Download
chemical equilibrium n.
Skip this Video
Loading SlideShow in 5 Seconds..
Chemical Equilibrium PowerPoint Presentation
Download Presentation
Chemical Equilibrium

Chemical Equilibrium

0 Views Download Presentation
Download Presentation

Chemical Equilibrium

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Chemical Equilibrium

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

  3. Basic Concepts • Chemical equilibrium exists when two opposing reactions occur simultaneously at the same rate. • Forward reaction rate is equal to the reverse reaction rate. • Chemical equilibria are dynamic equilibria. • Molecules are continually reacting.

  4. aA+bB↔cC+dD

  5. Basic Concepts

  6. Basic Concepts

  7. 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

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

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

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

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

  12. The Equilibrium Constant

  13. 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 The Equilibrium Constant

  14. 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

  15. 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?

  16. The Reaction Quotient • The mass action expression or reaction quotient has the symbol Q. • The concentrations used in Q are not necessarily equilibrium values.

  17. 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.

  18. The Reaction Quotient

  19. 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?

  20. Practice-Due Today 2. At 700 K, carbon monoxide reacts with water to form carbon dioxide and hydrogen. The equilibrium constant for this reaction is 5.10. Consider an experiment in which 1.00 mol of CO and 1.00 mole of water are mixed together in a 1.00 L flask. Calculate the concentrations of all the species at equilibrium.

  21. Partial Pressures and the Equilibrium Constant • For gas phase reactions the equilibrium constants can be expressed in partial pressures rather than concentrations. • For gases, the pressure is proportional to the concentration. • We can see this by looking at the ideal gas law. • PV = nRT • P = nRT/V • n/V = M • P= MRT and M = P/RT

  22. Partial Pressures and the Equilibrium Constant • For convenience we may express the amount of a gas in terms of its partial pressure rather than its concentration. • To derive this relationship, we must solve the ideal gas equation.

  23. Partial Pressures and the Equilibrium Constant • Consider this system at equilibrium at 5000C.

  24. Partial Pressures and the Equilibrium Constant

  25. Relationship Between Kp and Kc • From the previous slide we can see that the relationship between Kp and Kc is:

  26. Relationship Between Kp and Kc Example 17-15: Nitrosyl bromide, NOBr, is 34% dissociated by the following reaction at 25oC, in a vessel in which the total pressure is 0.25 atmosphere. What is the value of Kp?

  27. Relationship Between Kp and Kc • The numerical value of Kc for this reaction can be determined from the relationship of Kp and Kc.

  28. Example 17-16: Kc is 49 for the following reaction at 450oC. If 1.0 mole of H2 and 1.0 mole of I2 are allowed to reach equilibrium in a 3.0-liter vessel, (a) How many moles of I2 remain unreacted at equilibrium? You do it! Relationship Between Kp and Kc

  29. Relationship Between Kp and Kc (b) What are the equilibrium partial pressures of H2, I2 and HI? You do it!

  30. Relationship Between Kp and Kc (c) What is the total pressure in the reaction vessel? You do it!

  31. 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.

  32. 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.

  33. 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.

  34. 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.

  35. 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:

  36. Disturbing a System at Equilibrium : Predictions

  37. Disturbing a System at Equilibrium : Predictions

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

  39. 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

  40. 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?

  41. 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?

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

  43. Practice-Due Today • NO2(g) ↔N2(g) + O2(g) which is exothermic. • Predict how each of the following stresses will affect the conc. of O2 and the value of K: • NO2 is added • H2 is removed • The volume is halved • He (g) is added • Temperature is increased • Catalyst is added

  44. Solving Equilibrium Problems • Write the balanced equation. • Set up the equilibrium expression (no values yet). • If you can’t tell which way the rxn will shift, solve for Q. • Set up a chart that includes the equation, initial concentrations, changes in concentration (in terms of x), and final concentrations. **ICE** • Substitute the final concentrations and solve for x.

  45. Disturbing a System at Equilibrium: Calculations • To help with the calculations, we must determine the direction that the equilibrium will shift by comparing Q with Kc. • Example 17-12: An equilibrium mixture from the following reaction was found to contain 0.20 mol/L of A, 0.30 mol/L of B, and 0.30 mol/L of C. What is the value of Kc forthis reaction?

  46. Disturbing a System at Equilibrium: Calculations

  47. Disturbing a System at Equilibrium: Calculations Example 17-13: Refer to example 17-12. If the initial volume of the reaction vessel were halved, while the temperature remains constant, what will the new equilibrium concentrations be? Recall that the original concentrations were: [A]=0.20 M, [B]=0.30 M, and [C]=0.30 M.

  48. Disturbing a System at Equilibrium: Calculations An additional 0.80 mole of Cl2 is added to the vessel at the same temperature. Calculate the molar concentrations of CO, Cl2, and COCl2 when the new equilibrium is established.

  49. Heterogeneous equilibria have more than one phase present. For example, a gas and a solid or a liquid and a gas. How does the equilibrium constant differ for heterogeneous equilibria? Pure solids and liquids have activities of unity. Solvents in very dilute solutions have activities that are essentially unity. The Kc and Kp for the reaction shown above are: Heterogeneous Equlibria

  50. Heterogeneous Equlibria