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 container 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,
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 container 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,