Le Chatelier's Principle Date: 02/07/2013
Purpose
The purpose of this lab is to observe the effects of a stress placed on an equilibrium system. Le Chatelier's Principle predicts how the system will shift in response to the stress placed on the equilibrium system.
Procedure
1. Put on goggles and a lab apron for protection from harmful chemical. Place 100mL of water on a hot plate at a moderate setting. The water should not boil.
2. Place the well plate on a white sheet of paper. This will help distinguish the solution colors at the end of the experiment. After placing the well plate on the paper, label the paper as shown n the diagram. The rows should be labeled from A-D and the columns should be labeled 1-6.
3. Put on latex gloves for protection from hydrochloric acid and silver nitrate. Using four micro-pipets, one for each solution, fill the 24 wells with the concentrations shown in the diagram. Place 5 drops of Co(Cl)2 in all 24 wells.
4. Add two drops of HCl to the first column, labeled "1." Place four drops of HCl in each well in the second column, six drops in column 3, eight drops in column 4, ten drops in column 5, and 12 drops in column 6. Stir solution in each of the wells. Handle HCl with caution, it is very corrosive. Record the color of the solution in each well of row A in the Data Table. Row A will be the control for the experiment.
5. Add one more drop of HCl to each of the wells in row B. Stir the solution in each well with a new toothpick. Do not contaminate the solution with the toothpick used to stir the previous solutions. Record the colors of row B in the Data Table.
6. Add 5 drops of distilled water in each well in row C. Stir with a toothpick and record the colors of row C in the Data Table.
7. Add 5 drops of silver nitrate to each well in row D. Make sure to stir the solution thoroughly. Record any color changes or precipitate formation to the Data Table.
8. Place 5mL of the cobalt solution in a test tube. Add HCl until the solution is purple. Place a white sheet of paper behind the test tube for more accuracy. Then, place the test tube in the beaker with hot water until a color change is visible.
9. Next, prepare an ice bath using 250mL of water and ice. Place the test tube from the hot water to the ice bath until a color change occurs.
10. Properly dispose of all chemicals used in the experiment.
2. Place the well plate on a white sheet of paper. This will help distinguish the solution colors at the end of the experiment. After placing the well plate on the paper, label the paper as shown n the diagram. The rows should be labeled from A-D and the columns should be labeled 1-6.
3. Put on latex gloves for protection from hydrochloric acid and silver nitrate. Using four micro-pipets, one for each solution, fill the 24 wells with the concentrations shown in the diagram. Place 5 drops of Co(Cl)2 in all 24 wells.
4. Add two drops of HCl to the first column, labeled "1." Place four drops of HCl in each well in the second column, six drops in column 3, eight drops in column 4, ten drops in column 5, and 12 drops in column 6. Stir solution in each of the wells. Handle HCl with caution, it is very corrosive. Record the color of the solution in each well of row A in the Data Table. Row A will be the control for the experiment.
5. Add one more drop of HCl to each of the wells in row B. Stir the solution in each well with a new toothpick. Do not contaminate the solution with the toothpick used to stir the previous solutions. Record the colors of row B in the Data Table.
6. Add 5 drops of distilled water in each well in row C. Stir with a toothpick and record the colors of row C in the Data Table.
7. Add 5 drops of silver nitrate to each well in row D. Make sure to stir the solution thoroughly. Record any color changes or precipitate formation to the Data Table.
8. Place 5mL of the cobalt solution in a test tube. Add HCl until the solution is purple. Place a white sheet of paper behind the test tube for more accuracy. Then, place the test tube in the beaker with hot water until a color change is visible.
9. Next, prepare an ice bath using 250mL of water and ice. Place the test tube from the hot water to the ice bath until a color change occurs.
10. Properly dispose of all chemicals used in the experiment.
Data Tables
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Cobalt Solution Observations
Co(Cl2) at room temperature
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Purple; blue and pink mix
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Co(Cl2) in hot water
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Changed from purple to dark blue
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Co(Cl2) in cold water
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Changed from dark blue to purple-pink
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Conclusion
The equilibrium system shifted in response to the stresses. The blue tint of the solutions with a higher concentration of HCl indicates the addition of HCl causes the reaction to shift towards the products. The light pink color shown in row C indicates the addition of water causes dilution of the solution. The precipitate formed and the solution change to light pink in row D indicates the addition of silver nitrate causes the reaction to shift towards the reactants. Temperature change always puts stress on the equilibrium system. When the cobalt solution was placed in the hot water, the solution turned dark blue from being purple. This indicates the reaction shifted towards the products. When the solution was placed in an ice bath, it returned to its original color; purple.
