Materials Included In Kit

Additional Materials Required

Prelab Preparation

• pH 7 Buffer: Dissolve the contents of the pH 7 buffer envelope in 500 mL of distilled or deionized water. Mix well to dissolve.
• Hydrochloric acid solution, 0.1 M: Dilute 100 mL of the 0.5 M HCl to a final volume of 500 mL.
• Hydrochloric acid solution, 0.2 M: Dilute 200 mL of the 0.5 M HCl to a final volume of 500 mL.
• Sodium hydroxide solution, 0.1 M: Dilute 100 mL of 0.5 M NaOH to a final volume of 500 mL.
• Sodium hydroxide solution, 0.2 M: Dilute 200 mL of the 0.5 M NaOH to a final volume of 500 mL.
• Sodium acetate solution, 0.1 M: Dissolve 6.8 g sodium acetate trihydrate in 500 mL distilled or deionized water.

Safety Precautions

Dilute acid and base solutions, including acetic acid, ammonia, citric acid, hydrochloric acid and sodium hydroxide, are skin and eye irritants. Acetic acid and ammonia solutions may be irritating to the respiratory tract. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron when handling these chemicals. Avoid exposure of all chemicals to eyes and skin and clean up all spills promptly. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines. 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 may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Sodium hydrogen phosphate salts are also known and sold commercially by their common names. Sodium dihydrogen phosphate, NaH₂PO₄, is called sodium phosphate, monobasic, while sodium hydrogen phosphate, Na₂HPO₄, is also referred to as sodium phosphate, dibasic. To avoid confusing students, we suggest using the names that specify the number of hydrogen atoms in the polyatomic ion. Note that by extension the tribasic salt of phosphoric acid is Na₃PO₄, also known as sodium phosphate, tribasic.
• Distilled water (test tube A) is the appropriate comparison for interpreting the behavior of the buffer solutions in the Introductory Activity. The pH of the control changes immediately when even one drop of HCl or NaOH is added. In contrast, the pH of the buffer solutions does not change until 3–6 mL of HCl or NaOH have been added.
• Flinn pH meters (Catalog No. AP8673) provide an inexpensive and convenient way to measure pH values of solutions directly in a test tube or microscale reaction plate.
• Note the concentrations of HCl and NaOH are 0.1 M in the Introductory Activity and 0.2 M in the Guided-Inquiry Activity.
• Remind students to take into account the formulas of the hydrates in calculating the formula weights of the salts for the buffer challenges.
• Equilibrium constants and pK_a values are temperature dependent. The pH of an ideal buffer will change by up to 0.1 pH unit per degree Celsius.

Further Extensions

Opportunities for Undergraduate Research
Buffers can be prepared by partial neutralization of a weak acid with sodium hydroxide. Carbonate buffers covering the pH range from 9.4 to 10.8 can be prepared by combining the appropriate volumes of 0.2 M sodium bicarbonate and 0.2 M NaOH. Predict the appropriate volumes of NaHCO₃ and NaOH to combine to make buffers with pH values equal to 9.7, 10.3 and 10.8.

Answers to Prelab Questions

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations.
   {12721_PreLabAnswers_Table_4}
2. Write balanced chemical equations for dissociation of the following weak acids and identify their conjugate bases: phosphoric acid (H₃PO₄), formic acid (HCO₂H) and boric acid (H₃BO₃).
   {12721_Answers_Equation_11}
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.
      {12721_Answers_Equation_12}
   1. [H₃O⁺] = 1.76 x 10^–5 x (0.67/0.33) = 3.57 x 10^–5; pH = 4.45.
   2. [H₃O⁺] = 1.76 x 10^–5 x (0.33/0.67) = 8.56 x 10^–6; pH = 5.06.

Sample Data

Introductory Activity

Analysis of Results

• Addition of even one drop of HCl or NaOH to water caused a drastic change in pH, from 6.7 to 3.0 with acid, and from 6.7 to 10.9 with base. All of the buffers were resistant to pH changes when HCl or NaOH was added. Most buffers were able to maintain pH within ±1 unit, even after 3 mL of acid or base (one-third of their original volume) had been added.
• Buffer B had a higher capacity than buffer C or D with respect to added base, but a lower capacity for added acid. Buffer D had a higher capacity than buffer B or C with respect to added acid, but a lower capacity for added base.
• Buffer C has the composition of the ideal buffer, that is, a 1:1 ratio of the weak acid and its conjugate base. The pH of this buffer was equal to the theoretical or calculated pH; pH = pK_a.
• The experimental pH values for buffers B and D were very close to the calculated values of 4.45 for buffer B and 5.06 for buffer D (see Prelaboratory Assignment).

