Materials Included In Kit

Additional Materials Required

Prelab Preparation

Activity B

• Potassium Thiocyanate, KSCN, 0.002 M: Dilute 5 mL of 0.10 M KSCN to 250 mL with distilled or deionized water. Mix well prior to dispensing.

Activity C
Two Water Baths: Prepare water baths at different temperatures as follows:

• Add crushed ice and water to a 250-mL beaker for a bath at about 0 ºC.
• Add 200 mL of water to a 250-mL beaker and heat it on a hot plate on to prepare a hot water bath at around 65–70 ºC.

Safety Precautions

Cobalt chloride solution is moderately toxic by ingestion. Iron(II) nitrate solution may be a skin/tissue irritant. Concentrated hydrochloric acid is highly toxic by ingestion or inhalation and is severely corrosive to skin and eyes; can cause severe body tissue burns. Instruct students to notify the teacher immediately in case of an acid spill. Dilute hydrochloric acid and sodium hydroxide solutions are skin and eye irritants. Potassium thiocyanate is toxic by ingestion and emits a toxic gas if strongly heated—do not heat this solution and do not add acid. Sodium phosphate monobasic is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash hands thoroughly with soap and water before leaving the lab.

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. Cobalt-containing wastes from Activity C may be combined and treated according to Flinn Suggested Disposal Method #27f. Other waste solutions from Activity C, Activity B and Activity D may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• For best results, set up two stations for each activity throughout the lab. This will allow up to eight groups of students to rotate through four activities in a 50-minute lab period, if needed. A double lab period (two 50-minute class periods) will allow time both for a review of basic equilibrium concepts before lab and for a collaborative class discussion after lab.
• The activities may be completed in any order. Also, since each activity is a self-contained unit, the experiment may be set up with as many or as few of the activities as the teacher desires. Students should need only 10 minutes per station—keep the pace fairly brisk to avoid dawdling. Questions in the Observations and Results section may be answered during downtime between stations.
• Prelab Preparation is an essential component of lab safety, and it is also critical for success in the lab. (Standing in front of the lab station is not a good time for students to be reading the activity for the first time.) Having students complete the written prelab assignment for this lab will help teachers ensure that students are prepared for and can work safely in the lab.
• In Activity A, the green transition color (step 12) 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 half-drop, 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.
• In Activity C, less concentrated hydrochloric acid may also be used, but the pink-to-blue transition will not be complete. Using 9 M HCl causes a color change from pink to blue-violet while 6 M HCl only causes a color change from pink to lavender.
• In Activity C, color changes 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. Always compare the color changes against an appropriate control—B versus A, C versus B, etc.
• The following information can be used to supplement the color chart for bromcresol green in Activity D.
  {12639_Hints_Table_1}

Teacher Tips

• The activity of hemoglobin, the main oxygen-binding protein in red blood cells, illustrates an application of complex-ion equilibrium. Hemoglobin (Hb) contains four iron(II) ions that bind to oxygen molecules. This must be a reversible reaction, since the hemoglobin must be able to release the oxygen molecules in cells and body tissues (Equation 8).
  {12639_Tips_Equation_8}
  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 8 is shifted in the reverse direction. Less oxygen is therefore available in the bloodstream to be transported to the cells. The physical symptoms of the 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.
• The changes observed when a solution at equilibrium is heated or cooled reveal that the equilibrium “constant” is, in fact, temperature dependent. Thus, a single solution, whether pink or blue, changes color when heated or cooled, respectively, even though no other reagents are added.
• Most textbooks use Le Chatelier’s Principle to predict and explain the effects of both concentration and temperature. Strictly speaking, however, the two effects are different, in that changes in concentration affect the position of equilibrium, while changes in temperature affect the value of the equilibrium constant. At a given temperature, there are an infinite number of possible equilibrium positions, but only a single equilibrium constant value.
• In Activity B, steps 8 and 9 were incorporated into the demonstration based on a student’s question: “Are the Fe³⁺ and SCN^– ions still present even though there is no color?”

Answers to Prelab Questions

In Activity A, the nature of equilibrium will be studied by observing the color changes for an acid–base indicator equilibrium when reactant and product concentrations are changed.

