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Chapter 15: Applications of Aqueous Equilibria

Chapter 15: Applications of Aqueous Equilibria. Buffers Common Ion Effect Henderson-Hasselbalch Equation Buffer Capacity Acid-Base Titrations & Titration Curves Strong Acid-Strong Base Titrations Weak Acid-Strong Base Titrations Strong Acid-Weak Base Titrations Acid-Base Indicators

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Chapter 15: Applications of Aqueous Equilibria

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  1. Chapter 15: Applications of Aqueous Equilibria • Buffers • Common Ion Effect • Henderson-Hasselbalch Equation • Buffer Capacity • Acid-Base Titrations & Titration Curves • Strong Acid-Strong Base Titrations • Weak Acid-Strong Base Titrations • Strong Acid-Weak Base Titrations • Acid-Base Indicators • Solubility Equilibria • Calculating Solubility • Common Ion Effect • Selective Precipitation • Qualitative Analysis • Complex Ion Equilibria

  2. Example 1 Find the pH of a solution that is 0.20M KCNO and 0.10 M HCNO. Ka for HCNO is 3.5 x 10-4. Answer: pH = 3.76

  3. Example 2 The ratio of [HCO3–] to [H2CO3] in human blood is 20:1. Find the pH of an aqueous solution with this composition. Answer: pH = 7.67

  4. Example 3 Find the pH of a buffer that is made of 0.20M KCNO and 0.10 M HCNO. (Note: This is the same as Example 1, but here we will double check our answer using the Henderson-Hasselbalch equation.) Ka for HCNO is 3.5 x 10-4. Answer: pH = 3.76

  5. Example 4 What pH change would result from the addition of 5.0 mL of 0.10M HCl to 50.0 mL of a buffer containing 0.10 M NH3 and 0.10 M NH4Cl? How much would the pH of 50.0 mL of pure water change if the same amount of acid was added to it? Answers: DpH=-.09; DpH=-4.96

  6. Example 5 How many grams of Na2CO3 should be added to 1.5 L of 0.20 M NaHCO3 to make a buffer of pH = 10.00? Ka2 of H2CO3 = 5.6 x 10-11 Answer: 18 g of Na2CO3

  7. Example 6 What is the pH of a buffer made by adding 5.0 mL of 0.20 M NaOH to 25.0 mL of 0.10 M HC2H3O2? Answer: pH = 4.56

  8. Example 7 Calculate the pH change that occurs if 1.0 mL of 0.10 M HCl is added to 50.0 mL of a buffer containing: • 0.30 M pyruvic acid (HC3H3O3) and 0.30 M potassium pyruvate? Ka for pyruvic acid is 1.4 x 10-4. • 0.0030 M pyruvic acid and 0.0030 M potassium pyruvate? Answers: a. DpH=0.00 b. DpH=-0.70

  9. Example 8 A 75.0 mL sample of 0.200 M HBr is titrated with 0.100 M KOH to a phenolphthalein endpoint. How much KOH solution is needed to reach the equivalence point? What is the pH of the solution at the equivalence point?

  10. Strong Acid-Strong Base Titration Curves: Figures 15.1 & 15.2

  11. Acid-Base Indicators

  12. Example 9 When a 50.0 mL sample of 0.250 M nitrous acid is titrated with 0.100 M NaOH, what volume of NaOH is needed to reach the equivalence point? What is the pH at the equivalence point? What is the pH at the halfway point? Ka for nitrous acid is 4.0 x 10-4.

  13. Weak Acid-Strong Base Titration Curves

  14. Example 10 Find the volume of 0.100 M HCl needed to reach the equivalence point in the titration of 25.0 mL of 0.100 M NH3. Also find the pH of the solution in the flask: • Initially • After 10.0 mL of HCl have been added • At the halfway point • At the equivalence point • After 35.0 mL of HCl have been added Kb for NH3 is 1.8 x 10-5

  15. Example 11 A 100.0 mL sample of a weak, monoprotic acid with a concentration of 0.200 M is titrated with 0.100 M NaOH. After 10.0 mL of NaOH have been added, the pH is 5.79. What is Ka for this acid?

  16. Example 12 Silver bromide has a solubility of 0.133 mg per 1.00 L of water. Find Ksp for silver bromide.

  17. Example 13 Mercury (I) chloride has a Ksp of 1.3 x 10-18. Find its solubility in units of mole/L and g/L.

  18. Example 14 Which of the following ionic compounds is more soluble in water? (i.e. Which will dissolve more moles per liter?) • CaSO4 or CaCO3 • CaSO4 or Ca(OH)2 Ksp for CaSO4 = 6.1 x 10-5 Ksp for CaCO3 = 8.7 x 10-9 Ksp for Ca(OH)2 = 1.3 x 10-6

  19. Example 15 Calcium oxalate has a solubility of 6.1x10-3 g/L in water. Find its solubility in 0.20 M CaCl2. Ksp for CaC2O4 = 2.3x10-9

  20. Example 16 One type of kidney stones is made of calcium phospate. If [Ca2+] in urine is 0.080 g/L, what is the minimum molarity of phosphate that will cause kidney stones to form? Ksp for calcium phosphate = 1.3x10-32

  21. Example 17 A 65.0 mL sample of 0.010 M Pb(NO3)2 was added to a beaker containing 40.0 mL of 0.035 M KCl. Will a precipitate form?

  22. Example 18 What percentage of Ca2+ ions remain in solution after CaCO3 precipitates when 25.0 mL of 0.10 M CaCl2 is added to 25.0 mL of 0.10 M Na2CO3? Ksp for CaCO3 is 8.7 x 10-9.

  23. Example 19 When 1.0 M AgNO3 is slowly added to a solution containing 0.015M Cl- and 0.015M CrO42-, what percent of Cl- remains in solution when the Ag2CrO4 begins to precipitate? (i.e. What is the maximum separation of Cl- from CrO42- that can be achieved?)

  24. Example 20 A solution contains 0.10 M Cd2+ and 0.10 M Ni2+. What concentration of S2- will precipitate a maximum amount of one cation without precipitating the other? Ksp NiS = 3.0 x 10-21 Ksp CdS = 1.0 x 10-28

  25. Qualitative Analysis Scheme

  26. Example 21 How much Zn2+ ion remains in solution in a mixture that is 0.010 M Zn(NO3)2 and 0.10 M NH3? Kf Zn(NH3)42+ = 2.9 x 109

  27. Example 22 Calculate the solubility of AgI in: • 0.10 M KCN [Kf Ag(CN)2- = 5.6 x 1018] • water Ksp for AgI = 1.5 x 10-16

  28. Example 23 Will nickel (II) hydroxide precipitate from in a solution that is 0.0020 M NiSO4, 0.010 M NaOH, and 0.10 M NH3? Kf for Ni(NH3)62+ = 5.6 x 108 Ksp for Ni(OH)2 = 2.0 x 10-15

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