Materials Included In Kit

Additional Materials Required

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulation that may apply, before proceeding. The sodium hydroxide solution may be disposed of according to Flinn Suggested Disposal Method #10. The hydrochloric acid solution may be disposed of according to Flinn Suggested Disposal Method #24b. The acetic acid solution may be disposed of according to Flinn Suggested Disposal Method #24a. The buffer solutions and the sodium acetate solution may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• The experimental work for this lab can be completed in two 50-minute lab periods. The prelab assignment should be completed prior to the lab. The calculation section can be assigned as a post-lab activity or performed as a class activity the day following the lab.
• Make sure students are familiar with the proper calibration and use of the pH meters in your classroom.
• Remind students to allow adequate time for the pH reading to stabilize before recording it.
• When calculating the pH upon adding either HCl or NaOH, the student needs to remember that when mixed, the concentrations of acetic acid and sodium acetate in the buffer are 0.05 M, not 0.1 M.
• Stress to students the importance of rinsing all glassware and the stir bar with deionized water before starting a new determination. Clearly label all solutions.

Teacher Tips

• The calculations of pH values for acid or base additions to the ammonia–ammonium chloride buffer test the student’s real understanding of the buffering concept. Many students will not calculate pH, but pOH instead. Check student answers to Prelab Question 3.
• 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^–). 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.
  {13894_Tips_Equation_10}
• It is only a coincidence the K_b for ammonia and K_a for acetic acid both equal 1.8 x 10^–5.

Further Extensions

AP^® Chemistry Standards 
The lab fulfills the requirement for the College Board recommended AP Experiment #19: Preparation and Properties of Buffer Solutions. In addition, this lab provides the recommended familiarity with the use of pH meters.

Answers to Prelab Questions

1. How many grams of sodium acetate (molar mass 82.03 g/mol) must be added to 1.00 L of 0.200 M acetic acid solution to form a buffer of 4.20? K_a value for acetic acid is 1.8 x 10^–5.
   {13894_PreLabAnswers_Equation_11}
2. Three milliliters of a 2.0 M solution of HCl are added to 1 liter of buffer solution containing 0.40 moles of a weak acid, propanoic acid (K_a = 1.4 x 10^–5) and 0.50 moles of its conjugate base, sodium propanate.
   1. What is the original pH of the buffer before the strong acid is added?
      {13894_PreLabAnswers_Equation_12}
   2. What is the pH of the buffer after the HCl is added? Assume negligible volume change.
      {13894_PreLabAnswers_Equation_13}
3. The weak base–conjugate acid buffer used in this laboratory consists of a weak base, ammonia, NH₃, and its conjugate acid, ammonium chloride, NH₄Cl. If the NH₃ concentration is 0.05 M and the NH₄Cl concentration is also 0.05 M, what is the pH of the buffer? K_b for NH₃ is 1.8 x 10⁵. 
   {13894_PreLabAnswers_Equation_14}

Sample Data

Part 1. pH of Acetic Acid–Acetate Buffer

Part 2. pH of Ammonia–Ammonium Chloride Buffer
Part 3. Preparation of a pH 5.00 Buffer Solution

Answers to Questions

Calculations

1. Using Equation 4 in the Background section, calculate the pH of the Part 1 acetic acid–sodium acetate buffer solution before and after 1.0 mL of 0.2 M HCl solution is added to the buffer. K_a of acetic acid equals 1.8 x 10^–5. Enter these values in the Part 1 Data Table.
   1. Before any acid is added, [HA] = [A^–] = 0.05 M

      Therefore,

      pH = pK_a – log 1 = pK_a
      pH = –log (1.8 x 10^–5) = 4.74

   2. 1.0 mL of 0.2 M HCl = 0.0002 moles of H₃O⁺ ion added.

      Initial moles of HA in solution: 0.025 L x 0.10 mol/L = 0.0025 moles
      Initial moles of A^– in solution: 0.025 L x 0.10 mol/L = 0.0025 moles

      H₃O⁺ + A^– → HA + H₂O

      HA moles after addition of acid = (0.0025 + 0.0002) mol = 0.0027 moles
      A^– moles after addition of acid = (0.0025 – 0.0002) mol = 0.0023 moles

      {13894_Answers_Equation_15}

      pH = 4.74 – 0.07 = 4.67

2. Repeat the pH calculation for each successive 1.0 mL increment of 0.2 M HCl added to the buffer. Enter these values in the Part 1 Data Table.

