Materials Included In Kit

Additional Materials Required

Safety Precautions

Dilute solutions (0.1 M) of acetic acid, hydrochloric acid and sodium hydroxide are body tissue irritants. Avoid exposure to eyes and skin. 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.

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 Procedure # 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 class period. The prelaboratory assignment should be completed prior to lab, and the data analysis, calculations and conclusions sections can be assigned as a post-lab activity or performed as a class activity the day following the lab.
• Point out to the students that the control solution—distilled water (solution 6)—is the appropriate comparison for interpreting the behavior of the buffer solutions when either HCl or NaOH is added. The control changes color immediately to the acid or base reference color (pH) when even 1 drop of HCl or NaOH, respectively, is added. In contrast, the indicator colors (pH) of the buffer solutions do not change to the reference color (pH) until 5–20 drops of HCl or NaOH have been added. The reference solutions are provided to show the limiting indicator colors and pH values for acetic acid (solution 1) and sodium acetate (solution 5).
• The students may misinterpret the observation that the indicator color for the acetic acid reference solution (Part 2, effect of acid) or the sodium acetate reference solution (Part 3, effect of base) does not change. This may confuse the students into thinking that it is the reference solutions that resist changes in pH upon addition of HCl or NaOH. To prevent confusion, have the students test the pH of solutions 1A –1D in Part 2 and solutions 5A– 5D in Part 3 with wide-range pH paper (optional). Students will observe that the pH changes that are occurring here are hidden by the pH limits of the indicators.
• Students who are visual learners will immediately notice that the indicator colors on the reaction plates in Parts 2 and 3 show striking diagonal patterns. In Part 2, highlight the fact that they see a gradual color change for buffers A, B, and C (solutions 2–4) from blue-green to yellow. The amount of HCl required to bring about this color change increases in a regular manner from solutions 2 through 4. In Part 3, the observed diagonal pattern points in the opposite direction. The amount of NaOH required to bring about the gradual color change from purple-brown to orange increases in reverse order from solutions 4 to 2. Ask students to hypothesize about the significance of these visual patterns.
• Checker™ pH meters provide an inexpensive and convenient way to measure pH values of solutions directly on a microscale reaction plate! Consider adding the measurement of pH using a pH meter as a valuable extension of the procedure in Part 1 of this laboratory kit.

Answers to Prelab Questions

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations.
   {13372_PreLabAnswers_Table_4}
2. Write the reaction equations for the dissociation of the following weak acids and identify their conjugate bases: phosphoric acid (H₃PO₄), formic acid (HCO₂H) and boric acid (H₃BO₃).
   {13372_PreLabAnswers_Reaction_1}
3. What would be the composition and pH of an ideal buffer prepared from lactic acid (CH₃CHOHCO₂H), where the hydrogen atom highlighted in boldface is the acidic hydrogen atom? The K_a value for lactic acid is 1.38 x 10^–4.

   An ideal buffer would contain an equal number of moles of the weak acid component (lactic acid) and its conjugate base (sodium lactate, NaCH₃CHOHCO₂).
   The [H₃O⁺] concentration in an ideal buffer is equal to the K_a value:
   [H₃O⁺] = 1.38 x 10^–4; pH = 3.86.

4. Use the buffer equation to calculate the pH of buffer solutions prepared by dissolving the following amounts of acetic acid and sodium acetate, respectively, in enough water to make 1 L of solution:
   1. 0.67 moles acetic acid and 0.33 moles of sodium acetate
   2. 0.33 moles acetic acid and 0.67 moles of sodium acetate.
      {13372_PreLabAnswers_Equation_10}

Sample Data

Data Table 1. Initial pH and Indicator Colors

Data Table 2. Effect of HCl Addition Data Table 3. Effect of NaOH Addition

Answers to Questions

1. Use the buffer equation to calculate the expected [H₃O⁺] and pH values for buffer solutions A, B and C. Compare the calculated values against the experimental data (Data Table 1) for Buffers A, B and C. (Hint: The concentrations of the acetic acid and sodium acetate solutions are the same, and the total volume of the buffer solutions is constant. Therefore, the volume of each component used to prepare the buffer can be substituted directly into the concentration ratio expression in Equation 5).
   {13372_Answers_Equation_4}
   {13372_Answers_Table_8}
2. Summarize from Data Table 2 the number of drops of HCl required to change the color of buffer solutions A, B, C and the distilled water control to yellow (the acid reference color)?
   {13372_Answers_Table_9}
3. Did all of the buffer solutions exhibit buffering activity with respect to added HCl in Part 2? Explain. Which buffer solution was most effective with respect to added acid? Explain.

