applications of equilibrium constants

Applications of Equilibrium Constants

Kc and Kp can be used to determine the concentration of reactants and/or products at equilibrium.


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.


CaF2 (s) Ca2+ (aq) + 2 F- (aq)


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:


Set up a table showing the initial concentrations, changes in concentration, and equilibrium concentrations.

Initial 0.0400 M 0.000 M 0.000 M

Change

Equil.

PCl5 (g) PCl3 (g) + Cl2 (g)


Substitute the equilibrium values into the expression for Kc


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.


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 processIncreases profits for the companyReduces the cost for the consumer


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 PrincipleIf 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.


How does changing the concentration of a reactant or product impact the equilibrium?


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”


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)?


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.


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.


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 (re-establish the equilibrium) the teeter totter?


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.


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


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


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


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)


Effects of Changing Pressure and Volume:

If Volume decreases, partial pressures of the reactants and products increase:

system shifts to reduce pressure


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.


Decrease Volumeshift towards fewer gas molecules

Increase Partial Pressureshift towards fewer gas molecules

Increase Volume:shift towards more gas molecules

Decrease Partial Pressure:shift towards more gas molecules


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)


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.


To understand the impact of increasing temperature, consider heat to be a reactant (endothermic) or product (exothermic).

Endothermic Reactions: absorb heatReactants + heat Products

Exothermic Reactions: produce heatReactants Products + heat


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


When the temperature is decreased the equilibrium shifts in the direction that produces heat.

Endothermic Rxn:As T decreases shift towards

reactantsExothermic Rxn:

As T decreases shift towards products


Example: Consider the following equilibrium: N2O4 (g) 2 NO2 H = 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?


Effect of CatalystAddition 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.

applications of equilibrium constants

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