Le Chatelier Lab- 2/14/13
Purpose
To observe how the system responds to various stresses, such as changes in the concentration of reactants or products, or in the temperature of the system.
Procedure
1. Wear goggles and lab apron for safety.
2. Place a beaker containing about 100 mL of water on a hot plate. While waiting for the water to boil, move on to the next step.
3. Put on latex gloves and place five drops of CoCl2 solution in each of the 24 wells of the labeled well plate.
4. Add two, four, six, eight, ten, and twelve drops of HCl to the CoCl2 solution in column 1, 2, 3, 4, 5, 6 of the well plate, respectively.
5. Mix the contents of each well with the plastic toothpick. Record observations.
6. In row B, add one more drop of HCl to each well and stir with the toothpick. Record observations.
7. In row C, add five drops of distilled water to each well and stir with the toothpick. Record observations.
8. In row D, add five drops of AgNO3 solution to each well and stir with the toothpick. Record observations.
9. Place 5 mL of cobalt solution in a test tube.
10. Add just enough HCl to get a purple color.
11. Place the test tube in the beaker of hot water from step 2 until a color change occurs. Record observations.
12. Place test tube in a 250-mL beaker containing ice cubes and water until a color change occurs. Record observations.
13. Dispose all chemicals and clean up the work area. Wash hands before leaving.
2. Place a beaker containing about 100 mL of water on a hot plate. While waiting for the water to boil, move on to the next step.
3. Put on latex gloves and place five drops of CoCl2 solution in each of the 24 wells of the labeled well plate.
4. Add two, four, six, eight, ten, and twelve drops of HCl to the CoCl2 solution in column 1, 2, 3, 4, 5, 6 of the well plate, respectively.
5. Mix the contents of each well with the plastic toothpick. Record observations.
6. In row B, add one more drop of HCl to each well and stir with the toothpick. Record observations.
7. In row C, add five drops of distilled water to each well and stir with the toothpick. Record observations.
8. In row D, add five drops of AgNO3 solution to each well and stir with the toothpick. Record observations.
9. Place 5 mL of cobalt solution in a test tube.
10. Add just enough HCl to get a purple color.
11. Place the test tube in the beaker of hot water from step 2 until a color change occurs. Record observations.
12. Place test tube in a 250-mL beaker containing ice cubes and water until a color change occurs. Record observations.
13. Dispose all chemicals and clean up the work area. Wash hands before leaving.
Data Table
1 2 3 4 5 6
A Dark Pink Dark Pink Light Pink Teal Blue Blue
B Dark Pink Dark Pink Indigo Teal Blue Blue
C Dark Pink Dark Pink Dark Pink Dark Pink Dark Pink Dark Pink
D Pink Pink Pink Pink Pink Pink
CoCl2 at room temperature Purple
CoCl2 in hot water Blue
CoCl2 in cold water Red
A Dark Pink Dark Pink Light Pink Teal Blue Blue
B Dark Pink Dark Pink Indigo Teal Blue Blue
C Dark Pink Dark Pink Dark Pink Dark Pink Dark Pink Dark Pink
D Pink Pink Pink Pink Pink Pink
CoCl2 at room temperature Purple
CoCl2 in hot water Blue
CoCl2 in cold water Red
Conclusion
The lab demonstrated Le Chatelier's Principle and the nature of equilibrium shifts.
To do so, 5 drops of CoCl2 solution was dropped in 24 slots of the well plate. 2, 4, 6, 8, 10, and 12 drops of HCl was dropped in column 1, 2, 3, 4, 5, and 6, respectively. In row B, one more drop of HCl was added. In row C, 5 drops of H2O were added. In row D, 5 drops of Ag were added. A plastic toothpick was used to stir the solution. The color of the solutions were observed.
5 mL of CoCl2 solution was added with HCl in a test tube until the color of the solution turned purple. When the test tube was heated in hot water, the reaction turned blue. When the test tube was dipped in ice bath, the reaction turned red.
To do so, 5 drops of CoCl2 solution was dropped in 24 slots of the well plate. 2, 4, 6, 8, 10, and 12 drops of HCl was dropped in column 1, 2, 3, 4, 5, and 6, respectively. In row B, one more drop of HCl was added. In row C, 5 drops of H2O were added. In row D, 5 drops of Ag were added. A plastic toothpick was used to stir the solution. The color of the solutions were observed.
5 mL of CoCl2 solution was added with HCl in a test tube until the color of the solution turned purple. When the test tube was heated in hot water, the reaction turned blue. When the test tube was dipped in ice bath, the reaction turned red.
