Materials Included In Kit

Additional Materials Required

Safety Precautions

Dilute solutions (0.1 M) of sodium bicarbonate, hydrochloric acid and sodium hydroxide are body tissue irritants. Wear chemical splash goggles and chemical-resistant gloves when handling these chemicals. Avoid exposure to eyes and skin. 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. All solutions can be disposed of according to Flinn Scientific Disposal Method # 26b.

Teacher Tips

• Enough materials are provided in this kit for 30 students working in pairs, or for 15 groups of students. The experimental work for this lab can reasonably be completed in one 50-minute lab period. The Prelaboratory Assignment should be completed prior to lab, and the Post-Lab Questions can be assigned as a homework assignment or performed as a class activity the day following lab.
• Checker™ pH meters provide an inexpensive and convenient way to measure pH values of solutions directly on a microscale reaction plate. Consider adding the more precise measurement of pH using a pH meter as a valuable extension of the procedures in Parts A and B of this laboratory kit.

Answers to Prelab Questions

1. Briefly describe the hazards associated with 0.1 M HCl and 0.1 M NaOH solutions.

Both HCl and NaOH solutions are body tissue irritants. Avoid exposure to eyes and skin and wear goggles and gloves when working with these chemicals.

2. 1. Using Equation 1 as a guide, write an equation for the reaction of acetic acid (CH₃COOH) with water. Acetic acid is a familiar weak acid that is the primary ingredient in vinegar.
      {11939_PreLab_Equation_1}
   2. Identify the conjugate base of acetic acid in the reaction equation. Hint: The H atom highlighted in boldface is the acidic or “active” hydrogen in acetic acid.

      The conjugate base of acetic acid has the formula CH₃COO^–; it is called acetate ion.

3. Acetic acid and its conjugate base, sodium acetate, form buffer solutions that are effective in the pH range 3.7–5.7.

1. What would be the composition and pH of an ideal buffer prepared from acetic acid and its conjugate base, sodium acetate?

   An ideal buffer contains equal numbers of molecules of both the weak acid and its conjugate base component. The pH of the ideal buffer is the middle value in the pH range of a given buffer system. An ideal acetic acid–sodium acetate buffer solution, therefore, would be prepared from equal volumes of acetic acid and sodium acetate solutions having the same concentrations and have a pH value of 4.7.

2. In resisting a pH change, which buffer component would react with NaOH?

   When NaOH, a strong base, is added to a buffer solution, it reacts with and is neutralized by the weak acid component of the buffer, in this case, acetic acid.

   {11939_PreLab_Equation_2}
3. When the buffer components react with added HCl or NaOH, does the ratio of acetic acid to sodium acetate remain unchanged? What happens to the buffer activity when one of the components is exhausted?

   When strong acid or strong base is added to a buffer solution, it is neutralized by one of the two component “partners” present in the buffer solution. This effect decreases the amount of one of the buffer components and increases the amount of its partner. For example, in the equation shown above for the reaction that occurs when strong base is added to an acetic acid–sodium acetate buffer, the amount of the weak acid component is reduced, while the amount of its conjugate base is increased. Although this changes the ratio of the two components in the buffer solution, as long as both components are still present, the buffer will continue to be active and the pH change will not be dramatic. However, if sufficient strong acid or strong base is added to the buffer to consume completely either one of the components, then the buffer will no longer be effective. Any additional acid or base added to the solution would then cause a large change in pH.

4. People with impaired lung function, due to emphysema, for example, are not able to exchange carbon dioxide efficiently between the lungs and air as they exhale. The result is an increase in the amount of CO₂ dissolved in the blood. How is this likely to affect the buffer balance in the blood? Which term, respiratory acidosis or respiratory alkalosis, would better describe the resulting condition?
   

   An increase in the amount of CO₂ dissolved in the blood would lead, in turn, to an increase in the amount of carbonic acid H₂CO₃ in the blood, due to the reaction of CO₂ with water.

   {11939_PreLab_Equation_3}

   Carbonic acid is the weak acid component of the carbonic acid–bicarbonate buffer that is responsible for maintaining blood pH. An increase in the amount of the carbonic acid component relative to the amount of bicarbonate would tend make the blood buffer more acidic. The correct medical term for this condition is “respiratory acidosis.”

