Unit XII: Equilibrium

1. Unit XII: Equilibrium … Chapter 16…

2. The Concept of Equilibrium • Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.

3. The Concept of Equilibrium • As a system approaches equilibrium, both the forward and reverse reactions are occurring. • At equilibrium, the forward and reverse reactions are proceeding at the same rate.

4. A System at Equilibrium • Once equilibrium is achieved, the amount of each reactant and product remains constant

5. Reactions are Reversible • A + B  C + D ( forward) • C + D  A + B (reverse) • Initially there is only A and B so only the forward reaction is possible • As C and D build up, the reverse reaction speeds up while the forward reaction slows down. • Eventually the rates are equal

6. What is at Equilibrium? • Rates are equal • Concentrations are not. • Rates are determined by concentrations and activation energy. • The concentrations do not change at equilibrium. • or if the reaction is verrrry slooooow.

7. N2O4 (g) 2 NO2 (g) Depicting Equilibrium • Since, in a system at equilibrium, both the forward and reverse reactions are being carried out, we write its equation with a double arrow.

8. Equilibrium Constant • Forward reaction: N2O4 (g) 2 NO2 (g) • Rate Law: Rate = kf [N2O4]

9. Equilibrium Constant • Reverse reaction: 2 NO2 (g) N2O4 (g) • Rate Law: Rate = kr [NO2]2

10. [NO2]2 [N2O4] = kf kr Equilibrium Constant • Therefore, at equilibrium Ratef = Rater kf [N2O4] = kr [NO2]2 • Rewriting this, it becomes

11. [NO2]2 [N2O4] Keq = kf kr = Equilibrium Constant • The ratio of the rate constants is a constant at that temperature, and the expression becomes

12. [C]c[D]d [A]a[B]b Kc = Equilibrium Constant aA + bB cC + dD The equilibrium expression becomes:

13. Equilibrium Constant

14. Equilibrium Constant

15. Equilibrium Constant Expression • In an equilibrium constant expression • All concentrations are equilibrium values • Product concentrations appear in the numerator, and reactant concentrations appear in the denominator • Each concentration is raised to the power of its stoichiometric coefficient in the balanced chemical equation • The value of the constant K depends on the particular reaction and on the temperature • Units are never given with K

16. Writing Equilibrium Constant Expressions • Reactions involving solids S(s) + O2(g) SO2(g) • Concentration of a solid is determined by its density… and density is fixed • Meaning the concentration of a solid is essentially a constant and not included in the expression • The concentrations of any solid reactants and products are not included in the equilibrium constant expression

17. Writing Equilibrium Constant Expressions

18. Writing Equilibrium Constant Expressions • Reactions in Aqueous solutions • The real question here isn’t about aqueous… it’s about pure liquids • Both the concentrations of liquids and solids can be obtained by multiplying the density of the substance by its molar mass — and both of these are constants at constant temperature.

19. Writing Equilibrium Constant Expressions

20. Writing Equilibrium Constant Expressions

21. (PCc) (PDd) (PAa) (PBb) Kp = Writing Equilibrium Constant Expressions • Reactions involving Gases: Kc and Kp • Since pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written

22. n V P = RT Writing Equilibrium Constant Expressions • Relationship between Kc and Kp • From the Ideal Gas Law we know that • PV = nRT • Rearranging it, we get

23. Writing Equilibrium Constant Expressions • Plugging this into the expression for Kp for each substance, the relationship between Kc and Kp becomes where Kp = Kc (RT)n n = (moles of gaseous product) - (moles of gaseous reactant)

24. Practice

25. Practice • Write the equilibrium constant expressions for the following reactions: • N2(g) + 3H2(g) 2NH3(g) • H2CO3(g) + H2O(l) HCO3-(aq) + H3O+(aq)