1. Equilibrium
SCH4U
organic photochromic molecules respond to the UV light

2. Learning Goals:
Explain the concept of dynamic equilibrium, using examples of physical and chemical equilibrium systems.
Explain the concept of chemical equilibrium and how it applies to the concentration of reactants and products in a chemical reaction at equilibrium.
Create and complete an ICE table for an equilibrium system.
Draw graphs of c vs. t to illustrate a chemical system approaching equilibrium.
Use appropriate terminology such as dynamic equilibrium, reversible reaction, closed system, equilibrium concentrations, phase equilibrium, solubility equilibrium, chemical equilibrium, equilibrium position
blue bottle
glucose reduces the methylene blue to a colourless form
the oxygen from the air oxidized the colourless form back to the blue
the unused glucose with then reduce the methylene blue to colourless

3. Quantitative vs. Reversible
So far we have studied quantitative reactions, reactions which proceed until the limiting reactant is consumed
However, many reactions are not quantitative, they are reversible meaning they can proceed in both forward and reverse directions
A + B C + D

4. How do the following images represent dynamic equilibrium?
When the rate of the forward reaction and the rate of the reverse reaction are equal, the system is in dynamic equilibrium
This does NOT mean that nothing is happening!
Dynamic Equilibrium?
two kids balanced on a teeter-totter
player substitutions during a hockey game
outdoor pool
indoor pool
your own idea?

5. Consider the following…
Over time, what happens to an open beaker of water?
Over time, what happens to the water level if we cover the beaker?
Water molecules build up and reach a high enough vapour pressure and return to liquid phase, set up an equilibrium, no macroscopic changes but a lot is happening on a molecular level, constant coming and going
We can study equilibrium under closed systems!
Write an equation to describe the equilibrium system.

6. Properties of Equilibrium Systems
1. Macroscopic properties (e.g. colour, pressure, concentration, pH) are constant
I.e. It appears as if nothing is changing
2. Can only be reached in a closed system
3. Forward rate = Reverse rate
4. Can be established from either direction
At equilibrium, the concentrations of all reactants and products will remain constant

7. 3 types of Equilibrium Solubility Equilibrium Phase Equilibrium
Chemical Equilibrium
NaCl(s) Na+ (aq) + Cl- (aq)

8. Types of Equilibrium Solubility Equilibrium: A dynamic
equilibrium between a solute and solvent
in a saturated solution in a closed system
Phase Equilibrium: a dynamic equilibrium between different physical states of a pure substance in a closed system
Ex: water/liquid in a sealed container
Chemical Reaction equilibrium: a dynamic equilibrium between reactants and products of a chemical reaction in a closed system

9. Practice
Sketch a graph of Concentration of Reactants and products vs. time for Experiment 1 and 2

10. More Practice 1. Consider the reaction:
Hydrogen gas and hydrogen iodide gas are both clear colourless, while iodine gas is purple.
Explain why when hydrogen gas and iodine gas are added to a closed flask and allowed to react, the colour never fades to colourless.
2. Sketch a graph to illustrate the changes in concentration of each component over time as equilibrium is reached for Experiment 1

11. 3. If experiment 1 did not reach equilibrium, what would the theoretical yield of HI be?
4. What is the actual yield of HI in experiment 1?

12. Solving Equilibrium Problems
Use ICE Tables
I: Initial Concentration (mol/L)
C: Change in concentration
+x if increasing in concentration
-x if decreasing in concentration
Multiply x by the # of moles
E: Equilibrium Concentration (mol/L)

13. Practice #1
In a 4.00L container, 2.50 mol of carbon dioxide gas is decomposed. At equilibrium, [CO2(g)]eq = mol/L. Use an ICE table to find [O2(g)]eq and [CO(g)]eq

14. Practice #2
0.500 mol of NOCl(g) is decomposed in a closed 2.00 L container. [NO(g)]eq = mol/L. Find [Cl2(g)]eq and [NOCl(g)]eq.

15. Practice #3
When 2.00 mol of ethene gas and 1.50 mol of bromine vapour come to equilibrium in a closed 1.00 L container, the equilibrium concentration of bromine vapour is measured at mol/L. Find [C2H4(g)]eq and [C2H4Br2(g)]eq.

