Le’ Chatelier’s Principle

Purpose: The purpose of this lab is to develop a deeper understanding of LeChatelier’s Principle by observing several systems at chemical equilibrium and interpreting the effects of varying concentrations and temperature. The principle states that if systems at equilibria are altered or disturbed in any form, the equilibria will shift to reduce the disturbing influence ( Catalyst, 186). In a 3 part experiment, we analyzed the outcome of changes in reactant and product concentrations, equilibrium involving sparingly soluble salts, and the effect of temperature on the equilibrium.In part 1 , we observed the shift in equilibria of two aqueous solutions of Copper and Ammonia then Nickel and Ammonia. In part 2, we focused on cobalt ions in the presence of chloride ions as well as the precipitation of silver nitrate and sodium carbonate. In the last part of the experiment we utilized a solution of Cobalt chloride and compared the color at room temperature and then again in a contain er of boiling water. Physical Data: No physical Data was applicable to the experiment. Chemical Equations: Part i: Changes in Reactant or Product Concentrations A. Copper and Nickel Ions [Cu(H2O)4]2+ (aq) + 4NH3(aq) [Cu(NH3)4]2+(aq) + 4H2O(l) blue dark blue †¢[Ni(H2O)6]2+(aq) + 6NH3(aq) [Ni(NH3)6]2+(aq) + 6H2O(l) green pale violet †¢H+(aq) + NH3(aq) NH4 +(aq) B. Cobalt Ions †¢[Co(H2O)6]2+(aq) + 4CL- (aq) [CoCl4]2-(aq) + 6H2O(l) Part ii: Equilibrium Involving Sparingly Soluble Salts †¢2AgNO3(aq) + Na2CO3(aq) Ag2CO3(s) + 2NaNO3(aq) †¢2Ag+(aq) + CO32-(aq) Ag2CO3(s) Net ionic equation ^ †¢2H+(aq) + CO32-(aq) H2CO3(aq); H2CO3(aq) > CO2(g) + H2O(l) Ag+(aq) + Cl-(aq)AgCl(s) †¢Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) †¢I-(aq) + Ag+(aq) AgI(s) Safety †¢Safety goggles are required to be worn throughout entire duration of the lab experiment. †¢Wear gloves, as the chemicals may cause serious damage to the skin skin. †¢Be sure to clean mat erials with soap and water before beginning any procedures. †¢When disposing wastes, be sure to do so in the appropriate receptacle. †¢Use precaution when handling all chemicals, careful not to inhale anything. Experimental Procedure and Observations Part i: Changes in Reactant or Product ConcentrationsA. Copper and Nickel Ions Procedure Copper 1. Place 1 mL of 0. 1 M CuSO4 in a clean test tube. 2. Add 15 M NH3 drop wise until a color change occurs. 3. Mix the solution in the test tube as you add the NH3. 4. Add 1 M HCl drop wise while mixing the solution, until the color changes. Nickel 1. Place about 1 mL of 0. 1 M NiCl2 in a clean test tube. 2. Add 15 M NH3 drop wise until a color change occurs. 3. Mix the solution in the test tube as you add the NH3. 4. Add 1 M HCl drop wise while mixing the solution, until the color changes. Observations Copper . The liquid is light blue in color. 2. The solution turned to royal blue. 3. Solution begins to slowly change to a more tran sparent blue. 4. We added 56 drops, the top of the solution remained royal blue as the bottom turned completely clear and colorless. After shaking it, it turned completely light blue. Nickel 1. The liquid is light/clear green in color. 2. The solution turned from green to blue to a lavender complex. 3. The solution turned to a clear lavender color. 4. The solution reverted back to clear green. B. Cobalt Ions Procedure 1. Place 0. mL of 1 M CoCl2 in a test tube. 2. Add 12 M HCl to test tube until a change is noticeable. 3. Slowly add water to the test tube while mixing. Observations 1. Exactly 10 drops are placed in the tube. The liquid is pale pink in color. 2. The solution turned to dark blue. 3. The solution slowly turns to purple, as little particles form on the bottom. A pale pink color began to form at the top and the color consumed the entire solution. Part ii: Equilibrium Involving Sparingly Soluble Salts Procedure 1. Add 10 drops of 0. 01 M AgNO3 to 0. 5 mL of 0. 1 M of Na2C O3. . With caution, add 6 M HNO3 drop wise until a change occurs. 3. Add . 1 M of HCl drop wise until a change is observed. 4. Add 15 M NH3 drop wise until a change occurs. 5. Add 6 M HNO3 drop wise until there is evidence of a chemical change. 6. While mixing the solution, add 15 M NH3 drop wise. 7. Add 0. 