Discussion of Theory
The lab discusses the equilibrium using Le Chetelier's Principle. Equilibrium equation shows when the rate of reactants equal to the rate of products. When temperature or concentration are changed, the stress is placed on either of the reactant or products side. To relieve the stress, the equilibrium shifts left or right. The rate on the either side occur equally at the same time, making equilibrium dynamic. Change in the concentration will affect the system, only if the added or subtracted amounts of solution is gaseous or aqueous; liquids and solids do not contribute to the equilibrium shifts. Addition of concentration in the reactants side and the removal of concentration in the products side will lead to the equilibrium shift to the right. Removal of concentrations in the reactants side and the addition of concentration in the products side will lead to the equilibrium shifts to the left. Temperature also affect the equilibrium shift. For endothermic reactions, increased temperatures would cause the system to shift to the right. Decreased temperatures would cause the system to shift to the left. For exothermic reactions, increased temperatures would cause the system to shift to the left. Decreased temperatures would cause the system to shift to the right. The equilibrium constant, which equals to moles of product over moles of reactants, is represented with unit-less K.
Equilibrium can also be affected by pressure as well, with its constant represented by Kp. Instead of K = moles of products/moles of reactants, Kp is equal to Pproducts/Preactants. Pressures react differently than temperature; increased pressure leads the reaction to shift to the side with fewer moles; decreased pressure leads the reaction to shift to the side with more moles. This can be explained by particles' tendency to spread out; molecules shift to less concentrated area in a crowded space; molecules shift to more concentrated area in an open space.
Equilibrium can also be affected by pressure as well, with its constant represented by Kp. Instead of K = moles of products/moles of reactants, Kp is equal to Pproducts/Preactants. Pressures react differently than temperature; increased pressure leads the reaction to shift to the side with fewer moles; decreased pressure leads the reaction to shift to the side with more moles. This can be explained by particles' tendency to spread out; molecules shift to less concentrated area in a crowded space; molecules shift to more concentrated area in an open space.
Sources of Error
A few sources of error should be noted.
1. The water used for the reaction was tapped water, and may have contained impurities that may have affected the outcome.
It takes too much ions to get the
2. Also, when the drops of CoCl2, HCl, H2O, and Ag were added, the amount for each drops were different, resulting in a different outcome.
3. Some reactions were given more time than others. A delay was caused by limited bottles of HCl and may have caused some nuanced changes in the outcome.
1. The water used for the reaction was tapped water, and may have contained impurities that may have affected the outcome.
It takes too much ions to get the
2. Also, when the drops of CoCl2, HCl, H2O, and Ag were added, the amount for each drops were different, resulting in a different outcome.
3. Some reactions were given more time than others. A delay was caused by limited bottles of HCl and may have caused some nuanced changes in the outcome.
Pre-Lab Questions
1. State Le Chetalier's Principle
When an equilibrium system is subjected to a stress, the system responds by attaining a new equilibrium condition that minimizes the imposed stress.
2. Explain what happens when equilibrium is reached.
When the equilibrium is reached, the rate of the forward and reverse reactions are equal.
3. List the stresses that will be studied in this experiment.
Changes in concentration
Changes in temperature
4. The formula for solid cobalt(II) chloride is CoCl2*6H2O. What is the name given to compounds such as this, which have water as part of their crystal structure?
Cobalt (II) chloride hexahydrate
5. What safety precautions must be observed with hydrochloric acid (HCl)? With silver nitrate (AgNO3)?
Goggles and lab apron should be worn at all times during the lab. Hydrochloric acid is extremely corrosive and will irritate skin. Silver nitrate is toxic and will stain skin and clothing.
6. Predict the effect on the following equilibrium system if you: (a) add HCl; (b) add H2O; (c) add NaOH.
2CrO4^2- + 2H+ = Cr2O2^2- + H2O
a. When HCl is added to the equilibrium system, the reaction would shift to the right. Increased H+ from added HCl causes the reactant side to increase, causing the equilibrium to shift to the right.
b. When H2O is added to the equilibrium system, the reaction would shift to the left, because increased H2O would cause the product side to increase. However, the equilibrium would not shift if H2O is liquid.
c. When NaOH is added to the equilibrium system, the reaction would shift to the left. Addition of OH- would neutralize H+ and lessen the amount of H+, causing the reaction to move leftward.
When an equilibrium system is subjected to a stress, the system responds by attaining a new equilibrium condition that minimizes the imposed stress.
2. Explain what happens when equilibrium is reached.
When the equilibrium is reached, the rate of the forward and reverse reactions are equal.
3. List the stresses that will be studied in this experiment.
Changes in concentration
Changes in temperature
4. The formula for solid cobalt(II) chloride is CoCl2*6H2O. What is the name given to compounds such as this, which have water as part of their crystal structure?
Cobalt (II) chloride hexahydrate
5. What safety precautions must be observed with hydrochloric acid (HCl)? With silver nitrate (AgNO3)?
Goggles and lab apron should be worn at all times during the lab. Hydrochloric acid is extremely corrosive and will irritate skin. Silver nitrate is toxic and will stain skin and clothing.
6. Predict the effect on the following equilibrium system if you: (a) add HCl; (b) add H2O; (c) add NaOH.
2CrO4^2- + 2H+ = Cr2O2^2- + H2O
a. When HCl is added to the equilibrium system, the reaction would shift to the right. Increased H+ from added HCl causes the reactant side to increase, causing the equilibrium to shift to the right.
b. When H2O is added to the equilibrium system, the reaction would shift to the left, because increased H2O would cause the product side to increase. However, the equilibrium would not shift if H2O is liquid.
c. When NaOH is added to the equilibrium system, the reaction would shift to the left. Addition of OH- would neutralize H+ and lessen the amount of H+, causing the reaction to move leftward.
