1. Chapter 16
Equilibrium

2. Collision Theory Collision theory:
1. atoms, ions, and molecules must collide in order to react.
Only a small number of collisions produce reactions

3. 2. Reacting substances must collide in the correct orientation
2. Reacting substances must collide in the correct orientation. CO + NO2  CO2 + NO

4. 3. Reacting substances must collide with sufficient energy to form an activated complex.
Activated complex (transition state) – temporary, unstable arrangement of atoms in which old bonds are breaking and new bonds are forming.

5. Activation energy (Ea) – the minimum amount of energy needed to form the activated complex and lead to a reaction
Collision > Ea
Collision < Ea

7. Exothermic Reaction

8. Endothermic Reaction

9. Factors Affecting Reaction Rates
Nature of reactants
Some things are more reactive than others (activity series)
Concentration of reactants
Temperature
Catalysts

10. Lowers activation energy required for a reaction to take place

12. Chemical Equilibrium
Equilibrium – the exact balancing of two processes, one of which is the exact opposite of the other

14. Reversible reactions and chemical equilibrium
Not all reactions go to completion
Reversible reaction – a reaction that occurs in both the forward and reverse direction
Forward:
Reverse:
Both occurring:

15. Chemical equilibrium – a dynamic state in which the forward and reverse reactions balance each other out because they take place at equal rates
Rate forward reaction = Rate reverse reaction
Concentration of all reactants and products remains constant

16. Equal numbers of moles of H2O and CO are mixed in a closed container.
The reaction begins to occur, and some products (H2 and CO2) are formed.
The reaction continues as time passes and more reactants are changed to products.
Although time continues to pass, the numbers of reactant and product molecules are the same as in (c). No further changes are seen as time continues to pass. The system has reached equilibrium.

17. Why does equilibrium occur?
As concentration of reactants decreases, forward reaction slows down
As concentration of products increases, reverse reaction speeds up

19. Equilibrium expressions
Law of chemical equilibrium – at a given temperature a chemical system reaches a state in which the ratio of reactant and product concentration has a constant value

20. The Equilibrium Constant
Law of chemical equilibrium
For a reaction of the type
aA + bB ↔ cC + dD
Equilibrium expression
Each set of equilibrium concentrations is called an equilibrium position.

21. Interpreting Keq
Keq > 1:
Keq < 1:

22. Homogeneous equilibria
All reactants and products are in the same physical state
Write the equilibrium expression for the following:
H2(g) + F2(g) ↔2HF(g)
N2(g) + 3H2(g) ↔ 2NH3(g)

23. N2O4(g) ↔ 2NO2(g)
CO(g) + 3H2(g) ↔ CH4(g) + H20(g)

24. Suppose that for the reaction:. 2SO2(g) + O2(g)↔2SO3(g)
Suppose that for the reaction: 2SO2(g) + O2(g)↔2SO3(g) it is determined that at a particular temperature the equilibrium concentrations are: [SO2] = 1.50M, [O2] = 1.25M, [SO3] = 3.50M. Calculate the value of K for this reaction.

25. Heterogeneous equilibria
Reactants and products are present in more than one state

26. Keq does not depend on amounts of solids or liquids because their concentration cannot change.

27. Write the equilibrium expression for the following:
2NaHCO3(s) ↔ Na2CO3(s) + CO2(g) + H2O(g)
2H2O(l) ↔ 2H2(g) + O2(g)

28. Using Equilibrium Constants
Equilibrium constant expressions can be used to calculate concentrations and solubilities

29. Calculating equilibrium concentrations
At a temperature of 1405 K, hydrogen sulfide decomposes to form hydrogen and diatomic sulfur according to the reaction:
2H2S(g) ↔ 2H2(g) + S2(g). The equilibrium constant for the reaction is 2.27 x 10-3.What is the concentration of hydrogen gas if [S2] = M and [H2S] = 0.184M

30. PCl5 ↔ PCl3 + Cl2 . In a certain experiment, at a temperature where K = 8.96 x 10-2, the equilibrium concentrations of PCl5 = 6.70 x 10-3 M and PCl3 = M. What was the equilibrium concentration of Cl2?