1 M KI drop wise until there is evidence of a chemical reaction. Observations 1. The original Na2CO3 solution is clear in color. The addition of AgNO3 turns it cloudy almost immediately. A small amount of precipitate is also visible in the solution. 2.Exactly 4 drops of HNO3 are added and the color of the solution reverts back to clear. 3. 4 drops of HCl are also added and the solution once again turns back cloudy with visible precipitate. 4. 15 drops of NH3 are used and the solution becomes colorless with the precipitate dissolving. 5. The solution remains colorless and a small gas cloud forms over the solution. 6. The solution is still clear and the gas above is still visible. 7. The solution turns white/ creamy in color. There is visible precipitate and the gas above the liquid is no longer visible. Part iii. Effect of Temperature on EquilibriaProcedure 1. Using a 250 ml beaker, heat 75ml of water until it begins to boil. 2. Place 1 mL of 1. 0 M CoCl2 in a test tube and place the test into the boiling water (Careful not to spill). Observations 1. The water heats to a temperature of about 135 °C. 2. The color of the CoCl2 at 20 °C is red. After placing it in the boiling water it changes to a deep pink/magenta color. Data/ Results Part i: A Part i: B Part ii Part iii Calculations: No mathematical calculations were applicable to the experiment. Discussion: Beginning with the first experiment, which consisted of the Copper, Nickel, and Ammonia.In both reactions, the strength of the ammonia is stronger than that of the water, causing each of them to dissociate. Once Hydrochloric acid is added to left of the equation, the ammonia binds to hydrogen forming ammonium and driving the reaction back in the direction that it came from. The equilibrium is therefore established by the Nickel ion and Ammonia and shifted by the hydrochloric acid once the hydrogen reacts with ammonia in a common acid-base reaction. The ammonia-metal bond in each of the reactions causes a precipitate to form because of the hydroxide ions that are left after the donation of the hydrogen.Part B of the experiment consisted of the aqueous Cobalt and chloride ions. The addition of the hydrochloric acid, once again induces an immediate change in color. The equilibrium of the equation is disturbed because of the acid, which lead to the left shift in the equation. Increasing the amount of water allowed H2O to act as a base forming H3O, allowing the reaction to move back to the right. In the second portion of the lab, the combination of silver and sodium carbonate leads to the formation of a precipitate. This is accounted for based on the silver+carbonate complex.Adding h ydrochloric acid forms an unstable carbonic acid which will later dissociate into carbon dioxide and water. This also has the effect of dissolving the silver carbonate and shifting the equation back to the left. Further removal of the silver on the left forces the reaction to move in the direction of the loss. Silver ions react with ammonia that is added and added more acid to this caused ammonium to form. Ammonia is added once more to reestablish the equilibrium. The final add-on of potassium iodide once again disrupts the balance because the silver reacts the iodide causing the reaction to move left.By manipulating the temperature, we were able to deduce information about the final reaction involving cobalt chloride. Starting near room temperature at exactly 20 °C the cobalt chloride started at a light pink color. After placing the solution in a heated water bath of exactly 135 °C, the contents of the test tube turned dark pink. The reaction is therefore endothermic as the coo l CoCl2 absorbed heat from it’s water bath before making a chemical change, therefore the reaction shifts to the right to absorb the heat. Conclusion:Conducting the experiment gave us the opportunity to learn about the effects of varying concentration and temperature in a system, hence the objectives were met because in performing each section of the lab, we were able to apply LeChatelier’s principle. The methods applied greatly aided in our understanding of the material as we had to apply previous knowledge to understand the behavior of the chemicals. Many of the solutions that were added drop wise had to be done that way as to not add too much because too much of a substance could prevent the reversal properties of the reaction.