1. Chapter 19 Acid-Base Theories

2. Objectives Define the properties of Acids and Bases Compare and contrast acids and bases as defined by the theories of Arrhenius, Bronsted-Lowry, and Lewis

3. Acids They taste sour Conduct electricity ◦ Strong or weak electrolytes React with metals to form H 2 gas Change color of indicators (shown in 2 slides) React with bases to form water and a salt (not necessarily NaCl)

4. Acids Continued pH less than 7 (more later) React with carbonates and bicarbonates to produce a salt, water and CO 2 How can we tell something is an acid? ◦ Usually starts with H ◦ HCl, HNO 3, H 2 SO 4 (not water!)

5. Acids and Indicators Blue litmus paper turns red

6. Acids have a pH less than 7

7. Acids and Metals React with active metals to form salts and hydrogen gas HCl (aq) + Mg (s)  MgCl 2(aq) + H 2(g) What type of reaction is this?

8. Acids React with Carbonates and Bicarbonates HCl + NaHCO 3 NaCl + H 2 O + CO 2 Hydrochloric acid + sodium bicarbonate salt + water + carbon dioxide An old-time home remedy for relieving an upset stomach

9. Effects of Acid Rain

11. Acids React with Bases HCl + NaOH  NaCl + H 2 O Neutralization ◦ Acid + Base = salt + water What type of reaction is this?

12. Bases React with acids to form salt and water Taste bitter Feel slippery Strong or weak electrolytes Change color of indicators

13. Bases Examples Sodium hydroxide ◦ Drain cleaner Potassium hydroxide ◦ Alkaline batteries Magnesium hydroxide ◦ Milk of magnesia

14. Bases and Indicators Red litmus turns blue Phenolphthalein turns purple

15. Bases have a pH greater than 7

16. Theories 3 Acid-Base theories: ◦ Arrhenius ◦ Bronsted-Lowry ◦ Lewis

17. Svante Arrhenius Swedish chemist Won Nobel Prize in Chemistry in 1903 One of the first chemists to explain chemical behavior of acids and bases

18. Svante Arrhenius (1859-1927)

19. Arrhenius Theory Acids produce hydrogen ions (H + ) in aqueous solutions ◦ HCl  H + + Cl - Bases produce hydroxide ions (OH - ) in aqueous solutions ◦ NaOH  Na + + OH - Limited to water solutions Only one kind of base – hydroxide

20. Polyprotic Acid Some compounds have more than one ionizable hydrogen to release HNO 3 – monoprotic H 2 SO 4 – diprotic H 3 PO 4 – triprotic More hydrogens does not mean stronger!

21. Acids Not all compounds that have hydrogen are acids ◦ Example: water Not all the hydrogen in an acid will be released as ions

22. Arrhenius Examples HCl CH 3 COOH H 2 SO 4 H 3 PO 4 NaOH Ca(OH) 2 Al(OH) 3

23. Johannes Brønsted Thomas Lowry (1879-1947) (1874-1936) Denmark England

24. Bronsted-Lowry Acid = H + donor Base = H + acceptor Acids and bases always come in pairs

25. Why is Ammonia a Base? NH 3 is the H + acceptor (base) and water is the H + donor (acid) This causes the OH - concentration to increase and the ammonium solution becomes basic

26. Acid-Base Pairs Conjugate base – the remainder of the original acid after it donates its hydrogen ion Conjugate acid – formed when the original base gains a hydrogen ion Conjugate acid-base pair is related by loss or gain of single hydrogen ion

27. Water Water is amphoteric ◦ Can act as both an acid and a base

28. Gilbert Lewis (1875-1946)

29. Lewis Theory Lewis acid – electron pair acceptor Lewis base – electron pair donor Most general definition ◦ Acids don’t even need hydrogen!

