Equilibrium Basic Concepts

1. Equilibrium Basic Concepts • Reversiblereactions do not go to completion. • They can occur in either direction • Chemical equilibrium exists when two opposing reactions occur simultaneously at the same rate. • A chemical equilibrium is a reversible reaction that the forward reaction rate is equal to the reverse reaction rate. • Chemical equilibria are dynamic equilibria. • Molecules are continually reacting, even though the overall composition of the reaction mixture does not change.

2. The Equilibrium Constant • Kc is the equilibrium constant . • Kc is defined for a reversible reaction at a given temperature as the product of the equilibrium concentrations (in M) of the products, each raised to a power equal to its stoichiometric coefficient in the balanced equation, divided by the product of the equilibrium concentrations (in M) of the reactants, each raised to a power equal to its stoichiometric coefficient in the balanced equation. • Equilibrium constants are dimensionless because they actually involve a thermodynamic quantity called activity. • Activities are directly related to molarity

3. The Equilibrium Constant

4. Basic Concepts • One of the fundamental ideas of chemical equilibrium is that equilibrium can be established from either the forward or reverse direction. • The rates of the forward and reverse reactions can be represented as: • When system is at equilibrium: • Ratef = Rater

5. The Equilibrium Constant Write equilibrium constant expressions for the following reactions at 500oC. All reactants and products are gases at 500oC.

6. The Equilibrium Constant

7. One liter of equilibrium mixture from the following system at a high temperature was found to contain 0.172 mole of phosphorus trichloride, 0.086 mole of chlorine, and 0.028 mole of phosphorus pentachloride. Calculate Kc for the reaction. Equil []’s 0.028 M 0.172 M 0.086 M The Equilibrium Constant

8. The Equilibrium Constant • At a given temperature 0.80 mole of N2 and 0.90 mole of H2 were introduced in an evacuated 1.00-liter container after equilibrium was established the equilibrium concentrations were determined top be 0.20 mole of NH3, 0.70 moles N2, and 0.06 moles H2. Calculate Kc for the reaction.

9. The Equilibrium Constant Product-favored K > 1 Reactant-favored K < 1 Products favored103 > K > 10-3 Reactants favored

10. The value of Kc depends upon how the balanced equation is written. From the prior examples we have: PCl5 PCl3 + Cl2 Variation of Kc with the Form of the Balanced Equation Equil []’s 0.028 M 0.172 M 0.086 M Kc = [PCl3][Cl2]/[PCl5] = 0.53 PCl3 + Cl2 PCl5 Equil. []’s 0.172 M 0.086 M 0.028 M Kc’= [PCl5]/[PCl3][Cl2] = 1.89 PCl3 + Cl2 PCl5 Equil. []’s 0.172 M 0.086 M 0.028 M Kc “= [PCl5]2/[PCl3]2[Cl2]2 = 3.56 2PCl5 2PCl3 + 2Cl2 Equil []’s 0.028 M 0.172 M 0.086 M Kc ’’’= [PCl3] 2[Cl2] 2/[PCl5] 2 = 0.28

11. Partial Pressures and the Equilibrium Constant • For gas phase reactions the equilibrium constants can be expressed in partial pressures rather than concentrations. • For gases, the pressure is proportional to the concentration. • We can see this by looking at the ideal gas law. • PV = nRT • P = nRT/V • n/V = M • P= MRT and M = P/RT

12. Partial Pressures and the Equilibrium Constant • Consider this system at equilibrium at 5000C. 2SO2(g) + O2(g) 2SO3(g)

13. Kc is 49 for the following reaction at 450oC. If 1.0 mole of H2 and 1.0 mole of I2 are allowed to reach equilibrium in a 3.0-liter vessel, (a) How many moles of I2 remain unreacted at equilibrium? Relationship Between Kp and Kc (b) What are the equilibrium partial pressures of H2, I2 and HI? (c) What is the total pressure in the reaction vessel?

14. (a) How many moles of I2 remain unreacted at equilibrium?

15. (b) What are the equilibrium partial pressures of H2, I2 and HI?

16. (c) What is the total pressure in the reaction vessel?

17. Relationship Between Kp and Kc • Nitrosyl bromide, NOBr, is 34% dissociated by the following reaction at 25oC, in a vessel in which the total pressure is 0.25 atmosphere. What is the value of Kp?

18. Heterogeneous equilibria have more than one phase present. For example, a gas and a solid or a liquid and a gas. How does the equilibrium constant differ for heterogeneous equilibria? Pure solids and liquids have activities of unity. Solvents in very dilute solutions have activities that are essentially unity. The Kc and Kp for the reaction shown above are: Heterogeneous Equlibria

19. Heterogeneous Equlibria

20. Heterogeneous Equlibria • What are Kc and Kp for this reaction?

21. Heterogeneous Equlibria • What are Kc and Kp for this reaction?

22. An aside: Similar to Hess’ Law we can add several reactions together to get to an overall reaction not listed in a table. To determine K for the sum of the reactions simply multiply all the values for each intermediate step to get the K for the overall reaction.

