General Chemistry 2 - Lab Report 5 PDF

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Abdul Rahim 1

General Chemistry 2 SCC202.239B Afsana Abdul Rahim Lab Partners: Sarah, Jacob Experiment 5: Studying the pH of Strong Acid, Weak Acid, and Buffer Solutions 04/17/2018 Professor Amit Aggarwal

Abdul Rahim 2 Objectives: The objective of this lab was to compare the calculated and measured pHs of a series of hydrochloric acid and acetic acid solutions. Students were expected to calculate the acid dissociation constant of acetic acid using pH measurements, measure the pH of various salt solutions, and calculate the hydrolysis constant of ammonium chloride using pH measurements. In addition, students were expected to compare measured and calculated pHs of a buffer solution; the same buffer solution mixed with HCl and mixed with NaOH solution; distilled water; and distilled water mixed with HCl solution and mixed with NaOH solution. Materials/chemicals required: 

Six 16 125 mm test tubes

 0.10 M sodium carbonate (Na2CO3)



0.10 M hydrochloric acid (HCl)

 0.10 M ammonium chloride (NH4Cl)



Two 50 mL graduated cylinder

 0.10 M ammonium acetate (NH4C2H3O2)



Distilled water

 Glass stirring rod



Two 10 mL cylinders

 6 M acetic acid (HC2H3O2)



400 mL beaker

 Solid anhydrous sodium acetate



pH meter

(NaC2H3O2)



6 M sodium hydroxide (NaOH)

 6 M hydrochloric acid (HCl)



0.10 M acetic acid (HC2H3O2)

 150 mL beaker



0.10 M sodium chloride (NaCl)

 Electronic balance



0.10 M sodium acetate (NaC2H3O2)

 Four 50 mL beakers



0.10 M sodium bicarbonate (NaHCO3)

Abdul Rahim 3 Methods: I. Preparing HCl Solutions and Determining pH 1. Four 16 125 mm test tubes were labeled “1,” “2,” “3,” and “4.” 2. 10 mL of 0.10 M HCl was put into a clean, dry 10 mL graduated cylinder. 3. 5.0 mL of this solution was transferred into test tube 1. 4. The remaining 5 mL of HCl solution was transferred into a clean, dry 50 mL graduated cylinder. The 10 mL cylinder was rinsed with small quantities of distilled water and then the rinses were added to the 50 mL cylinder. The solution in the 50 mL cylinder was diluted to exactly 50 mL with distilled water. The solution was thoroughly mixed with a glass stirring rod. 5. 5 mL of the solution prepared to Step 4 was transferred to test tube 2. 6. 5 mL of the solution prepared in Step 4 was transferred to a second 50 mL graduated cylinder. The remaining 40 mL of solution from the first 50 mL cylinder was transferred into a 400 mL beaker labeled “Discarded Solutions.” The empty cylinder was rinsed and the rinses were added to the discard beaker. 7. The solution in the 50 mL cylinder was diluted to exactly 50 mL with distilled water. 5 mL of this solution was transferred to test tube 3. 5 mL was transferred into a clean, 50 mL graduated cylinder. The remaining 40 mL of diluted solution was transferred to the “Discarded Solutions” container, as in Step 6. 8. 5 mL of the solution in the 50 mL cylinder was diluted to exactly 50 mL with distilled water. 5 mL of the diluted solution was transferred to test tube 4. The remaining 40 mL of solution in Step 7 was discarded. 9. A pH meter was used to measure the pH of the four solutions in the test tubes. The solution