16. Practice #3 Continued
Graph the equilibrium reaction for the addition of bromine to ethene.

17. At equilibrium, the rate of the forward and reverse reactions are equal

18. The Equilibrium Constant, Keq
For the reaction:
aA + bB  cC + dD
At equilibrium:
rfwd = rrev
kfwd[A]a[B]b = krev[C]c[D]d
kfwd = [C]c[D]d
krev [A]a[B]b
Keq = [C]c[D]d
[A]a[B]b
Sub in rate law equation
Rearrange
Sub in Keq for
kfwd/krev
Equilibrium constant

19. Keq Conditions
Note: the equilibrium constant expression only works with reactions that occur in a single step
Keq will remain the same as long as the temperature is kept constant (changing the temperature changes the forward and reverse reactions by different amounts and therefore change the Keq)

20. Examples
Write the equilibrium constant expression for the following two reactions

21. Heterogeneous Equilibria
Equilibrium systems can involve all states of matter
However, the concentration of a pure solid or liquid cannot change
Therefore, equilibrium constant expressions will not include solids and liquids

22. Magnitude of Keq
Keq = [C]c[D]d
[A]a[B]b
Keq >> 1 At equilibrium there is more products than reactants. The reaction is product favoured
Keq = 1 At equilibrium there is an equal amount of products and reactants
Keq << 1 At equilibrium there is more reactants than products. The reaction is reactant favoured

24. Example Calculate the value of Keq for the following system
At Equilibrium:
[CO2] = mol/L
[H2] = mol/L
[CO] = [H2O] = mol/L

25. Practice
For each reaction and their respective equilibrium constant, predict whether reactants or products are favoured.
N2(g) + O2(g)  2NO(g) Keq = 4.7 x 10-31
NO(g) + CO(g) 2N2(g) + CO2(g) Keq = 2.2 x 1059

26. Le Chatelier’s Principle
When a chemical system at equilibrium is disturbed by a stress, the system adjusts (shifts) to oppose the change
Stresses include:
Change in concentration
Change in pressure (or volume)
Change in temperature

27. Change in Concentration
A(g) + 3B(g)  2C(g) + heat
Increasing the concentration of the reactants OR
Decreasing the concentration of the products
Will favour the forward reaction, causing the equilibrium to shift to the RIGHT
Decreasing the concentration of the reactants OR
Increasing the concentration of the products
Will favour the reverse reaction, causing the equilibrium to shift to the LEFT
RECALL: Addition or removal of solid or liquids does not change the concentration. Therefore does not cause a shift. I.e. only applies to gases and aqueous solutions.

28. Change in Concentration

29. N2(g) + 3H2(g)  2NH3

30. Change in Pressure
 volume  pressure
 volume  pressure
A(g) + 3B(g)  2C(g) + heat
Increasing the volume of the container OR Decreasing the pressure
Will cause a shift to the side with MORE gas molecules
In our example, it will shift left (4 molreactants > 2 molproducts)
Decreasing the volume of the container OR Increasing the pressure
Will cause a shift to the side with LESS gas molecules
In our example, it will shift right (4 molreactants > 2 molproducts)

32. Change in Temperature In an exothermic reaction:
Increasing the temperature will cause a shift to the LEFT
Decreasing the temperature will cause a shift to the RIGHT

33. Change in Temperature In an endothermic reaction:
Increasing the temperature will cause a shift to the RIGHT
Decreasing the temperature will cause a shift to the LEFT

34. Change in Temperature
Recall: Keq is temperature dependent. Therefore, changes in temperature will also affect Keq
Shift right =  products,  Keq
Shift left =  reactants, Keq

35. DEMONSTRATION

36. Variables that do NOT Affect Equilibrium
Catalysts
Increases reaction rate by lowering activation energy (of BOTH the forward and the reverse reactions equally)
Decreases the time required to reach equilibrium but does not affect the final position of equilibrium
Inert Gases
Increases the pressure, which will increase reaction rate
Increases the probability of successful collisions for BOTH products and reactants equally

37. Practice

38. The Reaction Quotient (Q)
If a chemical system begins with reactants only, it is obvious that the reaction will shift right (to form products).
However, if BOTH reactants and products are present initially, how can we tell which direction the reaction will proceed?
Use a trial value called the reaction quotient, Q
When a reaction is NOT at equilibrium
Q=Keq  the system is at equilibrium
Q > Keq  the system shifts towards reactants to reach equilibrium
Q < Keq the system shifts towards products to reach equilibrium

39. Practice #1
(p. 464) In a container at 450°C, N2 and H2 react to produce NH3. K = When the system is analysed, [N2] = 4.0 mol/L, [H2] = 2.0 X 10-2 mol/L, and [NH3] = 2.2 X 10-4 mol/L. Is the system at equilibrium, if not, predict the direction in which the reaction will proceed.

