1. Acid-Base Equilibria and Solubility Equilibria Chapter 16

2. Objectives Define key terms and concepts Explain how a buffer works and how to prepare a buffer system for a specific pH. Calculate the pH of a buffer solution. Calculate the equivalence point for acid-base titrations. Calculate the solubility product for a compound. Calculate concentration and solubility using the common ion and complex ion effects. Develop and utilize a flow chart for use in qualitative analysis.

3. Conjugate Acids and Bases

4. Buffers Buffer solutions resist changes in pH when small amounts of acids or bases are added Buffers are made of a weak acid or base and it’s salt (acid-base conjugate pair) Buffering Capacity

5. Buffers

6. Instructions for making up a buffer say to mix 60mL of 0.1M NH 3 with 40mL of 0.1M NH 4 Cl. What is the pH of the buffer?

7. Calculate the pH of a buffer solution using ammonia and ammonium to which 3mL of 0.10M HNO 3 is added.

8. Instructions for making up a buffer say to mix 100mL of 0.25M HC 2 H 3 O 2 with 75mL of 0.3M NaC 2 H 3 O 2. What is the pH of the buffer?

9. Calculate the pH of a buffer solution using acetic acid/sodium acetate to which 9.5mL of 0.10M HCl is added.

10. Buffers Henderson-Hasselbalch Equation Relates pH of a buffer for different concentrations of the conjugate acid-base If the molar concentration of the acid ≈ base, then the pH ≈ pK a pH = pK a + log [base] [acid] pK a = -log K a

12. What must the ratio of the concentration of the HCO 3 - and H 2 CO 3 buffer system found in blood be for the pH to be 7.4?

13. What acid-base conjugate pair could be used to make a buffer with a pH of 3.7? How would you prepare the buffer solution? H 3 PO 4 ↔ H 2 PO 4 - + H 3 O + K a = 7.5x10 -3 H 2 S ↔ HS - + H 3 O + K a = 1x10 -7 HCO 2 H ↔ HCO 2 - + H 3 O + K a = 1.8x10 -4

14. Acid-Base Titrations Titration Addition of a solution of a known concentration (standard solution) to a solution of unknown concentration until the reaction is complete (reaches its equivalence point)

16. Acid-Base Titrations Strong Acid-Base Titrations Weak Acid-Base Titrations Strong Acid-Weak Base Titration Weak Acid-Strong Base

17. Calculate the pH of a solution in which 10mL of 0.1M NaOH is added to 25.0mL of 0.1M HCl.

18. Calculate the pH of a solution when 25mL of 0.1M nicotinic acid (HC 5 H 4 NO 2 ) is titrated by 20mL of 0.1M NaOH. The K a for nicotinic acid is 1.4x10 -5.

19. What is the pH of a solution when 35mL of 0.2M ammonia is titrated by28mL 0.12M HCl. The K b for ammonia is 1.8x10 -5.

20. What is the pH at the equivalence point when 25mL of 0.1M HF is titrated by 0.15M NaOH?

21. What are your questions?

22. Solubility Equilibria Solubility Product (K sp ) The product of the molar concentrations of the ions that make-up a compound, each raised to the power of it’s stoichiometric coefficient. The smaller the K sp value, the less soluble the compound is in water. BaCO 3 ↔ Ba 2+ + CO 3 2- K sp = [Ba 2+ ] [CO 3 2- ] BaF 2 ↔ Ba 2+ + 2F - K sp = [Ba 2+ ] [F - ] 2

23. Solubility Equilibria The saturation level of a solution can be determined using K sp and Q. If Q < K sp, the solution is unsaturated If Q > K sp, the solution is supersaturated If Q = K sp, the solution is saturated

24. Write the solubility product expression for the following salts: AgCl Hg 2 Cl 2 Pb 3 (AsO 4 ) 2

25. A liter of a solution saturated at 25°C with calcium oxalate, CaC 2 O 4, is evaporated to dryness, giving a 0.0061g residue of CaC 2 O 4. Calculate the solubility product constant for the salt at 25°C.

26. By experiment, it is found that 1.2x10 -3 mole of lead (II) iodide dissolves in 1L of aqueous solution at 25°C. What is the solubility product constant at this temperature?

