Le Chatelier's Principle Lab: Due February 15th, 2013
Purpose: The purpose of this experiment is to determine the effects of various stresses, such as changes in temperature or concentrations of reactants and products, imposed on a system on its equilibrium. These effects will be predicted and explained using Le Chatelier's principle.
Procedure:
1. 100 mL of water was poured into a beaker and placed on a hot plate. The water was heated on a moderate setting, for the water should not boil.
2. A piece of paper was placed under the well plate, and labeled 1-6 across the columns and A-D down the rows to indicate which solutions were mixed in each well.
3. Micropipets containing CoCl2 (solution), HCl (hydrochloric acid), H2O (tap water), and AgNO3 (silver nitrate) were obtained.
4. 5 drops of CoCl2 were placed in each of the 24 wells.
5. 2 drops of HCl were added to wells A1, B1, C1 and D1, and 4 to all of those in Column 2. 6 drops were added to the wells in column 3, 8 to those in column 4, 10 to column 5, and 12 to column 6. The contents of each well were mixed carefully with a stirring rod.
6. 1 more drop of HCl was added to the wells in row B, and its contents stirred well with a plastic stirring rod. The colors of the solutions in Row B were recorded.
7. 5 drops of tap water were added to each well in row C, and the solution was mixed. The colors of the new solutions were recorded.
8. 5 drops of AgNO3 were added to row D and were stirred with a toothpick. After discarding the toothpick, the colors of the solutions were recorded.
9. 5 mL of the cobalt solution were placed in a test tube, and HCl was added until the color of the solution was halfway between pink and blue.
10. Place in an ice bath consisting of ice cubes and water in an 250 mL beaker.
11. All chemicals were disposed of properly and the lab station cleaned up.
2. A piece of paper was placed under the well plate, and labeled 1-6 across the columns and A-D down the rows to indicate which solutions were mixed in each well.
3. Micropipets containing CoCl2 (solution), HCl (hydrochloric acid), H2O (tap water), and AgNO3 (silver nitrate) were obtained.
4. 5 drops of CoCl2 were placed in each of the 24 wells.
5. 2 drops of HCl were added to wells A1, B1, C1 and D1, and 4 to all of those in Column 2. 6 drops were added to the wells in column 3, 8 to those in column 4, 10 to column 5, and 12 to column 6. The contents of each well were mixed carefully with a stirring rod.
6. 1 more drop of HCl was added to the wells in row B, and its contents stirred well with a plastic stirring rod. The colors of the solutions in Row B were recorded.
7. 5 drops of tap water were added to each well in row C, and the solution was mixed. The colors of the new solutions were recorded.
8. 5 drops of AgNO3 were added to row D and were stirred with a toothpick. After discarding the toothpick, the colors of the solutions were recorded.
9. 5 mL of the cobalt solution were placed in a test tube, and HCl was added until the color of the solution was halfway between pink and blue.
10. Place in an ice bath consisting of ice cubes and water in an 250 mL beaker.
11. All chemicals were disposed of properly and the lab station cleaned up.
Data Table
Conclusion
After analyzing the results of the experiment, it was concluded that Le Chatelier's principle perfectly explains the observations of the shifts in equilibrium of the reaction as the system was subjected to various stresses. When the temperature was increased, the system shifted right to favor the products and when the temperature was decreased, it shifted left to favor the reactants. This was observed through the color changes of the solution. When the reaction favored the products, the solution appeared a blue color; and when it favored the reactants it appeared pink. These shifts in equilibrium are explained in Le Chatelier's principle, which states that a reaction will shift either right or left to restore equilibrium. The various stresses the system was subjected to resulted in shifts of equilibrium, which is exactly what is stated in Le Chatelier's principle.
Discussion of Theory
The idea of equilibrium is said to be the most important and central concept in all of chemistry, and that every chemical reaction can be explained by this principle. A reaction is considered to be in equilibrium when the rate of the forward reaction is equal to that of the reverse reaction. The concentrations of the species are not equal; however, they are held constant as the reactions are occurring at equal rates. Equilibrium is dynamic- when a reaction reaches equilibrium, the reaction does not cease; the forward and reverse reactions are occurring at the same rate. The equilibrium constant expression relates the concentrations or pressures (gases) of the products to the reactants, and is represented with a capital K. When the equilibrium expression is written, only gases and aqueous species are included. This is because the concentrations of solids and liquids are constant, so there is no change to represent in the equilibrium expression.