Abdul Rahim 4 was transferred from test tube 1 into a clean, dry 50 mL beaker. The electrode was thoroughly rinsed with distilled water. The electrode was carefully lowered into the beaker until the glass bulb was completely beneath the solution surface. The beaker was carefully swirled so that the solution made good contact with the electrode. The solution pH was read and recorded on Data Sheet 1. The above procedure was repeated using the solutions in the other three tubes. All HCl solutions and rinses were transferred into the “Discarded Solutions” beaker. The graduated cylinders and test tubes were rinsed and drained, transferring the rinses into the “Discarded Solutions” beaker. The electrode was thoroughly rinsed with distilled water. II. Preparing HC2H3O2 Solutions and Determining pH 10. 10 mL of 0.10 M HC2H3O2 was put into a clean, dry 10 mL graduated cylinder. 11. Steps 3-9 were repeated using the HC2H3O2 solution in place of the HCl solution. All pH readings were recorded on Data Sheet 1. III. Determining the pH of Various Salt Solutions 12. 5 mL of the following aqueous salt solutions were put into six clean, dry, appropriately labeled 16 125 mm test tubes: 0.10 M NaCl, 0.10 M NaC2H3O2, 0.10 M NaHCO3, 0.10 M Na2CO3, 0.10 M NH4Cl, and 0.10 MNH4C2H3O2. 13. The NaCl solution was transferred into a clean, dry 50 mL beaker. 14. Using the pH meter, the pH of the NaCl solution was measured. This pH was recorded on Data Sheet 2. The solution was transferred into the “Discarded Solutions” beaker. 15. The beaker was rinsed twice with distilled water. The rinses were transferred into the “Discarded Solutions” beaker. The electrode was washed with distilled water. 16. Steps 13-15 were repeated using each of the five remaining salt solutions. 17. All glassware was rinsed and drained.

Abdul Rahim 5 IV. Preparing a Buffer Solution 18. A clean 150 mL beaker was weighed to the nearest 0.1 g. 19. 2.0 g of solid anhydrous NaC2H3O2 was added to the beaker. 20. The beaker was removed from the balance. Using a 10 mL graduated cylinder, 4.0 mL of 6 M HC2H3O2 was added to the beaker. 21. Using a 50 mL graduated cylinder, 46 mL distilled water was added to the beaker. The solution was stirred until the NaC2H3O2 was completely dissolved. This solution contained 2.4 x 10-2 mol each of NaC2H3O2 and HC2H3O2 in 50 mL of solution. 22. Four clean, dry beakers were labeled “1,” “2,” “3,” and “4.” 23. Using a 50 mL graduated cylinder, 25 mL of distilled water was poured into beakers 1 and 3. 24. Using a graduated cylinder, 25 mL of the buffer solution was added each to beakers 2 and 4. 25. The pH of the solutions was measured in beakers 1 and 2. The electrode was rinsed with distilled water after each measurement. The data were recorded on Data Sheet 2. 26. 0.5 mL OF 6 M HCl was transferred each to beakers 1 and 2. The solution was thoroughly mixed with a glass stirring rod; the rod was rinsed with distilled water after mixing each solution. The pH of each solution was measured, and the data were recorded on Data Sheet 2. The solutions were transferred to the “Discarded Solutions” beaker. 27. 0.5 mL of 6 M NaOH was transferred each to beakers 3 and 4. The solution was thoroughly mixed with a glass stirring rod; the rod was rinsed with distilled water after mixing each solution. The pH of each solution was measured, and the data were recorded on Data Sheet 2. The solutions were transferred to the “Discarded Solutions” beaker.

Abdul Rahim 6 Results: I. Preparing HCl Solutions and Determining pH Table 1. Measured pH and theoretical pH for different HCl concentrations. Concentration of HCl, M

Measured pH

Theoretical pH

1.0 x 10-1

1.11

1.0

1.0 x 10-2

2.07

2.0

1.0 x 10-3

3.43

3.0

1.0 x 10-4

6.45

4.0

1. Calculation for theoretical pH using the 1.010-1 M sample of HCl pH = -log [H3O+] = -log (1.0 x 10-1) = 1.0

(Equation 1)

II. Preparing HC2H3O2 Solutions and Determining pH Table 2. Measured pH, theoretical pH, calculated Ka,, and literature Ka for different HC2H3O2 concentrations. Concentration of HC2H3O2, M

