Purpose of the Experiment
To examine equilibrium systems and the effects of Le Châtelier’s principle on those systems.
Background Required
This experiment will use basic laboratory techniques. The concepts of equilibrium, Le Châtelier’s principle, and endothermic/exothermic reactions are used in this experiment.
Background Information
Not all reactions go to completion, or use up all of a reactant. In some chemical reactions, there is always some products and reactants present. In these chemical systems, both the forward and reverse reactions are occurring simultaneously. When the rate of the forward reaction (reactants forming products) equals the rate of the reverse reaction (products forming reactants), the system is at equilibrium. The concentrations of the products and reactants remain constant. The arrow used in writing an equilibrium reaction is a double headed arrow (⇋). This indicates that both the forward and reverse reaction are occurring at the same time.
Le Châtelier’s principle states that an equilibrium system will react to a change in conditions (stress) in such a way to offset the change and reestablish the equilibrium. The system will have one reaction dominating until the offsetting changes allow the rates of the forward and reverse reaction to be equal again (reestablishing equilibrium). If the forward reaction dominates in order to offset the changes, we say the system “shifts to the right” or “shifts toward products” in order to reestablish equilibrium conditions. This will increase the concentrations of the products and decrease the concentrations of the reactants. However, if the reverse reaction dominates in order to offset the changes, we say the system “shifts to the left” or “shifts toward reactants” to reestablish equilibrium conditions. This will increase the concentrations of the reactants and decrease the concentrations of the products. The changes will not return the system to the original conditions, but to a new set of conditions that establish equilibrium. The condition changes to be examined are changes that increase or decrease the concentrations of reactants or products and changes in the temperature.
In endothermic reactions, heat can be considered a reactant, since heat is absorbed from the surroundings. In exothermic reactions, heat can be considered a product, since heat is released from the system. When heat is added to an endothermic reaction, the system will shift toward products. When heat is added to an exothermic reaction, the system will shift toward reactants.
Endothermic: heat + D + E → F + G
Exothermic: A + B → C + D + heat
In this experiment, two equilibrium systems will be examined. In the first system, the reaction is shown below. The Fe(SCN)+2 complex ion is a blood-red solution. We can monitor how the equilibrium system shifts to offset the change by observing the changes in intensity of the red color.
Fe+3(aq) + SCN–(aq) ⇋ Fe(SCN)+2(aq)
In the second equilibrium system, the reaction involves two colored complex ions. The [CoCl4]–2 ion is deep blue while the [Co(H2O)6]+2 ion is rose-pink. This reaction is shown below. By observing the color changes, we can tell if there is more reactants (blue) or more products (pink).
[CoCl4]–2(alcohol) + 6 H2O(ℓ) ⇋ [Co(H2O)6]+2(aq) + 4 Cl–(aq)
Example
Problem
In the acetic acid equilibrium reaction given below, how will the system react if NaC2H3O2(aq) is added to the system?
Will the concentration of H3O+(aq) increase or decrease with the shift? Explain your answers in terms of the Le Châtelier’s principle.
HC2H3O2(aq) + H2O(ℓ) ⇋ C2H3O2–(aq) + H3O+(aq)
Solution
(1) NaC2H3O2 is an ionic substance, which dissociates 100% in solution to form Na+(aq) and C2H3O2–(aq). The acetate ion, C2H3O2–, is a product in the equilibrium. Adding products will cause the equilibrium to shift toward reactants. This will “use up” the extra C2H3O2– ions and form more reactants, HC2H3O2. So the reverse reaction will dominate until equilibrium is reestablished.
(2) Since the system is shifting toward reactants (reverse reaction is dominating), then the concentration of products is decreasing and the concentration of reactants is increasing until equilibrium is established. So the concentration of H3O+ will decrease.
Always Wear Safety Goggles and Use Good Lab Practices
Chemical Alert:
Fe(NO3)3+ KSCN Equilibrium
1. Start a hot water bath. Fill a 250 mL beaker about half full with water and place on a hot plate. Set the hot plate heat setting to about 4. Refer to Exp. 2 for the illustration of a hot water bath.
Start an ice water bath. Fill a 150 mL beaker about half full with ice and add enough water to cover the ice. Refer to Exp. 3 for the illustration of an ice bath.
2. Pour 15 mL of 1.7 × 10–3 M KSCN into a 50 mL beaker. Observe the color of the solution. Using a disposable pipet, add 5 drops of 0.2 M Fe(NO3)3 solution to the beaker containing the KSCN. Observe the color of the solution. This is your equilibrium mixture.
