Bianca Colacin Lab Report 3 - Determination of Weak Acid in Soft Drinks by Potentiometric Titration and Computer Data Analysis PDF

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Experiment I.3 Determination of Weak Acid in Soft Drinks by Potentiometric Titration and Computer Data Analysis 03/04/2020 Bianca Colacin Desk 4

Abstract The main objective of this experiment was to determine the concentrations of a polyprotic acid that is found in most soft drinks. Phosphoric acid is able to donate more than one proton; in order to effectively conduct a potentiometric titration with this analyte sodium hydroxide was used as a titrant. Using 100 mL of Coca Cola soft drink, a pH electrode and a standardized solution of NaOH base, the concentration results of H₃PO₄ was 1.13 ˣ 10⁻⁴ M and for H₂PO₄ was 1.79 ˣ  10⁻⁴ M. The pKa  value of the titration curve is 3.10, with a percent error of 44.2%.

Introduction Potentiometric titration is the use of an electrode to record the voltage emitted by chemical species in solution. The dark color of Coca Cola does not allow the physical change that happens at the endpoint to be observed. The potentiometric method allows for the equivalence point of the solution to be accurately measured without the need of a physical color signal from the indicator neutralization reaction. In order to determine the equivalence points in this experiment the pH of the solution will be a function of titrant delivered. Graphing these data points will show the titration curve of the phosphoric acid that is in the soft drink. At the areas on the graph where the curve inflects, the number of moles of analyte is equal to the number of moles of titrant. In other words, the point of inflection indicates an equivalence point. Dealing with a polyprotic acid will create more than one inflection point in a titration curve. In fact, for each proton dissociated from the acid there is an inflection on the graph. Phosphoric acid, H₃PO₄, titration curve will have three points of inflection. However once the pH of the solution supasses ≅ +11.00, the electrode will begin picking up other ions such as Na⁺. So for this experiment the y-axis will not go above 10.00. At the first equivalence point one proton from the acid will have been deprotonated. This is demonstrated in the following equilibria: (1)

H₃PO₄ + OH⁻ ⇋ H₂PO₄⁻ +  H₂O

As the acid continues to be titrated more protons are removed, by the second equivalence point, the following equilibria occurs: (2) H₂PO₄⁻ + OH⁻ ⇋ H₂PO₄⁻² + H₂O The concentration changes of the acid changes as the composition changes. So the concentrations will be significantly different as the processes progress.

Experimental After creating a 0.10 M solution of NaOH, standardizing it required four trials with a potassium hydrogen phthalate mixture. The solution of KHP and the phenolphthalein indicator is a pink color, allowing for the physical change that happens at the endpoint to be visible. The first two standardizations were done with a traditional acid-base titration method. The pH meter was then

calibrated to a pH of 4.00 and for the last two trials, the titration was conducted with the electrodes inserted into the solution. The electrodes response was compared to the visible endpoint. Once the base solution was standardized, the experiment continued to the Coca Cola titration. 100 mL of the soft drink was poured into a beaker. No indicator was necessary, as mentioned earlier it is not possible to determine the endpoint with that method. With the electrodes in place, the titration proceeded as normal and stopped once the pH reached 10.00. After performing the potentiometric experiment, a Zn²⁺ solution was prepared for Experiment I.4: Determination of Zn in a Cold Relief Lozenge by EDTA Complexometric Titration. Taking a small piece of Zn about ≅0.1g, and transferring it to a volumetric flask. Added to the flask was 10 mL of 3.0 M HNO₃ and deionized water to the graduated mark. This solution was stored in the desks to dissolve.

Results and Discussion A pH meter is connected to two types of electrodes in addition to a glass bulb. The outermost layer of the tool is a silver silver chloride electrode. Serving as a reference electrode, the silver silver electrode wraps around a glass electrode. The glass electrode contains a salt bridge and glass membrane underneath it. Since glass cannot carry a lot of current the potential difference across the glass membrane is measured using this outermost electrode. The glass electrode is ion sensitive and responds specifically to H⁺ because the proton is the only ion able to bind to the gel layer within the membrane. However, when the pH of the solution reaches a certain point, i.e for this pointemetric experiment at a level of ≅10.00-11.00, the electrode starts to respond to the Na⁺ concentration creating a sodium error. This error happens because the concentration of H⁺ is so low. The sodium error lowers the pH than it really is. Being that NaOH is a strong base, it will dissociate completely, the titration curves that will be produced are susceptible to the sodium error. Part I: Standardization of NaOH Getting into the standardization of the base, the best trial was used with 0.5016 g of KHP. Dissolving the compound in 25 mL of deionized water. Converting grams to moles gave a concentration of 0.0982 M. At the equivalence point, where the base neutralizes the acid, 25.35 mL of NaOH have been delivered. To find the concentration of the base inside the flask the moles of acid neutralized were subtracted from the moles of NaOH delivered. The calculated moles was divided by the total volume.

Concentration of KHP: 0.0982 M 0.5016 g of KHP 1 mol KHP = 2.456 ˣ 10⁻³ mol Molarity = 2.456  ˣ 10⁻³ mol = 0.0982 M 204.22 g 0.0250 L Volume of NaOH at equivalence point: 25.35 mL Concentration of NaOH: 0.0016 M KHP + NaOH → KOH + NaHP 2. 456 ˣ 10⁻³ mol 2.535ˣ 10⁻³ mol 0 0 - 2. 456 ˣ 10⁻³ mol - 2. 456 ˣ 10⁻³ mol +2. 456 ˣ 10⁻³ mol +2. 456 ˣ 10⁻³ mol 0 0.000079 mol 2. 456 ˣ 10⁻³ mol 2. 456 ˣ 10⁻³ mol [NaOH] = 0.000079 mol / (0.0250 + 0.02535) L = 0.0016 M Being able to see the physical change allowed for the comparison of the pH at the endpoint alongside the pH at the equivalence point. At the physical color change, the pH recorded was 9.81 versus the point of inflection which was at a pH 6.95. This shows how necessary it is to run a blank for titrations that are done in a traditional manner since there is some error in the endpoint that is unaccounted for otherwise.