30. Lewis Acid-Base

32. Hydrogen Ions and Acidity

33. Objectives Describe how [H + ] and [OH - ] are related in an aqueous solution Classify a solution as neutral, acidic, or basic given the hydrogen-ion and hydroxide ion concentration Convert hydrogen ion concentrations into pH and hydroxide ion concentrations to pOH Describe the purpose of an acid-base indicator

34. Ions from Water Water ionizes into H + and OH - ◦ H 2 O  H + + OH - ◦ Self-ionization of water Occurs at a very small extent ◦ [H + ] = [OH - ] = 1x10 -7 M Since the concentrations are equal the solution is neutral

35. Ion Product Constant K w = ion product constant of water K w = [H + ][OH - ] K w = 1x10 -14 M 2 If [H + ] > 10 -7 – acidic solution ◦ More H + than OH - If [OH - ] > 10 -7 – basic solution ◦ More OH - than H +

36. Problem If the [H + ] of a cola solution is 2.0 x 10 -6 M, is the solution acidic, basic or neutral? What is the [OH - ] of this solution?

37. On Your Own Classify each solution as acidic, basic or neutral: ◦ [H + ] = 6.0 x 10 -10 M ◦ [OH - ] = 3.0 x 10 -2 M ◦ [H + ] = 2.0 x 10 -7 M ◦ [OH - ] = 1.0 x 10 -7 M

38. pH pH = pouvoir hydrogene = “hydrogen power” pH = -log[H + ] Neutral solution: pH = 7 Acidic solution: pH < 7 Basic solution: pH > 7

40. Measuring pH Steps: ◦ Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod. ◦ Compare the color to the chart on the vial – then read the pH value.

41. Acid-Base Indicator Limitations: ◦ Usually given for a certain temperature – may change at different temps ◦ The solution may already have a color (i.e. paint) ◦ The ability to the human eye to distinguish colors is limited

42. Acid-Base Indicator A pH meter may give more definitive results ◦ Measures voltage between 2 electrodes ◦ Typically accurate to within 0.01 pH unit of true pH ◦ Instruments must be calibrated

43. pOH pOH = -log[OH - ] pH + pOH = 14 Thus, a solution with pOH 7 is acidic Not commonly used

44. pH and Significant Figures For pH calculations, the hydrogen ion concentration is usually expressed in scientific notation [H 1+ ] = 0.0010 M = 1.0 x 10 -3 M, and 0.0010 has 2 significant figures the pH = 3.00, with the two numbers to the right of the decimal corresponding to the two significant figures

45. Problem What is the pH of a solution with a hydrogen-ion concentration of 4.2 x 10 -10 M? The pH of an unknown solution is 6.35. what is the hydrogen ion concentration?

46. Another What is the pH of a solution if [OH - ] = 4.0 x 10 -11 M?

47. On Your Own Calculate pH of each solution ◦ [H + ] = 5.0 x 10 -5 M ◦ [OH - ] = 4.5 x 10 -11 M What are the hydrogen ion concentrations for solutions with following pH values: ◦ 4.00 ◦ 11.55

48. Strengths of Acids and Bases

49. Objectives Define strong acids and weak acids Describe how an acid’s strength is related to the value of its acid dissociation constant Calculate an acid dissociation constant (K a ) from concentration and pH measurements Order acids by strength according to their acid dissociation constants (K a ) Order bases by strength according to their base dissociation constants (K b )

50. Strength Acid-base strength is classified according to how much they ionize in water ◦ Strong = completely ionize ◦ Weak = slightly ionize Strength ≠ concentration

51. Strength Strong means it will make many ions when dissolved Mg(OH) 2 is a strong base – it falls completely apart – nearly 100% ◦ But not much of it dissolves in the first place – not concentrated

52. Problem assuming 100% dissociation! A 1.45x10 -3 M aqueous solution of sulfuric acid is made. Calculate the pH.

53. Try it 1 g of phosphoric acid is added to water to make a 2 L solution. Calculate the pH of the solution.

54. Strong Acid Dissociation (makes 100 % ions)

55. Weak Acid Dissociation (only partially ionizes)

56. Measuring Strength

57. Same for Bases

58. Strength vs. Concentration The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume The words strong and weak refer to the extent of ionization of an acid or base

59. Practice Problem A 0.1000 M solution of acetic acid is only partially ionized. From measurements of the pH of the solution, [H + ] is determined to be 1.34 x 10 -3 M. What is the acid dissociation constant (K a ) of acetic acid?

