1. Chapter 19 More about ACID-BASES

2. Self-Ionization of Water Two water molecules produce a hydronium ion & a hydroxide ion by the transfer of a proton. H 2 O (l) + H 2 O (l)  H 3 O + (aq) + OH - (aq) In pure water, every time you make one H 3 O + you get one OH - That is, [H 3 O + ] = [OH - ]

3. Hydronium Concentration Symbolized [H 3 O + ]

4. In Water at 25°C [H 3 O + ] = 1.0 x 10 -7 M And [OH - ] = 1.0 x 10 -7 M

5. IONIZATION CONSTANT OF WATER at 25°C K w = [H 3 O + ] [OH - ] K w = (1.0 x 10 -7 M) x (1.0 x 10 -7 M) = 1.0 x 10 -14 M 2

6. ION PRODUCT CONSTANT FOR WATER at 25°C K w = [H + ] [OH - ] K w = (1.0 x 10 -7 M) x (1.0 x 10 -7 M) = 1.0 x 10 -14 M 2

7. Neutral solutions Pure water is Neutral and [H 3 O + ] = [OH - ] In any solution that is neutral, [H 3 O + ] = [OH - ]

8. In any solution that is acidic, [H 3 O + ] > [OH - ] [H 3 O + ] > 1.0 x 10 -7 M

9. In any solution that is basic, [H 3 O + ] < [OH - ] [OH - ] > 1.0 x 10 -7 M [H 3 O + ] < 1.0 x 10 -7 M

10. STRONG ACIDS & BASES NEARLY COMPLETELY IONIZE OR DISSOCIATE IN AQUEOUS SOLUTION

11. STRONG ACID Solutions HClO 4 (aq) H 2 SO 4 (aq) HNO 3 (aq) HCl (aq) HBr (aq) HI (aq) MEMORIZE!!!

12. Strong Bases Group 1 Hydroxides NaOH KOH LiOH RbOH CsOH Group 2 Hydroxides Ca(OH) 2 Ba(OH) 2 Sr(OH) 2

13. NaOH  Na + (aq) + OH - (aq) Therefore, 1 mole of NaOH will yield 1 mole of OH - in an aqueous solution.

14. Ca(OH) 2  Ca 2+ (aq) + 2OH - (aq) Therefore, 1 mole of Ca(OH) 2 will yield 2 mole of OH - in an aqueous solution.

15. HCl  H + (aq) + Cl - (aq) Therefore, 1 mole of HCl will yield 1 mole of H + in an aqueous solution.

16. H 2 SO 4  2H + (aq) + SO 4 2- (aq) Therefore, 1 mole of H 2 SO 4 will yield 2 mole of H + in an aqueous solution.

17. For any aqueous solution at 25°C K w = [H 3 O + ] [OH - ] = 1.0 x 10 -14 M 2 Using this equation, if the concentration of one of the ions is known, then, the concentration of the other can be calculated.

18. Exercise Determine the hydronium and hydroxide ion concentrations in a solution that is 1 x 10 -4 M HCl. Answer: [H 3 O + ] = 1 x 10 -4 M [OH - ] = 1 x 10 -10 M

19. Exercise Determine the hydronium and hydroxide ion concentrations in a solution that is 1.0 x 10 -3 M HNO 3. Answer: [H 3 O + ] = 1.0 x 10 -3 M [OH - ] = 1.0 x 10 -11 M

20. Exercise Determine the hydronium and hydroxide ion concentrations in a solution that is 1.0 x 10 -4 M Ca(OH) 2 Answer: [H 3 O + ] = 5.0 x 10 -11 M [OH - ] = 2.0 x 10 -4 M

21. pH Convenient way to express numbers that tend to be very small. pH = - log[H 3 O + ] [H 3 O + ] = 1 x 10 -pH If [H 3 O + ] = 1 x 10 -7 then pH = 7

22. Figure 14.8 The pH Scale and pH Values of Some Common Substances

23. pOH pOH = - log[OH - ] If [OH - ] = 1 x 10 -7 then pOH = 7 pH + pOH = 14

24. Calculating pH & pOH Determine the pH & pOH of the following solutions. 1) 1 x 10 -3 M HCl pH = 3.0 pOH = 11.0 2) 1 x 10 -5 M HNO 3 pH = 5.0 pOH = 9.0 3) 1 x 10 -4 M NaOH pH = 10.0 pOH = 4.0

