Materials Included In Kit

Additional Materials Required

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. Cobalt(II) chloride solution is moderately toxic by ingestion; causes blood damage. Silver nitrate solution is corrosive and will stain skin and clothing. Calcium chloride is slightly toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult the appropriate Safety Data Sheets for further safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Solutions containing silver nitrate and silver chloride can be disposed of according to Flinn Suggested Disposal Method #11. Solutions containing cobalt(II) chloride can be disposed of according to Flinn Suggested Disposal Method #27f. Alternatively, the solutions can be combined and filtered to remove insoluble silver chloride, which can be dried and packaged for landfill disposal according to Flinn Suggested Disposal Method #26a. The combined filtrate can then be neutralized according to Flinn Suggested Disposal Method #24b and saved in a disposal container reserved for heavy-metal waste.

Teacher Tips

• Enough materials are provided in this kit for 30 students working in pairs, or for 15 groups of students. The laboratory procedure can reasonably be completed in one 50-minute class period.
• The solution colors can be distinguished more easily (especially in the lavender/blue transition) if the test tubes are held against a white background. Use a notecard or a piece of paper.
• The pink and blue equilibrium due to cobalt chloride in aqueous solution is the basis of moisture-sensitive paper such as Hydrion Humidicator Paper, which is used to measure the relative humidity in air. The paper strips are impregnated with deep blue anhydrous CoCl₂ powder that changes to bright pink when it is hydrated. The intermediate colors between blue and pink are clear and definite, and a color chart is available to estimate relative humidity levels between 20 and 80%.
• Many students misinterpret chemical equilibrium as a static, rather than dynamic, condition. The standard definition of equilibrium, namely, that at equilibrium the concentrations of reactants and products remain unchanged, is frequently misunderstood to mean that the concentrations of reactants and products have constant values. It is the equilibrium constant, the ratio of product to reactant concentrations, governed by the stoichiometry of the balanced chemical equation, that is constant. The concentrations of individual reactants and products are affected by changes in the other terms in the equilibrium constant expression. And the equilibrium “constant” itself, of course, is also temperature dependent!
• When the blue solution corresponding to CoCl₄^2– is diluted by the addition of water, the effect is to shift the equilibrium back to the left, toward Co(H₂O)₆²⁺. Students will interpret this qualitatively as due to the addition of “excess” product (water). Technically, however, the concentration of water in an aqueous solution is constant. A more correct quantitative explanation involves the equilibrium constant expression (K_eq) for the reaction, which contains one concentration term in the numerator but two terms in the denominator.
  {11961_Tips_Equation_1}
  Reducing each concentration term in this expression by a factor of about one-third (due to dilution with water) means that the ratio becomes greater than K_eq. The reaction shifts back to reactants so that the system return to equilibrium and the ratio becomes equal to K_eq again.