40. Practice #2
In a container, carbon monoxide and water vapour are producing carbon dioxide and hydrogen at 900oC.
CO(g) + H2O(g)  H2(g) + CO2(g) Keq = 4.00 at 900oC
If the concentrations at one point in the reaction are: [CO(g)] = 4.00 mol/L, [H2O(g)] = 2.00 mol/L, [CO2(g)] = 4.00 mol/L, and [H2(g)] = 2.00 mol/L. Determine whether the reaction has reached equilibrium, and, if not, in which direction it will proceed to establish equilibrium.

41. Practice #2 Answer
products

42. Practice #3 Calculating Equilibrium Concentrations from Initial Concentrations
Carbon monoxide reacts with water vapour to produce carbon dioxide and hydrogen. At 900oC, Keq is Calculate the concentrations of all entities at equilibrium if mol of each entity are initially place in a 1.00L closed container.

43. Practice #4 Calculating Equilibrium Concentrations Involving a Quadratic Equation
If 0.50 mol of N2O4 is placed in a 1.0L closed container at 150oC, what will be the concentrations of N2O4 and NO2 at equilibrium? (Keq = 4.50)

44. Practice #5
Simplifying Assumption: 100 rule (for small K values) If: [initial reactant] > 100, you can simplify the Keq expression K Ex: 2CO2(g)   2CO(g) + O2 (g) If K = 6.40 x 10-7, determine the concentrations of all substances at equilibrium if it starts with [CO2] = mol/L

45. Solubility Equilibrium
Not all ionic compounds are equally soluble
Ionic compounds dissociate into individual ions in aqueous solutions
This can be a reversible system
Example: CaCl2(s) Ca2+(aq) Cl-(aq)
Equilibrium can be reached between the solid substance and its dissolved ions (saturation point)
The solution is saturated at equilibrium (no more ions can dissociate)

46. Solubility Product Constant (Ksp)
An equilibrium equation can be written for solubility reactions
Ex: AgCl (s) Ag+ (aq) + Cl- (aq)
Recall: Since AgCl is a solid, the concentration is not changing, so it is “built in” to the K value:
The new constant is the solubility product constant (Ksp)

47. Example
Eg: Lead (II) chloride has a molar solubility of 1.62x102 mol/L at 25oC. What is the Ksp of this salt?
PbCl2   Pb Cl-
Ksp = [Pb2+][Cl-]2
[Pb2+] = [PbCl2] = 1.62x10-2mol/L
[Cl-] = 2[PbCl2] = 2(1.62x10-2mol/L) = 3.24 x 10-2mol/L
Ksp = [1.62x10-2mol/L][3.24 x 10-2mol/L]2
= 1.7x10-5

48. Example 2 Ksp = [Ag+][Cl-] 1.8x10-10 = [X][X] 1.8x10-10 = X2
The Ksp of silver chloride at 25oC is 1.8x What is the molar solubility of AgCl?
AgCl   Ag Cl-
AgCl  
Ag+
Cl-
Initial -
Change
+ X
Equilibrium X
Ksp = [Ag+][Cl-]
1.8x10-10 = [X][X]
1.8x10-10 = X2
X = 1.34x10-5M

49. The size of Ksp depends on the solubility of the salt.
Large Ksp: [ions] at equilibrium is high, salt is very soluble
Small Ksp: [ions] at equilibrium is low, salt has low solubility
To determine whether a precipitate will form during a reaction, a trial solubility product constant can be determine which is denoted by the symbol Qsp.
Qsp < Ksp : Shifts right to equilibrium – all solid dissolving
Qsp > Ksp : Shifts left to equilibrium – precipitate forms
Qsp = Ksp : Equilibrium (saturated) – no precipitate

50. Example 1
Will a ppt form if a solution of CaSO4 is made such that [Ca2+] = mol/L and [SO42-]= mol/L?
Ksp =7.1x10-5
(Ans: Q = 8.5 X 10-5, ppt forms)

51. Example 2
Will a precipitate form when a solution is made by dissolving PbCl2 into water.
[Pb2+]=0.0001mol/L. [Cl-]=0.05mol/L. Ksp = 1.2x10-5