1. Equilibrium © 2009, Prentice-Hall, Inc. Chapter 7 Chemical Equilibrium Dr Imededdine Arbi Nehdi King Saud University Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten

2. Equilibrium © 2009, Prentice-Hall, Inc. The Concept of Equilibrium Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.

3. Equilibrium © 2009, Prentice-Hall, Inc. 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.

4. Equilibrium © 2009, Prentice-Hall, Inc. A System at Equilibrium Once equilibrium is achieved, the amount of each reactant and product remains constant.

5. Equilibrium © 2009, Prentice-Hall, Inc. Depicting Equilibrium Since, in a system at equilibrium, both the forward and reverse reactions are being carried out, we write its equation with a double arrow. N 2 O 4 (g) 2 NO 2 (g)

6. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant

7. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant Forward reaction: N 2 O 4 (g)  2 NO 2 (g) Rate Law: Rate = k f [N 2 O 4 ]

8. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant Reverse reaction: 2 NO 2 (g)  N 2 O 4 (g) Rate Law: Rate = k r [NO 2 ] 2

9. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant Therefore, at equilibrium Rate f = Rate r k f [N 2 O 4 ] = k r [NO 2 ] 2 Rewriting this, it becomes kfkrkfkr [NO 2 ] 2 [N 2 O 4 ] =

10. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant The ratio of the rate constants is a constant at that temperature, and the expression becomes K eq = kfkrkfkr [NO 2 ] 2 [N 2 O 4 ] =

11. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant Consider the generalized reaction The equilibrium expression for this reaction would be K c = [C] c [D] d [A] a [B] b aA + bBcC + dD

12. Equilibrium © 2009, Prentice-Hall, Inc. The Equilibrium Constant Since pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written K p = (P C c ) (P D d ) (P A a ) (P B b )

13. Equilibrium © 2009, Prentice-Hall, Inc. Relationship Between K c and K p From the Ideal Gas Law we know that Rearranging it, we get PV = nRT P = RT nVnV

14. Equilibrium © 2009, Prentice-Hall, Inc. Relationship Between K c and K p Plugging this into the expression for K p for each substance, the relationship between K c and K p becomes where K p = K c (RT)  n  n = (moles of gaseous product) - (moles of gaseous reactant)

15. Equilibrium © 2009, Prentice-Hall, Inc. Equilibrium Can Be Reached from Either Direction As you can see, the ratio of [NO 2 ] 2 to [N 2 O 4 ] remains constant at this temperature no matter what the initial concentrations of NO 2 and N 2 O 4 are.

16. Equilibrium © 2009, Prentice-Hall, Inc. Equilibrium Can Be Reached from Either Direction This is the data from the last two trials from the table on the previous slide.

17. Equilibrium © 2009, Prentice-Hall, Inc. Equilibrium Can Be Reached from Either Direction It doesn’t matter whether we start with N 2 and H 2 or whether we start with NH 3 : we will have the same proportions of all three substances at equilibrium.

18. Equilibrium © 2009, Prentice-Hall, Inc. What Does the Value of K Mean? If K>>1, the reaction is product-favored; product predominates at equilibrium.

19. Equilibrium © 2009, Prentice-Hall, Inc. What Does the Value of K Mean? If K>>1, the reaction is product-favored; product predominates at equilibrium. If K<<1, the reaction is reactant-favored; reactant predominates at equilibrium.

20. Equilibrium © 2009, Prentice-Hall, Inc. Manipulating Equilibrium Constants The equilibrium constant of a reaction in the reverse reaction is the reciprocal of the equilibrium constant of the forward reaction. K c = = 0.212 at 100  C [NO 2 ] 2 [N 2 O 4 ] N2O4 (g)N2O4 (g) 2 NO 2 (g) K c = = 4.72 at 100  C [N 2 O 4 ] [NO 2 ] 2 N2O4 (g)N2O4 (g) 2 NO 2 (g)

21. Equilibrium © 2009, Prentice-Hall, Inc. Manipulating Equilibrium Constants The equilibrium constant of a reaction that has been multiplied by a number is the equilibrium constant raised to a power that is equal to that number. K c = = 0.212 at 100  C [NO 2 ] 2 [N 2 O 4 ] N 2 O 4(g) 2 NO 2(g) K c = = (0.212) 2 at 100  C [NO 2 ] 4 [N 2 O 4 ] 2 2 N 2 O 4(g) 4 NO 2(g)

22. Equilibrium © 2009, Prentice-Hall, Inc. Manipulating Equilibrium Constants The equilibrium constant for a net reaction made up of two or more steps is the product of the equilibrium constants for the individual steps.

23. Equilibrium © 2009, Prentice-Hall, Inc. Heterogeneous Equilibrium

24. Equilibrium © 2009, Prentice-Hall, Inc. The Concentrations of Solids and Liquids Are Essentially Constant Both can be obtained by multiplying the density of the substance by its molar mass — and both of these are constants at constant temperature.

