PREPARATION OF A BUFFERED SOLUTION
BACKGROUND
Most enzymes can function properly only within a narrow pH range. For that reason, our blood
maintains a pH of 7.4, primarily due to its carbonic acid-bicarbonate buffer.
H+ + HCO3- ↔ H2CO ↔ H2O + CO2
A pH deviation may denature proteins such that they do not bind to the appropriate binding sites
and enzymes fail to perform their functions.
By buffering a solution, we can prevent or correct large shifts in pH that occur when H+ or OHions are added. Typically a buffered solution is made up of a weak acid and its conjugate base,
or a weak base and its conjugate acid. Buffers may also be prepared from a weak acid with a
small amount of a strong base, a weak base with a small amount of strong acid, or from an
amphoteric substance.
The Henderson-Hasselbach equation allows the pH of a buffered solution to be calculated:
pH = pKa + log
[𝐵𝑎𝑠𝑒]
[𝐴𝑐𝑖𝑑]
Adding protons or hydroxide ions to the solution alters the ratio of acid and conjugate base.
When the ratio changes, the stoichiometry of the new system must be calculated before any
further equilibrium calculations can be done. After those calculations are made, pH may be
determined.
The same equation is used in designing a buffer of a certain target pH. If the target pH is
inserted for the pH in the equation and the pKa of the acid (or base) is inserted into the equation,
the ratio of the acid (or base) to its salt can be calculated.
To prepare a buffered solution of a certain pH and buffer capacity:
1. Decide which conjugate acid-base pair to use for your solution. A buffer is most
effective when the acid and base forms are in approximately equal concentrations. For
this reason, it is best to select a compound with a pKa that is within 1 pH unit of the
desired pH. The buffering range is + 1 pH unit of the pKa.
2. Next, decide the concentration of the buffer. Since the concentration of the buffer
components determine its buffering capacity, you need to be aware of the amount of
added H+ or OH- the buffered solution is likely to encounter.
Creating an ICE table (Initial, Change, Equilibrium) helps in determining the concentration of
species when equilibrium shifts.
PRELAB QUESTIONS
1. The pKa of acetic acid is 4.74. What is the buffering range of the acid and its conjugate
base? Explain your answer.
2. For this experiment, you will need to prepare 100.00mL of a 0.10M HC2H3O2 solution.
Calculate the volume of 6M acetic acid needed to prepare this solution.
3. Calculate the mass of sodium acetate (NaC2H3O2) required to prepare 100.00mL of a 0.10M
solution.
4. Determine the volumes of 0.10M HC2H3O2 and 0.10M NaC2H3O2 required to prepare
10.0mL of buffer of each of the following pH values. pKa of HC2H3O2 = 4.74.
a. pH 3.7
c. pH 5.7
b. pH 4.7
5. Read the laboratory procedures and then diagram the procedures using illustrations
and/or a flow chart to describe the steps. Ensure that your diagrams are labeled with
quantities and descriptions.
MATERIALS
Universal indicator
6 M acetic acid
sodium acetate
0.1 M NaOH
0.1 M HCl
distilled water (bottle and wash bottle)
deionized water (quart container)
electronic balance
scoopula
weight boat
24-well plate
Extra 100mL graduated cylinder
One 250mL flask
8 paper cups
10-mL graduated cylinder
China marker (wax pencil)
Sharpie®
color chart
plain white paper
9 disposable pipets
PROCEDURE
Part One
1. If you are not familiar with the proper procedure for preparing solutions or using any of the
equipment in this lab, ask your teacher for assistance.
2. Prepare 100.00 mL of a 0.10 M acetic acid solution and 100.00 mL of a 0.10 M solution of
sodium acetate based on the calculations you performed in the pre-lab section in two 100.00mL graduated cylinders. Label each with china marker.
3. Obtain approximately 100 mL of distilled water in a flask and label the container with china
marker.
4. Add 20 drops of Universal indicator to both of the solutions you prepared and the water in
the flask.
5. Label two more flasks - one “NaOH” and the other “HCl.” Use china marker. Obtain about
10 mL of each of these chemicals in the labeled containers. Add one drop of universal
indicator to each.
6. Label the other three flasks—“3.7,” “4.7,” and “5.7” with the china marker indicating the pH
of the buffer that each will contain.
7. Each of the nine pipets is to be used with a single solution. They are labeled to prevent
confusion. They are 0.1 M acetic acid, 0.1 M sodium acetate, NaOH, HCl, pH 3.7 buffer, pH
4.7 buffer, and pH 5.7 buffer, and universal. The final pipet for (distilled) water is not
marked.
