Materials Included In Kit

Additional Materials Required

Safety Precautions

Hydrochloric acid and sodium hydroxide solutions may cause skin and eye burns. Potassium thiocyanate solution is irritating to skin and eyes. Iron(III) nitrate solution is also a possible skin and body tissue irritant; it will stain clothes and skin. Avoid contact of all chemicals with eyes and skin. Clean up all chemical spills immediately. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Remind students to wash hands thoroughly with soap and water before leaving the lab. Please review current Safety Data Sheets for additional 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. The leftover reaction mixtures may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b. Save all remaining solutions and reagents for future use.

Lab Hints

• The laboratory work for this experiment can reasonably be completed within a typical 2-hour lab period. The Prelab Assignment should be completed by students as preparation for lab, and should be reviewed during prelab discussion.
• Enough disposable pipets are recommended in the Materials section so that each pair of students has a separate pipet for each solution. Encourage students to label their pipets to avoid contamination and waste. The pipets may be colorcoded, for example, using colored tape. If the reagents are placed in dropping bottles, the number of pipets may be reduced. To avoid congestion in the materials dispensing area, stagger the starting points for the two parts of the experiment. (Half the student groups start with the complex-ion equilibrium, the other half with the acid–base equilibrium.)
• Reaction of iron(III) nitrate with potassium thiocyanate may also be viewed as a double replacement reaction. The products are a series of complex ions having the general formula Fe(SCN)_n^3–n, where n = 1– 4. Although the red product is generally represented as FeSCN²⁺, all of the products are possible and all are deep red. The neutral compound Fe(SCN)₃ can be extracted from aqueous solution using ether.
• The effect of temperature on the equilibrium in the Fe³⁺/FeSCN²⁺ system may be counterintuitive. Most students rightly expect that the rate of a reaction always increases with temperature. This is true, but the rates of the forward and the reverse reactions will increase by different amounts because they have different activation energies. In the Fe³⁺/FeSCN²⁺ example, the rate of the reverse reaction must increase more than the rate of the forward reaction when the temperature is raised.
• In contrast to the effect of adding or removing reactants and products on the position of equilibrium, the effect of temperature cannot be generalized without further information. The equilibrium constant for a reaction depends on temperature. When Le Chatelier’s principle is introduced, students can determine whether a reaction is exothermic or endothermic as written based on the effect of temperature on the position of equilibrium. In this exploratory introduction to equilibrium, the temperature effect may be discussed in terms of the rates of the forward and reverse reactions.
• The green transition color (step 16) in the reversible reaction of bromcresol green is easy to overshoot. Students should carefully add HCl drop by drop and gently swirl the solution between drops. Have students try to add half a drop at a time. (Squeeze out a small amount from the pipet, press the pipet tip against the side of the test tube to dislodge the halfdrop, and then swirl the test tube contents to mix the half-drop into solution.) The pH range for the color transition is 3.8–5.4.
• Moisture-sensitive test paper coated with cobalt chloride is available under the trade name Hydrion Humidicator Paper (see Flinn Catalog No. AP4656). The test strips are coated with anhydrous cobalt chloride that changes from blue to pink when it is hydrated. A color chart is available to estimate relative humidity levels between 20 and 80%.
• Equilibrium is one of the most challenging topics in the chemistry curriculum. The dynamic nature of equilibrium, in particular, seems to be prone to misunderstanding. The purpose of this experiment is to allow students to discover the key principles of equilibrium based on direct observations and logical reasoning. The reactions A and B can occur in both directions. Why then do the reactions not proceed to completion under conditions where reactants and products are clearly present? Is it possible that some molecules react, but that others do not? This is the hard part! Some students will probably not be able to “make the leap” and conclude that both reactions are occurring. Those students who accept the idea that both the forward and reverse reactions must take place simultaneously will then easily conclude that they must occur at the same rate (otherwise the color would change).
• The activity of hemoglobin, the main oxygen-binding protein in red blood cells, illustrates a biomedical application of complex-ion equilibrium. Hemoglobin (Hb) contains four iron(II) ions that bind to oxygen molecules. This is a reversible reaction, since hemoglobin must release the bound oxygen molecules in cells and body tissues (Equation 5).
  {14038_Hints_Equation_5}
  Students should be able to apply what they have learned in this experiment to explain the effects of high altitudes on humans. At high altitudes, where the concentration of oxygen is lower, the equilibrium shown in Equation 5 is shifted in the reverse direction. Less oxygen is therefore available in the bloodstream to be transported to the cells. The physical symptoms of reduced oxygen availability are fatigue and dizziness. The human body, however, is marvelous in its adaptability. People who live or train at high altitudes compensate for the reduced oxygen supply by synthesizing more red blood cells. Increasing the concentration of hemoglobin increases the rate of the forward reaction and thus increases the amount of available oxygen.
• An interesting application of acid–base equilibrium is found in the carbonic acid–bicarbonate buffering system in the blood (Equation 6). The “blood buffer” is able to maintain blood pH within a very narrow range despite the acidifying effect of carbon dioxide produced by metabolism (approximately 350 g of CO₂ are produced per 1000 Calories of food burned). People with impaired lung function, due to emphysema, for instance, are not able to exchange CO₂ efficiently between the lungs and air as they exhale (Equations 7 and 8). The result is an increase in the amount of carbonic acid in the blood, making the blood more acidic. This condition is called respiratory acidosis.
  {14038_Hints_Equation_6}
  {14038_Hints_Equation_7}
  {14038_Hints_Equation_8}
• A particulate model may help students visualize the presence of both reactants and products at equilibrium. Ask students to draw pictures of what they imagine the Fe³⁺–SCN^– system looks like at equilibrium. Then ask them to add more reactants or products to their picture and predict what a new equilibrium position would look like.
  {14038_Hints_Figure_1}