In Activity B, the effect of concentration on equilibrium will be investigated for the formation of complex ions.

In Activity C, the effect temperature changes have on equilibrium will be studied by observing the shifts in equilibrium colors in both an iron thiocyanate and a cobalt chloride complex ion equilibrium when the solutions are put at two different temperatures.

In Activity D, the effect changing pressure has on the position of equilibrium is viewed in the color changes that occur in a seltzer water–acid–base indicator solution when the pressure of the air above the solution is varied.

Sample Data

Activity A. Nature of Equilibrium

Activity B. Effect of Concentration
Activity C. Effect of Temperature
Activity D. Effect of Pressure

Answers to Questions

Activity A. Nature of Equilibrium

1. Write the chemical equation for the reversible reaction of bromcresol green with water. Label this Equation A. Hint: Refer to Equation 3 in the Background section.
   {12639_Answers_Equation_A}
2. Use the color changes observed for the indicator before and after adding HCl (steps 1 and 2) 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 A. 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 A. 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^–.

3. 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.

4. In step 3, hydroxide ions reacted with and removed H⁺ ions from solution (see Equation 3 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.

5. 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.

   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.

6. What form(s) of the indicator were most likely present when the transition color was observed in step 5? 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 at this point 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 must be present at equilibrium, that is, the reaction does not go to completion.

Activity B. Effect of Concentration

1. Use Equations 1 and 2 and Le Chatelier’s Principle to explain the observed color changes.
   1. Step 2

      The addition of Fe³⁺ ions in solution shifts the equilibrium to the right in Equation 1, increasing the concentration of the red FeSCN²⁺ complex ion.

   2. Step 4

      The addition of the reactant ion, SCN^–, also shifts the equilibrium to the right, producing a darker red color.

   3. Step 6

      The H₂PO₄^– ion removes Fe³⁺ ions to product FeH₂PO₄²⁺, according to Equation 2. Removing Fe³⁺ from solution causes Equation 1 to shift to the left, reducing the concentration of red FeSCN²⁺ ions near the dissolving crystal. The solution color around the crystal becomes a lighter shade of orange.

   4. Step 7

      Once the crystal dissolves, the H₂PO₄^– ions remove nearly all of the Fe³⁺ ions in solution to produce FeH₂PO₄²⁺. As Equation 1 shifts to the left, the excess H₂PO₄^– ions react with the newly formed Fe³⁺ ions. This cycle continues until very little Fe³⁺ ions remain in solution. The solution contains mostly the colorless FeH₂PO₄²⁺ ion.

   5. Step 8

      Adding more Fe³⁺ shifts Equation 1 to the right, increasing the concentration of the red complex ion FeSCN²⁺.

   6. Step 9

      Adding more SCN^– also shifts Equation 1 to the right, increasing the concentration of the red complex ion FeSCN²⁺.

Activity C. Effect of Temperature

1. Following is the net ionic equation for the reversible reaction involving cobalt complex ions. The colors of the complex ions are shown underneath their formulas.
   {12639_Procedure_Equation_9}
   Based on the initial color of the cobalt chloride solution (test tube A), what complex ions are present in this solution? Explain.

   The initial color of the cobalt chloride solution is violet. This suggests that the solution contains approximately equal amounts of both the pink and blue complex ions, Co(H₂O)₆²⁺ and CoCl₄^2–, respectively.

2. Which complex ion in net ionic Equation 9 was favored when the solution was heated (step 8)? Which complex ion was favored when the solution was cooled (step 9)? Use these results to determine whether heat should be included on the reactant or product side in the cobalt complex ion Equation 9. Rewrite the equation to include the energy term (heat) directly in the equation.

   Heating the solution changed the color of the solution from pink to blue due to the formation of CoCl₄^2– complex ions. Cooling the solution changed the color of the solution from blue to pink due to the formation of Co(H₂O)₀²⁺ complex ions. Since adding heat shifted the reaction shown in Equation 5 to make more products, it would appear that heat is a “reactant” and that the reaction is endothermic.

