Chemical Equilibrium

1. Chemical Equilibrium Chapter 14

2. Reversibility of Reactions • Reversible reaction - Reaction that can occur in both the forward and reverse direction. • Irreversible reaction - A reaction which proceeds only in the forward direction.

3. Equilibrium • A state that is reached when the concentrations of reactants and products remain constant.

7. Le Chatelier’s Principle • If a stress is applied to a reaction mixture at equilibrium, the reaction occurs in the direction that relieves stress.

8. 3 O2(g) + heat ↔ 2 O3(g) • What happens if we • 1. Raise the temperature? • 2. Remove oxygen gas? • 3. Add ozone? • 4. Increase the volume?

9. N2 + 3 H2↔ 2 NH3

10. Equilibrium Constant Expressions • The relative amounts of all components in an equilibrium reaction will equal a constant ratio. This ratio is called an equilibrium constant. • aA + bB ↔ cC + dD

14. Q is a K expression with non-equilibrium concentration values.

15. Lets practice writing some K’s • 2 H2S(g)  2 H2(g) + S2(g) • 2 N2O5(g)  4 NO2(g) + O2(g) • C2H4(g) + H2(g)  C2H6(g) • N2(g) + 2 O2(g)  N2O4(g)

17. Heterogeneous Equilibria • 2 Hg2O(s) ↔ 2 Hg(l) + O2(g) • N2(g) + 2 O2(g) ↔ N2O4(g) • CO2(g) + C(s) ↔ 2 CO(g) • PCl5(s) + H2O(l) ↔ 2 HCl(g) + POCl3(g)

18. Phosgene, a toxic gas used in the synthesis of a variety of organic compounds, decomposes according to the equation • COCl2(g) <=> CO(g) + Cl2(g) • Phosgene is heated at 527oC in a reaction vessel. At equilibrium [CO] = 0.0456 M, [Cl2] = 0.0456 M, and [COCl2] = 0.449 M. Calculate Kc and Kp at 527oC.

19. The decomposition of phosphorous pentachloride occurs via the reaction • PCl5(g) <=> PCl3(g) + Cl2(g) • A sample of PCl5 is placed in a reaction vessel held at 250oC at an initial concentration of 1.10 M. When equilibrium is attained, the concentration of PCl5 is 0.33M. Calculate Kc.

20. Sulfur trioxide decomposes at high temperature in a sealed container: • 2 SO3(g) <==> 2 SO2(g) + O2(g) • Initially the vessel is charged at 1000K with SO3(g) at a concentrations of 6.09 x 103M. At equilibrium, the SO3 concentration is 2.44 x 103M. Calculate the value for Kc at 1000K.

21. Nitrogen dioxide decomposes at high temperature according to the reaction • 2 NO2(g) <=> 2 NO(g) + O2(g) • Suppose initially we have pure NO2(g) at 1000K and 0.500 atm. If the total pressure is 0.732 atm. when equilibrium is reached, what is the value of Kp?

22. Gaseous hydrogen iodide is placed in a closed container at a concentration of 4.49 x 103 M. At 425oC, the HI partially decomposes to hydrogen and iodine • 2 HI(g) <==> H2(g) + I2(g) • At equilibrium 21.3% of the HI has decomposed, calculate the value of Kc for the reaction.

23. Determining equilibrium concentrations • For A B

24. Given S2(g) + C(s) <=> CS2(g) Kc = 9.40 at 900K • If the initial concentration of S2 is 2.00M, how many grams of CS2 will be present in a 5.00 L flask at equilibrium?

25. At 1000oC, Kp is 0.263 atm-1 for the reaction • C(s) + 2 H2(g) <=> CH4(g) • If the initial methane concentration is 0.0625M, what is the equilibrium pressure of methane?

26. The equilibrium constant for the reaction • 2 ICl(g) <=> I2(g) + Cl2(g) is Kc = 0.11 • Calculate the equilibrium concentrations of ICl, I2, and Cl2, when 0.33 mol of I2 and 0.33 mol of Cl2 are added to a 1.5L reaction vessel.

27. Given 2 NaHCO3(s) <=> Na2CO3(s) + CO2(g) + H2O(g) with Kp = 0.25 atm2 at 125oC. Calculate the equilibrium pressures of CO2 and H2O in a closed container of baking soda (sodium bicarbonate) at 125oC.

28. Properties of equilibrium constant expressions • kforward = 1/kreverse • ksum of reactions • = kequation1 * kequation2 * kequation3 * ...

29. C(s) + H2O(g) <=> CO(g) + H2(g) • If we increase the pressure, in which direction does the reaction proceed? •  If we decrease the volume, in which direction does the reaction proceed? •  If we increase the amount of CO, in which direction does the reaction proceed? •  How do the concentrations of the following compounds change?   CO? H2? H2O?

30. CH4(g) + C(s) + O2(g)2 H2CO(g) + heat

31. CH4(g) + C(s) + O2(g)2 H2CO(g) + heat

32. CH4(g) + C(s) + O2(g)2 H2CO(g) + heat

33. CH4(g) + C(s) + O2(g)2 H2CO(g) + heat

34. CH4(g) + C(s) + O2(g)2 H2CO(g) + heat

35. Heat + 2 CO(g)  C(s) + CO2(g)

36. Heat + 2 CO(g)  C(s) + CO2(g)

37. Heat + 2 CO(g)  C(s) + CO2(g)

38. Heat + 2 CO(g)  C(s) + CO2(g)

39. Heat + 2 CO(g)  C(s) + CO2(g)

40. Heat + 2 CO(g)  C(s) + CO2(g)

41. If we put 5.0 mol of CO and 2.5 mol of Cl2 into a 10L container, how many moles of each gas will be present at equilibrium? • CO(g) + Cl2(g) <=> COCl2(g) • Kc = 4.0 M-1

42. Given • 3 O2(g) <=> 2 O3(g) • Kp = 1.1 x 10-50 atm-1. • If .211 atm of O2 are introduced into a reaction vessel, calculate the final concentrations of all species.

43. Given • 2 NH3(g) <=> N2(g) + 3H2(g) • Kp = 2.7 x 106 atm2 . • If the initial pressures of the reactants are NH3 = 2.7 atm, N2 = 3.0 atm, and H2 = 4.6 atm. What are the equilibrium pressures of all species?

44. Given the reaction • A + 2B <=> 2C + 2D • with a Kc = 0.0111M at a certain temperature. Calculate the equilibrium concentrations of all species when 3.78 moles of A and 2.15 moles of B are placed in a 1 L sealed container at that temperature.

46. Special Types of equilibrium constants • Solubility constant Ksp • Ca(OH)2 Ca+2 + 2 OH-1 Ksp = [Ca+2][OH-1]2 • Water ionization Kw • 2H2O H3O++OH-1 Kw = [H3O+][OH-] • Acid ionization Ka • HCN+H2O  H3O++CN-1 Ka= [H3O+][CN-]/[HCN] • Base ionization Kb • NH3+H2ONH4++OH- Kb=[NH4+][OH-]/[NH3]