Materials Included In Kit

Additional Materials Required

Safety Precautions

Dilute solutions of sodium bicarbonate, hydrochloric acid and sodium hydroxide are body tissue irritants. Avoid exposure to eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Consult 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

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 remaining in well plates and test tubes after testing may be rinsed down the drain with water according to Flinn Scientific Disposal Method #26b. Hydrochloric acid and sodium hydroxide solutions must be neutralized prior to disposal.

Lab Hints

• The experimental work for this lab can reasonably be completed within a typical 2-hour lab period.
• Flinn pH meters (Catalog No. AP8673) 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 to the procedures for studying the properties of the model biological buffers.
• Sodium hydrogen phosphate salts are also known commercially by their common names. Sodium dihydrogen phosphate, NaH₂PO₄, is called sodium phosphate, monobasic, while sodium hydrogen phosphate, Na₂HPO₄, is referred to as sodium phosphate, dibasic. Na₃PO₄ is called sodium phosphate, tribasic.
• Equilibrium constants and pK_a values are temperature dependent. The pH of an ideal, equimolar buffer will change by up to 0.1 pH unit per degree Celsius. Reference or literature pK_a values are usually reported for 25 °C.

Answers to Prelab Questions

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations, and state whether the solutions are acidic, basic or neutral. Which solution is most acidic?
   {14042_PreLabAnswers_Table_3}
   Gastric juice is the most acidic solution in this list.
2. 1. Using Equation 1 as a guide, write an equation for the reaction of acetic acid (CH₃COOH) with water.
      {14042_PreLabAnswers_Reaction_1}
   2. Identify the conjugate base of acetic acid in the reaction equation.

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

3. Acetic acid and a salt containing its conjugate base, such as 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. Therefore, an ideal acetic acid–sodium acetate buffer solution would 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.

   3. What happens to the buffer activity when this component is exhausted?

      If sufficient strong base is added to the buffer to completely consume the weak acid component, 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.

Sample Data

Laboratory Report

Model Carbonate Blood Buffer

Effect of HCl on Biological Phosphate Buffers

*From color chart

Data Table C. Effect of NaOH on Biological Phosphate Buffers

*From color chart

Answers to Questions

1. Compare the measured pH value for the model carbonate blood buffer to the expected pH of an ideal carbonic acid–bicarbonate buffer and the actual pH of blood.

   The model carbonate blood buffer 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 blood, which is regulated at pH = 7.2 ±0.2.

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 compared to the bicarbonate component? Explain.

   The pH of the model carbonate blood buffer indicates that it is more acidic than the actual buffer present in blood. Therefore, it is more likely that the model buffer contains a greater amount of the weak acid component relative to the bicarbonate (conjugate base) component.

3. What is the effect of adding even one drop of HCl or NaOH on the pH value of water? Compare this to the pH changes observed when HCl or NaOH was added 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 one drop of HCl was sufficient to decrease the pH to the “acid” color (pH < 6). In contrast, the buffer solution was approximately 35 times more resistant to pH change, 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 20–25 times greater than that of water.

4. Which phosphate buffer corresponds to the composition of an ideal buffer solution? Compare its measured pH value with the calculated pH of the ideal buffer.

   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) appears to be slightly lower than the calculated pH (7.0) of the ideal phosphate buffer. This difference is probably not significant and may be attributed to temperature and the relative accuracy of pH paper.

5. Use the universal indicator color chart to compare the observed pH changes for phosphate buffers A and B and the control (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 one drop of HCl. Addition of one 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 to 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.

Introduction

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 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
• Conjugate base
• Weak acid

Background

Many chemical reactions in living organisms take place at neutral pH values. Even a small change in pH 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.

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 either a weak acid (HA) and its conjugate base (A^–), or a weak base and its conjugate acid. 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 buffer components HA and A^– are able to neutralize either strong acid or base that might be added to the solution. The weak acid component HA reacts with any base, such as sodium hydroxide (NaOH), to give water and the conjugate base component A^– (Equation 2). The conjugate base component A^– reacts with any acid, such as hydrochloric acid (HCl), to regenerate its acid partner HA and chloride ion (Equation 3). These neutralization reactions can be visualized as a cyclic process (see Figure 1). Buffer activity will continue as long as both components are present in solution. Once either component is consumed, the buffer capacity will be exhausted and the buffer will no longer be effective. A buffer composed of an equal number of moles of a weak acid and its conjugate base is sometimes 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 will be effective is called its buffer range. 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 its buffer range is pH 5.4–7.4.

