1. Ch 8: Introducing Acids and Bases

2. pH of precipitation in the United States 2001, and in Europe as reported in 2002.

3. What are Acids and Bases? acid = substance that increases the concentration of H 3 O + base = decreases the concentration of H 3 O + (by increasing the amount of OH - )

4. Bronsted-Lowry Acid-Base Theory acid = proton donor (H + ) base = proton acceptor HCl + H 2 O → HCl + NH 3 →

5. Conjugate Acid-Base Pairs

6. Relation Between [H + ], [OH - ], and pH H 2 O + H 2 OH 3 O + + OH - H 2 OH + + OH - K w = [H + ][OH - ] = 1.01 x 10 -14 at 25 o C equivalent

8. Example, p. 169: Concentration of H + and OH - in Pure Water at 25 o C Calculate the concentrations of H + and OH - in pure water at 25 o C.

9. As the concentration of H + increases, OH - must decrease and vica-versa Example, p. 169: Finding [OH-] when H+ is Known. What is the concentration of OH - if [H + ] = 1.0 x 10 -3 M at 25 o C?

10. pH - a measure of the aciity of a solution ("puissance d'hydrogen") pH = -log [H + ](approximately!) [H+] = 10 -3 M [H+] = 10.0 M [H+] = 10 -10 M

11. Strengths of Acids and Bases strong = complete (100%) dissociation MEMORIZE these strong acids and bases - all other acids and bases are weak

12. weak = incomplete dissociation HA H + + A- OR HA + H 2 O H 3 O + + A - B + H 2 OBH + + OH -

13. Classes of Weak Acids and Bases carboxylic acids = weak acids amines = weak bases RNH 2 primary R 2 NHsecondary R 3 Ntertiary : : : polyprotic acids and bases H 2 CO 3 CO 3 2- H 3 PO 4 PO 4 3- Ca(OH) 2

14. Relation Between K a and K b HA + H 2 O H 3 O + + A - salt = conjugate base undergoes hydrolysis A - + H 2 O HA + OH -

15. Example, p. 174: K a for acetic acid is 1.75 x 10 -5. Find K b for the acetate ion.

17. pH of solutions of strong acids and bases Case I: concentration of acid or base >> 10 -7 HA → H + + A - BOH → B + + OH - H 2 O H + + OH - [H + ] = [OH - ] = 1.0 x 10 -7 strong acids and bases completely dissociate

18. pH of a strong acid: Example p. 175 Find the pH of 4.2 x 10 -3 M HClO 4

19. Case II: concentration of acid or base  10 -7 Now the contribution of H + from water must be included -

21. pH of a strong base at a low concentration: “trick question” top of p. 176 Find the pH of 4.2 x 10 -9 M KOH

22. pH of solutions of weak acids and bases (Sec 8-6, 8-7) – the “ICE” table Calculate the pH of a 0.020 M benzoic acid solution. K a = 6.28 x 10 -5 I: Exact solution using quadratic equation

23. II. Approximate solution

24. Weak Base Equilibrium, Example p. 183 Find the pH of a 0.0372 M solution of the commonly encountered (?) weak base cocaine. K b = 2.6 x 10 -6

25. Ch 9: Buffers buffer = resists changes in pH; solution of a weak acid or base and their salts

26. Henderson-Hasselbach Equation derivation and assumptions:

27. Example, p.191: Using the H-H Equation Sodium hypochlorite (NaOCl) was dissolved in a solution buffered to pH = 6.20. Find the ratio [OCl - ]/[HOCl]

28. Example, p. 192: A Buffer Solution Find the pH of a solution prepared by dissolving 12.43 g of TRIS (FM = 121.136) plus 4.67 g TRIS hydrochloride (FM = 157.597) in 1.00 L of water.

29. A Buffer in Action Weak Acid & Salt e.g. CH 3 COO - / CH 3 COOH Weak Base & Salt e.g. NH 4 + / NH 3 If add H + If add OH -

30. How to Prepare a Buffer Solution 1. Consult a table of pKa's and pick the weak acid or base closest to the pH you need. 2. Solve for the ratio mol salt/(mol acid/base) 3. Choose a reasonable value for either mol salt or mol acid/base and solve for the other 4. After preparing the buffer, adjust the pH to the desired value (you never get exactly what you calculate because of the assumptions made in deriving the H-H equation

31. Example: buffer with pH = 4.8 acid/basepK a pK b acetic acid4.757 benzoic acid4.202 ammonia4.74 dimethylamine3.13

32. Buffer Capacity: How well a solution resists changes in pH when an acid or base is added: when the pH = pKa!

33. Example, p.198 HA = H + + A - mol A - = 0.0383, mol HA = 0.9617