How Pure is Aspirin - This report includes data collected from the lab. This report is followed by PDF

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This report includes data collected from the lab. This report is followed by the report "How Pure is Aspirin"...

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Chemistry Lab Report Experiment 6: How Pure is Aspirin

March 5, 2021 Spring 2021

Introduction A calorimetric analysis is a method for is used to determine the concentration of a solute in a solution when this is colored. This analysis is accompanied by Beer’s law, Beer’s law is the relation between the concentration that has the solute as well as the light absorbed. Since calorimetric is used to determine the concentration of a colored solution, this can be used by using Beer’s law equation c = A/m. This states that the absorbance is proportional to the concentration. A standard calibration curve is a graph that can be used to plot different samples of concentrations and also absorbance to determine the slope of the calibration curve because it is the same as the molar. The way that calibration curve and Beer’s law-related is by absorbance to concentration this is true because when the light of the absorbance increases, the concentration of the solute is proportional to it. The goal of this experiment was to determine the concentration of a substance using Beer’s law, graph a calibration curve for solutions, be able to used and apply molarity and dilution concepts, be able to determine the percent aspirin in a sample, and also calculate the percent yield of a synthesized sample of aspirin. This was used to determine the purity of the aspirin sample because it is extremely important to determine the purity of aspirin synthesized to make sure the sample is pure aspirin because if not, the aspirin can obtain other acids that can be bad if inherited. However, In this experiment, two lab techniques were used to determine the purity of synthesized aspirin, and one of the techniques was using the calorimetric, which was used to find the absorbance and then the dilution to find the exact salicylic acid solution in the sample. When using the calorimetric technique to determine the synthesis of the purity of aspirin, the samples were plotted in the graph of the absorbance and the concentration, from the graph, the absorbance of the aspirin sample was determined to later the percent in the aspirin sample. In this experiment, the calibration curve is important when determining the purity

of aspirin because it provides the slope of the equation where the absorbance can be plotted in the equation provided by it, which is useful to find the percent of aspirin in a sample. The percent of purity was calculated by subtracting the mass of salicylic acid from the mass of the aspirin sample and then dividing the mass of aspirin by the mass of the aspirin sample and multiplying by 100. And then the percent yield was determined by dividing the percent of an aspirin sample by 100 and multiplying by the percent yield which was originally gotten from the previous experiment. In this experiment, there were some safety experimental procedures that had to be taken into consideration since some chemicals are flammable like Ethanol and Acetic anhydride which could not be near the fire, also other chemicals that were toxic if inhaled, harmful if ingested, irritant if contacted with the skin, some others that may cause a fire. Some safety precautions that were considered in this experiment were wearing proper PPE and avoiding breathing any of these chemicals. PART I Calibration Curve of Salicylic Acid Standard Solutions A. Preparation of Stock Salicylic Acid Solutions Mass of Initial Salicylic Acid (D.1)

0.20g

Moles of Salicylic Acid (C.1)

1.45 x10-3moles

Molarity of Salicylic Acid (C.2)

5.8 x10-3M Table 1

Part 1 A Calculations

0.20g

1𝑚𝑜𝑙𝑒 𝑆𝐴 = 1.45x10-3moles 138𝑔𝑆𝐴

250mL

1𝐿 = 0.25mL 1000𝑚𝐿

M=

0.00145𝑚𝑜𝑙𝑒𝑠 = 5.8x 10-3M 0.25𝐿

B. Preparation and Calibration of Salicylic Acid Standard Solutions Part 1 B Calculations 1mL + 0mL + 9mL = 10ML 1mL x 0.0058 = M 1 x 10mL → M1= 5.80 x10-4M 0.75mL + 0.25mL + 9mL = 10ML 0.75mL x 0.0058 = M 2 x 10mL → M2= 4.35 x10-4M 0.5mL + 0.5mL + 9mL = 10ML 0.5mL x 0.0058 = M 3 x 10mL → M3 = 2.9 x10-4M 0.25mL + 0.75mL + 9mL = 10ML 0.25mL x 0.0058 = M 4 x 10mL → M4 = 1.45 x10-4M Standard Calibration Curve Data Trial