Post-Lab Questions
1. Refer to the net ionic equation below or in the Introduction to answer the following questions.
Co(H2O)6^2+ + 4Cl- = CoCl4^2- + 6H2O
In what direction was the equilibrium shifted by
a. the addition of HCl?
The equilibrium was shifted to the right.
b. the addition of water?
The equilibrium was shifted to the left.
c. the addition of AgNO3?
The equilibrium was shifted to the left.
d. increasing the temperature?
The equilibrium was shifted to the right.
e. decreasing the temperature?
The equilibrium was shifted to the left.
2. How do you explain the results described in answers 1a and 1b?
When HCl or water is added to the system, the ions cause stress on the system. To relieve the stress, the reaction move where ever is appropriate and balance the system. For part a, the equilibrium was shifted right because of the added Cl- ion. For part b, the equilibrium was shifted to the left because of the increase of water.
3. Explain the results observed when AgNO3 was added.
When AgNO3 is mixed with chlorine ions, AgCl is formed. However, AgCl is a solid, and does not contribute to the change in the equilibrium. Forming of AgCl will take away some amounts of Cl- on the reactants side, so the equilibrium would shift to the left.
4. Is the reaction shown in the Introduction exothermic or endothermic? How do you know?
The reaction is endothermic. When the temperature was increased, the system shifted to the right, which implies that temperature is on the reactants side. Because temperature remains on the reactants side, it can be found that the reaction is endothermic.
5. Write the equilibrium expression for the system studied.
K = [CoCl4^2-]/[Co(H2O)6^2+][Cl-]^4
Co(H2O)6^2+ + 4Cl- = CoCl4^2- + 6H2O
In what direction was the equilibrium shifted by
a. the addition of HCl?
The equilibrium was shifted to the right.
b. the addition of water?
The equilibrium was shifted to the left.
c. the addition of AgNO3?
The equilibrium was shifted to the left.
d. increasing the temperature?
The equilibrium was shifted to the right.
e. decreasing the temperature?
The equilibrium was shifted to the left.
2. How do you explain the results described in answers 1a and 1b?
When HCl or water is added to the system, the ions cause stress on the system. To relieve the stress, the reaction move where ever is appropriate and balance the system. For part a, the equilibrium was shifted right because of the added Cl- ion. For part b, the equilibrium was shifted to the left because of the increase of water.
3. Explain the results observed when AgNO3 was added.
When AgNO3 is mixed with chlorine ions, AgCl is formed. However, AgCl is a solid, and does not contribute to the change in the equilibrium. Forming of AgCl will take away some amounts of Cl- on the reactants side, so the equilibrium would shift to the left.
4. Is the reaction shown in the Introduction exothermic or endothermic? How do you know?
The reaction is endothermic. When the temperature was increased, the system shifted to the right, which implies that temperature is on the reactants side. Because temperature remains on the reactants side, it can be found that the reaction is endothermic.
5. Write the equilibrium expression for the system studied.
K = [CoCl4^2-]/[Co(H2O)6^2+][Cl-]^4
Critical Thinking
1. Predict how the addition of sodium chloride would affect the equilibrium. Explain your prediction in terms of Le Chetalier's principle.
Adding NaCl to the equation would shift the system to the right. Cl- ions would increase in the reactants side, and the equilibrium would be shifted to the right to relieve the stress.
2. Rewrite the net ionic equation including the energy term where appropriate. The Delta H for this reaction is +50 kJ/mol.
50 kJ/mol + Co(H2O)6^2+ + 4Cl- <-> CoCl4^2- + 6H2O
3. Silver chloride (AgCl) is a white solid. For the equilibrium reaction
Ag+(aq) + Cl-(aq) = AgCl(s)
K(eq) = 6 * 10^9. At equilibrium, would you expect to have more silver and chloride ions or more solid silver chloride? Explain.
Silver chloride would be produced more, because of the large K value. K is products/reactants, and higher K value implies higher value in products.
Adding NaCl to the equation would shift the system to the right. Cl- ions would increase in the reactants side, and the equilibrium would be shifted to the right to relieve the stress.
2. Rewrite the net ionic equation including the energy term where appropriate. The Delta H for this reaction is +50 kJ/mol.
50 kJ/mol + Co(H2O)6^2+ + 4Cl- <-> CoCl4^2- + 6H2O
3. Silver chloride (AgCl) is a white solid. For the equilibrium reaction
Ag+(aq) + Cl-(aq) = AgCl(s)
K(eq) = 6 * 10^9. At equilibrium, would you expect to have more silver and chloride ions or more solid silver chloride? Explain.
Silver chloride would be produced more, because of the large K value. K is products/reactants, and higher K value implies higher value in products.