Chapter 17 Additional Aspects of Aqueous Equilibria

1. Chapter 17Additional Aspects of Aqueous Equilibria Chapter 17

2. 17.1: The Common Ion Effect • The solubility of a partially soluble acid is decreased when a common ion is added • HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2-(aq) • Consider the addition of C2H3O2- • This is a common ion • From a salt such as NaC2H3O2 • Therefore, [C2H3O2-] increases and the system is no longer at equilibrium • So, [H+] must decrease (shift left…LeChâtelier!) Chapter 17

3. 17.2: Buffered Solutions • Composition and Action of Buffered Solutions • A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X-): • The Ka expression is Chapter 17

4. A buffer resists a change in pH when a small amount of OH- or H+ is added • When OH- is added to the buffer, the OH- reacts with HX to produce X- and water • The [HX]/[X-] ratio remains more or less constant, so the pH is not significantly changed • When H+ is added to the buffer, X- is consumed to produce HX • the pH does not change significantly Chapter 17

5. Text, P. 665

6. Buffer Capacity and pH • Buffer capacity is the amount of acid or base neutralized by the buffer before there is a significant change in pH • It depends on the composition of the buffer • The greater the amounts of conjugate acid-base pair (molar concentration), the greater the buffer capacity • The pH of the buffer depends on Ka Chapter 17

7. Recall: • If Ka is small (the equilibrium concentration of the undissociated acid is close to the initial concentration), then • . the Henderson-Hasselbalch Equation! Chapter 17

8. Addition of Strong Acids or Bases to Buffers • The amount of strong acid or base added results in a neutralization reaction: • X- + H3O+ HX + H2O • HX + OH-  X- + H2O Text, P. 668 Chapter 17

9. Problems 3, 5, 9, 15, 17, 19 Chapter 17

10. 17.3: Acid-Base Titrations • Strong Acid-Strong Base Titrations • A plot of pH versus volume of acid (or base) added is called a titration curve • Consider adding a strong base (NaOH) to a solution of a strong acid (HCl): Chapter 17

11. Text, P. 672 Appropriate indicator: dramatic color change in the desired range pH is determined by ? pH is determined by ? pH is determined by ? pH is determined by ?

12. The equivalence point in a titration is the point at which the acid and base are present in stoichiometric quantities • The end point in a titration is the observed point • The difference between equivalence point and end point is called the titration error Chapter 17

13. Strong Base-Strong Acid Titrations • Add HCl to NaOH Text, P. 674

14. Weak Acid-Strong Base Titrations • Consider the titration of acetic acid, HC2H3O2 and NaOH • Before any base is added, the solution contains only weak acid • As strong base is added, the strong base consumes a stoichiometric quantity of weak acid: • HC2H3O2(aq) + NaOH(aq)  C2H3O2-(aq) + H2O(l) Chapter 17

15. Text, P. 674 pH is determined by pH is determined by ? • There is an excess of acid before the equivalence point so there is a mixture of weak acid and its conjugate base • The pH is given by the buffer calculation • First the amount of C2H3O2- generated is calculated, as well as the amount of HC2H3O2consumed (Stoichiometry) • Then the pH is calculated using equilibrium conditions (H-H) Chapter 17

16. Text, P. 675 Chapter 17

17. pH is determined by ? pH is determined by ? Note that pH is above 7 …the acetate ion is a weak base Chapter 17

18. Weak Acid/Strong Base Curve Strong Acid/Strong Base Curve Compare pH values at eq. points Compare pH change near eq. points Compare initial pH values Chapter 17

19. Text, P. 676 The influence of acid strength on the shape of the curve for the titration with NaOH Chapter 17

20. Text, P. 677 The titration of a weak base with a strong acid

21. Titrations of Polyprotic Acids • In polyprotic acids, each ionizable proton dissociates in steps • Therefore, in a titration there are n equivalence points corresponding to each ionizable proton • In the titration of H3PO3 with NaOH, • The first proton dissociates to form H2PO3- • Then the second proton dissociates to form HPO32- Chapter 17

22. Text, P. 677

23. Problems 25, 27, 29, 31, 33 Chapter 17

24. 17.4: Solubility Equilibria • The Solubility-Product Constant, Ksp • Consider equilibria that are heterogeneous • Some common applications: • Tooth enamel and soda, salts and kidney stones, stalactites and stalagmites • Example: • for which • Ksp is the solubility product constant Chapter 17

25. In general: the solubility product is the molar concentration of ions raised to their stoichiometric powers • Solubility is the amount (grams) of substance that dissolves to form a saturated solution • Affected by • pH • concentrations of other ions in solution • Molar solubility is the number of moles of solute dissolving to form a liter of saturated solution Chapter 17

26. Solubility and Ksp • To convert solubility to Ksp: • Solubility needs to be converted into molar solubility (via molar mass) • Molar solubility is converted into the molar concentration of ions at equilibrium (equilibrium calculation) • Ksp is the product of equilibrium concentration of ions Text, P. 697 Chapter 17

27. Sample Problems # 37 & 39 Chapter 17

28. 17.5: Factors that Affect Solubility • The Common Ion Effect • Solubility is decreased when a common ion is added • Le Châtelier’s principle: • as F- is added (from NaF), the equilibrium shifts away from the increase • CaF2(s) is formed and precipitation occurs Chapter 17

29. Solubility and pH • If the F- is removed, then the equilibrium shifts right and CaF2 dissolves • F- can be removed by adding a strong acid: • As pH decreases, [H+] increases and solubility increases • The effect of pH on solubility is dramatic • The more basic the anion, the more solubility is influenced by pH Chapter 17

30. Formation of Complex Ions • The formation of Ag(NH3)2+: • The Ag(NH3)2+ is called a complex ion • NH3 (the attached Lewis base) is called a ligand • Lewis bases share their nonbonded electron pairs with vacant orbitals on the metal atom • The equilibrium constant for the reaction is called the formation constant, Kf: Chapter 17

31. Text, P. 687 Chapter 17

32. Amphoterism • Amphoteric oxides will dissolve in either a strong acid or a strong base • Examples: hydroxides and oxides of Al3+, Cr3+, Zn2+, and Sn2+ • The hydroxides generally form complex ions with four hydroxide ligands attached to the metal: Chapter 17

33. Hydrated metal ions act as weak acids • Thus, the amphoterism is interrupted: Chapter 17

34. Problems 41, 43, 49 Chapter 17

35. 17.6: Precipitation and Separation of Ions • At any instant in time, Q = [Ba2+][SO42-] • If Q > Ksp, precipitation occurs until Q = Ksp • If Q = Ksp, equilibrium exists • If Q < Ksp, solid dissolves until Q = Ksp • Based on solubilities, ions can be selectively removed from solutions Chapter 17

36. Selective Precipitation of Ions • Ions can be separated from each other based on their salt solubilities • Example: if HCl is added to a solution containing Ag+ and Cu2+ • the silver precipitates (Ksp for AgCl is 1.8  10-10) while the Cu2+ remains in solution • Removal of one metal ion from a solution is called selective precipitation Chapter 17

37. Problems 51, 53, 55 Chapter 17

38. 17.7: Qualitative Analysis for Metallic Elements • Qualitative analysis is designed to detect the presence of metal ions • Quantitative analysis is designed to determine how much metal ion is present • See Text, P. 692-695