   {12639_Answers_Equation_10}
3. Use Le Chatelier’s Principle to explain the color changes that resulted from heating and cooling the solutions in steps 8 and 9, respectively.

   According to Le Chatelier’s Principle, adding heat should shift the reaction in the forward direction to use up some of the “excess” heat.

4. Write the chemical equation for the reversible reaction of iron(III) ions with thiocyanate ions. Use the information in the data table to write the color of each reactant and product underneath its formula.
   {12639_Answers_Equation_11}
5. How did the color of the solution in Part A change when it was cooled (step 14) or heated (step 15)? How do these results demonstrate that the reaction shown in Question 4 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 14), yellow when it was heated (step 15). 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).

Activity D. Effect of Pressure

1. What effect does decreasing the pressure have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 7?

   When the total pressure above the solution decreases, the reaction shifts to the left to reestablish equilibrium and the solubility of carbon dioxide is reduced.

2. What effect does increasing the pressure have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 7?

   Increasing the pressure increases the solubility of the carbon dioxide and the reaction shifts to the right.

3. Explain the results in terms of Le Chatelier’s Principle.

   In both cases, the change in reaction conditions causes the reaction to shift in such a way that the effect of the change will be counteracted. Equilibrium is then reestablished under these new conditions.

References

Special thanks to Patricia Mason (retired) Delphi Community H.S., Delphi, IN, and to Kathy Kitzmann, Mercy H.S., Farmington Hills, MI, for providing Flinn with the general idea and many specific activity suggestions for “activity stations” lab kits.

Introduction

The word equilibrium has two roots: æqui, 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
• Chemical equilibrium
• Equilibrium constant
• Le Chatelier’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, 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 for the reversible reaction of iodine molecules and iodide ions to give triiodide ions (Equation 1).

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 concentrations of reactants and products are equal. The forward and reverse reactions create an equal balance of opposing rates.

What happens when this balance is disturbed? Any factor that changes the rate of the forward or the reverse reaction will change the amounts of reactants and products that are present when equilibrium is reestablished. Reaction conditions that are known to affect the rates of chemical reactions include the concentration of reactants, the temperature and pressure.

If a reactant or product is added to the equilibrium, the system will react to reduce the effect of this addition. When a reactant is added, equilibrium shifts from left to right, consuming the reactant; if a product is added, the shift occurs from right to left. This continues until equilibrium is reestablished.

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.

With reversible reactions involving gases, changes in applied pressure can cause changes in the dynamic equilibrium. If the applied pressure is increased, the reaction balance will shift away from the side with the greater number of moles of gas. Conversely, if the applied pressure is decreased, the balance will shift toward the reaction that produces a greater number of moles of gas.

Activity A. Nature of Equilibrium
The properties of an indicator will be used to study the nature of equilibrium. An indicator is a dye that can gain or lose hydrogen ions to form substances that have different colors. Equation 2 summarizes the reversible reaction of the indicator bromcresol green (HIn). HIn represents an uncharged indicator molecule and In^– an indicator anion formed after the molecule has lost a hydrogen ion.

If the reaction conditions change, we can employ Le Chatlier’s Principle to predict 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.” The color of the indicator in the presence of either excess H⁺ or OH^– ions will show how changing the concentration of a product affects the equilibrium shown in Equation 2. Activity B. Effect of Concentration
An ion in solution can be part of, or common to, separate reversible reactions. In this activity, various solutions and solids will cause color changes when added to a solution of iron(III) nitrate and potassium thiocyanate.

Fe³⁺ and SCN^– ions in solution form the complex ion FeSCN²⁺, which is dark red in solution (Equation 4). The Fe³⁺ ion can also form a complex ion with the dihydrogen phosphate ion, H₂PO₄^– (Equation 5). A colorless solution becomes dark orange upon addition of a solution and then a solid. The dark orange color disappears after the addition of another solid but reappears again when more solution and the original solid are added. Interpret these color changes in terms of changes in concentrations.

Activity C. Effect of Temperature
The effect of temperature on equilibrium will be studied for two reversible reactions, the formation of cobalt complex ions and the formation of the complex ion iron(III) thiocyanate, FeSCN²⁺. 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 CoCl₄^2–, in which the metal ion is surrounded by four chloride ions.

