Le Chatlier's Principle Lab Write-up
(February 7,2013 - Lab Date)
(February 15, 2013 : Due Date)
Purpose
The purpose of this lab is to see how different stresses can have an affect on a system at equilibrium.
Procedures
1. Safety equipment (googles, latex gloves, and safety aprons) are equipped.
2. A beaker filled with 100mL of water is placed onto a hot plate. Beaker was carefully watched to ensure that it does not boil.
3. A well plate is placed on a blank sheet of paper. The well plate was labeled in rows (A-D) and columns (1-6). The paper is where the labels are placed.
4. Vials of CoCl2, HCl, H2O, and AgNO3 was obtained.
5. Five drops of CoCl2 was added to each of the 24 wells in the well plate. The color change is recorded and a photo is taken.
6. Two drops of HCl was added to each well in column 1. Four drops were added into column 2. Six drops were added to column 3. Eight drops were added to column 4. Ten drops were added to column 5. After all the drops were added, each well was stirred with a plastic toothpick until the solution is completely mixed. Observations on the color changes was recorded. The stirring mechanism was rinsed and washed to prevent any contamination.
7. One additional drop of HCl was added to row B and then mixed with the plastic toothpick. Observations was recorded. The plastic toothpick is rinsed.
8. Five drops of distilled water is added into each well in row C. The mixture is stirred with a plastic toothpick. Observations was recorded. Plastic toothpick is rinsed.
9. Five drops of AgNO3 solution was added to each well in row D. The mixture is stirred with a plastic toothpick. Observations was recorded. Plastic toothpick is rinsed.
10. 5mL of CoCl is added to a testtube. Hcl was added slowly until HCl turns a purple color in between pink and blue.
11. The test tube is placed into the beaker that was placed onto the hot plate. The color change is recorded.
12. An ice bath was made with a 250 mL beaker, ice and water. Place test tube in ice bath until color change. The color change is recorded.
13. Chemicals are disposed properly and safety equipment are put away. Lab station is cleaned up.
2. A beaker filled with 100mL of water is placed onto a hot plate. Beaker was carefully watched to ensure that it does not boil.
3. A well plate is placed on a blank sheet of paper. The well plate was labeled in rows (A-D) and columns (1-6). The paper is where the labels are placed.
4. Vials of CoCl2, HCl, H2O, and AgNO3 was obtained.
5. Five drops of CoCl2 was added to each of the 24 wells in the well plate. The color change is recorded and a photo is taken.
6. Two drops of HCl was added to each well in column 1. Four drops were added into column 2. Six drops were added to column 3. Eight drops were added to column 4. Ten drops were added to column 5. After all the drops were added, each well was stirred with a plastic toothpick until the solution is completely mixed. Observations on the color changes was recorded. The stirring mechanism was rinsed and washed to prevent any contamination.
7. One additional drop of HCl was added to row B and then mixed with the plastic toothpick. Observations was recorded. The plastic toothpick is rinsed.
8. Five drops of distilled water is added into each well in row C. The mixture is stirred with a plastic toothpick. Observations was recorded. Plastic toothpick is rinsed.
9. Five drops of AgNO3 solution was added to each well in row D. The mixture is stirred with a plastic toothpick. Observations was recorded. Plastic toothpick is rinsed.
10. 5mL of CoCl is added to a testtube. Hcl was added slowly until HCl turns a purple color in between pink and blue.
11. The test tube is placed into the beaker that was placed onto the hot plate. The color change is recorded.
12. An ice bath was made with a 250 mL beaker, ice and water. Place test tube in ice bath until color change. The color change is recorded.
13. Chemicals are disposed properly and safety equipment are put away. Lab station is cleaned up.
Data Table
CoCl2 at room temperature: purple (between blue and pink)
CoCl2 in cold water: pink
CoCl2 in hot water: dark blue
CoCl2 in cold water: pink
CoCl2 in hot water: dark blue
Conclusion
The addition of HCl to a mixture/solution of HCl and CoCl2 creates a dark blue color inside of the wells, indicating the reaction to shift towards the right. The addition of water to the solutions resulted in a light pink color, indifferent from the original purple produced by the HCl and CoCl2 mixture. This indicates a left shift in the reaction. The addition of AgNO3 gave out a precipitate, causing the concentration of the reactants to decrease, shifting the reaction to the left.
