Download Unit 4 Chemical Kinetics and Chemical Equilibrium

Applications of Equilibrium Constants
 Kc and Kp can be used to determine the
concentration of reactants and/or products
at equilibrium.
Applications of Equilibrium Constants
Example: Calculate the concentration of Ca2+
ions and F- ions that are present in a saturated
aqueous solution of CaF2 if Ksp = 3.90 x 10-11 at
25oC.
Applications of Equilibrium Constants
CaF2 (s)
Ca2+ (aq) + 2 F- (aq)
Applications of Equilibrium Constants
Example: At 250oC the reaction
PCl5 (g)
PCl3 (g) + Cl2 (g)
has an equilibrium constant Kc = 1.80. If the
initial concentration of PCl5 in a reactor at
250oC is 0.0400 M, what are the concentrations
of PCl5, PCl3, and Cl2 in the mixture at
equilibrium?
Write the expression
for Kc:
Applications of Equilibrium Constants
Set up a table showing the initial
concentrations, changes in concentration, and
equilibrium concentrations.
PCl5 (g)
Initial
Change
Equil.
0.0400 M
PCl3 (g) + Cl2 (g)
0.000 M
0.000 M
Applications of Equilibrium Constants
Applications of Equilibrium Constants
Substitute the equilibrium values into the
expression for Kc
Applications of Equilibrium Constants
Applications of Equilibrium Constants
Applications of Equilibrium Constants
Applications of Equilibrium Constants
Only one of the possible values of x will be
reasonable.
Determine which one is reasonable
substitute possible values of x into the
algebraic expression used to represent
the equilibrium concentration of
reactants and/or products.
Applications of Equilibrium Constants
Applications of Equilibrium Constants
Applications of Equilibrium Constants
 So:
[PCl5]equil =
[PCl3]equil =
[Cl2]equil =
 You can verify your answer by substituting
the concentrations into the expression for
Kc.
You should get the same (or very close to)
the value given for Kc.
Le Chatelier’s Principle
 Equilbrium reactions such as the one to form
NH3 tend to stop short of the maximum
(theoretical) yield of products.
 Industrial chemists are always looking for
ways to improve the yield of products in a
particular reaction.
Improves cost effectiveness of process
Increases profits for the company
Reduces the cost for the consumer
Le Chatelier’s Principle
 Le Chatelier’s Principle explains the way a
system at equilibrium will change in
response to changes made in the
temperature, pressure or concentration of
one of the components of a system at
equilibrium.
Le Chatelier’s Principle
 Le Chatelier’s Principle
If a system at equilibrium is disturbed
by a change in temperature, pressure,
or the concentration of one of the
components, the system will shift its
equilibrium position so as to counteract
the effect of the disturbance.
Le Chatelier’s Principle
 How does changing the concentration of a
reactant or product impact the equilibrium?
Le Chatelier’s Principle
 In order to visualize the impact of changing
the concentration of a reactant or product,
consider how adding or removing weight
from one side of a balanced teeter totter
impacts the balance.
Balanced:
At “equilibrium”
Le Chatelier’s Principle
 If you remove two blocks from the right side
of the teeter totter, what happens?
 What do you have to do to re-balance the
teeter totter (re-establish the equilibrium)?
Le Chatelier’s Principle
 To re-balance the teeter totter (re-establish
the equilibrium), you must move one of the
blocks from the left side over to the right
side.
Le Chatelier’s Principle
 When material is removed from one side of
the teeter totter, the teeter totter is no longer
balanced:
To re-balance:
shift material towards the side where it
was removed.
Le Chatelier’s Principle
 What happens if you add 2 blocks to the
right side of the original teeter totter?
 What do you have to do to re-balance (reestablish the equilibrium) the teeter totter?
Le Chatelier’s Principle
 To re-balance the teeter totter (re-establish
the equilibrium), you must move one of the
blocks from the right side to the left side.
Le Chatelier’s Principle
 When material is added to one side of the
teeter totter, the teeter totter is no longer
balanced.
To re-balance:
shift material away from where it was
added
Le Chatelier’s Principle
 Similar trends hold true for chemical
reactions at equilibrium:
If a reactant or product is added to a
system at equilibrium, the system will shift
away from the material added.
Add reactant
shift toward products
Add product
shift toward reactants
Le Chatelier’s Principle
 If a reactant or product is removed from a
system at equilibrium, the system will shift
toward the material removed.
 Remove reactant
shift toward reactants
 Remove product
shift toward products
Le Chatelier’s Principle
Example: Give three ways that the total
amount of ammonia produced in the following
reaction can be increased? (i.e. How can you
shift the reaction towards the products?)
N2 (g) + 3 H2 (g)
2 NH3 (g)
Le Chatelier’s Principle
 Effects of Changing Pressure and Volume:
If Volume decreases, partial pressures of
the reactants and products increase:
system shifts to reduce pressure
Le Chatelier’s Principle
 Reducing the volume (thereby increasing the
partial pressures) of a gaseous system at
equilibrium causes the reaction to shift in
the direction that reduces the total number
of moles of gas
 Increasing the volume (thereby decreasing
the partial pressure) of a gaseous
equilibrium mixture, cases a shift in the
direction that produces more gas molecules.
Le Chatelier’s Principle
 Decrease Volume
shift towards fewer gas molecules
 Increase Partial Pressure
shift towards fewer gas molecules
 Increase Volume:
shift towards more gas molecules
 Decrease Partial Pressure:
shift towards more gas molecules
Le Chatelier’s Principle
Example: What happens to the system
below if the total pressure is increased by
reducing the volume?
N2 (g) + 3 H2 (g)
2 NH3 (g)
Le Chatelier’s Principle
 Effect of Temperature Change:
The value of an equilibrium constant
depends on temperature.
The impact of increasing the temperature
of a reaction depends on whether the
reaction is exothermic or endothermic.
Le Chatelier’s Principle
 To understand the impact of increasing
temperature, consider heat to be a reactant
(endothermic) or product (exothermic).
 Endothermic Reactions: absorb heat
Reactants + heat
Products
 Exothermic Reactions: produce heat
Reactants
Products + heat
Le Chatelier’s Principle
 When the temperature is increased, the
equilibrium shifts in the direction that
absorbs heat (i.e uses up the heat).
Endothermic Rxn:
Increase T
shift towards products
Exothermic Rxn:
Increase T
shift towards reactants
Le Chatelier’s Principle
 When the temperature is decreased the
equilibrium shifts in the direction that
produces heat.
Endothermic Rxn:
As T decreases
reactants
Exothermic Rxn:
As T decreases
products
shift towards
shift towards
Le Chatelier’s Principle
Example: Consider the following
equilibrium: N2O4 (g)
2 NO2 DH = 58 kJ.
In what direction will the equilibrium shift if:
N2O4 is added:
 NO2 is removed:
 total pressure is increased
by adding N2 (g):
 volume is increased
 temperature is decreased?
Le Chatelier’s Principle
 Effect of Catalyst
Addition of a catalyst does not change the
equilibrium.
Addition of a catalyst simply increases the
rate at which equilibrium is attained by
reducing the activation energy for the
reaction.