Reaction of iron(III) nitrate with potassium thiocyanate will be used to study the effect of temperature changes on a complex-ion equilibrium. Iron(III) ions react with thiocyanate ions to form FeSCN²⁺ complex ions. Activity D. Effect of Pressure
When carbon dioxide gas dissolves in water, it forms a weakly acidic solution due to the following reversible reaction: The hydrogen ion concentration in solution depends on the amount of dissolved carbon dioxide. In this activity, the effect of pressure on the solubility of carbon dioxide and on the position of equilibrium for this reversible reaction will be studied.

Experiment Overview

The purpose of this activity-stations lab is to investigate various factors that may influence the balance of the equilibrium of reactants and products in reversible chemical reactions. Four mini-lab activities are set up around the classroom. Each activity focuses on the effect a particular reaction condition has on chemical equilibrium.

1. Nature of Equilibrium
2. Concentration Effect
3. Temperature Effect
4. Effect of Pressure

Materials

Prelab Questions

Read the Background material and Procedure for each activity A–D. Write a brief (one- or two-sentence) description of each experiment. Example: In Activity B, the effect of concentration on equilibrium will be investigated for the formation of complex ions.

Safety Precautions

Cobalt chloride solution is moderately toxic by ingestion. Iron(II) nitrate solution may be a skin/tissue irritant. Concentrated hydrochloric acid is highly toxic by ingestion or inhalation and is severely corrosive to skin and eyes; can cause severe body tissue burns. Notify the teacher immediately in case of an acid spill. Dilute hydrochloric acid and sodium hydroxide solutions are skin and eye irritants. Potassium thiocyanate is toxic by ingestion and emits a toxic gas if strongly heated—do not heat this solution and do not add acid. Sodium phosphate monobasic is moderately toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Activity A. Nature of Equilibrium

Acid–Base Equilibrium of Bromcresol Green

1. Obtain 2 mL of distilled water in the test tube. Add 5 drops of 0.04% bromcresol green. Swirl gently and record the color of the solution in the data table.
2. Add 3 drops of 0.1 M HCl solution to the test tube. Swirl gently and record the new color of the solution in the data table.
3. 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 in the data table.
4. Continue adding 0.1 M NaOH dropwise until a total of 5 drops of NaOH have been added in steps 3 and 4 combined.

   Can the process be reversed to obtain a color that is intermediate between that in steps 2 and 3?

5. Add 0.1 M HCl again dropwise very slowly until the solution reaches a “transition” color midway between the two colors recorded above (steps 2 and 3). Swirl gently between drops to avoid overshooting the transition color. Record the number of drops of HCl required and the color in the data table. Note: It may be necessary to add half a drop at a time.
6. Wash the contents of the test tube down the drain with excess water and rinse the test tubes with distilled or deionized water.

Activity B. Effect of Concentration

1. Using a 50-mL graduated cylinder, measure out 20 mL of potassium thiocyanate solution. Transfer the potassium thiocyanate solution to a Petri dish.
2. Add 5 drops of iron(III) nitrate solution into different spots in the Petri dish. Record any color changes in the data table.
3. Swirl the solution until the orange color is uniform throughout.
4. Add ½ pea size amount of the potassium thiocyanate crystals in one spot. Wait about 30 seconds and record any further changes to the solution.
5. Swirl the solution to dissolve the crystals until solution color becomes uniform throughout.
6. Add ¼ pea size amount of the sodium phosphate monobasic crystals in one spot. Wait about 60 seconds and record changes to the solution.
7. Swirl the solution to dissolve the crystals. Record the solution color.
8. Add one drop of the iron(III) nitrate solution in one spot off to the side. Don’t stir. Record any color change.
9. Add a pea size amount of the potassium thiocyanate crystals in a different spot. Wait about 30 seconds and record any changes to the solution around the crystal.
10. Wash the contents of the Petri dish in the wash beaker provided and rinse the Petri dish with distilled water.