The derivative is a slope of the tangent line. In order to calculate the derivative of the pH the point slope formula was used with the data from the standardization. Plotting the first derivative graph showed the peak to be at 25.35 mL with a solution at a pH of 6.95.

Table 1. Data for Standardization Titration and First Derivative Graph Volume (mL)

pH

dpH/dV

1.13

4.09

0.186

2.15

4.28

0.115

3.02

4.38

0.081

4.13

4.47

0.082

5.10

4.55

0.069

6.12

4.62

0.071

7.11

4.69

0.055

8.02

4.74

0.042

9.45

4.80

0.133

10.05

4.88

0.052

11.02

4.93

0.056

12.09

4.99

0.057

13.14

5.05

0.057

14.01

5.10

0.069

15.02

5.17

0.051

16.01

5.22

0.077

17.05

5.30

0.057

18.10

5.36

0.087

19.02

5.44

0.088

20.15

5.54

0.102

21.03

5.63

0.108

22.05

5.74

0.150

23.12

5.90

0.242

24.03

6.12

0.629

25.35

6.95

4.540

25.98

9.81

0.580

27.10

10.46

0.239

28.02

10.68

0.139

29.03

10.82

0.095

30.08

10.92

0.363

Part II: Coca Cola Titration

Table 2. Titration of Coca Cola Soft Drink Volume

pH

dpH/dV

1.48

2.50

0.1547

6.65

3.30

1.2245

7.63

4.50

1.7857

7.91

5.00

0.7519

9.24

6.00

0.2907

12.68

7.00

0.5000

14.68

8.00

0.4120

17.18

9.03

0.0748

30.15

10.00

0.3317

One of the errors with this lab is the lack of data points taken as a group. The first attempt at this lab, barely any data points were collected during the standardization. After coming to that realization more data points were taken for the coca cola titration but more could have been done. During the second attempt there was not enough time to redo the Coca Cola titration for a better collection of data.

Coca Cola’s titration curve has two equivalence points, one at a pH of 4.50 and another at 7.00. The volume of titrant needed for the first equivalence point is 7.63 mL, for the second equivalence point 12.68 mL. To compare the two volumes, it took a little less than twice the amount of the initial volume to reach the second equivalence point. Ve q2 < 2Ve q1 12.68 mL < (2)(7.63 mL) 12.68 mL < 15.26 mL Phosphoric acid is a polyprotic acid. The concentration of H⁺ increases as each proton is removed from the acid. The relationship between concentration and volume is inversely proportional, so as the concentration increases, the volume of analyte decreases. Ideally the volume of equivalence should be the same for every point, however if the experiment could continue to the third equivalence point the relationship between the three volumes would be, Ve q3 < (3)Ve q2 < (2)Ve q1. The CO₂ gas particles within the soft drink could affect the titration curve by reacting with the isolated protons forming formic acid. That could potentially throw the curve off since the concentration of H⁺ could fluctuate between being isolated or forming a weak acid.

The lack of data points taken ultimately hindered the way the first derivative graph came out. The second peak is not defined, additional data points would have resulted in a smoother graph. Using the first equivalence point of the Coca Cola titration, 7.63 mL, and the concentration calculated from the standardization process, 0.0016 M, concentration of phosphoric can be calculated. The moles of NaOH at equivalence are 1.22 ˣ 10⁻⁵ mol. -

0.0016 M (0.00763 L) = 1.22 ˣ 10⁻⁵ mol NaOH

The moles of NaOH are equal to the moles of H₃PO₄, the total volume is 0.10763 L. After calculations the molarity is 1.13ˣ 10⁻⁴ M. -

[H₃PO₄] = 1.22 ˣ 10⁻⁵ moles H₃PO₄/ 0.10763 L = 1.13 ˣ 10⁻⁴ M

Applying the same methods to determine the concentration of the intermediate form of phosphoric acid, H₂PO₄⁻. At the second equivalence point the volume of base delivered is 12.68 mL with the concentration of 0.0016 M. The moles of NaOH are 2.02ˣ 10⁻⁵ moles therefore the moles of the acid are 2.02ˣ 10⁻⁵ moles as well. With a total volume of 0.11268 L, the concentration resulted to 1.79 ˣ 10⁻⁴ M. Finding the pKₐ value involves finding the Kₐ constant first. Using the concentrations from  ₐ came out to be 7.84 above and the following equilibria H₃PO₄ + OH⁻ ⇋ H₂PO₄⁻ +  H₂O, the K ˣ 10⁻⁴. The pKₐ is then -log(7.84 ˣ 10⁻⁴) = 3.10. According to the literature the accepted value is 2.15, the percent error for this experiment is 44.2 %. (3.10-2.15)/2.15 = 0.442 ˣ 100 = 44.2% Summary of Coca Cola Titration V₁eq = 7.63 mL pH: 4.50

V₂eq =  12.68 mL pH: 7.00

Molarity of H₃PO₄ = 1.13 ˣ 10⁻⁴ M Molarity of H₂PO₄⁻ = 1.79 ˣ 10⁻⁴ M...