Equilibrium is affected mainly by changes in pressure, concentration and temperature. The only one of these factors that change the value of the equilibrium constant, however, is temperature. The effect of these changes is described in Le Chatelier's Principle, which states that a chemical reaction will attempt to restore equilibrium after affected by outside stresses by "undoing" what was changed. For example, when the concentration of a reactant is increased, the reaction will shift in favor of the products in an attempt to counteract the increase of reactants on the left side. This principle is often represented metaphorically as two piles of sand, one for the reactants and one for the reactants. When one pile is increased, the sand will overflow and flow into the other pile to make the two piles equal size once again. This is exactly what occurs during a chemical reaction attempting to reach equilibrium. After a stress is applied to a system, the system will try to overcome the change and return the rates of the forward and reverse reactions to equal. When the species of a reaction are gases, the same principle applies to its shifts of equilibrium. Gases are most affected by changes in pressure. When the pressure in the system is decreased, the gas molecules have a greater force on each other, forcing some of the molecules to be transfered to the side with fewer moles, and therefore fewer particles. This side is the side that is favored as a result of the decrease in pressure. A change in temperature is the only stress on a system that will alter the value of the equilibrium constant. Changing the temperature can be thought of as adding or subtracting heat, and the sign on this heat value shows whether the reaction is exothermic or endothermic. If the reaction is endothermic, the heat is written as a reactant, and the opposite for an exothermic reaction. In an endothermic reaction, an increase in temperature would cause a rightward shift on the reaction towards the products, because heat in this case is a reactant. Since heat is produced during the forward reaction, according to Le Chatelier's principle, some of that heat will be used in the reverse reaction to restore the system to equilibrium. The opposite is true for an exothermic reaction- an increase in temperature would cause the system to favor the reactants since the heat is written as a product. A change in temperature changes the value of the change in heat for the entire reaction, not just a single reactant or product. Enthalpy is a part of a chemical reaction that does not change easily, and is a defining characteristic of a reaction; therefore, any factor that changes its value also will change the equilibrium constant value as well. This is why an increase or decrease in temperature to a system will alter the equilibrium constant value.
Equilibrium and its accompanying principles explain the nature of chemical reactions at their core. Most processes that make the functioning of the universe can be related to equilibrium in some fashion. From homeostasis in the body to vital chemical reactions in nature, the attempt by to reach equilibrium is a very important idea to chemistry and other science fields as well.
Equilibrium is affected mainly by changes in pressure, concentration and temperature. The only one of these factors that change the value of the equilibrium constant, however, is temperature. The effect of these changes is described in Le Chatelier's Principle, which states that a chemical reaction will attempt to restore equilibrium after affected by outside stresses by "undoing" what was changed. For example, when the concentration of a reactant is increased, the reaction will shift in favor of the products in an attempt to counteract the increase of reactants on the left side. This principle is often represented metaphorically as two piles of sand, one for the reactants and one for the reactants. When one pile is increased, the sand will overflow and flow into the other pile to make the two piles equal size once again. This is exactly what occurs during a chemical reaction attempting to reach equilibrium. After a stress is applied to a system, the system will try to overcome the change and return the rates of the forward and reverse reactions to equal. When the species of a reaction are gases, the same principle applies to its shifts of equilibrium. Gases are most affected by changes in pressure. When the pressure in the system is decreased, the gas molecules have a greater force on each other, forcing some of the molecules to be transfered to the side with fewer moles, and therefore fewer particles. This side is the side that is favored as a result of the decrease in pressure. A change in temperature is the only stress on a system that will alter the value of the equilibrium constant. Changing the temperature can be thought of as adding or subtracting heat, and the sign on this heat value shows whether the reaction is exothermic or endothermic. If the reaction is endothermic, the heat is written as a reactant, and the opposite for an exothermic reaction. In an endothermic reaction, an increase in temperature would cause a rightward shift on the reaction towards the products, because heat in this case is a reactant. Since heat is produced during the forward reaction, according to Le Chatelier's principle, some of that heat will be used in the reverse reaction to restore the system to equilibrium. The opposite is true for an exothermic reaction- an increase in temperature would cause the system to favor the reactants since the heat is written as a product. A change in temperature changes the value of the change in heat for the entire reaction, not just a single reactant or product. Enthalpy is a part of a chemical reaction that does not change easily, and is a defining characteristic of a reaction; therefore, any factor that changes its value also will change the equilibrium constant value as well. This is why an increase or decrease in temperature to a system will alter the equilibrium constant value.
Equilibrium and its accompanying principles explain the nature of chemical reactions at their core. Most processes that make the functioning of the universe can be related to equilibrium in some fashion. From homeostasis in the body to vital chemical reactions in nature, the attempt by to reach equilibrium is a very important idea to chemistry and other science fields as well.
