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Chemical Equilibrium

Chemical Equilibrium. Chapter 6 pages 190 - 217. Reversible Reactions- most chemical reactions are reversible under the correct conditions. Equilibrium= . H 2 O(g) + CO(g)  H 2 (g) + CO 2 (g).

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Chemical Equilibrium

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  1. Chemical Equilibrium Chapter 6 pages 190 - 217

  2. Reversible Reactions- most chemical reactions are reversible under the correct conditions

  3. Equilibrium=

  4. H2O(g) + CO(g) H2(g) + CO2(g) The graph represents changes in concentrations with time for the reaction H2O(g) + CO(g)H2(g) + CO2(g) when equimolar quantities of H2O(g) and CO(g) are mixed.

  5. (a) H2O and CO are mixed in equal numbers and begin to react (b) to form CO2 and H2. After time has passed, equilibrium is reached (c) and the numbers of reactant and product molecules then remain constant over time (d).

  6. H2O(g) + CO(g)  H2(g) + CO2(g)orX  Y • At equilibrium, X forms Y at the same rate as Y forms X. • Therefore, the concentration of [X] and [Y] will remain constant. • Also, when equilibrium is reached, the rates of the forward reaction will equal the rates of the reverse reaction.

  7. The Equilibrium Constant, Keq(also called the mass action expression) • Two methods for describing the equilibrium of a reaction. • Kc- describes the equilibrium of a reaction where the concentrations of the materials is known. • Kp- describes the equilibrium of a gaseous reaction using partial pressures instead of concentrations.

  8. Kcc = concentration aA + bBcC + cD Coefficients in the chemical equation become exponents in the equilibrium constant expression. Include only substances in the gas or aqueous phase. Solid’s and liquid’s concentrations do not change during a chemical reaction.

  9. Write the equilibrium expression for the reaction below.Ni (s) + 4CO (g)  Ni(CO)4(g)

  10. Write the equilibrium expressions for the following reactions: • H2 (g) + I2 (g)  2HI (g) • 2NH3 (g) + 3 CuO(s)  3H2O (g) + N2 (g) + 3Cu (s) • NH4Cl (s)  NH3 (g) + HCl (g)

  11. Kpp= pressure • Equilibrium is described in terms of the partial pressures of the reactants and products. 2CO2 (g) 2CO (g) + O2 (g)

  12. Write the equilibrium constant for the following mixture of gases at equilibrium. 2NO (g) + O2 (g)  2NO2 (g)

  13. The Relationship Between Kc and Kp Where n = (total moles of gaseous products)-(total moles of gaseous reactants) R= ideal gas law constant (.08206 L atm / mol K T in Kelvin

  14. Example: A 3:1 starting mixture of hydrogen, H2, and nitrogen, N2, comes to equilibrium at 500oC. The mixture at equilibrium is 35.06% NH3, 96.143% N2 and 0.3506% H2 by volume. The total pressure in the reaction vessel was 50.0 atm. What is the value of Kp and Kc for this reaction? Kp = Kc =

  15. Example: 2NO (g) + Cl2(g)  2NOCl (g) In the above reaction, the total pressure of the mixture of gases at equilibrium is 1.55 atm. The percentages of each gas in the mixture areas follows: NOCl = 77.4%, NO = 3.20%, and Cl2 = 19.4%. Calculate Kp and Kc for the reaction.

  16. Interpretation of the Equilibrium Constant For K>1, we can state that the amount of product is greater than the amount of reactant. The equilibrium is favored to the right. Likewise, for K<1, the amount of reactants is much greater than the products. The equilibrium favored to the left.

  17. The Reaction Quotient- Q • There are times when a chemist may be given information about a chemical reaction that may or may not be at equilibrium. • A Q calculation can be used to determine the direction the reaction is heading to reach equilibrium. • If Q > K, then the reaction has too many products and is heading towards the reactant side. • If Q < K, then the reaction has too many reactants and is heading towards the product side. • If Q = K, then the reaction is at equilibrium.

  18. Example: N2 (g) + 3H2 (g)  2NH3 (g) In the reaction shown above, the value of Kc at 5000C is 6.0 X 10-2. At some point during the reaction, the concentrations of each material were measured. At this point, the concentrations of each substance were [N2] = 1.0 X 10-5 M, [H2] = 1.5 X 10-3M, and [NH3] = 1.5 X 10-3 M. Calculate the value of Q, and determine the direction that the reaction was most likely to proceed when the measurements were taken.

  19. Solving Problems When Not All Equilibrium Concentrations are Known H2 (g) + I2 (g)  2HI (g) When 2.00 mole each of hydrogen and iodine are mixed in an evacuated vessel, 3.50 mole of HI are produced. What is the value of the equilibrium constant, Kc? It is important to realize that not all of the reactants are converted to products in reversible reactions. Some reactants will remain at equilibrium.

  20. Solution: Careful Stoichiometry!

  21. Use the mole ratios to determine the change concentrations of the reactants.

  22. Since we are in 1.0 L, 0.25 mol = 0.25 M

  23. Use the final concentrations to calculate the value of Kc

  24. Example: Nitric oxide gas, NO, and oxygen gas, O2, react to form the poisonout gas nitrogen dioxide, NO2, in the reaction shown below: 2NO (g) + O2 (g)  2NO2 (g) 10.0 moles of NO and 6.00 moles of O2 are placed into an evacuated 1.00 L vessel, where they begin to react. At equilibrium, there are 8.80 moles of NO2 present. Calculate the value of Kc, assuming that the temperature remains constant throughout the reaction. Kc =

  25. Determining K When Only Initial Concentrations Are Known H2 (g) + I2(g)  2HI 0.500 mol of hydrogen and 0.500 mol of iodine are added to a 1.00 liter reaction vessel. The mixture is heated to 4980C and allowed to reach equilibrium according to the reaction shown above. At this temperature, Kc = 49.7. What is the composition of the reaction mixture at equilibrium in this system?

  26. H2 (g) + I2(g)  2HI

  27. Enter the Quadratic Formula

  28. Example: 0.500 moles of NO gas are placed into a 1.00 L reaction vessel. The gas is heated to an extremely high temperature, where it decomposes according to the reaction shown below: 2NO (g) N2 (g) + O2 (g) At equilibrium, Kc = 2.4 X 103. Determine the equilibrium composition of the mixture. 5.1 X 10-3 M NO, 0.25 M N2, 0.25 M O2

  29. Le Chatelier’s Principle • When a system at equilibrium is disturbed, it will react in a way that minimizes the disturbance. • Factors that can influence the position of an equilibrium. • 1. Concentration (aq or g) • 2. Temperature • 3. Pressure (gases) • 4. Addition of a catalyst

  30. Example: Describe how the reaction below adjusts its equilibrium in each of the scenarios. CO (g) + 2H2 (g)  CH3OH (g) H0 = -21.7 Kcal • Some of the methanol vapor is condensed and removed from the reaction vessel. • The pressure is increased by decreasing the volume of the reaction vessel. • The temperature is increased.

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