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2607 Buffers

  • Page ID
    440624
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    1.0 INTRODUCTION

    In this lab you will explore the concept of pH, a measure of acidity, by using acid-base indicators. You will also prepare and experiment with buffer solutions, solutions that resist changes to pH.

    1.1 OBJECTIVES

    • Determine the pH of various solutions using pH indicators
    • Create and study the properties of buffer solutions.

    1.2 BACKGROUND

    In the first part of today’s lab you will use five indicators to determine the pH of several solutions to within one pH unit. An acid-base indicator is a chemical species that changes color at a specific pH as the pH (or acidity) of the solution varies. Acid-base indicators are themselves weak acids where the color of the aqueous acid is different than the color of the corresponding conjugate base.

    The five indicators you will use in this experiment, their color transitions, and their respective values of pKa are given in Table 1 below.

    We can use the values in Table 1 to determine the approximate pH of a solution. For example, suppose we have a solution in which methyl violet is violet. This tells us that the pH of our unknown solution is greater than or equal to 2 because methyl violet turns violet at pH values of 2 or greater. Now suppose we add some congo red to a fresh sample of our solution and find that the color is violet. This tells us that the pH of our solution is less than or equal to 3 because congo red turns violet at pH values of 3 or less. From these two tests we know that the pH range of our solution must be between 2 and 3. Thus, we have determined the pH of our solution to within one pH unit. Proceeding in a similar manner, you will use the acid-base indicators in Table 1 to determine the pH range of four solutions to within one pH unit.

    5nH1R-Ab3f5vdH3ZTYPfU3LZH4ehIDG9Vq8sYJLy08ovAtkLQtgf8U1KTNTm0qL3vikw2-OCwwYJsFPoD-aqXS4EXGdIXPAbfwNjov1Mrr7rXiQ3pIYFFm8iPzD7ruC9gWIMtm3eCdDKTBL_AKWq1Q

    oJHuMb8GBARY5go0A-GPZxWUBryp4v-VmBnkB40Azk_Q8wJ4Qq1daaw9iG3fTwUy9pfsNYcGm8krceX9RER4s-vubqExIl6bBl3WZXq5bFdAED8Y37_z2JivQ9c5w0M54sEsvP3976wCNpMEHevRqw

    Figure 1. This shows an example of six of the indicators from about pH 4 – pH 10. From left to right: universal indicator, phenolphthalein, bromothymol blue, congo red, bromocresol green, and methyl orange.

    In the second part of lab, you will prepare and experiment with buffer solutions: solutions that resist changes to pH.

    The dissociation of a weak acid in water yields H+ ion and the conjugate base of the acid, where the extent of dissociation is defined by the acid equilibrium constant, Ka:

    HA(aq) H+(aq) + A-(aq) Ka = [H+][A-]/[HA] (1)

    The anion (A-) in the above reaction, referred to as the conjugate base of HA, gives rise to the following equilibrium when an ionic salt (i.e. NaA) of this ion is dissolved in water:

    A-(aq) + H2O(l) HA(aq) + OH-(aq)

    Note that the anion accepts H+ from water, thereby acting as a base in this equation, consistent with its name: conjugate base.

    When a weak acid is dissolved in a solution containing its conjugate base, the result is a mixture that is highly resistant to changes in [H+] upon addition of either strong acid or strong base. Added OH- ions consume H+ ions already present which are replaced by further dissociation of HA; while added H+ ions are consumed as the equilibrium shifts to the left producing more HA. These solutions are known as buffer systems. They are used to control the [H+] as is required in a wide variety of practical applications, including life! Our blood is buffered by bicarbonate ions.

    For buffer solutions, it is convenient to rearrange the equilibrium expression (1) into the following form:

    [H+] = Ka[HA]/[A-]

    Given that [H+] is usually expressed in units of pH (pH = -log [H+]), the above equation can be conveniently rewritten by taking (– log) of both sides, and substituting:

    – log[H+] = pH and –log Ka = pKa

    This results in the Henderson-Hasselbalch equation:

    pH = pKa + log ([A-]/[HA]) (2)

    In this lab you will use the Henderson-Hasselbalch equation in two ways. In one experiment you will measure the pH of a mixture of known concentrations of a weak acid ([HA]) and its conjugate base ([A-]) in order to determine the pKa. In another experiment you will use the Henderson-Hasselbalch equation in order to design and prepare a buffer solution of a specific pH.





