Equilibrium in a Syringe

Introduction

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 demonstration, the effect of pressure and temperature on the solubility of carbon dioxide and on the position of equilibrium for this reversible reaction will be studied.

Concepts

• Equilibrium
• Gas solubility
• Le Chatelier’s principle
• Acid–base indicator

Materials

Safety Precautions

Wear chemical splash goggles and chemical-resistant gloves and apron. Wear heat-resistant gloves when working with the boiling water bath. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Prelab Preparation

1. Prepare a boiling water bath and an ice-water bath for use in Part B. To prepare a boiling water bath, fill a 1000-mL beaker about ⅔-full with water, add a boiling stone, and heat at medium-high setting on a hot plate. To prepare an ice-water bath, fill a second 1000-mL beaker about ¾-full with a crushed ice/water mixture.
2. An optional student worksheet for testing student understanding of the demonstration is included. If using these worksheets, pass out copies of the worksheet to the students before beginning the demonstration.

Procedure

Part A. 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. Have students record the initial total volume of liquid plus gas in the syringe in the data table. (The initial volume should be about 50 mL.)
3. Compare the color of the seltzer/indicator solution with the bromcresol green color chart to determine the pH of the seltzer water. Have students record the initial color and pH of the solution in the data table. (The seltzer/indicator solution is yellow-green, corresponding to a pH of 4.0.)
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 may require two people—one person pulls the plunger out past the 100-mL mark and the other person inserts the nail (see Figure 1).
   {13973_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. Have students determine the pH of the solution and record the color, pH, and the total volume of liquid plus gas in the syringe. (The new pH is 4.4 at an expanded volume of 100 mL.)
7. Recall that pH and [H⁺] are inversely related—the higher the pH, the lower the hydrogen ion concentration. Write Equation 1 on the board or overhead projector. What effect does decreasing the pressure have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 1? (Decreasing the pressure shifts the equilibrium shown in Equation 1 to the left, reducing the solubility of carbon dioxide gas and decreasing the hydrogen ion concentration.)
8. 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).
   {13973_Procedure_Figure_2}
9. 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.
10. Determine the pH of the solution and record both the pH and the total volume of liquid plus gas in the syringe. (The new pH is 4.0 at a compressed volume of 45 mL.)
11. What effect does increasing the pressure have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 1? (Increasing the pressure shifts the equilibrium shown in Equation 1 to the right, increasing the solubility of carbon dioxide gas and increasing the hydrogen ion concentration.)
12. 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.
13. 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. (Note that the process can be repeated several times. Equilibrium will reestablish itself to give a pH of 4.4 at an expanded volume of 100 mL.)
14. Discuss the results in terms of Le Chatelier’s Principle.

Part B. Effect of Temperature

15. Dispose of the solution in the syringe and draw about 25 mL of fresh seltzer/indicator solution into the syringe. If the initial color of the solution has changed noticeably from Part A, prepare a new seltzer/indicator solution.
16. Holding the syringe with its open end upwards, draw carbon dioxide gas into the syringe from its source until the total volume of material in the syringe is about 70 mL. Seal the syringe by pushing the syringe tipcap firmly on its open end.
17. 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 pH of the solution, the total volume of all material in the syringe, and the temperature of the solution in the beaker. (The initial pH is 4.2 at a total volume of 75 mL and an initial temperature of 21 °C.)
18. Wearing heat-resistant gloves, remove the beaker of boiling water from the hot plate and place the syringe into the heated water.
19. Holding the syringe by the plunger, stir the hot water bath while shaking the mixture in the syringe up and down. Tap on the plunger a few times while shaking the solution to overcome the friction between the plunger and the sides of the syringe.
20. After several minutes, record the temperature of the water, the total volume of the mixture in the syringe, and the pH of the solution. (At a temperature of 100 °C, the volume increases to 139 mL and the pH increases to 4.6.)
21. What effect does increasing the temperature have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 1? (Increasing the temperature shifts the equilibrium shown in Equation 1 to the left, reducing the solubility of carbon dioxide gas and decreasing the hydrogen ion concentration.)
22. Remove the syringe from the hot water bath and place it into the ice-water bath. Holding the syringe by the plunger, stir the ice–water bath while shaking the solution in the syringe up and down. Tap on the plunger a few times while shaking the solution to overcome the friction between the plunger and the sides of the syringe.
23. After several minutes, record the temperature of the water, the total volume of the mixture in the syringe, and the pH of the solution. (At a temperature of 1.5 °C, the volume decreases to 70.5 mL and the pH decreases to 4.2.)
24. What effect does decreasing the temperature have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 1? (Decreasing the temperature shifts the equilibrium shown in Equation 1 to the right, increasing the solubility of carbon dioxide gas and increasing the hydrogen ion concentration.)
25. Discuss the results in terms of Le Chatelier’s Principle.

Student Worksheet PDF

13973_Student1.pdf

Teacher Tips

• This kit contains enough chemicals, with the exception of seltzer water, to perform the demonstration at least seven times.