25. Calculating the pH & pOH What is the pH of a solution if the hydronium ion conc. is 6.7 x 10 -4 M? pH = 3.17 What is the pH of a solution if the hydronium ion conc. is 2.5 x 10 -2 M? pH = 1.60 pOH = 12.40 Determine the pH of a 2.0 x 10 -2 M Sr(OH) 2 solution. pH = 12.60

26. Calculating the hydronium and hydroxide ion concentrations The pH of a solution is 5.0. What is the hydronium ion concentration? 1 x 10 -5 M The pH of an aqueous solution is measured to be 1.50. Calculate the hydronium ion and hydroxide ion concentrations. [H 3 O + ] = 3.2 x 10 -2 M [OH - ] = 3.2 x 10 -13 M

27. pH Calculations & the Strength of Acids and Bases For strong acids and bases, the hydronium and hydroxide ion concentrations can be directly calculated. For weak acids and bases, hydronium and hydroxide ion concentrations cannot be directly calculated because not all of the molecules are ionized.

28. Homework Complete worksheet

29. Titrations Technique used to measure the amount of acid or base present in a solution. Involves an acid/base neutralization reaction Equivalence Point: point at which the two solutions used in the titration are present in chemically equivalent amounts. Endpoint: The point in a titration at which an indicator changes color.

30. Acid-Base Indicator... marks the end point of a titration by changing color. The equivalence point is not necessarily the same as the end point.

31. Most common acid-base indicators are weak acids, HIn One color with H + ( HIn ) Another color without H + ( In - ) HIn  H + + In - Red Blue

32. Figure 15.6 The Acid and Base Forms of the Indicator Phenolphthalein

33. Figure 15.8 The Useful pH Ranges for Several Common Indicators

34. Titration (pH) Curve A plot of pH of the solution being analyzed as a function of the amount of titrant added. Equivalence point (stoichiometric point): Enough titrant has been added to react exactly with the solution being analyzed.

35. The pH Curve for the Titration of 100.0 mL of 0.10 M HCI with 0.10 M NaOH

36. The pH Curve for the Titration of 50 mL of 0.1 M HC 2 H 3 O 2 with 0.1 M NaOH

37. Titration Technique Go to page 615 in textbook

38. Titration Concentration of one solution is known precisely : Standard solution Concentration of other solution calculated from the chemically equivalent volumes.

39. Acid-Base Properties of Salts Salts = Ionic compounds Salts can behave as ACIDS or BASES.

40. 1.Salts that produce neutral solutions. Composed of cations from strong bases and anions from strong acids. Example: NaCl. NaNO 3, KCl

41. 2.Salts that produce basic solutions. Composed of cations with neutral properties and anions which are the conjugate base of a weak acid. Example: NaCH 3 COO Major species: Na + is neutral CH 3 COO - is conjugate base of weak acid H 2 O is weakly amphoteric

42. CH 3 COO 1- + H 2 O  CH 3 COOH + OH 1- CH 3 COO 1- in water produces OH 1- ions  Basic solution

43. 3.Salts that produce acidic solutions. Composed of cations which are the conjugate acid of a weak base and anions with neutral properties. Example: NH 4 Cl Major species: Cl -, H 2 O, & NH 4 + NH 4 1+ (aq)  NH 3 (aq) + H 1+ (aq)

44. Section 15.2 A Buffered Solution... resists change in its pH when either H + or OH  are added. 1.0 L of 0.50 M H 3 CCOOH –+ 0.50 M H 3 CCOONa pH = 4.74 Adding 0.010 mol solid NaOH raises the pH of the solution to 4.76, a very minor change.

45. One of the most practical applications of buffers is our blood. Most important one is the HCO 3 - and H 2 CO 3.

46. BUFFERS Weak acid and its Salt HF & NaF Or Weak base and its Salt NH 3 & NH 4 Cl

47. Key Points on Buffered Solutions 1.They are weak acids or bases containing a common ion..

48. Buffered Solution Characteristics 4 Buffers contain relatively large amounts of weak acid and corresponding base.  Added OH  reacts to completion with the weak acid. 4 OH - + HA  A - + H 2 O 4 The net result is the hydroxide is replaced by A -.

49. Buffered Solution Characteristics 4 Buffers contain relatively large amounts of weak acid and corresponding base. 4 Added H + reacts to completion with the weak base. 4 H + + A -  HA 4 Free H + converted to HA

50. Homework