Chapter 6 Problems

1. Chapter 6 Problems • 6-29, 6-31, 6-39, 6.41, 6-42, 6-48,

2. 6-29 • Distinguish between Lewis Acids/Bases & Bronsted-Lowry acids and bases. Give an example.

3. 6-31. • Why is the pH of water usually < 7? How can you prevent this from happening?

4. 6-39. • The equilibrium constant for autoprotolysis of water is 1.0 x 10-14 at 25oC. What is the value of K for 4 H2O  4H+ + 4OH- • K= [H+]4 [OH]4 • K = (1x10-7)4(1x10-7)4 • K = 1 x 10-56

5. 6-41 • Use Le Chatelier’s principle and Kw in Table 6-1 to decide whether the autoprotolysis of water is exothermic or endothermic at • 25oC • 100oC • 300oC

6. 6-42 • Make a list of strong acid and strong bases. Memorize this list.

7. 6-48 • Which is the stronger acid? Dichloracetic acid Chloroacetic acid Ka = 8 x 10-2 Ka = 1.36 x 10-3 Stronger Base? Hydrazine Urea Kb = 1.1 x 10-6 Kb= 1.5 x 10-14

8. Chapter 8 Activity

9. HomeworkChapter 8 - Activity • 8.2, 8.3, 8.6, 8.9, 8.10, 8.12

10. 8-1 Effect of Ionic Strength on Solubility of Salts • Consider a saturated solution of Hg2(IO3)2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg2(IO3)2(s)D Hg22+ + 2IO3- Ksp=1.3x10-18 some - - -x +x +2x some-x +x +2x I C E A seemingly strange effect is observed when a salt such as KNO3 is added. As more KNO3 is added to the solution, more solid dissolves until [Hg22+] increases to 1.0 x 10-6 M. Why?

11. Increased solubility • Why? • LeChatelier’s Principle? • NO – not a product nor reactant • Complex Ion? • No • Hg22+ and IO3- do not form complexes with K+ or NO3-. • How else?

12. - 2+ The Explanation • Consider Hg22+ and theIO3- Electrostatic attraction

13. The Explanation • Consider Hg22+ and theIO3- Electrostatic attraction - 2+ Hg2(IO3)2(s) The Precipitate!!

14. The Explanation • Consider Hg22+ and theIO3- Electrostatic attraction - 2+ NO3- K+ Add KNO3

15. The Explanation • Consider Hg22+ and theIO3- NO3- K+ NO3- K+ NO3- K+ NO3- K+ K+ - K+ 2+ K+ NO3- K+ NO3- NO3- K+ NO3- K+ NO3- K+ K+ NO3- Add KNO3

16. The Explanation • Consider Hg22+ and theIO3- NO3- K+ NO3- K+ NO3- K+ NO3- K+ K+ - K+ 2+ K+ NO3- K+ NO3- NO3- NO3- K+ NO3- K+ K+ K+ NO3- Hg22+ and IO3- can’t get CLOSE ENOUGH to form Crystal lattice Or at least it is a lot “Harder” to form crystal lattice

18. The potassium hydrogen tartrate example

19. Alright, what do we mean by Ionic strength? • Consider Hg22+ and theIO3- NO3- K+ NO3- K+ NO3- K+ NO3- K+ K+ - K+ 2+ K+ NO3- K+ NO3- NO3- K+ NO3- K+ NO3- K+ K+ NO3- Low Ionic Strength High Ionic Strength Higher Ionic Strength Add KNO3

20. Alright, what do we mean by Ionic strength? • Ionic strength is a measure of the total concentration of ions in solution. • Ionic strength is dependent on the number of ions in solution and their charge. • Not dependent on the chemical nature of the ions Ionic strength (m) = ½ (c1z12+ c2z22 + …) Or Ionic strength (m) = ½ S cizi2

21. Examples • Calculate the ionic strength of (a) 0.1 M solution of KNO3and (b) a 0.1 M solution of Na2SO4 (c) a mixture containing 0.1 M KNO3 and 0.1 M Na2SO4. (m) = ½ (c1z12+ c2z22 + …)

22. Alright, that’s great but how does it affect the equilibrium constant? A + B  C + D • Activity = Ac = [C]gc • AND

23. Relationship between activity coefficient and ionic strength Debye-Huckel Equation m = ionic strength of solution g = activity coefficient Z = Charge on the species x a = effective diameter of ion (nm) 2 comments • What happens to g when m approaches zero? • Most singly charged ions have an effective radius of about 0.3 nm We generally don’t need to calculate g – values are tabulated

24. Concept Test List at least three properties of activity coefficients • Dimensionless • Depends on size of the ions (ex. Hg22+ and IO3-) • Depends on the Ionic Strength of the Solution (K+ & NO3-) • Depends on the charge of the ions (ex. Hg22+ and IO3-) • In dilute solutions, where ionic strength is minimal, the activity coefficient -> 1, and has little effect on equilibrium constant

25. Activity coefficients are related to the hydrated radius of atoms in molecules

26. Relationship between m and g

29. Back to our original problem • Consider a saturated solution of Hg2(IO3)2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg2(IO3)2(s)D Hg22+ + 2IO3- Ksp=1.3x10-18 1 1 At low ionic strengths g -> 1

30. Back to our original problem • Consider a saturated solution of Hg2(IO3)2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg2(IO3)2(s)D Hg22+ + 2IO3- Ksp=1.3x10-18 In 0.1 M KNO3 - how much Hg22+ will be dissolved?

33. Back to our original problem • Consider a saturated solution of Hg2(IO3)2 in 0.1 M KNO3. Calculate the concentration of mercurous ions. Hg2(IO3)2(s)D Hg22+ + 2IO3- Ksp=1.3x10-18

34. Consider a saturated solution of Hg2(IO3)2 Calculate the concentration of mercurous ions in: • In Pure water = 6.9 x 10-7 • in 0.1 M KNO3 = 1.15 x 10-6

35. pH revisited

36. Definition of pH • pH = -log AH • or • pH = - log [H+] gH

37. pH of pure water H2O (l)  H+ (aq) + OH- (aq)Kw =1.0 x 10-14 Kw = AH AOH Kw = [H+] gH [OH-] gOH

39. pH of pure water H2O (l)  H+ (aq) + OH- (aq)Kw =1.0 x 10-14 Kw = AH AOH Kw = [H+] gH [OH-] gOH Kw = x2 x = 1.0 x 10-7 M I C E - - +x +x +x +x 1 1

40. pH of pure water x = 1.0 x 10-7 M Therefore [H+] = 1.0 x 10-7 M pH = -log AH = -log [H+] gH = - log [1.0 x 10-7] = 7.00 1

41. pH of pure water containing salt • Calculate the pH of pure water containing 0.10 M KCl at 25oC.