Sources of Error
When conducting an experiment involving the preparing of many different solutions, there is the potential for sources of error. The experiment originally called for distilled water, but tap water was used instead. Distilled water does not contain any minerals or other impurities that could interfere with the chemical reactions that took place during the lab. These minerals and impurities in tap water might have interfered with the reactions of the experiment, causing the results of the reactions observed to be different than the actual results. In addition to the lack of distilled water, each of the solutions were added in unequal amounts. For example, 5 drops of CoCl2 was added to each of the wells. Each of these five drops were not of equal amount, so the amount of this solution in each well was slightly different. This might have resulted in the reactions being slightly different than intended. Furthermore, the same stirring rod was used to stir each solution after the different chemicals were added. While the stirring rod was rinsed and dried in between uses, there was likely cross-contamination from the other wells and from water remaining on the stirring rod. In addition, the reactions were not all stirred for the same amount of time. One reaction might have needed more stirring time and another less stirring time, perhaps preventing some reactions from progressing properly.
Pre-Lab Questions
1. Le Chatelier's principle states that when a system undergoes stresses, such as changes in temperature, pressure, volume or concentration of the reactants or products, the system will shift either left or right to restore equilibrium. This shift will either cause the reaction to favor the reactants (leftward shift) or the products (rightward shift), restoring equal rates for the forward and reverse reactions.
2. When equilibrium is reached, the rates of the forward and reverse reactions are equal. The concentrations of reactants and products are not equal, but they are constant.
3. The stresses that wil be studied in this experiment are changes in temperature (in the form of heat), concentration of HCl, H2O and AgNO3.
4. The name given to compounds that have water as part of their crystal structure, such as CoCl2*H2O, are hydrates.
5. When hydrochloric acid and silver nitrate are used, certain safety precautions should be taken to prevent injury. Goggles must be worn at all times to prevent chemicals from splashing in the eyes, which would likely cause permanent blindness. In addition, gloves and aprons should be worn by all who handle the chemicals, to protect the hands and the rest of the body from chemical burns that occur when HCl comes in contact with skin and from semi-permanent stains on the skin from AgNO3. Furthermore, when piping the chemicals into their respected solutions, one should not inhale towards the open chemicals- HCl has a very strong odor that can cause severe burning of the inside of the nasal cavity, lasting for several minutes.
6. a) If HCl is added to the system, it would shift right to favor the products to restore equilibrium, due to the increased amount of HCl on the reactant side of the reaction.
b) If water is added to the system, the reaction would shift left to favor the reactants to restore equilibrium, because of the increased amount of water on the product side of the reaction.
c) If NaOH is added to the system, it would shift left, favoring the reactants and restoring equilibrium. This is because the basic NaOH would neutralize the H+ ions on the reactant side, causing a decrease in the number of H+ ions and thus a leftward shift of the system.
2. When equilibrium is reached, the rates of the forward and reverse reactions are equal. The concentrations of reactants and products are not equal, but they are constant.
3. The stresses that wil be studied in this experiment are changes in temperature (in the form of heat), concentration of HCl, H2O and AgNO3.
4. The name given to compounds that have water as part of their crystal structure, such as CoCl2*H2O, are hydrates.
5. When hydrochloric acid and silver nitrate are used, certain safety precautions should be taken to prevent injury. Goggles must be worn at all times to prevent chemicals from splashing in the eyes, which would likely cause permanent blindness. In addition, gloves and aprons should be worn by all who handle the chemicals, to protect the hands and the rest of the body from chemical burns that occur when HCl comes in contact with skin and from semi-permanent stains on the skin from AgNO3. Furthermore, when piping the chemicals into their respected solutions, one should not inhale towards the open chemicals- HCl has a very strong odor that can cause severe burning of the inside of the nasal cavity, lasting for several minutes.
6. a) If HCl is added to the system, it would shift right to favor the products to restore equilibrium, due to the increased amount of HCl on the reactant side of the reaction.
b) If water is added to the system, the reaction would shift left to favor the reactants to restore equilibrium, because of the increased amount of water on the product side of the reaction.
c) If NaOH is added to the system, it would shift left, favoring the reactants and restoring equilibrium. This is because the basic NaOH would neutralize the H+ ions on the reactant side, causing a decrease in the number of H+ ions and thus a leftward shift of the system.