• 250 mL of seltzer water is included in kit. For subsequent demonstrations, additional fresh seltzer water will be required. Any unflavored, unbuffered seltzer will work as a source of dissolved carbon dioxide. Do not use soda water or club soda, which contains sodium bicarbonate. Club soda is essentially a buffered solution—its pH will not change.
• The following information can be used to supplement the color chart for bromcresol green:
  {13973_Tips_Table_1}

• This demonstration requires about 40 mL of carbon dioxide gas. There are many common methods for generating and dispensing the gas. See the Supplementary Information in the Further Extensions section for a sample procedure.
• Steps 5 and 9 may take several minutes for equilibrium to be reestablished.

Further Extensions

Supplementary Information: Carbon Dioxide Generator

Carbon dioxide gas may be generated by the reaction of marble chips with 3 M hydrochloric acid in an Erlenmeyer flask equipped with a gas delivery tube. A lecture bottle of carbon dioxide can also be used as a generator. Use a 1-quart plastic freezer bag as a gas delivery apparatus, as described in Figure 3.

To construct the bag delivery apparatus, use a large, sharpened cork borer to cut a large hole in a 10, one-hole rubber stopper. The result is a stopper with a plug that can be removed (the plug has the original single hole in it). Remove the plug from the stopper and push the plastic bag through the hole in the stopper, leaving about 1 inch of the bag sticking out of the opening. Place the smaller, one-hole stopper “plug” into the freezer bag opening. The freezer bag should now be held tightly between the walls of the two stoppers. Carefully insert the tapered end of a medicine dropper through the hole in the stopper plug. Attach a short piece of latex tubing over the wide end of the medicine dropper and place a pinch clamp over the latex tubing.

To fill the bag with carbon dioxide, first evacuate the plastic bag by attaching the latex tubing to an aspirator. When the bag has been evacuated, replace the pinch clamp on the tubing and then attach the end of the tubing to the gas delivery tube on the carbon dioxide generator. Remove the pinch clamp and slowly fill the bag assembly with carbon dioxide. The bag should be taut when filled, but not ready to burst. Remove the tubing from the gas generator and replace the pinch clamp. The bag contains a slightly pressurized sample of carbon dioxide gas.

Sample Data

Part A. Effect of Pressure

Part B. Effect of Temperature

Answers to Questions

Part A. 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 1?

   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 1?

   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 so that the effect of the change will be counteracted. Equilibrium is then reestablished under these new conditions.

Part B. Effect of Temperature 

4. What effect does increasing the temperature have on the solubility of carbon dioxide gas and on the position of equilibrium for Equation 1?

   Increasing the temperature shifts the equilibrium shown in Equation 1 to the left, reducing the solubility of carbon dioxide gas and decreasing the hydrogen ion concentration.

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

   Decreasing the temperature shifts the equilibrium shown in Equation 1 to the right, increasing the solubility of carbon dioxide and increasing the hydrogen ion concentration.

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

   When the temperature increases, the volume of the gas phase increases. At increased volume, the concentration of carbon dioxide in the gas phase is reduced, causing the reaction to shift to the left to counteract this change. At lower temperatures, the volume decreases, the concentration of carbon dioxide in the gas phase increases, and the reaction shifts to the right to counteract the change.

Discussion

This activity demonstrates the effect of pressure and temperature on three reversible reactions: the solubility of carbon dioxide in water, the reaction of aqueous carbon dioxide and water to form H₂CO₃, and the weak acid ionization of H₂CO₃ to give HCO₃^– and H⁺ ions. For simplicity sake, in terms of classroom discussion, these reactions are combined in Equation 1. The position of equilibrium for this overall reaction can be determined by measuring the concentration of H⁺ ions in solution. Seltzer water is used as a source of dissolved carbon dioxide, and the concentration of H⁺ ions is estimated using bromcresol green as an indicator. The indicator is yellow when the pH is less than 3.8, blue when the pH is greater than 5.2 and various shades of green in the pH range 3.8–5.2. A sealed syringe is used to provide a closed system.

The effects of pressure and temperature on the solubility of carbon dioxide gas can be explained in terms of Le Chatelier’s Principle:

“If a system at equilibrium is disturbed by a change in temperature, pressure or the concentration of one of its components, the system will tend to shift its equilibrium position so as to counteract the effect of this disturbance.”

In Part A, the total pressure above the seltzer/indicator solution is reduced (by increasing the applied volume), the indicator changes from yellow-green to green, corresponding to a pH change from 4.0 to 4.4. Since a pH increase corresponds to a decrease in the hydrogen ion concentration, the results indicate that the equilibrium shown in Equation 1 is shifted to the left as the pressure decreases—the solubility of carbon dioxide decreases. This is in agreement with Le Chatelier’s Principle. Carbon dioxide gas bubbles out of solution, back into the gas phase. The reverse effect is observed when the pressure is increased. The relationship between pressure and the amount of dissolved carbon dioxide is an example of Henry’s Law, which states that the amount of a gas dissolved in solution is proportional to the pressure of the gas above the solution. In Part B, the effect of changing the temperature on the position of equilibrium for Equation 1 can also be explained in terms of a pressure effect. Increasing the temperature of a gas will increase both the volume and the temperature. In general, the solubility of a gas will be greatest at higher pressures and lower temperatures.

References

This demonstration has been adapted from Flinn ChemTopic™ Labs, Volume 15, Equilibrium, Cesa, I., Ed., Flinn Scientific, Batavia, IL, 2003.