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E q u i l i b r i u m

E q u i l i b r i u m. Chapter 13. What Is It?. The state where concentrations of all reactants and products remain constant with time. At the molecular level, the reaction continues. Macroscopically, the reaction appears static. Why Is It?. Reaction rate depends on concentration.

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E q u i l i b r i u m

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  1. Equilibrium Chapter 13

  2. What Is It? The state where concentrations of all reactants and products remain constant with time. At the molecular level, the reaction continues. Macroscopically, the reaction appears static.

  3. Why Is It? • Reaction rate depends on concentration. • Collision Theory says more collisions = faster reactions. • As the reaction progresses, the concentration of products begins to increase.

  4. Equilibrium Position is Determined By: • Initial Concentrations • Energies of reactants and products • Organization (disorder) of reactants and products The Haber Process N2(g) + 3H2(g) 2NH3(g) Nothing appears to change because the rates of reaction are so naturally slow. A catalyst is used to begin the process for commercial production of ammonia. Haber applied Le Chatelier’s Principle to maximize the forward reaction.

  5. Le Chatelier’s Principle When a stress is applied to a system at equilibrium, the reaction shifts to relieve the stress.

  6. What stresses out a Reaction? • Heat (heat is measured in kJ or kcal) • Pressure (related to # of moles on each side of the reaction) • Concentration (how much of each component is added) …When any of these increase or decrease from original conditions, that is stress.

  7. The Equilibrium Constant • The equilibrium constant is represented by K Where: wA + xB yC + zD

  8. Try Me! Write the equilibrium expression for the following reaction: 4 NH3 + 7O2 4NO2 + 6H2O Calculate the value of the equilibrium constant if the concentrations of the reactants and products are as follows: [NH3] = 3.1 x 10-2 mol/L [O2] = 5.4 x 10-2 mol/L [NO2] =3.1 x 10-2 mol/L [H2O] = 4.7 x 10-2 mol/L

  9. Equilibrium Position • A set of concentrations that indicate whether products or reactants dominate while the rxn is at equilibrium. Number of equilibrium positions available Number of values for K 1

  10. Equilibrium and Pressures • Pressure and concentration are interchangeable PV = nRT P = (n/V)RT n/V = concentration K Kp vs In General: Kp = K(RT)Dn Do Now! Figure out the relationship between K and Kp for the Haber Process. Dn = Sproducts -Sreactants

  11. Heterogeneous Equilibria • Pure liquids and pure solids are not included in the equilibrium expression for a reaction. • Concentrations of PURE liquids and solids cannot change.

  12. Applications of the Equilibrium Constant • The Equilibrium constant gives information that will allow us to… • Decide how likely it is that the reaction will occur. • Determine if a reaction is at equilibrium given a set of concentrations. • Determine which direction the reaction must shift in order to reach an equilibrium position. If K is greater than 1 The reaction is much More likely to occur Spontaneously (the equilibrium lies farthest to the right) If K is very small The reaction is not likely to be spontaneous (the equilibrium is near the reactants) Spontaneous Does Not Mean Fast!

  13. The Reaction Quotient Q • Obtained by replacing initial concentrations into the concentrations of the equilibrium expression. • Three possible cases: Q > K Shift Left Q = K No Shift Q < K Shift Right

  14. Try Me!! • For the synthesis of ammonia at 500oC, the equilibrium constant is 6.0 x 10-2. Predict the direction in which the system will shift to reach equilibrium: [NH3]o = 1.0 x 10-3 M [N2]o = 1.0 x 10-5 M [H2]o = 2.0 x 10-3 M Q = 1.3 x 107 Shift to the Left

  15. Try Me Again and Again! • Same problem, different conditions: [NH3]o = 2.0 x 10-4 M [N2]o = 1.5 x 10-5 M [H2]o = 3.54 x 10-1 M [NH3]o = 1.0 x 10-4 M [N2]o = 5.0 M [H2]o = 1.0 x 10-2 M Q = 6.01 x 10-2 No Shift Q = 2.0 x 10-3 Shift to the Right

  16. Solving for Concentrations and Pressures • Several Types of problems and Solving methods. • Plug and Chug • ICE • ICE with Stoichiometry • ICE with the quadratic equation

  17. Type 1: Plug and Chug Consider an experiment in which gaseous N2O4 was placed in a flask and allowed to reach equilibrium at a temperature where Kp = 0.133. At equilibrium, the pressure of N2O4 was found to be 2.71 atm. Calculate the equilibrium pressure of NO2. N2O4(g) 2NO2(g)

  18. Type 2: Initial, Change,Equilibrium • At a certain temperature a 1.00L flask initially contained 0.298 mol PCl3(g) and 8.70 x10-3 mol PCl5(g). After the system had reached equilibrium, 2.00 x10-3 mol Cl2(g) was found in the flask. Gaseous PCl5 decomposes according to the reaction PCl5(g) PCl3(g) + Cl2(g) Calculate the concentrations of all species and the value of k.

  19. Type 3: ICE with Stoichiometry • Carbon Monoxide reacts with steam to produce carbon dioxide and hydrogen. At 700K the equilibrium constant is 5.10. Calculate the equilibrium concentrations of all species if 1.00 mol of each component is mixed in a 1.00L flask.

  20. Type 4: ICE with Quadratic • A 1.00L flask is filled with 1.0 mol H2 gas and 2.0 mol I2 gas at 448oC. The value of the equilibrium constant Kc for the reaction H2 + I2 2HI at 448oC is 50.5. What are the equilibrium concentrations of H2, I2 and HI in mol/L?

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