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Chemical Equilibrium

Chemical Equilibrium

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Chemical Equilibrium

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  1. 17 Chemical Equilibrium

  2. Chapter Goals • Basic Concepts • The Equilibrium Constant • Variation of Kc with the Form of the Balanced Equation • The Reaction Quotient • Uses of the Equilibrium Constant, Kc • Disturbing a System at Equilibrium: Predictions

  3. Chapter Goals • The Haber Process: A Commercial Application of Equilibrium • Disturbing a System at Equilibrium: Calculations • Partial Pressures and the Equilibrium Constant • Relationship between Kp and Kc • Heterogeneous Equilibria

  4. Chapter Goals • Relationship between Gorxn and the Equilibrium Constant • Evaluation of Equilibrium Constants at Different Temperatures

  5. Basic Concepts • Reversiblereactions do not go to completion. • They can occur in either direction • Symbolically, this is represented as:

  6. Basic Concepts • Chemical equilibrium exists when two opposing reactions occur simultaneously at the same rate. • A chemical equilibrium is a reversible reaction that the forward reaction rate is equal to the reverse reaction rate. • Chemical equilibria are dynamic equilibria. • Molecules are continually reacting, even though the overall composition of the reaction mixture does not change.

  7. Basic Concepts • One example of a dynamic equilibrium can be shown using radioactive 131I as a tracer in a saturated PbI2 solution.

  8. Basic Concepts • This movie depicts a dynamic equilibrium.

  9. Basic Concepts • Graphically, this is a representation of the rates for the forward and reverse reactions for this general reaction.

  10. Basic Concepts • One of the fundamental ideas of chemical equilibrium is that equilibrium can be established from either the forward or reverse direction.

  11. Basic Concepts

  12. Basic Concepts

  13. For a simple one-step mechanism reversible reaction such as: The rates of the forward and reverse reactions can be represented as: The Equilibrium Constant

  14. The Equilibrium Constant • When system is at equilibrium: Ratef = Rater

  15. The Equilibrium Constant • Because the ratio of two constants is a constant we can define a new constant as follows :

  16. Similarly, for the general reaction: we can define a constant The Equilibrium Constant

  17. The Equilibrium Constant • Kc is the equilibrium constant . • Kc is defined for a reversible reaction at a given temperature as the product of the equilibrium concentrations (in M) of the products, each raised to a power equal to its stoichiometric coefficient in the balanced equation, divided by the product of the equilibrium concentrations (in M) of the reactants, each raised to a power equal to its stoichiometric coefficient in the balanced equation.

  18. The Equilibrium Constant Example 17-1: Write equilibrium constant expressions for the following reactions at 500oC. All reactants and products are gases at 500oC.

  19. The Equilibrium Constant

  20. The Equilibrium Constant

  21. The Equilibrium Constant • Equilibrium constants are dimensionless because they actually involve a thermodynamic quantity called activity. • Activities are directly related to molarity

  22. Example 17-2: One liter of equilibrium mixture from the following system at a high temperature was found to contain 0.172 mole of phosphorus trichloride, 0.086 mole of chlorine, and 0.028 mole of phosphorus pentachloride. Calculate Kc for the reaction. Equil []’s 0.028 M 0.172 M 0.086 M You do it! The Equilibrium Constant

  23. The Equilibrium Constant Example 17-3: The decomposition of PCl5 was studied at another temperature. One mole of PCl5 was introduced into an evacuated 1.00 liter container. The system was allowed to reach equilibrium at the new temperature. At equilibrium 0.60 mole of PCl3 was present in the container. Calculate the equilibrium constant at this temperature.

  24. The Equilibrium Constant Example 17-4: At a given temperature 0.80 mole of N2 and 0.90 mole of H2 were placed in an evacuated 1.00-liter container. At equilibrium 0.20 mole of NH3 was present. Calculate Kc for the reaction. You do it!

  25. The value of Kc depends upon how the balanced equation is written. From example 17-2 we have this reaction: This reaction has a Kc=[PCl3][Cl2]/[PCl5]=0.53 Variation of Kc with the Form of the Balanced Equation

  26. Example 17-5: Calculate the equilibrium constant for the reverse reaction by two methods, i.e, the equilibrium constant for this reaction. Equil. []’s 0.172 M 0.086 M 0.028 M The concentrations are from Example 17-2. Variation of Kc with the Form of the Balanced Equation

  27. The Reaction Quotient

  28. The Reaction Quotient

  29. The Reaction Quotient • The mass action expression or reaction quotient has the symbol Q. • Q has the same form as Kc • The major difference between Q and Kc is that the concentrations used in Q are not necessarily equilibrium values.