23. Uses of the Equilibrium Constant, Kc • The equilibrium constant, Kc, is 3.00 for the following reaction at a given temperature. If 1.00 mole of SO2 and 1.00 mole of NO2 are put into an evacuated 2.00 L container and allowed to reach equilibrium, what will be the concentration of each compound at equilibrium?

24. Uses of the Equilibrium Constant, Kc • The equilibrium constant is 49 for the following reaction at 450oC. If 1.00 mole of HI is put into an evacuated 1.00-liter container and allowed to reach equilibrium, what will be the equilibrium concentration of each substance?

25. The Reaction Quotient • The mass action expression or reaction quotient has the symbol Q. • Q has the same form as Kc • The major difference between Q and Kc is that the concentrations used in Q are not necessarily equilibrium values. • Why do we need another “equilibrium constant” that does not use equilibrium concentrations? • Q will help us predict how the equilibrium will respond to an applied stress. • To make this prediction we compare Q with Kc.

26. The Reaction Quotient • The equilibrium constant for the following reaction is 49 at 450oC. If 0.22 mole of I2, 0.22 mole of H2, and 0.66 mole of HI were put into an evacuated 1.00-liter container, would the system be at equilibrium? If not, what must occur to establish equilibrium?

27. Disturbing a System at Equlibrium: Predictions • LeChatelier’s Principle - If a change of conditions (stress) is applied to a system in equilibrium, the system responds in the way that best tends to reduce the stress in reaching a new state of equilibrium. • We first encountered LeChatelier’s Principle in Chapter 14. • Some possible stresses to a system at equilibrium are: • Changes in concentration of reactants or products. • Changes in pressure or volume (for gaseous reactions) • Changes in temperature.

28. Disturbing a System at Equlibrium: Predictions • Changes in Concentration of Reactants and/or Products • Also true for changes in pressure for reactions involving gases.

29. Disturbing a System at Equlibrium: Predictions • Changes in Volume • (and pressure for reactions involving gases) • Predict what will happen if the volume of this system at equilibrium is changed by changing the pressure at constant temperature:

30. Disturbing a System at Equlibrium: Predictions • Changing the Reaction Temperature

31. Introduction of a Catalyst Catalysts decrease the activation energy of both the forward and reverse reaction equally. Catalysts do not affect the position of equilibrium. The concentrations of the products and reactants will be the same whether a catalyst is introduced or not. Equilibrium will be established faster with a catalyst. Disturbing a System at Equlibrium: Predictions

32. Disturbing a System at Equlibrium: Predictions • Given the reaction below at equilibrium in a closed container at 500oC. How would the equilibrium be influenced by the following?

33. Disturbing a System at Equlibrium: Predictions • How will an increase in pressure (caused by decreasing the volume) affect the equilibrium in each of the following reactions?

34. Disturbing a System at Equlibrium: Predictions • How will an increase in temperature affect each of the following reactions?

35. A 2.00 liter vessel in which the following system is in equilibrium contains 1.20 moles of COCl2, 0.60 moles of CO and 0.20 mole of Cl2. Calculate the equilibrium constant. Disturbing a System at Equilibrium: Calculations

36. Disturbing a System at Equilibrium: Calculations • An additional 0.80 mole of Cl2 is added to the vessel at the same temperature. Calculate the molar concentrations of CO, Cl2, and COCl2 when the new equilibrium is established.

37. The Haber Process: An Application of Equilibrium • The Haber process is used for the commercial production of ammonia. • This is an enormous industrial process in the US and many other countries. • Ammonia is the starting material for fertilizer production. Fritz Haber 1868-1934 Nobel Prize, 1918 Carl Bosch 1874-1940 Nobel Prize, 1931

38. ∆G is change in free energy at non-standard conditions. ∆G is related to ∆G˚ ∆G = ∆G˚ + RT ln Q where Q = reaction quotient When Q < K or Q > K, reaction is spontaneous. When Q = K reaction is at equilibrium When ∆G = 0 reaction is at equilibrium Therefore, ∆G˚ = - RT ln K ∆G, ∆G˚, and Keq

39. Relationship BetweenDGorxn and the Equilibrium Constant • The relationships among DGorxn, K, and the spontaneity of a reaction are:

40. ∆G, ∆G˚, and Keq But systems can reach equilibrium when reactants have NOT converted completely to products. In this case ∆Grxn is < ∆Gorxn , so state with both reactants and products present is MORE STABLE than complete conversion. Product Favored, ∆G˚ negative, K > 1

41. ∆G, ∆G˚, and Keq • Product-favored • 2 NO2 ---> N2O4 • ∆Gorxn = – 4.8 kJ • State with both reactants and products present is more stable than complete conversion. • K > 1, more products than reactants.

42. ∆G, ∆G˚, and Keq • Reactant-favored • N2O4 --->2 NO2 ∆Gorxn = +4.8 kJ • State with both reactants and products present is more stable than complete conversion. • K < 1, more reactants than products

43. Thermodynamics and Keq • Keq is related to reaction favorability. • When ∆Gorxn < 0, reaction moves energetically “downhill” • ∆Gorxn is the change in free energy when reactants convert COMPLETELY to products.

44. Relationship Between Gorxn and the Equilibrium Constant • The relationship for K at conditions other than thermodynamic standard state conditionsis derived from this equation.