Restoring Balance: LeChatelier's Principle and Equilibrium

Restoring Balance
LeChâtelier’s Principle and Equilibrium

Introduction Chemical equilibrium is a true balancing act. What happens when the balance is disturbed? The purpose of this lab is to observe the effects of concentration and temperature on equilibrium and to visualize how balance can be restored based on LeChâtelier’s Principle.

Background
Not all chemical reactions proceed to completion, that is, to give 100% yield of products. In fact, most chemical reactions are reversible. In the forward direction, reactants interact to make products, while in the reverse direction the products revert back to reactants. This idea is represented symbolically using double arrows.

In a closed system, any reversible reaction will eventually reach a dynamic balance between the forward and reverse reactions. A system is said to reach chemical equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. At this point, no further changes will be observed in the amounts of either the reactants or products. Chemical equilibrium can be further defined, therefore, as the state where the concentrations of reactants and products remain constant with time. This does not mean the concentration of reactants and products are equal. The forward and reverse reactions create an equal balance of opposing rates.

What happens when the balance is disturbed—due to the addition of more reactants or products or due to changes in the temperature or pressure? LeChâtelier’s Principle predicts how equilibrium can be restored:
“If an equilibrium system is subjected to a stress, the system will react in such a way as to reduce the stress.”

Any change that is made to a system at equilibrium is considered a stress—this includes adding or removing reagents or changing the temperature. To reduce the stress, one of two things can happen. A reversible reaction can shift in the forward direction and make more products, thus using up some of the reactants. Alternatively, the reaction can shift in the reverse direction and re-form the reactants, thus using up some of the products.

The effect of temperature on a system at equilibrium depends on whether a reaction is endothermic (absorbs heat) or exothermic (produces heat). If a reaction is endothermic, heat appears on the reactant side in the chemical equation. Increasing the temperature of an endothermic reaction shifts the equilibrium in the forward direction, to consume some of the excess energy and make more products. The opposite effect is observed for exothermic reactions. In the case of an exothermic reaction, heat appears on the product side in the chemical equation. Increasing the temperature of an exothermic reaction shifts the equilibrium in the reverse direction.

Experiment Overview
The purpose of this experiment is to investigate the effect of reaction conditions on the reversible formation of cobalt complex ions. When cobalt chloride hexahydrate (CoCl2•6H2O) is dissolved in ethyl alcohol, three different solute species are present: Co2+ cations, Cl- anions, and water molecules. These can react to form two different complex ions, Co(H2O)62+ , where the cobalt ion is surrounded by six water molecules, and CoCl42-, in which the metal ion is surrounded by four chloride ions.

Pre-Lab Questions
1. (a) Adding a species which appears on the right side of an equation will shift the equilibrium to which side of the equation?

(b) Adding a species which appears on the left side of an equation will shift the equilibrium to which side of the equation?

2. (a) Removing or decreasing the concentration of a species which appears on the right side of an equation will shift the equilibrium to which side of the equation?

(b) Removing or decreasing the concentration of a species which appears on the left side of an equation will shift the equilibrium to which side of the equation?

3. Dissolving ammonium chloride in water is an endothermic reaction. Use LeChâtelier’s Principle to predict whether ammonium chloride will be more soluble in hot or cold water.

NH4Cl(s) + heat  NH4+(aq) + Cl-(aq)

Materials Cobalt chloride hexahydrate,
CoCl2•6H2O, 1% solution in alcohol Acetone, small dropper bottle Beaker, 50-mL, 250-mL (2) Calcium chloride, CaCl2, 4 pellets Graduated cylinder, 10mL Hydrochloric acid, HCl, 12M (conc.) Pipet Silver nitrate solution, AgNO3, 0.1 M Metal spatula Distilled water, small dropper bottle Stirring rod 3 x 5 card Test tubes, small, 6 Hot plate Thermometer Ice Masking tape

Safety Precautions
Concentrated hydrochloric acid is highly toxic by ingestion or inhalation and is severely corrosive to skin and eyes; can cause severe body tissue burns. Ethyl alcohol and acetone are flammable solvents. Cobalt (II) chloride solution is moderately toxic by ingestion. Silver nitrate solution is corrosive and will stain skin and clothing.

Procedure Preparation
1. Prepare hot-water and ice-water baths for Part B: Fill a 250-mL beaker half full with distilled water. Place it on a hot plate and heat to 65-70◦C for use in step 13. In a second 250-mL beaker, add water and ice to prepare an ice-water bath for use in step 14.

2. Thoroughly dry a 50-mL beaker with a paper towel, then use the markings on the side of the beaker to obtain about 20 mL of a 1% solution of cobalt chloride in alcohol.

3. Label six dry test tubes A-F and place them in a test tube rack. Use masking tape.

Part A. Effect of Concentration

4. Using a pipet, add about 2 mL of the cobalt chloride solution to each test tube A-F. Note: The exact volume is not important, but try to keep the volume of solution approximately equal in each test tube. Use a graduated cylinder to measure one volume, then fill each test tube to same amount.

5. Set aside test tube A as control. Record the color and appearance of the control solution in the data table.

6. To test tube B, add 4 drops of distilled water, one drop at a time. Record the color of the solution after each drop.

7. Add 4 drops of distilled water to each of the next 3 test tubes C, D, and E. Note: The color of the solutions should be the same in test tubes B-E at this point.

8. Take test tube C to the fume hood. Use the dropper provided on the acid bottle to carefully add 6 drops of concentrated hydrochloric acid to the test tube.

9. Gently swirl test tube C to mix the contents, then return the test tube to the test tube rack on your lab bench. Record the color of the solution in the data table.

10. To test tube D, add 4 small pellets of solid calcium chloride and gently stir the solution to dissolve the solid. Record the color and appearance of the solution in the data table.

11. To test tube E, add about 1 ml of acetone. Gently swirl the test tube to mix the contents and record the color of the solution in the data table.

12. To the last test tube F, add 5 drops of 0.1 M silver nitrate and gently swirl the test tube to mix the contents. Record the color and appearance of the mixture in the data table.

Part B. Effect of Temperature

13. Place test tube B from Part A in the hot-water bath at 65-70◦C for 2-3 minutes. Record the initial and final color of the solution in the data table.

14. Place test tube C from Part A in the ice-water bath at 0-5◦C for 5 minutes. Record the initial and final color of the solution in the data table.

NOTE: All final solutions should be one color. If two colors are present (such as top pink and bottom blue) then stir the contents thoroughly until the solution is one consist color.

Cleanup and Disposal
Dispose of solutions in the waste container in the fume hood. Thoroughly wash all glassware.

Data Table

Part A. Effect of Concentration
Test Tube Reagents Observations
A CoCl2 in alcohol (control)
B CoCl2 in alcohol + water 1 drop 2 drops 3 drops 4 drops
C CoCl2 in alcohol + water + HCl
D CoCl2 in alcohol + water + CaCl2
E CoCl2 in alcohol + water + acetone
F CoCl2 in alcohol + silver nitrate
Part B. Effect of Temperature
B CoCl2 in alcohol + water Initial Color Final color after heating to 75-80◦C
C CoCl2 in alcohol + water + HCl Initial color Final color after cooling to 0-5◦C