Measured pH

Theoretical pH

Calculated Ka of HC2H3O2 based on pH data

Literature Ka of HC2H3O2 b

1.0 x 10-1

2.91

2.87

1.51 x 10-5

1.8 x 10-5

1.0 x 10-2

3.48

3.37

1.89 x 10-5

1.8 x 10-5

1.0 x 10-3

4.42

3.87

1.44 x 10-6

1.8 x 10-5

1.0 x 10-4

6.46

4.37

1.22 x 10-9

1.8 x 10-5

The dissociation of HC2H3O2 in water: HC2H3O2 (aq) + H2O (l) ⇌ C2H3O2- (aq) + H3O+ (aq)

(Equation 2)

Sample calculations using 1.0 x 10-1 M HC2H3O2 2. Theoretical pH Ka = ([C2H3O2-] [H3O+]) / [ HC2H3O2 ] = [H3O+]2 / [ HC2H3O2 ] 1.8 x 10-5 = [H3O+]2 / [ 1.0 x 10-1]

(Equation 3)

Abdul Rahim 7 1.8 x 10-6 = [H3O+]2 0.00134 = [H3O+] pH = -log [H3O+] = -log (0.00134) = 2.87 3. Calculation for Ka using measured pH values for solution [H3O+] = 10-pH = 10-2.91 = 1.23 x 10-3 M

(Equation 4)

Ka = ([C2H3O2-] [H3O+]) / [ HC2H3O2] = [H3O+]2 / [ HC2H3O2] Ka = (1.23 x 10-3)2 / (1.0 x 10-1) = 1.51 x 10-5 4. Percent errors for Ka values Percent error, % = (calculated Ka - literature Ka) / (literature Ka) = ¿

(1.51 x 10−5−1.8 x 10−5 ) ∨x 100 = 16.1% (1.8 x 10−5 )

The same calculation was applied to the other concentrations. The percent errors for the Ka values of the 1.0 x 10-1 M, 1.0 x 10-2 M, 1.0 x 10-3 M, and 1.0 x 10-4 M solutions were 16.1%, 5%, 92%, and 99.9% respectively. III. Determining the pH of Various Salt Solutions Table 3. The pH of salt solutions. Explanation for difference between pH and 7.0

Salt

Salt type Experimental/Theoretica l

Measured pH

NaCl

Acidic/Neutral

3.62

Little hydrolysis was expected to occur. Low pH may either be from contaminants or mis calibrated pH meter.

NaC2H3O2

Basic/Basic

8.92

Only the anion C2H3O2- hydrolyzed, so the solution was basic

NaHCO3

Basic/Basic

8.95

Only the anion HCO3- hydrolyzed, so the solution was basic

Abdul Rahim 8

Na2CO3

Basic/Basic

10.95

Only the anion CO32hydrolyzed, so the solution was basic

NH4Cl

Acidic/Acidic

6.67

Only the cation NH4+ hydrolyzed, so the solution was acidic

NH4C2H3O2

Basic/Neutral

7.56

Both ions can hydrolyze, but more of the anion C2H3O2- hydrolyzed, so the solution was basic

5. Calculated Kh NH4Cl, based on measured pH: [H3O+] = 10-pH = 10-6.67 = 2.14 x 10-7 M Kh = [NH3] [H3O+] / [NH4+] = (2.14 x 10-7)2 / 0.10 = 4.58 x 10-13

(Equation 5)

IV. pH Changes in Water and Buffer Solutions upon Addition of HCl and NaOH Solutions. Table 4. Measured pH, theoretical [H3O+], and calculated pH for distilled water and buffer. Solution

Measured pH

Theoretical [H3O+]

Calculated pH

Distilled water

4.7

1.0 x 10-7

7.0

Original buffer solution

4.6

1.8 x 10-5

4.7

Distilled water with added 6 M HCl

1.45

1.2 x 10-1

0.92

Buffer solution with added 6 M HCl

4.23

2.3 x 10-5

4.6

Distilled water with added 6 M NaOH

12.13

8.5 x 10-14

13

Buffer solution with added 6 M NaOH

4.88

1.4 x 10-5

4.9

Sample calculations using original buffer solution and buffer solution with added 6 M HCl 6. Theoretical [H3O+] and calculated pH of original buffer solution: 1.8 10-5 = [H3O+] [C2H3O2-] / [HC2H3O2] = [H3O+] [6] / [6] [H3O+] = 1.8 x 10-5 pH = -log [H3O+] = -log (1.8 x 10-5) = 4.7 7. Theoretical [H3O+] and calculated pH of buffer solution with added 6 M HCl

