Experiment 3 Post-Lab Informal Report PDF

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Experiment 3: Spectroscopic Determination of the pKa of the Acid-Base Indicator Bromothymol Blue Trevor Du 10/6/2021 Mikaylin Nogler, CHEM 241L, Section 417

"I pledge that I have not used someone else’s old or current lab report when writing this lab report. I pledge that I did not collaborate with any other students, except where allowed, and that the report I submitted here contains my own ideas, thoughts, observations, calculations, data, conclusions and answers. Lastly, I pledge that the data presented in this report was my own collected during lab or provided to me by my TA.”

Introduction: Acids and bases have many applications in everyday life ranging from digestion of the food we eat to the products we use to clean our bathrooms. Measuring pH is highly important in the study of freshwater ecosystems. Compounds such as sulfur dioxide and nitrogen oxides originating from power stations and industrial plants mix with water vapor to cause acid rain, lowering the pH of freshwater biomes. This increased acidity can interfere with female fishes’ reproductive cycles, halting the growth of fish populations. The acidity can also cause the leaching of aluminum (Al3+) ions out of soil and into fresh water, which are highly toxic to fish (1). Acid-base indicators are used to determine whether a solution is acidic, neutral, or basic. They are often organic weak acids that change color upon deprotonation. Bromothymol blue is one such acid-base indicator. UV-visible spectroscopy is a powerful tool in finding the concentration of a given species in a solution. A UV-Vis spectrophotometer measures the intensity of light transmitted through a sample compared to a reference measurement of the incident light source. Using UV-visible spectroscopy, we can measure and plot the absorbance of bromothymol blue at various pH levels against wavelength to find its concentration. The goal of this experiment is to determine the pKa of the acid-base indicator bromothymol blue using UVVisible spectroscopy and a plot of dAbsorbance/dpH against average pH.

Results and Calculations: 1.4 1.2

Absorbance

1 0.8 pH 4.98

0.6

pH 10.01

0.4 0.2 0 380

480

-0.2

580

680

780

880

Wavelength (nm)

Figure 1. Plot of absorbance vs. wavelength for the bromothymol blue indicator at pH 4.98 and pH 10.01. The λmax values are 429.9 nm for the HIn form and 613.5 nm for the In- form. Table 1. Wavelength of Maximum Absorbance (λmax) pH

λmax (nm)

Absorbance at λmax

pH 4.98

429.9

0.508

pH 10.01

613.5

1.203

Table 2. Absorbance at each λmax for various pH values pH 4.98

pH 6.03

pH 7.00

pH 7.31

pH 7.62

pH 7.97

pH 9.00

pH 10.01

429.9 nm

0.508

0.491

0.389

0.324

0.227

0.211

0.106

0.331

613.5 nm

-0.038

0.079

0.545

0.843

0.848

1.053

1.105

1.203

1.4 1.2

Absorbance at λmax

1 0.8 429.9 nm

0.6

613.5 nm

0.4

0.2 0

4.98 -0.2

5.98

6.98

7.98

8.98

9.98

pH

Figure 2. Plot of absorbance at the λmax of HIn and In- forms of BTBlue at varying pH 1.2 1

dAbsorbance/dpH

0.8 0.6 429.9 nm

0.4

613.5 nm

0.2 0 5.5

6.5

7.5

8.5

9.5

-0.2 -0.4

Average pH

Figure 3. Plot of dAbsorbance/dpH vs. Average pH for both λmax values

1.2 1

Absorbance

0.8 0.6 pH 4.98 0.4

pH 10.01

0.2 0 380

480

-0.2

580

680

780

880

980

Wavelength (nm)

Figure 4. Plot of absorbance vs. wavelength for both extraction samples Sample Calculations At the 613.5 nm wavelength, between pH 7.00 and pH 7.31: 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑝𝐻 = 7.31 − 7.00 = 0.31 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑎𝑏𝑠𝑜𝑟𝑏𝑎𝑛𝑐𝑒 = 0.843 − 0.545 = 0.298 1 𝑎𝑣𝑒𝑟𝑎𝑔𝑒 𝑝𝐻 = (7.00 + 7.31) = 7.16 2 ∆𝐴𝑏𝑠 0.298 𝑑𝐴𝑏𝑠 = = = 0.96 𝑑𝑝𝐻 ∆𝑝𝐻 0.31