   2.0 mL of 0.2 M = 0.0004 moles of H₃O⁺ ion added.
   HA moles after addition of acid = (0.0025 + 0.0004) mol = 0.0029 moles
   A^– moles after addition of acid = (0.0025 – 0.0004) mol = 0.0021 moles

   {13894_Answers_Equation_16}

   pH = 4.74 – 0.14 = 4.60

   Repeat this calculation for each 1.0 mL increment addition of acid to the solution.

3. When strong base is added to a buffer of a weak acid–conjugate base, the acid reacts with the base to form water and its conjugate base.

   HA(aq) + OH^–(aq) → H₂O(l) + A^–(aq)

   Calculate the pH of the Part 1 acetic acid–sodium acetate buffer solution after 1.0 mL of the 0.2 M NaOH solution is added to the buffer. Enter this value in the Part 1 Data Table.

   1.0 mL of 0.2 M HCl = 0.0002 moles of OH^– ion added.
   Initial moles of HA in solution: 0.025 L x 0.10 mol/L = 0.0025 moles
   Initial moles of A^– in solution: 0.025 L x 0.10 mol/L = 0.0025 moles

   OH^– + HA → A^– + H₂O

   HA moles after addition of base = (0.0025 – 0.0002) mol = 0.0023 moles
   A^– moles after addition of base = (0.0025 + 0.0002) mol = 0.0027 moles
   {13894_Answers_Equation_17}

   = 4.74 + 0.07 = 4.81

4. Repeat the pH calculation for each successive 1.0 mL increment of 0.2 M NaOH added to the buffer. Enter these values in the Part 1 Data Table.

   2.0 mL of 0.2 M NaOH = 0.0004 moles of OH^– ion added.
   HA moles after addition of base = (0.0025 – 0.0004) mol = 0.0021 moles
   A^– moles after addition of base = (0.0025 + 0.0004) mol = 0.0029 moles

   {13894_Answers_Equation_18}

   pH = 4.74 + 0.14 = 4.88

   Repeat this calculation for each 1.0 mL increment addition of base to the solution.

5. The ammonia–ammonium chloride buffer solution is a weak base–conjugate acid buffer solution. K_b for NH₃ equals 1.8 x 10^–5. Using Equation 4 in the Background section and the relationship: pH = 14.00 – pOH calculate the pH of the ammonia–ammonium chloride buffer solution after 1.0 mL of 0.2 M HCl is added to the buffer solution. The initial moles of both NH₃ and NH₄Cl in 50 mL of the buffer solution are 0.0025 moles. Record the pH value in the Part 2 Data Table. [NH₃(aq) + H₃O⁺(aq) → NH₄⁺(aq) + H₂O(l)]

   1.0 mL of 0.2 M HCl = 0.0002 moles of H₃O⁺ ion added.
   Initial moles of B (weak base) in solution: 0.050 L x 0.05 mol/L = 0.0025 moles
   Initial moles of BH⁺ (weak base) in solution: 0.050 L x 0.05 mol/L = 0.0025 moles

   B + H₃O+ → BH⁺ + H₂O

   B moles after addition of acid = (0.0025 – 0.0002) mol = 0.0023 moles
   BH⁺ moles after addition of acid = (0.0025 + 0.0002) mol = 0.0027 moles
   for a weak base–conjugate acid buffer

   {13894_Answers_Equation_19}

   pOH = 4.74 + 0.07 = 4.81
   pH = 14.00 – pOH = 14.00 – 4.81 = 9.19

6. Repeat the pH calculation for each successive 1.0 mL increment of 0.2 M HCl added to the buffer. Enter these values in the Part 2 Data Table.

   2.0 mL of 0.2 M HCl = 0.0004 moles of H₃O⁺ ion added.
   B moles after addition of acid = (0.0025 – 0.0004) mol = 0.0021 moles
   BH⁺ moles after addition of acid = (0.0025 + 0.0004) mol = 0.0029 moles

   {13894_Answers_Equation_20}

   pOH = 4.74 + 0.14 = 4.88
   pH = 14.00 – pOH = 14.00 – 4.88 = 9.12

   Repeat this calculation for each 1.0 mL increment addition of acid to the solution.

7. Repeat the pH calculations for each 1.0 mL increment of 0.2 M NaOH added to the ammonia–ammonium chloride buffer solution. Enter these values in the Part 2 Data Table.