Discussion of Theory
Le Chatelier's Principle was investigated in the experiment. Le Chatelier's Principle states if a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the change and a new equilibrium is established. Liquids and solids are not expressed in the equilibrium expression. Therefore, liquids and solids do not affect the equilibrium of the system. This was demonstrated when the water was added to row C in the experiment. Water is a liquid and should not have affected the system. The impurities in the tap water and uncleaned wells can explain why the system shifted. The two variables that were tested in this lab include temperature and concentration of reactants and products. When the concentration of reactants increase the equilibrium system shifts towards the products. When the concentration of products increase the equilibrium system shifts towards the reactants. This idea was observed when HCl was added to the cobalt chloride solution. The addition of HCl caused the system to shift to the right. HCl increased the concentration on the reactant side, so in order to counteract this imbalance, the reaction shifted to the right. It was also observed when the silver nitrate solution was added. The addition of silver nitrate produced AgCl, a solid. This was evident because the solution in row D was cloudy and murky. This is proof that a precipitate was formed. Temperature is the second stress that was tested. When temperature changes the Kc or Kp value changes as well. Temperature is the only stress that affects Kc and Kp. The Kc value can be determined by raising the concentrations of products to their coefficients over the concentrations of reactants raised to their coefficients. When the cobalt chloride solution was placed in the hot water, the solution turned dark blue. The dark blue indicated that there were more products in the solution (cobalt chloride). When the solution was placed in the ice bath it return to its original color, purple. The purple indicated the reaction shifted towards the reverse reaction. This indicated the reaction was endothermic because the heat is on the reactants side, so the added heat causes the system to shift to the right to counteract this.
Sources of Error
One of the major potential sources of error in this lab was contamination. The wells in the well plate are too small to be able to clean all the sides thorough. There may have been residue from the previous experiments performed with the well plate that was used. The residue could have reacted with the solution and caused the reaction to shift when it was predicted to remain constant. For example, in well 6c. The addition of water should not have shifted the reaction, however, the system shifted to the left. Another major source of error in this lab was the addition of tap water, instead of distilled water. Tap water contains many impurities that could have reacted with the solution. This could have also contributed to the shift that occurred when water was added.
Pre-lab questions
1. Le Chatelier's Principle states if a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the change and a new equilibrium is established.
2. Equilibrium is reached when the rate of the forward reaction equals the reverse reaction.
3. The stresses that will be studied in the experiment are temperature and concentrations of products or reactants. The hot water and ice bath will account for the effects of temperature change and the addition of silver nitrate and HCl account for effects of varying concentration of reactants or products.
4. A hydrate is a compound that has water as a part of its crystal structure.
5. Goggles and gloves must be worn when handling hydrochloric acid. Also, strong concentrations of HCl should not be ingested or inhaled. Gloves and goggles must be worn when handling silver nitrate. Silver nitrate will stain the skin and must be washed off immediately if it comes in contact.
6. Predict the effect of HCl, water, and NaOH on the equilibrium system.
a.) If HCl is added the equilibrium system will shift to the right, towards the products because the addition of HCl increase the concentration of
H+ in the reactants.
b.) The addition of water will not affect the equilibrium system because H2O is a liquid and is not a part of the equilibrium reaction. Water dilutes both sides of the equation and therefore, does not affect the system.
c.) The addition of NaOH will cause the system to shift to the left, towards the reactants because the OH- ions from NaOH will neutralize the H+ ions. The decrease in H+ ions results in a lower concentration, which cause the system to shift to the left.
2. Equilibrium is reached when the rate of the forward reaction equals the reverse reaction.
3. The stresses that will be studied in the experiment are temperature and concentrations of products or reactants. The hot water and ice bath will account for the effects of temperature change and the addition of silver nitrate and HCl account for effects of varying concentration of reactants or products.
4. A hydrate is a compound that has water as a part of its crystal structure.
5. Goggles and gloves must be worn when handling hydrochloric acid. Also, strong concentrations of HCl should not be ingested or inhaled. Gloves and goggles must be worn when handling silver nitrate. Silver nitrate will stain the skin and must be washed off immediately if it comes in contact.
