Chemical Equilibrium

1. Chemical Equilibrium • Chapter 13

2. Chemical Equilibrium • The state where the concentrations of all reactants and products remain constant with time. • On the molecular level, there is frantic activity. Equilibrium is not static, but is a highly dynamic situation.

3. Equilibrium & Kinetics • At equilibrium, Rf = Rr • but kf is not necessarily equal to kr. • K= the equilibrium constant

4. Magnitude of K • A K value much larger than 1 means that the equilibrium system contains mostly products -- equilibrium lies far to the right. • A very small K value means the system contains mostly reactants -- equilibrium lies far to the left. • The size of K and the time required to reach equilibrium are not directly related!

5. Reactions That Appear to Run to Completion • 1. Formation of a precipitate. • 2. Formation of a gas. • 3. Formation of a molecular substance such as water. • These reactions appear to run to completion, but actually the equilibrium lies very far to the right. All reactions in closed vessels reach equilibrium.

6. The Law of Mass Action • For • jA + kB lC + mD • The law of mass action (Cato Guldberg & Peter Waage) is represented by the equilibrium expression:

7. Equilibrium Expression • 4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(g)

8. Equilibrium Expressions • N2(g) + 3H2(g) ---> 2NH3(g) • 2NH3(g) ---> N2(g) + 3H2(g) • K´ = 1/K

9. Equilibrium Position • For a given reaction at a given temperature, there is only one equilibrium constant (K), but there are an infinite number of equilibrium positions. • Where the equilibrium position lies is determined by the initial concentrations of the reactants and products. The initial concentrations do not affect the equilibrium constant.

10. Three equilibrium positions but only one equilibrium constant (K).

11. Heterogeneous Equilibria • . . . are equilibria that involve more than one phase. • CaCO3(s)  CaO(s) + CO2(g) • K = [CO2] • The position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present.

12. Write the equilibrium expression for the following reactions . • Ba(OH)2(s) Ba+2 (aq) + 2OH- (aq) HCl(g) + NH3(g) NH4Cl (s)

13. Kp • K in terms of equilibrium partial pressures: • .

14. K versus Kp • For jA + kB lC + mD • Kp = K(RT)n • n = sum of coefficients of gaseous products minus sum of coefficients of gaseous reactants.

15. Calculate the value of K at 25°C if the value for Kp is 1.9x103 for the reaction 2NO(g) + Cl2(g) 2NOCl(g)

16. K versus Kp • When n = 0 and Kp = K(RT)o, then • Kp = K • K and Kp are equal numerically but do not have the same units.

17. [SO2] = 1.5M [O2] = 1.25M[SO3] = 3.5M • 2SO2(g) + O2 2SO3

18. Review book • p306 1,2,4-8,20

19. Reaction Quotient • Reaction Quotient is used for concentrations not at equilibrium . It helps to determine the direction of the move toward equilibrium. • The law of mass action is applied with initial concentrations.

20. Reaction Quotient • H2(g) + F2(g)  2HF(g)

21. Q versus K • Q = K System is at equilibrium. • Q > K System will shift to the left. • Q < K System will shift to the right. • .

22. Kp= .133 atm at a particular temp for the reaction below N2O4(g) 2NO 2 (g) Calculate the reaction Quotient if pN2O4 =. 048 atm and pNO2= .058 What will the system do to reach equilibrium ?

23. Solving Equilibrium Problems • 1. Balance the equation. • 2. Write the equilibrium expression. • 3. List the initial concentrations. • 4. Calculate Q and determine the shift to equilibrium.

24. ICE • H2(g) + F2(g) <---> 2HF(g) • [H2]o = [F2]o = [HF]o = 3.000 mol/1.500 L = 2.000M 1.000 Q < K  equilibrium shifts right.

25. ICE • H2(g) + F2(g) <---> 2HF(g) • [H2] [F2] [HF] • Initial (mol/L) 2.000 2.000 2.000 • Change (mol/L) - x - x + 2x • Equil. (mol/L) 2.000 - x 2.000 - x 2.000 + 2x

26. ICEContinued x = 1.528 [H2] = [F2] = 2.000 M - 1.528 M = 0.472 M [HF] = 2.000 M + 2(1.528 M) = 5.056 M Substitute into K expression to check validity!

27. ICE & the Quadratic Formula • See Sample Exercise 13.12 on pages 637-638. • Use the quadratic formula to solve for K. • Obtain the quadratic formula program from your instructor.

28. Systems With Small K’s • See example on pages 639-640. • When K is small, the change (x) is small compared to the initial concentration and a simplification can be made for the calculation. • This approximation must be checked to see if it is valid. When K is two powers of ten or more smaller than the initial concentration, the approximation should be OK.

29. Le Châtelier’s Principle • . . . if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change.

30. Le Chatelier’s Principle • If a reactant or product is added to a system at equilibrium, the system will shift away from the added component. • If a reactant or product is removed, the system will shift toward the removed component. • See Sample Exercise 13.13 on pages 642-643.

31. Effects of Changes on the System • 1. Concentration: The system will shift away from the added component. • 2. Temperature: K will change depending upon the temperature (treat heat as a reactant or product).

32. Effects of Changes on the System (continued) • 3. Pressure: • a. Addition of inert gas does not affectthe equilibrium position. • b. Decreasing the volume shifts the equilibrium toward the side with fewer gaseous moles.

33. The system of N2, H2, and NH3 are initially at equilibrium. When the volume is decreased, the system shifts to the right -- toward fewer molecules.

35. Equilibrium Constant, K • For an exothermic reaction, if the temperature increases, K decreases. • For an endothermic reaction, if the temperature increases, K increases.