Cobalt Complex Ions Lab

Cobalt Complex Ions Lab:

Purpose:

Examine factors that shift equilibrium.

Procedure:

1. Label 4 test tubes P1, P2, B1, and B2.
2. Measure 5 mL of pink 0.1 M CoCl2 aqueous solution into each of the four test tubes.
3. Add 5 mL of 12M HCl to test tubes B1 and B2.
4. Set aside test tubes P1 and B1 as control solutions.
5. Fill a 400-mL beaker with tap water until it is half full. Use a hot plate to heat the water to 80–85˚C.
6. Fill a second 400-mL beaker with tap water and cool the filled beaker in an ice bath.
7. Add 5 mL of 12M HCl drop wise to test tube P2 until the solution turns blue.
8. Add 5 mL of distilled water drop wise to test tube P2 until the solution reverts back to pink.
9. Place test tube P2 in the hot water bath. The solution will gradually change from pink to lavender blue.
10. Remove test tube P2 from the hot water bath and immerse it in the ice water bath. The solution should revert back to pink.
11. Add 2 mL of 0.1 M AgNO3 solution to test tube B2. The solution should turn pink and a white buoyant precipitate will form.

Discussion of Theory:

        Chemical equilibrium is dynamic and can be shifted towards the products or towards the reactants based on external factors (such as changing concentration, temperature, or pressure) acting on system. At equilibrium the concentrations of reactants and products are constant because the formation of the products and the dissociation of the products into the reactants are happening at equal rates. LeChâtelier’s Principle states that “if the conditions of a system, initially at equilibrium, are changed, the equilibrium will shift in such a direction as to tend to restore the original conditions.”

[pic 1][pic 2]

A solution of cobalt(II) ion in water is pink, but when hydrochloric acid is added to the solution, the color changes to blue. This color shift is due to the formation of complex ions  of Co2+ and chloride ions, or of Co2+ and water. The formation of complex ions of   Co2+ and water molecules or chloride ions, respectively, is reversible, and quickly reaches equilibrium. Adding HCl’s Cl– ion shifts the equilibrium to the right to consume some of the added reactant. Diluting the blue solution with water decreases the concentration of the products and effectively shifts the equilibrium back to the left. Because Equation 1 is an endothermic reaction, the heat can be considered a reactant in the equation. Adding heat to the system increases the “heat reactant” and shifts Equation 1’s equilibrium to the right (blue), while removing heat through the ice bath shifts it back to the left (pink). Adding AgNO3 to the blue CoCl4 2– solution forms white AgCl precipitate ( Ag+(aq) + Cl–(aq) → AgCl(s) ), and a pink solution of Co(H2O)6 2+. Forming a precipitate containing the Cl- ion removes some of the Cl-(aq) ion from the solution and so decreases the concentration of the Cl-(aq) ion, and shifts Equation 1’s equilibrium to the left, toward reactant formation.