Answers to Prelab Questions

1. True or False: At equilibrium, no more reactants are transformed into products. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.

   False. At equilibrium, the rate at which reactants are transformed into products is equal to the rate at which products are transformed back into reactants.

2. True or False: At equilibrium, the concentrations of reactants and products are equal. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.

   False. At equilibrium, the concentrations of reactants and products are constant. (Alternatively: At equilibrium, the forward and reverse reaction rates are equal.) Note: The concentrations of reactants and products may be equal, but that would be a special case.

3. Paper coated with cobalt chloride is sold commercially as moisture-sensitive test strips to estimate relative humidity levels between 20 and 80 percent in air. The following reversible reaction takes place with water:
   {14038_PreLabAnswers_Reaction_1}
   1. What color do you think the paper will be when the humidity is low (20%)? What color will it be when the humidity is high (80%)? Explain.

      The paper should be blue when the relative humidity is low (20%), because a low concentration of the reactant H₂O means that the rate of the forward reaction will also be low. When the relative humidity is high (80%), and the concentration of water needed to react with CoCl₂ is greater, the paper should turn pink.

   2. The test strips come with a color chart to estimate intermediate humidity levels. Predict the intermediate color that might be observed when the humidity is about 50%. State your reasoning.

      At intermediate humidity levels, appreciable amounts of both reactants and products should be present at equilibrium, and an intermediate or transition color should be observed that is midway between blue and pink—lavender.

4. The effect of adding more reactant to a system at equilibrium is often explained or predicted using Le Chatelier’s principle. Look up this principle in a textbook or online and restate it here to explain the color change predicted in Question 3a at high humidity.

   Le Chatelier’s principle states that if an equilibrium system is subjected to a stress, the system will react in such a way as to reduce the stress. In the context of Question 3a, adding or increasing the concentration of water vapor at high humidity is a stress. The cobalt chloride system reacts with the additional water vapor, thereby reducing its amount (tbe stress), forming the hydrate and turning pink.

Sample Data

Complex-Ion Equilibrium of Iron(III) and Thiocyanate Ions

Acid–Base Equilibrium of Bromcresol Green

Answers to Questions

1. Write the chemical equation for the reversible reaction of iron(III) ions with thiocyanate ions. Label this Equation A. Use the information in the data table to write the color of each reactant and product underneath its formula.
   {14038_Answers_Equation_A}
2. How did the color of the solution change when additional reactant—either Fe(NO₃)₃ in step 6 or KSCN in step 7—was added? Explain the observed color changes: Adding more reactant to an equilibrium mixture of reactants and products increases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.