Guided-Inquiry Activity

Answers to Questions

Guided-Inquiry Design and Procedure 

1. Recalling that pH = –log[H₃O₊] and pK_a = –logK_a, transform Equation 5 in the Background section to the Henderson-Hasselbach equation relating the pH of a buffer solution to the pK_a value of a weak acid and the concentrations of the weak acid and its conjugate base.

   The Henderson-Hasselbach equation is derived by taking the negative log of each term in the equation and applying the mathematical reasoning that log (A x B) = log A + log B.
   –log[H₃O⁺] = −log(K_a x [ HA]/[A⁻])
   pH = pK_a – log([HA]/[A^–]) and, since –log([HA]/[A^–]) = log([A^–]/[HA])
   pH = pK_a + log([A^–]/[HA])

2. Citric acid is a triprotic weak acid with three ionizable hydrogen atoms. See its structure and following formula. Write equations for the three stepwise deprotonation reactions of citric acid and its conjugate base anions. Hint: Salts of the conjugate bases are called sodium dihydrogen citrate, sodium hydrogen citrate and sodium citrate.
   {12721_Answers_Equation_13}
   The other two —CO₂H groups may be deprotonated in turn, as follows:
   C₆H₇O₇⁻ + H₂O + C₆H₆O₇²⁻ + H₃O⁺
   C₆H₆O₇²⁻ + H₂O → C₆H₅O₇³⁻ + H₃O⁺
3. The pK_a values for the three stepwise dissociations of citric acid are 3.1, 4.8 and 6.4. (a) What is the expected composition of a citrate buffer having a pH value = 3.1? (b) What is the expected composition of a citrate buffer having a pH value = 6.4?
   1. The expected composition of a citrate buffer with pH = 3.1 is an equimolar ratio of citric acid (H₃A) and the dihydrogen citrate anion H₂A⁻.
   2. A citrate buffer composed of an equimolar mixture of hydrogen citrate HA²⁻ and citrate anion A³⁻ should have a pH value = 6.4.
4. Basic buffers, that is, those with pH values > 7, are derived from weak acids and their conjugate bases having pK_a values greater than 7. Consider bicarbonate ion and its conjugate base carbonate ion (Equation 9). The value of the dissociation constant for this reaction is 4.7 x 10^–11. Describe the pH and composition of an ideal bicarbonate/carbonate buffer.
   {12721_Answers_Equation_9}
   An ideal bicarbonate/carbonate buffer would be prepared by combining equal moles of sodium bicarbonate and sodium carbonate and would have a pH = 10.3.
5. Ammonia is a weak base.
   1. Write an equation for the reaction of ammonia with water and identify its conjugate acid.

      Reaction of ammonia with water: NH₃ + H₂O → NH₄⁺ + OH⁻
      The conjugate acid of ammonia is the ammonium ion, NH₄⁺.

   2. The pK_b value for this reaction is 4.7. Recalling that the relationship between pK_a and pK_b for a conjugate acid–base pair is pK_a + pK_b = 14, predict the pH of an equimolar solution containing ammonia and its conjugate acid.

      The pK_b value for ammonia is 4.7, and the pK_a value for the ammonium ion = 14 − 4.7 = 9.3. A buffer solution consisting of equal moles of NH₄⁺ and NH₃ is expected to have a pH value = 9.3.

6. The purpose of this activity is to design a buffer for a specific consumer or biochemical application. The weak acids that are available to prepare the target buffers include acetic acid, ammonium chloride, carbonic acid, citric acid and sodium dihydrogen phosphate (0.1 M solutions of each are provided). Carbonic acid is present in seltzer water, which contains approximately 0.07 moles of carbon dioxide per liter. Write equations for the reactions of these weak acids with water, identify their conjugate bases, and determine the pK_a value for each.
   {12721_Answers_Table_6}
7. Each group should choose one of the buffer challenges from the following table and select an appropriate weak acid–conjugate base pair to prepare the buffer.
   {12721_Answers_Table_7}

   *The concentration of dissolved CO₂ in seltzer varies with pressure and temperature. At typical bottling pressures = 2.5 atm the [CO₂] concentration is reported to be about 0.08 M. We have found the equilibrium [CO₂] (and potential [H₂CO₃]) concentration to be 0.07 M. If students assume this value, they should obtain good results.

