Purpose.
To determine which way a reaction will shift when certain stresses are placed upon the system and to see the effects these stresses have on the equilibrium of the system.
Procedure.
1. Place a beaker with about 100mL of water onto a hot plate and turn the hot plate on to a medium setting so the water does not boil but is hot.
2. Label a well from left to right across the top "1-6" and down the left side "A-D". Place the well on a white sheet of paper so you can see the effects of the system trying to reach equilibrium.
3. Get the four pipets of CoCl2 solution, HCl, H2O, and Ag.
4. Put five drops of CoCl2 into each of the 24 wells on the white sheet of paper.
5. Add two drops of HCl to the CoCl2 in wells A1, B1, C1, and D1. Add four drops of HCl into the wells A2, B2, C2, D2. Add six drops into each well in column 3, eight drops into each well in column 4, ten drops into each well in column 5, and twelve drops into each well in column 6. Record all observations in a data table.
6. In row B, add one more drop of HCl to each well. Record all observations in a data table.
7. Add five drops of distilled water into each well in row C. Record all observations in a data table.
8. Add five drops of the AgNO3 solution to each well in row D. Record all observations in a data table.
9. Place 5mL of cobalt solution into a test tube. Add just enough HCl to turn the solution a purple color. Place the test tube into the beaker of hot water from step one until a color change occurs. Record all observations.
10. Prepare an ice bath by placing ice cubes in a 250mL beaker and adding water. Place the test tube of cobalt and HCl solution into the ice bath until a color change occurs. Record all observations.
11. Dispose of chemicals as instructed and clean the lab station.
2. Label a well from left to right across the top "1-6" and down the left side "A-D". Place the well on a white sheet of paper so you can see the effects of the system trying to reach equilibrium.
3. Get the four pipets of CoCl2 solution, HCl, H2O, and Ag.
4. Put five drops of CoCl2 into each of the 24 wells on the white sheet of paper.
5. Add two drops of HCl to the CoCl2 in wells A1, B1, C1, and D1. Add four drops of HCl into the wells A2, B2, C2, D2. Add six drops into each well in column 3, eight drops into each well in column 4, ten drops into each well in column 5, and twelve drops into each well in column 6. Record all observations in a data table.
6. In row B, add one more drop of HCl to each well. Record all observations in a data table.
7. Add five drops of distilled water into each well in row C. Record all observations in a data table.
8. Add five drops of the AgNO3 solution to each well in row D. Record all observations in a data table.
9. Place 5mL of cobalt solution into a test tube. Add just enough HCl to turn the solution a purple color. Place the test tube into the beaker of hot water from step one until a color change occurs. Record all observations.
10. Prepare an ice bath by placing ice cubes in a 250mL beaker and adding water. Place the test tube of cobalt and HCl solution into the ice bath until a color change occurs. Record all observations.
11. Dispose of chemicals as instructed and clean the lab station.
Data Table.
CoCl2 at room temperature: perfectly purple
CoCl2 in hot water: indigo blue turns to dark blue
CoCl2 in cold water: see-through magenta (light red)
CoCl2 in hot water: indigo blue turns to dark blue
CoCl2 in cold water: see-through magenta (light red)
Conclusion.
In conclusion, as more compounds were added to the initial amount of CoCl2 in each well, a color change took place. This was stress placed onto the system. The color changreturn e that took place was the system returning to equilibrium. When heat, AgNO3, and water were added, the system underwent further stress and again returned to equilibrium with changes in color.
Discussion of Theory.
Le Chatelier's principle states that when stressed, a system will return to equilibrium, minimizing the stress. The main stresses that can be placed on an equilibrium system are changes in concentration, changes in temperature, and changes in pressure or volume for gases. The imposed stress on the system will cause the system to shift left, right, or not at all to reach equilibrium. When stress is placed on a system, solids and liquids do not affect the point of equilibrium. When reactants are increased, the system will shift to the right. When reactants are decreased, the system will shift to the left. When changing the temperature of a system, an exothermic reaction will produce heat, therefore heat is a product, shifting the reaction to the left. If a reaction uses heat it is endothermic, shifting the reaction to the right. Changes in temperature are the only thing that is able to affect the constant K value. Gases are also affected by pressure. Increasing the volume of the container decreases the pressure, therefore the reaction shifts to the side with more moles. Decreasing the volume of the container increases the pressure, so the reaction shifts to the side with less moles. When adding an inert gas (gas that is not involved in the reaction) to a system, the shift will depend on the external pressure and volume. If the inert gas is added at a constant external pressure, the system will shift to the side with more moles to reach equilibrium. If the inert gas is added at a constant volume, the system will not shift. In this experiment, no catalysts were used. Catalysts speed up the process of reaching equilibrium but do not change the point of equilibrium for a system.