Sample Data

Answers to Questions

1. Write the chemical equation for the complex-ion equilibrium that results when excess chloride ion is added to an aqueous solution of cobalt chloride. Note the observed color of each complex ion underneath its chemical formula.
   {11961_Answers_Equation_1}
2. What is the likely composition of the solution (relative amounts of the two different complex ion forms) when the intermediate or transition color is observed in step 6? How does this observation provide visual proof of the idea that not all reactions “go to completion”? Explain.
   The intermediate (lavender) color of the solution suggests that at this point half of the total available cobalt ions are present in the form of Co(H₂O)₆²⁺ complex ions (pink) and half are in the form of CoCl₄^2– complex ions (blue). The color, midway between pink and blue, is visual proof that at this point both reactants and products must be present at equilibrium, that is, the reaction does not necessarily go to completion.
3. Use Le Chatelier’s Principle to explain the color changes observed upon addition of water and calcium chloride to an equilibrium mixture of the two complex ions in this reaction (steps 7 and 8).
   When chloride ion was added to an equilibrium mixture of the two complex ions, the solution turned blue, indicating that all of the cobalt ions were converted to the blue CoCl₄^2– complex ion. In terms of LeChâtelier’s Principle, addition of excess reactant shifted the equilibrium shown in Question #1 to the right in order to consume some of the added reactant and restore the equilibrium condition. Adding water to an equilibrium mixture of the two complex ion forms caused the solution to turn pink as more Co(H₂O)₆²⁺ was formed. This is consistent with the restoring effect of LeChâtelier’s Principle, whereby adding more of one of the products of the reaction drives the equilibrium to the left, to consume some of the excess product. (See Tip #5.)
4. What was the effect of adding AgNO₃ on the position of equilibrium for these two complex ions? Is this effect consistent with LeChâtelier’s Principle? Explain.
   The color of the blue solution (containing mainly CoCl₄^2–) reverted to pink and a white precipitate (AgCl) settled out when AgNO₃ was added. This effect is consistent with predictions made using Le Chatelier’s Principle—removing one of the reactants of a reaction, in this case chloride ion, shifts the equilibrium to the left, toward reactant formation, in order to restore the balance between reactants and products.
5. How was the composition of the solution (relative amounts of the two complex ions) affected when the solution was heated (step 10)? When the solution was cooled (step 11)?
   Heating the pink solution (consisting mainly of Co(H₂O)₆²⁺), caused the solution to turn blue. When the resulting blue solution was cooled the solution reverted to its original pink color.
6. Based on the observed effect of temperature on the position of equilibrium, is the forward reaction for the equation in Question 1 endothermic or exothermic? Explain, using Le Chatelier’s Principle.
   The forward reaction in Question 1, formation of the CoCl₄^2– complex-ion product, must be endothermic. In accordance with Le Chatelier’s Principle, increasing the temperature of an endothermic chemical reaction shifted the reaction in the forward direction to make more product, in order to “consume” some of the excess heat. Cooling the solution reversed this effect—it drove the reaction back to the left in order to replace some of the heat that was lost when the solution was cooled. The cool solution was pink.

References

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison, (1983); pp 338–343.
Vonderbrink, Sally Ann Laboratory Experiments for Advanced Placement Chemistry; Flinn Scientific: Batavia, IL, 1995; pp 109–113.

Introduction

Chemical equilibrium is a true chemical 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 Le Chatelier’s Principle.

Concepts

• Chemical equilibrium
• Le Chatelier’s principle
• Complex-ion equilibrium
• Exothermic vs. endothermic reactions

Background

Not all chemical reactions proceed to completion, that is, to give 100% yield of products. In fact, most chemical reactions are reversible, meaning they can go both ways. 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, as shown below for the reversible reaction of nitrogen and hydrogen gas to make ammonia (Equation 1).

Any reversible chemical reaction will eventually reach a dynamic balance between the forward and reverse directions. 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 are 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. The forward and reverse reactions create an equal balance of opposing forces.

What happens when the balance is disturbed—due to the addition of excess reactants and products or due to changes in the temperature or pressure? LeChâtelier’s Principle provides an intuitive explanation for how balance can be restored: To remove the stress, one of two things can happen. A reversible reaction can shift in the forward direction to make more products, thus using up reactants. Alternatively, the reaction can shift in the reverse direction to re-form the reactants, thus using up products. Adding extra reactant to an equilibrium mixture of reactants and products shifts a reversible chemical reaction in the forward direction. Removing a product has the same effect. Conversely, adding extra product or removing a reactant has the opposite effect, shifting a reaction in the reverse direction to use up the excess product and make more reactants.