(Equation 6)

Abdul Rahim 9 Moles of NaC2H3O2 = 2.0 g NaC2H3O2 (1 mol / 82.0337 g) = 0.024 mol NaC2H3O2 Moles of HC2H3O2 = 6 M 0.004 L = 0.024 mol HC2H3O2 Equation after the addition of 6 M HCl: H3O+ (aq) + C2H3O2- (aq)  HC2H3O2 (aq) + H2O (l)

(Equation 7)

Moles of HCl added = 6 M 0.0005 L = 0.003 mol HCl New [C2H3O2-] after the addition of HCl: 0.024 - 0.003 = 0.021 M New [HC2H3O2] after the addition of HCl: 0.024 + 0.003 = 0.027 M 1.8 10-5 = [H3O+] (0.021) / (0.027) [H3O+] = (1.8 x 10-5) (0.027) / (0.021) = 2.3 10-5 pH = -log ([H3O+] = -log (2.3 10-5) = 4.6 Discussion: As expected, the solutions in the first experiment increased in pH as the concentration of HCl decreased (Table 1). HCl is a strong base, so it completely dissociates to release H + ions and form H3O+ in an aqueous solution (Bishop et al., 2016, p. 72). Since the value for pH is the negative logarithm of H3O+ concentration (Brown et al., 2014, p. 674), pH and H+ concentration are inversely proportional. The results from the first experiment were consistent with this observation. The measured values for pH were also close to the theoretical values, which were based on the pH formula (Equation 1) except for the 1.0 x 10-4 M concentration of HCl which had an increase in pH value compared to theoretical values. This could be due to contaminants such as excess acid or excess water or mis calibrated pH meter that caused an increase in pH value. As expected, the HC2H3O2 solutions in the second experiment had higher pH values than the HCl solutions in first experiment even though the same concentrations were used (Table 2).

Abdul Rahim 10 HC2H3O2 is a weak acid, so it is a weak electrolyte and only partially dissociates to release H+ ions in an aqueous solution (Brown et al., 2014, p. 680). As a result of the incomplete dissociation, the HC2H3O2 solutions had relatively lower H3O+ concentrations and higher pHs. Based on the measured pH, the acid-dissociation constant, Ka, for the 1.0 x 10-1 M, 1.0 x 10-2 M, 1.0 x 10-3 M, and 1.0 x 10-4 M solutions were 1.51 x 10-5, 1.89 x 10-5, 1.44 x 10-6, and 1.22 x 10-9 respectively (Table 2). Compared to the literature Ka, the calculated Ka for the values were 16.1%, 5%, 92%, and 99.9% respectively. The high percentage errors can be explained by the fact that the Ka calculations were derived from several exponential equations (Equations 3 and 4) that amplified small differences between the measured and theoretical pHs. The measured and theoretical pHs were very close in magnitude (Table 2), but small differences between them produced notably different Ka values. Since exponents are the number of times that a number is multiplied by itself, small differences between exponents can lead to dissimilar results. For example, 3 and 3.1 are very close in magnitude, but they produce very different results as exponents of the base 10 (103 is equal to 1000 whereas 103.1 is equal to 1259). In the experiment, measured pHs were used as negative exponents to calculate the H3O+ concentration. One possible explanation for the slight inaccuracy in measured pHs is that the solutions may have been contaminated with excess acid or excess water. The presence of excess water would explain why the measured pH was higher than expected the solutions. Another possible explanation for the inaccuracy is that the pH was not well calibrated or was not sensitive enough to accurately measure the pH. This would also explain the discrepancy between the measured pH values and the theoretical values. As expected, the basic salt solutions in the third experiment had pHs that were higher than 7 and the acidic salts had pHs that were lower than 7 (Table 3). Ions from salts can act as