Sample Data

Answers to Questions

1. Compare the measured pH value for the model carbonate blood buffer to

   a. The expected pH of an ideal carbonic acid–bicarbonate buffer

   The model carbonate blood buffer prepared in Part A has a pH value equal to 6.8–7.0. This is greater than the pH of an ideal carbonic acid–bicarbonate buffer (6.4). The pH of the model blood buffer is lower, however, than the actual pH of the buffer system present in blood, which is regulated at pH = 7.2 ±0.2.

   b. The actual pH of the carbonic acid–bicarbonate buffer system present in the blood.

   Student answers will vary.

2. Based on the pH comparisons in Question 1, which solution, the model carbonate blood buffer or an actual blood buffer, is more likely to contain a greater proportion of the carbonic acid component of the buffer compared to the bicarbonate component? Explain your reasoning.
   The pH of the model carbonate blood buffer indicates that it is more acidic than the actual carbonate blood buffer. Therefore, it is more likely that the model solution contains a greater amount of the weak acid component relative to the bicarbonate (conjugate base) component.
3. What are the effects of adding even a small amount of HCl or NaOH on the pH value of the control solution (water)? Compare this to the effect of adding HCl or NaOH to the model carbonate blood buffer.
   The pH of water was dramatically affected by the addition of even one drop of strong acid or strong base. For example, addition of 1 drop of HCl was sufficient to decrease the pH to the “acid” color. In contrast, the buffer solution was approximately 35 times as resistant to pH change upon addition of HCl, since 35–40 drops of HCl were necessary to change the pH of the buffer solution to the acid color. The model carbonate blood buffer was not quite as resistant to the effect of NaOH as it was to the effect of HCl. The buffer capacity with respect to NaOH addition, however, was still 15–20 times greater than that of water.
4. Which phosphate buffer solution in Part B corresponds to the composition of an ideal buffer solution? Compare its measured pH value with the calculated pH of the ideal buffer solution.
   Phosphate buffer B, containing equal amounts of the weak acid component (NaH₂PO₄) and its conjugate base (Na₂HPO₄), has the composition of an ideal buffer. Its measured pH (6.8) is slightly lower than the calculated pH (7.0) of the ideal phosphate buffer.
5. Use the universal indicator color chart to compare the observed pH changes for phosphate buffers A and B and the control solution (water) upon addition of HCl and NaOH. Were phosphate buffers A and B equally effective in resisting pH changes upon addition of either HCl or NaOH?
   Buffers A and B were both more resistant than the water control to pH change. The pH of water dropped from 7 (teal) to < 4 (red) upon addition of 1 drop of HCl. Addition of 1 drop of NaOH to water caused an equally steep pH change in the opposite direction, from pH 7 (teal) to > 11 (purple). The ideal phosphate buffer was able to stay within a narrow pH range from 6.0–8.0 (yellow to dark green) upon addition of 10 drops of either HCl or NaOH. Buffer A was also able to resist change upon addition of NaOH. It was not as effective, however, when HCl was added to the solution.
6. Do the results in Data Tables 2 and 3 confirm that the ideal buffer is optimally effective upon addition of either HCl or NaOH?
   Yes, the ideal buffer was equally resistant to pH change upon addition of either HCl or NaOH. It stayed within the expected phosphate buffer range (pH 6.0–8.0) at all times.

Introduction

One of the most interesting biological applications of acids and bases is that of buffers. A buffer protects against rapid changes in pH when acids or bases are added to it. Every living cell contains natural buffer systems to maintain the constant pH needed for proper cell function. Consumer products are also often buffered to safeguard their activity. What are buffers made of? How do buffers maintain the delicate pH balance needed for life and health?

Concepts

• pH
• Buffer
• Weak acid vs. conjugate base
• Neutralization

Background

Almost all chemical reactions in living organisms take place at pH values between 6 and 8. These chemical reactions are very sensitive to the presence of hydrogen ions (H⁺ or H₃O⁺). Even a small drop in pH due to an increase in hydrogen ion concentration can cause some of nature’s catalysts (the enzymes) to stop functioning. The pH level in blood, for example, must be maintained within extremely narrow limits—a pH change of as little as ±0.1–0.2 units can result in serious illness or even death. Constant pH ensures that the essential molecules of life—proteins, carbohydrates, lipids, and nucleic acids—are in their correct ionic states andwill function properly.

How Does a Buffer Work?

The ability of buffers to resist changes in pH upon addition of acid or base can be traced to their chemical composition. All buffers contain a mixture of both a weak acid (HA) and its conjugate base (A^–). The buffer components HA and A^– are related to each other by means of the following chemical reaction that describes the behavior of a weak acid in water (Equation 1).

