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Chemical Equilibrium

Chemical Equilibrium

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Chemical Equilibrium

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  1. Chemical Equilibrium Portions adapted from

  2. Chemical Equilibrium POGIL Partner (Work TOGETHER!) Group 1: Mackenzie, Austin and Odessa Group 2: Chetha, Paige and Cass Group 3: Libby, Cassie and Kyli Group 4: Shae, Manu and Madi Search for/identify “Key Ideas” Share out responses and Key Ideas 35 minutes

  3. Chemical Equilibrium • Reversible Reactions: • A chemical reaction in which the products can react to re-form the reactants • Chemical Equilibrium: • When the rate of the forward reaction equals the rate of the reverse reaction and the concentration of products and reactants remains unchanged A(g) B(g)

  4. The Equilibrium Condition • Dynamic Chemical Equilibrium • Reversible Reactions

  5. The Haber Process (Fritz Haber, Germany 1913) N2(g) + 3H2(g)  2NH3(g) ΔH°rxn = -91.8 kJ Nitrogen is very unreactive—why? Two reasons why the concentrations of reactants and products remain unchanged when mixed: • The system is at chemical equilibrium. • The forward and reverse reactions are so slow that the system moves toward equilibrium at a rate that cannot be detected.

  6. The Equilibrium Constant (Keq or Kc) aA + bB cC + dD • At a given temperature, Keqfor a reaction is always the same. • Keq has no units. (K, Kc, or Keqmay be used) • The numerical value of K indicates the extent to which the reactants are converted to products. • K>>1 Signifies the reaction is “product favored” • K=1 • K<<1 Signifies the reaction is “reactant favored”

  7. Write the equilibrium expression for the following reactions: 4NH3(g) + 7O2(g)  4NO2(g) + 6H20(g) C3H8(g) + O2(g)  CO2(g) + H2O(g) PCl5(g)  PCl3(g) + Cl2(g)

  8. The following equilibrium concentrations were observed for the Haber process at 127°C : [NH3] = 3.1 x 10-2mol/L [N2] = 8.5 x 10-1mol/L [H2] = 3.1 x 10-3mol/L a. Calculate the value of Keqat 127°C for this reaction. N2(g) + 3 H2(g) ↔ 2 NH3(g) b. Calculate the value of the equilibrium constant at 127°C for the reaction 2 NH3(g) ↔ N2(g) + 3 H2(g)

  9. Manipulating K Expressions • If all of the coefficients in a balanced equation are multiplied by a number, each term in the K expression must be raised to that power. Knew = (Koriginal)n • The K for a reverse reaction is the reciprocal of the K for the forward reaction. • For a reaction which is the sum of two or more reactions, the overall K is the product of the expressions for K for the individual steps: Koverall = K1 x K2 x K3 …

  10. Calculate the value of the equilibrium constant at 127°C for the reaction given by the equation ½ N2(g)+ 3/2 H2(g)  NH3(g)

  11. At 25°C, the following reactions have the equilibrium constants shown: 2 CO(g) + O2(g) ↔ 2 CO2(g) Kc = 3.3 x 1091 2 H2(g) + O2(g) ↔ 2 H2O(g) Kc = 9.1 x 1080 • Use this data to calculate Kc for the reaction CO(g) + H2O(g) ↔ CO2(g) + H2(g)

  12. The following results were collected for two experiments involving the reaction at 600°C between gaseous sulfur dioxide and oxygen to form gaseous sulfur trioxide: Experiment 1Experiment 2 • Show that the equilibrium constant is the same in both cases.

  13. Equilibrium Expressions Involving Gases & Pressure (Kp) PV=nRT For the gas phase reaction: 3H2(g) + N2(g)  2NH3(g) PNH3, PN2, PH2 are equilibrium partial pressures

  14. Example The reaction for the formation of nitrosyl chloride 2 NO(g) + Cl2(g) ↔ 2 NOCl(g) Was studied at 25°C. The pressures at equilibrium were found to be PNOCl = 1.2 atm PNO= 5.0 x 10-2atm PCl2= 3.0 x 10-1atm • Calculate the value of Kp for this reaction at 25°C.

  15. Relationship between Keq and Kp • NOT INTERCHANGEABLE!! Kp = Keq(RT)Δn • Δn= moles of gas product – moles of gas reactants • Sum of the products coefficients minus the sum of the reactants coefficients. • What if there is no change in the number of moles of gas?