Concentration (M)

Absorbance

1

(C.3)

5.80 x10-4M

(D.2) 0.705

2

(C.4)

4.35 x10-4M

(D.3)

0.537

3

(C.5)

2.90 x10-4M

(D.4)

0.332

4

(C.6)

1.45 x10-4M

(D.5)

0.193

Table 2 Regression Line Data

Graph 1 Best-line line or Linear regression equation

(D.6) 1201X - 0.0065

Slope

(D.7) 1201X Table 3

PART II and III Determination of Concentration and Purity of an Aspirin Sample and Percent Yield Part II and III Calculations y = 1201x -0.00650 0.216 =1201x - 0.0065 X=

0.216+ 0.0065 = 1.85x10-4 M 1201

1.85x10-4 x 0.010L = 1.85 x10 -6 SAmoles 1.85 x10-6 moles SAx

138.12𝑔𝑆𝐴 = 2.56x10-4g 1𝑚𝑜𝑙𝑒 𝑆𝐴

𝑥 0.040 → x = 8.0 x10-4 = 0.5𝑚𝐿 25𝑚𝐿 Mass of aspirin in a sample = 0.00080 - 0.000256 = 5.44x10-4 % in aspirin sample =

Final Percent yield =

0.000544 x 100 = 68% 0.00080 68 x 68.7 = 46.72% 100 Code C

Initial mass of aspirin sample (g) (D.8)

0.040g

Absorbance of an aspirin sample (D.9)

0.216

Moles of salicylic acid in an aspirin sample (C.8)

1.85x10-6 moles

Mass of salicylic acid in an aspirin sample (C.9)

2.56x10-4g

Mass of aspirin sample (C.10)

5.44x10-4g

Percent Aspirin in a sample (C.11)

68%

Final Percent Yield of your aspirin synthesis (C.12)

46.72%

Table 4 After conducting the experiment it is important to consider some of the important results from the experiment since a lot of calculations were made during the experiment. In part one, A of the experiment, 0.20g of salicylic acid was measured and then it was converted to moles in order to get the molarity when this was mixed with the solution. After calculating the moles, this was used to convert it to molarity with 250mL of solution that was given. The calculation resulted in 5.8x 10-3M as shown in “Table 1”. When preparing and calibrating the salicylic acid standard solution in part 1 B, different values for the salicylic acid solution and DI water were used as shown in Part 1 B calculations, first, 1mL of salicylic acid was used, and then, 0.75mL, 0.5mL, and finally 0.25mL. For each of these solutions, the concentration was found using the

following equation, M1V1=M2V2 This equation states that the values must be proportional to each other. The data was then collected and shown in “Table 2”. The absorbance of each of the trials was collected by using the calibration standard curve and the determination of the concentration. The data was then collected for each of the trials and collected in “Table 2”. The data collected in part one and part two was then used to plug it into the graph to get a standard curve to determine the percent of salicylic acid in aspirin and to further determine the purity of aspirin. After plugging the values, the graph resulted in a slope of 1201x, and the best-line line or Linear regression equation resulted in y = 1201x- 0.0065 as shown in “Graph 1” And then y was replaced by the absorbance which was collected by the specific code which was 0.216 and from there, it was solved for x as shown in part II and II calculations. Then the moles of SA in aspirin were found, this was accomplished by using the molarity equation which stated that molarity = moles of solute divided by liters of solution. Since 0.0l0L where used, this was multiplied by the molarity to get the moles. Once the moles were calculated, the mass of SA was then calculated to get it in grams. After that to get the mass of aspirin in a sample, was calculated. After the mass of the aspirin sample was calculated, the percent of aspirin in a sample was calculated and recorded in Table 4. After that, the final Percent yield of the aspirin synthesis was calculated using the percent yield that was calculated in the previous experiment. The final percent yield was 46.72% as shown in Table 4. Part I Calibration Curve of Salicylic Acid Standard Solutions The main purpose of this part of the experiment was to prepared different solutions with different amounts of salicylic acid to then dilute each of the amounts. Also, the calorimeter was used to collect the absorbance concentration for each of the solutions, and then graph a calibration curve to determine an equation of absorbance. The way the graph was created was by

