Ch. 14: Chemical Equilibrium

1. Ch. 14: Chemical Equilibrium Dr. Namphol Sinkaset Chem 201: General Chemistry II

2. I. Chapter Outline • Introduction • The Equilibrium Constant (K) • Values of Equilibrium Constants • The Reaction Quotient (Q) • Equilibrium Problems • Le Châtelier’s Principle

3. I. Introduction • Equilibrium will be the focus for the next several chapters. • Most reactions are reversible, meaning they can proceed in both forward and reverse directions. • This means that as products build up, they will react and reform reactants. • At equilibrium, the forward and backward reaction rates are equal.

4. I. Example Equilibrium

5. II. Equilibrium Concentrations • Equilibrium does not mean that concentrations are all equal!! • However, we can quantify concentrations at equilibrium. • Every equilibrium has its own equilibrium constant.

6. II. The Equilibrium Constant • equilibrium constant: the ratio at equilibrium of the [ ]’s of products raised to their stoichiometric coefficients divided by the [ ]’s of reactants raised to their stoichiometric coefficients. • The relationship between a balanced equation and equilibrium constant expression is the law of mass action.

7. II. The Equilibrium Constant • For a general equilibrium aA + bB  cC + dD, the equilibrium expression is:

8. II. Sample “Problem” • Write the equilibrium constant expression for the reaction: 2H2(g) + O2(g) 2H2O(g).

9. II. Physical Meaning of K • Large values of K mean that the equilibrium favors products, i.e. there are high [ ]’s of products and low [ ]’s of reactants at equilibrium. • Small values of K mean that the equilibrium favors reactants, i.e. there are low [ ]’s of products and high [ ]’s of reactants at equilibrium.

10. II. Rules for Manipulating K • If the equation is reversed, the equilibrium constant is inverted.

11. II. Rules for Manipulating K • If the equation is multiplied by a factor, the equilibrium constant is raised to the same factor.

12. II. Rules for Manipulating K • When chemical equations are added, their equilibrium constants are multiplied together to get the overall equilibrium constant.

13. II. Sample Problem • Predict the equilibrium constant for the first reaction given the equilibrium constants for the second and third reactions. CO2(g) + 3H2(g) CH3OH(g) + H2O(g) K1 = ? CO(g) + H2O(g)  CO2(g) + H2(g) K2 = 1.0 x 105 CO(g) + 2H2(g)  CH3OH(g) K3 = 1.4 x 107

14. II. K in Terms of Pressure • Up to this point, we’ve been using concentration exclusively in the equilibrium expressions. • Partial pressures are proportional to concentration via PV = nRT. • Thus, for gas reactions, partial pressures can be used in place of concentrations.

15. II. Two Different K’s • For the reaction 2SO3(g) 2SO2(g) + O2(g), we can write two equilibrium expressions.

16. II. Relationship Between Concentration and Pressure • To be able to convert between Kc and Kp, we need a relationship between concentration and pressure.

17. II. Converting Between Kc and Kp

18. II. Converting Between Kc and Kp • The Δn is the change in the number of moles of gas when going from reactants to products. • When does Kp equal Kc?

19. II. Sample Problem • Methanol can be synthesized via the reaction CO(g) + 2H2(g) CH3OH(g). If Kp of this reaction equals 3.8 x 10-2 at 200 °C, what’s the value of Kc?

20. II. Heterogeneous Equilibria • If an equilibrium contains pure solids or pure liquids, they are not included in the equilibrium constant expression.

21. III. Values of K • Values of K are most easily calculated by allowing a system to come to equilibrium and measuring [ ]’s of the components. • For the equilibrium H2(g) + I2(g) 2HI(g), let’s say equilibrium [ ]’s at 445 °C were found to be 0.11 M, 0.11 M, and 0.78 M for molecular hydrogen, molecular iodine, and hydrogen iodide, respectively.

22. III. Kc for a H2/I2 Mixture • Note that units are not included when calculating K’s. • Thus, equilibrium constants are unitless.

23. III. Equilibrium [ ]’s Vs. K • For any reaction, the equilibrium [ ]’s will depend on the initial [ ]’s of reactants or products. • However, no matter how you set up the reaction, the value of the equilibrium constant will be the same if the temperature is the same.

24. III. Equilibrium [ ]’s Vs. K

25. IV. The Reaction Quotient • What happens when we mix reactants together and wait? • Can we predict what will happen when we have a mixture of reactants and products? • The reaction quotient, Qc or Qp, is used to predict in which direction an equilibrium will move.

26. IV. Formula for Qc or Qp • You already know the formula because it’s the same as for Kc or Kp!! • The difference is, we don’t know if the reaction is at equilibrium, thus, we cannot set the ratio equal to K! • For the reaction aA + bB  cC + dD:

27. IV. Using Q • The value of Q relative to K tells you whether the reaction will form more products or more reactants to reach equilibrium. • Q < K means reaction forms products. • Q > K means reaction form reactants. • Q = K means reaction is at equilibrium.

28. IV. Sample Problem • Consider the reaction N2O4(g) 2NO2(g) with Kc = 5.85 x 10-3. If a reaction mixture contains [NO2] = 0.0255 M and [N2O4] = 0.0331 M, which way will the reaction proceed?