   Adding extra Fe(NO₃)₃ or extra KSCN produced the same effect—the color of the solution changed from orange to dark red. The dark red color indicates that more product was formed under these conditions. Adding more reactant to an equilibrium mixture of reactants and products increases the rate of the forward reaction and thus increases the amount of product.

3. How do the results obtained in steps 6 and 7 demonstrate that both reactants and products are present at equilibrium?

   Adding either reactant alone increased the amount of product. Since additional product formed when either reactant was added, the other reactant must already be present in solution. Both reactants and products are present at equilibrium.

4. How did the color of the solution change when it was cooled (step 8) or heated (step 9)? How do these results demonstrate that the reaction shown in Equation A does indeed occur in both the forward and reverse directions?

   Opposite color changes were observed when the control solution was cooled or heated. The original orange solution turned red-orange when it was cooled (step 8), yellow when it was heated (step 9). These results indicate that the reaction can indeed “go both ways.” The stock solution must contain equilibrium concentrations of reactants and products. When the solution was cooled, the concentration of the red product increased (net reaction in forward direction). When the solution was heated, the concentration of the red product decreased and more of the reactants were formed (net reaction in reverse direction).

5. In step 10, H₂PO₄^– ions combined with iron(III) ions and removed them from solution. How did the color of the solution change when NaH₂PO₄ was added? Explain the observed color change: Removing one of the reactants from an equilibrium mixture of reactants and products decreases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.

   Adding sodium phosphate decolorized the solution—the red color disappeared and the solution turned light yellow and cloudy. The amount of product decreased. Removing one of the reactants from an equilibrium mixture of reactants and products decreases the rate of the forward reaction and thus decreases the amount of product.

6. After observing the effect of NaH₂PO₄ on the equilibrium mixture in step 10, a student was skeptical that both Fe³⁺ and SCN^– ions were still present. Suggest additional experiments that could be done to prove that both reactants are still present at this point.

   Try adding more of either reactant separately to the reaction mixture at this point. If Fe³⁺ is still present in solution, it will react with the added SCN^–. Similarly, if SCN^– is still present in solution, it will react with the added Fe³⁺. In either case, the solution should turn red again. Note: Try it! It works.

7. Write the chemical equation for the reversible reaction of bromcresol green with water. Label this Equation B. Hint: Refer to Equation 3 in the Background section.
   {14038_Answers_Equation_B}
8. Use the color changes observed for the indicator before and after adding HCl (steps 12 and 13) to predict the colors of the HIn and In^– forms of bromcresol green. Write the colors of HIn and In^– underneath their formulas in Equation B. Explain your reasoning. Hint: Adding HCl increases the concentration of H⁺ ions. Which reaction, forward or reverse, would that increase?

   The colors of HIn and In^– are shown in Equation B. The colors can be inferred based on the color change observed when HCl was added to the initial indicator solution. The initial indicator color was blue; when HCl was added, the indicator turned yellow. Adding H⁺ ions (in the form of HCl) should increase the rate of the reverse reaction. When a new equilibrium is re-established, there will be a greater concentation of HIn. The yellow color must be due to HIn, the blue color to In^–.

9. Explain the observed color change: Adding more product to an equilibrium mixture of reactants and products increases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.

   Adding more product to an equilibrium mixture of reactants and products increases the rate of the reverse reaction and thus decreases the amount of product.

10. In step 14, hydroxide ions reacted with and removed H⁺ ions from solution (see Equation 4 in the Background section). What color change was observed when NaOH was added? Explain the observed color change: Removing one of the products from an equilibrium mixture of reactants and products decreases the rate of the (forward/reverse) reaction and thus (increases/decreases) the amount of product.

    The indicator color changed from yellow to blue when NaOH was added. Adding NaOH increased the amount of product present at equilibrium. Removing one of the products from an equilibrium mixture of reactants and products decreases the rate of the reverse reaction and thus increases the amount of product.

