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Aspirin Synthesis Lab Report
Transcript of Aspirin Synthesis Lab Report
4.0 mL of acetic anhydride and 7 drops of 85% phosphoric acid were added to the flask. The flask was placed into the 70°C water bath and heated for 10 minutes. Step 6
The flask was removed from the water bath and 2 mL of ice-cold distilled water was added in a dropwise manner. The flask was swirled after the addition of each drop. Step 7 3 ice cubes were added to the Erlenmeyer flask and the flask was placed in an ice bath. The flask was allowed to cool for approximately 10 minutes. Step 8
Filter paper was placed into the center of the Buchner funnel so that all of the holes were covered. The paper was moistened with distilled water. Then, a vacuum filtration apparatus was set up. Step 9 The product was collected in the Buchner funnel via vacuum filtration. Step 10
The crystals were washed four times with approximately 25 mL of ice-cold distilled water. Step 11 The aspirator was left running for about 10 minutes to remove as much water as possible. Data Aspirin was first synthesized by a German chemist named Felix Hoffman in 1897. Hoffman developed the process for the Bayer Company. He also synthesized Aspirin because his father, who had severe arthritis, could not tolerate the salicylic acid he was taking for pain relief.
The Kolbe Synthesis for the production of salicylic acid was created by the German chemist Hermann Kolbe. In this process, sodium phenoxide is heated with carbon dioxide under pressure (100 atm) and the mixture is then acidified to yield salicylic acid.
An analgesic is a member of the group of drugs used to achieve analgesia, which is the relief from pain. They are also known as painkillers, and act in various ways on the peripheral and central nervous systems. Non-opioid analgesics including aspirin, ibuprofen, acetaminophen, naproxen and other anti-inflammatory drugs, act at the site of the injury. These analgesics prevent inflammation and block the release of chemicals that stimulate nociceptors (a receptor for pain caused by injury, physical, or chemical, to body tissues). Aspirin belongs to non-opioid analgesics. Ibuprofen Acetaminophen Sources: C H O C H NO Taking aspirin daily can significantly lower the risk of heart attacks. People with diabetes or risk factors like high blood pressure who are at a high risk for heart attacks can experience the benefits from taking aspirin daily. It also helps people with diabetes who have had heart attacks or strokes before, who also have heart disease.
Reye’s syndrome is the malfunctioning of cells in the liver and the brain, a very rare disorder. This causes drowsiness, nausea, and severe vomiting, usually seen in children ages 3-12. This occurs when someone has the chicken pox, the flu, or a cold (occurs in these viral illness settings). There is an association of using aspirin to treat colds, chicken pox, etc. and the development of Reye’s syndrome. It is recommended that aspirin should not be given to children or teenagers as pain relievers if they have the flu, chicken pox, or a cold. Ibuprofen is sold under the brand name Advil and Acetaminophen is sold under the brand name Tylenol. Data Table: Test Tube # Color 1 2 3 Dark Purple Very Light Purple Clear and Colorless Balanced Equation: Part 2: NaOH Preparation and Standardization of Sodium Hydroxide Solution Objective 2 The second objective of this lab was to prepare two liters of approximately .1 M NaOH solution to be used in following experiments. Another objective of part 2 was to measure the exact molarity of the NaOH solution by doing an acid-base titration. Materials Phenolphthalein indicator
reagent bottle for NaOH
25 mL pipette
100 mL beaker
125 mL Erlenmeyer flask Calculations: 1. volume x concentration = total moles NaOH for hydrolysis
.0455 L x .0958 mole = 0.00436 moles NaOH 1 L 2. volume x concentration = moles HCl
.0154 L x .0932 mole = 0.00144 moles HCl 1 L 3. moles HCl = moles NaOH (excess)
0.00144 moles HCl = 0.00144 moles NaOH
4. total moles NaOH - excess moles NaOH = moles NaOH used
0.00436 moles NaOH - 0.00144 moles NaOH excess = 0.00292 moles NaOH
5. moles NaOH used = moles ASA hydrolyzed
0.00292 moles NaOH used = 0.00292 moles ASA hydrolyzed
6. moles ASA grams ASA 0.00292 moles ASA x 176.13 g ASA = 0.514 g ASA 1 1 mole ASA 7. our sample mass
8. grams ASA
grams sample x 100 0.514 g ASA
0.516 g ASA sample x 100 = 99.6 % Step 12
The product was allowed to air dry over night, and weighed during the next lab period. (Distilled water and 6 drops of 1% ferric chloride solution) "Analgesic." Google. Google, n.d. Web. 13 Apr. 2013.
"Analgesic." Wikipedia. Wikimedia Foundation, n.d. Web. 13 Apr. 2013.
"Chemistry 104: Synthesis of Aspirin." Chemistry 104: Synthesis of Aspirin. N.p., n.d. Web. 13 Apr.
