Lab 1-Spectrophotometer PDF

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For Dr. Lu...

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Spectrophotometric Analysis of Commercial Aspirin By: Hayden Casassa and Jonathan Baugh Section 0903 February 7th, 2016

Learning Objectives: 1. To learn about the absorption of light by using a spectrophotometer 2. To create a Beer’s Law curve and use it to calculate the concentration of an unknown substance 3. To learn how to prepare stock solutions and the importance behind them. 4. To practice pipetting and weighing small samples which were previously learned

Data and Results: Table 1: Data and Results for Stock Solution of pure aspirin in Part 1 Mass of Acetylsalicylic acid

0.161 grams

Moles of Acetylsalicylic acid

8.93 x 10-4 mol

Concentration of Acetylsalicylic acid in 100.0 ml volumetric flask

8.93 x 10-3 M

Mass of aspirin sample 1

0.162 grams

Mass of aspirin sample 2

0.164 grams

Table 2: Data and Results for Five standard aspirin solutions in Part 1 Solution

Concentration

Absorbance

A

4.465 x 10-4 M

0.450

B

3.572 x 10-4 M

0.355

C

2.679 x 10-4 M

0.272

D

1.786 x 10-4 M

0.189

E

8.93 x 10-5 M

0.100

Sample 1

1.87 x 10-4 M

0.230

Sample 2

M

0.245

Table 3: Data and Results Table for Commercial Aspirin Tablet in Part 2 Sample 1

Sample 2

Mass of Sample

0.162 grams

0.164 grams

Absorbance from Graph

0.187

0.245

Concentration in diluted solution

0.000187 M

0.000251 M

Concentration in Sample Stock solution

0.006233 M

0.00837 M

Moles of acetylsalicylic acid in stock solution

0.0006233 moles

0.000837 M

Mass of acetylsalicylic acid in stock solution

0.112 grams

0.150

Percent acetylsalicylic acid in sample

69.1%

92.0%

Average Percent acetylsalicylic acid

80.6%

Sample Calculations: Part I: Moles of Acetylsalicylic Acid (Standard): 0.161g/ 180.2 g (mm) =0.000893 moles Concentration of Acetylsalicylic Acid in 100.0 mL volumetric flask (Standard): .000893mol/ 0.1L= =0.00893 Part II: Concentration of solution A: MStockVStock=MDiluteVDilute

(.00893 M)(0.50 mL) = (MDilute)(10.0 mL) MDilute = .000446 M

Concentration of commercial aspirin in diluted solution using trend line (Sample 1): y(absorbance) = mx(concentration) + b y= 1.306x - 0.0137 .230 = 1.306(x) - 0.0137 x (concentration) = 0.000187 M Concentration of Sample 1 Stock Solution: MDiluteVDilute=MStockVStock (0.000187 M)(10.0 mL) = (MStock)(0.30 mL) Stock 0.006233 M Moles of Acetylsalicylic Acid in Stock Solution (Sample 1): 0.006233 M =(Moles) / .1 L Moles in Stock Solution= 0.0006233 moles

Mass of Acetylsalicylic Acid in Stock Solution (Sample 1): 0.0006233 moles x 180.2g/mol= =0.112 grams Percent of Acetylsalicylic Acid in Sample 1: 0.112 grams /0.162 grams x 100 =69.13 %

Average percent Acetylsalicylic Acid:69.1% + 92.0% / 2= =80.6% Percent Error: 86% - 80.6%/86%= = 6.3 % Standard Deviation (69.1-86)2=285.61 (92.0-86)2= 36 (285.61+36)1/2= 17.9

Discussion:

The Spectrophotometric analysis of a commercial aspirin utilized the Beer-Lambert Law. Beer’s Law equation is known as the molar absorptivity of the solution multiplied by the path width of the cell and the concentration of the solution. Multiplied together, they equal the measured absorbance of the solution. Beer’s Law generates a Linear line with the equation y=mx+b where y is the absorbance and x is the concentration. After obtaining the absorbance and concentrations of samples A-E in Part 1, it gave us the data we needed to calculate the linear equation. The trend line that was subsequently generated enabled us to find the concentration of the Aspirin complex by plugging the absorbance in for y. The spectrophotometer can be used to determine the concentration of a color solution. A higher absorbency means that the solution has a higher concentration because the solution absorbs more of the light. Ideally, pure water or any other transparent liquid should have 100% transmittance and a completely opaque liquid would have 0% transmittance, or an absorbance of 1. Sodium Hydroxide when combined with Iron (III) Chloride as an endothermic reaction produces the Iron Aspirin Complex (NaOH + FeCL3 + heat —> Iron Aspirin Complex). Finally, we were able to calculate the Mass % of the aspirin by determining molarity of the solution and using stoichiometry to convert it to grams. We then took the determined grams and divided it by the initial mass to determine the percent of aspirin in the pill. The calculated percent of acetylsalicylic acid in samples 1 and 2 were 69.1% and 92% respectively. The calculated standard deviation for the lab was 17.9 %. Our overall precision for the lab was not very accurate as our percent error was nearly 20% different for the 2 trials. However, one of our values was significantly lower then the theoretical value of 86% (69.1%) while the other one was slightly higher, enabling the accuracy to have less error then the precision. Our percent error was 6.3% which is moderately high proving that there were errors made in the lab. Two sources of error that were prevented in the lab was the wiping of the glassware before inserting into the spectrophotometer along with the use of stock solution. However, one source of error could be the lack of all the aspirin complex being transferred over from the bowl when being ground up along with residue being stuck to the inside of the pipette and the weigh boat. Nevertheless, the results were still accurate enough to form the Beer’s Law graph with the trend line hitting all 5 of the points with no major outliers. In the future, this lab could have avoided more sources of error by not only making sure all the residue was transferred from one container to the next, but by also performing two trials of part one of the lab and using the average absorption and concentration of the two trials to generate a more accurate curve of Beer’s Law curve....