Post-Lab Questions
1. a) Equilibrium will be shifted right with the addition of HCl. The additional reactant will increase the rate of the forward reaction to cause this shift.
b) With the addition of liquid water to the system, equilibrium should not shift at all, for liquids and solids do not change the status of equilibrium. However, it was observed that after the addition of water, the solution turned slightly pink. This means that equilibrium was shifted left. This observation is a result of systematic and other sources of error throughout the experiment.
c) Equilibrium will be shifted left with the addition of AgNO3 because a precipitate forms.
d) Equilibrium will be shifted right when the temperature increases. This is because the reaction is endothermic, and heat is written as a reactant. An increase in temperature will result in some of that heat being transfered to the reverse reaction, causing the rightward shift.
e) Equilibrium will be shifted left when the temperature decreases. The lack of heat on the reactants side results in the reverse reaction's rate to increase and cause the leftward shift.
2. Equilibrium shifts with the addition of HCl and AgNO3 in order to restore the system to equal rates of the forward and reverse reactions. The system will shift to favor the side that the compound was not added to, as an attempt to "balance" the reaction. For example, in (a), HCl was added to the reactants side, so the system will shift right towards the products side to restore equilibrium. In (b), water was added to the products side, so the system shifted left to restore equilibrium.
3. AgNO3 causes a leftward shift of the system due to the nature of the compound in the reaction. Ag reacts with Cl to form the solid AgCl, which does not change the status of equilibrium. But the formation of this solid reduces the number of chlorine ions on the reactant side, causing the leftward shift of the system to restore equilibrium.
4. The reaction shown in the introduction is endothermic. This is because when the reaction was heated, it changed colors from pink to blue, representing the progression of the forward reaction. If the reaction had turned pink from blue, that would indicate that the reverse reaction was taking place and some of the heat released from the forward reaction is being used to complete the reverse reaction.
5. The equilibrium expression for the system studied is as follows: K= [COCl2]
[Cl]^4[Co(H2O)6]
b) With the addition of liquid water to the system, equilibrium should not shift at all, for liquids and solids do not change the status of equilibrium. However, it was observed that after the addition of water, the solution turned slightly pink. This means that equilibrium was shifted left. This observation is a result of systematic and other sources of error throughout the experiment.
c) Equilibrium will be shifted left with the addition of AgNO3 because a precipitate forms.
d) Equilibrium will be shifted right when the temperature increases. This is because the reaction is endothermic, and heat is written as a reactant. An increase in temperature will result in some of that heat being transfered to the reverse reaction, causing the rightward shift.
e) Equilibrium will be shifted left when the temperature decreases. The lack of heat on the reactants side results in the reverse reaction's rate to increase and cause the leftward shift.
2. Equilibrium shifts with the addition of HCl and AgNO3 in order to restore the system to equal rates of the forward and reverse reactions. The system will shift to favor the side that the compound was not added to, as an attempt to "balance" the reaction. For example, in (a), HCl was added to the reactants side, so the system will shift right towards the products side to restore equilibrium. In (b), water was added to the products side, so the system shifted left to restore equilibrium.
3. AgNO3 causes a leftward shift of the system due to the nature of the compound in the reaction. Ag reacts with Cl to form the solid AgCl, which does not change the status of equilibrium. But the formation of this solid reduces the number of chlorine ions on the reactant side, causing the leftward shift of the system to restore equilibrium.
4. The reaction shown in the introduction is endothermic. This is because when the reaction was heated, it changed colors from pink to blue, representing the progression of the forward reaction. If the reaction had turned pink from blue, that would indicate that the reverse reaction was taking place and some of the heat released from the forward reaction is being used to complete the reverse reaction.
5. The equilibrium expression for the system studied is as follows: K= [COCl2]
[Cl]^4[Co(H2O)6]
Critical Thinking
1. The addition of sodium chloride would result in a shift in equilibrium to the right. This is because the number of chloride ions would increase on the forward reaction, resulting in a shift of equilibrium to the right to restore equal rates of the forward and reverse reactions.
2. The new net ionic equation with the energy included is: Heat+Co(H2O)6 + 4Cl DOUBLE ARROW CoCl4 + 6H2O. Since the change in heat is positive, this means the reaction needs heat added for the reaction to occur.
3. At equilibrium, there would be more solid silver chloride than silver and chloride ions. This is because the large K value indicates that there are more products present than reactants at equilibrium. A larger number on top of the K equation makes the K value a large, whole number. So a larger amount of solid silver chloride exists in equilibrium than those ions.
2. The new net ionic equation with the energy included is: Heat+Co(H2O)6 + 4Cl DOUBLE ARROW CoCl4 + 6H2O. Since the change in heat is positive, this means the reaction needs heat added for the reaction to occur.
3. At equilibrium, there would be more solid silver chloride than silver and chloride ions. This is because the large K value indicates that there are more products present than reactants at equilibrium. A larger number on top of the K equation makes the K value a large, whole number. So a larger amount of solid silver chloride exists in equilibrium than those ions.