60. Try it In an exactly 0.1 M solution of HCOOH, [H + ] = 4.2 x 10 -3 M. Calculate the K a of the HCOOH.

61. A Bit Harder Calculate [H + ], pH, and % dissociation in 0.1 M solution of weak acid HNO 2. ◦ Ka = 5.0 x 10 -5

62. On Your Own Calculate the pH and % dissociation of a 0.98 M solution of acetic acid. ◦ Ka = 1.77 x 10 -5

63. HARDER! (You will see it in AP) The K a of HF is 6.46 x 10 -4. Calculate the pH of 0.0100 M solution of HF.

64. Neutralization Reactions

65. Objectives Define the products of an acid-base reaction Explain how acid base titration is used to calculate the concentration of an acid or a base Explain the concept of equivalence in neutralization reactions Describe the relationship between equivalence point and the end point of a titration

66. Neutralization Reaction A reaction in which an acid and a base react in aqueous solution to produce a salt and water HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l) H 2 SO 4(aq) + 2KOH (aq)  K 2 SO 4(aq) + 2 H 2 O (l)

67. Titration Process of adding a known amount of solution of known concentration to determine the concentration of another solution ◦ Balanced equation gives mole ratio Equivalence point – when the moles of hydrogen ions equals moles of hydroxide ions (neutralized)

68. Sample Problem How many moles of sulfuric acid are required to neutralize 0.50 M sodium hydroxide solution? The solution is 2 L. ◦ What is the concentration of sulfuric acid?

69. Another What concentration of 15 mL of potassium hydroxide are needed to completely neutralize 25 mL of 1.56 M phosphoric acid?

70. Titration The concentration of an acid (or base) in solution can be determined by performing a neutralization reaction ◦ Indicator is used – shows neutralization occurs  Phenolphthalein often used because colorless in acid and pink in base

71. Neutralization Steps 1. Measured volume of unknown acid concentration is added to flask 2. Several drops of indicator added 3. A base of known concentration is added SLOWLY until the indicator changes color 4. Measure volume of base added

72. Neutralization The solution of known concentration is called the standard solution ◦ added by using a buret Continue adding until the indicator changes color ◦ called the “end point” of the titration

73. Salts in Solution

74. Objectives Describe when a solution of salt is acidic or basic Demonstrate with equations how buffers resist change in pH

75. Salt Salt – ionic compound ◦ Comes from anion of an acid ◦ Comes from cation of base ◦ Formed during neutralization reaction ◦ Some are neutral while others are acidic or basic

76. Salt Hydrolysis Salt hydrolysis – a salt that reacts with water to produce an acid or base Hydrolyzing salts come from: ◦ Strong acid + weak base ◦ A weak acid + strong base Strong – degree of ionization

78. Salt Hydrolysis To see if the resulting salt is acidic or basic you need to look at the “parent” acid and base Let’s look… ◦ HCl + NaOH  NaCl + H 2 O ◦ H 2 SO 4 + NH 4 OH  (NH 4 ) 2 SO 4 + H 2 O ◦ CH 3 COOH + KOH  CH 3 COOK + H 2 O

79. Predict whether the following salts will from a solution that is acidic, basic or neutral. a. FeCl 3 e. NaF b. NH 4 CN*f. LiClO 4 c. K 2 CO 3 d. CaBr 2 h. NH 4 Br

81. Buffers Buffer – solution in which the pH remains constant even when acid and bases are added to them ◦ Made from a pair of chemicals:  Weak acid and one of its salts  Weak base and one of its salts

82. Buffers Better able to resist change in pH than pure water Since it is a pair of chemicals: ◦ One chemical neutralizes any acid added while the other chemical would neutralize any base added

83. Buffer Capacity Buffers resist change in pH HOWEVER if you add enough acid or base the pH WILL change Buffer capacity – the amount of acid or base that can be added before a significant change in pH

84. Buffers The 2 buffers crucial to maintaining the pH of human blood: ◦ Carbonic acid (H 2 CO 3 ) & Hydrogen carbonate (HCO 3 -1 ) ◦ Dihydrogen phosphate (H 2 PO 4 -1 ) & monohydrogen phosphate (HPO 4 -2 )