   1.0 mL of 0.2 M NaOH = 0.0002 moles of OH^– ion added.
   Initial moles of B and BH⁺ in solution are 0.0025 moles

   OH^– + BH⁺ → B + H₂O

   B moles after addition of base = (0.0025 + 0.0002) mol = 0.0027 moles
   BH⁺ moles after addition of base = (0.0025 – 0.0002) mol = 0.0023 moles

   {13894_Answers_Equation_21}

   pH = 14.00 – pOH = 14.00 – 4.67 = 9.33

   Repeat this calculation for each 1.0 mL increment addition of base to the solution.

Post-Lab Questions

1. Calculate the pH change when 1 mL of 0.2 M HCl is added to 50 mL of deionized water. How does this pH value change compare to those obtained when 1.0 mL of 0.2 M HCl is added to the buffers?

   pH = –log [H₃O⁺]
   
   For 1 mL of 0.2 M HCl added to 50 mL of deionized water, moles of H₃O⁺ = 0.2 moles/L x 0.001 L = 0.0002 moles
   
   [H₃O⁺] = 0.0002 moles/L x 0.0051 L = 0.0039 moles/L
   
   pH = –log (0.0039) = 2.40
   
   Change in pH = 7.00 – 2.40 = 4.60 units
   
   For the acetic acid–sodium acetate buffer, from data tables,
   Initial pH = 4.74 and after addition of 1 mL of 0.2 M HCl,
   pH = 4.67
   Change in pH = 0.07 units
   
   For the ammonia–ammonium chloride buffer, the change in pH equals
   Change in pH = 9.26 – 9.19 = 0.07 units

2. At what point did each of the buffers lose their effectiveness? Explain.

   When the ratio of [HA]/[A^–] and [B]/[BH⁺] nears 10:1, the buffering capacity of the solution is exceeded and the solution can no longer maintain the pH.

Introduction

One of the most important applications of acids and bases in chemistry and biology is that of buffers. A buffer solution resists rapid changes in pH when acids and bases are added to it. Every living cell contains natural buffer systems to maintain the constant pH needed for proper cell function. Many consumer products are also buffered to safeguard their activity. What are buffers made of? How do buffers maintain the delicate pH balance needed for life and health?

Concepts

• Buffer
• Conjugate acid–base pair
• Ideal buffer
• Weak acid–weak base
• Dissociation constant

Background

The ability of buffers to resist changes in pH when acid or base is added is a result of their chemical composition. All buffers contain a mixture of a conjugate acid–base pair; either a weak acid (HA) and its conjugate base (A^–), or a weak base (B), and its conjugate acid (BH⁺). Weak acids and weak bases both dissociate slightly in water (Reactions 1 and 2).

These reactions are reversible and both the weak acid and its conjugate base or the weak base and its conjugate acid are present in solution.

The equilibrium constant expressions for these dissociation reactions are: Buffers control pH because the two buffering components, either HA and A^– or B and BH⁺, are able to neutralize both acids and bases added to the solution. The actual pH of a buffer solution depends on the concentration of the conjugate acid–base pair in solution. If Equation 1 is rearranged, the concentration of hydronium ions in solution is: and the pH is: If the concentrations of the acid–base pair are equal, [HA] = [A^–]. is equal to zero, and the pH of the buffer is equal to pK_a. By varying the amounts of HA and A^– in solution, the pH of the buffer solution can be changed. For a buffer made up of a weak base (B) and its conjuugate acid (BH⁺), the solution pH calculations are similar. If Equation 2 is rearranged, the concentration of hydroxide ions (OH^–) in solution is: and the pOH is: If pOH is known, then pH can be calculated using Equation 7:

pH + pOH = 14.00

Once the buffer is made, how does the pH remain constant when strong acid or base is added? Acetic acid is a weak acid, with K_a equal to 1.8 x 10^–5. If a buffer solution is made with 0.5 moles of acetic acid and 0.5 moles of its conjugate base sodium acetate, the initial pH of the solution will be equal to pK_a, or 4.74. Now, if 0.05 moles of a strong acid is added to the buffer, the H₃O⁺ will react with 0.05 moles of the sodium acetate to form 0.05 moles of acetic acid. This produces a solution with 0.55 moles of acetic acid and 0.45 moles of its conjugate base sodium acetate. If the solution volume change is slight, then the new pH of the solution is: The pH difference is only 0.09 units! 

For buffers to be effective, noticeable amounts of both the conjugate acid–base pair must be present in solution. This limits the concentration ratios for HA:A^– or B:BH⁺ to between 10:1 and 1:10 and the pH range for the buffering action of any weak acid to pK_a ±1. An ideal buffer is a solution that contains equal numbers of moles of the conjugate acid–base pair.