   Yes, all of the buffers exhibited some buffering activity with respect to added HCl. Buffers A, B and C were more effective than the control solution in accepting some HCl before the pH dropped to that of the acid reference solution. Buffer C was most effective in Part 2; it was able to neutralize 20 drops (approximately equal to its total volume) of 0.1 M HCl before the indicator color reverted to yellow.

4. The behavior of solution 5 in Data Table 2 demonstrates that a buffer can be made by partial neutralization of the basic component A^– with HCl. (a) Write an equation for the neutralization reaction of sodium acetate in solution 5 upon addition of HCl. (b) Describe and explain the significance of the color changes observed for solution 5 in Data Table 2.
   1. {13372_Answers_Equation_11}
   2. The sodium acetate reference solution (0.1 M) started out at pH 7.5. The bromcresol green indicator color for this pH was blue. As HCl was added in small increments to the solution it underwent a gradual color change and pH drop. The gradual color change reflects the fact that as a portion (but not all) of the sodium acetate was neutralized to give some acetic acid, the composition and pH of the solution dropped down to the buffer range. The solution then acted as an excellent buffer, undergoing a very gradual color transition from bluegreen to green-blue upon addition of excess HCl.
5. Summarize from Data Table 3 the number of drops of NaOH required to change the color of buffer solutions A, B, C and the distilled water control to orange (the base reference color)?
   {13372_Answers_Table_10}
6. Did all of the buffer solutions exhibit buffering activity with respect to added NaOH in Part 3? Explain. Which buffer solution was most effective with respect to added base in Part 3? Explain.

   Yes, all of the buffers exhibited some buffering activity with respect to added NaOH. Buffers A, B and C were more effective than the control solution in accepting some NaOH before the pH increased to that of the base reference solution. Buffer A was most effective in Part 3; it was able to neutralize more than 20 drops (approximately equal to its total volume) of 0.1 M NaOH before the indicator color reverted to orange.

7. The behavior of solution 1 in Data Table 3 demonstrates that a buffer can be made by partial neutralization of the acidic component HA with NaOH. (a) Write an equation for the neutralization reaction of acetic acid in solution 1 upon addition of NaOH. (b) Describe and explain the significance of the color changes observed for solution 1 in Data Table 3.
   1. CH₃CO₂H + NaOH → CH₃CO₂^–Na⁺ + H₂O
   2. The acetic acid reference solution (0.1 M) started out at pH 2.8. The congo red indicator color for this pH was purple. As NaOH was added in small increments to the solution it underwent a gradual color change and pH increase. The gradual color change reflects the fact that as a portion (but not all) of the acetic acid was neutralized to give some sodium acetate, the composition and pH of the solution were converted to the buffer range. The solution then acted as an excellent buffer, undergoing a very gradual color transition from purple-brown to orange-brown upon addition of excess NaOH.
8. Which buffer solution has the composition of a so-called “ideal buffer?” Do the results in Data Tables 2 and 3 support the conclusion that this buffer is an ideal buffer? Explain.

   Solution 3 (buffer solution B) was prepared from equal amounts of the two necessary buffer components, acetic acid and sodium acetate. This is the chemical composition of an ideal buffer, an equimolar mixture of the weak acid and its conjugate base. Yes, the data in Tables 2 and 3 support the conclusion that buffer B is an ideal buffer, since it was equally effective in resisting changes in pH upon addition of either acid or base. While buffer C was more effective than buffer B when HCl was added to the solution, it was much less effective than B when NaOH was added to the solution. The reverse is true for buffer A: it was better than B when NaOH was added, but much less effective when HCl was added to the buffer solution. The ideal pH has optimum activity under both conditions.

