounds with reported melting points. By analyzing your unknown’s melting point, you can determine which of the two compounds has been substituted for acetaminophen in the counterfeit tablets. Procedure This lab takes place over two weeks. In the first week, you will separate the three components of the painkiller. In the second week of the experiment, you will recrystallize your unknown compound and determine the identity of your unknown compound by melting point comparisons. Separation of Sucrose Accurately weigh approximately 1.5 g of the painkiller and transfer it to a clean, dry 100 mL beaker, then add 35 mL of dichloromethane. Stir the mixture thoroughly to dissolve as much of the solid as possible, using a stirring rod to break up any lumps or granules. Flute and preweigh a piece of filter paper, then filter the mixture into a 125 mL Erlenmeyer flask by gravity to collect the sucrose. When not actively pouring the mixture, cover your funnel with a small watch glass to deter solvent evaporation. Set the filter paper aside in a small beaker and allow it to dry. At the end of the lab period, determine the weight of the sucrose. If some sucrose remains in the original flask, allow it to dry, then transfer it into the filter paper before obtaining the weight at the end of the lab period. Note: If a significant amount of crystals collects at the top of the filter paper as you are performing the gravity filtration, you will need to wash the crystals with small portions of dichloromethane to get them to dissolve again. Some dichloromethane evaporates from the top of the filter paper as the gravity filtration is taking place. As the solvent evaporates, it leaves the aspirin and unknown component at the top of the paper. Placing the watch glass over the filter paper immediately after pouring your solution can inhibit evaporation and crystal formation. If you leave significant crystals behind on the filter paper, your sucrose weight will be too high, the other two component weights will be too low, and you will not effectively separate the components of your mixture. Separation of Aspirin Transfer the filtrate to a separatory funnel and extract the aspirin with a 15 mL portion of aqueous 1 M NaOH by mixing the layers thoroughly. The density of CH2Cl2 is 1.34 g/mL. Stop and think, which layer will be on top? Collect each layer separately, then return the dichloromethane layer to the separatory funnel and extract it again with a fresh 15 mL portion of aqueous NaOH. Collect the aqueous layer with the first one. Save the dichloromethane solution for the next part of the experiment. To the combined aqueous extracts, add up to 5 mL of 6 M HCl slowly, with stirring. Once you see significant aspirin precipitate that does not redissolve upon stirring, test the pH of the solution with pH paper to confirm a pH of 2 or lower. Do not dip the pH paper directly into your solution; use a glass rod to transfer a drop of the solution to the pH paper. If your pH is above 2, add more acid dropwise, and retest the pH. Cool the mixture in an ice bath, collect the aspirin by vacuum filtration and use a minimal amount of ice-cold water to aid in the transfer. Allow the aspirin to dry while pulling vacuum. Once it is completely dry, determine the weight of the aspirin. Isolation of the Unknown Component Dry the dichloromethane solution with anhydrous sodium sulfate (Na2SO4), then decant the solution into a dry, preweighed 125 mL Erlenmeyer flask. Evaporate the solvent from the dichloromethane solution with minimal heat using a hot plate in your hood (remember to add a boiling stick!). If the unknown is quite impure, it may remain a liquid after all the solvent is removed, but it should solidify upon cooling. This is your crude painkiller (acetanilide or phenacetin), which you will save for the following week. Purification of the unknown component Weigh your flas ...
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Thomas574
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