1. Equilibrium Chap. 18

2. I.Introduction: did you know that...

3. I.Introduction: A.Reactions don’t always completely use up reactants did you know that... A + 2B C Consider a reaction between 50 molecules of A and 100 molecules of B. What’s left at the end?

4. I.Introduction: A.Reactions don’t always completely use up reactants B.Reactions are reversible did you know that... Can you ‘un-rust’ or ‘un-combust’?

5. I.Introduction: A.Reactions don’t always completely use up reactants B.Reactions are reversible C.Arrows represent the forward did you know that... A + 2B C

6. I.Introduction: A.Reactions don’t always completely use up reactants B.Reactions are reversible C.Arrows represent the forward and reverse reactions did you know that... A + 2B C

7. I.Introduction: A.Reactions don’t always completely use up reactants B.Reactions are reversible C.Arrows represent the forward and reverse reactions D.Reaction rates depend on the concentration of reactants did you know that...

8. I.Introduction: A.Reactions don’t always completely use up reactants B.Reactions are reversible C.Arrows represent the forward and reverse reactions D.Reaction rates depend on the concentration of reactants E.The symbols, [ ] mean concentration (molarity) did you know that...

9. II. An example

10. N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

11. II. An example A.Initially the rate of the reverse reaction is because… N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

12. II. An example A.Initially the rate of the reverse reaction is zero because… N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

13. II. An example A.Initially the rate of the reverse reaction is zero because… B.Initially the rate of the forward reaction is relatively ___ because… N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

14. II. An example A.Initially the rate of the reverse reaction is zero because… B.Initially the rate of the forward reaction is relatively fast because… N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

15. II. An example C.Over time, the rate of the reverse reaction _______ and the rate of the forward reaction. N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

16. II. An example C.Over time, the rate of the reverse reaction increases and the rate of the forward reaction decreases. D.Eventually the two rates are ____. N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

17. II. An example C.Over time, the rate of the reverse reaction increases and the rate of the forward reaction decreases. D.Eventually the two rates are equal. N 2 + 3H 2 2NH 3 Imagine putting equal amounts of H 2 and N 2 in an empty container.

18. III. Describing Equilibrium

19. A.A situation in which the forward and reverse reaction rates.

20. III. Describing Equilibrium A.A situation in which the forward and reverse reaction rates are equal. B.A situation where the amounts of reactants/ products.

21. III. Describing Equilibrium A.A situation in which the forward and reverse reaction rates are equal. B.A situation where the amounts of reactants/ products remain constant.

22. III. Describing Equilibrium C.Equilibrium is dynamic. Although the amount of products and reactants remains constant the reaction doesn’t ‘stop’ at equilibrium

23. IV. Quantifying Equilibrium

24. A.At equilibrium the ratio of product concentrations to reactant concentrations is a constant

25. IV. Quantifying Equilibrium A.At equilibrium the ratio of product concentrations to reactant concentrations is a constant B.The symbol for this constant is: K eq

26. IV. Quantifying Equilibrium A.At equilibrium the ratio of product concentrations to reactant concentrations is a constant B.The symbol for this constant is: K eq C.The equilibrium expression: K eq = [products] [products]

27. IV. Quantifying Equilibrium D.Equations with coefficients aA + bB cC + dD equation: expression:

28. IV. Quantifying Equilibrium D.Equations with coefficients K eq = [C] c x [D] d [A] a x [B] b aA + bB cC + dD equation: expression:

29. IV. Quantifying Equilibrium D.Equations with coefficients E.Significance of K eq

30. IV. Quantifying Equilibrium D.Equations with coefficients E.Significance of K eq K eq = expression: 1.A large K eq

31. IV. Quantifying Equilibrium D.Equations with coefficients E.Significance of K eq K eq = [Products] [Reactants] expression: 1.A large K eq

32. IV. Quantifying Equilibrium D.Equations with coefficients E.Significance of K eq K eq = expression: 1.A large K eq 2.A small K eq

33. IV. Quantifying Equilibrium D.Equations with coefficients E.Significance of K eq K eq = expression: 1.A large K eq 2.A small K eq [Products] [Reactants]

34. V. Types of Equilibria

35. A.Homogeneous All substances in the same physical state 2NO 2 (g) N 2 O 4 (g)

36. V. Types of Equilibria A.Homogeneous B.Heterogeneous Substances not all in the same physical state C (s) + H 2 O (g) CO (g) + H 2 (g)

37. VI.Writing Equilibria Expressions

38. A.K eq = [products] x [reactants] y

39. VI.Writing Equilibria Expressions A.K eq = B.Use balanced equation to decide on exponents. [products] x [reactants] y

40. VI.Writing Equilibria Expressions A.K eq = B.Use balanced equation to decide on exponents. C.Don’t include (s) or (l) physical states in expression. [products] x [reactants] y

41. Self Check – Ex. 1 Write the equilibrium expression for the reaction below. NH 3 (g) + HCl (g) NH 4 Cl (s)

42. Self Check – Ex. 2 Write the equilibrium expression for the reaction below. 2H 2 S (g) + SO 2 (g) 3S (l) + 2H 2 O (g)