Buffers control pH because the two buffer components are able to react with and therefore neutralize either excess strong acid or excess strong base that are added to the solution. The weak acid component HA reacts with any strong base, such as sodium hydroxide (NaOH), added to the solution to yield water and the conjugate base component A^– (Equation 2). The conjugate base component A^– reacts with any acid, such as hydrochloric acid (HCl), added to the solution to yield its acid partner HA and chloride ion (Equation 3).

These complementary neutralization reactions can be visualized as a cyclic process (see Figure 1). Buffer activity will continue as long as both components remain present in solution. Once either component, A^– or HA, is completely consumed by Reaction 2 or 3 above, the buffer capacity will be exhausted and the buffer will no longer be effective.

Ideal Buffers and Buffer Range

A buffer composed of an equal number of molecules of a weak acid and its conjugate base is called an ideal buffer because it is equally effective in resisting pH changes upon addition of either acid or base. The pH range in which a buffer system will be effective is called its buffer range. Since a buffer solution must always contain noticeable amounts of both a weak acid and its conjugate base, the buffer range is usually limited to 2 pH units centered around the pH of the ideal buffer solution. An ideal carbonic acid–bicarbonate buffer, for example, has a pH of 6.4 and the buffer range for this system is pH 5.4–7.4.

Properties of Biological Buffers

The body is able maintain proper pH due to the presence of chemical buffer systems in cells and in the blood. The major buffer present in blood, for example, is composed of the weak acid, carbonic acid (H₂CO₃), and its conjugate base, bicarbonate ion (HCO₃^–) (Equation 4). The normal pH of blood (7.4) is at the upper limit of the effective range for the carbonic acid–bicarbonate buffer system. The buffer activity is kept in balance, however, by a reserve supply of gaseous CO₂ in the lungs, which can replenish H₂CO₃ in the blood by dissolving and reacting with water in the blood (Equation 5). The second most important biological buffer system involves dihydrogen phosphate (H₂PO₄^–) as the weak acid and its conjugate base hydrogen phosphate (HPO₄^2–) (Equation 6). An ideal phosphate buffer having the above composition has a pH range of 6.8–7.2, an optimum value for physiological pH! This is the most prominent buffer within cells. It is also the buffer of choice when studying proteins and enzymes, because it perfectly simulates physiological conditions.

Overview of the Biological Buffer Experiments
In Part A, carbonated seltzer water is used as a source of carbonic acid to prepare a model biological carbonic acid–bicarbonate buffer that is effective at neutral pH (pH = 7). The effects of added acid and base on the pH and buffer capacity of this model biological buffer will be examined. The pH value of the buffer solution will be estimated using bromthymol blue, an acid–base indicator that changes color in the pH range 6.0 to 7.6. Bromthymol blue is yellow when the pH is less than 6.0, blue when the pH is greater than 7.6, and green (the intermediate color) in the transition range 6.0–7.6.

In Part B, two different phosphate buffers that reflect the physiological role of buffers within cells will be prepared. The pH of the buffer solutions and the pH range over which phosphate buffers are effective will be measured. The pH changes will be followed using universal indicator solution, an acid–base indicator system that can be used over the pH 4–10 range. Consult the universal indicator color chart to determine the pH value corresponding to a given color of the solution.

Materials

Prelab Questions

Read the background information and answer the following questions on a separate sheet of paper.

1. Briefly describe the hazards associated with 0.1 M HCl and 0.1 M NaOH solutions.
2. Anwer the following questions
   1. Using Equation 1 as a guide, write an equation for the reaction of acetic acid (CH₃COOH) with water. Acetic acid is a familiar weak acid that is the primary ingredient in vinegar.
   2. Identify the conjugate base of acetic acid in the reaction equation. Hint: The H atom highlighted in boldface is the acidic or “active” hydrogen in acetic acid.
3. Acetic acid and its conjugate base, sodium acetate, form buffer solutions that are effective in the pH range 3.7–5.7.
   1. What would be the composition and pH of an ideal buffer prepared from acetic acid and its conjugate base, sodium acetate?
   2. In resisting a pH change, which buffer component would react with NaOH?
   3. When the buffer components react with added HCl or NaOH, does the ratio of acetic acid to sodium acetate remain unchanged? What happens to the buffer activity when one of the components is exhausted?
4. People with impaired lung function, due to emphysema, for example, are not able to exchange carbon dioxide efficiently between the lungs and air as they exhale. The result is an increase in the amount of CO₂ dissolved in the blood. How is this likely to affect the buffer balance in the blood? Which term, respiratory acidosis or respiratory alkalosis, would better describe the resulting condition?