8. Prepare each of the buffer solutions in the labeled flasks according to your pre-lab
calculations. The pipets are helpful for transferring the stock solution to the graduated
cylinder for precise measurement.
Note: Each of these solutions should be a different color. Refer to the color chart for the
indicator solution to confirm the pH of your prepared buffers.
9. Place the 24-well plate on a sheet of white paper to make it easier to compare the colors of
the solutions.
10. Use the appropriate graduated pipet to transfer 0.1 M acetic acid to each of the four wells in
the first row of the 24-well plate. Fill each well ½ full. Repeat with water in the second row,
0.1 M sodium acetate in the third row, 3.7 buffer in the fourth row, 4.7 buffer in the fifth row,
and 5.7 buffer in the sixth row. See Figure One.
11. Observe and record the color of the solution in the first well of each row on your data table.
Match the color to the color chart to determine the approximate pH of the solution, and
record this on your data table.
12. Leave the first well of each row as a reference.
13. In the second well of each row, add 1 drop of 0.10 M HCl using the pipet. Gently swirl the
solutions while resting the plate on the table. Record the resulting color and approximate pH
of each solution on your data table.
14. In the third well of each row, add 1 drop of 0.10 M NaOH. Gently swirl the solution while
resting the plate on the table. Record the resulting color and approximate pH of each solution
on your data table.
15. In the fourth well of each row, add 10 drops of 0.10 M NaOH, again recording the color and
corresponding pH on your data table.
If you have trouble seeing all the colors, add one drop of Universal Indicator to EACH well and
then record the colors and pH.
16. Bring the distilled water with universal indicator and any NaOH that you have left to your
teacher. You will need it for Part Two and can get it from her later.
17. Clean the well plate under running water. Clean the flask, graduated cylinders, and cups
also. Rinse with distilled water and dry.
Part Two
1. Turn the well plate so that it is four rows by six columns.
2. In one of the flasks, prepare 20.0 mL of the 4.7 buffer by doubling the calculations from the
Pre-laboratory Questions.
3. Add 2.0 mL of this solution to the six wells of the first row. See the following figure.
Figure 2.
Arrangement of
24-well plate for
Part Two.
4. Dilute 5.0 mL of this buffer solution to 20.0 mL to prepare the first dilution, using the water
from your flask which already contains indicator. Use the universal indicator colored distilled
water for this. Put in a paper cup marked #1. Add 2.0 mL of this dilution to each of the
wells in the second row.
5. Dilute 5.0 mL of the first dilution to 20.0 mL to prepare the second dilution. Use the
universal indicator colored distilled water for this. Put in a paper cup marked #2. Add 2.0
mL to each of the wells in the third row.
6. Dilute 5.0 mL of the second dilution to 20.0 mL to prepare the third dilution. Use the
universal indicator colored distilled water for this. Put in a paper cup marked #3 Add 2.0
mL of this solution to each of the wells in the fourth row.
7. Record the color of the first well of each row. You will not add any NaOH to these wells.
8. To the second well of each row add one drop of the 0.10M NaOH.
9. Add two drops NaOH to the third well of each row. Repeat, with three drops in the fourth
well, four drops in the fifth well, and five drops in the sixth well.
10. Gently swirl the solutions until the color within each well is uniform. Record the color and
corresponding pH in your data table.
11. Clean the well plate under running water. Clean all the cups also. Rinse with distilled water
and dry.
DATA AND OBSERVATIONS
Part One
Initial
Color
pH
1 drop HCl
Color
pH
1 drop NaOH
Color
pH
10 drops NaOH
Color
pH
0.10M
acetic acid
Water
0.1M
sodium
acetate
3.7 buffer
4.7 buffer
5.7 buffer
Part Two
Initial
color
Buffer
1st
dilution
2nd
dilution
3rd
dilution
pH
1 drop NaOH
2 drops NaOH
3 drops NaOH
4 drops NaOH 5 drops NaOH
Color
Color
Color
Color pH Color pH
pH
pH
pH
CALCULATIONS
Calculate the concentration of the acetate ion in each sample. Show your work for each.
Record the answer in the table.
Acetate ion
concentration
Buffer
1st
dilution
2nd
dilution
3rd
dilution
CONCLUSION
1. Write the balanced equation for the dissociation equilibrium of acetic acid.
2. Assuming the pipet dispenses 20 drops per milliliter, calculate the final pH of the fullstrength buffer after the addition of five drops of 0.10M NaOH. Ka = 1.8 × 10–5
3. Calculate the final pH of the second dilution buffer after the addition of one drop of 0.10M
NaOH.
LAB REPORT
Prelab questions from the group
Data and Observations from the group
Calculations from individuals
Conclusions from individuals.
Lab is due _________________________