"DrugBank: Ibuprofen (DB01050)." DrugBank: Ibuprofen (DB01050). Genome Alberta & Genome
Canada, n.d. Web. 25 Mar. 2013.
Freudenrich, Craig C. "How Do Analgesics Work on Pain? - Curiosity." Curiosity. Discovery
Communications, n.d. Web. 25 Mar. 2013.
"Ibuprofen - PubChem." Ibuprofen - PubChem. PubChem, n.d. Web. 25 Mar. 2013.
"Living with Diabetes." Aspirin. American Diabetes Association, n.d. Web. 25 Mar. 2013.
"Nociceptor." The Free Dictionary. Farlex, Inc., n.d. Web. 13 Apr. 2013.
Owens, Lisa, MD. "What Is Reye's Syndrome?" ABC News. ABC News, 29 Dec. 2009. Web. 13 Apr.
"Tower Medical Center of Nederland." Tower Medical. N.p., n.d. Web. 13 Apr. 2013.
Tylenol (Acetaminophen) Drug Information: Description, User Reviews, Drug Side Effects,
Interactions - Prescribing Information at RxList." RxList. RxList Inc., n.d. Web. 13 Apr. 2013.
"The University of York." The University of York. Universtiy of York, n.d. Web. 13 Apr. 2013.
"What Is the Role of Aspiring in Triggering Reye's?" National Reye's Syndrome Foundation. National
Reye's Syndrome Foundation, n.d. Web. 25 Mar. 2013. (Distilled water, our product, and 6 drops of 1% ferric chloride solution) (Distilled water, salicylic acid, and 6 drops of 1% ferric chloride solution) Explanation of Calculations: 1. We multiplied the volume of of NaOH added for hydrolysis by the concetration of the NaOH in order to get the total moles of NaOH needed for hydrolysis.
2. We multiplied the volume of HCl used in back titration by the concetration of the HCl in order to get the moles of HCl used in the back titration.
3. We concluded that the moles of HCl, which was 0.00144 moles, equals moles of excess NaOH.
4. We subtracted the excess moles of NaOH from the total moles of NaOH in order to get the moles NaOH used.
5. We concluded that the moles of NaOH used, which was 0.00202 moles, equals the moles of ASA hydrolyzed.
6. We converted the moles of ASA to grams of ASA by multiplying the moles by the molar mass of ASA over one mole of ASA, and got 0.514g.
7. Our sample mass was 0.516 grams.
8. We calculated the purity of the sample by dividing the grams of the ASA sample by our sample of ASA, and multiplied it by 100. Our purity percentage was 99.6%. Chemical Reagents: 2.0 g of salicylic acid (C H O )
4.0 mL of acetic anhydride (C H O )
7 drops of 85% phosphoric acid (H PO )
100 ml of ice-cold distilled water (H O)
3 ice cubes (H O)
18 drops of 1% ferric chloride solution (FeCl ) Chemical Reagents: Potassium Hydrogen Phthalate (KHP) (KHC H O )
Sodium Hydroxide (NaOH)
distilled water (H O) Data Table: Trial (KHP titration) 1 2 3 Mass of KHP (g) Initial Burette Reading (mL) Final Burette Reading (mL) Volume Used (mL) .436 .473 .527 0.00 0.00 0.00 22.50 24.10 26.80 22.50 24.10 26.80 Balanced Equation: Procedure Step 2
A solution of approximately 0.1 Molar NaOH was prepared by weighing out 4 grams of NaOH into a 1 liter plastic bottle. The bottle was filled with distilled water to the seam near the top of the bottle. The bottle was shaken until all the solid NaOH dissolved, and then was shaken for an additional 2 minutes. Step 1
The potassium hydrogen phthalate (KHP) was stored in the oven to keep it dry. It was then removed from the oven and placed in a desiccator to cool. (This was done prior to the start of part 2 of this lab.) Step 3
A burette was prepared by rinsing it twice with the NaOH solution. It was then filled close to the zero mark. Step 5
The starting volume on the burette was recorded. The KHP solution was titrated to the first appearance of faint pink color that persists for at least 20 seconds. The final volume on the burette was recorded. The flask was rinsed with distilled water and dried as much as possible. Step 4
Three samples of KHP were weighed out ranging from .4g to .6g. Each sample was transferred into a clean and dry 125mL Erlenmeyer flask before each titration. Exactly 50mL of distilled water was pipetted into the flask, which was then swirled to dissolve the KHP. Then 2 drops of phenolphthalein indicator were added to the flask. The color we don't want! Calculations: Trial 1:
1. mass KHP moles KHP (KHC H O )
.436 g KHC H O x 1 mol KHC H O = 0.00213 mole KHP 1 204.23 g KHC H O Steps 3 - 5 were repeated twice more for two more titrations. 2. moles KHP = moles NaOH
0.00213 mole KHP = 0.00213 mole NaOH 3. Molarity NaOH = moles NaOH = 0.00213 mole NaOH = 0.0947 M NaOH L solution .02250 L soln. 22.50 mL = .02250 L 1. We converted mass of KHP to moles of KHP by multiplying by the molar mass of KHP.