9. Based on the results and analysis obtained in this lab, estimate the effective ,em>pH range of acetic acid–sodium acetate buffer solutions.

   Buffer Solutions A, B and C all provided effective buffering activity upon addition of either acid or base (or both) to the solution. Buffers prepared from acetic acid–sodium acetate were confirmed to be effective in the pH range 4.1–5.1. According to the buffer equation and the calculated pH range for effective buffering, acetic acid–sodium acetate buffers should be effective in the range 4.75 ±1. The outer limits of this pH range were not tested in this lab activity.

Introduction

One of the most important applications of acids and bases in chemistry and biology is that of buffers. A buffer protects against rapid changes in pH when acids or bases are added to it. Every living cell is buffered to maintain constant pH and proper cell function. Consumer products are often buffered to safeguard their activity. The purpose of this lab activity is to investigate how buffers are made and the pH range in which they are effective.

Concepts

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

Background

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^–), which are related to each other by means of the dissociation reaction shown in Equation 1. An important feature of the dissociation reaction is the equilibrium arrow (⇄), which indicates that the reaction is reversible and that both the weak acid and the conjugate base are present in solution.

Buffers control pH because the two buffer components (HA and A^–) are able to neutralize either acid or base added to the solution. The weak acid component HA reacts with any base added to the solution to give its conjugate base A^–. The conjugate base component A^– reacts with any acid added to the solution to regenerate its acid partner HA. These competing reactions can be visualized as a cyclic process (see Figure 1). Buffer activity will continue as long as neither component A^– or HA is completely consumed or overwhelmed by the amount of strong acid or base. Properties of Weak Acids and Bases
The properties of weak acids and their conjugate bases determine why buffers behave as they do. The key difference between a weak acid and a strong acid is that dissociation of a weak acid is reversible and occurs to only a very limited degree in water. One familiar weak acid is acetic acid (CH₃COOH), which is the main ingredient in vinegar. A 0.1 M solution of acetic acid has a hydronium ion concentration [H₃O⁺] equal to 0.0013 M, giving an observed pH of 2.8–2.9. (Recall the definition and mathematical relationship between [H₃O⁺] and pH: pH = –log[H₃O⁺].) The observed pH value suggests that only about 1% of the acetic acid molecules are dissociated to the conjugate base form, acetate ion, under these conditions. In contrast, a strong acid such as hydrochloric acid (HCl) undergoes complete and irreversible 100% dissociation in water.

The degree to which a weak acid is ionized in aqueous solution is governed by the equilibrium constant K_a for its reversible dissociation reaction (Equations 2 and 3). The equilibrium constant K_a is also referred to as the dissociation constant of the weak acid. The K_a value for acetic acid, for example, is 1.76 x 10^–5. The Buffer Equation
Generalization of Equation 3 for any weak acid HA and its conjugate base A^– gives Equation 4, which can, in turn, be rearranged to solve for the [H₃O⁺] concentration (Equation 5). Equation 5 is known as the buffer equation; it provides the key to calculating the properties of buffer solutions. When the amount of the weak acid is equal to the amount of its conjugate base, the concentration ratio in Equation 5 will be equal to one and the [H₃O⁺] concentration will be equal to the dissociation constant K_a for the weak acid. Careful selection of the identity of the weak acid component makes it possible to prepare a buffer solution with almost any initial pH value. In the case of acetic acid, for example, a buffer solution consisting of a 1:1 molar mixture of acetic acid and its conjugate base sodium acetate will have a hydronium ion concentration equal to 1.76 x 10^–5 M, and the pH of the solution will be 4.75. Carbonic acid (H₂CO₃) has a K_a value equal to 4.30 x 10^–7; thus, a buffer prepared from equal moles of carbonic acid and its conjugate base bicarbonate ion (HCO₃^–) will have an [H₃O⁺] concentration equal to 4.30 x 10^–7 M and a pH value equal to 6.37.