Experiment Overview

The purpose of this experiment is to study the properties of buffer solutions. Two ideal buffer solutions, one consisting of a weak acid and its conjugate base and the other, a weak base and its conjugate acid, are made. The initial pH of each buffer is determined. Strong acid and strong base, are added to each buffer in a series of steps and the pH is determined after each addition. The resulting pH values after each addition will be compared to the calculated pH values for each buffer.

Materials

Prelab Questions

1. How many grams of sodium acetate (molar mass 82.03 g/mol) must be added to 1.00 L of a 0.200 M acetic acid solution to form a buffer of 4.20? K_a value for acetic acid is 1.8 x 10^–5.
2. Three milliliters of a 2.0 M solution of HCl are added to 1 liter of buffer solution containing 0.40 moles of the weak acid, propanoic acid (K_a = 1.4 x 10^–5) and 0.50 moles of its conjugate base, sodium propanate.
   1. What is the original pH of the buffer before the strong acid is added?
   2. What is the pH of the buffer after the HCl is added? Assume negligible volume change.
3. The weak base–conjugate acid buffer used in this laboratory consists of a weak base ammonia, NH₃, and its conjugate acid ammonium chloride, NH₄Cl. If the NH₃ concentration is 0.05 M and the NH₄Cl concentration is 0.05 M, what is the pH of the buffer? K_b for NH₃ is 1.8 x 10^–5.

Safety Precautions

The acetic acid solution is slightly corrosive. The hydrochloric acid solution is toxic by ingestion and inhalation and is corrosive to skin and eyes. The sodium hydroxide solution is corrosive to eyes, skin and other tissue. The ammonia/ammonium chloride buffer is strongly basic and is corrosive to skin, eyes and other tissues. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and an apron. Have students thoroughly wash their hands with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

Part 1. Buffering Properties of a Weak Acid–Conjugate Base Buffer

1. Obtain 25.0 mL of the 0.1 M acetic acid solution in a 25-mL graduated cylinder. Transfer the acetic acid solution to a clean 100-mL beaker.
2. Rinse the 25-mL graduated cylinder with deionized water. Obtain 25.0 mL of the 0.1 M sodium acetate solution with the 25-mL graduated cylinder.
3. Transfer the sodium acetate solution to the 100-mL beaker containing 25 mL of the acetic acid solution. Stir the solution.
4. Set-up a pH meter and electrode (or a pH sensor). Calibrate the pH meter using a standard pH 7 buffer solution.
5. Remove the pH 7 buffer solution, place a 100-mL beaker under the electrode, and rinse the electrode well with deionized water.
6. Set the 100-mL beaker containing the acetic acid–acetate buffer solution on a magnetic stirrer, if one is available. Add a stir bar to the solution. Gently stir the buffer solution.
7. Place the pH electrode in the solution. Record the pH of the solution in the Part 1 Data Table. (0 mL of 0.2 M HCl added.)
8. Obtain approximately 30 mL of 0.2 M HCl solution in a clean 50-mL beaker and label the beaker, 0.2 HCl.
9. Using a graduated Beral-type pipet, transfer 1.0 mL of the 0.2 M HCl solution to the acetic acid–acetate buffer solution.
10. Record the pH of the solution in the Part 1 Data Table. (1 mL of 0.2 M HCl added)
11. Using the same graduated pipet, transfer another 1.0 mL of the 0.2 M HCl solution to the acetic acid–acetate buffer solution.
12. Record the pH of the solution in the Part 1 Data Table.
13. Repeat steps 11 and 12, recording the pH in the Part 1 Data Table after each 1.0 mL of HCl is added, until a total of 10.0 mL of HCl has been added to the solution.
14. Remove the electrode from the solution, place a 100-mL beaker under the electrode, and rinse the electrode with deionized water.
15. Dispose of the acetic acid–sodium acetate solution as directed by the instructor, and rinse the 100-mL beaker with deionized water. Do not dispose of the HCl.
16. Obtain 25.0 mL of the 0.1 M acetic acid solution in a clean 25-mL graduated cylinder. Transfer the acetic acid solution to the rinsed 100-mL beaker.
17. Rinse the 25-mL graduated cylinder with deionized water. Obtain 25.0 mL of the 0.1 M sodium acetate solution in the 25-mL graduated cylinder.
18. Transfer the sodium acetate solution to the 100-mL beaker containing 25 mL of the acetic acid solution. Stir the solution.
19. Place the 100-mL beaker on the magnetic stirrer. Add a stir bar and gently stir the buffer solution.
20. Place the pH electrode in the solution. Record the pH of the solution in the data table. (0 mL of 0.2 M NaOH added.)
21. Obtain approximately 30 mL of 0.2 M NaOH solution in a clean 50-mL beaker. Label the beaker, 0.2 M NaOH.
22. Using a clean graduated Beral-type pipet, transfer 1.0 mL of the 0.2 M NaOH solution to the acetic acid–acetate buffer solution.
23. Record the pH of the solution in the Part 1 Data Table.
24. Repeat steps 22 and 23, recording the pH in the data table after each 1.0 mL of NaOH solution is added, until a total of 10.0 mL of NaOH has been added.
25. Remove the pH electrode from the solution, place a 100-mL beaker under the electrode, and rinse the electrode with deionized water.
26. Dispose of the acetic acid–sodium acetate solution as directed by the instructor, and rinse the 100-mL beaker with deionized water. Do not dispose of the NaOH solution.

