Le Chateliers Principle in a Cobalt Complex Worksheet PDF

Preview

Summary

Second term lab of chemistry sequence...

Description

Le Chatelier’s Principle in a Cobalt Complex Worksheet Name: Hope Robertson

TA: Sam Mumford

Table: Observations of manipulation of the cobalt complex Color Dark blue, very rich viscous solution [CoCl4]2-  (aq) [Co(H2O)6] 2+  (aq)

Pink lemonade color; translucent light pink/orange color

Step 3 Color change and observations (addition of HCl)

Pink color to light blue to dark blue.

Step 5 Color change and observations (addition of H₂O) Step 6 Color change and observations (Heating) Step 7 Color change and observations (Cooling)

causes the dark blue to become more clear and translucent

Step 8 Color Change and observations (Addition of AgNO₃ 0.1M)

solution with H₂O, turned very dark blue solution of just HCl and CoCl₄, very dark blue solution with H₂O, light blue → to clear/transparent → Light purple color solution of just HCl and CoCl₄, very dark blue → lighter blue but not very big change dark blue turns to light purple/pink and becomes more translucent. a white/pink precipitate forms

Questions: 1. What was the effect of adding excess chloride ions. Use Le Chatlier’s principle and provide evidence. The addition of excess chloride ions saturates the [Co(H 2 O)6 ] 2+ ”pushing” the equilibrium of the reaction to the left (see equation 1) creating the blue color (that is the CoCl4). 2. Based upon the heating and cooling of the two equilibrium mixtures, propose if the reaction is endothermic or exothermic. Use Le Chatlier’s principle and provide evidence (would heat be considered as a reactant or product?). 1) [CoCl 4]2− + 6H 2 O ↔ [Co(H 2 O)6 ]2+ + 4Cl − When heat was added the complex turned a deep shade of blue, pushing the equilibrium to the left. When solutions were cooled the diluted solution turned almost completely clear/purple and the other a lighter shade of blue (see table 1) pushing the equilibrium to the right side of the reaction.. Heat is considered a product of this reaction, an exothermic reaction.

3. How did the addition of silver nitrate affect the equilibrium if neither silver ions nor nitrate ions are in the equilibrium expression? Use Le Chatlier’s principle and provide evidence. Additionally, write a net ionic equation to describe the precipitation reaction. The addition of silver nitrate forms a precipitate, the silver binds with the chloride ions and the cobalt with the nitrate ions. With the precipitate formation there is a removal of chloride ions, changing the concentration and then moving the equilibrium to the left. The new net ionic equation: AgN O3 + C oCl4 = AgCl + CoN O3 + 3Cl 4. When perturbing the equilibrium with heating and cooling, how many times do you think the equilibrium can be shifted before it stops working? Why? How about modification of the equilibrium through changes in concentration? Heat changes the internal energy of the substances in the solution, therefore the equilibrium “pushes” either way infinitely. If we changed the concentrations (like we did with the addition of the silver nitrate) it would stop until it is diluted completely....