1. Chapter 15 Chemical Equilibrium
BLB 11th

2. 2 NO2(g) → 2 NO(g) + O2(g)

3. 15.1 The Concept of Equilibrium
N2O4(g)
2 NO2(g) ⇌
Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.

4. The Concept of Equilibrium
As a system approaches equilibrium, both the forward and reverse reactions are occurring.
At equilibrium, the forward and reverse reactions are proceeding at the same rate.

5. A System at Equilibrium
Once equilibrium is achieved, the amount of each reactant and product remains constant.
N2O4(g)
2 NO2(g) ⇌

6. The Concept of Equilibrium
The reaction system is closed.
Opposing reactions occur at equal rates, i.e. the rates of forward and reverse reactions are equal.
Dynamic process, i.e. never stops.
Concentrations of reactants and products are constant.
Equilibrium can be reached from either direction, reactants or products.

7. 15.2 The Equilibrium Constant
aA + bB ⇌ cC + dD
K is constant for a reaction, regardless of the [A]0.
K depends on stoichiometry, not on the mechanism.
Pure solid or liquid reactant & products are not included for heterogeneous equilibria. (15.4)
Water (or other pure solvent) is not included if the reactant and product concentrations are low. (p. 642)
Kc – concentrations (M) of solutions
Kp – partial pressures of gases; p. 636 to convert between Kp & Kc

8. Write the expression for K.
N2(g) + O2(g) ⇌ 2 NO(g)
2 SO2(g) + O2(g) ⇌ 2 SO3(g)

9. 15.3 Interpreting and Working with Equilibrium Constants
The Magnitude of K
K >> 1 Essentially all products; lies to right
K > 1 product-favored
K < 1 reactant-favored
K << 1 Essentially all reactants; lies to left

10. Manipulating K (pp. 638-39) Reverse reaction: Adding reactions:
Multiplying reaction by some factor:
#1: C(s) + ½ O2(g) ⇌ CO(g)
#2: 2 C(s) + O2(g) ⇌ 2 CO(g)

11. 15.4 Heterogeneous Equilibria
Reactant components are in different phases.
When a pure solid or liquid is involved in a heterogeneous equilibrium, its concentration is not included in the K expression.
Write the K expression:
SnO2(s) + 2 CO(g) ⇌ Sn(s) + 2 CO2(g)

12. 15.5 Calculating Equilibrium Constants
Must plug in the equilibrium concentrations of reactants and products
A concentration table with initial, change, and equilibrium concentrations is set up.
See Sample Exercise 15.9, pp. 643.
More in lab and in Chapter 16 with weak acids and bases.

13. Calculating K N2(g) + 3 H2(g) ⇌ 2 NH3(g) p. 642 @ 472°C
← eq. partial pressures (atm)

14. 15.6 Applications of Equilibrium Constants
Predicting the direction of a reaction
The reaction quotient, Q (p. 644)
Q has the same form as K, but for non-equilibrium conditions.
Comparing Q and K:
Q < K achieves equilibrium by shifting to right
Q = K @ equilibrium
Q > K achieves equilibrium by shifting to left

16. Sample problem 38a
N2(g) + 3 H2(g) ⇌ 2 NH3(g) KP = 4.51 x 10-5 at 450°C
45 atm 55 atm atm

17. 15.7 Le Châtelier’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 position to counteract the effect of the disturbance. (p. 648)

19. Le Châtelier’s Principle
Disturbance
Effect
More reactant or product
Shifts to consume
Volume decrease
(total pressure increase)
Shifts to reduce pressure (fewer moles)
Inert substance added
None
Temperature increase
If exothermic, shifts left and K decreases.
If endothermic, shifts right and K increases.

20. 2 SO2(g) + O2(g) ⇌ 2 SO3(g) ΔH < 0
Sample problem 51
2 SO2(g) + O2(g) ⇌ 2 SO3(g) ΔH < 0
Effect on equilibrium mixture when…
(a) O2(g) is added?
(b) the reaction is heated?
(c) the volume of container is doubled?
(d) a catalyst is added?
(e) the total P is increased by adding a noble gas?
(f) SO3(g) is removed?

21. 6 CO2(g) + 6 H2O(l) ⇌ C6H12O6(s) + 6 O2(g) ΔH = 2816 kJ
Write the K expression.
Effect on equilibrium mixture when…
(a) PCO2 is increased?
(b) the reaction is heated?
(c) some CO2(g) is removed?
(d) the total P is decreased?
(e) part of C6H12O6(s) is removed?
(f) a catalyst is added?