    2.0 SAFETY PROCEDURES AND WASTE DISPOSAL

    3.0 CHEMICALS AND SolutionS

    Chemical

    Approximate Amount

    Notes

    0.1 M HCl

    5 mL

     

    0.1 M NaH2PO4

    5 mL

     

    0.1 M CH3COOH

    5 - 25 mL

     

    0.1 M ZnSO4

    5 mL

     

    Methyl red solution

    Several drops

     

    Thymol blue solution

    Several drops

     

    Bromothymol Blue

    Several drops

     

    Methyl orange solution

    Several drops

     

    Congo red solution

    Several drops

     

    Phenolphthalein

    Several drops

     

    Bromocresol green

    Several drops

     

    Universal indicator solution

    Several drops

    optional

    0.1 M NaC2H3O2

    20 mL

     

    0.1 M NH4Cl

    20 mL

     

    0.1 M NH3

    20 mL

     

    0.1 M NaHCO3

    20 mL

     

    0.1 M Na2CO3

    20 mL

     

    0.1 M NaOH

    10 mL

     

    4.0 GLASSWARE AND APPARATUS

    pH meter and electrode (fragile!)

    5 small test tubes (or a well-plate)

    6 beakers, any sizes 50 - 150 mL

    Glass stir rod

    10 mL graduated cylinder

    Disposable pipets

    50 mL graduated cylinder

     

    5.0 PROCEDURE

    5.1 Determination of pH using Acid-Base Indicators

    1. Rinse several small test tubes using laboratory water (there is no need to dry these). To each test tubes add approximately 1 mL of 0.1 M HCl(aq). (Alternatively, if using a well-plate, place 2-3 drops of HCl to each well.)
    2. Into each of these test tubes add one of the eight indicators listed in Table 1. (On a well-plate, use only 1 drop of the indicator.) Record the color of each solution in the appropriate data table below.
    3. Repeat steps 1-2 for the following solutions and record your color observations in the data table:

    0.1 M sodium hydrogen phosphate, NaH2PO4(aq)

    0.1 M acetic acid, CH3COOH(aq)

    0.1 M zinc sulfate, ZnSO4(aq)

    5.2 Determination of Ka

    1. From among the reagents provided by the stockroom, select one buffer system that you can prepare from two of the provided solutions. (You need to understand the definition of buffers in order to do this step correctly! See 1.1 BACKGROUND) Use a graduated cylinder to transfer 20 mL of your chosen acid component of your buffer into a 150 mL beaker and then transfer 20 mL of your chosen base component of your buffer into the same 150 mL beaker. Stir this mixture well using a glass stir rod. Record on your data chart the buffer system you have selected.
    2. Calibrate a pH meter as directed by your instructor. Rinse the pH electrode well with laboratory water prior to any measurements. Do not allow the tip of the pH electrode to touch any solid object—it is very fragile!
    3. Measure the pH of your buffer solution and record it in the data table below.
    4. Fill a clean beaker with 30 mL laboratory water. Measure and record the pH of pure water.
    5. Pour the 30 mL laboratory water into your buffer solution. Stir well. Measure and record the pH of your diluted buffer.
    6. Divide your diluted buffer solution into two portions in separate beakers.
    7. Add 1-2 mL 0.1 M HCl to one of the portions, and 3 mL 0.1 M NaOH to the other portion. Stir each solution well, then measure and record the pH of each solution.
    8. Dispose of your buffer solutions, then clean and rinse the beakers well.
    9. Fill a clean beaker with 30 mL laboratory water, then add 1-2 mL 0.1 M HCl. Measure and record the pH.
    10. Fill a clean beaker with 30 mL laboratory water, then add 3 mL 0.1 M NaOH. Measure and record the pH.
    11. In a clean beaker prepare a new buffer of the same system; however, this time use 2 mL of your selected acid and 20 mL of your selected base. Measure and record the pH of this 2:20 buffer solution.
    12. Add 3 mL 0.1 M NaOH to the 2:20 buffer solution. Measure and record the pH.


    5.3 Designing and Preparing a Buffer of a Specific pH:

    1. You will be assigned a specific pH by your instructor which will not be the same as any of your classmates. Record the pH you have been assigned (your “target pH”). Determine which buffer system available would be appropriate to design a buffer of the pH you were assigned. Check with your instructor to make sure you’ve selected the correct system for your target pH before proceeding further.
    2. Showing your work in 7.0 DATA ANALYSIS, calculate the volume of each stock solution you need to prepare 30 – 50 mL buffer solution of your target pH. (Hint, choose either acid or base to begin with 20 mL, and determine the volume you will need of the other solution to have the correct base/acid ratio for your target pH).
    3. Prepare the buffer solution, then measure and record the pH.
    4. Allow for some small experimental error between your target pH and your measured pH. However, if the experimental error is large, you should recalculate and prepare another buffer solution.