plotting the values parameter of concentration and absorbance for each of the solutions that were calculated and collected. The name of the formula of the standard solution was Iron (III) Salicylate complex with its chemical formula which is presented as C21H15FeO9. In this experiment, it was necessary to calculate the concentration of the standard solution because it needed to be plotted in the graph as the concentration for the calibration curve. One of the calculations for the dilutions was, 0.75mL x 0.0058 = M x 10mL→ M= 4.35 x10-4M, which in this case was solved for M. The standard calibration in this experiment was useful to show the relationship between the concentration and the light absorbance, and in this experiment, the best fit line of the calibration curve was y = 1201x - 0.00650. The trend line was then used to determine the purity of aspirin by finding the mass of salicylic acid mass in the sample of aspirin. To later find the mass of aspirin as well as the percent of aspirin in a sample and with the percentage, determine the purity of aspirin.

Part IIDetermination of Concentration and Purity of an Aspirin Sample The amount of salicylic acid present in a sample was related to the purity of the aspirin because it was used to find the mass of aspirin by plugging it into the graph to later be able to calculate the percent of the aspirin sample. In the final sample aliquot, there were 8.0 x10-4 g of aspirin, this was accomplished by using the equation

0.040 𝑥 = and solving for x. The 25𝑚𝐿 0.5𝑚𝐿

absorbance of the aspirin sample was 0.216, the sample of absorbance fell within the calibration curve, it was not calculated again. When using the calibration curve from part I, the concentration of the salicylic acid in the aspirin sample was 1.85x10-4 M. By using the concentration it was more accessible to calculate the get the moles and from there calculate the grams of salicylic acid in the aspirin sample which were 2.56x10-4g this was accomplished by

calculating the moles and then converting it to grams as follows 1.85 x10-6 moles SAx

138.12𝑔𝑆𝐴 = 2.56x10-4g. After that, the grams of aspirin were calculated in a sample which 1𝑚𝑜𝑙𝑒 𝑆𝐴 resulted in 5.44x10-4. This was accomplished by first getting subtracting the mass of salicylic acid from the mass of the aspirin sample. These calculations were done as follows, Mass of aspirin = 0.00080 - 0.000256 = 5.44x10-4. After getting the mass of aspirin, the percent purity of the aspirin sample was calculated which resulted in 68%, this was accomplished by calculating the following, % in aspirin sample =

0.000544 x 100 = 68% 0.00080

Part III Determination of Percent Yield of Aspirin In this part of the experiment, the percent yield that was first calculated was needed to get the final percent yield of the aspirin and the percent yield that was first calculated was 67.8%. Now, when using the percent purity of the aspirin sample that was collected in part II the final percent yield of the aspirin synthesis was 46.72% which is much lower than expected but it is pure aspirin. The calculations that indicated this value was, final Percent yield =

68 x 68.7 = 100

46.72%. Some of the reasons for the low actual yield is because the aspirin was not as pure and once the real purity of the aspirin was calculated, the final percent yield went even lower. The aspirin synthesis percent yield was 68.7% this is because the actual yield was 2.620 and the actual yield was 3.811 this was then calculated as follows, Percent yield =