2. We concluded that the moles of KHP were equal to the moles of NaOH.
3. We calculated the molarity of NaOH by dividing the moles of NaOH by the liters of solution, which was the volume used in the titration.
Each of the calculations were the same for each trial.
The average was calculated by adding the 3 molarities of NaOH from each trial, and dividing their sum by 3. Trial 2:
1. mass KHP moles KHP
.473 g KHC H O x 1 mol KHC H O = 0.00232 mol KHP 1 Explanation of Calculations: 204.23 g KHC H O 2. moles KHP = moles NaOH
0.00232 mole KHP = 0.00232 mole NaOH 3. Molarity NaOH = moles NaOH = 0.00232 mole NaOH = 0.0963 M NaOH L solution .02410 L soln. Trial 3:
1. mass KHP moles KHP
.527 g KHC H O x 1 mol KHC H O = 0.00258 mol KHP 24.10 mL = .02410 L 1 204.23 g KHC H O 2. moles KHP = moles NaOH
0.00258 mole KHP = 0.00258 mole NaOH 3. Molarity NaOH = moles NaOH = 0.00258 mole NaOH = 0.0963 M NaOH L solution .02680 L soln. 26.80 mL = .02680 L Part 3: Preparation and Standardization of HCl Objective 3 The third objective of the lab was to determine the concentration of the HCl. Materials reagent bottle for HCl
acid and base burettes
100 mL beaker
125 mL Erlenmeyer flask
bromthymol blue indicator
base beaker Chemical Reagents 3 M HCl
distilled water (H O)
0.0958 M NaOH Average:
.0947M + .0963M + .0963M Step 1
In order to prepare 500 mL of approximately 0.1 Molar HCl, 17 mL of the sock 3 M HCl was diluted to 500 mL in the provided reagent bottle. The bottle was half-filled with distilled water before the 3 M HCl was added, and then the bottle was topped off with distilled water to the 500 mL mark. 3 = .0958M Step 2
The acid and base burettes were prepared by washing with the respective solutions twice. Each burette was filled with the respective solution and the initial reading was recorded. Approximately 20 mL of acid was drained from the acid burette into the 100 mL beaker. 3 drops of bromthymol blue indicator was added to the flask. This was titrated with base to a faint green end point. The final burette readings were recorded. Data Step 3
A total of three titrations were done. The volume of the acid and base used for each trial were calculated. Step 4
The molarity of the acid solution for each trial was calculated. The average concentration of the HCl was determined. Calculations: 1. M = M V = (0.0958 M NaOH)(.01982 L) = 0.0940 M A A B B V Trial 1 Trial 2 (0.02021 L) Trial 3 Acid 2. M = M V = (0.0958 M NaOH)(.01931 L) = 0.0925 M V Acid Base A B B Base A Acid Base (0.02000 L) Initial burette reading (mL) 3. M = M V = (0.0958 M NaOH)(.01955 L) = 0.0932 M A A B B V (.02010 L) Average of M A 0.0940 M + 0.0925 M + 0.0932 M = 0.2797 M
0.2797 / 3 = 0.0932 M
HCl + NaOH H O +NaCl Balanced Equation M = [HCl] = A Final burette reading (mL) Volume used (mL) 20.21 8 19.82 4 20.00 0.00 20.21 19.82 0.00 19.31 0.00 38.86 0.00 19.55 20.00 20.10 19.31 40.10 20.00 19.31 8 8 8 4 4 4 4 4 4 4 8 8 8 8 8 8 4 4 4 4 4 4 4 4 4 4 4 4 Part 4: Quantitative Analysis of Aspirin Objective 13 18 2 8 9 The objective of this lab was to determine the purity of the synthesized aspirin by determining the percentage of acetylsalicylic acid present in the product. 2 Balanced Equation Step 1
Three test tubes were filled with 4 mL of distilled water and placed in a test tube rack. Ferric Chloride Test for Purity Chemical Reagents: .1M HCl, standardized
.1M NaOH, standardized
Phenolphthalein solution Step 1 Approximately .5g of the synthesized aspirin was weighed and placed into a 250mL Erlenmeyer flask. Step 2
A small amount of salicylic acid was dissolved in the first test tube. The second test tubed was used to dissolve a portion of the product. The third test tube, which contained only distilled water, served as the control. Step 2 Step 3
Six drops of 1% ferric chloride solution were added to all 3 test tubes. The colors of each test tube were recorded in a data table. Two burettes were prepared, one with HCl and one with NaOH. Initial burette readings were recorded. Step 4 25mL of 95% ethanol was cooled to 15 degrees Celcius and ethyl alcohol was added to the flask to dissolve. Step 5 Materials Synthesized aspirin
250mL Erlenmeyer flask
Balance Data Conc. of NaOH: .0958M
Conc. of HCL: .0932M
Sample Weight: .516g Volume of NaOH required to neutralize all acids present: Final buret reading: 45.50mL
Initial buret reading: 0mL
Volume of NaOH: 30.50mL
Approximate volume of NaOH to be added for hydrolysis: 45.50mL Actual volume of NaOH added for hydrolysis: Final buret reading: 15.4mL
Initial buret reading: 0mL
Volume of NaOH: 45.50mL 7 6 3 4 6 3 3 4 2 2 3 #1 #2 #3 8 4 4 2 2 Procedure 2 Explanation of Calculations (l) (aq) (aq) (aq) 0.0932 M For all three trials, we used the equation:
M V = M V
Then we rearranged the equation to solve for M by dividing by V . Next we plugged in our numbers and found the molarity of the acid.