Discussion of Theory
The Le Chatelier's principle is when an equilibrium system is subjected to a stress, the system responds by attaining a new equilibrium condition that minimizes the imposed stress. There are three different stresses that pertains to an equilibrium system. They are changes in concentration, changes in temperature, and changes in pressure or volume for a gaseous condition. Stress causes the equilibrium system to shift left, right, or not at all. It is important to remember that when looking at concentration, if a product or reactant contains a solid or liquid, that reactant or product should be omitted. This is because solid and liquid does not affect a concentration. Adding or removing a pure solid or liquid does not change the concentration, so it does not affect the equilibrium. The increase in the concentration of reactants or the decrease in the concentration of products would lead to the equilibrium shifting towards the right. The decrease in the concentration of reactants or the increase in the concentration of products will cause a shift to the left. When talking about heat, it is important to know that the heat is affecting the whole reaction, not just one side. When trying to determine what side the equilibrium shifts to,increasing temperature or decreasing temperature can be thought of as adding or subtracting "heat". If heat is a product, this shows that the reaction is exothermic. The addition of heats shifts the reaction to the left. If heat is a reactant, it shows that the reaction is endothermic. That addition of heat would shift the reaction to the right. Decreasing heat would work vice versa. Changes in temperature also causes a change in the value of Kc or Kp. The reaction would shift left when a reverse reaction is favored. The reaction would shift right when a forward reaction is favored. Increase in the volume of a container decreases pressure. In this case, the reaction would shift to the side with more moles. Decreasing the volume would do the opposite. Adding an inert gas to an equilibrium system depends on if it is added to constant pressure or volume. If inert gas is added at constant external pressure, then volume increases to accommodate the added gas, thereby decreasing overall pressure. Equilibrium shifts toe the side with more moles. If an inert gas is dded at constant volume, the concentrations and partial pressures of reactants and products do not change. There would be no change in equilibrium. There is a way to speed up the process of reaching equilibrium. It is possible by adding a catalysts. It is important to know that catalysts do not change the equilibrium point.
Source of Error
There were many sources of error in pertaining to this lab. One source of error is that when heating the beaker, the water has boiled in many of occasions. This was a minor error but it may have affected the reaction in some ways. Another source of error is the inconsistency in size of drops. The amount of solution in each drop could have changed how much concentration of each solution in each drop. That small inconsistency could have been what decided if the equilibrium shifts right or left. The biggest source of error may have been the fact that the water that was used in this lab was not distilled water. Instead, tap water was used. The minerals in the tap water may have reacted in the reaction, affecting the severity of the equilibrium shift.
Pre-Lab Discussion
1. The Le Chatelier's principle states that when an equilibrium system is subjected to a stress, the system responds by attaining a new equilibrium condition that minimizes the imposed stress.
2. When equilibrium is reached, the rate of the product and the rate of the reactants produced are the same. This also means that the forward and reverse reaction rates are equal.
3. The stress that will be studied in this experiment is the temperature and the concentration of H2O (water), HCl (hydrochloric acid), and COCl2 (cobalt (II) chloride).
4. A compound that has water as part of their crystal structure is called a hydrate.
5. Safety precautions that must be observed with hydrochloric acid is that gloves must be worn at all times when handling this acid due to how corrosive it is. Goggles should be worn to protect the eyes just in case the fumes makes contact with the eyes or if the acid is splashed. Caution when using this acid. Silver nitrate should be handled with caution. Gloves should be used. Avoid contact with the skin and ingestion.
6. a) With the addition of HCl to the reaction, the equilibrium system would shift to the right caused by the increase in the concentration of H+ in the reactants.
6. b) With the addition of water (H2O) to the reaction, the equilibrium system would not be affected because water is a liquid. Pure solids or liquids does not change the concentration, so it doesn't affect equilibrium.
6. c) With the addition of NaOH, the equilibrium system would shift to the left due to the neutralization of some H+ in the reactants. The decrease in its concentration would cause a shift to the left.
2. When equilibrium is reached, the rate of the product and the rate of the reactants produced are the same. This also means that the forward and reverse reaction rates are equal.
3. The stress that will be studied in this experiment is the temperature and the concentration of H2O (water), HCl (hydrochloric acid), and COCl2 (cobalt (II) chloride).
4. A compound that has water as part of their crystal structure is called a hydrate.
5. Safety precautions that must be observed with hydrochloric acid is that gloves must be worn at all times when handling this acid due to how corrosive it is. Goggles should be worn to protect the eyes just in case the fumes makes contact with the eyes or if the acid is splashed. Caution when using this acid. Silver nitrate should be handled with caution. Gloves should be used. Avoid contact with the skin and ingestion.
