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Understanding Chemical Equilibrium: Important Concepts and Constants

Explore the concept of equilibrium, equilibrium constants, and interpreting equilibrium constants. Learn about heterogeneous equilibrium and related properties. Equilibrium equations and equilibrium constants are discussed with examples.

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Understanding Chemical Equilibrium: Important Concepts and Constants

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  1. Chapter 15Chemical Equilibrium

  2. 15.1 The Concept of Equilibrium

  3. Forward reaction: N2O4 (g) 2 NO2 (g) Rate law: Rate = kf [N2O4] • Reverse reaction: 2 NO2 (g) N2O4 (g) Rate law: Rate = kr [NO2]2

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

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

  6. 15.2 The Equilibrium Constant • Opposing reactions naturally lead to equilibrium regardless of complexity or the nature of the kinetic processes for the forward and reverse reactions. • Ex: Haber Process N2(g) + 3 H2(g) 2 NH3(g) • An equilibrium mixture is obtained regardless of the starting quantities of any of the reactants or products:

  7. It does not matter whether we start with N2 and H2 or whether we start with NH3. We will have the same proportions of all three substances at equilibrium. • The equilibrium condition can be reached from either direction. • Law of Mass Action: Expresses, for any reaction, the relationship between the concentrations of the reactants and the products present at equilibrium

  8. cC + dD aA + bB • A, B, D and E are chemical species involved and a, b, d, and e are the coefficients in the balanced equation • The equilibrium-constant expression (or equilibrium expression) for this reaction would be • The equilibrium-constant is expressed in terms of the concentrations of reactant products according to the stoichiometry of the reaction, not its mechanism. • Suppose we have the following general equilibrium equation

  9. Illustration of law of mass action • As you can see, the ratio of [NO2]2 to [N2O4] remains constant at this temperature no matter what the initial concentrations of NO2 and N2O4 are.

  10. Results from Experiments 3 and 4 show that equilibrium can be achieved starting with either N2O4 or NO2

  11. Writing Equilibrium-Constant Expressions

  12. (PC)c (PD)d (PA)a (PB)b Kp = Equilibrium Constants in Terms of Pressure, Kp • Because pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written

  13. Relationship between Kc and Kp n = (moles of gaseous product) − (moles of gaseous reactant)

  14. 15.3 Interpreting and Working with Equilibrium Constants 1. The Magnitude of Equilibrium Constants • If K >> 1, the reaction is product-favored; product predominates at equilibrium. • The numerator of the equilibrium-constant expression must be much larger than the denominator.

  15. If K >> 1, the reaction is product-favored; product predominates at equilibrium. • If K << 1, the reaction is reactant-favored; reactant predominates at equilibrium.

  16. Kc = = 4.72 at 100C Kc = = 0.212 at 100C 2 NO2(g) N2O4(g) 2 NO2(g) N2O4(g) [NO2]2 [N2O4] [N2O4] [NO2]2 1 0.212 = 2. The Direction of the Chemical Equation and K • The direction in which we write the chemical equation for an equilibrium is arbitrary. • The equilibrium constant of a reaction in the reverse reaction is the reciprocal of the equilibrium constant of the forward reaction.

  17. Kc = = 0.212 at 100C Kc = = (0.212)2 = 0.0449 at 100C 2 NO2(g) N2O4(g) 4 NO2(g) 2 N2O4(g) [NO2]2 [N2O4] [NO2]4 [N2O4]2 3. Relating Chemical Equations and Equilibrium Constants • The equilibrium constant of a reaction that has been multiplied by a number is the equilibrium constant raised to a power that is equal to that number.

  18. The equilibrium constant for a net reaction made up of two or more steps is the product of the equilibrium constants for the individual steps.

  19. Equilibria in the same phase are called homogeneous equilibria. • In other cases the substances in equilibrium are in different phases, giving rise to heterogeneous equilibria. • Whenever a pure solid or a pure liquid is involved in a heterogeneous equilibrium, its concentration is not included in the equilibrium-constant expression for the reaction 15.4 Heterogeneous Equilibrium

  20. Even though a solid does not appear in the equilibrium-constant expression, it must be present for equilibrium to occur. • Why are they excluded? • The concentration of a pure solid or liquid has a constant value, regardless of the amount present.

  21. CaCO3 (s) CO2 (g) + CaO(s) • As long as some CaCO3 or CaO remain in the system, the amount of CO2 above the solid will remain the same.

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