To calculate the average concentration of the acid, we added up the three molarities and divided by 3. A A B B A A 2 drops of phenolphthalein were added and the sample was titrated to a faint pink endpoint. The volume of NaOH used was recorded Step 7 Step 6 The amount of base used for the hydrolysis was determined by adding 15 mL to the volume of base base required in the previous titration. Step 3 One burette was filled with standard NaOH solution and the second was filled with standard HCl solution. Then the initial burette readings were recorded. The burette was refilled with more standard base and the initial reading was recorded. To hydrolyze the aspirin, additional NaOH was added to the flask from the burette. Step 8 Step 10 The mixture was heated for 15 minutes in a bath of boiling water. If the solution was not pink, two more drops of phenolphthalein indicator were added. Step 9 The flask was cooled to room temperature with an ice bath. Step 11 The initial volume of HCl was recorded and the excess base was back titrated with the standard HCl standard until the pink color disappeared. The volume of the HCl used was recorded. Further Analysis THE END Presented by MY LIFE: YOU LIGHT IT UP Conclusion The overall objectives of this experiment were to synthesize aspirin, to perform a crude purity test on said aspirin, to standardize and prepare solutions of both NaOH and HCL, and to qualitatively determine the purity of the aspirin by finding the percentage of acetylsalicylic acid in the aspirin. All objectives were performed successfully. First, aspirin was synthesized from salicylic acid and acetic anhydride, yielding a white, crystalline product. This product was then crudely tested for purity using Ferric Chloride. The test tube containing the product only changed to a very slight purple color after this test, indicating a very low level of salicylic acid and a high purity. A solution of NaOH was then prepared by dissolving solid NaOH in water and then titrating a known quantity of KHP with it to find the molarity, which turned out to be .0958 M. A solution of HCl was then prepared by diluting stock 3 M HCl and titrating the solution with the previously made NaOH solution to find the molarity, which turned out to be .0932 M. Then, the aspirin product was dissolved in water and titrated with both solutions to find the percent purity of the aspirin, which was found to be 99.6% pure. This is an extremely high purity, which coincides with the quantitative analysis done in Part 1. Thus, it is doubtful that there are very many sources of error. However, there are possible sources of error in this experiment. In part 1, the beaker in which the product was heated had to be kept close to 70 degrees Celsius or else different reactions would start taking place. It is possible that we lapsed in our efforts to maintain this temperature, allowing other, unknown reactions to take place, contaminating our sample. Additionally, it is possible that the aspirin never fully dried after the experiment, making the recorded mass of the aspirin higher and making the purity lower than the actual purity in the end. This is unlikely, though, considering that it was left for weeks to dry. It is, however, still a possibility. Additionally, in Part 2, by the time the indicator we used changes to pink the solution is already slightly basic, and not neutral, making the molarity of the NaOH solution in reality lower than it was recorded to be, eventually translating into a skewed HCl solution and likely inflating the purity of the final result. There were likely no sources of error in part 3 except for the possible error carried forward from Part 2. In Part 4, again, the issue where phenolphthalein only turns pink after the solution turns basic makes an appearance in the initial neutralization of the aspirin solution. The solution was slightly basic already before an excess of base was added, ultimately making the purity of the aspirin slightly higher than our result. It is additionally possible that we over-titrated when back-titrating, as there is no color to indicate the difference between an acidic solution and a neutral one. This would slightly increase the final purity of the aspirin. The only major flaw in this experiment is the indicator used to titrate most of the solutions, as there is no color for a strictly neutral solution. As a result, if we were to improve this experiment, we would use an indicator such as bromthymol in all titrations that has a specific color for a neutral solution to increase the accuracy of the titrations.