  16. Using the value of Kp obtained in previous example problem, calculate the value of K at 25°C for the reaction: 2 NO(g) + Cl2(g) ↔ 2 NOCl(g)

  17. Heterogeneous Equilibria • The position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present • The concentrations of pure liquids and solids are constant. • Write the equilibrium expression for the reaction: PCl5(s)  PCl3(l) + Cl2(g)

  18. Write the expressions for K and Kp for the following processes: a. CaCO3(s) ↔ CaO(s) + CO2(g) b. 2 CO(g) ↔ CO2(g) + C(s) c. 2 Hg(l) + Cl2(g) ↔ Hg2Cl2(s) d. NH3(g) + HCl(g) ↔ NH4Cl(s) e. Ag2CrO4(s) ↔ 2 Ag+(aq) + CrO42-(aq) – solubility of a salt f. 2 NaHCO3(s) ↔ Na2CO3(s) + H2O(g) + CO2(g) g. CuSO4∙5H2O(s) ↔ CuSO4(s) + 5 H2O(g)

  19. Calculate Kc for the following reaction: CaCO3(s) ↔ CaO(s) + CO2(g) Kp = 2.1 x 10-4 at 1000 K

  20. Individual Practice Part 1 • Unit 6 Review Packet

  21. Applications of the Equilibrium Constant (Reaction Quotient, Q) • Reaction Quotient, Q aA + bB ↔ cC + dD Qc= [C]c[D]d / [A]a[B]b Qp= PCcPDd / PAaPBb • Q < K The systems shifts to the right, consuming reactants and forming products until equilibrium is achieved • Q > K The system shifts to the left, consuming products and forming reactants until equilibrium is achieved • Q = K The system is at equilibrium

  22. Example • For the synthesis of ammonia at 500°C, the equilibrium constant is 6.0 x 10-2. Predict the direction in which the system will shift to reach equilibrium in each of the following cases: a. [NH3]0 = 1.0 x 10-3 M; [N2]0 = 1.0 x 10-5 M; [H2]0 = 2.0 x 10-3 M A: shifts left b. [NH3]0 = 2.00 x 10-4 M; [N2]0 = 1.50 x 10-5 M; [H2]0 = 3.54 x 10-1 M A: no shift c. [NH3]0 = 1.0 x 10-4 M; [N2]0 = 5.0 M; [H2]0 = 1.0 x 10-2 M A: shifts right

  23. Solving for Equilibrium Concentration (ICE) • Step #1: Write the K or equilibrium expression • Step #2: “ICE”, begin with the Initial concentrations • Step#3: Then plug equilibrium concentrations into our equilibrium expression, and solve for x • Step#4: Substitute x into our equilibrium concentrations to find the actual concentrations

  24. Equilibrium Calculations (ICE) • 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(g). • Answer: 0.600atm

  25. Equilibrium Calculations (ICE) • At a certain temperature in a 1.00-L flask initially contained 0.298 mol PCl3(g) and 8.70 x 10-3mol PCl5(g). After the system had reached equilibrium, 2.00 x 10-3mol Cl2(g) was found in the flask. Gaseous PCl5 decomposes according to the following reaction PCl5(g) ↔ PCl3(g) + Cl2(g) Calculate the equilibrium concentrations of all species and the value of K. • A: [Cl2] = 2.00 x 10-3 M • [PCl3] = 0.300 M • [PCl5] = 6.70 x 10-3 M • K = 8.96 x 10-2

  26. Equilibrium Calculations (ICE) • Carbon monoxide reacts with steam to produce carbon dioxide and hydrogen. At 700 K the equilibrium constant is 5.10. Calculate the equilibrium concentrations of all the species if 1.000 mol of each component is mixed in a 1.000-L flask. • A: [CO] = [H2O] = 0.613 M • [CO2] = [H2] = 1.387 M

  27. Equilibrium Calculations (ICE) • Assume that the reaction for the formation of gaseous hydrogen fluoride from hydrogen and fluorine has an equilibrium constant of 1.15 x 102 at a certain temperature. In a particular experiment, 3.000 mol of each component was added to a 1.500-L flask. Calculate the equilibrium concentrations of all species. • A: [H2] = [F2] = 0.472 M • [HF] = 5.056 M

  28. Equilibrium Calculations (ICE) • What are the steps to solving equilibrium problems?

  29. Equilibrium Calculations • Assume that gaseous hydrogen iodide is synthesized from hydrogen gas and iodine vapor at a temperature where the equilibrium constant is 1.00 x 102. Suppose HI at 5.000 x 10-1atm, H2 at 1.000 x 10-2atm, and I2 at 5.000 x 10-3atm are mixed in a 5.000-L flask. Calculate the equilibrium pressures of all species. • A: PHI = 4.29 x 10-1atm • PH2 = 4.55 x 10-2atm • PI2 = 4.05 x 10-2atm

  30. Le Chatelier’s Principle POGIL Partner (Work TOGETHER!) Search for/identify “Key Ideas” Share out responses and Key Ideas _________minutes Video Review and Video ReviewVideo Review Part 2