6. Predict the effect of HCl, water, and NaOH on the equilibrium system.
a.) If HCl is added the equilibrium system will shift to the right, towards the products because the addition of HCl increase the concentration of
H+ in the reactants.
b.) The addition of water will not affect the equilibrium system because H2O is a liquid and is not a part of the equilibrium reaction. Water dilutes both sides of the equation and therefore, does not affect the system.
c.) The addition of NaOH will cause the system to shift to the left, towards the reactants because the OH- ions from NaOH will neutralize the H+ ions. The decrease in H+ ions results in a lower concentration, which cause the system to shift to the left.
Post-lab questions
1. In what direction was the equilibrium shifted by
a. The addition of HCl caused the equilibrium to shift to the right.
b. The addition of water caused the equilibrium to shift to the left
c. The addition of AgNO3 caused the equilibrium to shift to the left.
d. Increasing the temperature caused the equilibrium to shift to the right.
e. Decreasing the temperature caused the equilibrium to shift to the left.
2. Explain the results from 1a and 1b.
a) The addition of HCL causes the equilibrium to shift right because HCl splits into H+ and Cl- ions. This causes an increase in Cl- ions and because Cl- is on the reactant side, it will cause it to shift right towards the products.
b) The addition of water would not shift the equilibrium because water is a liquid. Liquids and solids do not affect the equilibrium because they do not have measurable concentrations, however, adding water decreases the overall concentrations of the solutions (dilution). Tap water was used instead of distilled water. The impurities in the tap water could have caused the reaction to shift to the left.
3. When silver nitrate was added, a precipitate formed. The addition of silver nitrate caused a solid to form (AgCl), which decreased the amount of chloride ions in the solution and shifted the system to the right.
4. The reaction is endothermic because adding the cobalt chloride solution to hot water shifted the reaction to the right. If the reaction is endothermic the heat would be on the reactant side and if the reaction is exothermic the heat would be on the product side. Increasing the heat on the reactant side would cause the reaction to shift right, therefore, it can be concluded that the reaction shifted right because the solution turned dark blue.
5. The equilibrium expression for the system studied is
Ka = [Co(Cl)4]
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[ Co(H2O)6] [Cl-]^4
a. The addition of HCl caused the equilibrium to shift to the right.
b. The addition of water caused the equilibrium to shift to the left
c. The addition of AgNO3 caused the equilibrium to shift to the left.
d. Increasing the temperature caused the equilibrium to shift to the right.
e. Decreasing the temperature caused the equilibrium to shift to the left.
2. Explain the results from 1a and 1b.
a) The addition of HCL causes the equilibrium to shift right because HCl splits into H+ and Cl- ions. This causes an increase in Cl- ions and because Cl- is on the reactant side, it will cause it to shift right towards the products.
b) The addition of water would not shift the equilibrium because water is a liquid. Liquids and solids do not affect the equilibrium because they do not have measurable concentrations, however, adding water decreases the overall concentrations of the solutions (dilution). Tap water was used instead of distilled water. The impurities in the tap water could have caused the reaction to shift to the left.
3. When silver nitrate was added, a precipitate formed. The addition of silver nitrate caused a solid to form (AgCl), which decreased the amount of chloride ions in the solution and shifted the system to the right.
4. The reaction is endothermic because adding the cobalt chloride solution to hot water shifted the reaction to the right. If the reaction is endothermic the heat would be on the reactant side and if the reaction is exothermic the heat would be on the product side. Increasing the heat on the reactant side would cause the reaction to shift right, therefore, it can be concluded that the reaction shifted right because the solution turned dark blue.
5. The equilibrium expression for the system studied is
Ka = [Co(Cl)4]
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[ Co(H2O)6] [Cl-]^4
Critical Thinking
1. The addition of sodium chloride would shift equilibrium to the right. Le Chatelier's Principle states that when stress is put on an equilibrium system, the system will response by reaching a new equilibrium to reduce the stress. NaCl dissociates into Na+ and Cl- ions. Since chloride ions are on the reactant side, the addition of NaCl will increase the concentration of the reactants, which causes the system to shift to the right.
2. Co(H20)6^2+ + 4Cl- + 50 kJ/mol <---------> CoCl4^2- + 6H2O
3. There will be more silver chloride than silver and chloride ons because K value is greater than 1. This indicates that the concentration of the products (AgCl) is greater than the concentration of the reactants.
2. Co(H20)6^2+ + 4Cl- + 50 kJ/mol <---------> CoCl4^2- + 6H2O
3. There will be more silver chloride than silver and chloride ons because K value is greater than 1. This indicates that the concentration of the products (AgCl) is greater than the concentration of the reactants.