43. VII.Calculating Equilibria Constant Values

44. A.Write equilibrium expression

45. VII.Calculating Equilibria Constant Values A.Write equilibrium expression B.Plug in equilibrium concentrations * a set of equilibrium concentrations is called an equilibrium position

46. VIII.Determining if reaction is at equilibrium

47. A.Write equilibrium expression using Q Q is the reaction quotient, a value that is compared to K eq

48. VIII.Determining if reaction is at equilibrium A.Write equilibrium expression using Q B.Plug in concentrations

49. VIII.Determining if reaction is at equilibrium A.Write equilibrium expression using Q B.Plug in concentrations 1.If Q > K eq then reaction must ‘go to the left’

50. VIII.Determining if reaction is at equilibrium A.Write equilibrium expression using Q B.Plug in concentrations 1.If Q > K eq then reaction must ‘go to the left’ 2.If Q < K eq then reaction must ‘go to the right’

51. VIII.Determining if reaction is at equilibrium A.Write equilibrium expression using Q B.Plug in concentrations 1.If Q > K eq then reaction must ‘go to the left’ 2.If Q < K eq then reaction must ‘go to the right’ 3.If Q = K eq it’s at equilibrium

52. Self Check – Ex. 3 Is this reaction at equilibrium? N 2 (g) + 3H 2 (g) 2NH 3 (g) K eq = 0.105 [N 2 ] = 0.0020 M [H 2 ] = 0.10 M [NH 3 ] = 0.15 M

53. Self Check – Ex. 4 Is the following reaction at equilibrium? 2CO (g) + O 2 (g) CO 2 (g) K eq = 0.0021 [CO] = 0.28 M [O 2 ] = 0.42 M [CO 2 ] = 1.21 M

54. IX.Shifting Equilibrium: LeChatelier’s Principle

55. LeChatelier’s Principle When a change is imposed on a system at equilibrium, the equilibrium position shifts to minimize that change

56. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations Only affects equilibrium for gases and aqueous substances.

57. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations B.Changing Volume When volume decreases equilibrium shifts to the side with the fewest gas particles.

58. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations B.Changing Volume C.Changing Temperature Decreasing temperature shifts equilibrium toward side with ‘heat’ written on it.

59. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations B.Changing Volume C.Changing Temperature 1.Endothermic reactions: heat is on the left side

60. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations B.Changing Volume C.Changing Temperature 1.Endothermic reactions: heat is on the left side 2.Exothermic reactions: heat is on the right side

61. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s)

62. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s) [CO] is increased

63. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s) [CO] is increased Water vapor is added

64. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s) [CO] is increased Water vapor is added Carbon is added

65. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s) [CO] is increased Water vapor is added Carbon is added Volume is decreased

66. Self Check – Ex. 5 How do these changes shift equilibrium for this exothermic reaction? CO (g) + H 2 (g) H 2 O (g) + C (s) [CO] is increased Water vapor is added Carbon is added Volume is decreased Temperature is increased

67. IX.Shifting Equilibrium: LeChatelier’s Principle A.Changing Concentrations B.Changing Volume C.Changing Temperature D.Haber’s Process

68. X.Solubility Equilibria

69. A.Terms

70. X.Solubility Equilibria A.Terms 1.Dissolution Process in which an ionic solid dissolves into a liquid, separating into its ions, and ‘entering the solution’.

71. X.Solubility Equilibria A.Terms 1.Dissolution 2.Precipitation Process in which dissolved ions rejoin to form an ionic compound and they ‘leave the solution’.

72. X.Solubility Equilibria A.Terms 1.Dissolution 2.Precipitation 3.Solubility The amount of solute that dissolves in a given volume of solvent.

73. Self Check – Ex. 6 If a substance had a solubility of zero we’d say that substance is in water.

74. Self Check – Ex. 6 If a substance had a solubility of zero we’d say that substance is insoluble in water.

75. X.Solubility Equilibria A.Terms B. Solubility Equilibrium Conditions in which the rate of dissolution equals the rate of precipitation.

76. X.Solubility Equilibria A.Terms B. Solubility Equilibrium C. The Solubility Product constant (K sp ) Expression Remember to only include aqueous substances. CaCO 3 (s) Ca 2+ (aq) + CO 3 2- (aq)

77. Self Check – Ex. 7 Write the solubility product expression for calcium hydroxide, Ca(OH) 2. *hint – first write the solubility equation.

78. X.Solubility Equilibria D.Calculations

79. X.Solubility Equilibria D.Calculations 1.Finding K sp

80. Self Check – Ex. 8 When Mg(OH) 2 reaches equilibrium the concentration of Mg 2+ ions is 1.1 x 10 -4 mol/L. Determine K sp for this reaction.

81. X.Solubility Equilibria D.Calculations 1.Finding K sp 2.Finding solubility Find the moles/liter of solid that dissolves.

82. Self Check – Ex. 9 Using the following table find the solubility of PbF 2.

84. X.Solubility Equilibria D.Calculations 1.Finding K sp 2.Finding solubility 3.Equilibrium concentrations

85. Self Check – Ex. 10 What are the equilibrium concentrations of Al 3+ and OH - in a solution containing the slightly soluble Al(OH) 3 ?

86. X.Solubility Equilibria D.Calculations E.Predicting precipitates If Q ≥ K sp, then a precipitate forms.

87. the end