3. Prepare 6 aliquots by pouring approximately the same volume of the mixture into 6 small test tubes. Label the test tubes 1–6. The first test tube is your control or reference test tube.
4. Add a few crystals of KSCN to test tube #2 and stir well. Observe any color changes.
5. Add 1–2 drops of the Fe(NO3)3 solution to test tube #3 and stir well. Observe any color changes.
6. Add enough NaF solid (just enough to barely cover the tip of the spatula) to test tube #4 and stir well. Observe any color changes.
Fluoride ions react with Fe 3+ ions to remove the Fe 3+ ions from solution.
Fe3+(aq) + 6 F–(aq) → [FeF6]3–(aq)
7. Place test tube #5 in the hot water bath. (The water bath does not have to be boiling.)
Place test tube #6 in an ice water bath.
Record your observations for both test tubes after 5 minutes.
8. Pour the contents of all solutions into the waste container. Keep the hot water and ice water baths going for the next part.
CoCl2/Co(H2O)62+ Equilibrium
9. Pour 15 mL of the CoCl2(alc) into a clean, DRY, 50 mL beaker. (The beaker must be dry!) Observe the color of the solution. Using a disposable pipet, add 30 drops of deionized water to the beaker and stir well. Observe the color of the solution. This is your equilibrium mixture.
10. Prepare 6 aliquots by pouring approximately the same volume of the mixture into 6 small test tubes. Label the test tubes 1–6. The first test tube is your control or reference test tube.
11. Add one drop of 12 M HCl to test tube #2 and stir well. Observe any color changes.
12. Add 10 drops of the deionized water to test tube #3 and stir well. Observe any color changes. Since the CoCl2 solution has ethanol as the solvent, water is considered to be a reactant and not the solvent.
13. Add 1 drop of 0.1 M AgNO3 to test tube #4 and stir well. Observe any color changes.
Silver ions react with Cl– ions to remove the Cl– from solution by forming the white solid AgCl.
Ag+(aq) + Cl–(aq) → AgCl(s)
14. Place test tube #5 in the hot water bath. (The water bath does not have to be boiling.)
Place test tube #6 in an ice water bath.
Record your observations for both test tubes after 5 minutes.
15. Pour the contents of all solutions into the waste container.
Dismantle the hot water and ice water baths.
Determination
1. Summarize the results of this experiment.
2. What is the quality of the results? Do the results seem reasonable?
3. What were the errors or possible errors in the experiment?
4. How would the errors affect the results of the experiment?
1. Give the definitions of endothermic and exothermic.
2. Give a definition of Le Châtelier’s Principle in your own words. Give an example.
3. In the example equilibrium system, a drop of methyl orange indicator was added. For this indicator, the solution will be red in high concentrations of H3O+, and the solution will turn yellow in low concentrations of H3O+.
a. Initially the equilibrium solution was orange. What color is expected after the addition of the NaC2H3O2? Explain your reasoning.
b. Initially, the equilibrium solution was orange. What color is expected after the addition of a strong acid which donates H3O+ ions? Explain your reasoning.
c. Initially the equilibrium solution was orange. After heating the solution it changed to yellow-orange. Which way did the system shift? Explain your answer.
d. Initially the equilibrium solution was orange. Based on c., which color is expected if the mixture was placed in an ice bath? Explain your answer.
4. What would you have observed if you had added a concentrated solution of NaCl to the CoCl2/ Co(H2O)6+2 equilibrium system? Explain your answer in terms of shifts or Le Châtelier’s principle.
5. What would you have observed if you had added a reagent that removes water to the CoCl2 / Co(H2O)6+2 equilibrium system? Explain your answer in terms of shifts or Le Châtelier’s principle.
6. Equilibrium systems that include gas phases are also affected by Le Châtelier’s principle. One difference is the effect of pressure. If the total pressure of the system is increased, then the system will shift toward the reaction side that has fewer moles of gas. Of course, if the total pressure of the system is decreased, then the system will shift toward the reaction side that has more moles of gas.
For the equilibrium system below, state which way the system is expected to shift to offset the given changes. Then explain what happens to the partial pressure (or concentration) of SO2during the shifts.
2 SO2(g) + O2(g) ⇋ 2 SO3(g)
a. Increase the partial pressure of O2 (increase the concentration of O2).
b. Increase the partial pressure of SO3 (increase the concentration of SO3).
c. Increase the overall total pressure of the system.