1. Chapter 15
Chemical equilibria

2. Chemical Equilibrium Closed vial of NO2: NO2(g) + NO2(g)  N2O4(g)
brown colorless
Brown color becomes less intense, then, after some time period, color change stops.
Concentrations of products and reactants remain constant with time (equilibrium)
Applies to all reactions

3. Dynamic Equilibrium H2O(g) + CO(g)  H2(g) + CO2(g)
Start off with closed flask of CO and H2O. Equilibrium sets up (chemical reaction occurring).
What happens to amounts of H2O(g) + CO(g) ?
What happens to amounts of H2(g) + CO2(g) ?
Reaction still occurring past dotted line (although concentrations don’t change): Dynamic Equilibrium

4. Dynamic Equilibrium
H2O(g) + CO(g)  H2(g) + CO2(g)

5. Equilibrium Expression and K
Law of mass action
jA + kB  lC + mD
K = [C]l [D]m Equilibrium Expression
[A]j [B]k
Concentrations of species at equilibrium
K = equilibrium constant
H2O(g) + CO(g)  H2(g) + CO2(g)
2NO2(g)  N2O4(g)

6. Value of Equilibrium Constant
3H2(g) + N2(g)  2NH3(g)
Value for K is always the same for a particular reaction,

7. Magnitude of K
Equilibrium position (does equilibrium lie towards products or reactants?); is this important (Synthesis of Aspirin)
Large K
K1
Small K

8. Writing equilibrium expressions
Always predicted from the balanced equation:
4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g)
I2(g) + H2(g)  2HI(g)
C4H10(g) + O2(g)  CO2(g) + H2O(g)

9. Calculating values of K

10. Eq. laws for gaseous reactions
3H2(g) + N2(g)  2NH3(g)
K = [NH3]2 Equilibrium constant found in terms of
[H2]3 [N2] concentrations of species (Kc)
K can also be found in terms of partial pressures (Kp)
PV = nRT
Kp = PNH32
(PH23)(PN2)
Relationship between Kp and Kc:
Kp = Kc(RT)ng

11. Heterogeneous Equilibria
More than one phase exists in reaction mixture
Thermal decomposition of CaCO3(s)  CaO(s) + CO2(g)
Concentrations of pure solids and pure liquids are always constant (thus, can be removed from eq. expression)

12. Heterogeneous Equilibria
Write equilibrium expressions (in terms of both K and Kp for each of the following):
CaO(s) + SO2(g)  CaSO3(s)
Decomposition of solid phosphorous pentachloride to liquid phosphorous trichloride and chlorine gas
Deep blue solid copper(II)sulfate pentahydrate is heated to drive off water vapor to form white solid copper(II)sulfate

13. Reaction Quotient, Q
In which direction will a particular reaction shift to reach equilibrium?
3H2(g) + N2(g)  2NH3(g)
If [NH3]0 = 0, shift to right to achieve equilibrium.
If [H2]0 or [N2]0 = 0, shift to left to achieve equilibrium.
If initial concentrations of all three species are nonzero, which way will shift occur to achieve equilibrium? More difficult to predict.
Use Reaction Quotient, Q
Q is obtained by applying law of Mass Action to initial concentrations of species involved.

14. Reaction Quotient, Q 3H2(g) + N2(g)  2NH3(g) Q = [NH3]o2 [H2]o3 [N2]o
Then, to determine in which direction a system will shift to reach equilibrium, compare values of Q and K:
Q = K System is at equilibrium
Q > K Ratio of initial conc. of products to initial conc. of reactants is too large. System shifts to the left to reach equilibrium.
Q < K Ratio of initial conc. of products to initial conc. of reactants is too small. System shifts to the right to reach equilibrium

15. Calculating Equilibrium Concentrations
‘ICE’ Tables

16. Le Chatelier’s Principle
‘If an outside influence upsets an equilibrium, the system undergoes a change in the direction that counteracts the disturbing influence, and, if possible, returns the system to equilibrium.’
Outside influences?
Adding / removing a reactant / product
Changing volume/pressure of gaseous reactions
Changing T

17. Adding/removing a reactant or product
Eq. shifts in direction that will partially consume a reactant or product added
Eq. shifts in direction that will partially replace a reactant or product removed
3H2(g) + N2(g)  2NH3(g
Add some N2 to above equilibrium; which way will it shift to re-establish equilibrium?

18. Changing volume/pressure
Reducing volume (what happens to the pressure?) – eq. shifts to side with smaller # of gas molecules
3H2(g) + N2(g)  2NH3(g
Increasing volume? What would happen here?

19. Changing volume/pressure
2NO2(g)  N2O4(g)
brown colorless

20. Changing T
Increasing T shifts an equilibrium in direction that produces an endothermic change (need to know energy involved in reaction to predict which direction this is)
3H2(g) + N2(g)  2NH3(g) H= kJ/mol
Exothermic or endothermic reaction?
Increasing T:
Decreasing T:

21. Changing T Is this reaction exo- or endothermic? 2NO2(g)  N2O4(g)
brown colorless
Is this reaction exo- or endothermic?