Part 2. Buffering Properties of a Weak Base–Conjugate Acid Buffer

1. Obtain 50 mL of the ammonia–ammonium chloride buffer solution in a 50-mL graduated cylinder.
2. Transfer the 50 mL of buffer solution to a clean 100-mL beaker.
3. Set the 100-mL beaker on the magnetic stirrer. Add a stir bar to the solution. Gently stir the buffer solution.
4. Place the pH electrode in the solution. Record the pH of the solution in the Part 2 Data Table. (0 mL of 0.2 M HCl added.)
5. Using a clean graduated Beral-type pipet, transfer 1.0 mL of 0.2 M HCl solution to the ammonia–ammonium chloride buffer solution. Record the pH in the Part 2 Data Table.
6. Repeat step 5, recording the pH in the Part 2 Data Table after each 1.0 mL of HCl is added, until a total of 10.0 mL of HCl has been added.
7. Remove the electrode from the solution, place a 100-mL beaker under the electrode and rinse the electrode with deionized water.
8. Dispose of the ammonia–ammonium chloride solution as directed by the instructor and rinse the 100-mL beaker with deionized water.
9. Obtain another 50 mL of the ammonia–ammonium chloride buffer solution in a clean 50-mL graduated cylinder.
10. Transfer the 50 mL of buffer solution into the rinsed 100-mL beaker.
11. Set the beaker on a magnetic stirrer. Add a stir bar to the beaker and gently stir the buffer solution.
12. Place the pH electrode in the solution. Record the pH of the solution in the Part 2 Data Table. (0 mL of 0.2 M NaOH added.)
13. Using a clean graduated Beral-type pipet, transfer 1.0 mL of the 0.2 M NaOH solution to the ammonia–ammonium chloride buffer solution. Record the pH in the Part 2 Data Table.
14. Repeat step 13, recording the pH in the Part 2 Data Table after each 1.0 mL of NaOH is added, until a total of 10.0 mL of NaOH has been added.
15. Remove the electrode from the solution, place a 100-mL beaker under the electrode, and rinse the electrode with deionized water.
16. Dispose of the ammonia–ammonium chloride solution as directed by the instructor.

Part 3. Preparation of a pH 5.00 Buffer Solution

1. Calculate the correct volumes of the 0.1 M acetic acid solution and the 0.1 M sodium acetate solution needed to make 50 mL of a buffer solution with a pH value of 5.00. Record these amounts in the data table. K_a = 1.8 x 10^–5.
2. Obtain the calculated volume of 0.1 M acetic acid solution in a clean 50-mL graduated cylinder and transfer this volume to a clean 100-mL beaker.
3. Rinse the graduated cylinder with deionized water. Obtain the calculated volume of 0.1 M sodium acetate in the graduated cylinder and transfer the volume to the 100-mL beaker containing the acetic acid solution. Stir the solution.
4. Place the beaker under the pH electrode.
5. Place the pH electrode in the acetic acid–soidium acetate buffer solution. Record the pH in the Part 3 Data Table.
6. Remove the electrode from the solution, place a 100-mL beaker under the electrode, and rinse the electrode with deionized water.
7. Dispose of the acetic acid–acetate buffer solution as directed by the instructor.
8. Store and put away the pH electrode as directed by the instructor.
9. Dispose off all solutions as directed by the instructor.

Student Worksheet PDF

13894_Student1.pdf