Chemistry Post-lab Questions 6z4q6w

[Restoring Balance: LeChâtelier’s Principle and Equilibrium POST-LAB QUESTIONS] 1. Write the chemical equation for the complex-ion equilibrium that results when excess chloride ion is added to an aqueous solution of cobalt chloride. Note the observed color of each complex ion underneath its chemical formula. Co(H2O)62+(aq) + 4Cl-(aq) ↔ CoCl42-(aq) + 6H2O(l) Pink Blue 2. What is the likely composition of the solution (relative amounts of the two different complex ion forms) when the intermediate or transition color is observed in step 6? How does this observation provide visual proof of the idea that not all reactions “go to completion”? Explain. When the intermediate or transition color of lavender/purple was observed in step 6, we can qualitatively see that half of the total available Cobalt ions are in the form of Co(H2O)62+ (which takes on the pink color) and the other half of the Cobalt ions are in the form of CoCl42- (which takes on the blue color). What this also means is that at this point, half of the Co(H2O)62+ broke down to create half of the total amount of CoCl42-. Visually, we see that this makes sense because at this phase, the color seems to have reached a midpoint of red and blue (since it takes on a lavender/purple color). Because of this, we know that both the reactants and products must be present at equilibrium which then in turn also means that the reaction does not necessarily “go to completion”. 3. Use LeChâtelier’s Principle to explain the color changes observed upon addition of water and calcium chloride to an equilibrium mixture of the two complex ions in this reaction (steps 7 and 8). For the calcium chloride, because an excess reactant (CaCl2) was added, the equilibrium equation of Co(H2O)62+(aq) + 4Cl-(aq) ↔ CoCl42-(aq) + 6H2O(l) was shifted to the right so that more reactants could be consumed and therefore restore equilibrium. Algebraically, we see that this makes sense because the increase in reactants creates a “Q” value that is less than “Keq”. Because Keq > Q, the equation must shift right in accordance to LeChâtelier’s Principle. For the water, because an excess product (H2O) was added, the equilibrium equation of Co(H2O)62+(aq) + 4Cl-(aq) ↔ CoCl42-(aq) + 6H2O(l) was shifted to the left so that more of the excess products could be consumed and therefore restore equilibrium. Algebraically, we see that this makes sense because the increase in reactants creates a “Q” value that is greater than “Keq”. Because Keq < Q, the equation must shift right in accordance to LeChâtelier’s Principle. 4. What was the effect of adding AgNO3 on the position of equilibrium for these two complex ions? Is this effect consistent with LeChâtelier’s Principle? Explain. When AgNO3 was added, it broke down into ions and the Ag+ ions formed a bond with the Cl- ions from the equilibrium equation to form a precipitate. It is important to note that the precipitate that was formed through the addition of AgNO 3 is AgCl. Because the Ag+ ions formed a bond with the Cl- ions, we see that you have now removed reactants from the equilibrium equation, in this case the Cl - ion. As mentioned before in an earlier problem, because reactants have been removed, the equilibrium equation shifts left toward the formation of reactants so that a balance of order between the reactants and products can be formed. This is indeed consistent with LeChâtelier’s Principle. 5. How was the composition of the solution (relative amounts of the two complex ions) affected when the solution was heated (step 10)? When the solution was cooled (step 11)?

[Restoring Balance: LeChâtelier’s Principle and Equilibrium POST-LAB QUESTIONS] When the pink solution was heated (which contained the Co(H 2O)62+), we saw that the solution itself turned blue. When the blue solution (which now contained CoCl 42-) was cooled, the entire solution turned back to its original pink color. 6. Based on the observed effect of temperature on the position of equilibrium, is the forward reaction for the equation in Question # 1 endothermic or exothermic? Explain, using LeChâtelier’s Principle. Based off of these qualitative observations, it is quite obvious to conclude that the forward reaction is endothermic. However, in of LeChâtelier’s Principle, we see that increasing the temperature of this endothermic reaction shifted the entire equilibrium equation to the right to create more products so that more heat could be consumed. Contrarily, cooling down the solution reversed this effect.