1. Chemical Equilibrium
Chapter 18

2. Equilibrium Few chemical reactions proceed in only one direction
Most reactions are reversible
N2 + 3H2  2NH3
N2 + 3H2  2NH3
At first the reaction proceeds towards the formation of product
As soon as some product is formed the reverse reaction begins (reactants are produced)
N2 + 3H2 ⇌ 2NH3

3. Chemical Equilibrium Chemical equilibrium is reached when:
the rates of the forward and reverse reactions are equal
the concentrations of the reactants and products remain constant
Equilibrium is a state in which there are no observable changes as time goes by

4. Physical equilibrium Chemical equilibrium
H2O (l) ⇌
H2O (g)
Physical equilibrium
The changes that occur are physical processes
colorless dark brown
N2O4 (g) ⇌
2NO2 (g)
Chemical equilibrium
The changes that occur are chemical

6. N2O4 (g) ⇌ 2NO2 (g)
equilibrium

7. Calculating Equilibrium Constant K
For the generic reaction: aA + bB ⇌ cC + dD
A, B, C, D are products and reactants
a, b, c, d are the coefficients
Kc = [Product C ]c [Product D]d
[Reactant A ]a [Reactant B]b
Kc subscript c indicates the concentration expressed in molarity (M)
[ ] mean concentration (M)
K has no units
How do we use K?
If K>1 - More products than reactants at equilibrium
If K <1 - More reactants than products at equilibrium
equilibrium expression

8. N2O4 (g) ⇌ 2NO2 (g)
Equilibrium
constant
(Kc)
14.1

9. Calculating Kc for a Homogenous Equilibrium
In a homogenous equilibrium reactants and products are all in the same state (gas, liquid, solid):
N2(g) + 3H2(g) ⇌ 2 NH3 (g) -all gases
For the above reaction write the equilibrium expression and calculate the equilibrium constant (Kc) using the following information:
[NH3] = M
[N2] = M
[H2] = M
Are there more products or reactants at equilibrium?

10. Calculating KP
K can also be calculated using the partial pressures of the reactants and products (only works for gaseous reactions)
This is because at a constant temperature pressure of a gas is directly related to its concentration in mol/L:
PV = nRT P= nRT V
aA(g) ⇌ bB(g)
KP = PBb
PAa

11. Calculating KP aA(g) ⇌ bB(g) KP = PBb PAa N2O4(g) ⇌ 2NO2(g) KP = PNO22

12. ∆n = moles of gaseous products – moles of gaseous reactants
Kc and Kp
In general Kc and Kp are not equal
The relationship between Kc and Kp:
Kp = Kc(RT)∆n or
Kp = Kc(0.0821T)∆n
∆n = moles of gaseous products – moles of gaseous reactants
= (c + d) – (a + b) or ∑ product coeff. – ∑ reactant coeff.
when ∆n = 0, Kp = Kc

13. Calculating K for a Heterogeneous Equilibrium
In a heterogeneous equilibrium reactants and products are NOT in the same state:
2NaHCO3(s) ⇌ NaCO3(s) + CO2(g) + H2O(g)
Solids and liquids do not need to be included in equilibrium expressions and/or Kc calculations (Kp too)
For the above reaction write the equilibrium expression and calculate the equilibrium constant (Kc) using the following information:
[CO2] = M
[H2O] = M
Are there more products or reactants at equilibrium?

14. Practice
The equilibrium concentrations for the reaction between carbon monoxide and molecular chlorine to form COCl2 (g) at 740C are [CO] = M, [Cl2] = M, and [COCl2] = 0.14 M. Calculate the equilibrium constants Kc and Kp.
The equilibrium constant Kp for the following reaction is 158 at 1000K. What is the equilibrium pressure of O2 if the PNO2 = atm and PNO = atm
2NO2 (g) ⇌ 2NO (g) + O2 (g)
Consider the equilibrium below at 295 K. The partial pressure of each gas is atm. Calculate Kp and Kc for the reaction.
NH4HS (s) ⇌ NH3 (g) + H2S (g)

15. Calculating Equilibrium Concentrations using ICE
Initially, 0.40 mol of nitrogen, and 0.96 mole of hydrogen are placed in a 2.0 L container at a constant temperature. The mixture is allowed to react and at equilibrium, the molar concentration of ammonia is found to be 0.14 M. Calculate the equilibrium constant, Kc, for the reaction.
N2(g) + 3H2 (g) ⇌ 2NH3(g)
the Kc for the following system is 24.0 at 200oC. Initially only cis-stilbene is present at mol/L. Calculate the concentrations of cis-stilbene and trans-stilbene
cis-stilbene ⇌ trans-stilbene
A mixture of mol H2 and mole I2 as placed in a 1.00-L stainless-steel flask at 430oC. The equilibrium constant Kc for the reaction is 54.3 at this temperature. Calculate the concentrations of H2, I2, and HI at equilibrium.
H2 + I2 ⇌ 2HI

16. Le Châtelier’s Principle
If a stress is applied to a system at equilibrium, the equilibrium will react (or “shift”) to relieve the stress and reestablish equilibrium
“Stresses” include changes in:
Concentration
Temperature
Volume (pressure)

17. Changes in Concentration
If product or reactant is added to a system at equilibrium the system will react by shifting in the direction opposite the addition
If product or reactant is removed the system shifts in the direction of the removal
The system reacts to restore the concentration of the substance that is removed
Change
Shift in equilibrium
Increase concentration of reactant
Shift to the right (towards products)
Increase concentration of product
Shift to the left (towards reactants)
Remove reactant
Shift to the left
Remove product
Shift to the right

18. Changes in Volume/Pressure
Change in volume affects gaseous equilbria only
Recall:
As volume increases, pressure decreases
As volume decreases, pressure increases
A decrease in volume (pressure increases) causes the reaction to shift to the side with the fewer moles of gas
An increase in volume (pressure decreases) causes the reaction to shift to the side with more moles of gas
**There is no shift when the number of moles of gas is the same on each side

19. Changes in Temperature
Exothermic reactions:
N2(g) + 3H2(g) ⇌ 2NH3(g) + 92 kJ ∆Ho= -92kJ
Energy is released
Energy is a product
Endothermic reaction:
556kJ + CaCO3(s) ⇌ CaO(s) + CO2(g) (∆Ho= 556 kJ
Energy is absorbed
Energy is a reactant

20. Changes in Temperature
Treat energy as a reactant/product-Rules for changes in concentration apply
Increasing the temperature: Energy is being added
the system will react by shifting in the direction opposite the addition
equilibrium shifts in the direction that uses up energy (heat)
Decreasing temperature: Energy is being removed
System to shifts restore the energy lost
equilibrium shifts in the direction that creates energy (heat)

21. Changes in Temperature
Value of K changes with changes in temperature
(K does not change with changes in pressure or concentration--only position of equilibrium changes)
Exothermic reactions (energy is a product)
K decreases when temperature increases
Endothermic reactions (energy is a reactant)
K increases when temperature is increased

22. Addition of Catalyst Adding a Catalyst does not change K
does not shift the position of an equilibrium system
system will reach equilibrium sooner
uncatalyzed
catalyzed

23. Le Châtelier’s Principle in summary
Change Equilibrium
Constant
Change
Shift Equilibrium
Concentration
yes no
Pressure
yes no
Volume
yes no
Temperature
yes
yes
Catalyst no no