8. The specifications for each buffer are that the pH should be within ±0.5 pH units of the target pH, and 25 mL of the buffer should be able to maintain the desired pH within ±1 pH unit after the addition of either 10 mL of 0.2 M HCl or 10 mL 0.2 M NaOH. Predict the ratio of weak acid/conjugate base to meet the buffer challenge, and determine the amounts of HA and A^– to provide the desired buffer capacity.

   See the table for the calculated ratios.

Post-Laboratory Review

1. The major buffer in blood is composed of the weak acid carbonic acid (H₂CO₃) and its conjugate base, bicarbonate ion (HCO₃⁻). The normal pH of blood is 7.2−7.4, which is very far removed from the pK_a value. The pH is kept in check by the lungs, which remove CO₂ via exhalation, and by the kidneys, which excrete acid (H₃O⁺) in the urine. People with impaired lung function are not able to exchange carbon dioxide efficiently between the lungs and air. The result is an increase in the amount of CO₂ dissolved in the blood.
   1. How does this affect the buffer balance in the blood?

      Increasing the amount of dissolved CO₂ in blood will shift the equilibrium to contain more of the acid component H₂CO₃.

   2. Which term, respiratory acidosis or respiratory alkalosis, would better describe the resulting condition?

      The pH of blood will decrease—this is called respiratory acidosis.

2. Explain why a mixture of the strong acid HCl and its conjugate base NaCl does not provide buffering action.

   The conjugate base of the strong acid HCl is the chloride anion, which is neutral. Chloride ion will not react with any excess strong acid added to it. In a buffer the acid component neutralizes excess base and the basic component neutralizes excess acid.

3. Forensic analysis of DNA by electrophoresis requires the use of a pH 8.3 buffer to ensure that the DNA phosphate groups remain negatively charged. The major constituent of electrophoresis buffers is called TRIS, which stands for tris(hydroxymethyl) aminomethane. Its structure is shown. What weak acid/weak base combination used in this activity is TRIS most analogous to? Identify the basic functional group in TRIS that is protonated to give a weak acid.
   {12721_Answers_Figure_4}
   TRIS is analogous to the weak base ammonia, NH₃. It contains a basic nitrogen atom with a lone pair of electrons. The protonated form is a substituted ammonium cation (shown).
4. Many soft drinks contain phosphate buffers. Calculate the pH of an 8 oz. soft drink containing 4.4 g of sodium dihydrogen phosphate (formula weight = 120 g/mole) and 5.4 g of sodium hydrogen phosphate (formula weight = 142 g/mole).

   8 oz x 30 mL/oz = 240 mL (0.24 L)

   {12721_Answers_Equation_14}

Introduction

A buffer protects against rapid changes in pH when acids or bases are added. 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, the pH range in which they are effective and their buffer capacity.

Concepts

• pH
• Buffer
• Weak acids and bases
• Dissociation constant
• Neutralization
• Conjugate acid–base pairs

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 the dissociation reaction shown in Equation 1. The double arrow (⇄) indicates that the reaction is reversible and that both the weak acid and the conjugate base are present at equilibrium.

Buffers control pH because the buffer components HA and A^– are able to neutralize either acid or base added to the solution. The weak acid component HA reacts with base to give its conjugate base A^–. The conjugate base component A^– reacts with acid to regenerate its acid partner HA. These 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 by the amount of added 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. Dissociation of a weak acid is reversible and occurs to only a very limited degree in water. Consider acetic acid (CH₃COOH), 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 K_a value for acetic acid 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 be rearranged to solve for the [H₃O⁺] concentration (Equation 5). Equation 5 is sometimes known as the buffer equation; it provides the key to calculating the properties of buffer solutions. When the concentrations of the weak acid and its conjugate base are equal, the 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 4.4 x 10^–7. A buffer prepared from equal moles of carbonic acid and its conjugate base bicarbonate ion (HCO₃^–) will have an [H₃O⁺] concentration equal to 4.4 x 10^–7 M and a pH value equal to 6.4.