Abdul Rahim 11 acids or bases (Zumdahl & Zumdahl, 2008, p. 671). In the experiment, NaC2H3O2, NaHCO3, and Na2CO3 produced basic solutions. This was consistent with expectations because they are basic salts, which means that they are the products of a reaction between a strong base and a weak acid. The anion of a basic salt hydrolyzes and reacts with H+ in water, making the solution more basic. NH4Cl produced more acidic solution. This was consistent with expectations because NH4Cl is an acidic salt, which means that it is the product of a reaction between a strong acid and a weak base. The cation of an acidic salt hydrolyzed and contributes H+ to water, making the solution more acidic. NH4C2H3O2 produced basic solution. NH4C2H3O2 is a salt that is the product of a weak base and a weak acid and it can be either acidic, basic, or neutral since both of its ions hydrolyze (Bishop et al., 2016, p. 75). In the experiment, more of its anions, C2H3O2- hydrolyzed since the solution was slightly basic. Based on the measured pH, the hydrolysis equilibrium constant was 4.58 x 10-13. Finally, NaCl produced a acidic solution, which was inconsistent with expectations. NaCl is the product of a strong base and a strong acid, so it is a neutral salt. The ions in neutral salts do not measurably hydrolyze, so the solution was expected to be neutral at a pH of 7. It is possible that the solution was contaminated with traces of acids from previous experiments, which would have skewed the pH reading to be more acidic. In the last experiment, the buffer solutions experienced less pH changes compared to distilled water when HCl and NaOH were added (Table 4). This was consistent with expectations because buffer solutions resist large changes in pH when small amounts of a strong acid or a strong base is added to them (Brown et al., 2014, p. 721). Buffers are able to resist changes in pH because they are comprised of a weak acid-base conjugate pair that can neutralize both OH- ions and H+ ions that are added. In the experiment, the buffer contained both C2H3O2- and HC2H3O2.

Abdul Rahim 12 When 6 M HCl was added to the buffer, the pH of the buffer only decreased by 0.37. In contrast, the pH of the distilled water decreased by 5.45 after the addition of HCl. For the addition of 6 M NaOH, the pH of the buffer increased by only by 0.28 while the pH of distilled water increased by 5.23. Conclusion: In the first experiment, pH and H3O+ of HCl solutions were inversely proportional; pH decreased as the concentration of H3O+ increased. In the second experiment, the HC2H3O2 solutions had relatively higher pHs compared to the HCl solutions of the same concentrations. The acid-dissociation constant, Ka, was calculated based on the measured pHs of the HC2H3O2 solutions. The Ka values had significant percent errors because there were some discrepancies between the measured pHs and the theoretical pHs. The discrepancies can be attributed to either the presence of contaminants in the solutions or to a mis calibrated pH meter. In the third experiment, basic salts produced solutions with pHs above 7 and the acidic salt produced a solution with a pH below 7. NH4C2H3O2, which has the capacity to be an acidic, basic, or neutral salt produced a slightly basic solution. Based on the measured pH, the hydrolysis equilibrium constant was 4.58 x 10-13. In the last experiment, distilled water underwent large changes in pH after the addition of HCl and NaOH. In contrast, the buffers resisted drastic changes in pH after the addition of HCl and NaOH.

Abdul Rahim 13

References Bishop, C. B., Bishop, M. B., Whitten, K. W., Seager, S. L., Slabaugh, M. R. (2016). Fundamentals of Chemistry II. (pp. 71-82). Boston, MA: Cengage Learning. Brown, T. E., LeMay, H. E., Bursten, B. E., Murphy, C., Woodward, P., Stoltzfus, M. E. (2014). Chemistry: The Central Science. (13th ed., pp. 674-680, 721). Upper Saddle River, NJ: Prentice Hall. Zumdahl, S. S. and Zumdahl, S. A. (2008) Chemistry. (8th ed. pp. 671). Belmont, CA: Brooks Cole....