Properties of Biological Buffers
The body is able to 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 (7.2) 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 involves dihydrogen phosphate ion (H₂PO₄^–) as the weak acid and its conjugate base hydrogen phosphate ion (HPO₄^2–) (Equation 6). An ideal buffer composed of equal moles of H₂PO₄^– and HPO₄^2– has a pH range of 6.8–7.2. This is an optimum value for physiological pH. Phosphate buffers are the most prominent buffers within cells.

Experiment Overview

This experiment is divided into two parts to study a model carbonate blood buffer and sample phosphate cell buffers, respectively.

Seltzer water will be used as a source of carbonic acid to prepare a model biological carbonic acid–bicarbonate buffer. The effects of added acid and base on the pH and buffer capacity of this model 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 the second activity, two different phosphate buffers will be prepared. These buffers reflect the physiological role of buffers within cells. The pH of the buffers and the pH range over which they are effective will be measured. The pH changes will be followed using universal indicator, an acid–base indicator 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

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations, and state whether the solutions are acidic, basic or neutral. Which solution is most acidic?
   {14042_PreLab_Table_1}
2. 1. Using Equation 1 as a guide, write an equation for the reaction of acetic acid (CH₃COOH) with water.
   2. Identify the conjugate base of acetic acid in the reaction equation.
3. Acetic acid and a salt containing its conjugate base, such as 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. What happens to the buffer activity when this component is exhausted?

Safety Precautions

Dilute solutions of sodium bicarbonate, hydrochloric acid and sodium hydroxide are body tissue irritants. Avoid exposure to eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron.

Procedure

Model Carbonate Blood Buffer

1. Set up six medium-size test tubes in a rack. Label them 1–6.
2. With the following table as a guide, use a graduated cylinder to measure and add the indicated volumes of the required solutions to each test tube.
3. Thoroughly mix the contents of each test tube by gentle shaking or swirling.
   {14042_Procedure_Table_2}
4. Add 5 drops of bromthymol blue to each test tube 1–6. Shake to mix and record the initial color of each solution.
5. Measure the initial pH of each solution using a piece of narrow-range pH paper, pH 6.0–8.0 and record the results.
6. Using a Beral-type pipet, add 0.1 M hydrochloric acid solution dropwise to the model carbonate blood buffer in test tube 2. Be sure to swirl the contents after each drop to ensure thorough mixing.
7. Count the number of drops of HCl required to change the color to the same shade as the carbonic acid reference solution in test tube 1. Record the result.
8. Using a Beral-type pipet, add 0.1 M hydrochloric acid solution dropwise to the water (control) in test tube 3. Count the number of drops of HCl required to change the color to the same shade as the carbonic acid reference solution in test tube 1. Record the result.
9. Using a different Beral-type pipet, add 0.1 M sodium hydroxide solution dropwise to the model carbonate blood buffer in test tube 4. Be sure to swirl the contents after each drop to ensure thorough mixing.
10. Count the number of drops of NaOH required to change the color to the same shade as the sodium bicarbonate reference solution in test tube 6. Record the result.
11. Using a Beral-type pipet, add 0.1 M sodium hydroxide solution dropwise to the water (control) in test tube 5. Count the number of drops of NaOH required to change the color to the same shade as the sodium bicarbonate reference solution in test tube 6. Record the result.

Biological Phosphate Buffers

12. Obtain two 50-mL 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 beaker A. (This is Buffer A.)
14. Use clean graduated cylinders to add 8 mL of NaH₂PO₄ solution and 8 mL of Na₂HPO₄ solution to beaker B. (This is Buffer B.)
15. Stir each buffer solution with a stirring rod to ensure thorough mixing.
16. Use the following layout plan to fill each indicated well in a 24-well reaction plate with 1.5 mL of distilled water (the control), Buffer A or Buffer B, respectively.
    {14042_Procedure_Figure_2}
17. Add 3 drops of universal indicator to each filled well.
18. Record the initial indicator colors for the water control (well A1), Buffer A (well A2) and Buffer B (well A3).
19. Estimate the initial pH of wells A1, A2 and A3 using narrow-range (6.0–8.0) pH paper.

Effect of HCl Addition

20. Using a clean Beral-type pipet, add 1 drop of HCl to each well B1, B2 and B3. Record the indicator colors.
21. Add 5 drops of HCl to each well C2 and C3 and record any indicator color changes.
22. Add 10 drops of HCl to each well D2 and D3 and again record any indicator color changes.

Effect of NaOH Addition

23. Using a clean Beral-type pipet, add 1 drop of NaOH to each well B4, B5 and B6. Record the indicator colors.
24. Add 5 drops of NaOH to each well C4 and C5 and record any indicator color changes.
25. Add 10 drops of NaOH to each well D4 and D5 and again record any indicator color changes.
26. Rinse the contents of the reaction well plate down the drain under running water.

Student Worksheet PDF

14042_Student1.pdf