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 25 o C. What is the value of K for 4 H 2 O  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 K w in Table 6-1 to decide whether the autoprotolysis of water is exothermic or endothermic at 25 o C 100 o C 300 o C

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

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

8. Chapter 8 Activity

9. Homework Chapter 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 Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg 2 2+ + 2IO 3 - K sp =1.3x10 -18 seemingly strange effect A seemingly strange effect is observed when a salt such as KNO 3 is added. As more KNO 3 is added to the solution, more solid dissolves until [Hg 2 2+ ] increases to 1.0 x 10 -6 M. Why? ICE some- - -x+x+2x some-x+x+2x

11. Increased solubility Why? LeChatelier’s Principle? NO – not a product nor reactant Complex Ion? No Hg 2 2+ and IO 3 - do not form complexes with K + or NO 3 -. How else?

12. The Explanation Consider Hg 2 2+ and the IO 3 - - 2+ Electrostatic attraction

13. The Explanation Consider Hg 2 2+ and the IO 3 - 2+ Electrostatic attraction - The Precipitate!! Hg 2 (IO 3 ) 2 (s) The Precipitate!!

14. The Explanation Consider Hg 2 2+ and the IO 3 - - 2+ Electrostatic attraction Add KNO 3 NO 3 - K+

15. The Explanation Consider Hg 2 2+ and the IO 3 - - 2+ Add KNO 3 K+ NO 3 - K+

16. The Explanation Consider Hg 2 2+ and the IO 3 - 2+ NO 3 - Hg 2 2+ and IO 3 - can’t get CLOSE ENOUGH to form Crystal lattice Or at least it is a lot “Harder” to form crystal lattice - K+

18. The potassium hydrogen tartrate example

19. Alright, what do we mean by Ionic strength? Consider Hg 2 2+ and the IO 3 - - 2+ Add KNO 3 K+ NO 3 - K+ Low Ionic Strength Higher Ionic StrengthHigh Ionic Strength

20. Alright, what do we mean by Ionic strength? measure of the total concentration of ions in solution 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 (  ) = ½ (c 1 z 1 2 + c 2 z 2 2 + …) Or Ionic strength (m) = ½  c i z i 2

21. Examples Calculate the ionic strength of (a) 0.1 M solution of KNO 3 and (b) a 0.1 M solution of Na 2 SO 4 (c) a mixture containing 0.1 M KNO 3 and 0.1 M Na 2 SO 4. (  ) = ½ (c 1 z 1 2 + c 2 z 2 2 + …)

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

23. Relationship between activity coefficient and ionic strength Debye-Huckel Equation 2 comments  = ionic strength of solution  = activity coefficient Z = Charge on the species x  = effective diameter of ion (nm) (1)What happens to  when  approaches zero? (2)Most singly charged ions have an effective radius of about 0.3 nm We generally don’t need to calculate  – values are tabulated

24. Concept Test List at least three properties of activity coefficients Dimensionless Depends on size of the ions (ex. Hg 2 2+ and IO 3 - ) Depends on the Ionic Strength of the Solution (K + & NO 3 - ) Depends on the charge of the ions (ex. Hg 2 2+ and IO 3 - ) 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  and 

29. Back to our original problem Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg 2 2+ + 2IO 3 - K sp =1.3x10 -18 At low ionic strengths  -> 1 1 1

30. Back to our original problem Consider a saturated solution of Hg 2 (IO 3 ) 2 in ‘pure water’. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg 2 2+ + 2IO 3 - K sp =1.3x10 -18 In 0.1 M KNO 3 - how much Hg 2 2+ will be dissolved?

33. Back to our original problem 0.1 M KNO 3. Consider a saturated solution of Hg 2 (IO 3 ) 2 in 0.1 M KNO 3. Calculate the concentration of mercurous ions. Hg 2 (IO 3 ) 2(s)  Hg 2 2+ + 2IO 3 - K sp =1.3x10 -18

34. Consider a saturated solution of Hg 2 (IO 3 ) 2 Calculate the concentration of mercurous ions in: 6.9 x 10 -7 In Pure water = 6.9 x 10 -7 0.1 M KNO 3 = 1.1 5 x 10 -6 in 0.1 M KNO 3 = 1.1 5 x 10 -6

35. pH revisited

36. Definition of pH pH = -log A H or pH = - log [H + ]  H

37. pH of pure water H 2 O (l)  H + (aq) + OH - (aq) K w =1.0 x 10 -14 K w = A H A OH K w = [H + ]  H [OH - ]  OH

39. pH of pure water H 2 O (l)  H + (aq) + OH - (aq) K w =1.0 x 10 -14 K w = A H A OH K w = [H + ]  H [OH - ]  OH K w = x 2 x = 1.0 x 10 -7 M 1 1 -- ICEICE +x

40. pH of pure water x = 1.0 x 10 -7 M Therefore [H + ] = 1.0 x 10 -7 M pH = -log A H = -log [H + ]  H = - 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 25 o C.