1. Chapter 14 Acids and Bases

2. Acid/Base Theories Arrhenius Theory –Acids produce H + ions in solution –Bases produce OH - ions in solution –Downside Must be in solution and must have those ions Bronsted-Lowry Theory –Acids are H + donors (Proton donors) –Bases are H + acceptors (Proton acceptors)

3. Vocabulary H + is the hydrogen ion –Just a proton H 3 O + is a hydronium ion –It is the way H + exists in water Water accepts a hydrogen ion and becomes H 3 O + Either way is fine the first is just easier

4. Conjugates Acid/Base Pairs Conjugate Acid is formed when a base gains a proton Conjugate Base is what remains after an acid donates a proton Ex – HNO 3 + H 2 O  H 3 O + + NO 3 -

5. General Form for Acids and Bases HA + B  A - + BH + HA is an acid B is a base A - is the conjugate base –Just the negative ion of the acid BH + is the conjugate acid –Just the base plus a hydrogen

6. Strong and Weak Acids Strong acids completely ionize in solution –Nitric, Perchloric, Sulfuric, Hydrochloric, Hydrobromic, Hydroiodic –Weak Conjugate bases Weak Acids only partially ionize in solution –Every other acid –Equilibrium is established in the ionization –Weak acids have Ka values –Strong Conjugate bases

7. Acid Dissociation Constant, Ka HA (aq) + H 2 O  A - + H 3 O + For weak acids equilibrium is established Equilibruim constant is Ka Ka=[A - ][H 3 O + ]/[HA] Values tend to be small –Because CB is fairly strong Strong acids do not have Ka values.

8. Acid Terms Monoprotic – One acidic hydrogen Polyprotic – Many acidic hydrogens Diprotic – Two acidic hydrogens Triprotic – Three acidic hydrogens Oxyacid – Acid that has an acidic hydrogen attached to an oxygen Organic Acid – Acid that has the acidic hydrogen attached to the carboxyl group

9. Water As An Acid and Base Water is amphoteric – Both an acid and base H 2 O + H 2 O  H 3 O + + OH - Equilibrium system that always has the same value Called Kw

10. Autoionization of Water, Kw H 2 O + H 2 O  H 3 O + + OH - Kw = [H 3 O + ] [OH - ] In pure water the products have the same concentration, 1.00x10 -7 M Value of Kw = 1.00x10 -14 at 25ºC Concentrations can change is acid or base is added

11. Acid, Base, or Neutral If the concentration of H + = OH - –Neutral If the concentration of H + > OH - –Acidic If the concentration of H + < OH - –Basic

12. What is the hydrogen ion concentration when the hydroxide ion concentration is 1.00x10 -5 M

13. What is the hydroxide ion concentration when the when the concentration of nitric acid is 0.0010M?

14. Homework P. 704 30, 32, 33, 35, 39ab,40ab

15. Logarithms The logarithm of a number to a given base (commonly 10) is the power or exponent to which the base must be raised in order to produce the number. SAY WHAT!

16. Examples If your question says log 100 = x It is saying to what power must 10 be raised to equal 100 10 x = 100 10 2 = 100 So x = 2

17. Examples log 1 = x x = 0 log 10 = x x = 1 log 1000 = x x = 3 log 1x10 6 = x x = 6

18. Examples Logarithms can also be used for numbers smaller than 1 log 0.1 = x x = -1 log 0.01 = x x = -2 log 1x10 -5 x = -5

19. Examples If they are not easy to calculate you can do it on your calculator log 15 = x You can approximate it between... –Type log 15 on your calculator x = 1.18

20. Examples If your question says log x = 7 It is saying 10 to the 7 th power is what number 10 7 = 1x10 7 log x = 3 x = 1000 log x = -3 x = 0.001

21. Tougher Examples log 234 = x x = 2.37 log x = -3.3 x = 5.0x10 -4 -log 9.1x10 -5 = x x = 4.0 -log x = 12.1 x = 7.9x10 -13

22. pH Negative logarithm of the concentration of hydrogen ions in solution pH = -log [H + ] pH means power of Hydrogen Measures how acidic or basic a solution is pH scale typically goes from 0 to 14 pH < 7 Acidic pH > 7 Basic

23. Highly acidic = low pH Highly basic = high pH

25. Significant Figures and pH Digits after the decimal are the only ones that are significant in pH values pH = 4.44 2 Significant Figures pH = 10.874 3 Significant Figures If your [H + ] is 0.088 M your pH is 1.06

26. Determining [H + ] in solution The concentration of a strong acid is equal to the H + concentration. 0.010 M HCl has an [H + ] of 0.010 M To obtain the [H + ] for weak acids you must use equilibrium –Need Ka data Discuss hydroxide later

30. Other Info Turn Kw into a log equation pH + pOH = 14.00

31. pH’s You MUST Know When the [H + ] is _______ the pH is _____ 0.10 M 1.00 0.0010M 3.00 1.0x10 -6 M 6.00 1.0x10 -10 M 10.00

32. Find the pOH, [H + ], [OH - ] of lemon juice that has a pH of 2.48

33. Determine the pH of 0.150M HCl

34. Determine the pH of 2.3x10 -3 M Hydrocyanic Acid HCN. Ka = 6.2x10 -10

35. Homework Page 705 #’s 45,47,50,52,53,58

36. Mixtures of Weak Acids When there are mixtures of weak acids in solution determining the pH could be a difficult problem. However –The acid with the largest Ka will control the pH of the solution

37. Solutions of 0.10M HF (Ka = 7.2x10 -4 ) and 0.10M HCN (Ka = 6.2x10 -10 ) are mixed A) Which acid will control the pH of the solution? Why? B) What is the pH of the resulting solution.