11. What form(s) of the indicator were most likely present when the transition color was observed in step 16? How does this observation provide visual proof that not all reactions “go to completion?”

    The transition color of bromcresol green is green. The green color, midway between yellow and blue, suggests that approximately half of the available indicator molecules are present in the uncharged form HIn (yellow) and half in the ionic form In^– (blue). The green color offers visual proof that both reactants and products are present.

Introduction

The word equilibrium has two Latin roots: aequi, meaning equal, and libra, meaning weight or balance. Our physical sense of equilibrium—in the motion of a seesaw or the swing of a pendulum—suggests an equal balance of opposing forces. How does this physical sense of equilibrium translate to chemical equilibrium? Let’s explore the nature and consequences of equilibrium in chemical reactions.

Concepts

• Reversible reactions
• Reaction rates
• Chemical equilibrium
• Complex-ion reaction
• Acid–base indicators
• Le Chatelier’s principle

Background

Physical changes, such as melting ice or dissolving sugar, are often introduced by noting that these processes can be easily reversed. Some common examples of chemical change, such as burning wood or spoiling food, generally cannot be reversed. A closer look at chemical change, however, reveals that many chemical reactions are also reversible.

Consider the following example of a reversible chemical reaction. At high pressures and in the presence of a catalyst, nitrogen and hydrogen react to form ammonia. If the temperature is high enough, however, ammonia decomposes to reform its constituent elements. The reaction can go both ways! This reversible reaction is represented symbolically using double arrows (Equation 1).

What happens when nitrogen and hydrogen are allowed to react? In a closed system, the concentrations of nitrogen and hydrogen will decrease and the concentration of ammonia will steadily increase as the reaction proceeds in the forward direction. Soon, however, the concentration of ammonia will be large enough that the reverse reaction will begin to take place at a significant rate as well. Eventually, as the reaction occurs in both the forward and the reverse directions, the number of ammonia molecules being formed will become equal to the number of ammonia molecules being consumed. At this point, no further changes will be observed in the overall concentrations of nitrogen, hydrogen, and ammonia. This is the point of chemical equilibrium.

Chemical equilibrium is defined as the state where the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of reactants and products remain constant with time. Note that this definition describes a dynamic picture of equilibrium. The reactions continue, but there is an equal balance of opposing reaction rates.

What happens when the equilibrium is disturbed? Any factor that changes the rate of either the forward or the reverse reaction will change the amounts of reactants and products that are present at equilibrium. Reaction conditions that are known to affect the rates of chemical reactions include the concentrations of reactants and the temperature. In this experiment, we will investigate how changes in reaction conditions affect the amounts of reactants and products present at equilibrium.

Experiment Overview

The purpose of this experiment is to explore the nature of chemical equilibrium and to identify conditions that affect the position of equilibrium. Two different reversible reactions will be studied.

Reaction of iron(III) nitrate with potassium thiocyanate will be used to study complex-ion equilibrium. Iron(III) ions react with thiocyanate ions to form FeSCN²⁺ complex ions (Equation 2). The effects of changing the concentrations of reactants and the reaction temperature will be investigated.

The properties of an indicator will be used to study acid–base equilibrium. An indicator is a dye that can gain or lose hydrogen ions (H⁺) to form substances that have different colors. Equation 3 summarizes the reversible reaction of a general indicator, such as bromcresol green. HIn represents an indicator molecule and In^– an indicator anion formed after the molecule has lost a hydrogen ion. The color of the indicator in the presence of either excess H⁺ or OH^– ions (see Equation 4) will show how changing the concentration of a product affects the equilibrium shown in Equation 3. 