31. Equilibrium Constants
For a reaction at a constant temperature Keq will always be the same regardless of initial concentration of reactants and products
N2 + H2 ↔ 2NH3
Initial Concentrations
Equilibrium Concentrations
Exp.
[N2]
[H2]
[NH3] K 1
1.00M
0.921M
0.763M
0.157M
0.0602 2
0.399M
1.197M
0.203M 3
2.00M
3.00M
2.59M
2.77M
1.82M

32. Equilibrium Characteristics
Reaction must take place in a closed system
Temperature must remain constant
All reactants and products are in constant dynamic motion

33. Factors Affecting Chemical Equilibrium
When changes are made to a system at equilibrium the system shifts to a new equilibrium position

34. Le Chatelier’s principle: if a stress is applied to a system at equilibrium, the system shifts in the direction that relieves that stress.

35. Applying Le Chatelier’s Principle
Change in Concentration
N2(g) + 3H2(g) ↔ 2NH3(g)

36. When a reactant or product is added to a system at equilibrium, the system shifts away from the added component
If a reactant or product is removed from a system at equilibrium, the system shifts toward the removed component

38. For the reaction:. As4O6(s) + 6C(s) ↔ As4(g) + 6CO(g)
For the reaction: As4O6(s) + 6C(s) ↔ As4(g) + 6CO(g) Predict the direction of the shift in the equilibrium position in response to each of the following:
Addition of carbon monoxide
Addition of As4O6(s)
removal of C(s)
Removal of As4(g)

39. Changes in volume and pressure (gas only)
Remember: for gasses there is an inverse relationship between volume and pressure

40. The system is initially at equilibrium.
The piston is pushed in, decreasing the volume and increasing the pressure. The system shifts in the direction that consumes CO2 molecules, lowering the pressure again.

41. Decreasing the volume – system shifts in direction that gives fewest number of gas molecules

42. Increasing the volume – system shifts in the direction to increase the number of gas molecules (increase the pressure)

43. Predict the direction of the equilibrium shift for each of the following when the volume of the container is decreased:
CO(g) + 2H2(g) ↔ CH3OH(g)
H2(g) + F2(g) ↔ 2HF(g)

44. Change in temperature: Exothermic reaction – heat is a product
Adding energy shifts equilibrium away from products
Endothermic reaction – heat is a reactant
Adding energy shifts equilibrium away from reactants
Predict the direction of the shift in equilibrium position the same way as if a product or reactant were added or removed

46. Decide whether higher or lower temperatures will produce more CH3CHO in the reaction:
C2H2(g) + H2O(g) ↔ CH3CHO ΔH = -151 kJ

47. For the exothermic reaction:. 2SO2(g) + O2(g) ↔ 2SO3 (g)
For the exothermic reaction: SO2(g) + O2(g) ↔ 2SO3 (g) Predict the equilibrium shift caused by each of the following changes:
SO2 is added
SO3 is removed
Volume is decreased
Temperature is decreased
Ne(g) is added

48. Catalysts – speeds up a reaction equally in both directions therefore does not affect equilibrium

49. The Solubility Product Constant
Upon dissolving, all ionic compounds dissociate into ions
NaCl(s)
BaSO4(s)
Mg(OH)2(s)

50. Solubility Equilibria
CaF2(s) ↔ Ca2+(aq) + 2F-(aq)
At equilibrium, solution is saturated
Ksp = [Ca2+][F-]2
Ksp = solubility product constant

51. Write the balanced equation describing the reaction for dissolving each of the following in water, also write the Ksp expression for each solid:
PbCl2(s)
Ag2CrO4(s)

52. Small Ksp = not very soluble compound

53. The Ksp value for solid AgI is 1. 5 x 10-16
The Ksp value for solid AgI is 1.5 x Calculate the solubility of AgI in water.

54. The Ksp value for solid CuCO3 is 2. 5 x 10-10
The Ksp value for solid CuCO3 is 2.5 x Calculate the solubility of AgI in water.

55. The solubility of CuBr is 2. 0 x 10-4 mol/L
The solubility of CuBr is 2.0 x 10-4 mol/L. Calculate the Ksp for copper (I) bromide.