  30. The Reaction Quotient • Why do we need another “equilibrium constant” that does not use equilibrium concentrations? • Q will help us predict how the equilibrium will respond to an applied stress. • To make this prediction we compare Q with Kc.

  31. The Reaction Quotient

  32. The Reaction Quotient Example 17-6: The equilibrium constant for the following reaction is 49 at 450oC. If 0.22 mole of I2, 0.22 mole of H2, and 0.66 mole of HI were put into an evacuated 1.00-liter container, would the system be at equilibrium? If not, what must occur to establish equilibrium?

  33. Uses of the Equilibrium Constant, Kc Example 17-7: The equilibrium constant, Kc, is 3.00 for the following reaction at a given temperature. If 1.00 mole of SO2 and 1.00 mole of NO2 are put into an evacuated 2.00 L container and allowed to reach equilibrium, what will be the concentration of each compound at equilibrium?

  34. Uses of the Equilibrium Constant, Kc Example 17-8: The equilibrium constant is 49 for the following reaction at 450oC. If 1.00 mole of HI is put into an evacuated 1.00-liter container and allowed to reach equilibrium, what will be the equilibrium concentration of each substance?

  35. Disturbing a System at Equilibrium : Predictions • LeChatelier’s Principle - If a change of conditions (stress) is applied to a system in equilibrium, the system responds in the way that best tends to reduce the stress in reaching a new state of equilibrium. • We first encountered LeChatelier’s Principle in Chapter 14. • Some possible stresses to a system at equilibrium are: • Changes in concentration of reactants or products. • Changes in pressure or volume (for gaseous reactions) • Changes in temperature.

  36. Disturbing a System at Equilibrium : Predictions • Changes in Concentration of Reactants and/or Products • Also true for changes in pressure for reactions involving gases. • Look at the following system at equilibrium at 450oC.

  37. Disturbing a System at Equilibrium : Predictions • Changes in Concentration of Reactants and/or Products • Also true for changes in pressure for reactions involving gases. • Look at the following system at equilibrium at 450oC.

  38. Disturbing a System at Equilibrium : Predictions • Changes in Concentration of Reactants and/or Products • Also true for changes in pressure for reactions involving gases. • Look at the following system at equilibrium at 450oC.

  39. Disturbing a System at Equilibrium : Predictions • Changes in Volume • (and pressure for reactions involving gases) • Predict what will happen if the volume of this system at equilibrium is changed by changing the pressure at constant temperature:

  40. Disturbing a System at Equilibrium : Predictions

  41. Disturbing a System at Equilibrium : Predictions

  42. Disturbing a System at Equilibrium : Predictions

  43. Disturbing a System at Equilibrium : Predictions • Changing the Temperature

  44. Disturbing a System at Equilibrium : Predictions • Changing the Reaction Temperature • Consider the following reaction at equilibrium:

  45. Introduction of a Catalyst Catalysts decrease the activation energy of both the forward and reverse reaction equally. Catalysts do not affect the position of equilibrium. The concentrations of the products and reactants will be the same whether a catalyst is introduced or not. Equilibrium will be established faster with a catalyst. Disturbing a System at Equilibrium : Predictions

  46. Disturbing a System at Equilibrium : Predictions Example 17-9: Given the reaction below at equilibrium in a closed container at 500oC. How would the equilibrium be influenced by the following?

  47. Disturbing a System at Equilibrium : Predictions Example 17-10: How will an increase in pressure (caused by decreasing the volume) affect the equilibrium in each of the following reactions?

  48. Disturbing a System at Equilibrium : Predictions Example 17-11: How will an increase in temperature affect each of the following reactions?

  49. The Haber Process: A Practical Application of Equilibrium • The Haber process is used for the commercial production of ammonia. • This is an enormous industrial process in the US and many other countries. • Ammonia is the starting material for fertilizer production. • Look at Example 17-9. What conditions did we predict would be most favorable for the production of ammonia?

  50. The Haber Process: A Practical Application of Equilibrium