25. Equilibrium © 2009, Prentice-Hall, Inc. The Concentrations of Solids and Liquids Are Essentially Constant Therefore, the concentrations of solids and liquids do not appear in the equilibrium expression. K c = [Pb 2+ ] [Cl - ] 2 PbCl 2 (s) Pb 2+ (aq) + 2 Cl - (aq)

26. Equilibrium © 2009, Prentice-Hall, Inc. As long as some CaCO 3 or CaO remain in the system, the amount of CO 2 above the solid will remain the same. CaCO 3 (s) CO 2 (g) + CaO (s) K P = P CO2

27. Equilibrium © 2009, Prentice-Hall, Inc. Equilibrium Calculations

28. Equilibrium © 2009, Prentice-Hall, Inc. An Equilibrium Problem A closed system initially containing 1.000 x 10 -3 M H 2 and 2.000 x 10 -3 M I 2 at 448  C is allowed to reach equilibrium. Analysis of the equilibrium mixture shows that the concentration of HI is 1.87 x 10 -3 M. Calculate K c at 448  C for the reaction taking place, which is H 2 (g) + I 2 (g) 2 HI (g)

29. Equilibrium © 2009, Prentice-Hall, Inc. What Do We Know? [H 2 ], M[I 2 ], M[HI], M Initially1.000 x 10 -3 2.000 x 10 -3 0 Change At equilibrium1.87 x 10 -3

30. Equilibrium © 2009, Prentice-Hall, Inc. [HI] Increases by 1.87 x 10 -3 M [H 2 ], M[I 2 ], M[HI], M Initially1.000 x 10 -3 2.000 x 10 -3 0 Change+1.87 x 10 -3 At equilibrium1.87 x 10 -3

31. Equilibrium © 2009, Prentice-Hall, Inc. Stoichiometry tells us [H 2 ] and [I 2 ] decrease by half as much. [H 2 ], M[I 2 ], M[HI], M Initially1.000 x 10 -3 2.000 x 10 -3 0 Change-9.35 x 10 -4 +1.87 x 10 -3 At equilibrium1.87 x 10 -3

32. Equilibrium © 2009, Prentice-Hall, Inc. We can now calculate the equilibrium concentrations of all three compounds… [H 2 ], M[I 2 ], M[HI], M Initially1.000 x 10 -3 2.000 x 10 -3 0 Change-9.35 x 10 -4 +1.87 x 10 -3 At equilibrium6.5 x 10 -5 1.065 x 10 -3 1.87 x 10 -3

33. Equilibrium © 2009, Prentice-Hall, Inc. …and, therefore, the equilibrium constant. K c = [HI] 2 [H 2 ] [I 2 ] = 51 = (1.87 x 10 -3 ) 2 (6.5 x 10 -5 )(1.065 x 10 -3 ) aA + bBcC + dD n a /a = n b /b = n c /c = n d /d

34. Equilibrium © 2009, Prentice-Hall, Inc. The Reaction Quotient (Q) Q gives the same ratio the equilibrium expression gives, but for a system that is not at equilibrium. To calculate Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression.

35. Equilibrium © 2009, Prentice-Hall, Inc. If Q = K, the system is at equilibrium.

36. Equilibrium © 2009, Prentice-Hall, Inc. If Q > K, there is too much product, and the equilibrium shifts to the left.

37. Equilibrium © 2009, Prentice-Hall, Inc. If Q < K, there is too much reactant, and the equilibrium shifts to the right.

38. Equilibrium © 2009, Prentice-Hall, Inc. Le Châtelier’s Principle

39. Equilibrium © 2009, Prentice-Hall, Inc. 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 equilibrium position so as to counteract the effect of the disturbance.”

40. Equilibrium © 2009, Prentice-Hall, Inc. Change in Reactant or Product Concentrations “If a chemical system is at equilibrium and we add a substance (either a reactant or a product), the reaction system will shift so as to reestablish equilibrium by consuming part of the added substance. Conversely, removal of a substance will result in the reaction moving in the direction that forms more of this substance

41. Equilibrium © 2009, Prentice-Hall, Inc. Example: The Haber Process The transformation of nitrogen and hydrogen into ammonia (NH 3 ) is of tremendous significance in agriculture, where ammonia-based fertilizers are of utmost importance.

42. Equilibrium © 2009, Prentice-Hall, Inc. The Haber Process If H 2 is added to the system, N 2 will be consumed and the two reagents will form more NH 3.

43. Equilibrium © 2009, Prentice-Hall, Inc. The Haber Process This apparatus helps push the equilibrium to the right by removing the ammonia (NH 3 ) from the system as a liquid.

44. Equilibrium © 2009, Prentice-Hall, Inc. Effect of Temperature Change “If a chemical system is at equilibrium and when the temperature is increased, the reaction system will shift in the direction to absorbs heat ( endothermic reaction). Conversely, decreasing temperature, the system will move in the direction of the exothermic reaction

45. Equilibrium © 2009, Prentice-Hall, Inc. Example: The Effect of Changes in Temperature Co(H 2 O) 6 2+ (aq) + 4 Cl - (aq) (CoCl 4 ) 2- (aq) + 6 H 2 O (l) Forward reaction : endothermic Reverse reaction: exothermic Heating the solution Cooling the solution

46. Equilibrium © 2009, Prentice-Hall, Inc. Effect of Volume and Pressure Changes “At constant temperature, reducing the volume (increasing the total pressure) of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas. Conversely, increasing the volume (decreasing the total pressure) causes a shift in in the direction that produces more gas molecules.

47. Equilibrium © 2009, Prentice-Hall, Inc. Example: Effect of Volume and Pressure Changes What happens if the total pressure of this equilibrium mixture is incresed by decresing the volume ! According to Le Chatelier principle, the equilibrum will shifts to the side that reduces the total number of moles of gas,which is the reactant side in this case.