Sources of Error.
- There was no distilled water to use.
- The 'ice bath' was mostly water with only a little bit of ice.
- When adding HCl to CoCl2, the description of 'purple' could have been different depending on whom was asked.
- When adding drops of HCl to the wells, an extra drop could have easily been added accidentally.
- The 'ice bath' was mostly water with only a little bit of ice.
- When adding HCl to CoCl2, the description of 'purple' could have been different depending on whom was asked.
- When adding drops of HCl to the wells, an extra drop could have easily been added accidentally.
Pre-Lab Questions.
1. When stressed, a reaction will shift to restore equilibrium is Le Chatelier's principle.
2. When equilibrium is reached, the rate of the products is equal to the rate of the reactants.
3. The stresses studied in this experiment were temperature, a change in the concentration of the reactants, and a change in the concentration of the products.
4. Compounds such as CoCl2*6H2O that have water as part of their crystal structure are known as hydrates.
5. With both hydrochloric acid and silver nitrate the safety precautions included wearing goggles, gloves, an apron, and avoiding breathing near the compounds directly.
6. a. If HCl was added to the system, the system would shift to the right because H+ ions would be added.
b. If water was added to the system, nothing would happen because water is a liquid.
c. If NaOH was added, it would neutralize and reduce the amount of H+ ions, therefore shifting the system to the left.
2. When equilibrium is reached, the rate of the products is equal to the rate of the reactants.
3. The stresses studied in this experiment were temperature, a change in the concentration of the reactants, and a change in the concentration of the products.
4. Compounds such as CoCl2*6H2O that have water as part of their crystal structure are known as hydrates.
5. With both hydrochloric acid and silver nitrate the safety precautions included wearing goggles, gloves, an apron, and avoiding breathing near the compounds directly.
6. a. If HCl was added to the system, the system would shift to the right because H+ ions would be added.
b. If water was added to the system, nothing would happen because water is a liquid.
c. If NaOH was added, it would neutralize and reduce the amount of H+ ions, therefore shifting the system to the left.
Post-Lab Questions.
1. a. With the addition of HCl, the system would shift right.
b. With the addition of water, the system would shift left.
c. With the addition of AgNO3, the system would shift left.
d. When increasing the temperature, the system would shift right.
e. When decreasing the temperature, the system would shift left.
2. a. When HCl splits, the Cl- ions will increase.
b. When water is added, the reactants concentration will cause the system to shift to the right.
3. When AgNO3 was added, a precipitant formed in each of the wells in row D. This could be because the AgNO3 reacted with another one of the reactants and formed a solid as a product.
4. The reaction in the Introduction is endothermic because of both the color changes when exposed to heat and cold as well as the energy it will take to break apart H2O from Co.
5. The equilibrium expression for the system studied is stated as the concentration of products over the concentration of reactants with coefficients as exponents.
b. With the addition of water, the system would shift left.
c. With the addition of AgNO3, the system would shift left.
d. When increasing the temperature, the system would shift right.
e. When decreasing the temperature, the system would shift left.
2. a. When HCl splits, the Cl- ions will increase.
b. When water is added, the reactants concentration will cause the system to shift to the right.
3. When AgNO3 was added, a precipitant formed in each of the wells in row D. This could be because the AgNO3 reacted with another one of the reactants and formed a solid as a product.
4. The reaction in the Introduction is endothermic because of both the color changes when exposed to heat and cold as well as the energy it will take to break apart H2O from Co.
5. The equilibrium expression for the system studied is stated as the concentration of products over the concentration of reactants with coefficients as exponents.
Critical Thinking.
1. Adding sodium chloride would increase Cl- concentrations, therefore shifting the reaction to the right.
2. Net ionic equation:
50kj/mole + Co(H2O)6 +2 + 4Cl <- -> CoCl4 -2 + 6H2O
3. At equilibrium, more solid silver chloride would be expected because in the equilibrium reaction, K is equal to products over reactants. Therefore, increasing the amount of products would increase the K value.
2. Net ionic equation:
50kj/mole + Co(H2O)6 +2 + 4Cl <- -> CoCl4 -2 + 6H2O
3. At equilibrium, more solid silver chloride would be expected because in the equilibrium reaction, K is equal to products over reactants. Therefore, increasing the amount of products would increase the K value.