1. Acid-Base Equilibria and Solubility Equilibria Chapter 16 16.1-16.9

2. Acid-Base Equilibria and Solubility Equilibria We will continue to study acid-base reactions We will continue to study acid-base reactions Buffers Buffers Titrations Titrations We will also look at properties of slightly soluble compunds and their ions in solution We will also look at properties of slightly soluble compunds and their ions in solution

3. Homogeneous vs. Heterogeneous Homogeneous Solution Equilibria- a solution that has the same composition throughout after equilibrium has been reached. Homogeneous Solution Equilibria- a solution that has the same composition throughout after equilibrium has been reached. Heterogeneous Solution Equilibria- a solution that after equilibrium has been reached, results in components in more than one phase. Heterogeneous Solution Equilibria- a solution that after equilibrium has been reached, results in components in more than one phase.

4. The Common Ion Effect Acid-Base Solutions Acid-Base Solutions Common Ion Common Ion The Common Ion Effect- the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The Common Ion Effect- the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. pH pH

5. The Common Ion Effect The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider mixture of CH 3 COONa (strong electrolyte) and CH 3 COOH (weak acid). CH 3 COONa (s) Na + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) common ion

6. Henderson-Hasselbach Equation Relationship between pK a and K a. Relationship between pK a and K a. pK a = -log K a Henderson-Hasselbalch equation pH = pK a + log [conjugate base] [acid]

7. pH Calculations What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? Mixture of weak acid and conjugate base! HCOOH (aq) H + (aq) + HCOO - (aq) Initial (M) Change (M) Equilibrium (M) 0.30 -x-x 0.30 - x 0.000.52 +x+x+x+x x0.52 + x HCOOH pK a = 3.77

8. pH Calculations pH = pK a + log [HCOO - ] [HCOOH] pH = 3.77 + log [0.52] [0.30] Common ion effect 0.30 – x  0.30 0.52 + x  0.52 = 4.01

9. Buffer Solutions A buffer solution is a solution of: 1.A weak acid or a weak base and 2.The salt of the weak acid or weak base Both must be present! A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base.

10. Buffer Solutions IV solutions (~7.4) IV solutions (~7.4) Blood (~7.4) Blood (~7.4) Gastric Juice (~1.5) Gastric Juice (~1.5)

11. Buffer Solutions Consider an equal molar mixture of CH 3 COOH and CH 3 COONa Add strong acid H + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) Add strong base OH - (aq) + CH 3 COOH (aq) CH 3 COO - (aq) + H 2 O (l)

12. Buffer Solutions

17. Effectiveness of Buffer Solutions

18. Preparing a Buffer Solution with a Specific pH Work Backwards Work Backwards Choose a weak acid whose pKa is close to the desired pH Choose a weak acid whose pKa is close to the desired pH Substitute pKa and pH values into the Henderson-Hasselbach Equation Substitute pKa and pH values into the Henderson-Hasselbach Equation This will give a ratio of [conjugate Base]/[Acid] This will give a ratio of [conjugate Base]/[Acid] Convert ratio to molar quantities Convert ratio to molar quantities

21. Acid-Base Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Types of Titrations Those involving a strong acid and a strong base Those involving a weak acid and a strong base Those involving a strong acid and a weak base

22. Acid-Base Titrations Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL The indicator changes color (pink) Equivalence point – the point at which the reaction is complete

23. Acid-Base Titrations

24. Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) Strong Acid-Strong Base Titrations OH - (aq) + H + (aq) H 2 O (l)

25. Strong Acid-Base Titrations

26. Calculate the pH after the addition of 10.0mL of 0.100M NaOH to 25.0mL of 0.100M HCl. Calculate the pH after the addition of 10.0mL of 0.100M NaOH to 25.0mL of 0.100M HCl. # moles NaOH 10.0mL x (0.100mol NaOH/ 1L NaOH) x (1L/ 1000mL) = 1.00 x 10 -3 mol #moles HCl 25.0mL x (0.100mol HCl/ 1 L HCl) x (1L/ 1000mL) = 2.50 x 10 -3 mol Amount of HCl left after partial neutralization 2.50 x 10 -3 mol - 1.00 x 10 -3 mol = 1.5 x x 10 -3 mol

27. Strong Acid-Base Titrations Thus, [H + ] = 1.5 x x 10 -3 mol/ 0.035L [H + ] = 0.0429 M pH = -log [H + ] pH = -log 0.0429 pH = 1.37

28. Strong Acid-Base Titrations Calculate pH after the addition of 35.0mL of 0.100M NaOH to 25.0mL of 0.100M HCl. Calculate pH after the addition of 35.0mL of 0.100M NaOH to 25.0mL of 0.100M HCl. # moles NaOH 0.100 mol NaOH / 0.035 L NaOH = 3.50 x 10 -3 mol #moles HCl 0.100 mol HCl / 0.025 L HCl = 2.50 x 10 -3 mol Amount of NaOH left after full HCl neutralization 3.50 x 10 -3 mol – 2.50 x 10 -3 mol = 1.0 x x 10 -3 mol

29. Strong Acid-Base Titrations Thus, [NaOH] = 1.0 x x 10 -3 mol/ 0.06L [NaOH] = 0.0167 M [OH - ] = 0.0167 M pOH = -log [H + ] pOH = -log 0.0167 pOH = 1.78 pH = 14.0-pOH pH = 14.0-1.78 pH = 12.22

