Download presentation
Presentation is loading. Please wait.
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
Similar presentations
© 2025 SlidePlayer.com. Inc.
All rights reserved.