What happens when strong acid or base is added to a buffer? Reaction of the weak acid component HA with additional base, such as sodium hydroxide, converts the weak acid to its conjugate base form A^– (Equation 6). Similarly, reaction of the basic component A^– with added acid results in its neutralization to the conjugate acid form HA (Equation 7). The following example shows how the buffer equation (Equation 5) can be used to calculate the effect of adding a strong acid or base on the pH of a buffer solution. Assume a buffer contains 0.5 moles of acetic acid and 0.5 moles sodium acetate. If 0.1 moles of strong acid are added to the solution, the H₃O⁺ will react with 0.1 moles of the sodium acetate present to give 0.1 additional moles of acetic acid, according to Equation 7. The amount of acetic acid in the buffer solution after addition of strong acid will increase to 0.6 (0.5 + 0.1) moles, while the amount of the conjugate base component, sodium acetate, will be reduced to 0.4 (0.5 – 0.1) moles. Substituting these numbers into Equation 5 and solving for the [H₃O⁺] concentration (Equation 8) shows that under these conditions [H₃O⁺] increases to 2.64 x 10^–5 M (from an initial value of 1.76 x 10^–5 M), giving a new pH value of 4.58 (compared to its initial pH of 4.75). The pH difference is only 0.17 units!

pH = –log[H₃O⁺] = –log(2.64 x 10^–5) = 4.58

Ideal Buffers and Buffer Range
A buffer composed of an equal number of moles 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. As shown in the example above, in an ideal buffer solution the [H₃O⁺] concentration is equal to the dissociation constant (K_a) for the weak acid.

The pH range in which a buffer solution will be effective is called the 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 concentration ratios of HA:A^– between 1:10 and 10:1. Substituting these concentration ratios in Equation 5 reveals that the effective pH range for a given buffer is plus or minus one unit on either side of the pH value of the ideal buffer. An ideal acetic acid–sodium acetate buffer system has a pH of 4.75 and its buffer range is 3.75–5.75. Equation 9 shows the calculation for the lower pH limit of an acetic acid–sodium acetate buffer solution (when the concentration ratio of the weak acid component to the conjugate base component is equal to 10:1).

pH = –log(1.76 x 10^–4) = 3.75

Overview of the Buffer Experiments and Their Indicators
Three different acetic acid–sodium acetate buffer solutions will be prepared and their initial pH values measured using pH paper and indicators. Strong acid and strong base will then be added to the buffers and the pH changes that result will be estimated by observing changes in the indicator colors. The acid–base indicators are bromcresol green (to study the effect of acid) and congo red (to study the effect of base), respectively. Bromcresol green changes color in the pH range 3.8–5.2. It is yellow when the pH is less than 3.8, blue when the pH is above 5.2, and intermediate shades of green in the transition range 3.8–5.2. Congo red undergoes more complex color changes but is a useful indicator for estimating pH values between 3.0–6.0. The color of congo red is purple at a pH around 3.0, changes to shades of brown and orange in the pH range 4–6, and finally reaches a stable red-orange color when the pH is greater than 7.

The properties of the buffer solutions will be compared against both reference (pure acetic acid and sodium acetate) and control (distilled water) solutions. The results will be analyzed to compare observed and calculated pH values and to calculate the buffer range of an ideal buffer.

Materials

Prelab Questions

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

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations.
   {13372_PreLab_Table_1}
2. Write the equations for the dissociation of the following weak acids and identify their conjugate bases: phosphoric acid (H₃PO₄), formic acid (HCO₂H) and boric acid (H₃BO₃).
3. What would be the composition and pH of an ideal buffer prepared from lactic acid (CH₃CHOHCO₂H), where the hydrogen atom highlighted in boldface is the acidic hydrogen atom? The K_a value for lactic acid is 1.38 x 10^–4.
4. Use the buffer equation to calculate the pH of buffer solutions prepared by dissolving the following amounts of acetic acid and sodium acetate, respectively, in enough water to make 1 L of solution:
   1. 0.67 moles acetic acid and 0.33 moles of sodium acetate
   2. 0.33 moles acetic acid and 0.67 moles of sodium acetate.