Materials

Prelab Questions

1. True or False: At equilibrium, no more reactants are transformed into products. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.
2. True or False: At equilibrium, the concentrations of reactants and products are equal. If false, rewrite the statement so that it correctly describes the nature of chemical equilibrium.
3. Paper coated with cobalt chloride is sold commercially as moisture-sensitive test strips to estimate relative humidity levels between 20 and 80 percent in air. The following reversible reaction takes place with water:
   {14038_PreLab_Reaction_1}
   1. What color do you think the paper will be when the humidity is low (20%)? What color will it be when the humidity is high (80%)? Explain.
   2. The test strips come with a color chart to estimate intermediate humidity levels. Predict the intermediate color that might be observed when the humidity is about 50%. State your reasoning.
4. The effect of adding more reactant to a system at equilibrium is often explained or predicted using Le Chatelier’s principle. Look up this principle in a textbook or online and restate it here to explain the color change predicted in Question 3a at high humidity.

Safety Precautions

Hydrochloric acid and sodium hydroxide solutions may cause skin and eye burns. Potassium thiocyanate solution is irritating to skin and eyes. Iron(III) nitrate solution is also a possible skin and body tissue irritant; it will stain clothes and skin. Avoid contact of all chemicals with eyes and skin. Notify the instructor and clean up all chemical spills immediately. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Complex-Ion Equilibrium of Iron(III) and Thiocyanate Ions

1. Fill two beakers (250- or 400-mL) half-full with tap water. Add ice to one beaker to prepare an ice-water bath (0–5 °C) for use in step 8. Heat the second beaker on a hot plate to prepare a hot water bath (70–80 °C) for use in step 9. Do not boil the water.
2. Observe and record the initial colors of the Fe(NO₃)₃ and KSCN solutions.
3. Prepare a stock solution of FeSCN²⁺: In a clean 50-mL beaker, measure 40 mL of distilled water. Using separate Beral-type pipets for each solution, add 1 mL of 0.1 M Fe(NO₃)₃ and 2 mL of 0.1 M KSCN. Mix thoroughly with a stirring rod.
4. Label six clean test tubes 1–6. Using a graduated, Beral-type pipet or eyedropper, add 1 mL of the FeSCN²⁺ stock solution to each test tube 1–6.
5. Add 10 drops of distilled water to test tube 1. Gently swirl the test tube to mix the solution and record the color of the solution in the data table. Test tube 1 will be used as the control solution for color comparison purposes in steps 6–10.
6. Add 10 drops of 0.1 M Fe(NO₃)₃ to test tube 2. Gently swirl the test tube to mix the solution and compare the color of the resulting solution to the control in test tube 1. Record the color of the solution.
7. Add 10 drops of 0.1 M KSCN to test tube 3. Gently swirl the test tube to mix the solution and compare the color of the resulting solution to the control. Record the color.
8. Add 10 drops of distilled water to test tube 4 and place the sample in an ice-water bath. After 3–5 minutes, remove the test tube from the ice bath and compare the color to the control. Record the color.
9. Add 10 drops of distilled water to test tube 5 and place the sample in a hot water bath at 70–80 °C. After 2–3 minutes, remove the tube from the hot water bath and compare the color of the solution to the control. Record the color.
10. To test tube 6, add 10 drops of 0.1 M NaH₂PO₄. Record the color and appearance of the mixture.
11. Wash the contents of the test tubes down the drain with excess water and rinse with distilled water.

Acid–Base Equilibrium of Bromcresol Green

12. Obtain 2 mL of distilled water in a clean test tube and add 5 drops of 0.04% bromcresol green. Swirl gently and record the color of the solution in the data table.
13. Add 3 drops of 0.1 M HCl solution to the test tube. Swirl gently and record any color change(s).
14. Add 0.1 M NaOH dropwise to the solution until the original color is restored. Shake gently and record the number of drops of NaOH added and the color of the solution.
15. Continue adding 0.1 M NaOH dropwise until a total of 5 drops of NaOH have been added in steps 14 and 15 combined.

Can the process be reversed to obtain a color that is intermediate between that in steps 13 and 14?

16. Add 0.1 M HCl again dropwise very slowly until the solution reaches a “transition” color midway between the two colors recorded above (steps 13 and 14). Swirl gently between drops to avoid overshooting the transition color. Record the number of drops of HCl required and the color. Note: It may be necessary to add half-drops.
17. Wash the contents of the test tube down the drain with excess water and rinse with distilled water.

Student Worksheet PDF

14038_Student1.pdf