Activity C. Effect of Temperature

Preparation

1. Prepare hot-water and ice-water baths: Fill a 250-mL beaker half full with tap water. Place it on a hot plate and heat to 65–70 °C for use in step 9. In a second 250-mL beaker, add water and ice to prepare an ice-water bath for use in step 10.
2. Thoroughly dry a 50-mL beaker with a paper towel, then use the markings on the side of the beaker to obtain about 10 mL of a 1% solution of cobalt chloride in alcohol.
3. Label six clean, dry test tubes A–F and place them in a test tube rack.
4. Using a graduated, Beral-type pipet, add about 2 mL of the cobalt chloride solution to each test tube A–C. Note: The exact volume is not important, but try to keep the volume of solution approximately equal in each test tube.
5. Set test tube A aside as a control. Record the color and appearance of the control solution in the data table. Add 4 drops of distilled water, one drop at a time, to each test tube B and C.
6. Take test tube C to the fume hood. Use the dropper provided on the acid bottle to carefully add 3 drops of 12 M hydrochloric acid to the test tube. Caution: Wear gloves and do not breathe the acid fumes!
7. Gently swirl test tube C to mix the contents, then return the test tube to the test tube rack on your lab bench.
8. Place test tube B 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.
9. Place test tube C 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.
10. Observe and record the initial colors of the Fe(NO₃)₃ and KSCN solutions.
11. 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.
12. Using a graduated, Beral-type pipet, add 1-mL of the FeSCN²⁺ stock solution to each test tube D–F.
13. Add 10 drops of distilled water to test tube D. Gently swirl the test tube to mix the solution and record the color of the solution in the data table. Test tube D will be used as the control solution for comparison purposes in steps 14 and 15.
14. Add 10 drops of distilled water to test tube E 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 of the solution to the control in test tube D. Record the color comparison in the data table.
15. Add 10 drops of distilled water to test tube F 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 in test tube D. Record the color comparison in the data table.
16. Dispose of the contents of the test tubes as directed by your instructor.

Activity D. Effect of Pressure

1. Obtain about 50 mL of fresh seltzer water in a 100-mL beaker and add 4 mL of bromcresol green indicator using a graduated, Beral-type pipet. Swirl to mix the solution.
2. Draw about 25 mL of the seltzer/indicator solution into a 140-mL syringe and seal the syringe by pushing a tip cap firmly on its open end. Record the initial total volume of liquid plus gas in the syringe in the data table.
3. Compare the color of the seltzer/indicator solution with the bromcresol green color chart to determine the pH of the seltzer water. Record the initial color and pH of the solution in the data table.
4. Expand the volume of gas in the syringe: Withdraw the plunger to the 100-mL mark and then insert the nail in the prepared hole so that the syringe plunger will stay at the 100-mL mark. This step will require two people—one person pulls the plunger out past the 100-mL mark and the other person inserts the nail (see Figure 1). Caution: Be very careful when working with the nail—do not push too hard!
   {12639_Procedure_Figure_1}
5. While the plunger is in the withdrawn position, shake the solution until it no longer effervesces and the color no longer changes. Note: According to Boyle’s Law, increasing the applied volume should decrease the pressure of the gas in the syringe.
6. Determine the pH of the solution and record the color, pH and the total volume of liquid plus gas in the syringe.
7. Note the volume of carbon dioxide gas in the syringe. Carefully remove the nail from the syringe. Compress the mixture in the syringe to the 45-mL mark and insert the nail in the prepared holes in both the barrel and plunger of the syringe. This step may also require two people—one person pushes the syringe past the 45-mL mark and the other person inserts the nail (see Figure 2).
   {12639_Procedure_Figure_2}
8. Shake the syringe until both the color and volume no longer change. Note: According to Boyle’s Law, decreasing the applied volume should increase the pressure of the gas in the syringe.
9. Determine the pH of the solution and record both the pH and the total volume of liquid plus gas in the syringe in the data table.
10. Remove the nail, then repeat step 4 by withdrawing the plunger to the 100-mL mark and reinserting the nail. While holding the plunger in the withdrawn position, shake the solution until the color no longer changes.
11. Determine the pH of the solution at the increased volume and record both the pH and the total volume of liquid plus gas in the syringe.

Student Worksheet PDF

12639_Student1.pdf