38. Percent Dissociation Ratio of the concentration of the dissociated ions to the initial concentration Found just like the doing the 5% check Can be used to find pH and Ka

39. A 0.25M solution of HClO is 20.% dissociated. A) What is the pH? B) What is the Ka value of the acid?

40. Strong Bases Strong bases are any compound containing the hydroxide ion –Group 1 hydroxides are very soluble –Group 2 less soluble but still strong Group 2 have two hydroxides per mole –Be Careful

41. Determine the pH of 0.022M Sr(OH) 2

42. Weak Bases Organic Bases are weak bases. (Ammonia) –Contain Nitrogen (Amines) CH 3 NH 2 – Methylamine –Must have a lone pair of electrons –The hydroxide will come from water (CH 3 ) 3 N + H 2 O 

43. Cont. Weak bases have Kb values B + H 2 O  BH + + OH - The conjugate acid of a weak base is stronger than water.

44. Determine the pH of 15.0 M ammonia. The Kb is 1.8x10 -5

45. Homework Page 705 #’s 62,63a,72,73,77ab,87

46. Polyprotic Acids Acids with more than one acidic hydrogen Dissociate in a “stepwise” process H 3 PO 4  H + + H 2 PO 4 - Ka 1 =7.5x10 -3 H 2 PO 4 -  H + + HPO 4 -2 Ka 2 =6.2x10 -8 HPO 4 -2  H + + PO 4 -3 Ka 3 =4.8x10 -13 Ka 1 > Ka 2 > Ka 3 Successive dissociation do not effect pH (except sulfuric acid)

47. Determine the pH of 5.0M H 3 PO 4 (Ka=7.5x10 -3 ) and the [H 2 PO 4 - ], [HPO 4 -2 ], [HPO 4 -3 ]

48. Sulfuric Acid Sulfuric acids has two dissociations –The first is strong –The second is weak The second dissociation is quite strong, but it is not complete H 2 SO 4  H + + HSO 4 - (Strong) HSO 4 -  H + + SO 4 -2 Ka = 1.2x10 -2

49. Continued When you write sulfuric acid in net ionic equations only use the first equation H 2 SO 4  H + + HSO 4 - The second dissociation only needs to be considered in dilute solutions. A 1.0M solution of sulfuric acid will have a lower pH than a 1.0M solution of HCl

50. Acid / Base Properties of Salts Some salts have acid base properties –Make a solution have a pH below or above 7 Some have no acid base properties –Make a solution with a pH of 7

51. Neutral Salts Contain the following A metallic ions –(Except Aluminum) A conjugate base of a strong acid –Such AS.... Conjugate base of strong acids are weaker than water.

52. Basic Salts Contain the following A metallic ion A conjugate base of a weak acid –Such AS... Conjugate base of weak acids are stronger than water.

53. Acidic Salts Contain the following A conjugate base of a strong acid A conjugate acid of a weak base –Such AS... Conjugate acid of weak bases are stronger than water.

54. Acidic Aluminum Ions Highly charged cations polarizes the O-H bond in water Hydrogens in water become acidic Ion becomes hydrated Al(H 2 O) 6 +3

55. OMGosh I Have Both Types of Ions If a salt has acidic and basic ions –Such AS... If Ka > Kb the solution is acidic If Ka < Kb the solution is basic If Ka = Kb the solution is neutral

56. Converting Ka to Kb If you know an acids Ka value you can find its Kb value as a salt. If you know a bases Kb value you can find its Ka value as a salt Kw = Ka * Kb

57. Calculate the pH of 0.33 M NaHCO 3

58. Homework P 707 #’s 94,95b, 98,99,102,105a, 107,112

59. Effect of Structure On Acids Different bonding patterns lead to differences in the strength of acids Two different types of acids –Hydrohalic –Oxyacids

60. Hydrohalic Acids Acids that contain a hydrogen and a halogen The weaker the bond the stronger the acid BondEnthalpy H—F 567 H—Cl 433 H—Br 366 H—I 299

61. Oxyacids Tend to be of the form H – O – X The greater the electronegativity of X the stronger the acid Pulls electron density away and weakens the H – O bond Which is a stronger acid? H – O – Cl or H – O – Br –H – O – Cl because Cl has a greater E.N.

62. Oxyacids Bases tend to be of the form X – O – H too NaOH Why Metal has a low electronegativity and the O – H bond is not weakened

63. Oxyacids Acids with the same X element can have different number of oxygens As the number of oxygens increases so does the strength of the acid The extra oxygens weaken the H – O bond

65. Oxides Nonmetalic oxide in water gives an acid CO 2 + H 2 O  H 2 CO 3 Acidic anhydrides (Acid w/o water) Nonmetal must keep its oxidation # Metallic oxides in water gives a base Na 2 O + H 2 O  2NaOH Basic anhydrides

66. Lewis Acids and Bases Lewis acids are electron pair acceptors H + for example has no electrons Lewis bases are electron pair donors NH 3 has an electron pair to share

67. Homework P 708 #’s 113,114,115ab,116ab,119