30. Weak Acid-Strong Base Titrations CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq) CH 3 COO - (aq) + H 2 O (l) At equivalence point (pH > 7): CH 3 COO - (aq) + H 2 O (l) OH - (aq) + CH 3 COOH (aq)

31. Weak Acid-Strong Base Titrations

35. Strong Acid-Weak Base Titrations HCl (aq) + NH 3 (aq) NH 4 Cl (aq) H + (aq) + NH 3 (aq) NH 4 Cl (aq) At equivalence point (pH < 7): NH 4 + (aq) + H 2 O (l) NH 3 (aq) + H + (aq)

36. Strong Acid-Weak Base Titrations

39. Acid-Base Indicators Equivalence point occurs when OH- = H+ originally present. Equivalence point occurs when OH- = H+ originally present. Indicators Indicators End Point- Occurs when indicator changes color End Point- Occurs when indicator changes color End point ~ Equivalence point End point ~ Equivalence point

40. Acid-Base Indicators

43. pH

44. Solubility Equilibria Reactions that produce precipitates Reactions that produce precipitates Importance Importance Tooth Enamel + Acid = tooth decay Tooth Enamel + Acid = tooth decay Barium Sulfate = used in x-rays Barium Sulfate = used in x-rays Fudge!!! Fudge!!!

45. Solubility Equilibria Solubility Product Constant (K sp )- the product of the molar concentrations of the constituent ions, each raised to the power of its stoichiometric coefficient in the equilibrium equation. Solubility Product Constant (K sp )- the product of the molar concentrations of the constituent ions, each raised to the power of its stoichiometric coefficient in the equilibrium equation. AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ]

46. Solubility Products What are the correct solubility products of the following equations? What are the correct solubility products of the following equations? MgF 2 (s) Mg 2+ (aq) + 2F - (aq) K sp = [Mg 2+ ][F - ] 2 Ag 2 CO 3 (s) 2Ag + (aq) + CO 3 2 - (aq) K sp = [Ag + ] 2 [CO 3 2 - ] Ca 3 (PO 4 ) 2 (s) 3Ca 2+ (aq) + 2PO 4 3 - (aq) K sp = [Ca 2+ ] 3 [PO 3 3 - ] 2

47. Solubility Product Constants

48. Solubility Constants What does a large K sp mean? What does a large K sp mean? What does a small value mean? What does a small value mean?

49. Molar Solubility and Solubility Molar solubility (mol/L)- is the number of moles of solute dissolved in 1 L of a saturated solution. Solubility (g/L)- is the number of grams of solute dissolved in 1 L of a saturated solution.

50. Molar Solubility and Solubility

52. The Ksp of silver bromide is 7.7 x 10 -13. Calculate the molar solubility. The Ksp of silver bromide is 7.7 x 10 -13. Calculate the molar solubility. AgBr (s) ↔ Ag + (aq) + Br - (aq) Initial (M) 0 0 Change (M) -s +s+s Equilibrium (M) s s K sp = [Ag + ][Br - ] 7.7 x 10 -13 = [Ag + ][Br - ] 7.7 x 10 -13 = s 2 S= 8.8 x 10 -7 M

54. K sp vs. Q Dissolution of an ionic solid in aqueous solution: Q < K sp Q = K sp Q > K sp Unsaturated solution Saturated solution Supersaturated solution No precipitate Precipitate will form

55. Predicting Precipitation Reactions Calculate Q for the reaction Calculate Q for the reaction Is Q larger, smaller or equal to K sp ? Is Q larger, smaller or equal to K sp ?

58. Separation of Ions by Fractional Precipitation Removal of ions from solution Removal of ions from solution Useful in preparation of prescription medications Useful in preparation of prescription medications Ions can be removed by filtration Ions can be removed by filtration

59. Fractional Precipitation Ions + proper reagent Ions + proper reagent Smallest → Largest K sp Smallest → Largest K sp If AgNo3 is added to a solution containing Cl-, Br- and I- ions, which compound will precipitate out first? If AgNo3 is added to a solution containing Cl-, Br- and I- ions, which compound will precipitate out first? CompoundKsp AgCl 1.6 x 10 -10 AgBr 1.7 x 10 -13 AgI 8.3 x 10 -17

62. The Common Ion Effect and Solubility What is the molar solubility of AgBr in (a) pure water and (b) 0.0010 M NaBr? AgBr (s) Ag + (aq) + Br - (aq) K sp = 7.7 x 10 -13 s 2 = K sp s = 8.8 x 10 -7

63. The Common Ion Effect and Solubility AgBr (s) Ag + (aq) + Br - (aq) NaBr (s) Na + (aq) + Br - (aq) [Br - ] = 0.0010 M [Ag + ] = s [Br - ] = 0.0010 + s  0.0010 K sp = 0.0010 x s s = 7.7 x 10 -10

64. pH and Solubility The solubility of many substances depends on the pH of the solution. The solubility of many substances depends on the pH of the solution. ↑ pH ↑ pH ↓ pH ↓ pH Insoluble bases dissolve in acidic solutions Insoluble bases dissolve in acidic solutions Insoluble acids dissolve in basic solutions Insoluble acids dissolve in basic solutions