Chemical Equilibrium Answer Key

Name: _________________________KEY_______________________________________________ Recitation Section: __________________________ Group Number: _________________________ Chemical Equilibrium Instructions: • Listen to an introduction from your recitation instructor. • Work through the entire activity with your group. Each group member is responsible for contributing equally to this activity, and each person will turn in an assignment following the activity. • No one is allowed to leave before the recitation period is over. If your group completes the activity early, use the extra time to help other groups. • The recitation instructors will wrap up the session with an end-of-class 5-minute discussion. • At the end of the class, one student from each table should gather the individual activity packets and turn them in to the recitation instructor as a set. Each activity packet should be complete. Concepts and Skills: Chemical Equilibrium The Equilibrium Constant Le Chatelier’s Principle Buffered Solutions For Further Practice Chapter 12, Section 12.4 Chapter 12, Section 12.5 Chapter 12, Section 12.6 Chapter 13, Section 13.6 -1- Part 1: Defining Equilibrium • • A state of equilibrium is reached when the concentrations of reactants and products no longer change. The equilibrium constant (Keq) is defined as the ratio of the concentrations of products to reactants when the system is at equilibrium. For the generic equation: aA + bB ⇌ cC + dD [𝐂𝐂]𝒄𝒄 [𝐃𝐃]𝒅𝒅 The equilibrium constant expression is: 𝑲𝑲eq = [𝐀𝐀]𝒂𝒂[𝐁𝐁]𝒃𝒃 [equilibrium concentrations] 1. Consider the reaction of nitrogen dioxide to form dinitrogen tetroxide. In a closed system, this reaction establishes an equilibrium: 2NO2(g) ⇌ N2O4(g) The data table below shows measured concentration of NO2 over a period of time in a system that starts out with only NO2. (a) Use stoichiometry to calculate the concentrations of N2O4 at each time. Then calculate the ratio of [N2O4]/[NO2]2. Measurement # Concentration of NO2(g), (M) 1.00 Concentration of N2O4(g), (M) 0.00 [N2O4]/[NO2]2 (no units) 0.00 0.80 0.10 0.16 3 0.50 0.25 1.0 4 0.40 0.30 1.9 5 0.30 0.35 3.9 6 0.20 0.40 10 7 0.15 0.425 19 8 0.10 0.45 45 9 0.10 0.45 45 10 0.10 0.45 45 1 (Time = 0.00 minutes) 2 (b) At what measurement number(s) is equilibrium reached? Explain your reasoning. Equilibrium is reached at measurement #8 and continues through measurement #10 (c) What is the value of the equilibrium constant, Keq, for this reaction? 45 -2- Part 2: Le Chatelier’s Principle: Predicting Shifts in Equilibrium Le Chatelier’s Principle: When a system at equilibrium is stressed, so that it is no longer is at equilibrium, the reaction system shifts in the forward or reverse direction to reestablish equilibrium. Stresses include: • Adding or removing a reactant or product from the system • Changing the temperature (which changed Keq value) 2. Consider the following chemical equilibrium that is endothermic as written: Heat + Co(H2O)62+(aq) + 4Cl–(aq) ⇌ [CoCl4]2–(aq) + 6H2O(l) Pink Blue At equilibrium at a particular temperature, the concentration of Co(H2O)62+ is slightly greater than the concentration of [CoCl4]2–. The color of the solution is therefore a light purple. Predict what will happen when NaCl is added to the solution. Remember that NaCl(aq) consists of Na+(aq) and Cl–(aq) ions. (a) The reaction will shift _____to the products______________________________________________ (b) Concentration of Co(H2O)62+(aq) will ______decrease______________________________________ (c) Concentration of [CoCl4]2–(aq) will ______increase_______________________________________ (d) The color of the solution will become more ____blue______________________________________ 3. Consider again the following chemical equilibrium that is endothermic as written: Heat + Co(H2O)62+(aq) + 4Cl–(aq) ⇌ [CoCl4]2–(aq) + 6H2O(l) Pink Blue Predict what will happen when the temperature of the equilibrium system is reduced. This is equivalent to removing “heat” from the system. (a) The reaction will shift _____to the reactants_____________________________________________ (b) Concentration of Co(H2O)62+(aq) will _____increase_______________________________________ (c) Concentration of [CoCl4]2–(aq) will _____decrease________________________________________ (d) The color of the solution will become more _____pink____________________________________ (e) Does the value of Keq increase or decrease? Remember that Keq is a ratio of products over reactants at equilibrium. _________at equilibrium the constant will remain the same_______________ -3- 4. Some equilibrium systems involve H3O+(aq), which is the same as H+(aq) in solution, or OH–(aq). When this is the case, shifts in equilibrium can affect the acidity, or pH, of the system. • The greater the concentration of H3O+, the more acidic the solution and the lower the pH. • When the concentration of H3O+is reduced, the solution is less acidic and the pH is higher. Consider the following equilibrium which involves the reaction of a weak acid with water: H3PO4(aq) + H2O(l)(aq) ⇌ H2PO4–(aq) + H3O+(aq) Use Le Chatelier’s Principle to predict what will happen when the concentration of H2PO4–(aq) is increased by adding the salt Na+H2PO4–. (a) The reaction will shift _____to the reactants_____________________________________________ (b) Concentration of H3O+(aq) will ___decrease__________________________________________ (c) The acidity of the solution will ______decrease___________________________________________ (d) The pH will _____increase________________________________________ 5. Consider the following set of two equilibria that occur in our blood: H3O+(aq) + HCO3–(aq) ⇌ H2CO3(aq) + H2O(l) ⇌ 2H2O(l) + CO2(aq) Use Le Chatelier’s Principle to predict what will happen after additional H3O+ enters our blood from excessive exercise. (a) Both reactions will shift ______to the products_______________________________________ (b) At the new equilibrium, the concentration of H3O+(aq) will be __greater___________ than before the addition of extra H3O+. (c) At the new equilibrium, the concentration of CO2(aq) will be _____greater_________ than before the addition of extra H3O+. 6. A buffer system consists of a weak acid and its conjugate base (such as H2CO3 and HCO3–). A buffer system is one that prevents large changes in H3O+ concentration (and therefore pH). The concentration of H3O+ in our blood must be maintained at a relatively constant value so that the pH stays at about 7.4. An increase in H3O+ will lower the pH, and a decrease in H3O+ will raise the blood pH. Death can result if the pH falls below 6.8 or rises above 7.8.  Explain, in terms of Le Chatelier’s principle, how this set of equilibria involving a buffer system prevents large changes in pH (or H3O+ concentration). If the concentration of H3O+ increases in our blood, the equilibrium will shift away from the excess H3O+ to maintain the correct pH; and, if the concentration of H3O+ decreases in our blood, the equilibrium will shift toward the H3O+ to maintain the correct pH. -4-