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 effect of adding a strong acid or base on the pH of a buffer solution can be predicted using Le Chatelier’s principle. Consider the equimolar acetic acid–acetate buffer (Equation 2). Adding HCl to the buffer solution, with its equilibrium pH = 4.75, increases the concentration of H₃O⁺ ions, one of the products of the reversible reaction. This shifts the equilibrium to the left, increasing the concentration of acetic acid and decreasing the concentration of acetate ions. The ratio of [HA] to [A^–] in Equation 5 increases as well, and [H₃O⁺] is larger—the pH decreases. The opposite effect is observed when NaOH is added to the buffer solution. OH^– ions neutralize some of H₃O⁺ ions, which shifts the equilibrium to the right, increasing the concentration of acetate ions relative to acetic acid molecules. The ratio of [HA] to [A^–] decreases, and [H₃O⁺] is smaller—the pH increases. In either case, however, as long as the [HA]/[A^–] ratio stays within certain limits, the pH remain relatively constant.

Buffer Range and Buffer Capacity
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. As shown in the example, 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 into 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 has a pH of 4.75 and its buffer range is 3.75–5.75. Equation 8 shows the calculation for the lower pH limit of an acetic acid–sodium acetate buffer where the concentration ratio of the weak acid component to the conjugate base component is 10:1.

pH = –log(1.76 x 10^–4) = 3.75

The effectiveness of a buffer in resisting pH changes is called the buffer capacity. Consideration of Equation 5 reveals that the pH of a buffer prepared from a weak acid HA and its conjugate base A^– should be independent of their total concentration as long as the ratio [HA] to [A^–] is the same. Thus, an acetic acid–acetate buffer prepared from 0.1 mole HA and 0.1 mole A^– should have the same theoretical pH as a buffer containing 1 mole HA and 1 mole A–. The buffer capacity of the two buffers, however, will be very different. The capacity of the 0.1 moles HA/0.1 moles A^– buffer will be overwhelmed when approximately 0.09 moles of HºCl or NaOH have been added. The 1 M buffer will withstand almost 10X as much strong acid or strong base before either HA or A^– is consumed.

Experiment Overview

The purpose of this inquiry lab is to design an effective buffer with a specific pH value for a consumer or experimental biochemistry application. The investigation begins with an introductory activity to compare the properties of three acetate buffers containing varying ratios of HA and A^–. The results provide a model for guided-inquiry design of an experiment to prepare a desired buffer and verify its properties and performance. Five different buffer “challenges” are presented—each student group chooses one. The specifications for each buffer challenge are that (a) the pH should be within ±0.5 pH units of the desired pH, and (b) 25 mL of the buffer should maintain the desired pH ±1 after 10 mL of 0.02 M HCl or 10 mL of 0.2 M NaOH have been added. Preparation of a buffer by partial neutralization of a weak acid or a weak base offers additional opportunities for inquiry.

Materials

Prelab Questions

1. Calculate the pH value in each of the following solutions, given their [H₃O⁺] concentrations.
   {12721_PreLab_Table_1}
2. Write balanced chemical equations for 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 acid and base solutions, including acetic acid, ammonia, citric acid, hydrochloric acid and sodium hydroxide, are skin and eye irritants. Acetic acid and ammonia solutions may be irritating to the respiratory tract. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Avoid exposure of all chemicals to eyes and skin and notify the teacher of any spills. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

Introductory Activity

Acetate Buffers

1. Using a clean graduated cylinder for each solution, measure and add the indicated volumes of the following solutions to five test tubes labeled A−E. Mix the contents of test tubes B, C and D by gentle stirring, swirling or shaking.
   Table 1.
   {12721_Procedure_Table_2}
2. Measure and record the initial pH of the water in test tube A using either a pH meter or a combination of wide-range and narrow-range pH paper or indicators.
3. Add 1 drop of 0.1 M HCl to the water in test tube A and measure the pH.
4. If using a pH meter, rinse the electrode with distilled water and blot dry.
5. Measure and record the initial pH of the buffer in test tube B.
6. Using a graduated pipet or buret, add 1 mL of 0.1 M HCl to the buffer in test tube B. Measure and record the pH.
7. Add an additional 2 mL of 0.1 M HCl to test tube B, and again measure the pH.
8. Add an additional 3 mL of 0.1 M HCl to test tube B, and again measure the pH.
9. Rinse the pH meter with distilled water (if applicable).
10. Repeat steps 5–9 three more times to test the buffer solutions in test tubes C, D and E.
11. Dispose of the solutions in test tubes A–E. Rinse the test tubes with distilled water and blot dry with a paper towel.
12. Refill all test tubes A–E with the designated solutions shown in Table 1.
13. Measure the initial pH of the water in test tube A. Add 1 drop of 0.1 M NaOH, and again measure the pH.
14. If using a pH meter, rinse the electrode with distilled water and blot dry.
15. Measure and record the initial pH of buffer B.
16. Using a graduated pipet or buret, add 1 mL of 0.1 M NaOH to buffer B. Measure and record the pH.
17. Add an additional 2 mL of 0.1 M NaOH to buffer B, and again measure the pH.
18. Add an additional 3 mL of 0.1 M NaOH to buffer B, and again measure the pH.
19. If using a pH meter, rinse the electrode with distilled water and blot dry.
20. Repeat steps 15–19 three more times to test the buffer solutions in test tubes C, D and E.
21. Dispose of the solutions in test tubes A–E.

