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2.4: Buffer Preparation

  • Page ID
    114279
  • RELATED READINGS: Pages 13-15, 17-19, 278-279, (Appendix B)

    OBJECTIVES

    Upon completion of this exercise, appropriate discussions, and related reading, the student will be able to:

    1. Prepare buffer solutions from aqueous standards.
    2. Measure pH using a pH meter.
    3. Calculate pH values using the Henderson-Hasselbach equation.

    PRINCIPLE

    Buffers are used in almost all biochemical reactions to maintain optimum pH conditions. The buffer, composed of a weak acid (HA) and its conjugate base (A-), is chosen to provide maximal buffering capacity at a given pH. The calculation of the quantity of each buffer component needed to achieve the desired pH is determined by use of the Henderson-Hasselbach equation:

    \[pH=pK_{a} + \log \frac{[A^{-}]}{[HA]}\]

    where pKa is the dissociation constant for the weak acid. The maximum buffering capacity occurs when the pH of the solution is equal to the pKa of the weak acid. Knowledge of how to prepare a buffered solution is critical for accurate biochemical analysis.

    Materials

    • pH calibrating solution
    • 50 mL Volumetric Flasks
    • pH meter Pipets
    • O.2M Na2HPO4
    • Distilled Water
    • O.2M NaH2PO4 (pKa = 6.8)

    pH METER CALIBRATION (Optional)

    1. Remove electrode from storage solution.
    2. Check that the internal KCl solution is filled to the proper level.
    3. If the KCl level is incorrect, check with you instructor.
    4. Carefully blot the electrode dry (do not wipe the electrode).
    5. Immerse the tip of the electrode in the first calibration buffer (usually pH 7.00).
    6. Turn the instrument to “measure”.
    7. After allowing for equilibration (approx. 30 seconds), adjust pH reading to 7.00 according to your instructor’s directions.
    8. Turn the instrument to “standby”.
    9. Remove the electrode from the calibrating solution.
    10. Rinse the electrode with distilled water and blot dry.
    11. Repeat steps 5 and 6 using a second calibration buffer. This is usually pH 4.00 for a pH measurement <7.0, or pH 10.0 for a pH measurement >7.0.
    12. After calibration is completed turn the instrument to “standby”.
    13. Leave the electrode standing in neutral buffer or distilled water.

    PROCEDURE

    1. Carefully add the volumes of stock Na2HPO4 solution shown on the data sheet into clean labeled 50 mL volumetric flasks.
    2. Bring to volume with NaH2PO4 and water as indicated on data sheet.
    3. Mix each combination thoroughly.
    4. Following the directions given by your instructor, measure the pH of each buffer solution and record the value on the data sheet.
    5. Calculate the expected pH of each buffer using the Henderson-Hasselbach equation and record the value on the data sheet.
    6. Seal each flask with parafilm and store as indicated by your instructor. These buffers will be used in the titration exercise (Exercise #5).
    DATA SHEET, EXERCISE #4

    NAME: _______________

    DATE: _______________

    Solution Stock
    Solution
    Stock
    Solution
    H2O Measured pH Calculated pH
    0.2 M Na2HPO4 0.2 M Na2HPO4 H2O mL
    A 4.0 mL 46.0 mL 0
    B 12.3 mL 37.7 mL 0
    C 30.5 mL 19.5 mL 0
    D 43.6 mL 6.4 mL 0
    E 46.0 mL 4.0 mL 0
    F 2.0 mL 23.0 mL 25 mL

    CALCULATIONS

    Calculate the theoretical pH of each buffer solution using the Henderson-Hasselbach equation

    \[pH = pK_{a} + \log \frac{[A^{-}]}{[HA]}\]

    Think! Which of the stock buffer solutions contains the conjugate base (A-) and which the weak acid (HA)?

    Discussion Questions

    1. Describe how a pH electrode measures pH.
    2. Will buffer “E” that you prepared, be more effective against acid or alkaline changes? Why?
    3. How much 0.2 M sodium acetate and 0.2 M acetic acid would be required to make 100 mL of pH 4.7 acetate buffer? (pKa = 4.75).
    0.2 M sodium acetate _______________mL
    0.2 M acetic acid _______________mL
    1. Solution F has half the concentration of conjugate base and weak acid as solution A but has the same pH. How can you explain this?
    2. How do the measured and calculated pH of solutions B & 0 compare to the pHs expected from Appendix B, in Kaplan and Pesce?