Safety Precautions

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

Procedure

Prelab Preparation

1. Measure approximately 33 mL each of 0.1 M acetic acid and 0.1 M sodium acetate solutions into two separate and labeled 50-mL beakers.
2. Obtain 10 mL each of 0.1 M hydrochloric acid and 0.1 M sodium hydroxide and place in separate, labeled, medium-size test tubes.
3. Place 5 mL of bromcresol green and congo red indicators in separate, labeled, small test tubes.

Reference, Control and Buffer Solutions

4. Set up six medium-size test tubes in a rack. Label them 1–6.
5. 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.
   {13372_Procedure_Table_2}
6. Mix the contents of each test tube thoroughly by gentle shaking or swirling.
7. Place a clean microscale 24-well reaction plate on a piece of white paper.
8. Label six graduated Beral-type pipets 1–6.
9. Using the appropriate clean graduated Beral-type pipet for each solution 1–6, fill each well in the microscale reaction plate with 1 mL of the designated solution, according to the following layout plan (numbers refer to vertical columns of wells and letters refer to horizontal rows of wells).
   {13372_Procedure_Table_3}

Part 1. pH Measurement

10. Measure the pH of each solution 1–4 in row A using a piece of narrow range pH paper (3.0–5.5). The pH values for solutions 5 and 6 are outside of this range and have already been entered in Data Table 1.
11. Record the estimated pH for each solution 1–4 in Data Table 1.
12. Add 1 drop of bromcresol green indicator to each of the 24 wells in the reaction plate.
13. Record the initial indicator color for each solution 1A–6A in Data Table 1. (Note: All of the colors in each vertical column should be identical at this point. If this is not the case, an error may have been made in filling the plate—start again.)

Part 2. Effect of HCl Addition

14. Use a clean pipet to add the indicated amounts (steps 15–17) of 0.1 M HCl dropwise to all of the wells in rows B, C and D only (row A will not be altered—it will serve as a color reference guide to compare observed color changes).
15. To wells 1B–5B in row B add 5 drops of 0.1 M HCl. Add only 1 drop of HCl to the last well in this row, well 6B, the control. Use a clean toothpick to stir each solution and to ensure thorough mixing.
16. Add 10 drops of 0.1 M HCl to each well in row C (wells 1C–6C). Stir each solution.
17. Add 20 drops of 0.1 M HCl to each well in row D (wells 1D–6D). Stir each solution.
18. Complete Data Table 2 by recording the color of solution in each well after HCl addition.
19. Wash the contents of the reaction well plate down the drain under running water. Rinse the empty well plate with distilled water and dry it.

Part 3. Effect of NaOH Addition

20. Repeat steps 7–9 to prepare a microscale reaction plate to study the effect of NaOH addition on buffer activity.
21. Add 1 drop of congo red indicator to each of the 24 wells in the reaction plate.
22. Record the initial congo red indicator color for each solution 1–6 in Data Table 1. (Note: All of the colors in each vertical column should be identical at this point. If this is not the case, an error may have been made in filling the plate—start again.)
23. Use a clean pipet to add the indicated amounts (steps 24–26) of 0.1 M NaOH dropwise to all of the wells in rows B, C and D only (row A will not be altered—it will serve as a color reference guide to note when color changes have occurred).
24. To wells 1B–5B in row B add 5 drops of 0.1 M NaOH. Add only 1 drop of NaOH to the last well in this row, well 6B, the control. Use a clean toothpick to stir each solution and to ensure thorough mixing.
25. Add 10 drops of 0.1 M NaOH to each well in row C (wells 1C–6C). Stir each solution.
26. Add 20 drops of 0.1 M NaOH to each well in row D (wells 1D–6D). Stir each solution.
27. Complete Data Table 3 by recording the color of solution in each well after NaOH addition.
28. Wash the contents of the reaction well plate down the drain under running water. Rinse thoroughly.

Student Worksheet PDF

13372_Student1.pdf