Analyze the Results

• Compare the observed pH changes for distilled water versus the buffers in test tubes B–E when either HCl or NaOH was added.
• Which acetate buffer (B, C or D) was most effective with respect to added HCl? Explain.
• Which acetate buffer (B, C or D) was most effective with respect to added NaOH? Explain.
• Which acetate buffer (B, C or D) has the composition of an “ideal buffer?” Do the results support this description? Explain.

Guided-Inquiry Design and Procedure
Form a working group with other students and discuss the following questions.

1. Recalling that pH = –log[H₃O⁺] and pK_a = –logK_a, transform Equation 5 in the Background section to the Henderson-Hasselbach equation relating the pH of a buffer solution to the pK_a value of a weak acid and the concentrations of the weak acid and its conjugate base.
2. Citric acid is a triprotic weak acid with three ionizable hydrogen atoms. See its structure and following formula. Write equations for the three stepwise deprotonation reactions of citric acid and its conjugate base anions. Hint: Salts of the conjugate bases are called sodium dihydrogen citrate, sodium hydrogen citrate and sodium citrate.
   {12721_Procedure_Equation_10}
3. The pK_a values for the three stepwise dissociations of citric acid are 3.1, 4.8 and 6.4.
   1. What is the expected composition of a citrate buffer having a pH value = 3.1?
   2. What is the expected composition of a citrate buffer having a pH value = 6.4?
4. Basic buffers, that is, those with pH values > 7, are derived from weak acids and their conjugate bases having pK_a values greater than 7. Consider bicarbonate ion and its conjugate base carbonate ion (Equation 9). The value of the dissociation constant for this reaction is 4.7 x 10^–11. Describe the pH and composition of an ideal bicarbonate/carbonate buffer.
   {12721_Procedure_Equation_9}
5. Ammonia is a weak base.
   1. Write an equation for the reaction of ammonia with water and identify its conjugate acid.
   2. The pK_b value for this reaction is 4.7. Recalling that the relationship between pK_a and pK_b for a conjugate acid–base pair is pK_a + pK_b = 14, predict the pH of an equimolar solution containing ammonia and its conjugate acid.
6. The purpose of this activity is to design a buffer for a specific consumer or biochemical application. The weak acids that are available to prepare the target buffers include acetic acid, ammonium chloride, carbonic acid, citric acid and sodium dihydrogen phosphate (0.1 M solutions of each are provided). Carbonic acid is present in seltzer water, which contains approximately 0.07 moles of carbon dioxide per liter. Write equations for the reactions of these weak acids with water, identify their conjugate bases, and determine the pK_a value for each.
7. Each group should choose one of the buffer challenges from the following table and select an appropriate weak acid–conjugate base pair to prepare the buffer.
   Table 2.
   {12721_Procedure_Table_3}
8. The specifications for each buffer are that the initial pH should be within ±0.5 pH units of the target pH, and 25 mL of the buffer should be able to maintain the desired pH within ±1 pH unit after the addition of either 10 mL of 0.2 M HCl or 10 mL 0.2 M NaOH. Predict the ratio of weak acid/conjugate base to meet the buffer challenge, and determine the amounts of HA and A^– to provide the desired buffer capacity.
9. Write a detailed step-by-step procedure for preparing the selected buffer and testing its buffer capacity. Include all materials, glassware and equipment that will be needed, safety precautions that must be followed, amounts of each reactant, etc.
10. Review additional variables that may affect the reproducibility or accuracy of the investigation, and how these variables will be controlled.
11. Carry out the investigation and record results in appropriate data tables and graphs.
12. Each group should present evidence in the form of titration curves to show that the respective buffer specifications have been achieved.

Analyze the Results
Organize a collaborative class research meeting to discuss the results.

Student Worksheet PDF

12721_Student1.pdf