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Steps 3 - 5 were repeated twice more for two more titrations.
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.
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.
Step 3
A burette was prepared by rinsing it twice with the NaOH solution. It was then filled close to the zero mark.
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
4
8
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.
4
8
4
8
204.23 g KHC H O
1
4
8
22.50 mL = .02250 L
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.)
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.
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
4
8
Mass of KHP (g)
Trial (KHP titration)
1
204.23 g KHC H O
Volume Used (mL)
Final Burette Reading (mL)
Initial Burette Reading (mL)
4
8
24.10 mL = .02410 L
2. moles KHP = moles NaOH
0.00232 mole KHP = 0.00232 mole NaOH
1
.436
22.50
0.00
3. Molarity NaOH = moles NaOH = 0.00232 mole NaOH = 0.0963 M NaOH
.02410 L soln.
L solution
2
.473
0.00
24.10
Trial 3:
1. mass KHP moles KHP
.527 g KHC H O x 1 mol KHC H O = 0.00258 mol KHP
3
.527
26.80
0.00
4
8
4
204.23 g KHC H O
1
4
8
4
26.80 mL = .02680 L
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.
Average:
.0947M + .0963M + .0963M
3
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.
Step 1
A 70 degree Celsius water bath was prepared by filling a 400 mL beaker with about 200 mL of water. The beaker was placed on the hot plate.
Step 2
2.0g of salicylic acid was placed in a 125-mL Erlenmeyer Flask.
Step 3
4.0 mL of acetic anhydride and 7 drops of 85% phosphoric acid were added to the flask.
Step 4
The flask was swirled thoroughly to mix all reagents.
Data Table:
Test Tube #
Color
(Distilled water, salicylic acid, and 6 drops of 1% ferric chloride solution)
1
Dark Purple
(Distilled water, our product, and 6 drops of 1% ferric chloride solution)
2
Very Light Purple
3
Clear and Colorless
(Distilled water and 6 drops of 1% ferric chloride solution)
Distilled Water
Buchner Funnel Assembly
Goggles
Beakers
Balanced Equation:
Wire Gauze
Ring stand
Step 5
The flask was placed into the 70°C water bath and heated for 10 minutes.
125 mL Erlenmeyer Flask
Test Tubes and Test Tube Rack
Rubber Tubing
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.
Hot Plate
Calculator...I mean Thermometer
Clamp
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.
The aspirator was left running for about 10 minutes to remove as much water as possible.
Step 11
Step 12
The product was allowed to air dry over night, and weighed during the next lab period.
Step 1
Three test tubes were filled with 4 mL of distilled water and placed in a test tube rack.
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 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.
The first objective of this lab was to synthesize aspirin and perform a purity test on the product.
"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.
2013.
"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.
2013.
"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.
<http://www.york.ac.uk/>.
"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.
Ibuprofen is sold under the brand name Advil and Acetaminophen is sold under the brand name Tylenol.
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.
Step 3
A total of three titrations were done. The volume of the acid and base used for each trial were calculated.
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.
Step 4
The molarity of the acid solution for each trial was calculated. The average concentration of the HCl was determined.
HCl + NaOH H O +NaCl
2
(aq)
(l)
1. M = M V = (0.0958 M NaOH)(.01982 L) = 0.0940 M
A
B
V
(0.02021 L)
A
2. M = M V = (0.0958 M NaOH)(.01931 L) = 0.0925 M
B
A
V
(0.02000 L)
A
3. M = M V = (0.0958 M NaOH)(.01955 L) = 0.0932 M
B
A
V
(.02010 L)
A
Average of M
A
0.0940 M + 0.0925 M + 0.0932 M = 0.2797 M
0.2797 / 3 = 0.0932 M
Trial 2
Trial 1
Trial 3
Acid
Base
Acid
M = [HCl] =
A
Initial burette reading (mL)
Final burette reading (mL)
A
B
Volume used (mL)
A
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.
The third objective of the lab was to determine the concentration of the HCl.
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.
Conc. of NaOH: .0958M
Conc. of HCL: .0932M
Sample Weight: .516g
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
Final buret reading: 15.4mL
Initial buret reading: 0mL
Volume of NaOH: 45.50mL
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
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%.
0.00292 moles ASA x 176.13 g ASA = 0.514 g ASA
1 1 mole ASA
7. our sample mass
0.516 grams
8. grams ASA
grams sample
x 100
x 100 = 99.6 %
0.514 g ASA
0.516 g ASA sample