2.620 × 100 = 3.811

68.7% III. Conclusions The colorimetric technique was used in determining the purity of the synthesis of aspirin

because it was used to determine the concentration of the calibration curve. Once the curve was accomplished, then the absorbance was determined which was then used to calculate the percent of aspirin. In this particular experiment, salicylic acid was the impurity on the aspirin sample that allowed the colorimetric technique for the calculations. Using the purity of aspirin and the limiting and excess reagent that was performed in the previous experiment, the probable reason for the low percent yield was because there was too little of SA and too much of AA which resulted in making a low percent yield, and by using the percent of purity on the aspirin it appeared that there was not 100 percent of the aspirin in the sample. One of the limitations that the calorimetry technique may have is that the sample has to be diluted and or concentrated for the right calculations to be obtained Some of the possible causes of error in this experiment may involve a systematic error, where we cannot make the exact measurements or exact calculations that are necessary, another cause of the error can be a human error when transferring liquids or other substances or when and also when the numbers were rounded to certain decimal places. These errors could be avoided by having better systems to perform the experiment, carefully transfer liquids or solutions, and measuring exact amounts without rounding decimals, and having more accurate results repeat the experiment multiple times. A practical application of this experiment by applying the calorimetric techniques as well as the purity can be used to test the quality of water. With calorimetric techniques, the chemicals that are in the water can be found and can be compared, and analyzed to see if it is drinkable.

RESULTS ANALYSIS QUESTIONS (RAQ) 1. What is Beer’s Law? Beer’s law is the relation between the concentration that has the solute as well as the light absorbed. Since calorimetric is used to determine the concentration of a colored solution, this can be used by using Beer’s law equation c = A/m 2.

What is a colorimeter and why is it used in determining the concentration of the salicylic acid impurity in the aspirin sample? A colorimeter is a visible spectrometer that was used to determine the absorbance of a concentrarion in this case the absorbance of salicylic to determine how much salicylic acid was in the sample of aspirin.

3. What wavelength will I use to determine the concentration of salicylic acid in an aspirin sample? The wavelength that was used to determine the concentration of salicylic acid in an aspirin sample was 565 nanometer. 4. Why is it necessary to follow the colorimetric determination procedures carefully? It is necessary to follow the colorimetric determinations procedures carefully to avoid any erroneous measurement because if it absorbs more light then it will appear as black and if it transmits all the wavelength then it will appear as colorless therefore it must be determined depending on the color of the substance. 5. What are the concentrations of my four standard solutions? What are the absorbance values? The concentration of the four standard solutions was 5.80 x10-4M, 4.35 x10-4M, 2.90 x10-4M, and 1.45 x10-4Mand the four absorbance values were 0.705, 0.537, 0.332, and 0.193

6. What is the relationship between my standard solution concentrations and their light absorbance values as shown in the calibration curve? The relationship between the standard solution and concentration and their light observance values is that as shown in the calibration curve is that the light that id passed by the color solution can be then converted to the absorbance this is true because it is proportional to concentration or molarity of the solution. 7. What was the mass of the salicylic acid impurity, the % purity of aspirin, and the percent yield of the given synthesis condition in the handout? The mass of salicylic acid impurity was 2.56x10-4g, the percent purity was 68%, and the percent yield of the given synthesis condition was 68.7% and the final purity percent yield was 47.72%

Bibliography Encyclopædia Britannica, Inc. (2018). Beer’s Law. Retrieved October 15, 2018, from. https://www.britannica.com/science/Beers-law Randall, J. (2007). Advanced Chemistry with Vernier, third edition. Beaverton, OR: Vernier Software & Technology. Tro, N. (2017). Chemistry: A Molecular Approach. 4th Edition. Boston: Pearson. Chapter 4, Section 4.4, and Chapter 7, Section 7.2. Tro, N. (2017). Chemistry: A molecular approach, 4 th Edition. Boston, MA: Pearson. Vernier Software and Technology, LLC. (2018). Determining the Concentration of a Solution: Beer's Law. Retrieved October 15, 2018, from. https://www.vernier.com/video/vernier-colorimeter-beers-law-tech-tips/...