The effect of temperature on a system at equilibrium depends on whether a reaction is endothermic (requires heat) or exothermic (produces heat). If a reaction is endothermic, heat may be considered as a reactant in the forward reaction. Increasing the temperature of an endothermic reaction shifts the position of equilibrium in the forward direction, to use up the excess heat and make more products. The opposite effect is observed for exothermic reactions. In the case of an exothermic reaction, heat may be thought of as a product of the forward reaction. Increasing the temperature of an exothermic reaction shifts the position of equilibrium in the reverse direction, to use up the excess heat. For example, the ammonia synthesis reaction is exothermic (Equation 2). Increasing the temperature of this reaction shifts the equilibrium in the reverse direction, to use up some of the excess heat. Thus, ammonia decomposition becomes more significant at higher temperatures.

In this experiment, the nature of complex-ion equilibria will be examined by observing the effect of a variety of reaction conditions on the reversible reaction of aqueous cobalt ion with chloride ion. In aqueous solution cobalt ion exists in the form of Co(H₂O)6²⁺, a complex ion, with six water molecules tightly held by the cobalt ion. Adding chloride ions in the form of concentrated hydrochloric acid solution results in the reversible formation of a different complex ion, CoCl₄^2–, in which the metal ion is surrounded by four chloride ions (Equation 3). What are the predicted effects of adding more water, chloride ion, or silver ion to an equilibrium mixture consisting of both complex ions shown in Equation 3? (Hint: Silver ion reacts with chloride ion to form insoluble silver chloride via Equation 4. Adding Ag⁺ has the effect, therefore, of removing chloride ion from the equilibrium mixture.) The effect of temperature on this equilibrium reaction will also be investigated in order to determine whether the reaction is endothermic or exothermic in the forward direction.

Materials

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. Cobalt(II) chloride solution is moderately toxic by ingestion; causes blood damage. Silver nitrate solution is corrosive and will stain skin and clothing. Calcium chloride is slightly toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron.

Procedure

1. Prepare hot-water and ice-water baths: Fill a 250-mL beaker half full with water. Place it on a hot plate and heat to 80–85 °C. In a second 250-mL beaker prepare an ice-water bath for use in Step 11.
2. Obtain four graduated Beral-type pipets and label them as follows: CoCl₂, H₂O, HCl, and AgNO₃. Use the appropriate labeled pipet to add each solution in Steps 3–9 below.

Reversible Reaction of Cobalt(II) with Water and Chloride Ion

3. Label four test tubes A–D and add 20 drops of aqueous 0.1 M CoCl₂ solution to each.
4. Add 1 mL of distilled water to test tube A. Record the color of the solution in the data table. This is the characteristic color of the Co(H₂O)₆²⁺ complex ion. Set this tube aside as a control solution.
5. Slowly and carefully, add 1 mL of concentrated HCl to each test tube B, C and D. (Dispense concentrated HCl in a fume hood or a well-ventilated lab only.) Use a glass stirring rod to mix each solution. Record the color of the solutions in the data table. This is the characteristic color of the CoCl₄^2– complex ion. Set test tube B aside as a reference solution.
   Can the process be reversed to obtain a color that is intermediate between the colors of the control and reference solutions?
6. Add distilled water dropwise to each test tube C and D until the solutions reach an intermediate or transition color midway between that observed in steps 4 and 5. Record the amount of water added and the transition color in the data table. Effect of Concentration
7. Continue adding distilled water dropwise to test tube C until the color of the solution is stable and no longer changing. Record the color of the solution in the data table.
8. Add a few pieces of solid calcium chloride to test tube D. Stir gently and record the color of the solution in the data table.
9. Add 1 mL of 0.1 M AgNO₃ solution to the bottom of the solution in test tube D. Record the color and appearance of the solution in the data table. Effect of Temperature
10. Place test tube C in a hot water bath at 80–85 °C for 2–3 minutes. Record the observed color changes in the data table.
11. Using a test-tube holder, remove the hot test tube from the hot water bath. Cool it briefly and then immerse it directly in an ice water bath for 2–3 minutes. Record the final color of the solution in the data table.
12. Dispose of the solutions as directed by your instructor.

Student Worksheet PDF

11961_Student1.pdf