1. Equilibrium

2. Introduction:
did you know that. . .

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

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

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

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

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

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

9. An example

10. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.

11. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Initially the rate of the reverse reaction is because…

12. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Initially the rate of the reverse reaction is zero because…

13. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Initially the rate of the reverse reaction is zero because…
Initially the rate of the forward reaction is relatively ___ because…

14. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Initially the rate of the reverse reaction is zero because…
Initially the rate of the forward reaction is relatively fast because…

15. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Over time, the rate of the reverse reaction _______ and the rate of the forward reaction

16. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Over time, the rate of the reverse reaction increases and the rate of the forward reaction decreases.
Eventually the two rates are ____.

17. An example
N H NH3
Imagine putting equal amounts of H2 and N2 in an empty container.
Over time, the rate of the reverse reaction increases and the rate of the forward reaction decreases.
Eventually the two rates are equal.

18. Describing Equilibrium

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

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

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

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

23. Quantifying Equilibrium

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

25. Quantifying Equilibrium
At equilibrium the ratio of product concentrations to reactant concentrations is a constant
The symbol for this constant is: Keq

26. Quantifying Equilibrium
At equilibrium the ratio of product concentrations to reactant concentrations is a constant
The symbol for this constant is: Keq
The equilibrium expression:
[products]
Keq =
[reactants]

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

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

29. Quantifying Equilibrium
Equations with coefficients
Significance of Keq

30. Keq = Quantifying Equilibrium Equations with coefficients
Significance of Keq
A large Keq
Keq =
expression:

31. Keq = [Products] Quantifying Equilibrium Equations with coefficients
Significance of Keq
A large Keq
[Products]
Keq =
expression:
[Reactants]

32. Quantifying Equilibrium
Equations with coefficients
Significance of Keq
A large Keq
A small Keq
Keq =
expression:

33. [Reactants] Quantifying Equilibrium Equations with coefficients
Significance of Keq
A large Keq
A small Keq
[Products]
Keq =
expression:
[Reactants]

34. Types of Equilibria

35. Types of Equilibria 2NO2 (g) N2O4 (g) Homogeneous
All substances in the same physical state

36. C (s) + H2O (g) CO (g) + H2 (g)
Types of Equilibria
Homogeneous
Heterogeneous
C (s) + H2O (g) CO (g) + H2 (g)
Substances not all in the same physical state

37. Writing Equilibria Expressions

38. Writing Equilibria Expressions
[products]x
Keq =
[reactants]y

39. Writing Equilibria Expressions
[products]x
Keq =
Use balanced equation to decide on exponents.
[reactants]y

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

41. Self Check – Ex. 1
Write the equilibrium expression for the reaction below.
NH3 (g) + HCl (g) NH4Cl (s)

42. 2H2S (g) + SO2 (g) 3S (l) + 2H2O (g)
Self Check – Ex. 2
Write the equilibrium expression for the reaction below.
2H2S (g) + SO2 (g) S (l) + 2H2O (g)

43. Calculating Equilibria Constant Values

44. Calculating Equilibria Constant Values
Write equilibrium expression

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

46. Determining if reaction is at equilibrium

47. Determining if reaction is at equilibrium
Write equilibrium expression using Q
Q is the reaction quotient, a value that is compared to Keq

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

49. Determining if reaction is at equilibrium
Write equilibrium expression using Q
Plug in concentrations
If Q > Keq then reaction must ‘go to the left’

50. Determining if reaction is at equilibrium
Write equilibrium expression using Q
Plug in concentrations
If Q > Keq then reaction must ‘go to the left’
If Q < Keq then reaction must ‘go to the right’

51. Determining if reaction is at equilibrium
Write equilibrium expression using Q
Plug in concentrations
If Q > Keq then reaction must ‘go to the left’
If Q < Keq then reaction must ‘go to the right’
If Q = Keq it’s at equilibrium

52. Self Check – Ex. 3 Is this reaction at equilibrium?
N2 (g) + 3H2 (g) NH3 (g) Keq = 0.105
[N2] = M
[H2] = 0.10 M
[NH3] = 0.15 M

53. Self Check – Ex. 4 Is the following reaction at equilibrium?
2CO (g) + O2 (g) CO2 (g) Keq =
[CO] = 0.28 M
[O2] = 0.42 M
[CO2] = 1.21 M

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

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

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

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

60. Shifting Equilibrium: LeChatelier’s Principle
Changing Concentrations
Changing Volume
Changing Temperature
Endothermic reactions: heat is on the left side
Exothermic reactions: heat is on the right side

61. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)

62. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)
[CO] is increased

63. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)
[CO] is increased
Water vapor is added

64. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)
[CO] is increased
Water vapor is added
Carbon is added

65. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)
[CO] is increased
Water vapor is added
Carbon is added
Volume is decreased

66. Self Check – Ex. 5 CO (g) + H2 (g) H2O (g) + C (s)
How do these changes shift equilibrium for this exothermic reaction?
CO (g) + H2 (g) H2O (g) + C (s)
[CO] is increased
Water vapor is added
Carbon is added
Volume is decreased
Temperature is increased

67. Shifting Equilibrium: LeChatelier’s Principle
Changing Concentrations
Changing Volume
Changing Temperature
Haber’s Process

68. Solubility Equilibria

69. Solubility Equilibria
Terms

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

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

72. Solubility Equilibria
Terms
Dissolution
Precipitation
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. Solubility Equilibria
Terms
Solubility Equilibrium
Conditions in which the rate of dissolution equals the rate of precipitation.

76. Solubility Equilibria
Terms
Solubility Equilibrium
The Solubility Product constant (Ksp) Expression
CaCO3 (s) Ca2+ (aq) + CO32- (aq)
Remember to only include aqueous substances.

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

78. Solubility Equilibria
Calculations

79. Solubility Equilibria
Calculations
Finding Ksp

80. Self Check – Ex. 8
When Mg(OH)2 reaches equilibrium the concentration of Mg2+ ions is 1.1 x 10-4 mol/L. Determine Ksp for this reaction.

81. Solubility Equilibria
Calculations
Finding Ksp
Finding solubility
Find the moles/liter of solid that dissolves.

82. Self Check – Ex. 9
Using the following table find the solubility of PbF2.

84. Solubility Equilibria
Calculations
Finding Ksp
Finding solubility
Equilibrium concentrations

85. Self Check – Ex. 10
What are the equilibrium concentrations of Al3+ and OH- in a solution containing the slightly soluble Al(OH)3?

86. Solubility Equilibria
Calculations
Predicting precipitates
If Q ≥ Ksp, then a precipitate forms.

87. the end