Ch. 13: Chemical Equilibrium

1. Ch. 13: Chemical Equilibrium 13.1 The Equilibrium Condition

2. Equilibrium • dynamic equilibrium • may seem like no changes are occurring but there are changes • no NET changes • When did this reaction reach it?

3. reactants products H2O(g) + CO(g)  H2(g) + CO2(g)

4. At equilibrium, forward and reverse reaction rates are ____________

5. Equilibrium • equilibrium position of a reaction is determined by • initial _________________ • ____________ of reactants and products • degree of __________________ of reactants and products • GOAL: • _ • _

6. lies “to the left” more _________ less ___________ lies “to the right” less __________ more _________ If reactants are mixed and concentrations do not change could already be at equilibrium reaction rates are so ________ that change is too difficult to detect Equilibrium Position

7. Ch. 13: Chemical Equilibrium 13.2 Equilibrium Constant

8. Law of Mass Action • created in 1864 by Guldberg and Waage (Norweigen) • For a reaction: jA + kB ⇄ lC + mD • equilibrium constant: K

9. Law of Mass Action • because • Rateforward = • Ratereverse = • so if Ratef = Rater • kf[A]j[B]k =kr[C]l[D]m

10. Example • Write the equilibrium expression for: 4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g) • What would it be for the reverse reaction?

11. Equilibrium Constant • will always have the same value at a certain temperature • no matter what amounts are added • ratio at equilibrium will always be same

12. Equilibrium Position • each set of equilibrium concentrations • depends on initial concentrations

13. Ch. 13: Chemical Equilibrium 13.3 Equilibrium Expressions with Pressure

14. Equilibrium with Gases • equilibria involving gases can be described using __________ instead of _______________ N2(g) + 3H2(g)  2NH3(g)

15. Equilibrium with Gases N2(g) + 3H2(g)  2NH3(g)

16. Calculating K from KP • KP = KC : RT can cancel out if total of coefficients are same on each side • where ∆n is the difference in moles of gas on either side of the equation • ∆n = (l+m) – (j+k) N2(g) + 3H2(g)  2NH3(g) • ∆n =

17. Example • Setup the expression for KP in terms of KC, R and T 2NO(g) + Cl2(g)  2NOCl(g)

18. Ch. 13: Chemical Equilibrium 13.4: Heterogeneous Equilibria

19. Heterogeneous Equilibria • involve more than one phase • position of heterogeneous equilibria does NOT depend on amounts of: • _ • _ • because their concentrations stay constant (since they are PURE)

20. Heterogeneous Equilibria • do not include liquids or solids in equilibrium expression • only include ________ and ______________

21. Example 1 • 2H2O(l) ⇄ 2H2(g) + O2(g) • 2H2O(g) ⇄ 2H2(g) + O2(g)

22. Ch. 13: Chemical Equilibrium 13.5/6: Applications of Equilibrium Constant (K)

23. Equilibrium Constant • if we know the value of K, we can predict: • tendency of a reaction to occur • if a set of concentrations could be at equilibrium • equilibrium position, given initial concentrations

24. Equilibrium Constant • If you start a reaction with only reactants: • concentration of reactants will decrease by a certain amount • concentration of products will increase by a same amount

25. Example 2 • The following reaction has a K of 16. You are starting reaction with 9 O3 molecules and 12 CO molecules. • Find the amount of each species at equilibrium. O3(g) + CO(g)  CO2(g) + O2(g)

26. Example 2

27. Example 2

28. Example 2

29. If _________ mostly products goes essentially to completion lies far to right If _________ mostly reactants reaction is negligible lies far to left size of K and time needed to reach equilibrium are NOT related time required is determined by reaction rate (Ea) Extent of a Reaction

30. Reaction Quotient • Q: equal to equilibrium expression but ___________ have to be at equilibrium • used to tell if a reaction is at equilibrium or not • relationship between Q and K tells which way the reaction will shift • _______: at equilibrium, no shift • _______: too large, forms reactants, shift to left • ______: too small, forms products, shift to right

31. Example 3 • For the synthesis of ammonia at 500°C, the equilibrium constant is 6.0 x 10-2. Predict the direction the system will shift to reach equilibrium in the following case: N2(g) + 3H2(g)  2NH3(g)

32. Example 3 • [NH3]0 = 1.0x10-3 M, [N2]0=1.0x10-5 M [H2]0=2.0x10-3 M Q __ K so forms _________, shifts to _____

33. Example 4 • In the gas phase, dinitrogen tetroxide decomposes to gaseous nitrogen dioxide: N2O4(g) ⇄ 2NO2(g) • Consider an experiment in which gaseous N2O4 was placed in a flask and allowed to reach equilibrium at a T where KP = 0.133. At equilibrium, the pressure of N2O4 was found to be 2.71 atm. • Calculate the equilibrium pressure of NO2.

34. Example 4

35. Example 5 • At a certain temperature a 1.00 L flask initially contained 0.298 mol PCl3(g) and 8.70x10-3 mol PCl5(g). After the system had reached equilibrium, 2.00x10-3 mol Cl2(g) was found in the flask. PCl5(g)  PCl3(g) + Cl2(g) • Calculate the equilibrium concentrations of all the species and the value of K.

36. Example 5

37. Approximations • If K is very small, we can assume that the change (x) is going to be negligible • can be used to cancel out when adding or subtracting from a “normal” sized number • to simplify algebra 0

38. Example 6 • At 35°C, K=1.6x10-5 for the reaction 2NOCl(g) ⇄ 2NO(g) + Cl2(g) • Calculate the concentration of all species at equilibrium for the following mixtures • 2.0 mol NOCl in 2.0 L flask • 1.0 mol NOCl and 1.0 mol NO in 1.0 L flask • 2.0 mol NOCl and 1.0 mol Cl2 in 1.0 L flask

39. Example 6 • 2.0 mol NOCl in 2.0 L flask • [NOCl]= • [NO]= [Cl2]=

40. Example 6 • 1.0 mol NOCl and 1.0 mol NO in 1.0 L flask • [NOCl]= • [NO]= [Cl2]=

41. Example 6 • 2.0 mol NOCl and 1.0 mol Cl2 in 1.0 L flask • [NOCl]= • [Cl2]= [NO]=

42. Ch. 13: Chemical Equilibrium 13.7: Le’ Chatlier’s Principle

43. Le Châtlier’s Principle • can predict how certain changes in a reaction will affect the position of equilibrium

44. Changing Concentration • system will shift away from the added component or towards a removed component • Ex: N2 + 3H2 2NH3 • if more N2 is added, then equilibrium position shifts to right • if some NH3 is removed, then equilibrium position shifts to right

45. Change in Pressure • adding or removing gaseous reactant or product is same as changing conc. • adding inert or uninvolved gas • increase the ___________________ • ___________effect the equilibrium position

46. Change in Pressure • changing the volume • decrease V • decrease in # gas molecules • shifts towards the side of the reaction with _____ gas molecules • increase V • increase in # of gas molecules • shifts towards the side of the reaction with _____ gas molecules

47. Change in Temperature • all other changes alter the concentration at equilibrium position but don’t actually change value of K • value of K does change with temperature

48. Change in Temperature • if energy is added, the reaction will shift in direction that consumes energy • treat energy as a • __________: for endothermic reactions • __________: for exothermic reactions

49. add CO add C remove C add As4O6 remove As4O6 remove As4 decrease volume add Ne gas As4O6(s) + 6C(s) ⇄ As4(g) + 6CO(g)

50. decrease volume increase volume add P4 remove Cl2 add Kr gas add PCl3 P4(s) + 6Cl2(g) ⇄ 4PCl3(l)