6. a) With the addition of HCl to the reaction, the equilibrium system would shift to the right caused by the increase in the concentration of H+ in the reactants.
6. b) With the addition of water (H2O) to the reaction, the equilibrium system would not be affected because water is a liquid. Pure solids or liquids does not change the concentration, so it doesn't affect equilibrium.
6. c) With the addition of NaOH, the equilibrium system would shift to the left due to the neutralization of some H+ in the reactants. The decrease in its concentration would cause a shift to the left.
Post-Lab Questions
1. a) The equilibrium was shifted to the right with the addition of HCl.
b) The equilibrium was shifted to the left with the addition of water.
c) The equilibrium was shifted to the left with the addition of AgNO3.
d) The equilibrium was shifted to the right with the increase in temperature.
e) The equilibrium was shifted to the left with the decrease in temperature.
2. When HCl dissociates, it will dissociates into H+ and Cl- ions. This will increase the amount of Cl- ions, causing an increase in concentration. This stress will cause the equilibrium to shift to the right due to the increase in the concentration of the product. The addition of water will cause the H2O in Co(H2O)6^2 to become liquid water. Liquid and solids do not count towards concentration. The equilibrium would shift to the left to remove the stress.
3. When AgNO3 was added to the solutions, it would successfully react with the Cl-. This would form AgCl, which appeared to be in the form of a solid. Because it was a solid, the concentration of Cl- is decreased, causing the equilibrium system to shift to the left.
4. The reaction shown in the Introductions was endothermic. This was determined after increasing the solution of CoCl2 and HCl at equilibrium. The color changed from purple to dark blue, indicating a shift to the right. When that solution is added into the ice bath, it turned to pink, indicating a shift to the left. A shift to the left causes a decrease in temperature and a shift to the right causes a increase in temperature. These shifts indicates that the reaction is indeed endothermic.
5. Ka = ([CoCl4][H2O]^6) / ([Co(H2O)6][Cl-]^4
b) The equilibrium was shifted to the left with the addition of water.
c) The equilibrium was shifted to the left with the addition of AgNO3.
d) The equilibrium was shifted to the right with the increase in temperature.
e) The equilibrium was shifted to the left with the decrease in temperature.
2. When HCl dissociates, it will dissociates into H+ and Cl- ions. This will increase the amount of Cl- ions, causing an increase in concentration. This stress will cause the equilibrium to shift to the right due to the increase in the concentration of the product. The addition of water will cause the H2O in Co(H2O)6^2 to become liquid water. Liquid and solids do not count towards concentration. The equilibrium would shift to the left to remove the stress.
3. When AgNO3 was added to the solutions, it would successfully react with the Cl-. This would form AgCl, which appeared to be in the form of a solid. Because it was a solid, the concentration of Cl- is decreased, causing the equilibrium system to shift to the left.
4. The reaction shown in the Introductions was endothermic. This was determined after increasing the solution of CoCl2 and HCl at equilibrium. The color changed from purple to dark blue, indicating a shift to the right. When that solution is added into the ice bath, it turned to pink, indicating a shift to the left. A shift to the left causes a decrease in temperature and a shift to the right causes a increase in temperature. These shifts indicates that the reaction is indeed endothermic.
5. Ka = ([CoCl4][H2O]^6) / ([Co(H2O)6][Cl-]^4
Critical Thinking
1. With the addition of NaCl, there would be an increase in Cl- ions due to NaCl separating into ions, adding to the amount of Cl- ions. This would increase the concentration in Cl-, causing the forward rate of reaction to increase in order to obtain equilibrium.
2. Co(H2O)6^2+ + 4Cl- + 50 kj/mol ←→ CoCl4^2- + 6H2O
3. Silver Chloride should be expected to be more present . Given that Keq = 6 x 10^9 for the reaction, it tells the inspectors that the concentration of the products is very high and the concentration of the reactants is low. It is shown by how large the value of Keq is. Because AgCl(s) is the product, it is expected that silver chloride will have more in amount.
2. Co(H2O)6^2+ + 4Cl- + 50 kj/mol ←→ CoCl4^2- + 6H2O
3. Silver Chloride should be expected to be more present . Given that Keq = 6 x 10^9 for the reaction, it tells the inspectors that the concentration of the products is very high and the concentration of the reactants is low. It is shown by how large the value of Keq is. Because AgCl(s) is the product, it is expected that silver chloride will have more in amount.