  31. Le Chatelier Activity and Lab Partners • Group 1: Paige and Cass • Group 2: Libby and Mackenzie • Group 3: Shae and Kyli • Group 4: Odessa and Austin • Group 5: Manu and Madeline • Group 6: Chetha and Cassie

  32. Stressing a system If a stress is applied to a system, the system will shift in the direction that relieves the stress Change in Concentration Change in Volume Change in Pressure Change in Temperature

  33. Le Châtelier’s Principle: • 1. The effect of a change in concentration • Example: Predict the direction of the shift of the equilibrium position in response to each of the following changes in conditions: As4O6(s) + 6 C(s) ↔ As4(g) + 6 CO(g) a. Addition of carbon monoxide. b. Addition or removal of carbon or tetraarsenichexoxide. c. Removal of gaseous arsenic (As4).

  34. Changes in Concentration Option A: If you increase the reactant, the reaction shiftsto the rightto produce more product. Option B: If you increasethe product, the reaction shifts to the leftto produce more reactants.

  35. Le Châtelier’s Principle: • 2. The Effect of a Change in Pressure Three ways to change the pressure of a reaction system involving gaseous components: 1. Add or remove a gaseous reactant or product. 2. Add an inert gas (one not involved in the reaction) 3.Change the volume of the container

  36. Change in Volume/Pressure • Look at equation and pay attention to only gaseous substances (g) • If you increase the pressure (ie. decrease the volume) the reaction will shift to the side with fewest gas particles

  37. Pressure Example • Predict the shift in equilibrium position that will occur for each of the following processes when the volume is reduced. a. P4(s) + 6 Cl2(g) ↔ 4 PCl3(l) 3 b. PCl3(g) + Cl2(g) ↔ PCl5(g) c. PCl3(g) + NH3(g) ↔ P(NH2)3(g) + 3 HCl(g)

  38. Le Châtelier’s Principle: 3. The Effect of a Change in Temperature Remember, the value of K changes with temperature! Changes in concentration and pressure do not alter the equilibrium constant, K. N2(g) + 3 H2(g) ↔ 2 NH3(g) + 92 kJ 556 kJ + CaCO (s) ↔ CaO(s) + CO2(g)

  39. Changes in Temperature Determine if reaction is exothermic or endothermic. -Exothermic—heat is a product -Endothermic—heat is a reactant Exothermic -Add heat, left shift -Remove heat, right shift Endothermic -Add heat, right shift -Remove heat, left shift

  40. Refresher • The position of equilibrium also changes if you change the temperature. According to Le Chatelier's Principle, the position of equilibrium moves in such a way as to tend to undo the change that you have made. • If you increase the temperature, the position of equilibrium will move in such a way as to reduce the temperature again. It will do that by favoring the reaction which absorbs heat.

  41. Temperature Example For each of the following reactions, predict how the value of K changes as the temperature is increased. a. N2(g) + O2(g) ↔ 2 NO(g) ΔH° = 181 kJ b. 2 SO2(g) + O2(g) ↔ 2 SO3(g) ΔH° = –198 kJ

  42. Le Châtelier’s Principle: • 4. The effect of adding a catalyst • Example: The reaction N2O4(g)↔ 2 NO2(g) Is endothermic, with ΔH° = +56.9 kJ. How will the amount of NO2 at equilibrium be affected by: a. adding N2O4 b. increasing the volume of the container c. raising the temperature d. adding a catalyst to the system e. Which of these changes alters the value of Kc?

  43. Practice Catalyst Iron (III) oxide reacts with carbon monoxide in a blast furnace, reducing it to iron metal: Fe2O3(s) + 3 CO(g) ↔ 2 Fe(l) + 3 CO2(g) • Use Le Châtelier’s Principle to predict the direction of net reaction when an equilibrium mixture is disturbed by: a. removing CO2 b. removing CO c. adding CO d. adding Fe2O3 e. decreasing the volume of the container f. Which of these changes alters the value of Kc?

  44. Let’s Practice Concentration: N2 (g) + 3H2 (g) 2NH3 (g) Q: What type of shift will occur when the following changes are made to the system? • Removing hydrogen from the system • Adding ammonia to the system • Adding hydrogen to the system

  45. Let’s Practice…Again Pressure/Volume: 2SO2 (g) + O2 (g) 2SO3 (g) Q: What type of shift will occur when the following changes are made to the system? • How many gas particles are on the reactant and product side of the reaction? • You increase the volume? • You increase the pressure?

  46. Let’s Practice…Last One Temperature: C2H2 (g) + H2O (g) CH3CHO (g) ΔH = -151 kJ Q: Would you raise or lower the temperature to obtain the following results? • An increase in the amount of CH3CHO (g) • A decrease in the amount of C2H2 (g) • An increase in the amount of H2O (g)

  47. Individual Practice • Unit 6 Review Packet Equilibrium