𝑝𝐾𝑎 ≈

% 𝑒𝑟𝑟𝑜𝑟 =

7.16 + 7.47 ≈ 7.32 2

7.32 − 7.1 ∗ 100 = 3.1% 𝑒𝑟𝑟𝑜𝑟 7.1

Literature value of pKa of BTBlue = 7.1

End-of-Lab Report Questions: 1) How did the color change of the bromothymol blue solution as the solution became more basic? How does this relate to the structure of the molecule? The color of the bromothymol blue solution changed from yellow to a deep blue as the pH increased. This color change occurs because of the structural reorganization of the bonds in bromothymol blue as it loses a proton. The acidic form of the indicator contains a sulfur atom double bonded to two oxygens, while in the basic form the sulfur is bonded to 3 oxygens and has a -1 formal charge (2). 2) What is happening to the concentrations of the acidic and basic forms of bromothymol blue as the pH increases? How do you know? Referencing your data, support your answer. From Figure 2, we can see that at the λmax for the basic form, the absorbance increases as pH goes up. Meanwhile, at λmax for the acidic form, the absorbance decreases with increasing pH. Beer’s Law tells us that for a fixed path length (which we have), absorbance is directly proportional to concentration. This allows us to draw the conclusion that as the pH of the solution increases, the concentration of the basic form of bromothymol blue increases while the concentration of the acidic form decreases. 3) Does the first derivative analysis (Figure 3) yield the same pKa value when using the absorbance values at the λmax of both the acidic and basic forms? Does this make sense? Why or why not? No, the pKa value determined using the λmax of the acidic form was slightly (0.21) higher than the pKa determined using the λmax of the basic form. This does not make sense, and was the

result of experimental error. The pKa of an acid is a fixed value and does not depend on what forms of the acid are present nor the concentration of these forms. 4) What was the pKa you determined for bromothymol blue? How does your experimentallydetermined pKa compare to the literature value? What was the percent error? What error(s) contributed to the discrepancies? The pKa of BTBlue that was determined from the data was 7.32. The literature value for this pKa is 7.1 (3). These values differ by 0.22, and the percent error is 3.1%. A possible systematic error that could have contributed to this discrepancy is scratches present on the sides of the cuvette which may have skewed the measured values of absorbance. 5) Describe how the results obtained from the liquid-liquid extraction relate to the structure of the acidic and basic forms of bromothymol blue. The peak absorbance of the basic form of BTBlue (In-) was found to be much higher than the acidic form (HIn). This means the basic form showed a higher affinity than the acidic form for the aqueous phase. This makes sense structurally because In- is a charged molecule, while HIn is uncharged. We expect charged molecules to have a higher affinity to the aqueous phase because water is a very polar molecule, while organic phase solvents such as ethyl acetate are much less polar. 6) Was the experiment successful? What results were particularly significant? Were you able to calculate the pKa of bromothymol blue? Why is this important? Based on the error discussed, how could the experiment be improved? The experiment was somewhat successful; we were able to calculate the pKa of bromothymol blue within a 3.1% margin of error, and our data for the solvent-solvent extraction is consistent

with what we would expect given the chemical structure of the two forms of BTBlue. The color was significant because it gave us a way to qualitatively observe the change in concentration of the two forms of BTBlue as the pH was adjusted. It is important to know the pKa of acid-base indicators because they provide points of reference during titrations as to when the equivalence point is reached. The experiment could be improved by taking all measurements at more than just 8 pH values. Increasing the sample size like this helps to reduce random error in data analysis. References: 1. EPA, U. S.; OAR. Effects of Acid Rain. 2016. 2. Laboratory in Separations and Analytical Characterization of Organic and Biological Compounds. Chapel Hill, NC: UNC Student Stores Course Pack Publishing; 2021.

3. Sigma-Aldrich. Bromothymol blue ACS reagent, Dye content 95 %; 2021....