1. Chemical Equilibrium Chapter 13 4 out of 75 m/c Free Response: Required Every Year

2. What is it?? In a reversible reaction that does not go to completion, some balanced state exists at which the rate of the forward and reverse reactions perfectly balance each other.

3. Definition Chemical Equilibrium—the state where the concentrations of all reactants and products remain constant with time

4. Forward & reverse rxns occur at different rates until equilibrium is established.

5. Terminology Rxn that almost goes to completion—equilibrium position far to the right Rxn that forms very little product—equilibrium far to the left

6. Dynamic Equilibrium Forward and reverse reactions are occurring. Overall, the concentrations are unchanged, but products are constantly being broken down and reformed.

7. Achieving Equilibrium Rapid change at the start of the reaction Change slows as conc. nears equilibrium Change stops at equilibrium (no slope)

8. Achieving Equilibrium Nature tries to achieve minimum energy & maximum disorder Depends on: Initial concentrations Relative energies of reactants and products Degree of “organization”

9. Law of Mass Action Equilibrium expressions based on experimental data Equilibrium expression— mathematical representation of equilibrium condition

10. Equilibrium Expression For jA + kB  lC + mD: [C] l [D] m [A] j [B] k K = equilibrium constant [A], [B]…concentrations of reactants & products l, m, …coefficients in balanced equation K =

11. Writing an equilibrium expression Equal to K Products over reactants Raise to the power of the coefficients

12. Index Card: Write equilibrium expressions for: 1. 2NO + Cl 2  2NOCl 2. N 2 O 4  2NO 2 3. N 2 + 3H 2  2NH 3 4. H 2 SO 3  H + + HSO 3 -

13. Solving for K Substitute values for equilibrium concentrations of reactants and products into equilibrium expression K is temperature dependent—only true for given temp.

14. Index Card If equilibrium concentrations for the first reaction are as follows, what is the equilibrium constant? [NO] =.34 M; [Cl 2 ] =.68 M; [NOCl] = 5.6 M

15. Interpreting an Equilibrium Expression If K is large (much greater than 1), is the equilibrium far to the left or right (more product or reactant)? If K is very small (much less than one), is the equilibrium far to the left or right (more reactant or product)?

16. Manipulating K If you know K for a reaction: what is K for the reverse reaction?? what is K if you double all of the coefficients?

17. Equilibrium Position Each set of data that gives K defines an equilibrium position. Only one equilibrium constant, but infinite equilibrium positions

18. Gaseous Equilibrium Since pressure is analogous to concentration for gases, you can use pressure to calculate K Written as K p Same process/formula

19. Index Card Write Kp expressions for the following: 1. PCl 5 (g)  PCl 3 (g) + Cl 2 (g) 2. H 2 (g) + F 2 (g)  2HF (g)

20. Relating K and K p PV = nRT ; P = (n/v)(RT) ; n/v = C ; C = P / RT Substitute P / RT for C If the sum of coefficients on both sides of an equation are equal, then RT term cancels: K = K p

21. RT Term Cancels 2HCl(g) + F 2 (g)  2HF (g) + Cl 2 (g)

22. Relating K and K p If sum of coefficients is not equal, then the RT term must be included Sample 3

23. N 2 + 3H 2  2NH 3 If Kc at 25* C is 0.0602, what is Kp? Kp = K(RT) n

24. Homogeneous Equilibria All reactants and products are in the same phase (gaseous)

25. Heterogeneous Equilibria Involves more than one phase of matter Experimental evidence shows that equilibrium does not depend on the amount present of pure substances in solid or liquid phase.

26. Amount of gas does not depend on amount of solid.

27. Heterogeneous Equilibria Pure solids and pure liquids are not included in the equilibrium expression.

28. Uses of Eq. Constant Predict tendency for rxn. to occur (but NOT rate of rxn) Tell if a system is at equilibrium or if rxn is just too slow to detect Tell what eq. conc.will be

29. Initial Conditions See p. 622-623 Reactant conc. at 100% Product conc. at 0%

30. Equilibrium Conditions Reactant conc. at (100% - x) Product conc. at x Oversimplification— concentrations will depend on balanced equation but can be determined

31. Reaction Quotient (Q) Use law of mass action with initial conc. instead of eq. conc. to predict behavior If Q = K—no change (already at eq.) If Q > K—shifts to left If Q < K—shifts to right

32. Example Problem For the reaction N 2 O 4 2NO 2, K p =.133. Calculate Q an tell which way the reaction will proceed if the pressure of N 2 O 4 is 2.71 atm and the pressure of NO 2 is 1.68 atm.

33. Solution

34. Index Card: For synthesis of ammonia at 500* C, K = 6.0 x 10 -2. Calculate Q for each set of conditions below and tell which way the reaction will proceed. [NH 3 ][N 2 ][H 2 ] 1.0 x 10 -3 M1.0 x 10 -5 M2.0 x 10 -3 M 2.0 x 10 -4 M1.5 x 10 -5 M3.5 x 10 -1 M 1.0 x 10 -4 M5.0 M1.0 x 10 -2 M

35. Extent of Rxn Large K—greater than one—more products than reactants remain Small K—less than one—more reactants than products remain

36. Extent of Rxn The size of K and the time required to reach equilibrium are unrelated!!

37. Calculating Equilibrium Concentration or Pressure Sample Problems 13.8 – 13.9

38. Equilibrium Concentrations You can find the final concentrations of reactants and products for many different initial concentrations. ICE Chart

39. Summary Start with balanced equation Write eq. Expression (Law of Mass Action) List initial conc. (I) Calculate Q to determine shift of system

40. Summary Define the change in system (C) Substitute known conc. into expression and solve for unknown (E) Check by using solution to solve for K

41. Sample Problems 13.10-13.11

42. Le Chatelier’s Principle If a change is imposed on a system in equilibrium, the position of the equilibrium will shift in a direction that tends to reduce the change.

43. Le Chatelier’s Principle Allows prediction of the effects of changes in concentration, pressure, and temperature Useful to manipulate a reaction—maximize products

44. Concentration Change Shifts away from the side of the substance whose concentration increased Shifts toward the side of the substance whose concentration decreased Consider effect on Q

45. Concentration Change Example: N 2 +3H 2  2NH 3 Add N 2 or H 2 —Q decreases; more NH 3 produced—uses up N 2 & H 2 Take away NH 3 as it forms— Q decreases—more NH 3 forms K must stay the same.

46. Effect of adding N 2 to the equilibrium mixture

47. Causing Pressure Change Add or remove gaseous reactant or product Add an inert gas Change the volume of container

48. Pressure Change When adding or removing gaseous product or reactant, shift is same as conc. Change. When adding an inert gas, no effect on equilibrium because partial pressures of each reactant or product are the same

49. Pressure Change Decreasing total volume— increases pressure—favors rxn that produces fewer moles of gas Increasing total volume— decreases pressure—favors rxn that produces more moles of gas

50. Effect of increased pressure

51. Pressure Change Example: N 2 + 3H 2  2NH 3 Increased pressure favors product (2 moles of gas vs. 4 moles) Decreased pressure favors reactant (4 moles vs. 2 moles)

52. Temperature Change Changes value of K Increase in temperature favors process that uses energy Decrease in temperature favors process that produces energy