AP Chemistry
Name _______________________________
Chapter 15 Review
Section 15.2 – The Equilibrium Constant
1. Write the equilibrium expression (K =) for each of the following conditions:
a. 2 H2S(g) ↔ 2 H2(g) + S2(g)
b. 4 NO2(g) + 6 H2O(l) ↔ 4 NH3(g) + 7 O2(g)
c. Ag+(aq) + I-(aq) ↔ AgI(s)
2. Calculate the equilibrium constant, Kp, at 25°C for the reaction:
2 NO(g) + O2(g) ↔ 2 NO2(g)
If the equilibrium pressures are NO2 = 0.55 atm, NO = 6.5 x 10-5 atm and O2 = 4.5 x 10-5 atm.
3. Calculate the equilibrium constant, Kc, for the reaction:
3 H2(g) + N2(g) ↔ 2 NH3(g)
If the equilibrium concentrations are [H2] = 0.85 M, [N2] = 1.33 M and [NH3] = 0.22 M.
4. The dissociation of acetic acid, CH3COOH, has an equilibrium constant of 1.8 x 10-5.
CH3COOH(aq) ↔ CH3COO-(aq) + H+(aq)
If the equilibrium concentration of acetic acid is 0.46 mol in 0.500 L of water and that of
CH3COO- is 8.1 x 10-3 mol in the same 0.500 L, calculate [H+] at equilibrium for the reaction.
5. At 700 K (and at equilibrium), the measured values for the partial pressures of ammonia, hydrogen
and nitrogen are 0.400 atm, 7.20 atm and 2.40 atm, respectively. Calculate Kp and Kc.
N2(g) + 3 H2(g) ↔ 2 NH3(g)
6. Calculate the value for Kp at 25°C if the value for Kc is 3.7 x 109 for the reaction
CO(g) + Cl2(g) ↔ COCl2(g)
Section 15.3 – Understanding and Working with Equilibrium Constants
7. If K > 1, which side of the reaction is favored – the reactants or products?
8. The equilibrium constant at a certain temperature for 2 HI(g) ↔ H2(g) + I2(g) is 1.39 x 10-2.
Calculate the equilibrium constant for the reverse reaction at that same temperature.
9. Using your answer from #2, calculate the value for K for each of the following reactions:
a. ½ NO(g) + ¼ O2(g) ↔ ½ NO2(g)
b. 2 NO2(g) ↔ 2 NO(g) + O2(g)
10. If you were adding the two equations in #9, what would the resulting K value be for the net
equation?
Section 15.5 – Calculating Equilibrium Constants
11. Given the initial partial pressures of PCl5 = 0.050 atm, PCl3 = 0.15 atm and Cl2 = 0.25 atm for the
following reaction, what must each equilibrium partial pressure be?
PCl5(g) ↔ PCl3(g) + Cl2(g)
Kp = 2.15
12. The reaction 2 NO(g) ↔ N2(g) + O2(g) has a Kc = 2.4 x 103 at 2000 K. If 0.61 g of NO are
initially put into a 3.00 L vessel, calculate the equilibrium concentrations of NO, N2 and O2.
13. The following reaction has an equilibrium constant of 6.2 x 102 at a certain temperature. Calculate
the equilibrium concentrations of all species if 4.5 mol of each component were initially added to
a 3.0 L flask.
H2(g) + F2(g) ↔ 2 HF(g)
14. Calculate [NH3], [NH4+] and [OH-] in a solution initially 0.200 M in NH3.
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)
Kc = 1.8 x 10-5
Section 15.6 – Applications of Equilibrium Constants
15. Determine what the system will do to reach equilibrium given the following values:
a. K = 2.9 x 102 and Q = 3.1 x 101
b. K = 0.621 and Q = 6.21 x 10-1
c. K = 7.3 x 102 and Q = 8.2 x 102
16. Calculate the reaction quotient if the partial pressure of N2O4 = 0.048 atm and the partial pressure
of NO2 = 0.056 atm.
N2O4(g) ↔ 2 NO2(g)
If Kp = 0.133 atm, which direction will the reaction shift to reach equilibrium?
Section 15.7 – Le Chatelier’s Principle
17. The following reaction is at equilibrium.
2 H2(g) + O2(g) ↔ 2 H2O(g)
How would the reaction respond if the pressure were increased at a constant temperature?
18. Using the following reaction and Appendix C, calculate ∆H for the reaction.
CH4(g) + 4 Cl2(g) ↔ CCl4(g) + 4 HCl(g)
19. Using your answer to #18, predict the effect of the following changes to the system on the
direction of equilibrium:
a. Pressure is doubled
b. CCl4 is removed as it is generated
c. Heat is added to the system