    6.0 DATA RECORDING SHEET

    Last Name

    First Name

     

    Partner Name(s)

    Date

    Determination of pH using Acid-Base Indicators

    Record the colors of each solution tested. Then after lab ends, use these colors and Table 1 to estimate the pH range of each solution (for example, pH = 3 - 4).

    Indicator

    0.1 M HCl

    0.1 M NaH2PO4

    0.1 M CH3COOH

    0.1 M ZnSO4

    Methyl red

           

    Thymol blue

           

    Methyl orange

           

    Congo red

           

    Phenolphthalein

           

    Bromocresol green

           

    Bromothymol blue

           

    Universal indicator (optional)

           

    pH range =


           

    Determination of Ka

    List the buffer system you chose:

    Acid component = ____________________

    Base component = ____________________

    Record your pH measurements in the Table below. After lab ends, use your data and the Henderson-Hasselbalch equation as discussed in 1.1 BACKGROUND to calculate the pKa and the Ka, showing your work in 7.0 DATA ANALYSIS.

     

    Buffer System

    Pure Water

     

    Acid/Base = 1:1

    Acid/Base = 2:20

    (pure water)

    Observed pH

         

    pH after dilution with water

     

    N/A

    N/A

    pH after addition of HCl

     

    N/A

     

    pH after addition of NaOH

         

    Calculated pKa

       

    N/A

    Calculated Ka

       

    N/A

    7.0 DATA ANALYSIS

    7.1 Determination of pH using Acid-Base Indicators

    1. Choose one of the four solutions (0.1 M HCl, 0.1 M NaH2PO4, 0.1 M CH3COOH, or 0.1 M ZnSO4), describe the most helpful observations of that solution with the indicators, and explain how you used Table 1 and the colors of the indicators to determine the pH range for your chosen solution.












    1. Consider your results for the solutions of 0.1 M HCl and 0.1 M CH3COOH. Notice that the concentrations are the same. Which has the lower pH and why is its pH lower?









    1. Consider your results for the 0.1 M NaH2PO4 solution.
    1. Is the solution acidic or basic?

    1. Which ion, Na+ or H2PO4-, is causing the observed acidity or basicity?

    1. Write the net ionic equation below that shows why this ion is acidic or basic.


    1. Consider your results for the 0.1 M ZnSO4 solution.
    1. Is the solution acidic or basic?

    1. Which ion, Zn2+ or SO42-, is causing the observed acidity or basicity?

    1. Write the net ionic equation below that shows why this ion is acidic or basic.



    1. Determination of Ka
    1. Show your work for the calculations of pKa and Ka (see 1.1 BACKGROUND). Record your answers in the Table found in 6.0 DATA RECORDING SHEET.








    1. Compare your experimental value of pKa with the literature value of the pKa of the weak acid. Propose reasons for any differences.






    1. Consider the addition of 3 mL of NaOH to the 1:1 buffer and to the 2:20 buffer. Did the pH change by the same amount when NaOH was added to each buffer? Propose a reason for why or why not the pH changes were the same.





    1. Consider the addition of 30 mL of pure water to the 1:1 buffer. Did the pH of the buffer change upon dilution? Propose a reason for why or why not.










    1. Consider the addition of HCl to the 1:1 buffer and to pure water. Was the pH different? Propose a reason for why or why not. Include chemical reaction(s) in your explanation.









    7.3 Designing and Preparing a Buffer of a Specific pH

    Target pH _________________

    List the buffer system you chose:

    Acid component = ____________________

    Base component = ____________________

    pKa of acid component (look it up) = ____________________

    1. Show your work and calculate the volume of each stock solution you need to prepare 30 – 50 mL buffer solution of your target pH. (Hint, choose either acid or base to begin with 20 mL, and determine the volume you will need of the other solution to have the correct base/acid ratio for your target pH. See 1.1 BACKGROUND for more info).

    Volume of Acid component = ____________________

    Volume of Base component = ____________________

    Measured pH of your buffer solution _________________

    Percent error = (measured pH – target pH)/target pH x 100 = _________________





    1. Discuss your results and propose reasons for your percent error.

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