27. Solubility Equilibria Molar Solubility The number of moles of solute in 1 liter of saturated solution Moles per liter Solubility The number of grams of solute in 1 liter of saturates solution Grams per liter These both refer to solutions at a specific temperature

29. Solubility Product and Molar Solubility

30. Calculate the solubility in g/L of lead (II) chromate in water. The K sp for PbCrO 4 is 1.8x10 -14.

31. Calculate the solubility in g/L for lead (II) arsenate when dissolved in water. The K sp for Pb 3 AsO 4 ) 2 is 4.0x10 -36.

32. Calculate the solubility product for potassium nitrate if the solubility of the compound in water is 133g/L at 0°C.

33. If 0.17g of calcium fluoride is soluble in 1L of solvent, what is the solubility product of calcium fluoride?

34. Precipitation Reactions If the ion product exceeds the K sp, a precipitate will form. Soluble Compounds Almost all salts of Na +, K +, and NH 4 + Salts of NO 3 -, ClO 3 -, ClO 4 -, C 2 H 3 O 2 - Exceptions Almost all salts of Cl -, Br -, I - Halides of Ag +, Hg 2+, Pb 2+ Compounds containing F - Fluorides of Mg 2+, Ca 2+, Sr 2+, Ba 2+, and Pb 2+ Salts of SO 4 2- Sulfates of Ca 2+, Sr 2+, Ba 2+, Pb 2+ Insoluble CompoundsExceptions Most salts of CO 3 2-, PO 4 3-, C 2 O 4 2-, CrO 4 2- Salts of NH 4 + and alkali metal cations Most metal sulfides, S 2- Most metal hydroxides and oxides

35. Common Ion Effect and Solubility A compound can be precipitated out of solution if another compound containing a common ion is added to the solution. The saturation level of a solution can be determined using K sp and Q (ion product). If Q < K sp, the solution is unsaturated If Q > K sp, the solution is supersaturated If Q = K sp, the solution is saturated

36. Common Ion Effect and Solubility

37. o Which of the following compounds should most affect the solubility of lead (II) sulfate in water to which it has been added? a) NaCl b) Na 2 SO 4 c) PbS o Which of the following compounds should most affect the solubility of silver chloride in water to which it has been added? a) NaCl b) Na 2 SO 4 c) PbS

38. The concentration of calcium ion in blood plasma is 0.0025M. If the concentration of oxalate ion is 1.0x10 -7 M, do you expect calcium oxalate to precipitate? K sp for calcium oxalate is 2.3x10 -9.

39. AgCl will be dissolved into a solution with is ALREADY 0.0100 M in chloride ion. What is the solubility of AgCl? The K sp =1.77 x 10¯ 10

40. What is the solubility of Ca(OH) 2 in 0.0860 M Ba(OH) 2 ? The K sp for Ca(OH) 2 is known to be 4.68 x 10¯ 6.

41. A.If 1.1 x 10 -4 g of Cr(OH) 3 is added to 120 L of water at 25 °C, will all of the solute dissolve? The of chromium (III) hydroxide is K sp = 6.7 x 10 -31. B.If 4.0 x 10 -4 g of NaOH is added to the solution described above, will a precipitate form? C.Calculate the molar solubility of Cr(OH) 3 in a solution buffered at pH = 11.00.

42. Complex Ion Equilibria and Solubility Complex Ion An ion containing a central metal cation bonded to one or more molecules or ion. Commonly formed by transition metals Formation Constant (K f ) The equilibrium constant for complex ion formation The larger the K f, the more stable the complex

43. Complex Ions

46. Cu 2+ + 4NH 3 2+  Cu(NH 3 ) 4 2+ (solutions turns a deep blue) Ni 2+ + 6NH 3  Ni(NH 3 ) 6 2+ (Solutions is blue)

47. Will silver chloride precipitate from a solution that is 0.01M AgNO 3 and 0.01M NaCl? The K sp for silver chloride is 8.3x10 -17

48. Will silver iodide precipitate from a solution that is 0.0045M AgNO 3 and 0.15M NaI? The K sp for silver iodide is 1.8x10 -10.

49. Sulfate ion in solution is often determined quantitatively by precipitating it as barium sulfate. The sulfite ion may have been formed from a sulfur compound. Analysis for the amount of sulfate ion then indicates the percentage of sulfur in the compound. Is a precipitate expected to form at equilibrium when 50mL of 0.0010M barium chloride is added to 50mL of 0.00010M sodium sulfate? The solubility product constant for barium sulfate is 1.1x10 -10.

50. What are your questions?