Safety Precautions

Dilute (0.1 M) solutions of sodium bicarbonate, hydrochloric acid, and sodium hydroxide are body tissue irritants. Wear chemical splash goggles and chemical-resistant gloves when handling these chemicals. Avoid exposure to eyes and skin. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part A. Model Carbonate Blood Buffer

1. Set up six medium-size test tubes in a rack. Label them 1–6.
2. With Table 1 as a guide, use a graduated cylinder to measure and add the indicated volumes of the required solutions to each test tube.
3. Mix the contents of each test tube thoroughly by gentle shaking or swirling.
   {11939_Procedure_Table_1}
4. Add 5 drops of bromthymol blue indicator solution to each test tube 1–6. Shake to mix. Record the initial color for each solution 1–6 in Data Table 1.
5. Measure the initial pH of each solution using a piece of narrow-range pH paper, pH 6.0–8.0. Record the results in Data Table 1.
6. Use a Beral-type pipet to add 0.1 M hydrochloric acid solution dropwise to the biological buffer solution in test tube 2. Be sure to swirl the contents periodically to ensure thorough mixing.
7. Count the number of drops of HCl required to change the color to the same shade of yellow as the carbonic acid reference solution in test tube 1. Record the number of drops in Data Table 1.
8. Use a Beral-type pipet to add 0.1 M hydrochloric acid solution dropwise to the water control solution in test tube 3. Count the number of drops of HCl required to change the color to the same shade of yellow as the carbonic acid reference solution in test tube 1. Record the number of drops in Data Table 1.
9. Use a Beral-type pipet to add 0.1 M sodium hydroxide solution dropwise to the biological buffer solution in test tube 4. Be sure to swirl the contents periodically to ensure thorough mixing.
10. Count the number of drops of NaOH required to change the color to the same shade of blue as the sodium bicarbonate reference solution in test tube 6. Record the number of drops in Data Table 1.
11. Use a Beral-type pipet to add 0.1 M sodium hydroxide solution dropwise to the water control solution in test tube 5. Count the number of drops of NaOH required to change the color to the same shade of blue as the sodium bicarbonate reference solution in test tube 6. Record the number of drops in Data Table 1.

Part B. Biological Phosphate Buffers

12. Obtain 2 large test tubes or small beakers and label them A and B.
13. Use clean graduated cylinders to add 12 mL of NaH₂PO₄ solution and 3 mL of Na₂HPO₄ solution to test tube A. (This is buffer A.)
14. Use clean graduated cylinders to add 8 mL of NaH₂PO₄ solution and 8 mL of Na₂HPO₄ solution to test tube B. (This is buffer B.)
15. Mix the contents of each test tube by gentle shaking or swirling.
16. Use the following layout plan to fill each indicated well in a microscale 24-well reaction plate with 1.5 mL of distilled water (the control), phosphate buffer A, or phosphate buffer B, respectively.
    {11939_Procedure_Figure_1}
17. Add 3 drops of universal indicator solution to each filled well.
18. Record the initial indicator colors for the water control (Well 1A), buffer solution A (Well 2A) and buffer solution B (Well 3A) in Data Table 2.
19. Estimate the initial pH of each of these samples 1–3 in Row A using a small (1 ) piece of narrow-range (6.0–8.0) pH paper.
20. Using a clean Beral-type pipet, add 1 drop of HCl to Wells 1, 2, and 3 in Row B. Record the indicator colors in Data Table 2.
21. Add 5 drops of HCl to Wells 2C and 3C and record the new indicator colors in Data Table 2.
22. Add 10 drops of HCl to Wells 2D and 3D and again record the indicator colors in Data Table 2.
23. Using a clean Beral-type pipet, add 1 drop of NaOH to Wells 4B, 5B, and 6B.
24. Record the indicator colors in Data Table 3.
25. Add 5 drops of NaOH to Wells 4C and 5C and record the indicator colors in Data Table 3.
26. Add 10 drops of NaOH to Wells 4D and 5D and again record the indicator colors in Data Table 3.
27. Wash the contents of the reaction well plate down the drain under running water.

Student Worksheet PDF

11939_Student1.pdf